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If g1 and g2 are two elements of a group G, they are called conjugate if there exists an element g3 in G such that:
g3g1g3-1 = g2.
Conjugacy is an equivalence relation and therefore partitions G into equivalence classes: every element of the group belongs to precisely one conjugacy class
The equivalence class that contains the element g1 in G is
Cl(g1) = { g3g1g3-1| g3 ∈ G}
and is called the conjugacy class of g1. The class number of G is the number of conjugacy classes.
The classes Cl(g1) and Cl(g2) are equal if and only if g1 and g2 are conjugate, and disjoint otherwise.
For Abelian groups the concept is trivial, since each element forms a class on its own.
Conventional cell
For each lattice, the conventional cell is the cell obeying the following conditions:
• its basis vectors define a right-handed axial setting;
• its edges are along symmetry directions of the lattice;
• it is the smallest cell compatible with the above condition.
Crystals having the same type of conventional cell belong to the same crystal family.
Coset
If G is a group, H a subgroup of G, and g an element of G, then
gH = { gh : h ∈ H } is a left coset of H in G
Hg = { hg : h ∈ H } is a right coset of H in G.
The decomposition of a group into cosets is unique. Left coset and right cosets however in general do not coincide, unless H is a normal subgroup of G.
Any two left cosets are either identical or disjoint: the left cosets form a partition of G, because every element of G belongs to one and only one left coset. In particular the identity is only in one coset, and that coset is H itself; this is also the only coset that is a subgroup. The same holds for right cosets.
All left cosets and all right cosets have the same order (number of elements, or cardinality), equal to the order of H, because H is itself a coset. Furthermore, the number of left cosets is equal to the number of right cosets and is known as the index of H in G, written as [G : H] and given by Lagrange's theorem:
|G|/|H| = [G : H].
Cosets are also sometimes called associate complexes.
Example
The coset decomposition of the twin lattice point group with respect to the point group of the individual gives the different possible twin laws. Each element in a coset is a possible twin operation.
Crystal
Those solids in which atoms, ions or molecules are arranged in a definite three dimensional pattern are called crystalline solids. A material is a crystal if it has essentially a sharp diffraction pattern. The word essentially means that most of the intensity of the diffraction is concentrated in relatively sharp Bragg peaks, besides the always present diffuse scattering. In all cases, the positions of the diffraction peaks can be expressed by
$H = \sum _{i=1}^n h_i a_i^* \, (n \ge 3) \nonumber$
Here $\textbf{a}_{i}^{*}$ and $h_i$ are the basis vectors of the reciprocal lattice and integer coefficients respectively and the number $n\0 is the minimum for which the positions of the peaks can be described with integer coefficient \(h_i$. The conventional crystals are a special class, though very large, for which $n = 3$.
See also
Acta Cryst. (1992), A48, 928 where the definition of a crystal appears in the Terms of reference of the IUCr commission on aperiodic crystals
Crystal family
A crystal family is the smallest set of space groups containing, for any of its members, all space groups of the Bravais flock and all space groups of the geometric crystal class to which this member belongs.
Crystal pattern
An object in the n-dimensional point space En is called an n-dimensional crystallographic pattern or, for short, crystal pattern if among its symmetry operations:
1. there are n translations, the translation vectors t1, ... , tn of which are linearly independent;
2. all translation vectors, except the zero vector 0, have a length of at least d > 0.
When the crystal pattern consists of atoms, it takes the name of crystal structure. The crystal pattern is thus the generalization of a crystal structure to any pattern, concrete of abstract, in any dimension, which obeys the conditions of periodicity and discreteness expressed above.
Crystal system
A crystal-class system , or crystal system for short, contains complete geometric crystal classes of space groups . All those geometric crystal classes belong to the the same crystal system which intersect exactly the same set of Bravais flocks.
Crystallographic basis
A basis of n vectors e1, e2, ... , en of the vector space Vn is a crystallographic basis of the vector lattice L if every integral linear combination t = u1e1 + u2e2 + ... + unen is a lattice vector of L. It may or may not be a primitive basis. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Conjugacy_class.txt |
In mathematics , an orbit is a general group-theoretical term describing any set of objects that are mapped onto each other by the action of a group. In crystallography, the concept of orbit is used to indicate a point configuration in association with its generating group.
From any point of the three-dimensional Euclidean space the symmetry operations of a given space group G generate an infinite set of points, called a crystallographic orbit. The space group G is called the generating space group of the orbit. Two crystallographic orbits are said configuration-equivalent if and only if their sets of points are identical.
Crystallographic symmetry
A motion is called a crystallographic symmetry operation if a crystal pattern exists for which it is a symmetry operation.
D centered cell
The D centered cell is the used for the rhombohedral description of the hexagonal lattice. Six right-handed D cell with basis vectors of equal length are obtained from the hP cell by means of one of the following transformation matrices:
D1: 10-1/01-1/111 D2: -101/0-11/111
the other four D cells are obtained by cyclic permutation of the basis vectors.
The resulting hD cell has centering nodes at 1/3,1/3,1/3 and 2/3,2/3,2/3
Direct Lattice
The direct lattice represents the triple periodicity of the ideal infinite perfect periodic structure that can be associated to the structure of a finite real crystal. To express this periodicity one calls crystal pattern an object in point space En (direct space) that is invariant with respect to three linearly independent translations, t1, t2 and t3. One distinguishes two kinds of lattices, the vector lattices and the point lattices.
Any translation t = ui ti (ui arbitrary integers) is also a translation of the pattern and the infinite set of all translation vectors of a crystal pattern is the vector lattice L of this crystal pattern.
Given an arbitrary point P in point space, the set of all the points Pi deduced from one of them by a translation PPi = ti of the vector lattice L is called the point lattice.
A basis a, b, c of the vector space Vn is a crystallographic basis of the vector lattice L if every integral linear combination t =u a + v b + w c is a lattice vector of L. It is called a primitive basis if every lattice vector t of L may be obtained as an integral linear combination of the basis vectors, a, b, c. Referred to any crystallographic basis the coefficients of each lattice vector are either integral or rational, while in the case of a primitive basis they are integral. Non-primitive bases are used conventionally to describe centered lattices.
The parallelepiped built on the basis vectors is the unit cell. Its volume is given by the triple scalar product, V = (a, b, c).
If the basis is primitive, the unit cell is called the primitive cell. It contains only one lattice point. If the basis is non-primitive, the unit cell is a multiple cell and it contains more than one lattice point. The multiplicity of the cell is given by the ratio of its volume to the volume of a primitive cell.
The generalization of the notion of point and vector lattices to n-dimensional space is given in Section 8.1 of International Tables of Crystallography, Volume A | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Crystallographic_orbit.txt |
In group theory, direct product of two groups (G, *) and (H, o), denoted by G × H is the as set of the elements obtained by taking the cartesian product of the sets of elements of G and H: {(g, h): g in G, h in H};
For abelian groups which are written additively, it may also be called the direct sum of two groups, denoted by \(G \oplus H\).
The group obtained in this way has a normal subgroup isomorphic to G (given by the elements of the form (g, 1)), and one isomorphic to H (comprising the elements (1, h)).
The reverse also holds: if a group K contains two normal subgroups G and H, such that K= GH and the intersection of G and H contains only the identity, then K = G x H. A relaxation of these conditions gives the semidirect product.
Direct space
The direct space (or crystal space) is the point space, En, in which the structures of finite real crystals are idealized as infinite perfect three-dimensional structures. To this space one associates the vector space, Vn, of which lattice and translation vectors are elements. It is a Euclidean space where the scalar product of two vectors is defined. The two spaces are connected through the following relations:
(i) To any two points P and Q of the point space En a vector PQ = r of the vector space Vn is attached
(ii) For each point P of En and for each vector r of Vn there is exactly one point Q of En for which PQ = r holds
(iii) If R is a third point of the point space, PQ + QR = PR
Displacive modulation
For a displacively modulated crystal phase, the positions of the atoms are displaced from those of a basis structure with space group symmetry (an ordinary crystal). The displacements are given by the atomic modulation function uj(r), where j indicates the jth atom in the unit cell of the basic structure.
$r( n,j)~=~ n+ r_j+ u_j( n+ r_j).$
The modulation function has a Fourier expansion
$u_j( r)~=~\sum_ k \hat{ u}( k) \exp (2\pi i k. r),~with~ k=\sum_{i=1}^n h_i a_i^*,$
with finite value of n. If n=1, the modulated structure is one-dimensionally modulated. A special case of a one-dimensionally modulated structure is
$r(n,j)_{\alpha}~=~ n_{\alpha}+ r_{j\alpha}+A_{j\alpha} \sin \left(2\pi i q. n+ r_j)+\phi_{j\alpha}\right), (\alpha=x,y,z).$
Domain of influence
The domain of influence of a lattice point P (Delaunay 1933), or Dirichlet domain or Voronoi domain, consists of all points Q in space that are closer to this lattice point than to any other lattice point or at most equidistant to it, namely such that OP ≤|t - OP| for any vector t belonging to the vector lattice L. It is the inside of the Wigner-Seitz cell.
Double coset
Let G be a group, and H and K be two subgroups of G. One says that the two elements g1 ∈ G and g2 ∈ G belong to the same double coset of G relative to H and K if there exist elements hi ∈ H and kj ∈ K such that:
g2 = hig1kj
The complex Hg1K is called a double coset
The partition of G into double cosets relative to H and K is a classification, i.e. each gi ∈ G belongs to exactly one double coset. It is also a generalization of the coset decomposition, because the double coset Hg1K contains complete left cosets of K and complete right cosets of H.
Eigensymmetry
The eigensymmetry, or inherent symmetry, of a crystal is the point group or space group of a crystal, irrespective of its orientation and location in space. For instance, all individuals of a twinned crystal have the same (or the enantiomorphic) eigensymmetry but may exhibit different orientations. The orientations of each of two twin components are related by a twin operation which cannot be part of the eigensymmetry.
In morphology, the eigensymmetry is the full symmetry of a crystalline form, considered as a polyhedron by itself. The eigensymmetry point group is either the generating point group itself or a supergroup of it.
Euclidean mapping
The Euclidean mapping or isometry is a special case of affine mapping that, besides collinearity and ratios of distances, keeps also distances and angles. Because of this, a Euclidean mapping is also called a rigid motion.
Euclidean mappings are of three types:
• translations
• rotations
• reflections.
A special case of Euclidean mapping is a symmetry operation.
Factor group
Let N be a normal subgroup of a group G. The factor group or quotient group G/N is the set of all left cosets of N in G, i.e.:
\(G/N = \{ aN : a \isin G \}.\)
For each aN and bN in G/N, the product of aN and bN is (aN)(bN), which is still a left coset. In fact, because N is normal:
(aN)(bN) = a(Nb)N = a(bN)N = (ab)NN = (ab)N.
The inverse of an element aN of G/N is a-1N.
Family structure
By superposing two or more identical copies of the same polytype translated by a superposition vector (i.e. a vector corresponding to a submultiple of a translation period) a fictitious structure is obtained, which is termed a superposition structure. Among the infinitely possible superposition structures, that structure having all the possible positions of each OD layers is termed a family structure: it exists only if the shifts between adjacent layers are rational, i.e. if they correspond to a submultiple of lattice translations.
The family structure is common to all polytypes of the same family. From a group-theoretical viewpoint, building the family structure corresponds to transforming (“completing”) all the local symmetry operations of a space groupoid into the global symmetry operations of a space-group. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Direct_product.txt |
Space groups with no special Wyckoff positions (i.e. with no special crystallographic orbits) are called fixed-point-free space groups or torsion-free space groups or Bieberbach groups. In fixed-point-free space groups group every element other than the identity has infinite order.
Flack parameter
The Flack parameter is the molar fraction x in the defining equation $C=(1-x)X + x\bar X$, where C represents an oriented two-domain-structure crystal, twinned by inversion, consisting of an oriented domain structure X and an oriented inverted domain structure $\bar X$. In reciprocal space, the Flack parameter x is defined by the structure-amplitude equation
$G^2(h,k,l,x)=(1-x)|F(h,k,l)|^2 + x|F({\bar h}, {\bar k}, {\bar l})|^2$.
For a multidomain-structure twin of a chiral crystal structure, an equivalent Flack parameter may be calculated according to the method of Flack and Bernardinelli (1999).
Form
For a point group P a form is a set of all symmetrically equivalent "elements", namely:
• in vector space, a crystal form or face form is a set of all symmetrically equivalent faces;
• in point space, a point form is a set of all symmetrically equivalent points.
The polyhedron or polygon of a point form is dual to the polyhedron of the corresponding face form, where "dual" means that they have the same number of edges but the number of faces and vertices is interchanged. The inherent symmetry of a form is a point group C which either coincides with the generating point group P or is a supergroup of it.
Forms in point groups correspond to crystallographic orbits in space groups.
Friedel's law
Friedel's law, or rule, states that the intensities of the h, k, l and ${\bar h}, {\bar k}, {\bar l}$ reflections are equal. This is true either if the crystal is centrosymmetric or if no resonant scattering is present. It is in that case not possible to tell by diffraction whether an inversion center is present or not. The apparent symmetry of the crystal is then one of the eleven Laue classes.
The reason for Friedel's rule is that the diffracted intensity is proportional to the square of the modulus of the structure factor, |Fh|2, according to the geometrical, or kinematical theory of diffraction. It depends similarly on the modulus of the structure factor according to the dynamical theory of diffraction. The structure factor is given by:
$F_h = \Sigma_j f_j {\rm exp - 2 \pi i} {\bold h} . {\bold r_j}$
where fj is the atomic scattering factor of atom j, h the reflection vector and ${\bold r_j}$ the position vector of atom j. There comes:
$|F_h|^2 = F_h F_h^* = F_h F_{\bar h} = |F_{\bar h}|^2$
if the atomic scattering factor, fj, is real. The intensities of the h, k, l and ${\bar h}, {\bar k}, {\bar l}$ reflections are then equal. If the crystal is absorbing, however, due to resonant scattering, the atomic scattering factor is complex and
$F_{\bar h} \ne F_h^*$
The reflections h, k, l and ${\bar h}, {\bar k}, {\bar l}$ are called a Friedel pair. They are used in the resolution of the phase problem for the solution of crystal structures and in the determination of absolute structure.
History
Friedel's law was stated by G. Friedel (1865-1933) in 1913 (Friedel G., 1913, Sur les symétries cristallines que peut révéler la diffraction des rayons X., C.R. Acad. Sci. Paris, 157, 1533-1536.
See also
Absolute structure
Section 3.1 of International Tables of crystallography, Volume A
Geometric crystal class
Geometric crystal classes (or simply 'crystal classes') classify the symmetry groups of the external shape of macroscopic crystals, namely according to the morphological symmetry. There are 10 two-dimensional geometric crystal classes and 32 three-dimensional geometric crystal classes, in one to one correspondence with the 10 and 32 types of point groups in E2 and E3, respectively.
See also
• Section 8.2.4 of International Tables for Crystallography, Volume A
Geometric element
A geometric element is an element in space (plane, line, point, or a combination of these) about which asymmetry operation is performed. Geometric elements are classified on the basis of the dimensionality N of the space on which they act, the upper limit on the dimensionality of the symmetry element being N-1. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Fixed-point-free_space_groups.txt |
A set G equipped with a binary operation *: G x GG, assigning to a pair (g,h) the product g*h is called a group if:
1. The operation is associative, i.e. (a*b)*c = a*(b*c).
2. G contains an identity element (neutral element) e: g*e = e*g = g for all g in G
3. Every g in G has an inverse element h for which g*h = h*g = e. The inverse element of g is written asg -1.
Often, the symbol for the binary operation is omitted, the product of the elements g and h is then denoted by the concatenation gh.
The binary operation need not be commutative, i.e. in general one will have g*h ≠ h*g. In the case that g*h = h*g holds for all g,h in G, the group is an Abelian group.
A group G may have a finite or infinite number of elements. In the first case, the number of elements of G is the order of G, in the latter case, G is called an infinite group. Examples of infinite groups are space groups and their translation subgroups, whereas point groups are finite groups.
Group isomorphism
A group isomorphism is a special type of group homomorphism. It is a mapping between two groups that sets up a one-to-one correspondence between the elements of the groups in a way that respects the respective group operations. If there exists an isomorphism between two groups, then the groups are called isomorphic. Isomorphic groups have the same properties and the same structure of their multiplication table.
Let (G, *) and (H, #) be two groups, where "*" and "#" are the binary operations in G and H, respectively. A group isomorphism from (G, *) to (H, #) is a bijection from G to H, i.e. a bijective mapping f : GH such that for all u and v in G one has
f (u * v) = f (u) # f (v).
Two groups (G, *) and (H, #) are isomorphic if an isomorphism between them exists. This is written:
(G, *) \(\cong\) (H, #)
If H = G and the binary operations # and * coincide, the bijection is an automorphism.
Groupoid
A groupoid (G,*) is a set G with a law of composition * mapping of a subset of G x G into G. The properties of a groupoid are:
• if x, y, z ∈ G and if one of the compositions (x*y)*z or x*(y*z) is defined, so is the other and they are equal; (associativity);
• if x, x' and y ∈ G are such that x*y and x'*y are defined and equal, then x = x'; (cancellation property)
• for all x ∈ G there exist elements ex (left unit of x), ex' (right unit of x) and x-1 ("inverse" of x) such that:
• ex*x = x
• x* ex' = x
• x-1*x = ex'.
From these properties it follows that:
• x* x-1 = ex, i.e. that that ex is right unit for x-1,
• ex' is left unit for x-1
• ex and ex' are idempotents, i.e. ex* ex = ex and ex'* ex' = ex'.
The concept of groupoid as defined here was introduced by Brandt (1927). An alternative meaning of groupoid was introduced by Hausmann & Ore (1937) as a set on which binary operations act but neither the identity nor the inversion are included. For this second meaning nowadays the term magma is used instead (Bourbaki, 1998).
H centered cell
The H centered cell (triple hexagonal cell) is an alternative description of the hexagonal Bravais lattice. From the conventional hP cell one obtains the hH cell by taking the new basis vectors by means of one of the following transformation matrices, which give three possible orientations of the hH cell with respect to the hP cell:
H1: 110/-120/001 H2: 2-10/110/001 H3: 1-20/2-10/001
The resulting hH cell has centering nodes at 1/3,2/3,0 and 2/3,1/3,0.
Secondary and tertiary elements in the hP cell are exchanged in the hH cell. For example, the space-group symbol P3m1 become H31m when the triple cell is used.
Hemihedry
The point group of a crystal is called hemihedry if it is a subgroup of index 2 of the point group of its lattice. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Group.txt |
The point group of a crystal is called holohedry if it is identical to the point group of its lattice. In the three-dimensional space, there are seven holohedral geometric crystal classes: ${\bar 1}, 2/m, mmm, {\bar 3}m, 4/m mm, 6/m mm, m{\bar 3}m$
Image
Let X and Y be sets, f be the function f : XY, and x be some member of X. Then the image of x under f, denoted f(x), is the unique member y of Y that f associates with x.
The image of a subset AX under f is the subset of Y defined by f[A] = {yY | y = f(x) for some xA}.
Incommensurate composite crystal
An incommensurate composite crystal is a compound with two or more (N) subsystems that are themselves modulated structures, with basis structures that are mutually incommensurate. Each subsystem (numbered by ν) has a reciprocal lattice for its basic structure with three basis vectors $a_i^{*\nu}$. There is a basis of the vector module of diffraction spots that has at most 3N basis vectors $A_j^*$ such that
$a_i^{*\nu}~=~\sum_{j=1}^n Z_{ij}^{\nu} A_j^* ~~~(i=1,2,3), \nonumber$
where $Z_{ij}^{\nu}$ are integer coefficients. If n is larger than the dimension of space (three), the composite crystal is an aperiodic crystal. n is the rank of the vector module.
Incommensurate magnetic structure
An incommensurate magnetic structure is a structure in which the magnetic moments are ordered, but without periodicity that is commensurate with that of the nuclear structure of the crystal. In particular, the magnetic moments have a spin density with wave vectors that have at least one irrational component with respect to the reciprocal lattice of the atoms. Or, in the case of localized moments, the spin function S(n+rj) (where the jth atom has position rj in the unit cell) has Fourier components with irrational indices with respect to the reciprocal lattice of the crystal.
Incommensurate modulated structure
An incommensurate modulated crystal structure is a modulated crystal structure, for which the modulation function has a Fourier transform of sharp peaks at wave vectors that cannot all be expressed by rational coefficients in a basis of the reciprocal lattice of the basic structure. At least one of the components of the wave vectors of the modulation with respect to the basis structure should be irrational.
Lattice
A lattice in the vector space Vn is the set of all integral linear combinations t = u1a1 + u2a2 + ... + ukak of a system (a1, a2, ... , ak) of linearly independent vectors in Vn.
If k = n, i.e. if the linearly independent system is a basis of Vn, the lattice is often called a full lattice. In crystallography, lattices are almost always full lattices, therefore the attribute "full" is usually suppressed.
Lattice complex
A lattice complex is the set of all point configurations that may be generated within one type of Wyckoff set. All Wyckoff positions, Wyckoff sets and types of Wyckoff sets that generate the same set of point configurations are assigned to the same lattice complex.
Concretely, two Wyckoff positions are assigned to the same lattice complex if there is a suitable transformation that maps the point configurations of the two Wyckoff positions onto each other and if their space groups belong to the same crystal family. The 72 (in E2) or 1731 (in E3) Wyckoff positions are classified in 51 (E2) or 1128 (E3) types of Wyckoff sets. They are assigned to 30 (E2) or 402 (E3) lattice complexes.
The name lattice complex comes from the fact that an assemblage of points that are equivalent with respect to a group of symmetry operations including lattice translations can be visualized as a set of equivalent lattices.
Lattice system
A lattice system of space groups contains complete Bravais flocks. All those Bravais flocks which intersect exactly the same set of geometric crystal classes belong to the same lattice system.
All those Bravais flocks belong to the same lattice system for which the Bravais classes belong to the same (holohedral) geometric crystal class.
Laue classes
The Laue classes correspond to the eleven centrosymmetric crystallographic point groups. When absorption is negligible and Friedel's law applies, it is impossible to distinguish by diffraction between a centrosymmetric point group and one of its non-centrosymmetric subgroups.
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Laue classes Non-centrosymmetric groups
having the same Laue class
${\bar 1}$ 1
2/m 2, m
mmm 222, 2mm
${\bar 3}$ 3
${\bar 3}m$ 32, 3m
4/m 4, ${\bar 4}$
4/mmm 422, ${\bar 4}2m$, 42m
6/m 6, ${\bar 6}$
6/mmm 622, ${\bar 6}2m$, 62m
$m{\bar 3}$ 23
$m{\bar 3}m$ 432, ${\bar 4}$32
Limiting complex
A limiting complex is a lattice complex L1 which forms a true subset of a second lattice complex L2. Each point configuration of L1 also belongs to L2.
L2 is called a comprehensive complex of L1.
Local symmetry
A motion of En mapping onto itself a subdomain of a crystal pattern but not the whole crystal pattern is called a local symmetry operation. It may be crystallographic or noncrystallographic depending on whether or not it is possible to extend the subdomain to an n-dimensional crystal pattern invariant under the motion.
Mapping
The term mapping is often used in mathematics as a synonym of function. In crystallography it is particularly used to indicate a transformation. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Holohedry.txt |
Merohedral is the adjectival form of merohedry and indicates a crystal that does not possess the full point symmetry of its lattice.
Merohedry
The point group of a crystal is called merohedry if it is a subgroup of the point group of its lattice.
Mesh
In a two-dimensional pattern possessing rotational symmetry, the rotation points constitute the nodes of a net and divide the plane into regions that are called meshes. The number of meshes meeting at any rotation point equals twice the order of the rotation at that point.
N.B. The term mesh is sometimes improperly used to indicate a two-dimensional cell.
Modulated crystal structure
A modulated crystal structure is a density (or atom arrangement) that may be obtained from a density (or atom arrangement) with space-group symmetry by a finite density change (or finite displacement of each atom, respectively) that is (quasi)periodic. A function or a displacement field is periodic if it is invariant under a lattice of translations. Then its Fourier transform consists of δ-peaks on a reciprocal lattice that spans the space and is nowhere dense. A quasiperiodic function has a Fourier transform consisting of δ-peaks on a vector module of finite rank. This means that the peaks may be indexed with integers using a finite number of basis vectors . If the modulation consists of deviations from the basic structure in the positions, the modulation is displacive (displacive modulation). When the probability distribution deviates from that in the basic structure the modulation is occupational.
See also
Model for a displacively modulated crystal structure . The basic structure is two-dimensional rectangular, with lattice constants a and b, the modulation wave vector is in the b-direction, the wavelength of the periodic modulation is λ such that λ/b is an irrational number.
Normal subgroup
A subgroup H of a group G is normal in G (H \(\triangleleft\) G) if gH = Hg for any g ∈G. Equivalently, H ⊂ G is normal if and only if gHg-1= H for any g ∈G, i.e., if and only if each conjugacy class of G is either entirely inside H or entirely outside H. This is equivalent to say that H is invariant under all inner automorphisms of G.
The property gH = Hg means that left and rights cosets of H in G coincide. From this one sees that the cosets form a group with the operation g1H * g2H = g1g2H which is called the factor group or quotient group of G by H, denoted by G/H.
In the special case that a subgroup H has only two cosets in G (namely H and gH for some g not contained in H), the subgroup H is always normal in G.
Normalizer
Given a group G and one of its supergroups S, they are uniquely related to a third, intermediated group NS(G), called the normalizer of G with respect to S. NS(G) is defined as the set of all elements S ∈ S that map G onto itself by conjugation:
NS(G) := {S ∈S | S-1GS = G}
The normalizer NS(G) may coincide either with G or with S or it may be a proper intermediate group. In any case, G is a normal subgroup of its normalizer.
OD structure
OD structures consist of slabs with their own symmetry, containing coincidence operations constituting a diperiodic group(layer group) only within individual slabs. For the entire structure these coincidence operations are only local (partial), i.e. they are valid only in a subspace of the crystal space. The ambiguity (= existence of more than one equivalent possibilities) in the stacking of slabs arises from the existence of this local symmetry, which does not appear in the space group of the structure. The resulting structure can be "ordered" (periodic) or "disordered" (non-periodic), depending on the sequence of local symmetry operations relating pairs of slabs. The set of all the operations valid in the whole crystal space constitutes a space group; by adding the set of all the operations valid in a subspace of it, one obtains a space groupoid.
In the OD theory, a central role is played by the vicinity condition (VC), which states the geometrical equivalence of layer pairs. The vicinity condition consists of three parts:
• VC α: VC layers are either geometrically equivalent or, if not, they are relatively few in kind
• VC β: translation groups of all VC layers are either identical or they have a common subgroup
• VC γ: equivalent sides of equivalent layers are faced by equivalent sides of adjacent layers so that the resulting pairs are equivalent.
If the position of a layer is uniquely defined by the position of the adjacent layers and by the VC, the resulting structure is fully ordered. If, on the other hand, more than one position is possible that obeys the VC, the resulting structure is an OD structure and the layers are OD layers. VC structures may thus be either fully ordered structures or OD structures. All OD structures are polytypic; the reverse may or may not be true. Equivalency depends on the choice of OD layers and also on the definition of polytypism.
Ogdohedry
The point group of a crystal is called ogdohedry if it is a subgroup of index 8 of the point group of its lattice.
In the three dimensional space there is only one ogdohedry: it corresponds to the geometric crystal class 3 of crystals belonging to the hexagonal lattice system (in case rhombohedral crystals, it corresponds instead to a tetartohedry). | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Merohedral.txt |
If G is a group consisting of a finite number of elements, this number of elements is the order of G. For example, the point group `m3m` has order 48.
For an element g of a (not necessarily finite) group G, the order of g is the smallest integer n such that gn is the identity element of G. If no such integer exists, g is of infinite order. For example, the rotoinversion `3` has order 6 and a translation has infinite order. An element of order 2 is called an involution.
Partial symmetry
The symmetry operations of a space group are isometries operating on the whole crystal pattern and are also called total operations or global operations. More generally, the crystal space can be divided in N components S1 to SN, and a coincidence operation φ(Si)→Sj can act on just the i-th component Si to bring it to coincide with the j-th component Sj. Such an operation is not one of the operations of the space group of the crystal because it is not a coincidence operation of the whole crystal space; it is not even defined, in general, for any component k different from i. It is called a partial operation: from the mathematical viewpoint, partial operations are space-groupoid operations.
When i = j, i.e. when the operation is φ(Si)→Si and brings a component to coincide with itself, the partial operation is of special type and is called local. A local operation is in fact a symmetry operation, which is defined only on a part of the crystal space: local operations may constitute a subperiodic group.
Patterson methods
The family of methods employed in structure determination to derive relationships between the scattering centers in a crystal lattice when the diffraction phases are unknown. They depend upon interpretation of the Patterson function
$P(uvw) = \dfrac{1}{V} \sum_h\sum_k\sum_l { | F(hkl) | ^2\cos[2\pi(hu + kv + lw)]} \nonumber$
to reveal interatomic vectors within the unit cell.
Patterson vector
A real-space vector representing the difference between two position vectors locating scattering centers in a diffracting crystal lattice .
Point configuration
The concept of point configuration is closely related to that of crystallographic orbit, but differs from it by the fact that point configurations are detached from their generating space groups. The concept of point configuration is the basis for the definition of lattice complexes.
Two crystallographic orbits are said configuration-equivalent if and only if their sets of points are identical. A point configuration is the set of all points that is common to a class of configuration-equivalent crystallographic orbits.
This definition uniquely assigns crystallographic orbits to point configurations but not vice versa.
The inherent symmetry of a point configuration is the most comprehensive space group that maps the point configuration onto itself. One crystallographic orbit out of each class of configuration-equivalent ones stands out because its generating space group coincides with the inherent symmetry of its point configuration.
Point group
A point group is a group of symmetry operations all of which leave at least one point unmoved. A crystallographic point group is a point group that maps a point lattice onto itself: in three dimensions rotations and rotoinversions are restricted to 1, 2, 3, 4, 6 and $\bar 1$, $\bar 2$ (= m), $\bar 3$, $\bar 4$, $\bar 6$ respectively.
Point space
A mathematical model of the space in which we live is the point space. Its elements are points. Objects in point space may be single points; finite sets of points like the centers of the atoms of a molecule; infinite discontinuous point sets like the centers of the atoms of an ideal crystal pattern; continuous point sets like straight lines, curves, planes, curved surfaces, etc.
Objects in point space are described by means of a coordinate system referred to point chosen as the origin O. An arbitrary point P is then described by its coordinates x, y, z.
The point space used in crystallography is a Euclidean space, i.e. an affine space where the scalar product is defined.
Crystal structures are described in point space. The vector space is a dual of the point space because to each pair of points in point space a vector in vector space can be associated.
Point symmetry
The point symmetry of a position is its site symmetry. The point symmetry, or point group of a lattice is the group of linear mappings (symmetry operations, isometries) that map the vector lattice L onto itself. Those geometric crystal classes to which point symmetries of lattices belong are called holohedries.
Polytypism
An element or compound is polytypic if it occurs in several structural modifications, each of which can be regarded as built up by stacking layers of (nearly) identical structure and composition, and if the modifications differ only in their stacking sequence. Polytypism is a special case of polymorphism: the two-dimensional translations within the layers are essentially preserved.
The complete definition is given in the Report of the International Union of Crystallography Ad-Hoc Committee on the Nomenclature of Disordered, Modulated and Polytype Structures: Acta Cryst. A40, 399-404(1984), "Nomenclature of Polytype Structures".
Primitive basis
A primitive basis is a crystallographic basis of the vector lattice L such that every lattice vector t of L may be obtained as an integral linear combination of the basis vectors, a, b, c.
In mathematics, a primitive basis is often called a lattice basis, whereas in crystallography the latter has a more general meaning and corresponds to a crystallographic basis.
Primitive cell
A primitive cell is a unit cell built on the basis vectors of a primitive basis of the direct lattice, namely a crystallographic basis of the vector lattice L such that every lattice vector t of L may be obtained as an integral linear combination of the basis vectors, a, b, c.
It contains only one lattice point and its volume is equal to the triple scalar product (a, b, c).
Non-primitive bases are used conventionally to describe centered lattices. In that case, the unit cell is a multiple cell and it contains more than one lattice point. The multiplicity of the cell is given by the ratio of its volume to the volume of a primitive cell. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Order.txt |
When more than one kind of symmetry element occurs for a given symmetry direction, the choice for defining the appropriate Hermann–Mauguin symbol for the space group is made in order of descending priority:
• m, e, a, b, c, n, d;
• rotation axes before screw axes.
There are a few exceptions to this choice:
1. For glide planes in centered monoclinic space groups, the priority rule is purposely not followed in order to bring out the relations between the three ‘cell choices’ given for each setting.
2. For orthorhombic space groups, the priority rule is applied only to the ‘standard symbol’. The symbols for the other five settings are obtained from the standard symbol by the appropriate transformations, without invoking the priority rule again.
3. Space-group types I222 and I212121 are two distinct types. Both contain parallel twofold rotation and screw axes and thus would receive the same symbol according to the priority rule. In I222, the three rotation axes and the three screw axes intersect, whereas in I212121 neither the three rotation axes nor the three screw axes intersect.
4. For space groups of type No. 73, the standard symbol Ibca was adopted, instead of Ibaa according to the rule, because Ibca displays the equivalence of the three symmetry directions clearly.
5. In tetragonal space groups with both a and b glide planes containing the [001] direction, the preference was given to b, as in P4bm.
6. In cubic space groups where tertiary symmetry planes with glide components 1/2, 0, 0; 0, 1/2, 0; 0, 0, 1/2 and 1/2, 1/2 , 1/2 coexist, the tertiary symmetry element was called n in P groups (instead of a, b or c) but c in F groups, because these symmetry elements intersect the origin.
7. Space groups of type I23 and I213 (199) are two distinct types of space groups. For this pair, the same arguments apply as given above for I222 and I212121.
Pseudo symmetry
A crystal space can in general be divided in N components S1 to SN. When a coincidence operation φ(Si)→Sj brings the i-th component Si to coincide with the j-th component Sj, for any i and j, φ is a symmetry operation of the space group.
Sometimes, φ brings Si close to, but not exactly on, the position and orientation of Sj: in this case the operation mapping Si onto Sj is not crystallographic but the linear and/or rotational deviation from a space group operation is limited. For this reason, it is preferable to describe the crystallographic operation φ as a pseudo symmetry operation.
Pseudo symmetry operations for the lattice play an important role in twinning, namely in the case of twinning by pseudomerohedry and twinning by reticular pseudomerohedry.
Reciprocal Space
The basis vectors a*, b*, c* of the reciprocal space are related to the basis vectors a, b, c of the direct space (or crystal space) through either of the following two equivalent sets of relations:
(1)
a*. a = 1; b*. b = 1; c*. c = 1;
a*. b = 0; a*. c = 0; b*. a = 0; b*. c = 0; c*. a = 0; c*. b = 0.
(2)
a* = (b × c)/ (a, b, c);
b* = (c × a)/ (a, b, c);
c* = (b × c)/ (a, b, c);
where (b × c) is the vector product of basis vectors b and c and (a, b, c) = V is the triple scalar product of basis vectors a, band c and is equal to the volume V of the cell constructed on the vectors a, b and c.
The reciprocal and direct spaces are reciprocal of one another, that is the reciprocal space associated to the reciprocal space is the direct space. They are related by a Fourier transform and the reciprocal space is also called Fourier space or phase space.
The vector product of two direct space vectors, ${\bold r_1} = u_1 {\bold a} + v_1 {\bold b} + w_1 {\bold c}$ and ${\bold r_2} = u_2 {\bold a} + v_2 {\bold b} + w_2 {\bold c}$ is a reciprocal space vector,
${\bold r*} = {\bold r_1} \times {\bold r_2} = V (v_1 w_2 - v_2 w_1) {\bold a*} + V (w_1 u_2 - w_2 u_1) {\bold b*} + V (u_1 v_2 - u_2 v_1) {\bold c}.$
Reciprocally, the vector product of two reciprocal vectors is a direct space vector.
As a consequence of the set of definitions (1), the scalar product of a direct space vector r = u a + v b + w c by a reciprocal space vector r* = h a* + k b* + l c* is simply:
r . r* = uh + vk +wl.
In a change of coordinate system, The coordinates of a vector in reciprocal space transform like the basis vectors in direct space and are called for that reason covariant. The vectors in reciprocal transform like the coordinates in direct space and are called contravariant. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Priority_rule.txt |
The reciprocal lattice is constituted by the set of all possible linear combinations of the basis vectors a*, b*, c* of the reciprocal space. A point (node), H, of the reciprocal lattice is defined by its position vector:
OH = rhkl* = h a* + k b* + l c*.
If H is the nth node on the row OH, one has:
OH = n OH1 = n (h1 a* + k1 b* + l1 c*),
where H1 is the first node on the row OH and h1 , k1 , l1 are relatively prime.
The generalization of the reciprocal lattice in a four-dimensional space for incommensurate structures is described in Section 9.8 of International Tables of Crystallography, Volume C.
Semidirect product
In group theory, a semidirect product describes a particular way in which a group can be put together from two subgroups, one of which is normal.
Let G be a group, N a normal subgroup of G (i.e., NG) and H a subgroup of G. G is a semidirect product of N and H if there exists a homomorphism GH which is the identity on H and whose kernel is N. This is equivalent to say that:
• G = NH and NH = {1} (where "1" is identity element of G )
• G = HN and NH = {1}
• Every element of G can be written as a unique product of an element of N and an element of H
• Every element of G can be written as a unique product of an element of H and an element of N
One also says that "G splits over N".
Site symmetry
The site-symmetry group (often called point symmetry) of a point is the finite group formed by the set of all symmetry operations of the space group of the crystal that leave that point invariant. It is isomorphic to a (proper or improper)subgroup of the point group to which the space group under consideration belongs. In general, the origin is a point of highest site symmetry.
Space group
The symmetry group of a three-dimensional crystal pattern is called its space group. In E2, the symmetry group of a two-dimensional crystal pattern is called its plane group. In E1, the symmetry group of a one-dimensional crystal pattern is called its line group.
To each crystal pattern belongs an infinite set of translations T, which are symmetry operations of that pattern. The set of all T forms a group known as the translation subgroup T of the space group G of the crystal pattern. T is an Abelian group.
Stabilizer
Let G be a group which acts on a set A by a composition law *, and let a be a given element of A. Then the set: Ga = {g ∈ G | a*g = a} is called the stabilizer of A. Ga is the set of all elements of G which leave a unchanged or 'stable'. Ga is a subgroup of G.
Statistical descriptors
A separate online document Statistical Descriptors in Crystallography prepared for the International Union of Crystallography provides an authoritative statement on some aspects of the use of statistics and statistical techniques in crystallography, principally that of least-squares refinement of diffraction data against an atomic model. The following topics are treated:
• Glossary
• Basic notions
• Uncertainty of measurement
• Refinement
• Refinement on I, |F|2 or |F|?
• Defects in the model
• Weighting schemes
• Recommendations | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Reciprocal_lattice.txt |
A cell C' obtained from a cell C by adding one or more translation vectors of fractional periodicity in C is called a subcell of C. The translation subgroup T' of the lattice L' built on C' is a supergroup of the translation subgroup T of the lattice L built on C and corresponds therefore to superlattice of L.
Subgroup
Let G be a group and H a non-empty subset of G. Then H is called a subgroup of G if the elements of H obey the group postulates, i.e. if
1. the identity element 1G of G is contained in H;
2. H is closed under the group operation (inherited from G);
3. H is closed under taking inverses.
The subgroup H is called a proper subgroup of G if there are elements of G not contained in H.
A subgroup H of G is called a maximal subgroup of G if there is no proper subgroup M of G such that H is a proper subgroup of M.
Sublattice
A lattice L' obtained by another lattice L by removing one or more sets of nodes is called a sublattice of L. The translation subgroup T' of L' is a subgroup of the translation subgroup T of L. The unit cell of L' is larger than the unit cell of L and is therefore called a supercell.
Subperiodic group
A subperiodic group is a group of Euclidean mappings such that its translations form a lattice in a proper subspace of the space on which it acts.
A crystallographic subperiodic group in n-dimensional space is a subperiodic group for which the group of linear parts is a crystallographic point group of n-dimensional space. The crystallographic subperiodic groups in two and three-dimensional space are classified in:
• frieze groups: 7 two-dimensional groups with one-dimensional translations;
• rod groups: 75 three-dimensional groups with one-dimensional translations;
• layer groups: 80 three-dimensional groups with two-dimensional translations.
Supercell
A cell C' obtained from a cell C by removing one or more translation vectors is called a supercell of C. The translation subgroup T' of the lattice L' built on C' is a subgroup of the translation subgroup T of the lattice L built on C and corresponds therefore to sublattice of L.
Supergroup
If G is a group and H is a subgroup of G, then G is a supergroup of H.
If H is a maximal subgroup of G, then G is a minimal supergroup of H.
Superlattice
A lattice L' obtained by another lattice L by adding one or more sets of nodes is called a superlattice of L. The translation subgroup T' of L' is a supergroup of the translation subgroup T of L. The unit cell of L' is smaller than the unit cell of L and is therefore called a subcell.
Symmetry element
A symmetry element (of a given crystal structure or object) is defined as a concept with a double meaning, namely the combination of a geometric element with the set of symmetry operations having this geometric element in common (termed its element set). Together with the identity and the translations, for which a geometric element is not defined, the element sets cover all symmetry operations. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Subcell.txt |
A symmetry operation is an isometry, i.e. a transformation under which two objects, or two configurations or an object, are brought to coincide. A symmetry operation is a Euclidean mapping: to each point of the first configuration there corresponds a point of the second configuration, the distances between two points are kept by the transformation, as are the angles.
The two configurations/objects can be either congruent or enantiomorphous. Correspondingly, the symmetry operations are classed into two kinds:
• symmetry operations of first kind: they relate congruent objects and consist of translations, rotations and screw rotations;
• symmetry operations of second kind: they relate enantiomorphous objects and consist of inversion, reflections, rotoinversions, and glide reflections. There exist a 1:1 correspondence between rotoinversion and rotoreflections: the latter are more used in Schoenflies notation, whereas rotoinversions are preferred in Hermann-Mauguin notation.
A symmetry operation is performed about a symmetry element.
Symmorphic space groups
A space group is called ‘symmorphic’ if, apart from the lattice translations, all generating symmetry operations leave one common point fixed. Permitted as generators are thus only the point-group operations: rotations, reflections, inversions and rotoinversions. The symmorphic space groups may be easily identified because their Hermann-Mauguin symbol does not contain a glide or screw operation. The combination of the Bravais lattices with symmetry elements with no translational components yields the 73 symmorphic space groups, e.g. P2, Cm, P2/m, P222, P32, P23. They are in one to one correspondence with the arithmetic crystal classes.
A characteristic feature of a symmorphic space group is the existence of a special position, the site-symmetry group of which is isomorphic to the point group to which the space group belongs. Symmorphic space groups have no zonal or serial reflection conditions, but may have integral reflection conditions (e.g. C2, Fmmm).
Tetartohedry
The point group of a crystal is called tetartohedry if it is a subgroup of index 4 of the point group of its lattice.
Unit cell
The unit cell is the parallelepiped built on the vectors, a, b, c, of a crystallographic basis of the direct lattice. Its volume is given by the scalar triple product, V = (a, b, c) and corresponds to the square root of the determinant of the metric tensor.
If the basis is primitive, the unit cell is called the primitive cell. It contains only one lattice point. If the basis is non-primitive, the unit cell is a multiple cell and it contains more than one lattice point. The multiplicity of the cell is given by the ratio of its volume to the volume of a primitive cell.
Vector module
A vector module is the set of vectors spanned by a number n of basis vectors with integer coefficients. The basis vectors should be independent over the integers, which means that any linear combination
miai
i
with mi integers is equal to zero if, and only if, all coefficients mi are zero. The term Z-module is sometimes used to underline the condition that the coefficients are integers. The number of basis vectors is the rank of the vector module.
Vector space
For each pair of points X and Y in point space one can draw a vector r from X to Y. The set of all vectors forms a vector space. The vector space has no origin but instead there is the zero vector which is obtained by connecting any point X with itself. The vector r has a length which is designed by |r| = r, where r is a non–negative real number. This number is also called the absolute value of the vector. The maximal number of linearly independent vectors in a vector space is called the dimension of the space.
An essential difference between the behavior of vectors and points is provided by the changes in their coefficients and coordinates if a different origin in point space is chosen. The coordinates of the points change when moving from an origin to the other one. However, the coefficients of the vector r do not change.
The point space is a dual of the vector space because to each vector in vector space a pair of points in point space can be associated.
Face normals, translation vectors, Patterson vectors and reciprocal lattice vectors are elements of vector space. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Symmetry_operation.txt |
The Voronoi domain (or 'cell', or 'region') is the name given in mathematics to the Wigner-Seitz cell. Voronoi domains are in the form of polyhedra and are classified according to their topological properties ; one distinguishes five types of Voronoi polyhedra (see Section 9.1.6 of International Tables of Crystallography, Volume A.
Weissenberg complex
A Weissenberg complex is a lattice complex for which the multiplicity does not decrease for any special values of the coordinates. Weissenberg complexes can simulate invariant lattice complexes as limiting complexes.
Wigner-Seitz cell
The Wigner-Seitz cell is a a polyhedron obtained by connecting a lattice point P to all other lattice points and drawing the planes perpendicular to these connecting lines and passing through their midpoints (Figure 1). The polyhedron enclosed by these planes is the Wigner-Seitz cell. This construction is called the Dirichlet construction. The cell thus obtained is a primitive cell and it is possible to fill up the whole space by translation of that cell.
The Wigner-Seitz cell of a body-centered cubic lattice I is a cuboctahedron (Figure 2) and the Wigner-Seitz cell of a face-centered cubic lattice F is a rhomb-dodecahedron (Figure 3). In reciprocal space this cell is the first Brillouin zone. Since the reciprocal lattice of body-centered lattice is a face-centered lattice and reciprocally, the first Brillouin zone of a body-centered cubic lattice is a rhomb-dodecahedron and that of a face-centered cubic lattice is a cuboctahedron.
The inside of the Wigner-Seitz cell has been called domain of influence by Delaunay (1933). It is also called Dirichlet domain or Voronoi domain. The domain of influence of lattice point P thus consists of all points Q in space that are closer to this lattice point than to any other lattice point or at most equidistant to it (such that OP≤ |t - OP| for any vector tL).
Wyckoff position
A Wyckoff position of a space group G consists of all points X for which the site-symmetry groups are conjugate subgroups of G.
Each Wyckoff positon of a space group is labeled by a letter which is called the Wyckoff letter.
The number of different Wyckoff positions of each space group is finite, the maximal numbers being 9 for plane groups (realized in p2mm) and 27 for space groups (realized in Pmmm).
There is a total of 72 Wyckoff positions in plane groups and 1731 Wyckoff positions in space groups.
The transfer of Wyckoff positions from individual space groups to space-group types is not unique because Wyckoff positions with the same type of site-symmetry group may be exchanged in different space groups of the same type. This is no longer true when one makes use of Wyckoff sets.
Wyckoff set
A Wyckoff set with respect to a space group G is the set of all points X for which the site-symmetry groups are conjugate subgroups of the normalizer N of G in the group of all affine mappings.
Any Wyckoff position of G is transformed onto itself by all elements of G, but not necessarily by the elements of N. Any Wyckoff set of G is instead transformed onto itself also by those elements of N that are contained in G.
Lepton
Crystallography (from the Greek words crystallon = cold drop / frozen drop, with its meaning extending to all solids with some degree of transparency, and graphein = write) is the experimental science of determining the arrangement of atoms in solids. In older usage, it is the scientific study of crystals.
History of crystallography
Lepton is a term introduced by Friedrich Rinne in his book Crystals and the Fine-Structure of Matter(London: Methuen & Co., 1924, English translation by S. W. Stiles) to indicate electrons, atoms, ions, radicals, and molecules building the unit formula of a crystalline compound. The crystal structure of a solid was sometimes called "leptonic structure".
Today, the term lepton is used to indicate spin-1/2 particles and Rinne's use is only of historical interest. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Fundamental_Crystallography/Voronoi_domain.txt |
The law of rational indices states that the intercepts, OP, OQ, OR, of the natural faces of a crystal form with the unit-cell axes a, b, c (see Figure 1) are inversely proportional to prime integers, h, k, l. They are called the Miller indices of the face. They are usually small because the corresponding lattice planes are among the densest and have therefore a high interplanar spacing and low indices.
The law of rational indices was deduced by Haüy (1784, 1801) from the observation of the stacking laws required to build the natural faces of crystals by piling up elementary blocks, for instance cubes to construct the {110} faces of the rhomb-dodecahedron observed in garnets or the ½{210} faces of the pentagon-dodecahedron observed in pyrite, or rhombohedrons to construct the {21.1} (referred to an hexagonal lattice, $\{2\bar{1}0\}$, referred to a rhombohedral lattice) scalenohedron of calcite.
(Models from Haüy's Traité de Minéralogie (1801) - the crystal forms have been redrawn in red).
Law of the constancy of interfacial angles
The law of the constancy of interfacial angles (or 'first law of crystallography') states that the angles between the crystal faces of a given species are constant, whatever the lateral extension of these faces and the origin of the crystal, and are characteristic of that species. It paved the way for Haüy's law of rational indices.
Polar Lattice
The polar lattice is a lattice dual of the direct lattice, which is the ancestor of the reciprocal lattice. It was introduced by Auguste Bravais in a " mémoire" presented to the Académie de Sciences de Paris on 11 December 1848.
The construction of the polar lattice is essentially the same as that of the reciprocal lattice, but the parameter along a row of the polar lattice is V2/3/d(hkl) instead of 1/d(hkl). The polar lattice has thus the same dimensions as the direct lattice, namely Ångstroms, instead of Ångstroms-1, like the reciprocal lattice.
• The unit cell of the polar lattice has the same volume as that of the direct lattice.
• The scalar product of the basis vectors of the direct and polar lattice is V2/3δij, where δ is Kroneker's tensor and the indices i and j point to the basis vectors.
The polar lattice was introduced to facilitate the morphological study of crystals.
Zone axis
A zone axis is a lattice row parallel to the intersection of two (or more) families of lattices planes. It is denoted by [u v w]. A zone axis [u v w] is parallel to a family of lattice planes of Miller indices (hkl) if:
$uh + vk + wl = 0 \nonumber$
This is the so-called Weiss law.
The indices of the zone axis defined by two lattice planes (h1,k1,l1), (h2,k2,l2) are given by:
$\frac{u}{\begin{vmatrix} k_1 &l_1 \ k_2 &l_2 \end{vmatrix}}=\frac{v}{\begin{vmatrix} l_1 &h_1 \ l_2 &h_2 \end{vmatrix}}=\frac{w}{\begin{vmatrix} h_1 &k_1 \ h_2 &k_2\ \end{vmatrix}} \nonumber$
Conversely, any crystal face can be determined if one knows two zone axes parallel to it. It is the zone law, or Zonenverbandgesetz.
Three lattice planes have a common zone axis (are in zone) if their Miller indices (h1,k1,l1), (h2,k2,l2), (h3,k3,l3) satisfy the relation:
$\begin{vmatrix} h_1 & k_1 & l_1\ h_2 & k_2 & l_2\ h_3 & k_3 & l_3\ \end{vmatrix} = 0$
$\begin{vmatrix} h_1 & k_1 & l_1\h_2 & k_2 & l_2\h_3 & k_3 & l_3\\end{vmatrix}=0 \nonumber$ | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Morphological_crystallography/Law_of_Rational_Indices.txt |
A measure of the scattering power of an isolated atom (also known as the atomic form factor). The scattering factor depends on the scattering amplitude of an individual atom and also the Bragg angle of scattering. It depends on the type of radiation involved .
X-ray scattering
The scattering from a crystal of an X-ray beam results from the interaction between the electric component of the incident electromagnetic radiation and the electrons in the crystal. Tightly bound electrons scatter coherently (Rayleigh scattering); free electrons scatter incoherently (Compton scattering). The scattering process from atomic electrons in a crystal lattice has both coherent and incoherent components, and is described as Thomson scattering.
The scattering amplitude from a neutral atom depends on the number of electrons (\Z\) = the atomic number) and also on the Bragg angle $\theta$ – destructive interference among waves scattered from the individual electrons reduces the intensity at other than zero scattering angle. For $\theta=0$ the scattering amplitude is normally equal to $Z$. However, the scattering factor is modified by anomalous scattering if the incident wavelength is near an absorption edge of the scattering element.
The X-ray scattering factor is evaluated as the Fourier transform of the electron density distribution of an atom or ion, which is calculated from theoretical wavefunctions for free atoms.
See also
1. Electron diffraction. C. Colliex, J. M. Cowley, S. L. Dudarev, M. Fink, J. Gjønnes, R. Hilderbrandt, A. Howie, D. F. Lynch, L. M. Peng, G. Ren, A. W. Ross, V. H. Smith Jr, J. C. H. Spence, J. W. Steeds, J. Wang, M. J. Whelan and B. B. Zvyagin. International Tables for Crystallography(2006). Vol. C, ch. 4.3, pp. 259-429
2. Intensity of diffracted intensities. P. J. Brown, A. G. Fox, E. N. Maslen, M. A. O'Keefe and B. T. M. Willis. International Tables for Crystallography (2006). Vol. C, ch. 6.1, pp. 554-595
3. Neutron techniques. I. S. Anderson, P. J. Brown, J. M. Carpenter, G. Lander, R. Pynn, J. M. Rowe, O. Schärpf, V. F. Sears and B. T. M. Willis. International Tables for Crystallography (2006). Vol. C, ch. 4.4, pp. 430-487
Curie laws
Curie extended the notion of symmetry to include that of physical phenomena and stated that:
• the symmetry characteristic of a phenomenon is the highest compatible with the existence of the phenomenon;
• the phenomenon may exist in a medium which possesses that symmetry or that of a subgroup of that symmetry.
and concludes that some symmetry elements may coexist with the phenomenon but that their presence is not necessary. On the contrary, what is necessary is the absence of certain symmetry elements: ‘asymmetry creates the phenomenon’. Noting that physical phenomena usually express relations between a cause and an effect (an influence and a response), P. Curie restated the two above propositions in the following way, now known as Curie laws, although they are not, strictly speaking, laws (Curie himself spoke about 'the principle of symmetry'):
• the asymmetry of the effects must pre-exist in the causes;
• the effects may be more symmetric than the causes.
Cylindrical system
The cylindrical system contains non-crystallographic point groups with one axis of revolution (or isotropy axis). There are five groups in the spherical system:
Hermann-Mauguin symbol Short Hermann-Mauguin symbol Schönfliess symbol order of the group general form
$A_{\infty} \nonumber$ $\infty \nonumber$ $C_{\infty} \nonumber$ $\infty \nonumber$ rotating cone
$\frac{A_{\infty}}{M}C \nonumber$ $\bar\infty \nonumber$
$C_{\infty\,h}\equiv\,S_{\infty}\equiv\,C_{\infty\,i} \nonumber$
$\infty \nonumber$
rotating finite cylinder
$A_{\infty}\infty\,A_2 \nonumber$ $\infty2 \nonumber$ $D_{\infty} \nonumber$ $\infty \nonumber$ finite cylinder
submitted to equal and
opposite torques
$A_{\infty}M \nonumber$ $\infty\,m \nonumber$
$C_{\infty\,v} \nonumber$
$\infty \nonumber$ stationary cone
$\frac{A_{\infty}}{M}\frac{\infty\,A_2}{\infty\,M}C \nonumber$
$\bar\infty\,m\equiv\bar\infty\frac{2}{m} \nonumber$
$D_{\infty\,h}\equiv\,D_{\infty\,d} \nonumber$
$\infty \nonumber$ stationary finite cylinder
Note that $A_{\infty}M$ represents the symmetry of a force, or of an electric field and that $\frac{A_{\infty}}{M}C$ represents the symmetry of a magnetic field (Curie 1894), while $\frac{A_{\infty}}{M}\frac{\infty\,A_2}{\infty\,M}C$ represents the symmetry of a uniaxial compression.
Dual basis
The dual basis is a basis associated to the basis of a vector space. In three-dimensional space, it is isomorphous to the basis of the reciprocal lattice. It is mathematically defined as follows:
Given a basis of n vectors ei spanning the direct space En, and a vector x = x i ei, let us consider the n quantities defined by the scalar products of x with the basis vectors, ei:
xi = x . ei = x j ej . ei = x j gji,
where the gji 's are the doubly covariant components of the metric tensor.
By solving these equations in terms of x j, one gets:
x j = xi gij
where the matrix of the gij 's is inverse of that of the gij 's (gikgjk = δij). The development of vector x with respect to basis vectors ei can now also be written:
x = x i ei = xi gij ej
The set of n vectors ei = gij ej that span the space En forms a basis since vector x can be written:
x = xi ei
This basis is the dual basis and the n quantities xi defined above are the coordinates of x with respect to the dual basis. In a similar way one can express the direct basis vectors in terms of the dual basis vectors:
ei = gij ej
The scalar products of the basis vectors of the dual and direct bases are:
gij = ei . ej = gik ek . ej = gikgjk = δij.
One has therefore, since the matrices gik and gij are inverse:
gij = ei . ej = δij.
These relations show that the dual basis vectors satisfy the definition conditions of the reciprocal vectors. In a three-dimensional space the dual basis and the basis of reciprocal space are identical.
Electrocaloric effect
The electrocaloric effect is the converse of the pyroelectric effect: it describes the variation of entropy δσ of a material submitted to an applied electrical field Ei :
δσ = piT Ei
where piT is the electrocaloric coefficient at constant stress. It is equal to the pyroelectric coefficient. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Physical_Properties_of_Crystals/Atomic_Scattering_Factor.txt |
A metric tensor is used to measure distances in a space. In crystallography the spaces considered are vector spaces with Euclidean metrics, i.e.ones for which the rules of Euclidean geometry apply. In that case, given a basis ei of a Euclidean space, En, the metric tensor is a rank 2 tensor the components of which are:
gij = ei . ej = ej.ei = gji.
It is a symmetrical tensor. Using the metric tensor, the scalar product of two vectors, x = xi ei and y = yj ej is written:
x . y = xi ei . yj ej = gij xi yj.
In a three-dimensional space with basis vectors a, b, c, the coefficients gij of the metric tensor are:
g11 = a2; g12 = a . b; g13 = a . c;
g21 = b . a; g22 = b2; g23 = b . c;
g31 = c . a; g32 = c . b; g33 = c2;
Because the metric tensor is symmetric, g12 = g21, g13 = g31, and g13 = g31. Thus there are only six unique elements, often presented as
g11 g22 g33
g23 g13 g12
or, multiplying the second row by 2, as a so-called G6 ("G" for Gruber) vector
( a2, b2, c2, 2 b . c, 2 a . c, 2 a . b )
The inverse matrix of gij, gij, (gikgkj = δkj, Kronecker symbol, = 0 if ij, = 1 if i = j) relates the dual basis, or reciprocal space vectors ei to the direct basis vectors ei through the relations:
ej = gij ej
In three-dimensional space, the dual basis vectors are identical to the reciprocal space vectors and the components of gij are:
g11 = a*2; g12 = a* . b*; g13 = a* . c*;
g21 = b* . a*; g22 = b*2; g23 = b* . c*;
g31 = c* . a*; g32 = c* . b*; g33 = c*2;
with:
g11 = b2c2 sin2 α/ V2; g22 = c2a2 sin2 β/ V2; g33 = a2b2 sin2 γ/ V2;
g12 = g21 = (abc2/ V2)(cos α cos β - cos γ); g23 = g32 = (a2bc/ V2)(cos β cos γ - cos α); g31 = g13 = (ab2c/ V2)(cos γ cos α - cos β)
where V is the volume of the unit cell (a, b, c).
Change of basis
In a change of basis the direct basis vectors and coordinates transform like:
e'j = Aj i ei; x'j = Bi j x i,
where Aj i and Bi j are transformation matrices, transpose of one another. According to their definition, the components gij, of the metric tensor transform like products of basis vectors:
g'kl = AkiAljgij.
They are the doubly covariant components of the metric tensor.
The dual basis vectors and coordinates transform in the change of basis according to:
e'j = Bi j ei; x'j = Aj ixi,
and the components gij transform like products of dual basis vectors:
g'kl = Aik Ajl gij.
They are the doubly contravariant components of the metric tensor.
The mixed components, gij = δij, are once covariant and once contravariant and are invariant.
Properties of the metric tensor
• The tensor nature of the metric tensor is demonstrated by the behaviour of its components in a change of basis. The components gij andgij are the components of a unique tensor.
• The squares of the volumes V and V* of the direct space and reciprocal space unit cells are respectively equal to the determinants of thegij 's and the gij 's:
V 2 = Δ (gij) = abc(1 - cos 2 α - cos 2 β - cos2 γ + 2 cos α cos α cos α)
V*2 = Δ (gij) = 1/ V 2.
• One changes the variance of a tensor by taking the contracted tensor product of the tensor by the suitable form of the metric tensor. For instance:
gimt ij..kl.. = t j..klm..
Multiplying by the doubly covariant form of the metric tensor increases the covariance by one, multiplying by the doubly contravariant form increases the contravariance by one.
See also
• Section 1.1.3 of International Tables of Crystallography, Volume B
• Section 1.1.2 of International Tables of Crystallography, Volume D | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Physical_Properties_of_Crystals/Metric_tensor.txt |
Neumann's principle, or principle of symmetry, states that, if a crystal is invariant with respect to certain symmetry elements, any of its physical properties must also be invariant with respect to the same symmetry elements, or otherwise stated, the symmetry elements of any physical property of a crystal must include the symmetry elements of the point group of the crystal. It is generalized to physical phenomena by Curie laws.
Piezoelectricity
Piezoelectricity is the property presented by certain materials that exhibit an electric polarization when submitted to an applied mechanical stress such as a uniaxial compression. Conversely, their shape changes when they are submitted to an external electric field; this is the converse piezoelectric effect. The piezoelectric effect and the converse effect are described by third-rank tensors:
• For a small stress, represented by a second-rank tensor, Tij, the resulting polarization, of components Pk , is given by:
Pk = dkijTij
where dkij is a third-rank tensor representing the direct piezoelectric effect.
• For a small applied electric field, of components Ek, the resulting strain, represented by a second-rank tensor, Sij, is given by:
Sij = dijkEk + QijklEkEl
where the first-order term, dijk, represents the inverse piezoelectric effect and the second-order term, Qijkl, a symmetric fourth-rank tensor, the electrostriction effect. The sense of the strain due to the piezoelectric effect changes when the sign of the applied electric field changes , while that due to electrostriction, a quadratic effect, does not.
The matrices associated to the coefficients dkij and dkij of the direct and converse piezoelectric effects, respectively, are transpose of one another.
Pyroelectricity
Pyroelectricity is the property presented by certain materials that exhibit an electric polarization Pi when a temperature variation δΘ is applied uniformly:
Pi = piT δΘ
where piT is the pyroelectric coefficient at constant stress. Pyroelectric crystals actually have a spontaneous polarization, but the pyroelectric effect can only be observed during a temperature change. If the polarization can be reversed by the application of an electric field, the crystal is ferroelectric.
If the crystal is also piezoelectric, the polarization due to an applied temperature variation is also partly due to the piezoelectric effect. The coefficient describing the pure pyroelectric effect is the pyroelectric coefficient at constant strain, piS. The two coefficients are related by:
piT = cijkldklnαjn + piS
where the cijkl are the elastic stiffnesses, the dkln the piezoelectric coefficients and the αjn the linear thermal expansion coefficients.
The converse effect is the electrocaloric effect. If a pyroelectric crystal is submitted to an electric field, it will undergo a change of entropy Δσ:
Δσ = pi Ei
and will release or absorb a quantity of heat given by Θ V Δσ where Θ is the temperature of the specimen and V its volume.
Spherical Systems
The spherical system contains non-crystallographic point groups with more than one axis of revolution. These groups, therefore, contain an infinity of axes of revolution (or isotropy axis). There are two groups in the spherical system:
Hermann-Mauguin symbol Short Hermann-Mauguin symbol Schönfliess symbol order of the group general form
$\infty\,A_{\infty}$ $2\infty$ K $\infty$ sphere filled with
an optically active liquid
$\infty\frac{A_{\infty}}{M}C$ $m\bar\infty,\frac{2}{m}\bar\infty$ Kh $\infty$ stationary sphere
History
The groups containing isotropy axes were introduced by P. Curie (1859-1906) in order to describe the symmetry of physical systems (Curie P. (1884). Sur les questions d'ordre: répétitions. Bull. Soc. Fr. Minéral., 7, 89-110; Curie P. (1894). Sur la symétrie dans les phénomènes physiques, symétrie d’un champ électrique et d’un champ magnétique. J. Phys. (Paris), 3, 393-415.).
See also
Section 10.1.4 of International Tables of Crystallography, Volume A
Section 1.1.4 of International Tables of Crystallography, Volume D
Preferred orientation
Powder diffraction uses X-ray, neutron, or electron diffraction on powder or microcrystalline samples for structural characterization.
Powder Diffraction
Preferred orientation arises when there is a stronger tendency for the crystallites in a powder or a texture to be oriented more one way, or one set of ways, than all others. An easily visualized case of preferred orientation is that which results when a material with a strong cleavage or growth habit is packed into a specimen or when a metal sheet is obtained by rolling.
Preferred orientation should not be confused with 'graininess' or 'inadequate powder average', in which there are so few crystallites being irradiated that the number of correctly oriented crystallites varies significantly from reflection to reflection of different types. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Physical_Properties_of_Crystals/Neumann%27s_Principle.txt |
The R factor measuring the agreement between the reflection intensities calculated from a crystallographic model and those measured experimentally.
\(R_B =
\).
In the Rietveld method RB is useful because it depends on the fit of the structural parameters and not on the profile parameters.
Constrained refinement
A refinement is said to be constrained if one or more parameters in the refinement are held fixed or are determined by the value of one or more refined parameters. Constraints related to space group symmetry are not usually counted among the constraints applied to a given refinement as they are always present where applicable.
Direct methods
The family of methods for solving the phase problem in crystal structure determination. The phases of scattered diffraction beams cannot be directly observed. However, they can be estimated from probability relationships applied to the phases of the most intense diffraction peaks. The facts that scattering centers in a crystal are discrete atoms (i.e. sources of electron density) and that the electron density must be non-negative are the types of constraints that restrict the possible values of the phases, and allow initial estimates of some of them.
Free R factor
A residual function calculated during structure refinement in the same way as the conventional R factor, but applied to a small subset of reflections that are not used in the refinement of the structural model. The purpose is to monitor the progress of refinement and to check that the R factor is not being artificially reduced by the introduction of too many parameters.
Harker section
In Patterson methods of structure solution, relationships between symmetrically related atoms produce peaks in the Patterson function on certain planes or along certain lines determined by the known crystallographic symmetries. Harker sections are portions of the Patterson map that contain a large proportion of the readily interpretable information because they contain many such Harker peaks (vectors between space-group equivalent atoms).
Heavy-Atom Method
An application of Patterson methods in crystal structure determination. For a compound containing a heavy atom (i.e. one with a significantly higher atomic scattering factor than the others present) the diffraction phases calculated from the position of the heavy atom are used to compute a first approximate electron density map. Further portions of the structure are recognizable as additional peaks in the map. Successive approximate electron density maps may then be calculated to solve the entire structure.
Phase problem
Waves diffracted from a primitive lattice of simple scatterers obey Bragg's law, which allows ready determination of interplanar distances and thus the easy recovery of a description of the crystal lattice. Where the scattering objects are complex (e.g. in molecular crystals) the diffracted radiation suffers a phase shift arising from the spatial distribution of individual scatterers. The amplitudes of the resulting structure factors are directly derivable from the experimental measured intensities of the diffracted beams, but the phases are not. Without a knowledge of the phases, it is not possible to reconstruct the individual atomic positions. Estimating the phases is an essential step in successful structure determination.
R factor
The term R factor in crystallography is commonly taken to refer to the 'conventional' R factor
\(R =
\),
a measure of agreement between the amplitudes of the structure factors calculated from a crystallographic model and those from the original X-ray diffraction data. The R factor is calculated during each cycle of least-squares structure refinement to assess progress. The final R factor is one measure of model quality.
More generally, a variety of R factors may be determined to measure analogous residuals during least-squares optimization procedures. Where the refinement attempts to minimize the deviates of the squares of the structure factors (refinement against F2), the R factor based on F2 is used to monitor the progress of refinement:
\(R(F^2) =
\).
Likewise, refinement against I can be tracked using the Bragg R factor
\(R_B =
\).
Even for refinement against F2 or I, the 'conventional' R factor may be calculated and quoted as a measure of model quality, in order to compare the resulting quality of models calculated at different times and with different refinement strategies.
The R factor is sometimes described as the discrepancy index.
Refinement
In structure determination, the process of improving the parameters of an approximate (trial) structure until the best fit is achieved between an observed diffraction pattern and that calculated by Fourier transformation from the numerically parameterized trial structure.
Restrained refinement
A refinement is said to be restrained if the refinement is based on additional observations or pseudo-observations besides the observed structure factors.
Rietveld method
Method of analyzing powder diffraction data in which the crystal structure is refined by fitting the entire profile of the diffraction pattern to a calculated profile using a least-squares approach. There is no intermediate step of extracting structure factors, and so patterns containing many overlapping Bragg peaks can be analyzed. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Structure_Determination/Bragg_R_factor.txt |
Patterson methods of structure determination use the Patterson function
P(uvw) = (1 / V) { | F(hkl) | 2cos[2π(hu + kv + lw)]}
h k l
to generate a map of interatomic vectors within the unit cell. Better results can be obtained by artificially sharpening the peaks in the Patterson function, thereby enhancing the resolution of individual peaks.
One technique for doing so, introduced by Patterson in 1935, considers the effect of thermal motion on the broadening of electron-density peaks and consequently their Patterson peaks. The F2 coefficients can be corrected for thermal effects by simulating the atoms as point scatterers and using a modified set of coefficients $|\mathbf{F}_{\mathbf{h}, \mathrm{sharp}}|^{2} = |\mathbf{F}_\mathbf{h}|^{2} / \bar{f}^{2}$, where $\bar{f}$, the average scattering factor per electron, is given by
$\bar{f} = {\sum\limits_{i = 1}^{N}}\; f_{i} / {\sum\limits_{i = 1}^{N}} Z_{i}.$
A common formulation for this type of sharpening expresses the atomic scattering factors at a given angle in terms of an overall isotropic thermal parameter B as f(s) = f0exp( − Bs2). The Patterson coefficients then become
$\mathbf{F}_{\mathbf{h}, \, \mathrm{sharp}} = $.
More often nowadays normalized structure factors | E2 | − 1 are used in place of | F2 | . Normalized structure factors are used in direct methods techniques of structure solution. They are defined as
$|E_\mathbf{h}|^{2} = |F_\mathbf{h}|^{2}/\langle |F_\mathbf{h}|^{2}\rangle$,
where the squared observed structure-factor magnitudes on an absolute scale are divided by their expected values. Their use gives much greater weight to higher-resolution data and resolves some peaks in the vector map that would otherwise be continuous. On the other hand, they are less accurately known and are adjacent to data that have not been measured; they may therefore introduce spurious definition into the map.
A compromise that is often helpful is to use | EF | as the Patterson coefficients.
Structure-factor coefficient
The quantity $F$, $F^2$ or $I$ that is used in place of $Y$ in the function $\sum w(Y_o-Y_c)^2$ minimized during least-squares refinement of a crystal structure determination.
Structure amplitude
The magnitude of the structure factor.
Structure determination
Structure determination in crystallography refers to the process of elaborating the three-dimensional positional coordinates (and also, usually, the three-dimensional anisotropic displacement parameters) of the scattering centers in an ordered crystal lattice. Where a crystal is composed of a molecular compound, the term generally includes the three-dimensional description of the chemical structures of each molecular compound present.
Experimental techniques
Owing to the highly ordered arrangement of atoms as scattering centers in a crystal lattice, most structure determination techniques involve the diffraction of electromagnetic or matter waves of wavelengths comparable to atomic dimensions. Bragg's law specifies the condition for plane waves to be diffracted from lattice planes. The diffracted radiation passing through a crystal emerges with intensity varying as a function of scattering angle. This variation arises from constructive and destructive interference of scattered beams from the planes associated with the different atoms present in the lattice. The result is seen by an imaging detector as a pattern of diffraction spots or rings.
Among diffraction-based techniques are:
• single-crystal X-ray diffraction
• X-ray powder diffraction
• X-ray fiber diffraction
• neutron powder diffraction
• neutron single-crystal diffraction
• polymer electron diffraction
Other techniques for three-dimensional structure determination that are complementary to diffraction methods include
• electron microscopy
• nuclear magnetic resonance spectroscopy (used largely for biological macromolecules in solution)
Methodology
The following summary applies to single-crystal X-ray diffraction. A crystal, mounted on a goniometer, is illuminated by a collimated monochromatic X-ray beam, and the positions and intensities of diffracted beams are measured. The measured intensities Ihkl (corresponding to scattering from a lattice plane with Miller indices h,k,l) are reduced to structure amplitudes Fhkl by the application of a number of experimental corrections:
$F^2_{hkl} = I_{hkl}(k \mathrm{Lp} A)^{-1}$
where k is a scale factor, Lp the Lorentz–polarization correction, and A the transmission factor representing the absorption of X-rays by the crystal. The structure amplitude represents the amplitude of the diffracted wave measured relative to the scattering amplitude of a single electron.
However, the diffracted wave is completely described by the structure factor $\mathbf{F}_{hkl}$:
$\mathbf{F}_{hkl} = F_{hkl}\exp(i\alpha_{hkl}) = \sum_j f_j\exp[2\pi i (hx_j + ky_j + lz_j)]$
$\qquad = \sum_j f_j\cos[2\pi (hx_j + ky_j + lz_j)] + i\sum_{j} f_j\sin[2\pi (hx_j + ky_j + lz_j)]$
$\qquad = A_{hkl} + iB_{hkl}$
where the sum is over all atoms in the unit cell, xj,yj,zj are the positional coordinates of the jth atom, fj is the scattering factor of the jth atom, and αhkl is the phase of the diffracted beam.
The atomic scattering factor can be worked out from the physical properties of the atom species, but the phase cannot be determined by direct experimental observation. If the phases can be derived in some way, then the positional coordinates can be calculated from the expression above. The phase problem represents the major obstacle to constructing an initial structural model, and is addressed through a number of techniques, such as direct methods, Patterson synthesis, heavy-atom method, isomorphous replacement etc.
Once an initial structural model has been calculated, it is usually necessary to conduct an iterative refinement procedure to improve the agreement between the structural model and the experimental diffraction intensities. The most common approach is to perform a least-squares minimization between the experimental structure factors and those calculated by varying the adjustable parameters of the structural model. These normally include atomic positions, anisotropic displacement parameters, occupancies, chemical bond lengths and angles, and other geometric characteristics of a molecule. Some metric, such as the residual R factor
$R = $,
is used to indicate improvements or reductions in the quality of fit between model and observation. In the expression above (the 'conventional' R factor), Fobs and Fcalc are the observed and calculated structure amplitudes, and the deviations are summed over all experimentally recorded intensities. There is considerable discussion on the most appropriate statistical metric to use for this purpose.
Other refinement techniques, such as maximum likelihood and maximum entropy, are also used.
Superposition methods
A subset of Patterson methods of structure solution that involve analyzing the Patterson map, transforming the origin of the map in turn on to the known positions of certain atoms, and suitably combining the superposed maps. The degree of coincidence between the peaks of the superposing maps is assessed by optimizing one of a number of image-seeking functions, of which the symmetry minimum function is perhaps most often used. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/Structure_Determination/Sharpened_Patterson_function.txt |
An absorption edge is a sharp discontinuity in the absorption spectrum of X-rays by an element that occurs when the energy of the photon corresponds to the energy of a shell of the atom (K, LI, LII, LIII, etc.).
Anomalous absorption
Anomalous absorption takes place when radiation is dynamically diffracted by a perfect or nearly perfect crystal. The optical field in the crystal is then made up by several components, called wavefields, two in the two beam case (neglecting polarization in the X-ray case). One of them is absorbed more than normal and the other one less than normal. In the transmission, or Laue geometry, both wavefields propagate inside the crystal; one then speaks of anomalous transmission for the less absorbed wavefield (Borrmann effect). In the reflection, or Bragg geometry, one wavefield only propagates in the crystal, the more absorbed one for angles of incidence corresponding to one side of the total reflection rocking curve and the less absorbed one for the other side. This results in an asymmetry of the rocking curve that is calculated using dynamical theory.
Anomalous dispersion
The 'anomalous' dispersion corrections, which are not in fact anomalous, take into account the effect of absorption in the scattering of phonons by electrons. In the classic picture the electron is approximated by a damped harmonic oscillator. The scattering factor of the electron is then complex and the atomic scattering factor, or atomic form factor, is given by:
f + f' + i f"
where f' and f" are the real and imaginary parts of the anomalous dispersion correction. Their importance increases as one gets closer to an absorption edge (resonant scattering). Numerical calculations usually follow the Hartree-Fock approximations. For details on the non-relativistic and relativistic approaches, see Section 4.2.6 of International Tables of Crystallography, Volume C.
Anomalous scattering
The history of the description of the scattering of an atom when illuminated with X-rays is that initially wavelength dependencies were ignored. This was initially referred to as 'normal scattering'. The wavelength dependencies were then corrections to the normal scattering and also called anomalous. These had to describe changes in amplitude and phase, respectively initially given the symbols \(\Delta f\,'\) and \(\Delta f\,''\). Thus the X-ray scattering factor of an atom is described by the equation:-
\(f=f_o + \Delta f\,' + i\Delta f\,''\)
The nomenclature changed when tunable synchrotron sources became available and whereby the Δ prefixes were removed because changes between two wavlengths would then have required a double Δ label, which is cumbersome. Thus the \(\Delta f\,'\) now means the change in \(f\,'\)between two wavelengths. The Δ prefix to \(f\,''\) is dropped for consistency even though its use is based on its value at a single wavelength.
The values of \(f\,'\) and \(f\,''\) change most at the absorption edge of the element in question. Thus this resonance effect sometimes leads to the term being refererred to as 'resonant scattering'. However, since the off resonance \(f\,''\) effect is extensively used in crystal structure determination of the hand of a molecule (its chirality) 'anomalous scattering' is the best i.e. most widely embracing term. Another commonly used term is Multiple-wavelength Anomalous Dispersion ('MAD'), which involves measurements made at the resonance condition and at more than one wavelength obviously.
Borrmann Effect
Due to anomalous absorption, type 1 wavefields propagate in a perfect or nearly perfect crystal with a less than normal absorption. For details and the physical interpretation, see anomalous absorption.
Super-Borrmann effect
It is the enhancement of the Borrmann effect in a three-beam case, e.g. when the $111$ and $\overline{111}$ reflections are simultaneously excited in a silicon or germanium crystal.
History
The Borrmann effect was first discovered in quartz (Borrmann G., 1941, Über Extinktionsdiagramme der Röntgenstrahlen von Quarz. Physik Z., 42, 157-162) and then in calcite crystals (Borrmann G., 1950, Die Absorption von Röntgenstrahlen in Fall der Interferenz. Z. Phys., 127, 297-323), and interpreted by Laue (Laue, M. von, 1949, Die Absorption der Röntgenstrahlen in Kristallen im Interferenzfall. Acta Crystallogr. 2, 106-113).
The super-Borrmann effect was first observed by Borrmann G. and Hartwig W. (1965), Die Absorption der Röntgenstrahlen im Dreistrahlfall der Interferenz. Z. Krist., 121, 401-409.
See also
• Section 5.1 of International Tables of Crystallography, Volume B for X-rays
• Section 5.2 of International Tables of Crystallography, Volume B for electrons
• Section 5.3 of International Tables of Crystallography, Volume B for neutrons | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/X-rays/Absorption_edge.txt |
Bragg's law provides the condition for a plane wave to be diffracted by a family of lattice planes:
$2 d \sin θ = n λ. \nonumber$
where d is the lattice spacing, θ the angle between the wavevector of the incident plane wave, ko, and the lattice planes, λ its wave length and n is an integer, the order of the reflection. It is equivalent to the diffraction condition in reciprocal space and to the Laue equations.
Direct derivation of Bragg's law
• Reflection from the first plane
The scattered waves will be in phase whatever the distribution of the point scatterers in the first plane if the angle of the reflected wave vector, kh, is also equal to θ. This is Snell-Descartes' law of reflection.
• Reflection from the second plane
Since the phase of the reflected waves is independent of the position of the point scatterer in the plane, the phase difference between the waves reflected by two successive lattice planes is obtained by choosing arbitrarily a scattering point, A, on the first plane and a scattering point,b on the second plane such that AB is normal to the planes. If C and d are the projections of A on the incident and reflected wave vectors passing through B, it is clear from figure 1 that the path difference between the waves reflected at A and B, respectively, is:
$CB + BD = 2 d \sin θ \nonumber$
and that the two waves will be in phase if this path difference is equal to n λ where n is an integer.
• Reflection from the third, etc., planes
If Bragg's relation is satisfied for the first two planes, the waves reflected with wave vector kh will be in phase fo all the planes of the family.
Order of the reflection
Bragg's law may also be written:
$2 \left(\dfrac{d}{n}\right) \sin θ = λ. \nonumber$
One may then say that a Bragg reflection of order n on a family of lattice planes or order n is equivalent to reflection of order 1 on a family of fictitious, or imaginary, planes of lattice spacing:
$d_{hkl} = \dfrac{d}{n} \nonumber$
This fictitious family is associated to the reciprocal lattice vector OH where OH = n/d = 1/dhkl. The indices of the reflection are: hkl. For instance, the dashed blue lines in Figure 1 correspond to the fistitious planes associated to the second order, n = 2.
Extinctions, or systematic absences
If there is a glide plane or a screw axis normal to the lattice planes, the spacing of the actual reflecting planes isd/2 for a glide plane and (d p/q) for a qp screw axis. Bragg's law should then be written:
2 (d/2) sin θ = n λ ⇒ 2 d sin θ = 2n λ
for a glide plane and
2 (d p/q) sin θ = n λ ⇒ 2 d sin θ = (q/p)n λ
for a screw axis qp.
The reflections of odd order for a glide plane and of order different from (q/p)n for a screw axis are then absent. One speaks of extinctions or systematic absences related to the presence of glide or screw components.
As an example, the case of a 21 screw axis. Reflections of odd order will be systematically absent.
Influence of deformation
A deformation that leaves a family of lattice planes (hkl) undistorted and its lattice spacing d unchanged will not affect the Bragg angle of kklreflections, e.g. lattice planes parallel to a screw dislocations.
History
A deformation that leaves a family of lattice planes (hkl) undistorted and its lattice spacing d unchanged will not affect the Bragg angle of kkl reflections, e.g. lattice planes parallel to a screw dislocations.
Bragg angle
In Bragg's law describing the condition for a plane wave to be diffracted from a family of lattice planes, the angle θ between the wavevector of the incident plane wave, ko, and the lattice planes.
CromerMann coefficients
The set of nine coefficients $a_i, b_i, c\, (i=1,\dots, 4)$ in a parameterization of the non-dispersive part of the atomic scattering factor for neutral atoms as a function of (sinθ) / λ:
$f^0(\sin\theta/\lambda) = \sum_{i=1}^4 a_i \exp[-b_i(\sin\theta/\lambda)^2] + c$
for $0 < (\sin\theta)/\lambda < 2.0\,\mathrm{\AA}^{-1}$.
This expression is convenient for calculation in crystal structure software suites.
Dynamical diffraction
When a crystal is perfect or nearly perfect, the usual geometrical, or kinematical theory of diffraction is an insufficient approximation and the dynamical theory of diffraction must be used to describe the diffracted intensities. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/X-rays/Bragg%27s_law.txt |
In the geometrical, or kinematical theory, the amplitudes diffracted by a three-dimensional periodic assembly of atoms (Laue) or by a stack of planes (Darwin) is derived by adding the amplitudes of the waves diffracted by each atom or by each plane, simply taking into account the optical path differences between them, but neglecting the interaction of the propagating waves and matter. This approximation is not compatible with the law of conservation of energy and is only valid for very small or highly imperfect crystals. The purpose of the dynamical theory is to take this interaction into account. There are three forms of the dynamical theory:
Electron density map
A three-dimensional description of the electron density in a crystal structure, determined from X-ray diffraction experiments. X-rays scatter from the electron clouds of atoms in the crystal lattice; the diffracted waves from scattering planes h,k,l are described by structure factors $\mathbf{F}_{hkl}.$ The electron density as a function of position x,y,z is the Fourier transform of the structure factors:
$\rho(xyz) = {1\over V}\sum_{hkl} F(hkl)\exp[-2\pi i(hx + ky + lz)]$.
The electron density map describes the contents of the unit cells averaged over the whole crystal and not the contents of a single unit cell (a distinction that is important where structural disorder is present).
Three-dimensional maps are often evaluated as parallel two-dimensional contoured sections at different heights in the unit cell.
Units
Electron density is measured in electrons per cubic ångström, e Å-3.
Ewald sphere
The Ewald sphere, or sphere of reflection, is a sphere of radius 1/λ passing through the origin O of the reciprocal lattice. The incident direction is along a radius of the sphere, IO (Figure 1). A reflected direction, of unit vector sh, will satisfy the diffraction condition if the diffraction vector OH = IHIO = sh/λ – so/λ (so unit vector in the direction IO) is a reciprocal lattice vector, namely if H is a node of the reciprocal lattice (see Diffraction condition in reciprocal space) . If other reciprocal lattice nodes, such as G, lie also on the sphere, there will be reflected beams along IG, etc. This construction is known as the Ewald construction. When the wavelength is large, there are seldom more than two nodes, O and H, of the reciprocal lattice simultaneously on the Ewald sphere. When there are three or more, one speaks of multiple diffraction, multiple scattering or n-beam diffraction. This situation becomes increasingly frequent as the wavelength decreases and is practically routine for very short wavelengths such as those of γ-rays and electrons. The curvature of Ewald sphere then becomes negligible and it can often be approximated by a plane. Many reflections must then be taken into account at the same time.
When the wavelength changes, the radius of the Ewald sphere changes. If the incident beam is a white beam, with a wavelength range λmin ≤ λ ≤ λmax, there will be a nest of Ewald spheres of radii 1/λmax≤ 1/λ ≤ 1/λmin.
F(000)
The expression for a structure factor evaluated in the zeroth-order case h = k = l = 0 yields the result
$F(000) = [ (\sum f_{r} )^{\,2} + (\sum f_{i} )^{\,2} ]^{1/2}$
where fr is the real part of the scattering factors at $\theta = 0^\circ$, fi is the imaginary part of the scattering factors at $\theta = 0^\circ$, θ is the Bragg angle, and the sum is taken over each atom in the unit cell.
F(000) is computed without dispersion effects in electron-density calculation by Fourier inversion. In all cases, non-dispersive F(000) is a structure factor and not a structure amplitude: it has both magnitude and a sign.
For X-rays non-dispersive F(000) is positive definite and in many cases an integer (but it is not an integer for non-stoichiometric compounds). It counts the number of electrons in the cell.
For neutrons non-dispersive F(000) is either positive or negative and counts the total nuclear scattering power in the cell. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/X-rays/Dynamical_theory_of_Scattering.txt |
Friedel's law states that the intensities of the $h$, $k$, $l$ and $\overline{h}$, $\overline{k}$, $\overline{l}$ reflections are equal. This is true either if the crystal is centrosymmetric or if no resonant scattering is present. It is in that case not possible to tell by diffraction whether an inversion center is present or not. The apparent symmetry of the crystal is then one of the eleven Laue classes.
The reason for Friedel's rule is that the diffracted intensity is proportional to the square of the modulus of the structure factor, |Fh|2, according to the geometrical, or kinematical theory of diffraction. It depends similarly on the modulus of the structure factor according to the dynamical theory of diffraction. The structure factor is given by:
$F_h = \sum_j f_j (exp - 2 \pi i) \textbf {h} \cdot \textbf{r_j} \nonumber$
where $f_j$ is the atomic scattering factor of atom j, $\textbf{h}$ the reflection vector and $\textbf{r_j}$ the position vector of atom $j$. There comes:
$|F_h|^2 = F_h F_h^* = F_h F_{\bar h} = |F_{\bar h}|^2 \nonumber$
if the atomic scattering factor, fj, is real. The intensities of the $h$, $k$, $l$ and $\overline{h}$, $\overline{k}$, $\overline{l}$ reflections are then equal. If the crystal is absorbing, however, due to resonant scattering, the atomic scattering factor is complex and
$F_{\bar h} \ne F_h^* \nonumber$
The reflections h, k, l and ${\bar h}$, ${\bar k}$, ${\bar l}$ are called a Friedel pair. They are used in the resolution of the phase problem for the solution of crystal structures and in the determination of absolute structure.
History
Friedel's law was stated by G. Friedel (1865-1933) in 1913 (Friedel G., 1913, Sur les symétries cristallines que peut révéler la diffraction des rayons X., C.R. Acad. Sci. Paris, 157, 1533-1536.
See also
Section 3.1 of International Tables of crystallography, Volume A
Friedel pair
The couple of reflections h, k, l and ${\bar h}, {\bar k}, {\bar l}$ is called a Friedel pair, or Bijvoet pair. Their intensities are equal either if the crystal structure is centrosymmetric or if there is no resonant scattering, but differ otherwise. Friedel's law then does not hold. For crystals with a non-centrosymmetric structure and significant resonant scattering, equivalent reflections generated by the symmetry operations of the point group of the crystal have intensities different from those of equivalent reflections generated by the introduction of an additional inversion centre in normal scattering. Friedel, or Bijvoet pairs are used in the resolution of the phase problem for the solution of crystal structures and in the determination of absolute structure.
Integral reflection conditions
The integral reflections are the general reflection conditions due to the centering of cells. They are given in the table below:
Integral reflection conditions for centered lattices.
Reflection
condition
Centering type of cell Centering symbol
None Primitive P
R (rhombohedral axes)
h + k = 2n C-face centered C
k + l = 2n A-face centered A
l + h = 2n B-face centered B
h + k + l = 2n body centered I
h + k, h + l and
k + l = 2n or:
h, k, l all odd or all
even (‘unmixed’)
all-face centered F
h + k + l = 3n
rhombohedrally
centered, reverse
setting
R (hexagonal axes)
hk + l = 3n
rhombohedrally
centered, obverse
setting (standard)
hk = 3n hexagonally centered H
Kinematical theory
In the kinematical or geometrical theory, the amplitudes diffracted by a three-dimensional periodic assembly of atoms (Laue) or by a stack of planes (Darwin) is derived by adding the amplitudes of the waves diffracted by each atom or by each plane, simply taking into account the optical path differences between them, but neglecting the interaction of the propagating waves and matter. This approximation is not compatible with the law of conservation of energy and is only valid for very small or highly imperfect crystals. The purpose of the dynamical theory is to take this interaction into account .
Laue equations
The three Laue equations give the conditions to be satisfied by an incident wave to be diffracted by a crystal. Consider the three basis vectors, OA = a, OB = b , OC = c of the crystal and let $\vec{s}_o$ and $\vec{s}_h$ be unit vectors along the incident and reflected directions, respectively. The conditions that the waves scattered by O and A, O and B,O and C, respectively, be in phase are that
$\vec{a} \cdot (\ce{s}_h - \vec{s}_o) = h λ \nonumber$
$\vec{b} \cdot (\vec{s}_h - \vec{s}_o) = k λ \nonumber$
$\vec{c} \cdot (\vec{s}_h - \vec{s}_o) = l λ \nonumber$
If these three conditions are simultaneously satisfied, the incoming wave is reflected on the set of lattice planes of Miller indices h/n, k/n, l/n. h, k, l are the indices of the reflection.
The three Laue equations can be generalized by saying that the diffraction condition is satisfied if the scalar product
$\vec{r} \cdot (\vec{s}_h/λ - \vec{s}_o/λ) \nonumber$
is an integer for any vector
$\vec{r} = u \vec{a} + v \vec{b} + w \vec{c} \nonumber$
where (u, v, w integers) of the direct lattice. This is the case if
(sh/λ - so/λ) = h a* + k b* + l c*,
where h, k, l are integers, namely if the diffraction vector OH = sh,/λ - so/λ is a vector of the reciprocal lattice. This is the diffraction condition in reciprocal space.
History
The three Laue conditions for diffraction were first given in Laue, M. (1912). Eine quantitative Prüfung der Theorie für die Interferenz-Erscheinungen bei Röntgenstrahlen. Sitzungsberichte der Kgl. Bayer. Akad. der Wiss 363--373, reprinted in Ann. Phys. (1913), 41, 989-1002 where he interpreted and indexed the first diffraction diagram (Friedrich, W., Knipping, P., and Laue, M. (1912). Interferenz-Erscheinungen bei Röntgenstrahlen, Sitzungsberichte der Kgl. Bayer. Akad. der Wiss, 303--322, reprinted in Ann. Phys., (1913), 41, 971-988, taken with zinc-blende, ZnS. For details, see P. P. Ewald, 1962, IUCr, 50 Years of X-ray Diffraction, Section 4, page 52.
Lorentzpolarization correction
A multiplicative factor involved in converting diffracted radiation intensities to structure factors during the process of structure determination for X-ray diffraction experiments involving moving crystals.
Mosaic crystal
The mosaic crystal is a simplified model of real crystals proposed by C. G. Darwin. In this model, a real crystal is described as a mosaic of crystalline blocks with dimensions of 10-5 cm, tilted to each other by fractions of a minute of arc. Each block is separated from the surrounding blocks by faults and cracks.
In a diffraction experiment, interference between waves only occurs inside a block, whose dimensions satisfy the theoretical conditions of applicability of the kinematical theory. Because of the loss of coherence between the waves diffracted from different blocks, the diffracted intensity from the whole crystal is equal to the sum of the intensities diffracted by every block. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/X-rays/Friedel%27s_Law.txt |
The primary extinction is responsible for the loss of intensity due to dynamic effect inside every block of a mosaic crystal. At the Bragg angle, each incident wave can undergo multiple reflections from different atomic planes; each scattering introduced causes a phase difference of λ/4 = π/2 so that along each direction waves differing by an even number of scattering, i.e. by nπ in phase, interfere , whose intensity decreasing rapidly with the number of scatterings. Because of the decrease in intensity of waves multiply scattered, the effect of waves differing by more than two scatterings can normally be neglected: in each direction one observes then the destructive interference between waves having a significant difference in intensity and an overall reduction of the intensity with respect to the intensity given by the kinematical theory.
Reflection conditions
The reflection conditions describe the conditions of occurrence of a reflection (structure factor not systematically zero). There are two types of systematic reflection conditions for diffraction of crystals by radiation:
(1) General conditions. They apply to all Wyckoff positions of a space group, i.e. they are always obeyed, irrespective of which Wyckoff positions are occupied by atoms in a particular crystal structure. They are due to one of three effects:
• Centered cells.
The resulting conditions apply to the whole three-dimensional set of reflections hkl. Accordingly, they are called integral reflection conditions. They are given in Table 1.
• Glide planes.
The resulting conditions apply only to two-dimensional sets of reflections, i.e. to reciprocal-lattice nets containing the origin (such as hk0, h0l, 0kl, hhl). For this reason, they are called zonal reflection conditions. For instance, for a glide plane parallel to (001):
type of reflection reflection condition glide vector glide plane
0kl k = 2 n b/2 b
l = 2 n c/2 c
k + l = 2 n b/2 + c/2 n
k + l = 4 n
k, l = 2n
b/4 ± c/4 d
The zonal reflection conditions are listed in Table 2.2.13.2 of International Tables of Crystallography, Volume A.
• Screw axes.
The resulting conditions apply only to one-dimensional sets of reflections, i.e. reciprocal-lattice rows containing the origin (such as h00, 0k0, 00l). They are called serial reflection conditions. For instance, for a screw axis parallel to [001], the reflection conditions are:
type of reflection reflection condition screw vector screw axis
00l l = 2 n c/2 21; 42
l = 4 n c/4 41; 43
000l l = 2 n c/2 63
l = 3 n c/3 41; 31; 32; 62; 64
l = 6 n c/6 61;65
The serial reflection conditions are listed in Table 2.2.13.2 of International Tables of Crystallography, Volume A.
(2) Special conditions (‘extra’ conditions). They apply only to special Wyckoff positions and occur always in addition to the general conditions of the space group.
Resolution
In crystal structure determination, the term resolution is used to describe the ability to distinguish between neighboring features in an electron density map. By convention, it is defined as the minimum plane spacing given by Bragg's law for a particular set of X-ray diffraction intensities. The resolution improves with an increase in the maximum value of $(\sin \theta)/\lambda$ at which reflections are measured.
Resonant scattering
The elementary theory of the scattering of X-rays by atoms, leading to the real atomic scattering factor fo, applies only for X-radiation whose wavelength is far removed from that of any natural (resonant) frequency of the atom. When this condition does not hold, one needs to use as physical model for the scattering that of a forced damped harmonic oscillator. This leads to resonant-scattering terms in the full, now complex, scattering factor of an atom, represented by:-
\(f=f_o + f\,' + if\,''\) .
The real and imaginary terms, \(f\,', f\,''\), in the atomic scattering factor are independent of sin(θ)/λ and in general small compared to fo. The values of \(f\,'\) and \(f\,''\) change most at the absorption edge of the element in question.
In the older literature the term anomalous dispersion was used for resonant scattering. In macromolecular crystallography the term anomalous scattering is used widely in place of resonant scattering.
History
The resonant scattering of X-rays was theoretically predicted by Waller (Waller, I., 1928, Über eine verallgemeinerte Streuungsformel. Z. Phys. 51, 213-231.) and first calculated by Hönl (Hönl, H., 1933, Zur Dispersionstheorie der Röntgenstrahlen. Z. Phys. 84, 1-16; Hönl, H., 1933, Atomfactor für Röntgenstrahlen als Problem der Dispersionstheorie (K-Schale). Ann. Phys. (Leipzig), 18, 625-657.
Secondary extinction
The secondary extinction is responsible for the loss of intensity occurring when the incident beam crosses a crystal. Each plane of a family (hkl) satisfying Laue equations (or Bragg's law) diffracts the incident beam, and thus subtracts part of the intensity to the incident beam. Successive planes of the same family will then experience a weakening of the incident beam and as a consequence the diffracted beams will result from the positive interference of waves not having the same intensities, as it is instead considered by the kinematical theory.
Secondary extinction is equivalent to an increase of the linear absorption coefficient and is negligible for sufficiently small crystals.
Serial reflection conditions
The serial reflection conditions are the general reflection conditions due to the presence of screw axes. The resulting conditions apply only to one-dimensional sets of reflections, i.e. reciprocal-lattice rows containing the origin (such as h00, 0k0, 00l). For instance, for a screw axis parallel to [001], the reflection conditions are:
type of reflection reflection condition screw vector screw axis
00l l = 2 n c/2 21; 42
l = 4 n c/4 41; 43
000l l = 2 n c/2 63
l = 3 n c/3 41; 31; 32; 62; 64
l = 6 n c/6 61;65
The serial reflection conditions are listed in Table 2.2.13.2 of International Tables of Crystallography, Volume A. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/X-rays/Primary_extinction.txt |
The structure factor $\mathbf{F}_{hkl}$is a mathematical function describing the amplitude and phase of a wave diffracted from crystal lattice planes characterized by Miller indices h,k,l.
The structure factor may be expressed as
\begin{align} \mathbf{F}_{hkl} &= F_{hkl}\exp(i\alpha_{hkl}) = \sum_j f_j\exp[2\pi i (hx_j + ky_j + lz_j) \[4pt] &= \sum_j f_j\cos[2\pi (hx_j + ky_j + lz_j)] + i\sum_{j} f_j\sin[2\pi (hx_j + ky_j + lz_j)] \[4pt] &= A_{hkl} + iB_{hkl} \end{align} \nonumber
where the sum is over all atoms in the unit cell, xj,yj,zj are the positional coordinates of the jth atom, fj is the scattering factor of the jth atom, and $α_{hkl}$ is the phase of the diffracted beam.
The intensity of a diffracted beam is directly related to the amplitude of the structure factor, but the phase must normally be deduced by indirect means. In structure determination, phases are estimated and an initial description of the positions and anisotropic displacements of the scattering atoms is deduced. From this initial model, structure factors are calculated and compared with those experimentally observed. Iterative refinement procedures attempt to minimize the difference between calculation and experiment, until a satisfactory fit has been obtained.
Derivation
Consider Bragg's law for an array of atom scatterers in a primitive lattice with just one atom at each lattice point. An incident X-ray wave of wavelength λ diffracts strongly through an angle $2θ$ where the perpendicular distance between two lattice planes $d_{hkl}$ satisfies the relation
$2d_{hkl}\sin{θ} = nλ \nonumber$
It is seen that the path difference between waves diffracted from the two planes shown differs by just one wave cycle.
Now consider a second atom added to the unit cell. Each original atom is now accompanied by a companion atom of the new type, offset by a displacement vector r1. The incident X-ray beam will also diffract from these new scatterers (since they occupy planes parallel to those originally drawn). But now there is a phase difference φ1 between the waves scattered from the first and the new sets of atoms.
The amplitudes of the waves are proportional to the atomic scattering factors f0 and f1. The phases differ by the angle φ1 . The resultant vector represents the two-atom structure factor with amplitude Fhkl. Note that there is a net phase φ arising from the phase difference due to the offset in position between the two sets of diffracting atoms.
As in the case of two atoms, the resultant diffracted wave is obtained from the linear superposition of the wave vectors scattered from each different atom.
Units
The units of the structure-factor amplitude depend on the incident radiation. For X-ray crystallography they are multiples of the unit of scattering by a single electron ($2.82 \times 10^{-15}\;m$); for neutron scattering by atomic nuclei the unit of scattering length of $10^{-14}\; m$ is commonly used.
See also
The structure factor. P. Coppens. International Tables for Crystallography (2006). Vol. B, ch. 1.2, pp. 10-24
Systematic absences
One speaks of systematic absences or extinctions when the structure factor is zero, due either to the centering of the lattice or to the presence of glide or screw symmetry elements. Conversely, the conditions for a reflection to exist and not to be systematically absent, are called reflection conditions.
Zonal reflection conditions
The zonal reflection conditions are the general reflection conditions due to the presence of glide planes. The resulting conditions apply only to two dimensional sets of reflections, i.e. to reciprocal-lattice nets containing the origin (such as hk0, h0l, 0kl, hhl). For instance, for a glide plane parallel to (001):
type of reflection reflection condition glide vector glide plane
0kl k = 2 n b/2 b
l = 2 n c/2 c
k + l = 2 n b/2 + c/2 n
k + l = 4 n
k, l = 2n
b/4 ± c/4 d
The zonal reflection conditions are listed in Table 2.2.13.2 of International Tables of Crystallography, Volume A. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Crystallography/X-rays/Structure_Factor.txt |
Alkali metals are the chemical elements found in Group 1 of the periodic table. The alkali metals include: lithium, sodium, potassium, rubidium, cesium, and francium. Although often listed in Group 1 due to its electronic configuration, hydrogen is not technically an alkali metal since it rarely exhibits similar behavior. The word "alkali" received its name from the Arabic word "al qali," meaning "from ashes", which since these elements react with water to form hydroxide ions, creating alkaline solutions (pH>7).
• Group 1: Properties of Alkali Metals
This page discusses the trends in some atomic and physical properties of the Group 1 elements - lithium, sodium, potassium, rubidium and cesium. Sections below cover the trends in atomic radius, first ionization energy, electronegativity, melting and boiling points, and density.
• Group 1: Reactivity of Alkali Metals
Alkali metals are among the most reactive metals. This is due in part to their larger atomic radii and low ionization energies. They tend to donate their electrons in reactions and have an oxidation state of +1. These metals are characterized by their soft texture and silvery color. They also have low boiling and melting points and are less dense than most elements. All these characteristics can be attributed to these elements' large atomic radii and weak metallic bonding.
• Chemistry of Hydrogen (Z=1)
Hydrogen is one of the most important elements in the world. It is all around us. It is a component of water (H2O), fats, petroleum, table sugar (C6H12O6), ammonia (NH3), and hydrogen peroxide (H2O2)—things essential to life, as we know it. This module will explore several aspects of the element and how they apply to the world.
• Chemistry of Lithium (Z=3)
Chlorine is a halogen in Lithium is a rare element found primarily in molten rock and saltwater in very small amounts. It is understood to be non-vital in human biological processes, although it is used in many drug treatments due to its positive effects on the human brain. Because of its reactive properties, humans have utilized lithium in batteries, nuclear fusion reactions, and thermonuclear weapons.
• Chemistry of Sodium (Z=11)
Sodium is metallic element found in the first group of the periodic table. As the sixth most abundant element in the Earth's crust, sodium compounds are commonly found dissolved in the oceans, in minerals, and even in our bodies.
• Chemistry of Potassium (Z=19)
In its pure form, potassium has a white-sliver color but it quickly oxidizes upon exposure to air, tarnishing in minutes if it is not stored under oil or grease. Potassium is essential to several aspects of plant, animal, and human life and is thus mined, manufactured, and consumed in huge quantities around the world.
• Chemistry of Rubidium (Z=37)
Rubidium (Latin: rubidius = red) is similar in physical and chemical characteristics to potassium, but much more reactive. It is the seventeenth most abundant element and was discovered by its red spectral emission in 1861 by Bunsen and Kirchhoff. Its melting point is so low you could melt it in your hand if you had a fever (39°C). But that would not be a good idea because it would react violently with the moisture in your skin.
• Chemistry of Cesium (Z=55)
Cesium is so reactive that it will even explode on contact with ice! It has been used as a "getter" in the manufacture of vacuum tubes (i.e., it helps remove trace quantities of remaining gases). An isotope of cesium is used in the atomic clocks.
• Chemistry of Francium (Z=87)
Francium is the last of the known alkali metals and does not occur to any significant extent in nature. All known isotopes are radioactive and have short half-lives (22 minutes is the longest).
Group 1: The Alkali Metals
This page discusses the trends in some atomic and physical properties of the Group 1 elements - lithium, sodium, potassium, rubidium and cesium. Sections below cover the trends in atomic radius, first ionization energy, electronegativity, melting and boiling points, and density.
Trends in Atomic Radius
The chart below shows the increase in atomic radius down the group.
The radius of an atom is governed by two factors:
1. The number of layers of electrons around the nucleus
2. The attraction the outer electrons feel from the nucleus
Compare the electronic configurations of lithium and sodium:
• Li: 1s22s1
• Na: 1s22s22p63s1
In each element, the outer electron experiences a net charge of +1 from the nucleus. The positive charge on the nucleus is canceled out by the negative charges of the inner electrons. This effect is illustrated in the figure below:
This is true for each of the other atoms in Group 1. The only factor affecting the size of the atom is the number of layers of inner electrons which surround the atom. More layers of electrons take up more space, due to electron-electron repulsion. Therefore, the atoms increase in size down the group.
The first ionization energy of an atom is defined as the energy required to remove the most loosely held electron from each of one mole of gaseous atoms, producing one mole of singly charged gaseous ions; in other words, it is the energy required for 1 mole of this process:
$X(g) \rightarrow X^+ (g) + e^- \nonumber$
A graph showing the first ionization energies of the Group 1 atoms is shown above. Notice that first ionization energy decreases down the group. Ionization energy is governed by three factors:
• the charge on the nucleus,
• the amount of screening by the inner electrons,
• the distance between the outer electrons and the nucleus.
Down the group, the increase in nuclear charge is exactly offset by the increase in the number of inner electrons. As mentioned before, in each of the elements Group 1, the outermost electrons experience a net charge of +1 from the center. However, the distance between the nucleus and the outer electrons increases down the group; electrons become easier to remove, and the ionization energy falls.
Trends in Electronegativity
Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. It is usually measured on the Pauling scale, on which the most electronegative element (fluorine) is given an electronegativity of 4.0 (Table A2).
A graph showing the electronegativities of the Group 1 elements is shown above. Each of these elements has a very low electronegativity when compared with fluorine, and the electronegativities decrease from lithium to cesium.
Picture a bond between a sodium atom and a chlorine atom. The bond can be considered covalent, composed of a pair of shared electrons. The electron pair will be pulled toward the chlorine atom because the chlorine nucleus contains many more protons than the sodium nucleus. This is illustrated in the figure below:
The electron pair is so close to the chlorine that an effective electron transfer from the sodium atom to the chlorine atom occurs—the atoms are ionized. This strong attraction from the chlorine nucleus explains why chlorine is much more electronegative than sodium.
Now compare this with a lithium-chlorine bond. The net pull from each end of the bond is the same as before, but the lithium atom is smaller than the sodium atom. That means that the electron pair is going to be more strongly attracted to the net +1 charge on the lithium end, and thus closer to it.
In some lithium compounds there is often a degree of covalent bonding that is not present in the rest of the group. Lithium iodide, for example, will dissolve in organic solvents; this is a typical property of covalent compounds. The iodine atom is so large that the pull from the iodine nucleus on the pair of electrons is relatively weak, and a fully-ionic bond is not formed.
Summarizing the trend down the group
As the metal atoms increase in size, any bonding electron pair becomes farther from the metal nucleus, and so is less strongly attracted towards it. This corresponds with a decrease in electronegativity down Group 1. With the exception of some lithium compounds, the Group 1 elements each form compounds that can be considered ionic. Each is so weakly electronegative that in a Group 1-halogen bond, we assume that the electron pair on a more electronegative atom is pulled so close to that atom that ions are formed.
Trends in melting and boiling points
The figure above shows melting and boiling points of the Group 1 elements. Both the melting and boiling points decrease down the group.
When any of the Group 1 metals is melted, the metallic bond is weakened enough for the atoms to move more freely, and is broken completely when the boiling point is reached. The decrease in melting and boiling points reflects the decrease in the strength of each metallic bond.
The atoms in a metal are held together by the attraction of the nuclei to electrons which are delocalized over the whole metal mass. As the atoms increase in size, the distance between the nuclei and these delocalized electrons increases; therefore, attractions fall. The atoms are more easily pulled apart to form a liquid, and then a gas. As previously discussed, each atom exhibits a net pull from the nuclei of +1. The increased charge on the nucleus down the group is offset by additional levels of screening electrons. As before, the trend is determined by the distance between the nucleus and the bonding electrons.
Trends in Density
The densities of the Group 1 elements increase down the group (except for a downward fluctuation at potassium). This trend is shown in the figure below:
The metals in this series are relatively light—lithium, sodium, and potassium are less dense than water (less than 1 g cm-3). It is difficult to develop a simple explanation for this trend because density depends on two factors, both of which change down the group. The atoms are packed in the same way, so the two factors considered are how many atoms can be packed in a given volume, and the mass of the individual atoms. The amount packed depends on the individual atoms' volumes; these volumes, in turn, depends on their atomic radius.
Atomic radius increases down a group, so the volume of the atoms also increases. Fewer sodium atoms than lithium atoms, therefore, can be packed into a given volume. However, as the atoms become larger, their masses increase. A given number of sodium atoms will weigh more than the same number of lithium atoms. Therefore, 1 cm3 of sodium contains fewer atoms than the same volume of lithium, but each atom weighs more. Mathematical calculations are required to determine the densities.
Contributors and Attributions
Jim Clark (Chemguide.co.uk) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__1%3A_The_Alkali_Metals/1Group_1%3A_Physical_Properties_of_Alkali_Metal.txt |
Alkali metals are among the most reactive metals. This is due in part to their larger atomic radii and low ionization energies. They tend to donate their electrons in reactions and have an oxidation state of +1. These metals are characterized by their soft texture and silvery color. They also have low boiling and melting points and are less dense than most elements. Lithium, sodium, and potassium float on water because of their low density. All these characteristics can be attributed to these elements' large atomic radii and weak metallic bonding.
2Reactions of the Group 1 Elements
This page describes how to perform a flame test for a range of metal ions, and briefly discusses how the flame color arises. Flame tests are used to identify the presence of a relatively small number of metal ions in a compound. Not all metal ions give flame colors. For Group 1 compounds, flame tests are usually by far the easiest way of identifying which metal you have got. For other metals, there are usually other easy methods that are more reliable - but the flame test can give a useful hint as to where to look.
Practical Details in Carrying out Flame Tests
• Clean a platinum or nichrome (a nickel-chromium alloy) wire by dipping it into concentrated hydrochloric acid and then holding it in a hot Bunsen flame. Repeat this until the wire produces no color in the flame.
• When the wire is clean, moisten it again in the acid and then dip it into a small amount of the solid to be tested so that some sticks to the wire. Place the wire back in the flame.
• If the flame color is weak, it is often helpful to dip the wire back in the acid and put it back into the flame as if cleaning it. This should produce a very short but intense flash of color.
There will, in fact, always be a trace of orange in the flame if you use nichrome. Platinum is much better to use but is much, much more expensive. If you have a particularly dirty bit of nichrome wire, you can just chop the end off. You do not do that with platinum! Dilute hydrochloric acid can be used instead of concentrated acid for safety reasons, but does not always give such intense flame colors.
The colors in Table \(1\) are just a guide. Almost everybody sees and describes colors differently. I have, for example, used the word "red" several times to describe colors that can be quite different from each other. Other people use words like "carmine" or "crimson" or "scarlet", but not everyone knows the differences between these words - particularly if their first language is not English.
Table \(1\): Flame colors of common elements
Element flame color
Lithium red
Sodium strong, persistent orange
Potassium lilac (pink)
Rubidium red (red-violet)
Cesium blue/violet (see below)
Calcium orange-red
Strontium red
Barium pale green
Copper blue-green (often with white flashes)
Lead gray-white
What do you do if you have a red flame color for an unknown compound and do not know which of the various reds it is? Get samples of known lithium, strontium (etc) compounds and repeat the flame test, comparing the colors produced by one of the known compounds and the unknown compound side by side until you have a good match.
The Origin of Flame Colors
If you excite an atom or an ion by very strong heating, electrons can be promoted from their normal unexcited state into higher orbitals. As they fall back down to lower levels (either in one go or in several steps), energy is released as light. Each of these jumps involves a specific amount of energy being released as light energy, and each corresponds to a particular wavelength (or frequency). As a result of all these jumps, a spectrum of lines will be produced, some of which will be in the visible part of the spectrum. The color you see will be a combination of all these individual colors.
In the case of sodium (or other metal) ions, the jumps involve very high energies and these result in lines in the UV part of the spectrum which your eyes can't see. The jumps that you can see in flame tests come from electrons falling from a higher to a lower level in the metal atoms. So if, for example, you put sodium chloride which contains sodium ions, into a flame, where do the atoms come from? In the hot flame, some of the sodium ions regain their electrons to form neutral sodium atoms again. A sodium atom in an unexcited state has the structure 1s22s22p63s1, but within the flame there will be all sorts of excited states of the electrons. Sodium's familiar bright orange-yellow flame color results from promoted electrons falling back from the 3p1 level to their normal 3s1 level.
The exact sizes of the possible jumps in energy terms vary from one metal to another. That means that each different metal will have a different pattern of spectral lines, and so a different flame color. Flame colors are produced from the movement of the electrons in the metal ions present in the compounds. For example, a sodium ion in an unexcited state has the electron configuration 1s22s22p6. When heated, the electrons gain energy and can be excited into any of the empty higher-energy orbitals—7s, 6p, 4d, or any other, depending on the amount of energy a particular electron happens to absorb from the flame. Because the electron is now at a higher and more energetically unstable level, it falls back down to the original level, but not necessarily in one transition.
The electron transitions which produced lines in the visible spectrum involved atoms rather than ions. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__1%3A_The_Alkali_Metals/2Reactions_of_the_Group_1_Elements/Flame_Tests.txt |
This page discusses a few compounds of the Group 1 elements (lithium, sodium, potassium, rubidium and cesium), including some information about the nitrates, carbonates, hydrogen carbonates and hydrides of the metals.
The effect of heat on Group 1 compounds
Group 1 compounds are more resistant to heat than the corresponding compounds in Group 2. Lithium compounds often behave similarly to Group 2 compounds, but the rest of Group 1 act differently in various ways.
Heating the nitrates
Most nitrates tend to decompose on heating to the metal oxide, brown fumes of nitrogen dioxide, and oxygen. For example, a typical Group 2 nitrate like magnesium nitrate decomposes this way:
$2Mg(NO_3)_2 (s) \rightarrow 2MgO(s) + 4NO_2 (g) + O_2 (g) \nonumber$
In Group 1, lithium nitrate behaves in the same way, producing lithium oxide, nitrogen dioxide, and oxygen as shown:
$4LiNO_3 (s) \rightarrow 2Li_2O (s) + 4NO_2 (g) + O_2 (g) \nonumber$
The other Group 1 nitrates, however, do not decompose completely at regular laboratory temperatures. They produce the metal nitrite and oxygen, but no nitrogen dioxide:
$2XNO_3 (s) \rightarrow 2XNO_2(s) + O_2 (g) \nonumber$
Each of the nitrates from sodium to cesium decomposes in this way; the only difference is in the temperature required for the reaction to proceed. For larger metals, the decomposition is more difficult and requires higher temperatures.
Heating the carbonates
Most carbonates decompose on heating to the metal oxide and carbon dioxide. For example, a typical Group 2 carbonate like calcium carbonate decomposes like this:
$CaCO_3 (s) \rightarrow CaO(s) + CO_2 \nonumber$
In Group 1, lithium carbonate behaves in the same way, producing lithium oxide and carbon dioxide:
$Li_2CO_3 (s) \rightarrow Li_2O(s) + CO_2 \nonumber$
The rest of the Group 1 carbonates do not decompose at laboratory temperatures, although at higher temperatures this becomes possible. The decomposition temperatures again increase down the Group.
The thermal stability of the hydrogen carbonates
The Group 2 hydrogen carbonates such as calcium hydrogen carbonate are so unstable to heat that they only exist in solution. Any attempt to extract them from solution causes them to decompose to the carbonate, carbon dioxide and water as shown:
$Ca(HCO_3)_2 (aq) \rightarrow CaCO_3 (s) + CO_2 (g) + H_2O (l) \nonumber$
By contrast, the Group 1 hydrogen carbonates are stable enough to exist as solids, although they do decompose easily on heating. For example, this is the reaction for sodium hydrogen carbonate:
$2NaHCO_3 (s) \rightarrow Na_2CO_3 (s) + CO_2 (g) + H_2O (l) \nonumber$
Explaining the trends in thermal stability
Detailed explanations are given for the carbonates because the diagrams are easier to draw. Exactly the same arguments apply to the nitrates or hydrogen carbonates. There are two ways of explaining the increase in thermal stability down the Group. The hard way is in terms of the energetics of the process; the simple way is in terms of the polarizing ability of the positive ions.
Explaining the trend in terms of the polarizing ability of the positive ion
A small positive ion has a large amount of charge packed into a small volume of space; this is especially true if it has a charge greater than +1. An ion with a high charge density has a marked distorting effect on any negative ions which happen to be nearby. A larger positive ion has the same charge spread over a larger volume of space. Its charge density is therefore lower, and it causes less distortion to nearby negative ions.
The structure of the carbonate ion
The molecular structure of carbonate is given below:
This figure shows two carbon-oxygen single bonds and one double bond, with two oxygen atoms each carrying a negative charge. However, experimental data shows that all the carbonate bonds are identical, with the charge spread out over the whole ion (concentrated on the oxygen atoms). In other words, the charges are delocalized.
This is a more complicated version of the bonding in benzene or in ions like ethanoate. The next diagram shows the delocalized electrons. The shading shows electron density, implying a greater chance of finding electrons around the oxygen atoms than near the carbon.
Polarizing the carbonate ion
Imagine that this ion is placed next to a positive ion. The positive ion attracts the delocalized electrons in the carbonate ion towards itself. The carbonate ion becomes polarized. The diagram shows what happens with an ion from Group 2, carrying two positive charges:
If this system is heated, the carbon dioxide breaks free, leaving a metal oxide. The amount of heat required depends on how polarized the ion was. If it is highly polarized, less heat is required than if it is only slightly polarized. If the positive ion only has one positive charge, the polarizing effect is lessened. This is why the Group 1 compounds are more thermally stable than those in Group 2. The Group 1 compound must be heated more because the carbonate ion is less polarized by a singly-charged positive ion.
The smaller the positive ion, the higher the charge density, and the greater the effect on the carbonate ion. As the positive ions get bigger down the group, they have less effect on the carbonate ions near them. To compensate, the compound must be heated more in order to force the carbon dioxide to break off and leave the metal oxide.
In other words, carbonates become more thermally stable down the group.
Extension to nitrates and hydrogen carbonates
The polarization argument exactly the same for these compounds. The small positive ions at the top of the Group polarize the nitrate or hydrogen carbonate ions to a greater extent than the larger positive ions at the bottom. Again, the Group 1 compounds need more heat than those in Group 2 because the Group 1 ions are less polarizing.
The solubility of Group 1 compounds
Group 1 compounds are more soluble than the corresponding ones in Group 2.
The carbonates
Group 2 carbonates are virtually insoluble in water. Magnesium carbonate (the most soluble Group 2 carbonate) has a solubility of about 0.02 g per 100 g of water at room temperature. By contrast, the least soluble Group 1 carbonate is lithium carbonate. A saturated solution has a concentration of about 1.3 g per 100 g of water at 20°C. The other carbonates in the group are very soluble, with solubilities increasing to an astonishing 261.5 g per 100 g of water at this temperature for cesium carbonate.
Solubility of the carbonates increases down Group 1.
The hydroxides
The least soluble hydroxide in Group 1 is lithium hydroxide, but it is still possible to make a solution with a concentration of 12.8 g per 100 g of water at 20°C. The other hydroxides in the group are even more soluble. Solubility of the hydroxides increases down Group 1. In Group 2, the most soluble is barium hydroxide—it is only possible to make a solution of concentration around 3.9 g per 100 g of water at the same temperature of 20°C .
It is difficult to explain the trends in solubility. The discussion on Group 2 of the periodic table explains why the usual explanations for these trends are not accurate.
The Group 1 hydrides
Saline (salt-like) hydrides
Group 1 metal hydrides are white crystalline solids; each contains the metal ion and a hydride ion, H-. They have the same crystal structure as sodium chloride, which is why they are called saline or salt-like hydrides. Because they can react violently with water or moist air, they are normally supplied as suspensions in mineral oil.
Preparation of the Group 1 hydrides
Group 1 hydrides are made by passing hydrogen gas over the heated metal. For example, for lithium hydride:
$2Li + H_2 \rightarrow 2LiH \nonumber$
Reactions of the Group 1 hydrides
Two of the most common reactions include electrolysis and reactions with water.
Electrolysis
On heating, most of these hydrides decompose into the metal and hydrogen before they melt. It is, however, possible to melt lithium hydride and to electrolyze the melt. The metal is deposited at the cathode as expected. Hydrogen is given off at the anode (the positive electrode); this is convincing evidence for the presence of the negative hydride ion in lithium hydride.
The anode equation is:
$2H^- \rightarrow H_2 + 2e^- \nonumber$
The other Group 1 hydrides can be electrolyzed in solution in various molten mixtures such as a mixture of lithium chloride and potassium chloride. These mixtures melt at lower temperatures than the pure chlorides.
Reaction with water
Group 1 hydrides react violently with water releasing hydrogen gas and producing aqueous metal hydroxide. For example, sodium hydride reacts with water to produce sodium hydroxide and hydrogen gas:
$NaH + H_2O \rightarrow NaOH + H_2 \nonumber$ | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__1%3A_The_Alkali_Metals/2Reactions_of_the_Group_1_Elements/Group_1_Comp.txt |
All of Group 1 elements—lithium, sodium, potassium, rubidium and cesium react vigorously or even explosively with cold water. In each case, the aqueous metal hydroxide and hydrogen gas are produced, as shown:
$\ce{ 2X (s) + 2H_2O (l) \rightarrow 2XOH (aq) + H_2 (g)} \nonumber$
where $X$ is any Group 1 metal. In each of the following descriptions, a very small portion of the metal is dropped into a large container of water.
Details for the individual metals
• Lithium: Lithium's density is only about half that of water, so it floats on the surface, fizzing and giving off hydrogen gas. It gradually reacts and disappears, forming a colorless solution of lithium hydroxide. The reaction generates heat slowly, and lithium's melting point is too high for it to melt (this is not the case for sodium).
• Sodium: Sodium also floats in water, but enough heat is given off to melt the sodium (sodium has a lower melting point than lithium) and it melts almost at once to form a small silvery ball that moves rapidly across the surface. The ball leaves a white trail of sodium hydroxide, which soon dissolves to give a colorless solution of sodium hydroxide.
The sodium moves because it is pushed by the hydrogen produced during the reaction. If the sodium becomes trapped on the side of the container, the hydrogen may catch fire and burn with an orange flame. The color is due to contamination of the normally blue hydrogen flame with sodium compounds.
• Potassium: Potassium behaves like sodium except that the reaction is faster and enough heat is given off to ignite the hydrogen. This time the hydrogen flame is contaminated by potassium compounds, so the flame is lilac-colored.
• Rubidium: Rubidium sinks because it is denser than water. It reacts violently and immediately, with everything leaving the container. Rubidium hydroxide solution and hydrogen are formed.
• Cesium: Cesium explodes on contact with water, possibly shattering the container. Cesium hydroxide and hydrogen are formed.
Note: Summary of the trend in reactivity
The Group 1 metals become more reactive towards water down the group.
The Net Enthalpy Changes (Thermodynamics)
It is tempting to conclude that because the reactions get more dramatic down the group, the amount of heat given off increases from lithium to cesium. This is not the case. The table below gives estimates of the enthalpy change for each of the elements undergoing the reaction with water:
$\ce{X (s) + H_2O(l) \rightarrow XOH(aq) + 1/2 H_2 (g) } \nonumber$
Element $\Delta H$ (kJ / mol)
Li -222
Na -184
K -196
Rb -195
Cs -203
There is no consistent pattern in these values; they are all very similar, and counter intuitively, lithium releases the most heat during the reaction. The differences between the reactions are determined at the atomic level. In each case, metal ions in a solid are solvated, as in the reaction below:
$\ce{ X(s) \rightarrow X^+(aq) + e^{-}} \nonumber$
The net enthalpy change for this process can be determined using Hess's Law, and breaking it into several theoretical steps with known enthalpy changes.
$\ce{ X(s) \rightarrow X(g)} \nonumber$
$\ce{ X(g) \rightarrow X^+(g) + e^{-}} \nonumber$
• The final enthalpy change is the hydration enthalpy, or the heat released when the gaseous ion comes into contact with water.
$\ce{ X^+(g) \rightarrow X^{+}(aq)} \nonumber$
These values are tabulated below (all energy values are given in kJ / mol):
Element atomization energy 1st IE hydration enthalpy total
Li +161 +519 -519 +161
Na +109 +494 -406 +197
K +90 +418 -322 +186
Rb +86 +402 -301 +187
Cs +79 +376 -276 +179
There is no overall trend in the overall reaction enthalpy, but each of the component input enthalpies (in which energy must be supplied) decreases down the group, while the hydration enthalpies increase:
1. The atomization energy is a measure of the strength of the metallic bond in each element. This decreases as the size of the atoms and the length of the metallic bond increase. The delocalized electrons are further from the attraction of the nuclei in the larger atoms.
2. The first ionization energy decreases because the electron being removed is more distant from the nucleus with each progressive atom. The extra protons in the nucleus are screened by additional layers of electrons.
3. The hydration enthalpy is a measure of the attraction between the metal ions and lone pairs on water molecules. As the ions increase in size, the water molecules are farther from the attraction of the nucleus. The extra protons in the nucleus are again screened by the extra layers of electrons.
The summation of these effects eliminates any overall pattern. Knowing the atomization energy, the first ionization energy, and the hydration enthalpy, however, reveals useful patterns.
Activation Energies (Kinetics)
Consider the energy input terms:
atomization energy 1st IE total
Li +161 +519 +680
Na +109 +494 +603
K +90 +418 +508
Rb +86 +402 +488
Cs +79 +376 +455
A steady decrease down the group is apparent. From lithium to cesium, less energy is required to form a positive ion. This energy will be recovered (and overcompensated) later, but must be initially supplied. This process is related to the activation energy of the reaction.
The lower the activation energy, the faster the reaction.
Although lithium releases the most heat during the reaction, it does so relatively slowly—not in one short, sharp burst. Cesium, on the other hand, has a significantly lower activation energy, and so although it does not release as much heat overall, it does so extremely quickly, causing an explosion.
Explaining the increase in reactivity down the group
The reactions proceed faster as the energy needed to form positive ions falls. This is in part due to a decrease in ionization energy down the group, and in part to a decrease in atomization energy reflecting weaker metallic bonds from lithium to cesium. This leads to lower activation energies, and therefore faster reactions. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__1%3A_The_Alkali_Metals/2Reactions_of_the_Group_1_Elements/Reactions_of.txt |
Hydrogen is a colorless, odorless and tasteless gas that is the most abundant element in the known universe. It is also the lightest (in terms of atomic mass) and the simplest, having only one proton and one electron (and no neutrons in its most common isotope). It is all around us. It is a component of water (H2O), fats, petroleum, table sugar (C6H12O6), ammonia (NH3), and hydrogen peroxide (H2O2)—things essential to life, as we know it.
Hydrogen Facts
• Atomic Number: 1
• Atomic Symbol: H
• Atomic Weight: 1.0079
• Electronic Configuration: 1s1
• Oxidation States: 1, -1
• Atomic Radius: 78 pm
• Melting Point: -259.34°C
• Boiling Point: -252.87° C
• Elemental Classification: Non-Metal
• At Room Temperature: Colorless & Odorless Diatomic Gas
History of Hydrogen
Hydrogen comes from Greek meaning “water producer” (“hydro” =water and “gennao”=to make). First isolated and identified as an element by Cavendish in 1766, hydrogen was believed to be many different things. Cavendish himself thought that it was "inflammable air from metals", owing to its production by the action of acids on metals. Before that, Robert Boyle and Paracelsus both used reactions of iron and acids to produce hydrogen gas and Antoine Lavoisier gave hydrogen its name because it produced water when ignited in air. Others thought it was pure phlogiston because of its flammability. Hydrogen is among the ten most abundant elements on the planet, but very little is found in elemental form due to its low density and reactivity. Much of the terrestrial hydrogen is locked up in water molecules and organic compounds like hydrocarbons.
Properties of Hydrogen
Hydrogen is a nonmetal and is placed above group in the periodic table because it has ns1 electron configuration like the alkali metals. However, it varies greatly from the alkali metals as it forms cations (H+) more reluctantly than the other alkali metals. Hydrogen‘s ionization energy is 1312 kJ/mol, while lithium (the alkali metal with the highest ionization energy) has an ionization energy of 520 kJ/mol.
Because hydrogen is a nonmetal and forms H- (hydride anions), it is sometimes placed above the halogens in the periodic table. Hydrogen also forms H2 dihydrogen like halogens. However, hydrogen is very different from the halogens. Hydrogen has a much smaller electron affinity than the halogens.
H2 dihydrogen or molecular hydrogen is non-polar with two electrons. There are weak attractive forces between H2 molecules, resulting in low boiling and melting points. However, H2 has very strong intramolecular forces; H2 reactions are generally slow at room temperature due to strong H—H bond. H2 is easily activated by heat, irradiation, or catalysis. Activated hydrogen gas reacts very quickly and exothermically with many substances.
Hydrogen also has an ability to form covalent bonds with a large variety of substances. Because it makes strong O—H bonds, it is a good reducing agent for metal oxides. Example: CuO(s) + H2(g) → Cu(s) + H2O(g) H2(g) passes over CuO(s) to reduce the Cu2+ to Cu(s), while getting oxidized itself.
Reactions of Hydrogen
Hydrogen's low ionization energy makes it act like an alkali metal:
$H_{(g)} \rightarrow H^+_{(g)} + e^- \nonumber$
However, it half-filled valence shell (with a $1s^1$ configuration) with one $e^-$ also causes hydrogen to act like a halogen non-metal to gain noble gas configuration by adding an additional electron
$H_{(g)} + e^- \rightarrow H^-_{(g)} \nonumber$
Reactions of Hydrogen with Active Metals
Hydrogen accepts e- from an active metal to form ionic hydrides like LiH. By forming an ion with -1 charge, the hydrogen behaves like a halogen.
Group 1 metals
$2M_{(s)}+H_{2(g)} \rightarrow 2MH_{(s)} \nonumber$
with $M$ representing Group 1 Alkali metals
Examples:
• $2K_{(s)}+H_{2(g)} \rightarrow 2KH_{(s)}$
• $2K_{(s)}+Cl_{2(g)} \rightarrow 2KCl_{(s)}$
Group 2 metals
$M_{(s)}+H_{2(g)} \rightarrow MH_{2(s)} \nonumber$
with $M$ representing Group 2 Alkaline Earth metals
Example:
• $Ca_{(s)}+H_{2(g)} \rightarrow CaH_{2(s)}$
• $Ca_{(s)}+Cl_{2(g)} \rightarrow CaCl_{2(s)}$
Reactions of Hydrogen with Nonmetals
Unlike metals forming ionic bonds with nonmetals, hydrogen forms polar covalent bonds. Despite being electropositive like the active metals that form ionic bonds with nonmetals, hydrogen is much less electropositive than the active metals, and forms covalent bonds.
Hydrogen + Halogen → Hydrogen Halide
$H_{2(g)}+ Cl_{2(g)} \rightarrow HCl_{(g)} \nonumber$
Hydrogen gas reacting with oxygen to produce water and a large amount of heat: Hydrogen + Oxygen → Water
$(H_{2(g)}+O_{2(g)} \rightarrow H_2O_{(g)} \nonumber$
Reactions with Transition Metals
Reactions of hydrogen with Transition metals (Group 3-12) form metallic hydrides. There is no fixed ratio of hydrogen atom to metal because the hydrogen atoms fill holes between metal atoms in the crystalline structure.
Uses & Application
The vast majority of hydrogen produced industrially today is made either from treatment of methane gas with steam or in the production of "water gas" from the reaction of coal with steam. Most of this hydrogen is used in the Haber process to manufacture ammonia.
Hydrogen is also used for hydrogenation vegetable oils, turning them into margarine and shortening, and some is used for liquid rocket fuel. Liquid hydrogen (combined with liquid oxygen) is a major component of rocket fuel (as mentioned above combination of hydrogen and oxygen relapses a huge amount of energy). Because hydrogen is a good reducing agent, it is used to produce metals like iron, copper, nickel, and cobalt from their ores.
Because one cubic feet of hydrogen can lift about 0.07 lbs, hydrogen lifted airships or Zeppelins became very common in the early 1900s.However, the use of hydrogen for this purpose was largely discontinued around World War II after the explosion of The Hindenburg; this prompted greater use of inert helium, rather than flammable hydrogen for air travel.
Video Showing the explosion of The Hindenburg. (Video from Youtube)
Recently, due to the fear of fossil fuels running out, extensive research is being done on hydrogen as a source of energy.Because of their moderately high energy densities liquid hydrogen and compressed hydrogen gas are possible fuels for the future.A huge advantage in using them is that their combustion only produces water (it burns “clean”). However, it is very costly, and not economically feasible with current technology.
Combustion of fuel produces energy that can be converted into electrical energy when energy in the steam turns a turbine to drive a generator. However, this is not very efficient because a great deal of energy is lost as heat. The production of electricity using voltaic cell can yield more electricity (a form of usable energy). Voltaic cells that transform chemical energy in fuels (like H2 and CH4) are called fuel cells. These are not self-contained and so are not considered batteries. The hydrogen cell is a type of fuel cell involving the reaction between H2(g) with O2(g) to form liquid water; this cell is twice as efficient as the best internal combustion engine. In the cell (in basic conditions), the oxygen is reduced at the cathode, while the hydrogen is oxidized at the anode.
Reduction: O2(g)+2H2O(l)+4e- → 4OH-(aq)
Oxidation: H2(g) + 2OH-(aq) → 2H2O(l) + 2e-
Overall: 2H2(g) + O2(g) → 2H2O(l)
E°cell= Reduction- Oxidation= E°O2/OH- - E°H2O/H2 = 0.401V – (-0.828V) = +1.23
However, this technology is far from being used in everyday life due to its great costs.
Image of A Hydrogen Fuel Cell. (Image made by Ridhi Sachdev)
Natural Occurrence & Other Sources
Naturally Occurring Hydrogen
Hydrogen is the fuel for reactions of the Sun and other stars (fusion reactions). Hydrogen is the lightest and most abundant element in the universe. About 70%- 75% of the universe is composed of hydrogen by mass. All stars are essentially large masses of hydrogen gas that produce enormous amounts of energy through the fusion of hydrogen atoms at their dense cores. In smaller stars, hydrogen atoms collided and fused to form helium and other light elements like nitrogen and carbon(essential for life). In the larger stars, fusion produces the lighter and heavier elements like calcium, oxygen, and silicon.
On Earth, hydrogen is mostly found in association with oxygen; its most abundant form being water (H2O). Hydrogen is only .9% by mass and 15% by volume abundant on the earth, despite water covering about 70% of the planet. Because hydrogen is so light, there is only 0.5 ppm (parts per million) in the atmosphere, which is a good thing considering it is EXTREMELY flammable.
Other Sources of Hydrogen
Hydrogen gas can be prepared by reacting a dilute strong acid like hydrochloric acids with an active metal. The metal becomes oxides, while the H+ (from the acid) is reduced to hydrogen gas. This method is only practical for producing small amounts of hydrogen in the lab, but is much too costly for industrial production:
$Zn_{(s)} + 2H^+_{(aq)} \rightarrow Zn^{2+}_{(aq)} + H_{2(g)} \nonumber$
The purest form of H2(g) can come from electrolysis of H2O(l), the most common hydrogen compound on this plant. This method is also not commercially viable because it requires a significant amount of energy ($\Delta H = 572 \;kJ$):
$2H_2O_{(l)} \rightarrow 2H_{2(g)} + O_{2(g)} \nonumber$
$H_2O$ is the most abundant form of hydrogen on the planet, so it seems logical to try to extract hydrogen from water without electrolysis of water. To do so, we must reduce the hydrogen with +1 oxidation state to hydrogen with 0 oxidation state (in hydrogen gas). Three commonly used reducing agents are carbon (in coke or coal), carbon monoxide, and methane. These react with water vapor form H2(g):
$C_{(s)} + 2H_2O_{(g)} \rightarrow CO(g) + H_{2(g)} \nonumber$
$CO_{(g)} + 2H_2O_{(g)} \rightarrow CO2 + H_{2(g)} \nonumber$
Reforming of Methane:
$CH_{4(g)} + H_2O_{(g)} \rightarrow CO(g) + 3H_{2(g)} \nonumber$
These three methods are most industrially feasible (cost effective) methods of producing H2(g).
Isotopes
There are two important isotopes of hydrogen. Deuterium (2H) has an abundance of 0.015% of terrestrial hydrogen and the nucleus of the isotope contains one neutron.
• Protium (1H) is the most common isotope, consisting of 99.98% of naturally occurring hydrogen. It is a nucleus containing a single proton.
• Deuterium (2H) is another an isotope containing a proton and neutron, consisting of only 0.0156% of the naturally occurring hydrogen. Commonly indicated with symbol D and sometimes called heavy hydrogen, deuterium is separated by the fractional distillation of liquid hydrogen but it can also be produced by the prolonged electrolysis of ordinary water. Approximately 100,000 gallons of water will produce a single gallon of D2O, "heavy water". This special kind of water has a higher density, melting point, and boiling point than regular water and used as a moderator in some fission power reactors. Deuterium fuel is used in experimental fusion reactors. Replacing protium with deuterium has important uses for exploring reaction mechanisms via the kinetic isotope effect.
• Tritium (3H) contains two neutrons in its nucleus and is radioactive with a 12.3-year half-life, which is continuously formed in the upper atmosphere due to cosmic rays. It is can also be made in a lab from Lithium-6 in a nuclear reactor. Tritium is also used in hydrogen bombs. It is very rare (about 1 in every 1,018 atoms) and is formed in the environment by cosmic ray bombardment. Most tritium is manufactured by bombarding Li with neutrons. Tritium is used in thermonuclear weapons and experimental fusion reactors.
Problems
1. Write the reaction of Na(s) with H2(g).
2. What is the name of the radioactive isotope of hydrogen?
3. What characteristics of alkali metals does hydrogen display?
4. What characteristics of halogens does hydrogen display?
5. How does the electronegativity of hydrogen compare to that of the halogens?
6. What is the electron configuration of a neutral hydrogen atom.
Answers
1. 2Na(s) + H2(g)→ 2NaH(s)
2. Tritium
3. Hydrogen is placed above group in the periodic table because it has ns1 electron configuration like the alkali metals. However, it varies greatly from the alkali metals as it forms cations (H+) more reluctantly than the other alkali metals. Hydrogen‘s ionization energy is 1312 kJ/mol, while lithium (the alkali metal with the highest ionization energy) has an ionization energy of 520 kJ/mol.
4. Because hydrogen is a nonmetal and forms H- (hydride anions), it is sometimes placed above the halogens in the periodic table. Hydrogen also forms H2 dihydrogen like halogens. However, hydrogen is very different from the halogens. Hydrogen has a much smaller electron affinity than the halogens.
5. Hydrogen is less electronegative than the halogens.
6. 1s1
Contributors and Attributions
• Ridhi Sachdev (UC Davis)
Stephen R. Marsden
Z001 Chemistry of Hydrogen (Z1)
Hydrogen chloride (HCl) is a colorless gas which forms white fumes of hydrochloric acid when brought into contact with atmospheric humidity. Inhalation of the gas can cause severe burns of the nose, throat, and upper respiratory tract (which may lead to death in severe cases). Hydrogen chloride may also result in severe burns of the eyes.
The hydrogen and the chlorine atom are connected by a covalent single bond, which is highly polar, since the chlorine atom is much more electronegative than the hydrogen atom. Thus the molecule shows a large dipole moment with the negative charge at the chlorine atom.
When dissolved in water, the HCl gas dissociates and forms hydronium and chloride ions:
HCl + H2O H3O+ + Cl-
This solution is called hydrochloric acid(1), which is a strong acid with a high acid dissociation constant.
Hydrogen Pe
Hydrogen peroxide (H2O2) is a viscous liquid (mp.: -0.89°C, bp.: 152.1°C, density: 1.448 g/cm3 at 20°C) that has strong oxidizing properties.
It is commonly used (in concentrations typically around 5%) to bleach human hair, hence the phrases "peroxide blonde" and "bottle blonde". It burns the skin upon contact in sufficient concentration. In lower concentrations (3%), it is used medically for cleaning wounds and removing dead tissue. However, recent studies have indicated that hydrogen peroxide is toxic to new cells and is therefore not recommended for wound care. The same solution is often used by medical professionals to clean blood from cloth and equipment.
H2O2 is produced by a combination of electrolysis of sulfuric acid and subsequent hydrolysis of the resulting peroxo-disulfuric acid:
2 H2SO4 H2S2O8 + H2
H2S2O8 + 2 H2O H2O2 2 H2SO4
2 H2O H2O2 + H2
Hydrogen peroxide decomposes exothermically into water and oxygen gas (46.87 kcal/mol). However, at room temperature the rate of decomposition is very low, so that pure H2O2 is stable (metastable). Catalysts (like pulverized silver, gold, platinum, manganese dioxed, iron and copper salts, alkali salts, dust, activated carbon, etc.) dramatically increase the rate of decomposition. High concentrations of H2O2 may lead to explosions when it comes into contact with catalysts. H2O2 may be stabilized by adding phosphoric acid, sodium diphosphate, uric acid, or barbituric acid. Pure H2O2 without any stabilizers can only be stored in glass bottles which are coated with paraffin way, or in pure aluminum (better than 99.9 %) bottles. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__1%3A_The_Alkali_Metals/Z001_Chemistry_of_Hydrogen_%28Z1%29/Hydrogen_Ch.txt |
Hydrogen is a colorless, odorless and tasteless gas that is the most abundant element in the known universe. It is also the lightest (in terms of atomic mass) and the simplest, having only one proton and one electron (and no neutrons in its most common isotope). It is all around us. It is a component of water (H2O), fats, petroleum, table sugar (C6H12O6), ammonia (NH3), and hydrogen peroxide (H2O2)—things essential to life, as we know it.
Hydrogen Facts
• Atomic Number: 1
• Atomic Symbol: H
• Atomic Weight: 1.0079
• Electronic Configuration: 1s1
• Oxidation States: 1, -1
• Atomic Radius: 78 pm
• Melting Point: -259.34°C
• Boiling Point: -252.87° C
• Elemental Classification: Non-Metal
• At Room Temperature: Colorless & Odorless Diatomic Gas
History of Hydrogen
Hydrogen comes from Greek meaning “water producer” (“hydro” =water and “gennao”=to make). First isolated and identified as an element by Cavendish in 1766, hydrogen was believed to be many different things. Cavendish himself thought that it was "inflammable air from metals", owing to its production by the action of acids on metals. Before that, Robert Boyle and Paracelsus both used reactions of iron and acids to produce hydrogen gas and Antoine Lavoisier gave hydrogen its name because it produced water when ignited in air. Others thought it was pure phlogiston because of its flammability. Hydrogen is among the ten most abundant elements on the planet, but very little is found in elemental form due to its low density and reactivity. Much of the terrestrial hydrogen is locked up in water molecules and organic compounds like hydrocarbons.
Properties of Hydrogen
Hydrogen is a nonmetal and is placed above group in the periodic table because it has ns1 electron configuration like the alkali metals. However, it varies greatly from the alkali metals as it forms cations (H+) more reluctantly than the other alkali metals. Hydrogen‘s ionization energy is 1312 kJ/mol, while lithium (the alkali metal with the highest ionization energy) has an ionization energy of 520 kJ/mol.
Because hydrogen is a nonmetal and forms H- (hydride anions), it is sometimes placed above the halogens in the periodic table. Hydrogen also forms H2 dihydrogen like halogens. However, hydrogen is very different from the halogens. Hydrogen has a much smaller electron affinity than the halogens.
H2 dihydrogen or molecular hydrogen is non-polar with two electrons. There are weak attractive forces between H2 molecules, resulting in low boiling and melting points. However, H2 has very strong intramolecular forces; H2 reactions are generally slow at room temperature due to strong H—H bond. H2 is easily activated by heat, irradiation, or catalysis. Activated hydrogen gas reacts very quickly and exothermically with many substances.
Hydrogen also has an ability to form covalent bonds with a large variety of substances. Because it makes strong O—H bonds, it is a good reducing agent for metal oxides. Example: CuO(s) + H2(g) → Cu(s) + H2O(g) H2(g) passes over CuO(s) to reduce the Cu2+ to Cu(s), while getting oxidized itself.
Reactions of Hydrogen
Hydrogen's low ionization energy makes it act like an alkali metal:
$H_{(g)} \rightarrow H^+_{(g)} + e^- \nonumber$
However, it half-filled valence shell (with a $1s^1$ configuration) with one $e^-$ also causes hydrogen to act like a halogen non-metal to gain noble gas configuration by adding an additional electron
$H_{(g)} + e^- \rightarrow H^-_{(g)} \nonumber$
Reactions of Hydrogen with Active Metals
Hydrogen accepts e- from an active metal to form ionic hydrides like LiH. By forming an ion with -1 charge, the hydrogen behaves like a halogen.
Group 1 metals
$2M_{(s)}+H_{2(g)} \rightarrow 2MH_{(s)} \nonumber$
with $M$ representing Group 1 Alkali metals
Examples:
• $2K_{(s)}+H_{2(g)} \rightarrow 2KH_{(s)}$
• $2K_{(s)}+Cl_{2(g)} \rightarrow 2KCl_{(s)}$
Group 2 metals
$M_{(s)}+H_{2(g)} \rightarrow MH_{2(s)} \nonumber$
with $M$ representing Group 2 Alkaline Earth metals
Example:
• $Ca_{(s)}+H_{2(g)} \rightarrow CaH_{2(s)}$
• $Ca_{(s)}+Cl_{2(g)} \rightarrow CaCl_{2(s)}$
Reactions of Hydrogen with Nonmetals
Unlike metals forming ionic bonds with nonmetals, hydrogen forms polar covalent bonds. Despite being electropositive like the active metals that form ionic bonds with nonmetals, hydrogen is much less electropositive than the active metals, and forms covalent bonds.
Hydrogen + Halogen → Hydrogen Halide
$H_{2(g)}+ Cl_{2(g)} \rightarrow HCl_{(g)} \nonumber$
Hydrogen gas reacting with oxygen to produce water and a large amount of heat: Hydrogen + Oxygen → Water
$(H_{2(g)}+O_{2(g)} \rightarrow H_2O_{(g)} \nonumber$
Reactions with Transition Metals
Reactions of hydrogen with Transition metals (Group 3-12) form metallic hydrides. There is no fixed ratio of hydrogen atom to metal because the hydrogen atoms fill holes between metal atoms in the crystalline structure.
Uses & Application
The vast majority of hydrogen produced industrially today is made either from treatment of methane gas with steam or in the production of "water gas" from the reaction of coal with steam. Most of this hydrogen is used in the Haber process to manufacture ammonia.
Hydrogen is also used for hydrogenation vegetable oils, turning them into margarine and shortening, and some is used for liquid rocket fuel. Liquid hydrogen (combined with liquid oxygen) is a major component of rocket fuel (as mentioned above combination of hydrogen and oxygen relapses a huge amount of energy). Because hydrogen is a good reducing agent, it is used to produce metals like iron, copper, nickel, and cobalt from their ores.
Because one cubic feet of hydrogen can lift about 0.07 lbs, hydrogen lifted airships or Zeppelins became very common in the early 1900s.However, the use of hydrogen for this purpose was largely discontinued around World War II after the explosion of The Hindenburg; this prompted greater use of inert helium, rather than flammable hydrogen for air travel.
Video Showing the explosion of The Hindenburg. (Video from Youtube)
Recently, due to the fear of fossil fuels running out, extensive research is being done on hydrogen as a source of energy.Because of their moderately high energy densities liquid hydrogen and compressed hydrogen gas are possible fuels for the future.A huge advantage in using them is that their combustion only produces water (it burns “clean”). However, it is very costly, and not economically feasible with current technology.
Combustion of fuel produces energy that can be converted into electrical energy when energy in the steam turns a turbine to drive a generator. However, this is not very efficient because a great deal of energy is lost as heat. The production of electricity using voltaic cell can yield more electricity (a form of usable energy). Voltaic cells that transform chemical energy in fuels (like H2 and CH4) are called fuel cells. These are not self-contained and so are not considered batteries. The hydrogen cell is a type of fuel cell involving the reaction between H2(g) with O2(g) to form liquid water; this cell is twice as efficient as the best internal combustion engine. In the cell (in basic conditions), the oxygen is reduced at the cathode, while the hydrogen is oxidized at the anode.
Reduction: O2(g)+2H2O(l)+4e- → 4OH-(aq)
Oxidation: H2(g) + 2OH-(aq) → 2H2O(l) + 2e-
Overall: 2H2(g) + O2(g) → 2H2O(l)
E°cell= Reduction- Oxidation= E°O2/OH- - E°H2O/H2 = 0.401V – (-0.828V) = +1.23
However, this technology is far from being used in everyday life due to its great costs.
Image of A Hydrogen Fuel Cell. (Image made by Ridhi Sachdev)
Natural Occurrence & Other Sources
Naturally Occurring Hydrogen
Hydrogen is the fuel for reactions of the Sun and other stars (fusion reactions). Hydrogen is the lightest and most abundant element in the universe. About 70%- 75% of the universe is composed of hydrogen by mass. All stars are essentially large masses of hydrogen gas that produce enormous amounts of energy through the fusion of hydrogen atoms at their dense cores. In smaller stars, hydrogen atoms collided and fused to form helium and other light elements like nitrogen and carbon(essential for life). In the larger stars, fusion produces the lighter and heavier elements like calcium, oxygen, and silicon.
On Earth, hydrogen is mostly found in association with oxygen; its most abundant form being water (H2O). Hydrogen is only .9% by mass and 15% by volume abundant on the earth, despite water covering about 70% of the planet. Because hydrogen is so light, there is only 0.5 ppm (parts per million) in the atmosphere, which is a good thing considering it is EXTREMELY flammable.
Other Sources of Hydrogen
Hydrogen gas can be prepared by reacting a dilute strong acid like hydrochloric acids with an active metal. The metal becomes oxides, while the H+ (from the acid) is reduced to hydrogen gas. This method is only practical for producing small amounts of hydrogen in the lab, but is much too costly for industrial production:
$Zn_{(s)} + 2H^+_{(aq)} \rightarrow Zn^{2+}_{(aq)} + H_{2(g)} \nonumber$
The purest form of H2(g) can come from electrolysis of H2O(l), the most common hydrogen compound on this plant. This method is also not commercially viable because it requires a significant amount of energy ($\Delta H = 572 \;kJ$):
$2H_2O_{(l)} \rightarrow 2H_{2(g)} + O_{2(g)} \nonumber$
$H_2O$ is the most abundant form of hydrogen on the planet, so it seems logical to try to extract hydrogen from water without electrolysis of water. To do so, we must reduce the hydrogen with +1 oxidation state to hydrogen with 0 oxidation state (in hydrogen gas). Three commonly used reducing agents are carbon (in coke or coal), carbon monoxide, and methane. These react with water vapor form H2(g):
$C_{(s)} + 2H_2O_{(g)} \rightarrow CO(g) + H_{2(g)} \nonumber$
$CO_{(g)} + 2H_2O_{(g)} \rightarrow CO2 + H_{2(g)} \nonumber$
Reforming of Methane:
$CH_{4(g)} + H_2O_{(g)} \rightarrow CO(g) + 3H_{2(g)} \nonumber$
These three methods are most industrially feasible (cost effective) methods of producing H2(g).
Isotopes
There are two important isotopes of hydrogen. Deuterium (2H) has an abundance of 0.015% of terrestrial hydrogen and the nucleus of the isotope contains one neutron.
• Protium (1H) is the most common isotope, consisting of 99.98% of naturally occurring hydrogen. It is a nucleus containing a single proton.
• Deuterium (2H) is another an isotope containing a proton and neutron, consisting of only 0.0156% of the naturally occurring hydrogen. Commonly indicated with symbol D and sometimes called heavy hydrogen, deuterium is separated by the fractional distillation of liquid hydrogen but it can also be produced by the prolonged electrolysis of ordinary water. Approximately 100,000 gallons of water will produce a single gallon of D2O, "heavy water". This special kind of water has a higher density, melting point, and boiling point than regular water and used as a moderator in some fission power reactors. Deuterium fuel is used in experimental fusion reactors. Replacing protium with deuterium has important uses for exploring reaction mechanisms via the kinetic isotope effect.
• Tritium (3H) contains two neutrons in its nucleus and is radioactive with a 12.3-year half-life, which is continuously formed in the upper atmosphere due to cosmic rays. It is can also be made in a lab from Lithium-6 in a nuclear reactor. Tritium is also used in hydrogen bombs. It is very rare (about 1 in every 1,018 atoms) and is formed in the environment by cosmic ray bombardment. Most tritium is manufactured by bombarding Li with neutrons. Tritium is used in thermonuclear weapons and experimental fusion reactors.
Problems
1. Write the reaction of Na(s) with H2(g).
2. What is the name of the radioactive isotope of hydrogen?
3. What characteristics of alkali metals does hydrogen display?
4. What characteristics of halogens does hydrogen display?
5. How does the electronegativity of hydrogen compare to that of the halogens?
6. What is the electron configuration of a neutral hydrogen atom.
Answers
1. 2Na(s) + H2(g)→ 2NaH(s)
2. Tritium
3. Hydrogen is placed above group in the periodic table because it has ns1 electron configuration like the alkali metals. However, it varies greatly from the alkali metals as it forms cations (H+) more reluctantly than the other alkali metals. Hydrogen‘s ionization energy is 1312 kJ/mol, while lithium (the alkali metal with the highest ionization energy) has an ionization energy of 520 kJ/mol.
4. Because hydrogen is a nonmetal and forms H- (hydride anions), it is sometimes placed above the halogens in the periodic table. Hydrogen also forms H2 dihydrogen like halogens. However, hydrogen is very different from the halogens. Hydrogen has a much smaller electron affinity than the halogens.
5. Hydrogen is less electronegative than the halogens.
6. 1s1
Contributors and Attributions
• Ridhi Sachdev (UC Davis)
Stephen R. Marsden
Z001 Chemistry of Hydrogen (Z1)
Hydrogen chloride (HCl) is a colorless gas which forms white fumes of hydrochloric acid when brought into contact with atmospheric humidity. Inhalation of the gas can cause severe burns of the nose, throat, and upper respiratory tract (which may lead to death in severe cases). Hydrogen chloride may also result in severe burns of the eyes.
The hydrogen and the chlorine atom are connected by a covalent single bond, which is highly polar, since the chlorine atom is much more electronegative than the hydrogen atom. Thus the molecule shows a large dipole moment with the negative charge at the chlorine atom.
When dissolved in water, the HCl gas dissociates and forms hydronium and chloride ions:
HCl + H2O H3O+ + Cl-
This solution is called hydrochloric acid(1), which is a strong acid with a high acid dissociation constant.
Hydrogen Peroxi
Hydrogen peroxide (H2O2) is a viscous liquid (mp.: -0.89°C, bp.: 152.1°C, density: 1.448 g/cm3 at 20°C) that has strong oxidizing properties.
It is commonly used (in concentrations typically around 5%) to bleach human hair, hence the phrases "peroxide blonde" and "bottle blonde". It burns the skin upon contact in sufficient concentration. In lower concentrations (3%), it is used medically for cleaning wounds and removing dead tissue. However, recent studies have indicated that hydrogen peroxide is toxic to new cells and is therefore not recommended for wound care. The same solution is often used by medical professionals to clean blood from cloth and equipment.
H2O2 is produced by a combination of electrolysis of sulfuric acid and subsequent hydrolysis of the resulting peroxo-disulfuric acid:
2 H2SO4 H2S2O8 + H2
H2S2O8 + 2 H2O H2O2 2 H2SO4
2 H2O H2O2 + H2
Hydrogen peroxide decomposes exothermically into water and oxygen gas (46.87 kcal/mol). However, at room temperature the rate of decomposition is very low, so that pure H2O2 is stable (metastable). Catalysts (like pulverized silver, gold, platinum, manganese dioxed, iron and copper salts, alkali salts, dust, activated carbon, etc.) dramatically increase the rate of decomposition. High concentrations of H2O2 may lead to explosions when it comes into contact with catalysts. H2O2 may be stabilized by adding phosphoric acid, sodium diphosphate, uric acid, or barbituric acid. Pure H2O2 without any stabilizers can only be stored in glass bottles which are coated with paraffin way, or in pure aluminum (better than 99.9 %) bottles. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__1%3A_The_Alkali_Metals/Z001_Chemistry_of_Hydrogen_(Z1)/Hydrogen_Chlori.txt |
Lithium is a rare element found primarily in molten rock and saltwater in very small amounts. It is understood to be non-vital in human biological processes, although it is used in many drug treatments due to its positive effects on the human brain. Because of its reactive properties, humans have utilized lithium in batteries, nuclear fusion reactions, and thermonuclear weapons.
Introduction
Lithium was first identified as a component of of the mineral petalite and was discovered in 1817 by Johan August Arfwedson, but not isolated until some time later by W.T. Brande and Sir Humphry Davy. In its mineral forms it accounts for only 0.0007% of the earth's crust. It compounds are used in certain kinds of glass and porcelain products. More recently lithium has become important in dry-cell batteries and nuclear reactors. Some compounds of lithium have been used to treat manic depressives.
Lithium is an alkali metal with the atomic number = 3 and an atomic mass of 6.941 g/mol. This means that lithium has 3 protons, 3 electrons and 4 neutrons (6.941 - 3 = ~4). Being an alkali metal, lithium is a soft, flammable, and highly reactive metal that tends to form hydroxides. It also has a pretty low density and under standard conditions, it is the least dense solid element.
Properties
Lithium is the lightest of all metals and is named from the Greek work for stone (lithos). It is the first member of the Alkali Metal family. It is less dense than water (with which it reacts) and forms a black oxide in contact with air.
Table 1. Properties of lithium.
Atomic Number 3
Atomic Mass 6.941 g/mol
Atomic Radius 152 pm
Density 0.534 g/cm3
Color light silver
Melting point 453.69 K
Boiling point 1615 K
Heat of fusion 3.00 kJ/mol
Heat of vaporization 147.1 kJ/mol
Specific heat capacity 24.860 kJ/mol
First ionization energy 520.2 kJ/mol
Oxidation states +1, -1
Electronegativity 0.98
Crystal structure body-centered cubic
Magnetism paramagnetic
2 stable isotopes 6Li (7.5%) and 7Li (92.5%)
Periodic Trends of Lithium
Being on the upper left side of the Periodic Table, lithium has a fairly low electronegativity and electron affinity as compared to the rest of the elements. Also, lithium has high metallic character and subsequently lower nonmetallic character when compared with the other elements. Lithium has a higher atomic radius than most of the elements on the Periodic Table. In compounds lithium (like all the alkali metals) has a +1 charge. In its pure form it is soft and silvery white and has a relatively low melting point (181oC).
Reactivity
Lithium is part of the Group 1 Alkali Metals, which are highly reactive and are never found in their pure form in nature. This is due to their electron configuration, in that they have a single valence electron (Figure 1) which is very easily given up in order to create bonds and form compounds.
_↑ ↓_ _↑__
1s2 2s1
Reactions with Water
When placed in contact with water, pure lithium reacts to form lithium hydroxide and hydrogen gas.
$2Li (s) + 2H_2O (l) \rightarrow 2LiOH (aq) + H_2 (g) \nonumber$
Out of all the group 1 metals, lithium reacts the least violently, slowly releasing the hydrogen gas which may create a bright orange flame only if a substantial amount of lithium is used. This occurs because lithium has the highest activation energy of its group - that is, it takes more energy to remove lithium's one valence electron than with other group 1 elements, because lithium's electron is closer to its nucleus. Atoms with higher activation energies will react slower, although lithium will release more total heat through the entire process.
Reactions with Air
Pure lithium will form lithium hydroxide due to moisture in the air, as well as lithium nitride ($Li_3N$) from $N_2$ gas, and lithium carbonate $(Li_2CO_3$) from carbon dioxide. These compounds give the normally the silver-white metal a black tarnish. Additionally, it will combust with oxygen as a red flame to form lithium oxide.
$4Li (s) + O_2 (g) \rightarrow 2Li_2O \nonumber$
Applications
In its mineral forms it accounts for only 0.0007% of the earth's crust. It compounds are used in certain kinds of glass and porcelain products. More recently lithium has become important in dry-cell batteries and nuclear reactors. Some compounds of lithium have been used to treat manic depressives.
Batteries
Lithium is able to be used in the function of a Lithium battery in which the Lithium metal serves as the anode. Lithium ions serve in lithium ion batteries (chargeable) in which the lithium ions move from the negative to positive electrode when discharging, and vice versa when charging.
Heat Transfer
Lithium has the highest specific heat capacity of the solids, Lithium tends to be used as a cooler for heat transfer techniques and applications.
Sources and Extraction
Lithium is most commonly found combined with aluminum, silicon, and oxygen to form the minerals known as spodumene (LiAl(SiO3)2) or petalite/castorite (LiAlSi4O10). These have been found on each of the 6 inhabited continents, but they are mined primarily in Western Australia, China, and Chile. Mineral sources of lithium are becoming less essential, as methods have now been developed to make use of the lithium salts found in saltwater.
Extraction from minerals
The mineral forms of lithium are heated to a high enough temperature (1200 K - 1300 K) in order to crumble them and thus allow for subsequent reactions to more easily take place. After this process, one of three methods can be applied.
1. The use of sulfuric acid and sodium carbonate to allow the iron and aluminum to precipitate from the ore - from there, more sodium carbonate is applied to the remaining material allow the lithium to precipitate out, forming lithium carbonate. This is treated with hydrochloric acid to form lithium chloride.
2. The use of limestone to calcinate the ore, and then leaching with water, forming lithium hydroxide. Again, this is treated with hydrochloric acid to form lithium chloride.
3. The use of sulfuric acid, and then leaching with water, forming lithium sulfate monohydrate. This is treated with sodium carbonate to form lithium carbonate, and then hydrochloric acid to form lithium chloride.
The lithium chloride obtained from any of the three methods undergoes an oxidation-reduction reaction in an electrolytic cell, to separate the chloride ions from the lithium ions. The chloride ions are oxidized, and the lithium ions are reduced.
$2Cl^- - 2e^- \rightarrow Cl_2 \;\; \text{(oxidation)} \nonumber$
$Li^+ + e^- \rightarrow Li \;\; \text{(reduction)} \nonumber$
Extraction from Saltwater
Saltwater naturally contains lithium chloride, which must be extracted in the form of lithium carbonate, then it is re-treated, separated into its ions, and reduced in the same electrolytic process as in extraction from lithium ores. Only three saltwater lakes in the world are currently used for lithium extraction, in Nevada, Chile, and Argentina.
Saltwater is channeled into shallow ponds and over a period of a year or more, water evaporates out to leave behind various salts. Lime is used to remove the magnesium salt, so that the remaining solution contains a fairly concentrated amount of lithium chloride. The solution is then treated with sodium carbonate in order for usable lithium carbonate to precipitate out.
Problems
1. With which group of elements will lithium form compounds the most easily with?
2. What is the electron configuration of Li+?
3. What are some common uses of lithium?
4. For a lithium-ion battery containing LiCoO2, should the compound be placed in the anode or cathode?
5. Given that 7Li is 7.0160 amu and 6Li is 6.0151 amu, and their percent abundance is 92.58% and 7.42% respectively, what is the atomic mass of lithium?
Solutions
1. Group 17 Halogens (lithium forms strongly inic bonds with them, as halogens are highly electronegative and lithium has a free electron)
2. 1s2
3. Lithium-ion batteries, disposable lithium batteries, pyrotechnics, creation of strong metal alloys, etc.
4. Anode - lithium is oxidized (LiCoO2 → Li+ + CoO2)
5. 6.942 g/mol
Contributors and Attributions
• Katherine Szelong (UCD), Kevin Fan | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__1%3A_The_Alkali_Metals/Z003_Chemistry_of_Lithium_%28Z3%29.txt |
Lithium is a rare element found primarily in molten rock and saltwater in very small amounts. It is understood to be non-vital in human biological processes, although it is used in many drug treatments due to its positive effects on the human brain. Because of its reactive properties, humans have utilized lithium in batteries, nuclear fusion reactions, and thermonuclear weapons.
Introduction
Lithium was first identified as a component of of the mineral petalite and was discovered in 1817 by Johan August Arfwedson, but not isolated until some time later by W.T. Brande and Sir Humphry Davy. In its mineral forms it accounts for only 0.0007% of the earth's crust. It compounds are used in certain kinds of glass and porcelain products. More recently lithium has become important in dry-cell batteries and nuclear reactors. Some compounds of lithium have been used to treat manic depressives.
Lithium is an alkali metal with the atomic number = 3 and an atomic mass of 6.941 g/mol. This means that lithium has 3 protons, 3 electrons and 4 neutrons (6.941 - 3 = ~4). Being an alkali metal, lithium is a soft, flammable, and highly reactive metal that tends to form hydroxides. It also has a pretty low density and under standard conditions, it is the least dense solid element.
Properties
Lithium is the lightest of all metals and is named from the Greek work for stone (lithos). It is the first member of the Alkali Metal family. It is less dense than water (with which it reacts) and forms a black oxide in contact with air.
Table 1. Properties of lithium.
Atomic Number 3
Atomic Mass 6.941 g/mol
Atomic Radius 152 pm
Density 0.534 g/cm3
Color light silver
Melting point 453.69 K
Boiling point 1615 K
Heat of fusion 3.00 kJ/mol
Heat of vaporization 147.1 kJ/mol
Specific heat capacity 24.860 kJ/mol
First ionization energy 520.2 kJ/mol
Oxidation states +1, -1
Electronegativity 0.98
Crystal structure body-centered cubic
Magnetism paramagnetic
2 stable isotopes 6Li (7.5%) and 7Li (92.5%)
Periodic Trends of Lithium
Being on the upper left side of the Periodic Table, lithium has a fairly low electronegativity and electron affinity as compared to the rest of the elements. Also, lithium has high metallic character and subsequently lower nonmetallic character when compared with the other elements. Lithium has a higher atomic radius than most of the elements on the Periodic Table. In compounds lithium (like all the alkali metals) has a +1 charge. In its pure form it is soft and silvery white and has a relatively low melting point (181oC).
Reactivity
Lithium is part of the Group 1 Alkali Metals, which are highly reactive and are never found in their pure form in nature. This is due to their electron configuration, in that they have a single valence electron (Figure 1) which is very easily given up in order to create bonds and form compounds.
_↑ ↓_ _↑__
1s2 2s1
Reactions with Water
When placed in contact with water, pure lithium reacts to form lithium hydroxide and hydrogen gas.
$2Li (s) + 2H_2O (l) \rightarrow 2LiOH (aq) + H_2 (g) \nonumber$
Out of all the group 1 metals, lithium reacts the least violently, slowly releasing the hydrogen gas which may create a bright orange flame only if a substantial amount of lithium is used. This occurs because lithium has the highest activation energy of its group - that is, it takes more energy to remove lithium's one valence electron than with other group 1 elements, because lithium's electron is closer to its nucleus. Atoms with higher activation energies will react slower, although lithium will release more total heat through the entire process.
Reactions with Air
Pure lithium will form lithium hydroxide due to moisture in the air, as well as lithium nitride ($Li_3N$) from $N_2$ gas, and lithium carbonate $(Li_2CO_3$) from carbon dioxide. These compounds give the normally the silver-white metal a black tarnish. Additionally, it will combust with oxygen as a red flame to form lithium oxide.
$4Li (s) + O_2 (g) \rightarrow 2Li_2O \nonumber$
Applications
In its mineral forms it accounts for only 0.0007% of the earth's crust. It compounds are used in certain kinds of glass and porcelain products. More recently lithium has become important in dry-cell batteries and nuclear reactors. Some compounds of lithium have been used to treat manic depressives.
Batteries
Lithium is able to be used in the function of a Lithium battery in which the Lithium metal serves as the anode. Lithium ions serve in lithium ion batteries (chargeable) in which the lithium ions move from the negative to positive electrode when discharging, and vice versa when charging.
Heat Transfer
Lithium has the highest specific heat capacity of the solids, Lithium tends to be used as a cooler for heat transfer techniques and applications.
Sources and Extraction
Lithium is most commonly found combined with aluminum, silicon, and oxygen to form the minerals known as spodumene (LiAl(SiO3)2) or petalite/castorite (LiAlSi4O10). These have been found on each of the 6 inhabited continents, but they are mined primarily in Western Australia, China, and Chile. Mineral sources of lithium are becoming less essential, as methods have now been developed to make use of the lithium salts found in saltwater.
Extraction from minerals
The mineral forms of lithium are heated to a high enough temperature (1200 K - 1300 K) in order to crumble them and thus allow for subsequent reactions to more easily take place. After this process, one of three methods can be applied.
1. The use of sulfuric acid and sodium carbonate to allow the iron and aluminum to precipitate from the ore - from there, more sodium carbonate is applied to the remaining material allow the lithium to precipitate out, forming lithium carbonate. This is treated with hydrochloric acid to form lithium chloride.
2. The use of limestone to calcinate the ore, and then leaching with water, forming lithium hydroxide. Again, this is treated with hydrochloric acid to form lithium chloride.
3. The use of sulfuric acid, and then leaching with water, forming lithium sulfate monohydrate. This is treated with sodium carbonate to form lithium carbonate, and then hydrochloric acid to form lithium chloride.
The lithium chloride obtained from any of the three methods undergoes an oxidation-reduction reaction in an electrolytic cell, to separate the chloride ions from the lithium ions. The chloride ions are oxidized, and the lithium ions are reduced.
$2Cl^- - 2e^- \rightarrow Cl_2 \;\; \text{(oxidation)} \nonumber$
$Li^+ + e^- \rightarrow Li \;\; \text{(reduction)} \nonumber$
Extraction from Saltwater
Saltwater naturally contains lithium chloride, which must be extracted in the form of lithium carbonate, then it is re-treated, separated into its ions, and reduced in the same electrolytic process as in extraction from lithium ores. Only three saltwater lakes in the world are currently used for lithium extraction, in Nevada, Chile, and Argentina.
Saltwater is channeled into shallow ponds and over a period of a year or more, water evaporates out to leave behind various salts. Lime is used to remove the magnesium salt, so that the remaining solution contains a fairly concentrated amount of lithium chloride. The solution is then treated with sodium carbonate in order for usable lithium carbonate to precipitate out.
Problems
1. With which group of elements will lithium form compounds the most easily with?
2. What is the electron configuration of Li+?
3. What are some common uses of lithium?
4. For a lithium-ion battery containing LiCoO2, should the compound be placed in the anode or cathode?
5. Given that 7Li is 7.0160 amu and 6Li is 6.0151 amu, and their percent abundance is 92.58% and 7.42% respectively, what is the atomic mass of lithium?
Solutions
1. Group 17 Halogens (lithium forms strongly inic bonds with them, as halogens are highly electronegative and lithium has a free electron)
2. 1s2
3. Lithium-ion batteries, disposable lithium batteries, pyrotechnics, creation of strong metal alloys, etc.
4. Anode - lithium is oxidized (LiCoO2 → Li+ + CoO2)
5. 6.942 g/mol
Contributors and Attributions
• Katherine Szelong (UCD), Kevin Fan | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__1%3A_The_Alkali_Metals/Z003_Chemistry_of_Lithium_(Z3).txt |
Sodium is metallic element found in the first group of the periodic table. As the sixth most abundant element in the Earth's crust, sodium compounds are commonly found dissolved in the oceans, in minerals, and even in our bodies.
Contributors and Attributions
• Helen Min (University of California, Davis)
Z011 Chemistry of Sodium (Z11)
Sodium carbonate (also known as washing soda or soda ash), Na2CO3, is a sodium salt of carbonic acid and is a fairly strong, non-volatile base. It most commonly occurs as a crystaline heptahydrate which readily effloresces to form a white powder, the monohydrate. It has a cooling alkaline taste, and can be extracted from the ashes of many plants. It is produced artificially in large quantities from common salt.
Sodium Carbonate
Sodium is metallic element found in the first group of the periodic table. As the sixth most abundant element in the Earth's crust, sodium compounds are commonly found dissolved in the oceans, in minerals, and even in our bodies.
Contributors and Attributions
• Helen Min (University of California, Davis)
Z011 Chemistry of Sodium (Z11)
Sodium carbonate (also known as washing soda or soda ash), Na2CO3, is a sodium salt of carbonic acid and is a fairly strong, non-volatile base. It most commonly occurs as a crystaline heptahydrate which readily effloresces to form a white powder, the monohydrate. It has a cooling alkaline taste, and can be extracted from the ashes of many plants. It is produced artificially in large quantities from common salt.
Potassium
Potassium is a group 1 metal, abbreviated as K on the periodic table. In its pure form, potassium has a white-sliver color, but quickly oxidizes upon exposure to air and tarnishing in minutes if it is not stored under oil or grease. Potassium is essential to several aspects of plant, animal, and human life and is thus mined, manufactured, and consumed in huge quantities around the world.
The seventh most abundant element, potassium was discovered and isolated in 1807 by Sir Humphry Davy. Important compounds of potassium include potassium hydroxide (used in some drain cleaners), potassium superoxide, $KO_2$, which is used in respiratory equipment and potassium nitrate, used in fertilizers and pyrotechnics. Potassium, like sodium, melts below the boiling point of water (63 °C) and is less dense than water also. Like most of the alkali metals, potassium compounds impart a characteristic color to flames. In the case of the 19th element, the color is pale lavender. Like sodium ions, the presence of potassium ions in the body is essential for the correct function of many cells.
Table $1$: Basic Chemical and Physical Properties
Atomic Number 19
Atomic Mass 39.098 g/mol
Electronegativity 0.8
Density 0.862 g/cm3 (at 0o C) (floats on water)
Melting Point 63.65o C
Boiling Point 773.9o C
Atomic Radius 227 pm
Ionic Radius 0.133 (+1)
Isotopes 5
Electronic Shell [Ar] 4s1
1st Ionization Energy 418.8 kJ/mol
Electrode Potential -2.924
Hardness 0.5
Crystal Lattice body-centered cubic
Specific Heat 0.741 J/gK
Heat of Fusion 59.591 J/g
Heat of Vaporization 2075 kJ/g
Electron Configuration 1s22s22p63s23p64s1
Notable Reactions with Phosphorus
Potassium reacts so violently with water that it bursts into flame. The silvery white metal is very soft and reacts rapidly with the oxygen in air. Its chemical symbol is derived from the Latin word kalium which means "alkali". Its English name is from potash which is the common name for a compound containing it.
Table $1$: Key reactions of Potassium
Reactant Reaction Product
H2 begins slowly at ca 200°C; rapid above 300°C KH
O2 begins slowly with solid; fairly rapid with liquid K2O, K2O2, KO2
H2O extremely vigorous and frequently results in hydrogen–air explosions KOH, H2
C(graphite) 150–400°C KC4, KC8, KC24
CO forms unstable carbonyls (KCO)
NH3 dissolves as K; iron, nickel, and other metals catalyze in gas and liquid phase KNH2
S molten state in liquid ammonia K2S, K2S2, K2S4
F2, Cl2, Br2 violent to explosive KF, KCl, KBr
I2 ignition KI
CO2 occurs readily, but is sometimes explosive CO, C, K2CO3
Potassium in the Environment
Potassium has a 2.6% abundance by mass in the earth's crust and is found mostly in mineral form as part of feldspars (groups of minerals) and clays. Potassium easily leaches out of these minerals over time and thus has a relatively high concentration in sea water as well (0.75g/L). Today, most of the world's potassium is mined in Canada, the U.S., and Chile but was originally monopolized by Germany.
Potassium and Living Organisms
Plants, animals and humans all depend on potassium for survival and good health. The element is part of many bodily fluids and assists related functions of the human body. Most notably, potassium aids nerve functions and is found in several cell types (including skeletal cells, smooth muscle cells, endocrine cells, cardiac cells, and central neurons). Plants depend on potassium for healthy growth. Potassium found in animal excretions and dead plants easily binds to clay in the soil they fall on and is thus utilized by plants. The element helps maintain osmotic pressure and cell size and plays a role in photosynthesis and energy production.
Applications
95% of manufactured potassium is used in fertilizers and the rest is used to produce specific compounds of potassium, such as potassium hydroxide ($KOH$), which can then be turned into potassium carbonate ($K_2CO_3$). Potassium carbonate is used in glass manufacturing and potassium hydroxide is found in liquid soaps and detergents. Potassium chloride is used in many pharmaceuticals and other salts of potassium are used in baking, photography, tanning leather, and iodized salt. In these cases, potassium is utilized for its negative anion.
Potassium can be obtained through various known reactions, all of which require heat treatment:
$K_2CO_3+2C \overset{\Delta}{\longrightarrow} 3CO+2K \label{1}$
$2KCl+CaC_2 \overset{\Delta}{\longrightarrow} CaCl_2+2C+K \label{2}$
$2KN_3 \overset{\Delta}{\longrightarrow} 3N_2+2K \label{3}$
Due to expenses, these processes are not commercially adaptable. Therefore the element is commonly obtained through reduction at elevated heats (i.e., pyrometallurgy). Sodium is often combined with $KCl$, $KOH$, or $K_2CO_3$ to produce potassium sodium alloys and in the 1950's the Mine Safety Appliances Company developed a reduction process that yields high purity potassium:
$KCl+Na \overset{\Delta}{\longrightarrow} K+NaCl \label{4}$
The reaction is heated in a special device equipped with a furnace, heat-exchanger tubes, a fractionating column, a $KCl$ feed, a waste removal system, and a vapor condensing system. Because the reaction attains equilibrium quickly, potassium can be removed continuously as a product in order to shift equilibrium to the right and produce even more potassium in its place.
Alloys of potassium include $NaK$ (Sodium) and $KLi$ (Lithium). Both of these alloys produce metals of low vapor pressure and melting points.
Problems
1. Why is Potassium never found pure in nature?
2. Write out the chemical reaction between potassium and water.
3. Name 3 uses of potassium.
4. Where is Potassium on the periodic table. What are a few things you can deduce just from this location?
5. Name a common alloy of Potassium. What are the beneficial properties of this alloy?
Answers
1. It is too reactive. Potassium is a very strong reducing agent because of its desire to achieve an inert gas electron configuration (like the other alkali metals). This means that it will easily give up electrons, giving it the ability to reduce numerous other elements.
2. $K+H_2O \rightarrow KOH + H_{2(g)}$: Like other group 1 metals, potassium reacts readily with water to generate hydrogen gas.
3. Potassium is used in glass making and is found in fertilizers and soaps.
4. Potassium is in group one, and is the 4th element down in it's column. This tells us that it is an alkali metal. It is very reactive, has a low density, and is a good reducing agent.
5. Potassium can form an alloy with $Na$ that has a low vapor pressure and melting point.
Z019 Chemistry of Potassium (Z19)
Potassium carbonate, K2CO3, or potash(1), is highly soluble in water, forming an alkaline solution. The name potash gave the chemical element potassium its English and French name. K2CO3 was called potash since it was extracted from wood ash being leached out by water in a pot.
It is made by passing CO2 into a 50% potassium hydroxide solution:
$2 KOH + CO_2 \rightarrow K_2CO_3 + H_2O \nonumber$
Potassium carbonate is used in the production of soap and glass and as a mild drying agent, especially for organic solvents. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__1%3A_The_Alkali_Metals/Z011_Chemistry_of_Sodium_%28Z11%29/Sodium_Carbo.txt |
Potassium is a group 1 metal, abbreviated as K on the periodic table. In its pure form, potassium has a white-sliver color, but quickly oxidizes upon exposure to air and tarnishing in minutes if it is not stored under oil or grease. Potassium is essential to several aspects of plant, animal, and human life and is thus mined, manufactured, and consumed in huge quantities around the world.
The seventh most abundant element, potassium was discovered and isolated in 1807 by Sir Humphry Davy. Important compounds of potassium include potassium hydroxide (used in some drain cleaners), potassium superoxide, $KO_2$, which is used in respiratory equipment and potassium nitrate, used in fertilizers and pyrotechnics. Potassium, like sodium, melts below the boiling point of water (63 °C) and is less dense than water also. Like most of the alkali metals, potassium compounds impart a characteristic color to flames. In the case of the 19th element, the color is pale lavender. Like sodium ions, the presence of potassium ions in the body is essential for the correct function of many cells.
Table $1$: Basic Chemical and Physical Properties
Atomic Number 19
Atomic Mass 39.098 g/mol
Electronegativity 0.8
Density 0.862 g/cm3 (at 0o C) (floats on water)
Melting Point 63.65o C
Boiling Point 773.9o C
Atomic Radius 227 pm
Ionic Radius 0.133 (+1)
Isotopes 5
Electronic Shell [Ar] 4s1
1st Ionization Energy 418.8 kJ/mol
Electrode Potential -2.924
Hardness 0.5
Crystal Lattice body-centered cubic
Specific Heat 0.741 J/gK
Heat of Fusion 59.591 J/g
Heat of Vaporization 2075 kJ/g
Electron Configuration 1s22s22p63s23p64s1
Notable Reactions with Phosphorus
Potassium reacts so violently with water that it bursts into flame. The silvery white metal is very soft and reacts rapidly with the oxygen in air. Its chemical symbol is derived from the Latin word kalium which means "alkali". Its English name is from potash which is the common name for a compound containing it.
Table $1$: Key reactions of Potassium
Reactant Reaction Product
H2 begins slowly at ca 200°C; rapid above 300°C KH
O2 begins slowly with solid; fairly rapid with liquid K2O, K2O2, KO2
H2O extremely vigorous and frequently results in hydrogen–air explosions KOH, H2
C(graphite) 150–400°C KC4, KC8, KC24
CO forms unstable carbonyls (KCO)
NH3 dissolves as K; iron, nickel, and other metals catalyze in gas and liquid phase KNH2
S molten state in liquid ammonia K2S, K2S2, K2S4
F2, Cl2, Br2 violent to explosive KF, KCl, KBr
I2 ignition KI
CO2 occurs readily, but is sometimes explosive CO, C, K2CO3
Potassium in the Environment
Potassium has a 2.6% abundance by mass in the earth's crust and is found mostly in mineral form as part of feldspars (groups of minerals) and clays. Potassium easily leaches out of these minerals over time and thus has a relatively high concentration in sea water as well (0.75g/L). Today, most of the world's potassium is mined in Canada, the U.S., and Chile but was originally monopolized by Germany.
Potassium and Living Organisms
Plants, animals and humans all depend on potassium for survival and good health. The element is part of many bodily fluids and assists related functions of the human body. Most notably, potassium aids nerve functions and is found in several cell types (including skeletal cells, smooth muscle cells, endocrine cells, cardiac cells, and central neurons). Plants depend on potassium for healthy growth. Potassium found in animal excretions and dead plants easily binds to clay in the soil they fall on and is thus utilized by plants. The element helps maintain osmotic pressure and cell size and plays a role in photosynthesis and energy production.
Applications
95% of manufactured potassium is used in fertilizers and the rest is used to produce specific compounds of potassium, such as potassium hydroxide ($KOH$), which can then be turned into potassium carbonate ($K_2CO_3$). Potassium carbonate is used in glass manufacturing and potassium hydroxide is found in liquid soaps and detergents. Potassium chloride is used in many pharmaceuticals and other salts of potassium are used in baking, photography, tanning leather, and iodized salt. In these cases, potassium is utilized for its negative anion.
Potassium can be obtained through various known reactions, all of which require heat treatment:
$K_2CO_3+2C \overset{\Delta}{\longrightarrow} 3CO+2K \label{1}$
$2KCl+CaC_2 \overset{\Delta}{\longrightarrow} CaCl_2+2C+K \label{2}$
$2KN_3 \overset{\Delta}{\longrightarrow} 3N_2+2K \label{3}$
Due to expenses, these processes are not commercially adaptable. Therefore the element is commonly obtained through reduction at elevated heats (i.e., pyrometallurgy). Sodium is often combined with $KCl$, $KOH$, or $K_2CO_3$ to produce potassium sodium alloys and in the 1950's the Mine Safety Appliances Company developed a reduction process that yields high purity potassium:
$KCl+Na \overset{\Delta}{\longrightarrow} K+NaCl \label{4}$
The reaction is heated in a special device equipped with a furnace, heat-exchanger tubes, a fractionating column, a $KCl$ feed, a waste removal system, and a vapor condensing system. Because the reaction attains equilibrium quickly, potassium can be removed continuously as a product in order to shift equilibrium to the right and produce even more potassium in its place.
Alloys of potassium include $NaK$ (Sodium) and $KLi$ (Lithium). Both of these alloys produce metals of low vapor pressure and melting points.
Problems
1. Why is Potassium never found pure in nature?
2. Write out the chemical reaction between potassium and water.
3. Name 3 uses of potassium.
4. Where is Potassium on the periodic table. What are a few things you can deduce just from this location?
5. Name a common alloy of Potassium. What are the beneficial properties of this alloy?
Answers
1. It is too reactive. Potassium is a very strong reducing agent because of its desire to achieve an inert gas electron configuration (like the other alkali metals). This means that it will easily give up electrons, giving it the ability to reduce numerous other elements.
2. $K+H_2O \rightarrow KOH + H_{2(g)}$: Like other group 1 metals, potassium reacts readily with water to generate hydrogen gas.
3. Potassium is used in glass making and is found in fertilizers and soaps.
4. Potassium is in group one, and is the 4th element down in it's column. This tells us that it is an alkali metal. It is very reactive, has a low density, and is a good reducing agent.
5. Potassium can form an alloy with $Na$ that has a low vapor pressure and melting point.
Z019 Chemistry of Potassium (Z19)
Potassium carbonate, K2CO3, or potash(1), is highly soluble in water, forming an alkaline solution. The name potash gave the chemical element potassium its English and French name. K2CO3 was called potash since it was extracted from wood ash being leached out by water in a pot.
It is made by passing CO2 into a 50% potassium hydroxide solution:
$2 KOH + CO_2 \rightarrow K_2CO_3 + H_2O \nonumber$
Potassium carbonate is used in the production of soap and glass and as a mild drying agent, especially for organic solvents. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__1%3A_The_Alkali_Metals/Z019_Chemistry_of_Potassium_(Z19)/Potassium_Car.txt |
Rubidium (Latin: rubidius = red) is similar in physical and chemical characteristics to potassium, but much more reactive. It is the seventeenth most abundant element and was discovered by its red spectral emission in 1861 by Bunsen and Kirchhoff. Its melting point is so low you could melt it in your hand if you had a fever (39°C). But that would not be a good idea because it would react violently with the moisture in your skin.
Rubidium was once thought to be quite rare but recent discoveries of large deposits indicate that there is plenty to use. However at present it finds only limited application in the manufacture of cathode ray tubes.
Contributors and Attributions
Stephen R. Marsden
Z037 Chemistry of Rubidium (Z37)
Rubidium (Latin: rubidius = red) is similar in physical and chemical characteristics to potassium, but much more reactive. It is the seventeenth most abundant element and was discovered by its red spectral emission in 1861 by Bunsen and Kirchhoff. Its melting point is so low you could melt it in your hand if you had a fever (39°C). But that would not be a good idea because it would react violently with the moisture in your skin.
Rubidium was once thought to be quite rare but recent discoveries of large deposits indicate that there is plenty to use. However at present it finds only limited application in the manufacture of cathode ray tubes.
Contributors and Attributions
Stephen R. Marsden
Z055 Chemistry of Cesium (Z55)
Cesium is a bright silvery metal which is a liquid in a warm room (28oC). Its name is from the Latin caesius which is a description of a sky blue spectral emission by which it was discovered in 1860 by Bunsen and Kirchhoff.
Cesium is so reactive that it will even explode on contact with ice! It has been used as a "getter" in the manufacture of vacuum tubes (i.e., it helps remove trace quantities of remaining gases). An isotope of cesium is used in the atomic clocks.
Contributors and Attributions
• Stephen R. Marsden
Z055 Chemistry of Cesium (Z55)
Cesium is a bright silvery metal which is a liquid in a warm room (28oC). Its name is from the Latin caesius which is a description of a sky blue spectral emission by which it was discovered in 1860 by Bunsen and Kirchhoff.
Cesium is so reactive that it will even explode on contact with ice! It has been used as a "getter" in the manufacture of vacuum tubes (i.e., it helps remove trace quantities of remaining gases). An isotope of cesium is used in the atomic clocks.
Contributors and Attributions
• Stephen R. Marsden
Z087 Chemistry of Francium (Z87)
Francium is the last of the known alkali metals and does not occur to any significant extent in nature. All known isotopes are radioactive and have short half-lives (22 minutes is the longest).
The existence of Francium was predicted by Dmitri Mendeleev in the 1870's and he presumed it would have chemical and physical properties similar to cesium. That may well be, but not enough francium has been isolated to test.
Numerous historical claims to the discovery of element 87 were made resulting in the names russium, virginium, and moldavium. However, the confirmed discovery is credited to Marguerite Perey who was an assistant to Marie Curie at the Radium Institute in Paris. She named the element after her native country.
Contributors and Attributions
Stephen R. Marsden
Z087 Chemistry of Francium (Z87)
Francium is the last of the known alkali metals and does not occur to any significant extent in nature. All known isotopes are radioactive and have short half-lives (22 minutes is the longest).
The existence of Francium was predicted by Dmitri Mendeleev in the 1870's and he presumed it would have chemical and physical properties similar to cesium. That may well be, but not enough francium has been isolated to test.
Numerous historical claims to the discovery of element 87 were made resulting in the names russium, virginium, and moldavium. However, the confirmed discovery is credited to Marguerite Perey who was an assistant to Marie Curie at the Radium Institute in Paris. She named the element after her native country.
Contributors and Attributions
Stephen R. Marsden | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__1%3A_The_Alkali_Metals/Z037_Chemistry_of_Rubidium_%28Z37%29.txt |
The Group 2 alkaline earth metals include Beryllium, Magnesium, Calcium, Barium, Strontium and Radium and are soft, silver metals that are less metallic in character than the Group 1 Alkali Metals. Although many characteristics are common throughout the group, the heavier metals such as Ca, Sr, Ba, and Ra are almost as reactive as the Group 1 Alkali Metals. All the elements in Group 2 have two electrons in their valence shells, giving them an oxidation state of +2.
• Group 2: Chemical Properties of Alkali Earth Metals
Covers the elements beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr) and barium (Ba). Includes trends in atomic and physical properties, trends in reactivity, the solubility patterns in the hydroxides and sulfates, trends in the thermal decomposition of the nitrates and carbonates, and some of the atypical properties of beryllium.
• Group 2: Physical Properties of Alkali Earth Metals
This page explores the trends in some atomic and physical properties of the Group 2 elements: beryllium, magnesium, calcium, strontium and barium. Sections below cover the trends in atomic radius, first ionization energy, electronegativity, and physical properties.
• Chemistry of Beryllium (Z=4)
The name Beryllium comes from the Greek beryllos which is the name for the gemstone beryl. The element is a high-melting, silver-white metal which is the first member of the alkaline earth metals. It is not abundant in the environment and occurs mainly in the mineral beryl with aluminum and silicon.
• Chemistry of Magnesium (Z=12)
Magnesium is a group two element and is the eighth most common element in the earth's crust. Magnesium is light, silvery-white, and tough. Like aluminum, it forms a thin layer around itself to help prevent itself from rusting when exposed to air. Fine particles of magnesium can also catch on fire when exposed to air.
• Chemistry of Calcium (Z=20)
Calcium is the 20th element in the periodic table. It is a group 2 metal, also known as an alkaline-earth metal, and no populated d-orbital electrons. Calcium is the fifth most abundant element by mass (3.4%) in both the Earth's crust and in seawater. All living organisms (in fact, even dead ones) have and need calcium for survival.
• Chemistry of Strontium (Z=38)
Strontium is a group 2 element that does not occur as a free element due to its extreme reactivity with oxygen and water. It occurs naturally only in compounds with other elements such as strontianite. It is softer than calcium and decomposes water more vigorously. It has a silver appearance but then turns yellow with the formation of oxide. Strontium is named after the Scottish village on Strontian.
• Chemistry of Barium (Z=56)
Barium is a soft, silvery white metal, and has a melting point of 1000 K. Because of its reaction to air, barium cannot be found in nature in its pure form but can be extracted from the mineral barite.
• Chemistry of Radium (Z=88)
Radium takes its name from the Latin word radius or ray. All isotopes of radium are radioactive and many exhibit luminescence, reacting readily with oxygen and water. The metal was discovered and isolated in 1911 by Marie Curie.
Group 2 Elements: The Alkaline Earth Metals
Covers the elements beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr) and barium (Ba). Includes trends in atomic and physical properties, trends in reactivity, the solubility patterns in the hydroxides and sulfates, trends in the thermal decomposition of the nitrates and carbonates, and some of the atypical properties of beryllium.
2 Group 2: Physical Properti
This page explores the trends in some atomic and physical properties of the Group 2 elements: beryllium, magnesium, calcium, strontium and barium. Sections below cover the trends in atomic radius, first ionization energy, electronegativity, and physical properties.
Contributors and Attributions
Jim Clark (Chemguide.co.uk) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements%3A_The_Alkaline_Earth_Metals/1Group_2%3A_Chemical_Reactions.txt |
Beryllium is an element found in nature and is combined with other elements in minerals, including beryl and chrysoberyl. In its purest form, beryllium is a steel-gray and lightweight alkaline earth metal.
Introduction
Due to its physical properties, beryllium is useful as a hardening agent in alloys, making aerospace material, and used as a filter for radiation. Beryllium is not used for commercial use due to the harmful effects when it is inhaled through dust particles, causing berylliosis (a corrosive disease typically in the lungs). Beryllium is a rare element on Earth and in the universe and is not found to be necessary or helpful for plants or animals.
General Properties of Beryllium
Symbol Be
Color Steel Gray
Atomic Number 4
Category Alkali Earth Metal
Atomic Weight 9.012182
Group/Period 2/2
Electron Configuration 1s2 2s2
Valence Electrons 2
Phase (room temperature) Solid
Melting Point 1560 K, 1287°C
Boiling Point 2742 K, 2469 °C
Atomic Radius 105 pm
History
In 1798, N.L. Vauquelin discovered the element in beryl and emerald.Beryllium was first isolated in 1828 by Wöhler. It is used in specialty alloys such as spring metal in which it increases toughness. It was once known as glucinium because of the sweet taste of its compounds (which, alas, are toxic). During World War I, larger amounts of beryllium were made, but it was not until the early 1930s when mass quantities of beryllium were made. During World War II, the Brush Beryllium Company was popular as the demand for beryllium copper alloys and fluorescent material in lamps grew.
Characteristics
Physical
Despite how light it is for a metal, beryllium has a very high melting point. It also has a high modulus of elasticity that is 50% greater than steel, along with a low density giving it a fast sound conduction speed. At STP, beryllium resists oxidation and resists corrosion in the air.
Isotopes
Although beryllium has many isotopes, only 9Be is stable, classifying it as a monoisotopic element. 10Be is produced in the atmosphere when cosmic ray spallation of oxygen and nitrogen occurs. 10Be resides on top of soil, and has a long half-life which allows it to survive a long time before turning into 10B. This makes 10Be useful for examining soil and solar activity because solar activity is inversely correlated with 10Be production. Besides 10Be, many of beryllium's isotopes, especially 13Be, have very short half-lives.
Chemical
Beryllium is a steel-gray metal that tarnishes slowly in the air due to oxide forming around it. This thin layer of oxide allows beryllium to scratch glass. Common compounds containing beryllium are emerald and aquamarine. Due to its light, stiff, and stable structure, beryllium alloys are being used in industrial work more and more.
Like all elements in the 2nd group on the periodic table, beryllium has +2 oxidation state. With a small atomic radius, Be2+ has high polarization characteristics allowing it to form many covalent bonds. Beryllium forms an oxide layer making it not react with air or water even in extreme heat. However, when it is ignited, beryllium burns brightly making beryllium oxide and beryllium nitride. Beryllium dissolves easily in non-oxidizing acids, such as HCl, with the exception of nitric because it forms the oxide making it very similar to aluminum metal.
Beryllium combines with many non-metals to form binary compounds, such as beryllium oxide (BeO). BeO is a white solid that has a high melting point, making it useful in engines, radio equipment, and semiconductor devices.
Occurrence and Production
Beryllium is in approximately 100 of the 4000 known minerals, such as bertrandite, beryl, chrysoberyl, and phenakite. Beryllium is also present in precious gems such as aquamarine, bixbite, and emerald. Of the many beryllium minerals, only two are of commercial importance in the preparation of beryllium metal and its compounds. Bertrandite (Be4Si2O7(OH)2) contains less than 1% Be and is the main beryllium mineral mined in the U.S., while beryl (Be3Al2(SiO3)6 is mined in other countries and contains approximately 4% Be. In the U.S., beryllium is mainly mined at Gold Hill and Spor Mountain in Utah, and in Alaska on the Seward Peninsula.
Beryllium metal began commercial production in 1957, but did not live up to its expectation of expanding the industry. Beryllium is made by reducing beryllium fluoride with magnesium metal in the following equation:
$\ce{BeF2 + Mg -> MgF2 + Be} \nonumber$
Emerald is less common than diamond and more expensive than gold. Columbia produces the most emerald in the world, where the Muzo mine and eastern emerald belt are located.
Applications
Despite having problems of beryllium being brittle, pricey, and poisonous, it still has many valuable purposes. Its light weight, non-magnetic properties, and reluctance to spark is great for non-sparking tools. Beryllium is great for making aircraft and space ships due to its low density, high heat capacity, and high modulus of elasticity.
With a simple nuclei of just 4 protons and 5 neutrons, beryllium is great for tubes as it allows all radiation to pass through easily. On the contrary, beryllium atoms reflect neutrons making it great for reflectors, moderators, and control rods in research reactors. Beryllium oxide is a great electric insulator and heat conductor. It is transparent to microwaves making it useful in microwave communications systems. Beryllium oxide is also used in computers, lasers, and automotive ignition systems.
Toxicity
Beryllium is very toxic to people. The severity of toxicity depends upon how it enters the body, how long it remains in the body, how much enters the body, and how many times it enters the body. Inside the body, beryllium binds to phosphate-containing systems causing damage to ones health. Due to only small amounts of beryllium in the natural environment, there is no biological system creating a protection against this element, despite it being used more in today's industry.
If one is exposed to high levels of beryllium for an extended period of time, chronic beryllium disease (CBD) may develop. CBD symptoms include fatigue, weakness, difficulty in breathing, and chronic coughing. CBD develops when the body senses beryllium particles entering the body, which is typically inhaled into the lungs. The immune system responds by sending white blood cells to the organ, which later prevents the organ from performing its job efficiently.
Problems
1. What is the common chemical equation for making beryllium?
2. How many stable isotopes does beryllium have and what are they?
3. Why is beryllium able to scratch through glass?
4. Why is beryllium harmful to ones body? What disease can it lead to?
5. Name 2 industrial/commercial uses for beryllium?
Answers (Highlight over the section)
1. BeF2+Mg --> MgF2+Be
2. 9Be is the only stable beryllium isotope, making it a monoisotopic element.
3. Beryllium is able to scratch glass because a thin layer of oxide forms around beryllium when it is exposed to the air.
4. Beryllium is harmful to the body because it bonds to phosphate-containing systems and interferes with the organs job working productively. This can lead to illnesses such as chronic beryllium disease (CBD).
5. Beryllium is used in X-rays to allow radiation to enter and block other rays and for building aircraft and space parts due to its high heat capacity and light weight.
Contributors and Attributions
• Jean Kim (UC Davis) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements%3A_The_Alkaline_Earth_Metals/Z004_Chemistry_of_Beryllium_%2.txt |
Beryllium is an element found in nature and is combined with other elements in minerals, including beryl and chrysoberyl. In its purest form, beryllium is a steel-gray and lightweight alkaline earth metal.
Introduction
Due to its physical properties, beryllium is useful as a hardening agent in alloys, making aerospace material, and used as a filter for radiation. Beryllium is not used for commercial use due to the harmful effects when it is inhaled through dust particles, causing berylliosis (a corrosive disease typically in the lungs). Beryllium is a rare element on Earth and in the universe and is not found to be necessary or helpful for plants or animals.
General Properties of Beryllium
Symbol Be
Color Steel Gray
Atomic Number 4
Category Alkali Earth Metal
Atomic Weight 9.012182
Group/Period 2/2
Electron Configuration 1s2 2s2
Valence Electrons 2
Phase (room temperature) Solid
Melting Point 1560 K, 1287°C
Boiling Point 2742 K, 2469 °C
Atomic Radius 105 pm
History
In 1798, N.L. Vauquelin discovered the element in beryl and emerald.Beryllium was first isolated in 1828 by Wöhler. It is used in specialty alloys such as spring metal in which it increases toughness. It was once known as glucinium because of the sweet taste of its compounds (which, alas, are toxic). During World War I, larger amounts of beryllium were made, but it was not until the early 1930s when mass quantities of beryllium were made. During World War II, the Brush Beryllium Company was popular as the demand for beryllium copper alloys and fluorescent material in lamps grew.
Characteristics
Physical
Despite how light it is for a metal, beryllium has a very high melting point. It also has a high modulus of elasticity that is 50% greater than steel, along with a low density giving it a fast sound conduction speed. At STP, beryllium resists oxidation and resists corrosion in the air.
Isotopes
Although beryllium has many isotopes, only 9Be is stable, classifying it as a monoisotopic element. 10Be is produced in the atmosphere when cosmic ray spallation of oxygen and nitrogen occurs. 10Be resides on top of soil, and has a long half-life which allows it to survive a long time before turning into 10B. This makes 10Be useful for examining soil and solar activity because solar activity is inversely correlated with 10Be production. Besides 10Be, many of beryllium's isotopes, especially 13Be, have very short half-lives.
Chemical
Beryllium is a steel-gray metal that tarnishes slowly in the air due to oxide forming around it. This thin layer of oxide allows beryllium to scratch glass. Common compounds containing beryllium are emerald and aquamarine. Due to its light, stiff, and stable structure, beryllium alloys are being used in industrial work more and more.
Like all elements in the 2nd group on the periodic table, beryllium has +2 oxidation state. With a small atomic radius, Be2+ has high polarization characteristics allowing it to form many covalent bonds. Beryllium forms an oxide layer making it not react with air or water even in extreme heat. However, when it is ignited, beryllium burns brightly making beryllium oxide and beryllium nitride. Beryllium dissolves easily in non-oxidizing acids, such as HCl, with the exception of nitric because it forms the oxide making it very similar to aluminum metal.
Beryllium combines with many non-metals to form binary compounds, such as beryllium oxide (BeO). BeO is a white solid that has a high melting point, making it useful in engines, radio equipment, and semiconductor devices.
Occurrence and Production
Beryllium is in approximately 100 of the 4000 known minerals, such as bertrandite, beryl, chrysoberyl, and phenakite. Beryllium is also present in precious gems such as aquamarine, bixbite, and emerald. Of the many beryllium minerals, only two are of commercial importance in the preparation of beryllium metal and its compounds. Bertrandite (Be4Si2O7(OH)2) contains less than 1% Be and is the main beryllium mineral mined in the U.S., while beryl (Be3Al2(SiO3)6 is mined in other countries and contains approximately 4% Be. In the U.S., beryllium is mainly mined at Gold Hill and Spor Mountain in Utah, and in Alaska on the Seward Peninsula.
Beryllium metal began commercial production in 1957, but did not live up to its expectation of expanding the industry. Beryllium is made by reducing beryllium fluoride with magnesium metal in the following equation:
$\ce{BeF2 + Mg -> MgF2 + Be} \nonumber$
Emerald is less common than diamond and more expensive than gold. Columbia produces the most emerald in the world, where the Muzo mine and eastern emerald belt are located.
Applications
Despite having problems of beryllium being brittle, pricey, and poisonous, it still has many valuable purposes. Its light weight, non-magnetic properties, and reluctance to spark is great for non-sparking tools. Beryllium is great for making aircraft and space ships due to its low density, high heat capacity, and high modulus of elasticity.
With a simple nuclei of just 4 protons and 5 neutrons, beryllium is great for tubes as it allows all radiation to pass through easily. On the contrary, beryllium atoms reflect neutrons making it great for reflectors, moderators, and control rods in research reactors. Beryllium oxide is a great electric insulator and heat conductor. It is transparent to microwaves making it useful in microwave communications systems. Beryllium oxide is also used in computers, lasers, and automotive ignition systems.
Toxicity
Beryllium is very toxic to people. The severity of toxicity depends upon how it enters the body, how long it remains in the body, how much enters the body, and how many times it enters the body. Inside the body, beryllium binds to phosphate-containing systems causing damage to ones health. Due to only small amounts of beryllium in the natural environment, there is no biological system creating a protection against this element, despite it being used more in today's industry.
If one is exposed to high levels of beryllium for an extended period of time, chronic beryllium disease (CBD) may develop. CBD symptoms include fatigue, weakness, difficulty in breathing, and chronic coughing. CBD develops when the body senses beryllium particles entering the body, which is typically inhaled into the lungs. The immune system responds by sending white blood cells to the organ, which later prevents the organ from performing its job efficiently.
Problems
1. What is the common chemical equation for making beryllium?
2. How many stable isotopes does beryllium have and what are they?
3. Why is beryllium able to scratch through glass?
4. Why is beryllium harmful to ones body? What disease can it lead to?
5. Name 2 industrial/commercial uses for beryllium?
Answers (Highlight over the section)
1. BeF2+Mg --> MgF2+Be
2. 9Be is the only stable beryllium isotope, making it a monoisotopic element.
3. Beryllium is able to scratch glass because a thin layer of oxide forms around beryllium when it is exposed to the air.
4. Beryllium is harmful to the body because it bonds to phosphate-containing systems and interferes with the organs job working productively. This can lead to illnesses such as chronic beryllium disease (CBD).
5. Beryllium is used in X-rays to allow radiation to enter and block other rays and for building aircraft and space parts due to its high heat capacity and light weight.
Contributors and Attributions
• Jean Kim (UC Davis) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements%3A_The_Alkaline_Earth_Metals/Z004_Chemistry_of_Beryllium_(Z.txt |
Magnesium is a group two element and is the eighth most common element in the earth's crust. Magnesium is light, silvery-white, and tough. Like aluminum, it forms a thin layer around itself to help prevent itself from rusting when exposed to air. Fine particles of magnesium can also catch on fire when exposed to air. Magnesium is essential in nutrition for animals and plants. It is also used as an alloy to combine with other metals to make them lighter and easier to weld, for purposes in the aerospace industry along with other industries. It is also used in medicine, in the forms of magnesium hydroxides, sulfates, chlorides, and citrates.
General Information
• Symbol: Mg
• Atomic Number: 12
• Atomic/Molar Mass: 24.31
• Melting Point: 648.8°C, 921.8K
• Boiling Point: 1090°C, 1363K
• Density:1.738 g/cc
• Oxidation states: +2
• Electron Shell Configuration: [Ne]3s2
Characteristics
Magnesium takes it name from magnesite ore, named for the district Magnesia in Thessaly, Greece. Magnesium is a strong metal that is light and silvery-white. Recognized as a element as far back as 1775, it was first isolated in pure form by Davy in 1805. Magnesium has the ability to tarnish, which creates an oxide layer around itself to prevent it from rusting. It also has the ability to react with water at room temperature. When exposed to water, bubbles form around the metal. Increasing the temperature speeds up this reaction.
Magnesium Fire
One property of magnesium is high flammability. Like many other things, magnesium is more flammable when it has a higher surface area to volume ratio. An example of surface area to volume ratio is seen in the lighting of fire wood. It is easier to light kindling and smaller branches than a whole log. This property of magnesium is used in war, photography, and in light bulbs. Magnesium is used in war for incendiary bombs, flares, and tracer bullets. When these weapons are used, they ignite immediately and cause fires. The only way to extinguish a magnesium fire is to cover it with sand. Water does not extinguish the fire as water reacts with the hot magnesium and releases even more hydrogen.
Applications
Magnesium is one of the lightest metals, and when used as an alloy, it is commonly used in the automotive and aeronautical industries. The use of magnesium has increased and peaked in 1943. One reason the use of magnesium has increased is that it is useful in alloys. Alloys with magnesium are able to be welded better and are lighter, which is ideal for metals used in the production of planes and other military goods.
Another characteristic of magnesium is that it aids in the digestive process. Magnesium is commonly used in milk of magnesia and Epsom salts. These forms of magnesium can range from magnesium hydroxide, magnesium sulfate, magnesium chloride, and magnesium citrate. Magnesium not only aids in humans and animals, but also in plants. It is used to convert the sun's lights into energy for the plant in a process known as photosynthesis. The main component of this process is chlorophyll. This is a pigment molecule that is composed of magnesium. Without magnesium, photosynthesis as we know it would not be possible.
Isotopes
Magnesium has three stable isotopes, Mg-24, Mg-25, Mg-26. The most common isotope is Mg-24, which is 79% of all Mg found on Earth. Mg25 and Mg26 are used to study the absorption and metabolism of magnesium in the human body. They are also used to study heart disease.
Magnesium not only has stable isotopes, but also has radioactive isotopes, which are isotopes that have an unstable nuclei. These isotopes are Mg--22, Mg23, Mg-27, Mg-28, and Mg-29. Mg-28 was commonly used in nuclear sites for scientific experiments from the 1950s to 1970s.
Reactions With
Water: When exposed to steam, magnesium changes from magnesium to magnesium oxide and hydrogen.
$Mg(s) +H_2O(g) \rightarrow MgO(s) + H_2(g) \nonumber$
When exposed to cold water, the reaction is a bit different. The reaction does not stop because the magnesium hydroxide gets insoluble in water.
$Mg(s) +2H_2O(g) \rightarrow Mg(OH)_2(s) + H_2(g) \nonumber$
Oxygen: When exposed to oxygen, magnesium turns into magnesium oxide.
$2Mg(s) +O_2(g) \rightarrow 2MgO(s) \nonumber$
Hydrogen: When exposed to hydrogen, magnesium turns into magnesium hydride.
$Mg(s) + H_2(g) \rightarrow MgH_2(s) \nonumber$
Nitrogen: When reacted with nitrogen, magnesium turns into magnesium nitride.
$3Mg(s) + N_2(g) \rightarrow Mg_3N_2(s) \nonumber$
Halogens: When reacted with a halogen, magnesium is very reactive. An example will be with chloride. When reacted with chloride, the product is magnesium(II) chloride.
$Mg(s) + Cl_2(g) \rightarrow MgCl_2(s) \nonumber$
Acids: When reacted with acids, magnesium dissolves and forms solutions that have both the Mg(II) ion and hydrogen gas.
$Mg(s) + 2HCl(aq) \rightarrow Mg^{2+}(aq) + 2Cl^-(aq) + H_2(g) \nonumber$
Bases: When reacted with bases, magnesium react.
Problems
1. Why does magnesium not rust?
2. Why is it not possible to extinguish magnesium with water?
3. What isotopes are used for heart studies?
4. Why is magnesium commonly used to create automobiles and planes?
5. What is the common oxidation state for magnesium?
Solutions
1. Because it is able to tarnish the ability to create a thin oxide layer around the metal.
2. Because water releases hydrogen when exposed to hot magnesium.
3. Mg25 and Mg26
4. Because it can be combined with other metals to make them lighter and easier to weld.
5. +2
Contributors and Attributions
• Ryan Kim (UC Davis)
• Avneet Kahlon (UC Davis) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements%3A_The_Alkaline_Earth_Metals/Z012_Chemistry_of_Magnesium_%2.txt |
Magnesium is a group two element and is the eighth most common element in the earth's crust. Magnesium is light, silvery-white, and tough. Like aluminum, it forms a thin layer around itself to help prevent itself from rusting when exposed to air. Fine particles of magnesium can also catch on fire when exposed to air. Magnesium is essential in nutrition for animals and plants. It is also used as an alloy to combine with other metals to make them lighter and easier to weld, for purposes in the aerospace industry along with other industries. It is also used in medicine, in the forms of magnesium hydroxides, sulfates, chlorides, and citrates.
General Information
• Symbol: Mg
• Atomic Number: 12
• Atomic/Molar Mass: 24.31
• Melting Point: 648.8°C, 921.8K
• Boiling Point: 1090°C, 1363K
• Density:1.738 g/cc
• Oxidation states: +2
• Electron Shell Configuration: [Ne]3s2
Characteristics
Magnesium takes it name from magnesite ore, named for the district Magnesia in Thessaly, Greece. Magnesium is a strong metal that is light and silvery-white. Recognized as a element as far back as 1775, it was first isolated in pure form by Davy in 1805. Magnesium has the ability to tarnish, which creates an oxide layer around itself to prevent it from rusting. It also has the ability to react with water at room temperature. When exposed to water, bubbles form around the metal. Increasing the temperature speeds up this reaction.
Magnesium Fire
One property of magnesium is high flammability. Like many other things, magnesium is more flammable when it has a higher surface area to volume ratio. An example of surface area to volume ratio is seen in the lighting of fire wood. It is easier to light kindling and smaller branches than a whole log. This property of magnesium is used in war, photography, and in light bulbs. Magnesium is used in war for incendiary bombs, flares, and tracer bullets. When these weapons are used, they ignite immediately and cause fires. The only way to extinguish a magnesium fire is to cover it with sand. Water does not extinguish the fire as water reacts with the hot magnesium and releases even more hydrogen.
Applications
Magnesium is one of the lightest metals, and when used as an alloy, it is commonly used in the automotive and aeronautical industries. The use of magnesium has increased and peaked in 1943. One reason the use of magnesium has increased is that it is useful in alloys. Alloys with magnesium are able to be welded better and are lighter, which is ideal for metals used in the production of planes and other military goods.
Another characteristic of magnesium is that it aids in the digestive process. Magnesium is commonly used in milk of magnesia and Epsom salts. These forms of magnesium can range from magnesium hydroxide, magnesium sulfate, magnesium chloride, and magnesium citrate. Magnesium not only aids in humans and animals, but also in plants. It is used to convert the sun's lights into energy for the plant in a process known as photosynthesis. The main component of this process is chlorophyll. This is a pigment molecule that is composed of magnesium. Without magnesium, photosynthesis as we know it would not be possible.
Isotopes
Magnesium has three stable isotopes, Mg-24, Mg-25, Mg-26. The most common isotope is Mg-24, which is 79% of all Mg found on Earth. Mg25 and Mg26 are used to study the absorption and metabolism of magnesium in the human body. They are also used to study heart disease.
Magnesium not only has stable isotopes, but also has radioactive isotopes, which are isotopes that have an unstable nuclei. These isotopes are Mg--22, Mg23, Mg-27, Mg-28, and Mg-29. Mg-28 was commonly used in nuclear sites for scientific experiments from the 1950s to 1970s.
Reactions With
Water: When exposed to steam, magnesium changes from magnesium to magnesium oxide and hydrogen.
$Mg(s) +H_2O(g) \rightarrow MgO(s) + H_2(g) \nonumber$
When exposed to cold water, the reaction is a bit different. The reaction does not stop because the magnesium hydroxide gets insoluble in water.
$Mg(s) +2H_2O(g) \rightarrow Mg(OH)_2(s) + H_2(g) \nonumber$
Oxygen: When exposed to oxygen, magnesium turns into magnesium oxide.
$2Mg(s) +O_2(g) \rightarrow 2MgO(s) \nonumber$
Hydrogen: When exposed to hydrogen, magnesium turns into magnesium hydride.
$Mg(s) + H_2(g) \rightarrow MgH_2(s) \nonumber$
Nitrogen: When reacted with nitrogen, magnesium turns into magnesium nitride.
$3Mg(s) + N_2(g) \rightarrow Mg_3N_2(s) \nonumber$
Halogens: When reacted with a halogen, magnesium is very reactive. An example will be with chloride. When reacted with chloride, the product is magnesium(II) chloride.
$Mg(s) + Cl_2(g) \rightarrow MgCl_2(s) \nonumber$
Acids: When reacted with acids, magnesium dissolves and forms solutions that have both the Mg(II) ion and hydrogen gas.
$Mg(s) + 2HCl(aq) \rightarrow Mg^{2+}(aq) + 2Cl^-(aq) + H_2(g) \nonumber$
Bases: When reacted with bases, magnesium react.
Problems
1. Why does magnesium not rust?
2. Why is it not possible to extinguish magnesium with water?
3. What isotopes are used for heart studies?
4. Why is magnesium commonly used to create automobiles and planes?
5. What is the common oxidation state for magnesium?
Solutions
1. Because it is able to tarnish the ability to create a thin oxide layer around the metal.
2. Because water releases hydrogen when exposed to hot magnesium.
3. Mg25 and Mg26
4. Because it can be combined with other metals to make them lighter and easier to weld.
5. +2
Contributors and Attributions
• Ryan Kim (UC Davis)
• Avneet Kahlon (UC Davis) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements%3A_The_Alkaline_Earth_Metals/Z012_Chemistry_of_Magnesium_(Z.txt |
Calcium is the 20th element in the periodic table. It is a group 2 metal, also known as an alkaline-earth metal, and no populated d-orbital electrons. Calcium is the fifth most abundant element by mass (3.4%) in both the Earth's crust and in seawater. All living organisms require calcium for survival. Calcium is a silver-gray metal which takes its name from the Latin word calx, which means lime. It is the fifth most abundant element in the earth's crust and is widely distributed as limestone (CaCO3), quicklime (CaO) and calcium fluoride.
General Properties of Calcium
Symbol Ca
Color dull gray or silver
Atomic Number 20
Category alkaline earth metal
Atomic Weight 40.078 g•mol−1
Group,Period,Block 2,4,s
Electron Configuration [Ar]4s2
Valence Electrons 2
Phase (room temperature) solid
Melting Point 1115 K, 842°C
Boiling Point 1757 K, 1484 °C
Atomic Radius 197 pm
Oxidation States 2
Density at room temp 1.55 g•cm−3
Electronegativity 1.00 (Pauling)
First ionization energy 589.8 kJ•mol−1
Number of stable isotopes 4 (2 more are fairly stable)
Electronic Structure
Fig. 1: Calcium Atom
Discovery and Properties of Calcium
In 1808, British chemist Sir Humphry Davy first isolated elemental calcium using electrolysis. Calcium is the lightest of all metals, with a density 1.55 g/cm3. It reacts with both air and water, usually in reactions involving calcium carbonate, but this reaction is quite slow because calcium hydroxide, Ca(OH)2, is not very soluble in water.
Reactions
Reaction of Calcium with Halides
Calcium forms salts with halides, such as $CaCl_2$ or $CaF_2$. They have a variety of uses, but the most usage most familiar to chemistry students is the use of calcium chloride as a desiccant (drying agent).
$CaCl_2+ 2 H_2O(l) \rightarrow CaCl_2 \cdot 2H_2O \nonumber$
Reaction of Calcium with Carbonates
Calcium carbonate is important in the formation of cave stalactites and stalagmites. This reaction allows calcium carbonate to be dissolved into solution as calcium bicarbonate.
$CaCO_3(s)+ CO_2(l)+ H_2O(l) \rightarrow Ca(HCO_3)_2 \nonumber$
The reverse reaction then allows the solution to become solid calcium carbonate once again, forming spikes of limestone in caves as the calcium bicarbonate solution drips vertically for several millennia.
Reaction of Calcium with Water
Calcium metal is fairly reactive and combines with water at room temperature of produce hydrogen gas and calcium hydroxide
$Ca(s) + 2H_2O(g) \rightarrow Ca(OH)_2(aq) + H_2(g) \nonumber$
Product will reveal hydrogen bubbles on calcium metal's surface.
Reaction of Calcium with Acid
Calcium dissolves in acid to form dissociated ions of Ca and Cl along with Hydrogen gas.
$Ca(s) + 2HCl(aq) \rightarrow Ca^{2+}(aq) + 2Cl^-(aq) + H_2 (g) \nonumber$
Reaction of Calcium with Oxygen
Calcium metal slowly oxidizes in air, becoming encrusted with white $CaO$ and $CaCO_3$, which protect from attack by air. When ignited, Calcium burns to give calcium xxide.
$2Ca \,(s)+ O_2 \,(g)\rightarrow 2CaO\,(s) \nonumber$
Calcium in living organisms
Because calcium is essential for life, it can be found in all organisms, living or dead. Shells of aquatic organisms, snail shells, and egg shells are all composed of mostly calcium carbonate, which can be dissolved in acid. Besides skeletal functions, the Ca2+ ion in animals and many organisms also plays an essential role in signal transduction pathways, neurotransmission, muscle function, fertilization, and enzymatic function. In plants, calcium is also important in the cell wall, membrane, and vacuole.
One of the most important calcium deposits is in coral reefs, which are comprised of mostly calcium carbonate. Coral secrete calcium carbonate over the period of their life, then die to allow new coral to build on top of their calcium carbonate structure. Over massive amounts of time, these calcium deposits grow into gigantic reefs, some of which can be seen from space (like the Great Barrier Reef in Australia). With the waters rich in sunlight and minerals like calcium, photosynthesis in sea plants is highly favored, allowing fish and other marine life to flourish in these regions.
Human bones are made up of mostly calcium phosphate ($Ca_3(PO_4)_2$). Cow milk also contains a large amount of calcium phosphate, which is why human culture encourages children and those particularly susceptible to osteoporosis to drink milk.
Uses of Calcium
The first known uses of calcium were by the Romans in the first century to make calcium oxide. Other written documentation around 975 C.E. suggests that plaster of Paris and Calcium Sulfate were medically useful. Calcium is available in a wide variety of forms, from limestone and chalk (calcium carbonate) to marble (calcite) and pearls. It also has many mineral forms (see link provided in Outside Links section). Calcium carbonate, in moderate amounts, can also be used as an antacid or a calcium supplement. Calcium nitrate is also a common fertilizer. Pure $CaCo_3$ can be extracted from limestone in a series of three reactions:
• Calcination: $CaCO_3 \rightarrow CaO + CO_2 \nonumber$
• Slaking: $CaO + H_2O \rightarrow Ca(OH)_2 \nonumber$
• Carbonation: $Ca(OH)_2 + CO_2 → CaCO_3 + H_2O \nonumber$
Fig. 2: A Blue Starfish resting on hard Acropora coral. Lighthouse, Ribbon Reefs, Great Barrier Reef. from Wikipedia.
Hard water
Hard water, as opposed to soft water, has a high mineral content of calcium sulfate ($CaSO_4$) or calcium carbonate ($CaCO_3$). It also includes magnesium ions (Mg2+) and sometimes iron, aluminum, and manganese. When left to evaporate, white calcium minerals can be seen on sinks, showers, etc.
Calcium oxide (quicklime)
Calcium oxide, often referred to as quicklime, has many commercial functions, some of which include making mortar and pottery, food, construction, agriculture, pollution control, and in medicine. It also heats quite readily with water in this reaction:
$CaO + H_2O \rightarrow Ca(OH)_2 \;\;\;\; \Delta H = −63.7\; kJ/mol \nonumber$
Problems
1. Chalk (calcium carbonate) is a common compound of calcium. When an acid such as acetic acid is added to chalk, carbon dioxide is formed. Write and balance this reaction. What are the products of this reaction?
2. How can pure calcium carbonate be produced? What are the reactions called?
3. Which reaction is responsible for the formation of stalactites and stalagmites? Under what conditions?
4. Which compound of calcium is found in cow's milk, and how does this relate to overall bone health?
5. A sample of hard water has 156.9ppm Ca2+. An ion exchange column removes all Ca2+ and replaces it with H3O+. What is the final pH of the water as it leaves the column, assuming it began at pH of 7 and no other ions were produced?
Solutions
1) 2CaCo3 + HC2H3O2 Ca(C2H3O2)2 + CO2 + H2O
Calcium acetate, c
arbon dioxide, and water.
2) Purification of limestone through calcination, slaking, and carbonation.
CaCO3 CaO + CO2
CaO + H2O Ca(OH)2
Ca(OH)2 + CO2 CaCO3 + H2O
3) CaCO3 + CO2 + H2O Ca(HCO3)2
Under acidic conditions.
4) Calcium phosphate (Ca3(PO4)2). Bones are made from a mineral of this called hydroxylapatite with the formula Ca10(PO4)6(OH)2. More calcium phosphate means more mineral for bone growth, ensuring that enough calcium is available for the body to both make bones and have enough Ca2+ ions for other important signaling processes.
5) 156.9 mg Ca2+ x (1 mol Ca2+)/(40080 mg Ca2+) x (2 mol H3O+)/(1 mol Ca2+)= 7.83 E-3 M H3O+
-log(7.83 E-3)= pH 2.11
Contributors and Attributions
• Anna Zhu (UCD) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements%3A_The_Alkaline_Earth_Metals/Z020_Chemistry_of_Calcium_%28Z.txt |
Calcium is the 20th element in the periodic table. It is a group 2 metal, also known as an alkaline-earth metal, and no populated d-orbital electrons. Calcium is the fifth most abundant element by mass (3.4%) in both the Earth's crust and in seawater. All living organisms require calcium for survival. Calcium is a silver-gray metal which takes its name from the Latin word calx, which means lime. It is the fifth most abundant element in the earth's crust and is widely distributed as limestone (CaCO3), quicklime (CaO) and calcium fluoride.
General Properties of Calcium
Symbol Ca
Color dull gray or silver
Atomic Number 20
Category alkaline earth metal
Atomic Weight 40.078 g•mol−1
Group,Period,Block 2,4,s
Electron Configuration [Ar]4s2
Valence Electrons 2
Phase (room temperature) solid
Melting Point 1115 K, 842°C
Boiling Point 1757 K, 1484 °C
Atomic Radius 197 pm
Oxidation States 2
Density at room temp 1.55 g•cm−3
Electronegativity 1.00 (Pauling)
First ionization energy 589.8 kJ•mol−1
Number of stable isotopes 4 (2 more are fairly stable)
Electronic Structure
Fig. 1: Calcium Atom
Discovery and Properties of Calcium
In 1808, British chemist Sir Humphry Davy first isolated elemental calcium using electrolysis. Calcium is the lightest of all metals, with a density 1.55 g/cm3. It reacts with both air and water, usually in reactions involving calcium carbonate, but this reaction is quite slow because calcium hydroxide, Ca(OH)2, is not very soluble in water.
Reactions
Reaction of Calcium with Halides
Calcium forms salts with halides, such as $CaCl_2$ or $CaF_2$. They have a variety of uses, but the most usage most familiar to chemistry students is the use of calcium chloride as a desiccant (drying agent).
$CaCl_2+ 2 H_2O(l) \rightarrow CaCl_2 \cdot 2H_2O \nonumber$
Reaction of Calcium with Carbonates
Calcium carbonate is important in the formation of cave stalactites and stalagmites. This reaction allows calcium carbonate to be dissolved into solution as calcium bicarbonate.
$CaCO_3(s)+ CO_2(l)+ H_2O(l) \rightarrow Ca(HCO_3)_2 \nonumber$
The reverse reaction then allows the solution to become solid calcium carbonate once again, forming spikes of limestone in caves as the calcium bicarbonate solution drips vertically for several millennia.
Reaction of Calcium with Water
Calcium metal is fairly reactive and combines with water at room temperature of produce hydrogen gas and calcium hydroxide
$Ca(s) + 2H_2O(g) \rightarrow Ca(OH)_2(aq) + H_2(g) \nonumber$
Product will reveal hydrogen bubbles on calcium metal's surface.
Reaction of Calcium with Acid
Calcium dissolves in acid to form dissociated ions of Ca and Cl along with Hydrogen gas.
$Ca(s) + 2HCl(aq) \rightarrow Ca^{2+}(aq) + 2Cl^-(aq) + H_2 (g) \nonumber$
Reaction of Calcium with Oxygen
Calcium metal slowly oxidizes in air, becoming encrusted with white $CaO$ and $CaCO_3$, which protect from attack by air. When ignited, Calcium burns to give calcium xxide.
$2Ca \,(s)+ O_2 \,(g)\rightarrow 2CaO\,(s) \nonumber$
Calcium in living organisms
Because calcium is essential for life, it can be found in all organisms, living or dead. Shells of aquatic organisms, snail shells, and egg shells are all composed of mostly calcium carbonate, which can be dissolved in acid. Besides skeletal functions, the Ca2+ ion in animals and many organisms also plays an essential role in signal transduction pathways, neurotransmission, muscle function, fertilization, and enzymatic function. In plants, calcium is also important in the cell wall, membrane, and vacuole.
One of the most important calcium deposits is in coral reefs, which are comprised of mostly calcium carbonate. Coral secrete calcium carbonate over the period of their life, then die to allow new coral to build on top of their calcium carbonate structure. Over massive amounts of time, these calcium deposits grow into gigantic reefs, some of which can be seen from space (like the Great Barrier Reef in Australia). With the waters rich in sunlight and minerals like calcium, photosynthesis in sea plants is highly favored, allowing fish and other marine life to flourish in these regions.
Human bones are made up of mostly calcium phosphate ($Ca_3(PO_4)_2$). Cow milk also contains a large amount of calcium phosphate, which is why human culture encourages children and those particularly susceptible to osteoporosis to drink milk.
Uses of Calcium
The first known uses of calcium were by the Romans in the first century to make calcium oxide. Other written documentation around 975 C.E. suggests that plaster of Paris and Calcium Sulfate were medically useful. Calcium is available in a wide variety of forms, from limestone and chalk (calcium carbonate) to marble (calcite) and pearls. It also has many mineral forms (see link provided in Outside Links section). Calcium carbonate, in moderate amounts, can also be used as an antacid or a calcium supplement. Calcium nitrate is also a common fertilizer. Pure $CaCo_3$ can be extracted from limestone in a series of three reactions:
• Calcination: $CaCO_3 \rightarrow CaO + CO_2 \nonumber$
• Slaking: $CaO + H_2O \rightarrow Ca(OH)_2 \nonumber$
• Carbonation: $Ca(OH)_2 + CO_2 → CaCO_3 + H_2O \nonumber$
Fig. 2: A Blue Starfish resting on hard Acropora coral. Lighthouse, Ribbon Reefs, Great Barrier Reef. from Wikipedia.
Hard water
Hard water, as opposed to soft water, has a high mineral content of calcium sulfate ($CaSO_4$) or calcium carbonate ($CaCO_3$). It also includes magnesium ions (Mg2+) and sometimes iron, aluminum, and manganese. When left to evaporate, white calcium minerals can be seen on sinks, showers, etc.
Calcium oxide (quicklime)
Calcium oxide, often referred to as quicklime, has many commercial functions, some of which include making mortar and pottery, food, construction, agriculture, pollution control, and in medicine. It also heats quite readily with water in this reaction:
$CaO + H_2O \rightarrow Ca(OH)_2 \;\;\;\; \Delta H = −63.7\; kJ/mol \nonumber$
Problems
1. Chalk (calcium carbonate) is a common compound of calcium. When an acid such as acetic acid is added to chalk, carbon dioxide is formed. Write and balance this reaction. What are the products of this reaction?
2. How can pure calcium carbonate be produced? What are the reactions called?
3. Which reaction is responsible for the formation of stalactites and stalagmites? Under what conditions?
4. Which compound of calcium is found in cow's milk, and how does this relate to overall bone health?
5. A sample of hard water has 156.9ppm Ca2+. An ion exchange column removes all Ca2+ and replaces it with H3O+. What is the final pH of the water as it leaves the column, assuming it began at pH of 7 and no other ions were produced?
Solutions
1) 2CaCo3 + HC2H3O2 Ca(C2H3O2)2 + CO2 + H2O
Calcium acetate, c
arbon dioxide, and water.
2) Purification of limestone through calcination, slaking, and carbonation.
CaCO3 CaO + CO2
CaO + H2O Ca(OH)2
Ca(OH)2 + CO2 CaCO3 + H2O
3) CaCO3 + CO2 + H2O Ca(HCO3)2
Under acidic conditions.
4) Calcium phosphate (Ca3(PO4)2). Bones are made from a mineral of this called hydroxylapatite with the formula Ca10(PO4)6(OH)2. More calcium phosphate means more mineral for bone growth, ensuring that enough calcium is available for the body to both make bones and have enough Ca2+ ions for other important signaling processes.
5) 156.9 mg Ca2+ x (1 mol Ca2+)/(40080 mg Ca2+) x (2 mol H3O+)/(1 mol Ca2+)= 7.83 E-3 M H3O+
-log(7.83 E-3)= pH 2.11
Contributors and Attributions
• Anna Zhu (UCD) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements%3A_The_Alkaline_Earth_Metals/Z020_Chemistry_of_Calcium_(Z20.txt |
Strontium is a group 2 element that does not occur as a free element due to its extreme reactivity with oxygen and water. It occurs naturally only in compounds with other elements such as strontianite. It is softer than calcium and decomposes water more vigorously. It has a silver appearance but then turns yellow with the formation of oxide.
General Information
• Atomic Number: 38
• Symbol: Sr
• Atomic weight: 87.62
• Electron shell configuration: [Kr]5s^2
• Melting Point (K): 1042
• Boiling Point(K): 1657
Introduction
The element strontium is named for a Scottish town, Strontian. It was isolated in 1808 by Davy and is a silvery and malleable metal that reacts vigorously with water to produce hydrogen gas. It has the same relative abundance as carbon and sulfur but does not occur in pure form.
Characteristics
Strontium is softer than calcium and decomposes vigorously in water. It is a silvery color but rapidly oxidizes to yellow due to the formation of strontium oxide. Because of its propensity for oxidation and ignition, strontium is stored typically under kerosene. Finely powdered strontium metal is sufficiently reactive to ignite spontaneously in air. It reacts with water quickly (but not violently like group 1 elements) to produce strontium hydroxide and hydrogen gas. Strontium and its compounds burn with a crimson flame and are used in fireworks.
Applications
Strontium compounds are useful in pyrotechnic devices and signal flares because of the bright crimson coloring they give to flames.
Strontium is used for producing glass for color televisions. It is also used in producing ferrite ceramic magnets and in refining zinc. Strontium atoms helped develop the world's most accurate atomic clock, which is accurate to one second in 200 million years. Toothpaste for sensitive teeth uses strontium chloride, and strontium oxide is used to improve the quality of pottery glazes. The isotope 90Sr is one of the world's best long-lived, high energy beta emitters known. It is also used in cancer therapy.
Strontium-90, a radioactive isotope of the metal produced by fission reactions is a dangerous environmental menace because its chemistry is similar to calcium and it may take its place in bones. The strong radiation emitted by the isotope interferes with the production of new blood cells and can cause death.
Contributors and Attributions
• Stephen R. Marsden
Z038 Chemistry of Strontium (Z
Strontium is a group 2 element that does not occur as a free element due to its extreme reactivity with oxygen and water. It occurs naturally only in compounds with other elements such as strontianite. It is softer than calcium and decomposes water more vigorously. It has a silver appearance but then turns yellow with the formation of oxide.
General Information
• Atomic Number: 38
• Symbol: Sr
• Atomic weight: 87.62
• Electron shell configuration: [Kr]5s^2
• Melting Point (K): 1042
• Boiling Point(K): 1657
Introduction
The element strontium is named for a Scottish town, Strontian. It was isolated in 1808 by Davy and is a silvery and malleable metal that reacts vigorously with water to produce hydrogen gas. It has the same relative abundance as carbon and sulfur but does not occur in pure form.
Characteristics
Strontium is softer than calcium and decomposes vigorously in water. It is a silvery color but rapidly oxidizes to yellow due to the formation of strontium oxide. Because of its propensity for oxidation and ignition, strontium is stored typically under kerosene. Finely powdered strontium metal is sufficiently reactive to ignite spontaneously in air. It reacts with water quickly (but not violently like group 1 elements) to produce strontium hydroxide and hydrogen gas. Strontium and its compounds burn with a crimson flame and are used in fireworks.
Applications
Strontium compounds are useful in pyrotechnic devices and signal flares because of the bright crimson coloring they give to flames.
Strontium is used for producing glass for color televisions. It is also used in producing ferrite ceramic magnets and in refining zinc. Strontium atoms helped develop the world's most accurate atomic clock, which is accurate to one second in 200 million years. Toothpaste for sensitive teeth uses strontium chloride, and strontium oxide is used to improve the quality of pottery glazes. The isotope 90Sr is one of the world's best long-lived, high energy beta emitters known. It is also used in cancer therapy.
Strontium-90, a radioactive isotope of the metal produced by fission reactions is a dangerous environmental menace because its chemistry is similar to calcium and it may take its place in bones. The strong radiation emitted by the isotope interferes with the production of new blood cells and can cause death.
Contributors and Attributions
• Stephen R. Marsden | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements%3A_The_Alkaline_Earth_Metals/Z038_Chemistry_of_Strontium_%2.txt |
Barium takes it name from the Greek word barys for heavy. Barium was first discovered in 1774 by Carl Scheele, but was not isolated as a pure metal until 1808 when Sir Humphry Davy electrolyzed molten barium salts. The name Barium comes from the Greek word barys, which means heavy. Barium is a soft, silvery white metal, and has a melting point of 1000 K. Because of its reaction to air, barium cannot be found in nature in its pure form but can be extracted from the mineral barite.
• Atomic Number = 56
• Mass = 137.3 g mol-1
• Electrion Configuration = [Xe]6s2
• Density = 3.51 g cm-3
Reactivity
Like the lighter members of its family, barium reacts vigorously with water to produce hydrogen gas and so is commonly stored in oil.
Abundance and Extraction
The metal does not occur free in nature but chiefly as the sulfate and carbonate. The sulfate is used in X-ray diagnostics as a contrast medium (i.e., in soft tissue like the digestive tract).
Barium Isotopes
There are seven stable isotopes of naturally occurring barium: 130Ba, 132Ba, 134Ba, 135Ba, 136Ba, 137Ba, and 138Ba. In total, twenty-two isotopes are known to exist, but a majority of them are highly radioactive and have relatively short half-lives.
Barium Compounds
Barium sulfate (BaSO4), or barite, is the most common mineral abundant in barium. This mineral has a density of 4.5g/cm3 and is extremely insoluble in water. Uses of barium sulfate include being a radiocontrast agent for X-ray imaging of the digestive system. Barium carbonate (BaCO3) is also commonly used as a rat poison.
Barium compounds (which are toxic) are also useful in pyrotechnic devices where they impart a characteristic green color.
Contributors and Attributions
• Stephen R. Marsden
Z056 Chemistry of Barium (Z56)
Barium takes it name from the Greek word barys for heavy. Barium was first discovered in 1774 by Carl Scheele, but was not isolated as a pure metal until 1808 when Sir Humphry Davy electrolyzed molten barium salts. The name Barium comes from the Greek word barys, which means heavy. Barium is a soft, silvery white metal, and has a melting point of 1000 K. Because of its reaction to air, barium cannot be found in nature in its pure form but can be extracted from the mineral barite.
• Atomic Number = 56
• Mass = 137.3 g mol-1
• Electrion Configuration = [Xe]6s2
• Density = 3.51 g cm-3
Reactivity
Like the lighter members of its family, barium reacts vigorously with water to produce hydrogen gas and so is commonly stored in oil.
Abundance and Extraction
The metal does not occur free in nature but chiefly as the sulfate and carbonate. The sulfate is used in X-ray diagnostics as a contrast medium (i.e., in soft tissue like the digestive tract).
Barium Isotopes
There are seven stable isotopes of naturally occurring barium: 130Ba, 132Ba, 134Ba, 135Ba, 136Ba, 137Ba, and 138Ba. In total, twenty-two isotopes are known to exist, but a majority of them are highly radioactive and have relatively short half-lives.
Barium Compounds
Barium sulfate (BaSO4), or barite, is the most common mineral abundant in barium. This mineral has a density of 4.5g/cm3 and is extremely insoluble in water. Uses of barium sulfate include being a radiocontrast agent for X-ray imaging of the digestive system. Barium carbonate (BaCO3) is also commonly used as a rat poison.
Barium compounds (which are toxic) are also useful in pyrotechnic devices where they impart a characteristic green color.
Contributors and Attributions
• Stephen R. Marsden
Z088 Chemistry of Radium (Z8
Radium takes its name from the Latin word radius or ray. All isotopes of radium are radioactive and many exhibit luminescence, reacting readily with oxygen and water. The metal was discovered and isolated in 1911 by Marie Curie.
When first discovered, compounds of the metal were used on watch dials for self-luminescence and in early cancer therapy. It has been all but replaced now by safer alternatives.
Contributors and Attributions
Stephen R. Marsden
Z088 Chemistry of Radium (Z88)
Radium takes its name from the Latin word radius or ray. All isotopes of radium are radioactive and many exhibit luminescence, reacting readily with oxygen and water. The metal was discovered and isolated in 1911 by Marie Curie.
When first discovered, compounds of the metal were used on watch dials for self-luminescence and in early cancer therapy. It has been all but replaced now by safer alternatives.
Contributors and Attributions
Stephen R. Marsden | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/1_s-Block_Elements/Group__2_Elements%3A_The_Alkaline_Earth_Metals/Z056_Chemistry_of_Barium_%28Z5.txt |
The boron family contains elements in group 13 of the periodic talbe and include the semi-metal boron (B) and the metals aluminum (Al), gallium (Ga), indium (In), and thallium (Tl). Aluminum, gallium, indium, and thallium have three electrons in their outermost shell (a full s orbital and one electron in the p orbital) with the valence electron configuration ns2np1. The elments of the boron family adopts oxidation states +3 or +1. The +3 oxidation states are favorable except for the heavier elements, such as Tl, which prefer the +1 oxidation state due to its stability; this is known as the inert pair effect. The elements generally follow periodic trends except for certain Tl deviations:
• Group 13: Chemical Reactivity
The boron family contains the semi-metal boron (B) and metals aluminum (Al), gallium (Ga), indium (In), and thallium (Tl).
• Group 13: Physical Properties of Group 13
The boron family contains the semi-metal boron (B) and metals aluminum (Al), gallium (Ga), indium (In), and thallium (Tl).
• Chemistry of Boron (Z=5)
Boron is the fifth element of the periodic table (Z=5), located in Group 13. It is classified as a metalloid due it its properties that reflect a combination of both metals and nonmetals.
• Chemistry of Aluminum (Z=13)
Aluminum (also called Aluminium) is the third most abundant element in the earth's crust. It is commonly used in the household as aluminum foil, in crafts such as dyeing and pottery, and also in construction to make alloys. In its purest form the metal is bluish-white and very ductile. It is an excellent conductor of heat and electricity and finds use in some wiring. When pure it is too soft for construction purposes but addition of small amounts of silicon and iron hardens it significantly.
• Chemistry of Gallium (Z=31)
Gallium is the chemical element with the atomic number 31 and symbol Ga on the periodic table. It is in the Boron family (group 13) and in period 4. Gallium was discovered in 1875 by Paul Emile Lecoq de Boisbaudran. Boisbaudran named his newly discovered element after himself, deriving from the Latin word, “Gallia,” which means “Gaul.” Elemental Gallium does not exist in nature but gallium (III) salt can be extracted in small amounts from bauxite and zinc ores.
• Chemistry of Indium (Z=49)
Indium has the chemical symbol In and the atomic number 49. It has the electron configuration [Kr] 2s22p1 and may adopt the +1 or +3 oxidation state; however, the +3 state is more common. It is a soft, malleable metal that is similar to gallium. Indium forms InAs, which is found in photoconductors in optical instruments. The physical properties of indium include its silver-white color and the "tin cry" it makes when bent. Indium is soluble in acids, but does not react with oxygen at room tempera
• Chemistry of Thalium (Z=81)
Thallium has the chemical symbol Tl and atomic number 81. It has the electron configuration \([Xe] 2s^22p^1\) and has a +3 or +1 oxidation state. As stated above, because thallium is heavy, it has a greater stability in the +1 oxidation state (inert pair effect). Therefore, it is found more commonly in its +1 oxidation state. Thallium is soft and malleable.
• Chemistry of Nihonium (Z=113)
In studies announced jointly by the Joint Institute for Nuclear Research in Dubna, Russia, and the Lawrence Livermore National Laboratory in the U.S., four atoms of element 113 were produced in 2004 via decay of element 115 after the fusion of Ca-48 and Am-243.
Thumbnail: Crystals of 99.999% gallium. (CC-SA-BY 3.0; Foobar)
Group 13: The Boron Family
The boron family contains the semi-metal boron (B) and metals aluminum (Al), gallium (Ga), indium (In), and thallium (Tl). The boron family adopts oxidation states +3 or +1. The +3 oxidation states are favorable except for the heavier elements, such as Tl, which prefer the +1 oxidation state due to its stability; this is known as the inert pair effect. The elements generally follow periodic trends except for certain Tl deviations:
Boron
Boron tends to forms hydrides, the simplest of which is diborane, $B_2H_6$. Boron hydrides are used to synthesize organic compounds. One of the main compounds used to form other boron compounds is boric acid, which is a weak acid and is formed in the following two-step reaction:
$B_2O_{3 \;(s)} + 3 H_2O _{(l)} \rightarrow 2 B(OH)_{3 (aq)} \nonumber$
$B(OH)_{3 \;(aq)} + 2 H_2O_{(l)} \rightarrow H_3O^+_{(aq)} + B(OH)^-_{4\; (aq)} \nonumber$
Boron can be crystallized from a solution of hydrogen peroxide and borax to produce sodium perborate, a bleach alternative. The bleaching ability of perborate is due to the two peroxo groups bound to the boron atoms.
Aluminum
Aluminum is an active metal with the electron configuration [Ne] 2s22p1, and usually adopts a +3 oxidation state. This element is the most abundant metal in the Earth's crust (7.5-8.4%). Even though it is very abundant, before 1886 aluminum was considered a semiprecious metal; it was difficult to isolate due to its high melting point. Aluminum is very expensive to produce, because the electrolysis of one mole of aluminum requires three moles of electrons:
$Al^{3+} + 3e^- \rightarrow Al(l) \nonumber$
Aluminum can dissolve in both acids and bases—it is amphoteric. In an aqueous OH- solution it produce Al(OH)4-, and in an aqueous H3O+ solution it produce [Al(H2O)6]3+. Another important feature of aluminum is that it is a good reducing agent due to its +3 oxidation state. It can therefore react with acids to reduce H+(aq) to H2(g). For example:
$2Al (s) + 6H^+(aq) \rightarrow 2Al^{3+}(aq) + 3H_2(g) \nonumber$
Aluminum can also extract oxygen from any metal oxide. The following reaction, which is known as the thermite reaction, is very exothermic:
$Fe_2O_3(s) + 2 Al(s) \rightarrow Al_2O_3(s) +2 Fe(l) \nonumber$
Gallium
Gallium has the chemical symbol Ga and the atomic number 31. It has the electron configuration [Ar] 2s22p1 and a +3 oxidation state. Gallium is industrially important because it forms gallium arsenide (GaAs), which converts light directly into electricity. Gallium is also used in conjunction with aluminum to generate hydrogen. In a process similar to the thermite reaction, aluminum extracts oxygen from water and releases hydrogen gas. However, as mentioned above, aluminum forms a protective coat in the presence of water. Combining gallium and aluminum prevents the formation of this protective layer, allowing aluminum to reduce water to hydrogen.[7]
Indium
Indium has the electron configuration [Kr] 2s22p1 and may adopt the +1 or +3 oxidation state; however, the +3 state is more common. Indium is soluble in acids, but does not react with oxygen at room temperature. It is obtained by separation from zinc ores. Indium is mainly used to make alloys, and only a small amount is required to enhance the metal strength.
Thallium
Thallium has the electron configuration [Xe] 2s22p1 and has a +3 or +1 oxidation state. Because thallium is heavy, it has a greater stability in the +1 oxidation state (inert pair effect). Therefore, it is found more commonly in its +1 oxidation state. Thallium is soft and malleable. It is poisonous, but used in high-temperature superconductors.
Diagonal Relationship of Beryllium and Aluminum
Both Be2+ and Al3+ are hydrated to produce [Be(H2O)4]2+ and Al(H2O)63+, respectively. When reacted with water, both compounds produce hydronium ions, making them slightly acidic. Another similarity between aluminum and beryllium is that they are amphoteric, and their hydroxides are very basic. Both metals also react with oxygen to produce oxide coatings capable of protecting other metals from corrosion. Both metals also react with halides that can act as Lewis acids.
Summary of Boron Group Trends
1. The chemistry of boron is quite different from that of the heavier Group IIIA (Boron column) elements. It differs from aluminum in the following ways.
1. Its oxide and hydroxide are acidic, where as those of aluminum are amphoteric.
2. Boron is a semiconductor which has various polymorphs based on icosohedral boron cages, whereas aluminum is a metal with a close packed structure. Boron is very inert and only attacked by hot concentrated oxidizing acids.
3. No simple salts of B3+ are known, whereas those of Al3+ are numerous and well documented.
4. Boron forms a wide range of hydrides, which have cage structures. (AlH3)4 has a polymeric structure which resembles that of AlF3.
5. The stereochemistries of many boron compounds are based on trigonal sp2 and tetrahedral sp3 geometries. In the latter the octet rule is obeyed. Aluminum forms many compounds with tetrahedral, trigonal bipyramidal, and octahedral geometries.
6. Multiple pπ- pπ bonding in boron-nitrogen, boron-oxygen, and boron-fluorine compounds is more significant than for the corresponding aluminum compounds. {BN}x for example adopts a graphite structure.
2. Aluminum, gallium, indium and thallium all form a range of compounds in the +3 oxidation state and compounds in the +1 oxidation state become progressively more stable down group IIIA.
3. The oxides of aluminum and gallium are amphoteric and indium and thallium oxides are more basic.
4. The octahedral aqua-ions [M(OH2)6]3+ are acidic and the pKa values for the equillibria: [M(OH2)6]3+ ó [M(OH2)5(OH)]2+ + H+ are Al, ~5; Ga, 3; In, ~4; Tl, 1; showing the Al3+ aquo-ion is the least acidic and the Tl3+ ion the most acidic.
5. The MX3 compounds are Lewis acids and Lewis acid strengths decrease in the order: Al > Ga > In.
6. The stability of the hydrides decreases down the group and there are no stable Tl-H compounds. Extraordinary precautions are required to exclude air and moisture in order to isolate Ga2H6.
7. Aluminum is resistant to corrosion because of an impermeable oxide layer, but is soluble in non-oxidizing mineral acids. Gallium, indium, and thallium dissolve readily in acids, but thallium dissolves slowly in H2SO4 and HCl.
8. AlN, GaN, and InN have wurtzite structures analogous to cubic BN, but show no analogue of the graphite-type structure of BN. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_13%3A_The_Boron_Family/1Group_13%3A_Chemical_Reactivity.txt |
The boron family contains the semi-metal boron (B) and metals aluminum (Al), gallium (Ga), indium (In), and thallium (Tl).
Properties and Periodic Trends
These elements are found in Group 13 (XIII) of the p block in the Periodic Table of Elements. Aluminum, gallium, indium, and thallium are metallic. They each have three electrons in their outermost shell (a full s orbital and one electron in the p orbital) with the valence electron configuration ns2np1. The boron family adopts oxidation states +3 or +1. The +3 oxidation states are favorable except for the heavier elements, such as Tl, which prefer the +1 oxidation state due to its stability; this is known as the inert pair effect. The elements generally follow periodic trends except for certain Tl deviations:
• Atomic radius increases down the group (Tl has the largest atomic radius.)
• Electrode potential increases down the group (reactivity decreases down the group)
• Ionization Energy decreases going down the group (because the electrons are farther from the core and therefore are easier to remove; Tl does not fit this trend)*
Table 1: Properties & Trends of Group 3 Elements
Elemental Symbol Atomic Number (Z) Molecular Mass (g/mol) Melting Point °C Standard Reduction Potential (V) Ionization Energy (kJ/mol)
Boron B 5 10.811 2076 - 801
Aluminum Al 13 26,9815 660 -1.68 578
Gallium Ga 31 69.723 29.8 -0.56 558
Indium In 49 114.818 156 0.34 558
Thallium Th 81 205.383 303 +0.72 589
Boron
Boron is the first element of Group 13 and is the only metalloid of the group. Its chemical symbol is B, and it has an atomic number of 5. Boron has the electron configuration [He] 2s22p1and prefers an oxidation state of +3. Boron has no natural elemental form; it forms compounds which are abundant in the Earth's crust. Boron is an essential nutrient for plants. There are a few locations where boron ores, known as borax, are found in great concentrations. Due to its lack of a complete octet, boron is a Lewis acid. It tends to forms hydrides, the simplest of which is diborane, $B_2H_6$. Boron hydrides are used to synthesize organic compounds. One of the main compounds used to form other boron compounds is boric acid, which is a weak acid and is formed in the following two-step reaction:
$B_2O_{3 \;(s)} + 3 H_2O _{(l)} \rightarrow 2 B(OH)_{3 (aq)} \nonumber$
$B(OH)_{3 \;(aq)} + 2 H_2O_{(l)} \rightarrow H_3O^+_{(aq)} + B(OH)^-_{4\; (aq)} \nonumber$
Boron can be crystallized from a solution of hydrogen peroxide and borax to produce sodium perborate, a bleach alternative. The bleaching ability of perborate is due to the two peroxo groups bound to the boron atoms.
Aluminum
Aluminum is the most important metal in the boron family, with the chemical symbol Al and atomic number 13. It is used in lightweight alloys and is an active metal. It has the electron configuration [Ne] 2s22p1, and usually adopts a +3 oxidation state. This element is the most abundant metal in the Earth's crust (7.5-8.4%). Even though it is very abundant, before 1886 aluminum was considered a semiprecious metal; it was difficult to isolate due to its high melting point. Aluminum is very expensive to produce, because the electrolysis of one mole of aluminum requires three moles of electrons:
$Al^{3+} + 3e^- \rightarrow Al(l) \nonumber$
Aluminum is a soft, malleable metal that is silver or gray in color. It is highly reactive, and therefore found in nature in compounds. Aluminum does not appear to react with water because it is aluminum is protected by a layer of Al2O3; this effect is known as anodizing. The thickness of the Al2O3 layer varies based on galvanic reactions, but it prevents the metal from oxidizing further. Aluminum is used in many alloys to prevent corrosion.
• Aluminum Oxide (Al2O3): Commonly referred to as alumina, it has highly desirable metallic characteristics due to its strong ionic bonding. It is an excellent thermal insulator, and forms corondum upon crystallization. Corondum exist in several forms, including in sapphires and rubies. The differences in the colors of these gems are due to transition metal impurities in their corondum structure.
• Aluminum Sulfate Al2(SO4)3: Very important commercial compound. Used in sizing paper (paper in which waxes and glues are used to make the paper more water resistant.) Aluminum sulfate, however, has a acidic properties that may deteriorate the paper.
Aluminum can dissolve in both acids and bases—it is amphoteric. In an aqueous OH- solution it produce Al(OH)4-, and in an aqueous H3O+ solution it produce [Al(H2O)6]3+
Another important feature of aluminum is that it is a good reducing agent due to its +3 oxidation state. It can therefore react with acids to reduce H+(aq) to H2(g). For example:
$2Al (s) + 6H^+(aq) \rightarrow 2Al^{3+}(aq) + 3H_2(g) \nonumber$
Aluminum can also extract oxygen from any metal oxide. The following reaction, which is known as the thermite reaction, is very exothermic:
$Fe_2O_3(s) + 2 Al(s) \rightarrow Al_2O_3(s) +2 Fe(l) \nonumber$
Gallium
Gallium has the chemical symbol Ga and the atomic number 31. It has the electron configuration [Ar] 2s22p1 and a +3 oxidation state. The melting point is 29.8º C, slightly above room temperature. Gallium has the second lowest melting point (after mercury) and can remain in the liquid phase at a larger range of temperatures than any other substance. Gallium is industrially important because it forms gallium arsenide (GaAs), which converts light directly into electricity. Gallium is also used in conjunction with aluminum to generate hydrogen. In a process similar to the thermite reaction, aluminum extracts oxygen from water and releases hydrogen gas. However, as mentioned above, aluminum forms a protective coat in the presence of water. Combining gallium and aluminum prevents the formation of this protective layer, allowing aluminum to reduce water to hydrogen.[7]
Indium
Indium has the chemical symbol In and the atomic number 49. It has the electron configuration [Kr] 2s22p1 and may adopt the +1 or +3 oxidation state; however, the +3 state is more common. It is a soft, malleable metal that is similar to gallium. Indium forms InAs, which is found in photoconductors in optical instruments. The physical properties of indium include its silver-white color and the "tin cry" it makes when bent. Indium is soluble in acids, but does not react with oxygen at room temperature. It is obtained by separation from zinc ores. Indium is mainly used to make alloys, and only a small amount is required to enhance the metal strength. For example, indium is added to gold or platinum to make the metals more useful industrial tools.
Thallium
Thallium has the chemical symbol Tl and atomic number 81. It has the electron configuration [Xe] 2s22p1 and has a +3 or +1 oxidation state. As stated above, because thallium is heavy, it has a greater stability in the +1 oxidation state (inert pair effect). Therefore, it is found more commonly in its +1 oxidation state. Thallium is soft and malleable. It is poisonous, but used in high-temperature superconductors. Because of its toxicity, thallium was widely used in insecticide and rat poison until this usage was prohibited in 1975 in the U.S.
Diagonal Relationship of Beryllium and Aluminum
Both Be2+ and Al3+ are hydrated to produce [Be(H2O)4]2+ and Al(H2O)63+, respectively. When reacted with water, both compounds produce hydronium ions, making them slightly acidic. Another similarity between aluminum and beryllium is that they are amphoteric, and their hydroxides are very basic. Both metals also react with oxygen to produce oxide coatings capable of protecting other metals from corrosion. Both metals also react with halides that can act as Lewis acids.
Problems
( highlight the blue areas to find the answers)
1. T/F In reality, aluminum forms a protective layer and does not react with water.
True, This is known as anodizing.
2. Which statement about Gallium is false?
1. It melts on contact with human hands T
2. It can combine with aluminum to reduce water T
3. It is mainly found in the oxidation state +1 F
4. It can form a good source of hydrogen T
3. T/F Thallium is highly toxic and therefore it is commonly used for rat poisons and insecticides in the United States.
False Since 1975, thallium is prohibited from such usage since there is no warning if one digested it.
4. Boron:
1. has the electron configuration [Ne] 2s22p1 F
2. is the first metal of Group 13. F
3. has an atomic number 6. F
4. is an important element that we use in our daily lives. T
5. All of the above is correct. F
5. Aluminum Oxide:
1. has poor corrosion resistance F
2. is not a very good thermal insulator F
3. in its crystalline form it is called corondum T
4. is not a very reactive metal F
6. T/F Aluminum is amphoteric
True, aluminum can dissolve in both acids and bases
7. Which element is the only metalloid in the boron family?
Boron
8. When beryllium reacts with a halide, which of the following is true?
1. It acts as a Lewis base F
2. It forms a covalent bond F
3. It forms a Lewis acid T
4. It forms a neural molecule F
9. What is the electron configuration of thallium?
[Xe]2s22p1
10. Which statement is False?
1. Thallium is the heaviest element. T
2. Boron has the highest melting point. T
3. Electron potential increases going down the group. T
4. Thallium has the lowest ionization energy. F
5. All of the above are correct. T
Contributors and Attributions
• Stephanie Lee (UCD), Constantine La (CU-Boulder), Zoe Lim (UCD) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_13%3A_The_Boron_Family/1Group_13%3A_General_Properties_and_Reactions.txt |
Boron is the fifth element of the periodic table (Z=5), located in Group 13. It is classified as a metalloid due it its properties that reflect a combination of both metals and nonmetals.
Introduction
The name Boron comes from the Arabic and Persian words for borax, its principal ore. Although compounds of boron were known in ancient times, it was first isolated in 1808 by Gay-Lussac and Thénard and independently by Sir Humphry Davy (who has a lot of elements to his credit!).
Boron exists in the earth's crust to the extent of only about 10 ppm (about the same abundance as lead). The pure element is shiny and black. It is very hard and in extremely pure form is nearly as hard as diamond, but much too brittle for practical use. At high temperatures it is a good conductor but at room temperature and below is an insulator. This behavior as well as many of its other properties earn it the classification of a metalloid. In addition to the crystalline form of boron there is also an amorphous dark brown powder (as shown above).
The element can be prepared by the reduction of borax (\(Na_2B_4O_7\)) with carbon. High-purity boron can be produced by electrolysis of molten potassium fluoroborate. Common compounds of boron include borax and boric acid (\(H_3BO_3\)).
Atomic Mass 10.811 g/mol
Electronic Configuration [He]2s2 2p1
Melting Point 2349 K
Boiling Point 4200 K
Heat of Fusion 50.2 kJ/mol
Heat of Vaporization 480 kJ/mol
Specific Heat Capacity 11.087 J/mol·K
Oxidation States +4, +3, +2, +1
Magnetic Ordering diamagnetic
Electronegativity 2.04
Atomic Radius 90 pm
Stable Isotopes 10B, 11B
Boron is the only element in group 3 that is not a metal. It has properties that lie between metals and non-metals (semimetals). For example Boron is a semiconductor unlike the rest of the group 13 elements. Chemically, it is closer to aluminum than any of the other group 13 elements.
History
Boron was first discovered by Joseph-Louis Gay-Lussac and Louis-Jaques Thenard, and independently by Humphry Davy in the year 1808. These chemists isolated Boron by combining boric acid with potassium. Today, there are many ways of obtaining Boron but the most common way is by heating borax (a compound of sodium and boron) with calcium.
Boron and its Compounds
Many boron compounds are electron-deficient, meaning that they lack an octet of electrons around the central boron atom. This deficiency is what accounts for boron being a strong Lewis acid, in that it can accept protons (H+ ions) in solution. Boron-hydrogen compounds are referred to as boron hydrides, or boranes.
Boranes
In the molecule BH3, each of the 3 hydrogen atoms is bonded to the central boron atom. The boron atom has only six electrons in its outer shell, leading to an electron deficiency.
Diborane:
H H
I I
H - B ? B - H
I I
H H
This molecule has 12 valence shell electrons; 3 each from the B atoms, and 1 each from the six H atoms. To make this structure follow the rules required to draw any lewis structure model, then it must have 14 valence shell electrons; however it does not. According to this figure, the two B atoms and four H atoms lie in the same plane (sp3- perpendicular to the plane of the page). In these four bonds 8 electrons are involved. Four electrons bond the remaining H atoms to the two B atoms and the B atoms together. This is done when the two H atoms simultaneously bond to the two B atoms. This creates what is called an atom "bridge" because there are two electrons shared among three atoms. These bonds are also called three-center two-electron bonds. The bond between the H and the B atoms can be rationalized using molecular orbital theory.
Other Boron Compounds
Although boron compounds are widely distributed in Earth's crust, a few concentrated ores are located in Italy, Russia, Tibet, Turkey, and California. Borax is the most common ore found, and it can be turned into a variety of boron compounds. When a solution of borax and hydrogen peroxide is crystallized, sodium perborate (NaBO3 * 4 H2O) is formed. Sodium perborate is used in color-safe bleaches. The key to the bleaching ability of this compound is the presence of its two peroxo groups that bridge the boron atoms together. Another compound that other boron compounds can be synthesized from is boric acid (B(OH)3). When mixed with water, the weakly acidic and electron deficient boric acid accepts an OH- ion from water and forms the complex ion [B(OH)4]-.
Borate salts produce basic solutions that are used in cleaning agents. Boric acid is also used as an insecticide to kill roaches, and as an antiseptic in eyewash solutions. Other boron compounds are used in a variety of things, for example: adhesives, cement, disinfectants, fertilizers, fire retardants, glass, herbicides, metallurgical fluxes, and textile bleaches and dye.
Problems
1. What is the electronic configuration of boron?
2. What accounts for the formation of boron hydrides?
3. What are some uses of boron compounds?
4. Draw B4H10.
5. What is the molecular orbital theory and how is it used to rationalize the bonds in boron hydrides?
Answers
1. [He]2s2 2p1
2. Boron is highly electronegative, and wants to form compounds with hydrogen atoms.
3. Adhesives, cement, disinfectants, fertilizers, etc.
4. The molecular orbital theory treats compounds not as having individual bonds between atoms, but as sharing electrons with multiple atoms through their orbitals. In this way, hydrogen atoms are "bonded" between 2 other atoms at a time, in that a pair of electrons is shared between 3 atoms at once.
Contributors and Attributions
• Forogh Rahim (UCD)
• Stephen R. Marsden
Z005 Chemistry of Boron (Z5)
Boranes and the Bonding in boranes
Boranes are compounds consisting of boron and hydrogen. They were investigated systematically by the german scientist Alfred Stock at the beginning of the 19th century. The most basic example is diborane ($\ce{B2H6}$), all boranes are electron-deficient compounds. For $\ce{B2H6}$ usually 14 electrons are needed to form 2c,2e-bonds, but only 12 valence electrons are present. Because of this there are two B-H-B bonds, which have three centers, but only two electrons (3c, 2e bond). This can be interpreted as a molecular orbital that is formed by combining the contributed atomic orbitals of the three atoms. In more complicated boranes not only B-H-B bonds but also B-B-B 3c, 2e-bonds occur. In such a bond the three B-atoms lie at the corners of an equilateral triangle with their sp3 hybrid orbitals overlapping at its center. One of the common properties of boranes is, that they are flammable or react spontaneously with air. They burn with a characteristic green flame. And they are colorless, diamagnetic substances.
Nomenclature
In neutral boranes the number of boron atoms is given by a prefix and the number of Hydrogen-atoms is given in parentheses behind the name. example: $\ce{B5H11}$ -> pentaborane(11), $\ce{B4H10}$ -> tetraborane(10) For ions primarily the number of hydrogen-atoms and than the number of boron-atoms is given, behind the name the charge is given in parentheses. example: $\ce{[B6H6]^{2-}}$ -> hexahydrohexaborat(2-)
Wades rule, Structures of boranes
Wades rule helps to predict the general shape of a borane from its formula.
• count the number of B-H units
• every B-H unit contains 4 valence electrons, but two of them are needed to establish the bond between B and H, thus every B-H unit contributes two electrons to the skeletal electrons.
• every further H-Atom contributes a further electron to the skeletal electrons and
• charge contributes electrons
• the resulting number of electrons has to be divided by two to get the number of skeletal electron pairs within the borane. The general structure is defined by the number of skeletal electron pairs
Formula Skeletal Skeletal electron pairs type
$\ce{[B_{n} H_{n}]^{2-}}$ n+1 closo
$\ce{B_{n} H_{n + 4}}$ n+2 nido
$\ce{B_{n} H_{n + 6}}$ n+3 arachno
$\ce{B_{n} H_{n + 8}}$ n+4 hype
The polyhedra are always made up of triangular faces, so they are called deltahedra. Usually there are three possible structure types:
Closo-boranes
• closed deltahedra without B-H-B 3c,2e-bonds
• thermally stable and moderately reactive.
• example: $\ce{[B5H5]^{2-}}$: the ion builds up a trigonal, bipyramidal polyhedron
Nido-boranes
• closo borane with one corner less and addition of two hydrogen-atoms instead
• B-H-B-bonds and B-B-bonds are possible.
• thermally stability lies between closo- and arachno-boranes.
• example: $\ce{B5H9}$ its structure can be assumed as the octahedral deltahedron of $\ce{[B6H6]^{2-}}$ without one corner tetragonal pyramid
Arachno-boranes
• closo borane deltahedron but with two BH-units removed and two H-atoms added.
• it has to have B-H-B 3c, 2e-bonds.
• thermally unstable at room temperature and highly reactive.
• example: \{\ce{B4H10}\) the structure can be derived from $\ce{[B6H6]^{2-}}$ -> deltahedron with two corners less.
There exist also other structures like the hypho-boranes, but they are less important.
Synthesis of Boranes
Diborane can be synthesized by an exchange reaction (metathesis) of a boron halide with $\ce{LiAlH4}$ or $\ce{LiBH4}$ in ether, for example:
$\ce{3 LiAlH4 + 4 BF3 -> 2 B2H6 + 3 LiAlF4} \nonumber$
The reaction has to be done under vacuum or with exclusion of air, because diborane burns in contact with air. Higher boranes are obtained by controlled pyrolysis of diborane in the gas phase. example:
$\ce{H2B6 (g) -> 2BH3 (g)} \nonumber$
$\ce{B2H6 (g) + BH3(g) -> B3H7(g) + H2(g)} \nonumber$
$\ce{BH3 (g) + B3H7(g) -> B4H10(g)} \nonumber$
Sources
1. D. F. Shriver, P. W. Atkins, Inorganic Chemistry Third edition, Oxford University Press, 2001 | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_13%3A_The_Boron_Family/Z005_Chemistry_of_Boron_%28Z5%29/Boranes_and_Bor.txt |
Aluminum (also called Aluminium) is the third most abundant element in the earth's crust. It is commonly used in the household as aluminum foil, in crafts such as dyeing and pottery, and also in construction to make alloys. In its purest form the metal is bluish-white and very ductile. It is an excellent conductor of heat and electricity and finds use in some wiring. When pure it is too soft for construction purposes but addition of small amounts of silicon and iron hardens it significantly.
Facts
• Symbol: Al
• Atomic Number: 13
• Atomic Weight: 26.98154 amu
• Color: Silver
• Melting Point: 933.4 K
• Boiling Point: 2792 K
• Density: 2.70 g/cm3
• Number Oxidation States: 3
• Great reducing agent
• Has 13 electrons, 13 protons, and 14 neutrons
• Metal
• Good conductor
• Resists corrosion
• Non-magnetic
• Stable ion
• Forms dimers
• Group Number: 13
History of Aluminum
Aluminum ranks third on the list of the ten most abundant elements in the earth's crust, while its oxide is fourth among the ten most common compounds in the crust. It is the most abundant metal on the planet. Its name is taken from the Latin alumen for alum. Soft, lightweight and silvery, its existence was proposed by Lavoisier in 1787, it was named by Davy in 1807 and finally isolated by Ørsted in 1825. Before this, it was known for being part of alum, which is used as a mordant to help set dye on fabric. At this time it was known as a very expensive metal. In the late 1800s, two scientists, Charles Martin Hall and Paul L. T. Heroult, found that they could produce aluminum from aluminum oxide through electrolysis and a cryolite (molten mineral) solvent. This allowed the price to decrease and for aluminum to become available for commercial use.
Aluminum on Earth
Aluminum is the third most abundant element found on earth, and the most abundant metal. It makes up 8.1% of the earth's crust by mass, following oxygen and silicon. Naturally, it is found in chemical compounds with other elements like bauxite. It is not easily removed from natural ores because it must first be reduced. To see how alumina, which is used to make aluminum, is extracted from bauxite, read the Bayer Process in the refining aluminum section.
Electron Configuration of Aluminum
To find the electron configuration of an atom, you first need to know the number of electrons that it has. Since aluminum's atomic number is thirteen, it has thirteen electrons. You then split the electrons between the different orbitals. Aluminum's first two electrons fall in the 1s orbital, and the following two electrons go in the 2s orbital. The next six electrons fill the 2p orbital in the second shell (that's ten electrons so far, three more to go). Then electrons 11 and 12 fill the 3s orbital. Finally the last electron occupies the 3p orbital.
The electron configuration for Aluminum is 1s22s22p63s23p1. The ground state electron configuration is [Ne]3s23p1.
Oxidation States
Aluminum has three oxidation states. The most common one is +3. The other two are +1 and +2. One +3 oxidation state for Aluminum can be found in the compound aluminum oxide, Al2O3. In AlO, aluminum monoxide, it has a +2 oxidation state, and AlH has an oxidation state of +1.
Aluminum Compounds
Although it does not seem to be particularly reactive, aluminum is considered an active metal. Its behavior is deceptive because it reacts rapidly with the oxygen in the air to form aluminum oxide ($Al_2O_3$), or alumina, which is tightly bound to the metal and exists as a dense coating (unlike the oxides of iron). This coating protects it from further reaction. Clearly, however, this coating is not entirely foolproof since aluminum does not exist in native form.
Alumina is the refractory oxide of aluminum and is found in bauxite and corundum (sapphires and rubies). It has a very high melting point. One of the applications of this compound is used to produce different color light that can be used as a laser beam. It is also used in pottery, dyeing, antacid medicines, and in making chemicals.
Another compound containing aluminum is Al(OH)3, which is usually formed as a gelatinous precipitate when aluminum compounds are hydrolyzed in water.
Aluminum sulfate, Al2(SO4)3·18H2O is a very useful aluminum compound, made from the oxide and sulfuric acid. One use of this salt is in the dyeing of cotton fabrics.
Aluminum Reactions
Aluminum is easily oxidized to Al3+such as in this equation:
$2Al_{(s)} + 6H^+_{(aq)} \rightarrow 2Al^{3+} +3H_{2(g)} \nonumber$
In the welding of large objects, the thermite reaction is used:
$2Al_{(s)} + Fe_2O_{3(s)} \rightarrow Al_2O_{3(s)} + Fe_{(s)} \nonumber$
Reactions with Halogens
Aluminum Halides, like the boron halides, are reactive Lewis Acids, meaning that they readily accept a pair of electrons. For example an important halide complex for the production with aluminum is cryolite, NaAlF6.
$6 HF +Al(OH)_3 + 3NaOH \rightarrow Na_3AlF_6 +6 H_2O \nonumber$
Aluminum Oxide and Hydroxide
Aluminum oxide is often referred to as alumina or when crystallized, corundum. Aluminum oxide is relatively unreactive because the small Al3+ ions and the O2+ form a very stable ionic lattice in cubic closet structure with the ions occupying small octahedral holes. Aluminum is protected against corrosion due to thin coating of Al2O3which prevent further oxidation of the aluminum metal.
$2Al_{(s)} +3H_2O_{(l)} \rightarrow Al_2O_{3(s)} +6H^+ +6e^- \nonumber$
Aluminum hydroxide is amphoteric which means that it can react with acids or bases.
$\text{Acid:}\; Al(OH)_{3 (s)} +3H_3O^+_{(aq)} \rightarrow [Al(H_2O)_6]^{3+}_{(aq)} \nonumber$
$\text{Base:}\; Al(OH)_{3 (s)} +OH^-_{(aq)} \rightarrow [Al(OH)_4]^-_{ (aq)} \nonumber$
Refining Aluminum
Most of the aluminum today is produced by the Hall process which uses significant amounts of energy in the form of electricity to electrolyze aluminum metal from a molten salt mixture. The large initial outlay of energy is one important reason why recycling aluminum is such a good and cost-effective idea.
Since aluminum is found in compounds with other elements it needs to be reduced. The Bayer process was invented by Karl Bayer in 1887. It is essentially referring to the refining of bauxite, the most important aluminum ore, to produce alumina. From here, the intermediate alumina must be smelted into metallic aluminum through the Hall-Heroult Process.
Problems
1. Write the configuration of Aluminum, assuming that it has lost its valence electrons.
2. What happens to aluminum when it reacts with chlorine?
3. Balance these equations: a) 2Al(s) → Al3+(aq) + e-
b) Al(s) + Pb+ (aq)→ Al3+(aq)+ Pb(s)
4. What is the electron configuration of Aluminum?
5. What is the process by which alumina is extracted from bauxite?
6. Complete and balance the following reactions.
1. Al(OH)3 (s) + OH-(aq) →
2. Al(OH)3 (s) + H+ (aq) →
Solutions
1. 1s2 2s2 2p6, or [Ne]
2. It forms a dimer.
3. a) 2Al(s) → Al3+ (aq) + 3e-
b) 2Al(s) + 3Pb+ (aq) → 2Al3+ (aq) + 3Pb(s)
4. 1s22s22p63s23p1
5. Bayer Process
6.1 Al(OH)3 (s) +OH-(aq) → [Al(OH)4]-(aq)
6.2 Al(OH)3(s) + 3H+(aq) → Al(H2O)3]3+(aq)
Contributors and Attributions
• Lia D'Angelo, Hanna Towers, Daniel Ahrens
Stephen R. Marsden
Z013 Chemistry of Aluminum (Z13)
This page starts by looking at the extraction of aluminum from its ore, bauxite, including some economic and environmental issues. It finishes by looking at some uses of aluminum.
Introduction
Aluminum is too high in the electrochemical series (reactivity series) to extract it from its ore using carbon reduction. The temperatures needed are too high to be economic. Instead, it is extracted by electrolysis. The ore is first converted into pure aluminum oxide by the Bayer Process, and this is then electrolyzed in solution in molten cryolite - another aluminum compound. The aluminum oxide has too high a melting point to electrolyse on its own. The usual aluminum ore is bauxite. Bauxite is essentially an impure aluminum oxide. The major impurities include iron oxides, silicon dioxide and titanium dioxide.
The Bayer Process
Reaction with sodium hydroxide solution
Crushed bauxite is treated with moderately concentrated sodium hydroxide solution. The concentration, temperature and pressure used depend on the source of the bauxite and exactly what form of aluminum oxide it contains. Temperatures are typically from 140°C to 240°C; pressures can be up to about 35 atmospheres.
High pressures are necessary to keep the water in the sodium hydroxide solution liquid at temperatures above 100°C. The higher the temperature, the higher the pressure needed. With hot concentrated sodium hydroxide solution, aluminum oxide reacts to give a solution of sodium tetrahydroxoaluminate.
$Al_2O_3 + 2NaOH + 3H_2O \longrightarrow 2NaAl(OH)_4 \nonumber$
The impurities in the bauxite remain as solids. For example, the other metal oxides present tend not to react with the sodium hydroxide solution and so remain unchanged. Some of the silicon dioxide reacts, but goes on to form a sodium aluminosilicate which precipitates out. All of these solids are separated from the sodium tetrahydroxoaluminate solution by filtration. They form a "red mud" which is just stored in huge lagoons.
Precipitation of aluminum hydroxide
The sodium tetrahydroxoaluminate solution is cooled, and "seeded" with some previously produced aluminum hydroxide. This provides something for the new aluminum hydroxide to precipitate around.
$NaAl(OH)_4 \longrightarrow Al(OH)_3 + NaOH \nonumber$
Formation of pure aluminum oxide
Aluminum oxide (sometimes known as alumina) is made by heating the aluminum hydroxide to a temperature of about 1100 - 1200°C.
$2Al(OH)_3 \longrightarrow Al_2O_3 + 3H_2O \nonumber$
Conversion of the aluminum oxide into aluminum by electrolysis
The aluminum oxide is electrolyzed in solution in molten cryolite, Na3AlF6. Cryolite is another aluminum ore, but is rare and expensive, and most is now made chemically.
The electrolysis cell
The diagram shows a very simplified version of an electrolysis cell.
Although the carbon lining of the cell is labelled as the cathode, the effective cathode is mainly the molten aluminum that forms on the bottom of the cell. Molten aluminum is syphoned out of the cell from time to time, and new aluminum oxide added at the top. The cell operates at a low voltage of about 5 - 6 volts, but at huge currents of 100,000 amps or more. The heating effect of these large currents keeps the cell at a temperature of about 1000°C.
The electrode reactions
These are very complicated - in fact one source I've looked at says that they aren't fully understood. For chemistry purposes at this level, they are always simplified (to the point of being wrong! - see comment below).
This is the simplification:
Aluminum is released at the cathode. Aluminum ions are reduced by gaining 3 electrons.
$Al^3+ + 3e^- \longrightarrow Al \nonumber$
Oxygen is produced initially at the anode.
$2O^{2-} \longrightarrow O_2 + 4e^- \nonumber$
However, at the temperature of the cell, the carbon anodes burn in this oxygen to give carbon dioxide and carbon monoxide. Continual replacement of the anodes is a major expense.
Some economic and environmental considerations
This section is designed to give you a brief idea of the sort of economic and environmental issues involved with the extraction of aluminum. I wouldn't claim that it covers everything! Think about:
• The high cost of the process because of the huge amounts of electricity it uses. This is so high because to produce 1 mole of aluminum which only weighs 27 g you need 3 moles of electrons. You are having to add a lot of electrons (because of the high charge on the ion) to produce a small mass of aluminum (because of its low relative atomic mass).
• Energy and material costs in constantly replacing the anodes.
• Energy and material costs in producing the cryolite, some of which gets lost during the electrolysis.
Environmental problems in mining and transporting the bauxite
Think about:
• Loss of landscape due to mining, processing and transporting the bauxite.
• Noise and air pollution (greenhouse effect, acid rain) involved in these operations.
Extracting aluminum from the bauxite
Think about:
• Loss of landscape due to the size of the chemical plant needed, and in the production and transport of the electricity.
• Noise.
• Atmospheric pollution from the various stages of extraction. For example: carbon dioxide from the burning of the anodes (greenhouse effect); carbon monoxide (poisonous); fluorine (and fluorine compounds) lost from the cryolite during the electrolysis process (poisonous).
• Pollution caused by power generation (varying depending on how the electricity is generated.)
• Disposal of red mud into unsightly lagoons.
• Transport of the finished aluminum.
Recycling
Think about:
• Saving of raw materials and particularly electrical energy by not having to extract the aluminum from the bauxite. Recycling aluminum uses only about 5% of the energy used to extract it from bauxite.
• Avoiding the environmental problems in the extraction of aluminum from the bauxite.
• Not having to find space to dump the unwanted aluminum if it wasn't recycled.
• (Offsetting these to a minor extent) Energy and pollution costs in collecting and transporting the recycled aluminum.
Uses of aluminum
Aluminum is usually alloyed with other elements such as silicon, copper or magnesium. Pure aluminum isn't very strong, and alloying it adds to it strength. Aluminum is especially useful because it
• has a low density;
• is strong when alloyed;
• is a good conductor of electricity;
• has a good appearance;
• resists corrosion because of the strong thin layer of aluminum oxide on its surface. This layer can be strengthened further by anodizing the aluminum.
Anodizing essentially involves etching the aluminum with sodium hydroxide solution to remove the existing oxide layer, and then making the aluminum article the anode in an electrolysis of dilute sulphuric acid. The oxygen given of at the anode reacts with the aluminum surface, to build up a film of oxide up to about 0.02 mm thick. As well as increasing the corrosion resistance of the aluminum, this film is porous at this stage and will also take up dyes. (It is further treated to make it completely non-porous afterwards.) That means that you can make aluminum articles with the colour built into the surface.
Some uses include:
aluminum is used for because
aircraft light, strong, resists corrosion
other transport such as ships' superstructures, container vehicle bodies, tube trains (metro trains) light, strong, resists corrosion
overhead power cables (with a steel core to strengthen them) light, resists corrosion, good conductor of electricity
saucepans light, resists corrosion, good appearance, good conductor of heat | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_13%3A_The_Boron_Family/Z013_Chemistry_of_Aluminum_%28Z13%29/Aluminium_M.txt |
Aluminum oxide, with the chemical formula $Al_2O_3$, is an amphoteric oxide and is commonly referred to as alumina. Corundum (α-aluminum oxide), emery, sapphire, amethyst, topaz, as well as many other names are reflecting its widespread occurrence in nature and industry. Corundum is the most common naturally occurring crystalline form of aluminum oxide. Rubies and sapphires are gem-quality forms of corundum, which owe their characteristic colors to trace impurities. Rubies are given their characteristic deep red color and their laser qualities by traces of chromium. Sapphires come in different colors given by various other impurities, such as iron and titanium.
Its most significant use is in the production of aluminum metal, although it is also used as an abrasive due to its hardness and as a refractory material due to its high melting point.
Properties
Aluminum oxide is an electrical insulator but has a relatively high thermal conductivity ($30\, W m^{-1} K^{-1}$) for a ceramic material. It is thus used as insulating material in power electronics. Aluminum oxide is responsible for resistance of metallic aluminum to weathering. Since metallic aluminum is very reactive with atmospheric oxygen, a thin passivation layer of alumina (4 nm thickness) forms on any exposed aluminum surface, protecting the metal from further oxidation. The thickness and properties of this oxide layer can be enhanced using a process called anodizing.
Production
The production of aluminum oxide is mainly from bauxite (the main aluminum ore), which is a mixture of various minerals including gibbsite ($Al(OH)_3$, boehmite ($\gamma-AlO(OH)$), and diaspore ($\alpha-AlO(OH)$) along with impurities of iron oxides, quartz, and silicates.
Bauxite is purified by the Bayer process which is the principal industrial refining process. As bauxite contains only about 40 to 50% of alumina, the rest has been removed. This is achieved by washing bauxite with hot sodium hydroxide, which dissolves the alumina by converting it to aluminum hydroxide which forms a solution in a strong base:
$Al_2O_3 + 2 OH^- + 3 H_2O \rightarrow 2 [Al(OH)_4]^- \nonumber$
The other components of the bauxite do not dissolve and are filtered off (the residues usually form a red sludge which presents a disposal problem, since it contains, for example, arsen and cadmium(1)). Next the solution is cooled which causes precipitation of a fluffy solid (aluminum hydroxide). The aluminum hydroxide is then heated to 1050°C which causes it to decompose into aluminum oxide and water:
$2 Al(OH)_3 \rightarrow Al_2O_3 + 3 H_2O \nonumber$
Case Study%
Bauxite Ore Processing
Aluminum is found in varying amounts in nature as aluminosilicates (contains aluminum, silicon, and oxygen) in various types of clay. As the minerals are weathered they gradually breakdown into various forms of hydrated aluminum oxide, Al2O3.xH2O, known as bauxite. The bauxite is purified by the Bayer Process. First the ore is mixed with a hot concentrated solution of sodium hydroxide. The NaOH will dissolve the oxides of aluminum and silicon but not other impurities such as iron oxides, which remains insoluble. The insoluble materials are removed by filtration. The solution which now contains the oxides of aluminum and silicon are next treated by bubbling carbon dioxide gas through the solution. Carbon dioxide forms a weak acid solution of carbonic acid which neutralizes the sodium hydroxide from the first treatment. This neutralization selectively precipitates the aluminum oxide, but leaves the silicates in solution. Again filtration is used for the separation. After this stage the purified aluminum oxide is heated to evaporate the water. Aluminum in the metal form is very difficult to obtain by using some of the traditional chemical methods involving carbon or carbon monoxide as reducing agents to reduce the aluminum ions to aluminum metal. One of the earliest and costly methods in 1850 was to reduce aluminum chloride with sodium metal to obtain aluminum metal and sodium chloride. (Sodium metal is not easy to obtain either). As a result some of the earliest aluminum metal was made into jewelry.
Hall-Heroult Process
In 1886, Charles Hall, an American (23 yrs. old), and Paul Heroult, a Frenchmen (23 yrs old), simultaneously and independently developed the process still in use today to make aluminum metal. The purified aluminum oxide is mixed with cryolite, a mixture of sodium fluoride and aluminum fluoride, and heated to about 980 degrees Celsius to melt the solids. The mixture melts at a much lower temperature than aluminum oxide would by itself. The hot molten mixture is electrolyzed at a low voltage of 4-5 volts, but a high current of 50,000-150,000 amps. Aluminum ions are reduced to aluminum metal at the cathode (the sides and bottom of the electrolysis cell). At the anode, oxygen is produced from the oxide ions. The anode material is carbon in the form of graphite, which also is oxidized and must be replaced quite frequently. The electricity used to produce aluminum is relatively high. One pound of aluminum requires 6-8 kilowatt-hours of electrical energy. This amount of aluminum can be used to make 23 pop cans or one 300 watt light bulb burning for one hour is required to make one pop can. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_13%3A_The_Boron_Family/Z013_Chemistry_of_Aluminum_%28Z13%29/Aluminum_Ox.txt |
Gallium is the chemical element with the atomic number 31 and symbol Ga on the periodic table. It is in the Boron family (group 13) and in period 4. Gallium was discovered in 1875 by Paul Emile Lecoq de Boisbaudran. Boisbaudran named his newly discovered element after himself, deriving from the Latin word, “Gallia,” which means “Gaul.” Elemental Gallium does not exist in nature but gallium (III) salt can be extracted in small amounts from bauxite and zinc ores. Also, it is known for liquefying at temperatures just above room temperature.
Introduction
Gallium is one of the elements originally predicted by Mendeleev in 1871 when he published the first form of the periodic table. He dubbed it ekaaluminum, indicating that it should have chemical properties similar to aluminum. The actual metal was isolated and named (from the Latin Gallia, for France) by Paul-Emile Lecoq de Boisbaudran in 1875.
The detective work behind the isolation of gallium depended on the recognition of unexpected lines in the emission spectrum of a zinc mineral, sphalerite. Eventual extraction and characterization followed. Today, most gallium is still extracted from this zinc mineral.
Although once considered fairly obscure, gallium became an important commercial item in the '70s with the advent of gallium arsenide LEDs and laser diodes. At room temperature gallium is as soft as lead and can be cut with a knife. Its melting point is abnormally low and it will begin to melt in the palm of a warm hand. Gallium is one of a small number of metals that expands when freezing.
Basic Chemical and Physical Properties
Atomic Number 31
Atomic Mass 69.723 g/mol
Element Category Post-transition metal
Phase Solid
Electronegativity 1.6 (Pauling Scale)
Density (at 0oC) 5.91 g/cm3
Melting Point 29.7646oC
Boiling Point 2204oC
Atomic Radius 135 pm
Ionic Radius 62 pm
Isotopes 2 (69Ga; 60.11% & 71Ga; 39.89%)
1st ionization energy 578.8 kJ/mol
Electrode Potential -0.56 eo
Electrical Conductivity 9.1
Oxidation States +3,+2, +1
Hardness 1.5 (Mohs) 60 MPa (Brinell)
Crystal Structure Orthorhombic
Specific Heat 25.86 J/molK
Heat of Fusion 5.59 kJ/mol
Heat of Vaporization 254 kJ/mol
Electronic Configuration
1s22s22p63s23p64s23d104p1
[Ar]4s2 3d104p1
Characteristics
Gallium has a few notable characteristics which are summarized below:
• In its solid phase, Gallium is blue-grey in color
• It melts in temperatures warmer than room temperature; therefore, if you were to hold a chunk of gallium in your hand, it will start to liquefy.
• Solid gallium is soft and can easily be cut with a knife.
• It is stable in air and water, but reacts and dissolves in acids and alkalis.
• If solidifying, gallium expands by 3.1 percent and thus storage in glass or metal is avoided.
• It also easily to transform into an alloy with many metals and has been used in nuclear bombs to stabilize the crystal structure.
• Gallium is one of the few metals that can replace the use the mercury in thermometers because its melting point is close to room temperature.
Video 1: the video depicts the solidifying of liquid Gallium in 10x speed. Density of solid Gallium smaller than density of the liquid, so it's expanding during solidification and break the bottle.
Video 2: The video shows Gallium melting in your hands due to its melting point.
Occurrences
Gallium usually cannot be found in nature. It exists in the earth's crust, where its abundance is about 16.9 ppm. It is extracted from bauxite and sometimes sphalerite. Gallium can also be found in coal, diaspore and germanite.
Applications
Health: While Gallium can be found in the human body in very small amounts, there is no evidence for it harming the body. In fact, Gallium (III) salt is used in many pharmaceuticals, used as treatment for hypercalcemia, which can lead to growth of tumors on bones. Further, it has even been suggested that it can be used to treat cancer, infectious disease, and inflammatory disease. However, exposure to large amounts of Gallium can cause irritation in the throat, chest pains, and the fume it produces can lead to very serious conditions.
Semiconductors: Roughly 90-95% of gallium consumption is in the electronics industry. In the United States, Gallium arsenide (GaAs) and gallium nitride (GaN) represent approximately 98% of the gallium consumption. Gallium arsenide (GaAs) can convert light directly into electricity. Further, gallium arsenide is also used in LEDs and transistors.
Other applications of Gallium deal with wetting and alloy improvement:
Gallium has the property to wet porcelain and even glass surfaces. As a result, gallium can be used to create dazzling mirrors. Scientists employ an alloy with Gallium for the plutonium pits of nuclear weapons to stabilize the alloptropes of plutonium. As a result, some have issue with the element.
Problems
1. What is the electronic configuration of Gallium?
2. What do you think is one of the issues that people might have with usage of gallium?
3. Gallium is part of which group and period?
4. What are some applications of Gallium?
5. Name three properties of Gallium that make it different from any other element.
Answers
1. 1s22s22p63s23p64s23d104p1
2. The use of it in nuclear bombs.
3. Gallium is in group 13 (Boron family) and in period 4.
4. Semiconductors; cancer treatment; hypercalcemia treatment; stabilization in nuclear bombs. See section above on Application for more detail.
5. 5. See the section above on properties and characteristics for more detail.
1. Gallium is blue-grey in color in its solid phase.
2. Melts in temperatures warmer than room temperature
3. Stable in air and water, but reacts and dissolves in acids and alkalis.
Contributors and Attributions
• Angela Tang, Sarang Dave
Stephen R. Marsden | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_13%3A_The_Boron_Family/Z031_Chemistry_of_Gallium_%28Z31%29.txt |
Indium is the 49th element, abbreviated as In. Discovered in 1863, indium rarely is found as an isolated element. Alloys of indium have low melting points and are good semiconductors; it's use in LCD displays has recently increased the demand for Indium.
History
The element indium (named from the Latin indicum, for the color indigo) was discovered in 1863 by Reich and Richter. It is a rare metal, with an abundance similar to that of silver. It is generally found in deposits with zinc and refineries which produce this more common metal often sell indium as well. The pure metal is so soft that you can "wipe" it onto other materials in much the same way as lead (or even pencil graphite). It is corrosion resistant.
As with gallium, identification of indium involved the recognition of new emission spectrum lines (its name was chosen because of indigo lines in its spectrum). Curiously enough, Reich who did the initial chemical isolation work was color blind and had to turn over his experiment to an assistant (Richter) who was the first to observe the characteristic lines.
Like pure tin, pure indium emits a squealing sound when bent.
Properties
Indium has the chemical symbol In and the atomic number 49. It has the electron configuration [Kr] 2s22p1 and may adopt the +1 or +3 oxidation state; however, the +3 state is more common. It is a soft, malleable metal that is similar to gallium. Indium forms InAs, which is found in photoconductors in optical instruments. The physical properties of indium include its silver-white color and the "tin cry" it makes when bent. Indium is soluble in acids, but does not react with oxygen at room temperature. It is obtained by separation from zinc ores. Indium is mainly used to make alloys, and only a small amount is required to enhance the metal strength. For example, indium is added to gold or platinum to make the metals more useful industrial tools.
Production
Almost all of Indium is produced as a byproduct of zinc production. A small amount is produced from tin production. Indium is produced from the slag of zinc production (the waste matter separated from metals during the smelting or refining of ore) using a leaching process. Leaching is an technique used to extract metals which converts them to an aqueous state. Indium is leached using \(HCl\) or \(H_2SO_4\). There are many different processes to extract Indium from the zinc slag, and they vary from processor to processor.
Indium recovered from this process is metal of a low-grade. It is further refined to be a high purity metal in refineries. Indium is produced in a lot of different forms, such as foil, ribbon, ingot, plates, powder, shot and pellets, and wire.
Z081 Chemistry of Thalium (Z81)
Sir William Crookes discovered thallium in 1861, positively identifying it by a bright green line in its spectrum (hence the name, which is from the Greek, thallos, for "green twig"). Although in appearance thallium resembles lead, it does not have the corrosion resistance of lead and so has few commercial applications.
Thallium has the chemical symbol Tl and atomic number 81. It has the electron configuration [Xe] 6s26p1 and has a +3 or +1 oxidation state. As stated above, because thallium is heavy, it has a greater stability in the +1 oxidation state (inert pair effect). Therefore, it is found more commonly in its +1 oxidation state. Thallium is soft and malleable.
Thallium compounds are quite toxic and some have been used as rat poisons. A few compounds are used in glasses for special infra-red lenses. Because of its toxicity, thallium was widely used in insecticide and rat poison until this usage was prohibited in 1975 in the U.S.
Contributors and Attributions
Stephen R. Marsden
Z113 Chemistry of Nihonium (Z113)
In studies announced jointly by the Joint Institute for Nuclear Research in Dubna, Russia, and the Lawrence Livermore National Laboratory in the U.S., four atoms of element 113 were produced in 2004 via decay of element 115 after the fusion of Ca-48 and Am-243.
Contributors and Attributions
Stephen R. Marsden | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_13%3A_The_Boron_Family/Z049_Chemistry_of_Indium_%28Z49%29.txt |
Carbon is one of the most common elements on earth, and greatly influences everyday life. Common molecules containing carbon include carbon dioxide (CO2) and methane (CH4). Many scientists in a variety of fields study of carbon: biologists investigating the origins of life; oceanographers measuring the acidification of the oceans; and engineers developing diamond film tools. This article details the periodic properties of the carbon family and briefly discusses of the individual properties of carbon, silicon, germanium, tin, lead, and flerovium.
• Group 14: General Chemistry
Covers the Group 4 (IUPAC: Group 14) chemistry (carbon, silicon, germanium, tin and lead) and specifically the trend from non-metal to metal as you go down the group, and the increasing tendency towards an oxidation state of +2. Also a certain amount of chemistry of the chlorides and oxides.
• Group 14: General Properties and Reactions
Carbon is one of the most common elements on earth, and greatly influences everyday life. This article details the periodic properties of the carbon family and briefly discusses of the individual properties of carbon, silicon, germanium, tin, lead, and flerovium.
• Chemistry of Carbon (Z=6)
Organic chemistry involves structures and reactions of mainly carbon and hydrogen. Inorganic chemistry deal with interactions of all other pure elements besides carbon, amongst geo/biochemistry. So where does inorganic chemistry of carbon fit in? The inorganic chemistry of carbon also known as inorganic carbon chemistry, is the chemistry of carbon that does not fall within the organic chemistry zone.
• Chemistry of Silicon (Z=14)
Silicon, the second most abundant element on earth, is an essential part of the mineral world. Its stable tetrahedral configuration makes it incredibly versatile and is used in various way in our every day lives. Found in everything from spaceships to synthetic body parts, silicon can be found all around us, and sometimes even in us.
• Chemistry of Germanium (Z=32)
Germanium, categorized as a metalloid in group 14, the Carbon family, has five naturally occurring isotopes. Germanium, abundant in the Earth's crust has been said to improve the immune system of cancer patients. It is also used in transistors, but its most important use is in fiber-optic systems and infrared optics.
• Chemistry of Tin (Z=50)
Mentioned in the Hebrew scriptures, tin is of ancient origins. Tin is an element in Group 14 (The carbon family) and has mainly metallic properties. Tin has atomic number 50 and an atomic mass of 118.710 atomic mass units. Tin, or Sn (from the Latin name Stannum) has been known since ancient times, although it could only be obtained by extraction from its ore. Tin shares chemical similarities with germanium and lead. Tin mining began in Australia in 1872 and today Tin is used extensively.
• Chemistry of Lead (Z=82)
Although lead is not very common in the earth's crust, what is there is readily available and easy to refine. Its chief use today is in lead-acid storage batteries such as those used in automobiles. In pure form it is too soft to be used for much else. Lead has a blue-white color when first cut but quickly dulls on exposure to air, forming Pb2O, one of the few lead(I) compounds. Most stable lead compounds contain lead in oxidation states of +2 or +4.
• Chemistry of Flerovium (Z=114)
The synthesis of element 114 was reported in January of 1999 by scientists from the Joint Institute for Nuclear Research in Dubna (near Moscow) and Lawrence Livermore National Laboratory (in California). In an experiment lasting more than 40 days Russian scientists bombarded a film of Pu-244 supplied by Livermore scientists with a beam of Ca-48. One atom of element 114 was detected with a half-life of more than 30 seconds.
Group 14: The Carbon Family
Covers the Group 4 (IUPAC: Group 14) chemistry (carbon, silicon, germanium, tin and lead) and specifically the trend from non-metal to metal as you go down the group, and the increasing tendency towards an oxidation state of +2. Also a certain amount of chemistry of the chlorides and oxides.
• Chemistry of Aqueous Lead(II) Ions
This page discusses the precipitation of insoluble lead(II) compounds from aqueous lead(II) ions in solution. It describes the formation of lead(II) hydroxide, lead(II) chloride, lead(II) iodide and lead(II) sulfate. Because many lead(II) compounds are insoluble, a common source of aqueous lead(II) ions is lead(II) nitrate; this source is assumed in all following examples.
• Chlorides of Group 4 Elements
• Oxidation State Trends in Group 4
This page explores the oxidation states (oxidation numbers) adopted by the Group 4 elements (carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb)). It examines the increasing tendency of the elements to form compounds with +2 oxidation states, particularly for tin and lead.
• Oxides of Group 4 Elements
This page briefly examines the oxides of carbon, silicon, germanium, tin and lead. It concentrates on the structural differences between carbon dioxide and silicon dioxide, and on the trends in acid-base behavior of the oxides down Group 4.
• The Trend from Non-Metal to Metal in Group 4 Elements
1Group 14: General Chemistry
This page discusses the precipitation of insoluble lead(II) compounds from aqueous lead(II) ions in solution. It describes the formation of lead(II) hydroxide, lead(II) chloride, lead(II) iodide and lead(II) sulfate. Because many lead(II) compounds are insoluble, a common source of aqueous lead(II) ions is lead(II) nitrate; this source is assumed in all following examples.
Making lead(II) hydroxide
If a small amount of sodium hydroxide solution is added to colorless lead(II) nitrate solution, a white precipitate of lead(II) hydroxide is produced:
$Pb^{2+} + 2OH^- (aq) \rightarrow Pb(OH)_2(s) \nonumber$
If more sodium hydroxide solution is added, the precipitate redissolves, forming colorless sodium plumbate(II) solution:
$Pb(OH)_2 (s) + 2OH^- (aq) \rightarrow PbO_2^{2-} + 2H_2O \nonumber$
Making lead(II) chloride
Lead(II) chloride, a white precipitate, is formed by adding a chloride ions (in dilute hydrochloric acid) to lead(II) nitrate solution. The chemical equation is shown below:
$Pb^{2+}(aq) + 2Cl^- (aq) \rightarrow PbCl_2 (s) \nonumber$
Adding excess concentrated hydrochloric acid dissolves lead(II) chloride by forming soluble, complex ions such as PbCl42-.
Making lead(II) iodide
If you add colorless potassium iodide solution (or any other source of iodide ions in solution) to a solution of lead(II) nitrate, a bright yellow precipitate of lead(II) iodide is produced.
$Pb^{2+}(aq) + 2I^- (aq) \rightarrow PbI_2(s) \nonumber$
Making lead(II) sulfate
Adding aqueous sulfate ions to a solution of lead(II) nitrate results in a white precipitate of lead(II) sulfate. The most convenient source of sulfate ions is dilute sulfuric acid. The equation is given below:
$Pb^{2+} (aq) + SO_4^{2-} (aq) \rightarrow PbSO_4(s) \nonumber$
Contributors and Attributions
Jim Clark (Chemguide.co.uk) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14%3A_The_Carbon_Family/1Group_14%3A_General_Chemistry/Chemistry_of_Aqu.txt |
This page briefly examines the tetrachlorides of carbon, silicon, and lead, as well as lead(II) chloride. It considers the compounds' structures, stability, and reactions with water.
Structures
Carbon, silicon and lead tetrachlorides
Each of these compounds has the formula XCl4. They are simple covalent molecules with a typical tetrahedral shape. They are liquids at room temperature (although at room temperature, lead(IV) chloride will tend to decompose to give lead(II) chloride and chlorine gas—see the discussion below).
Lead(II) chloride, PbCl2
Lead(II) chloride is a white solid, melting at 501°C. It is slightly soluble in cold water, but its solubility increases with temperature. Lead(II) chloride is essentially ionic in character.
Stability
At the top of Group 4, the most stable oxidation state is +4. This is the oxidation state of carbon and silicon in CCl4 and SiCl4. These compounds have no tendency to break down into dichlorides. However, the relative stability of the +4 oxidation state decreases down the group, and the +2 oxidation state becomes the most stable for lead and below. Lead(IV) chloride decomposes at room temperature to form the more stable lead(II) chloride and chlorine gas.
Reaction with water (hydrolysis)
Carbon tetrachloride (tetrachloromethane)
Carbon tetrachloride has no reaction with water. When added to water, it forms a separate layer underneath the layer of water. If a water molecule were to react with carbon tetrachloride, the oxygen atom in the water molecule would need to attach itself to the carbon atom via the oxygen's lone pair. A chlorine atom would be displaced the process. There are two problems with this idea.
First, chlorine atoms are so bulky and the carbon atom so small that the oxygen atom is sterically hindered from attacking the carbon atom.
Even if this were possible, there would be considerable cluttering around that carbon atom before the chlorine atom breaks away completely, causing a lot of repulsion between the various lone pairs on all the atoms surrounding the carbon, as shown below:
This repulsion makes the transition state very unstable. An unstable transition state indicates a high activation energy for the reaction.
The other problem is that there is no appropriate empty carbon orbital the oxygen lone pair can occupy.
If it attaches before the chlorine starts to break away, there would be an advantage. Forming a bond releases energy, and that energy would be readily available for breaking a carbon-chlorine bond. In the case of a carbon atom, however, this is impossible.
Silicon tetrachloride
The situation is different with silicon tetrachloride. Silicon is larger, so there is more room for the water molecule to attack; the transition is less cluttered. Silicon has an additional advantage: there are empty 3d orbitals available to accept a lone pair from the water molecule. Carbon lacks this advantage because there are no empty 2-level orbitals available.
The oxygen atom can therefore bond to silicon before a silicon-chlorine bond breaks, makes the whole process easier energetically. In practice, silicon tetrachloride therefore reacts violently with water, forming white solid silicon dioxide and HCl gas.
$SiCl_4 + 2H_2O \rightarrow SiO_2 + 4HCl \nonumber$
Liquid SiCl4 fumes in moist air for this reason—it reacts with water vapor in the air.
Lead tetrachloride (lead(IV) chloride)
The reaction of lead(IV) chloride with water is just like that of silicon tetrachloride. Lead(IV) oxide is produced as a brown solid, and fumes of hydrogen chloride given off (this can be confused with the decomposition of the lead(IV) chloride, which gives lead(II) chloride and chlorine gas as mentioned above).
$PbCl_4 + 2H_2O \rightarrow PbO_2 + 4HCl \nonumber$
Lead(II) chloride
Unlike the tetrachlorides, lead(II) chloride can be considered ionic in nature. It is slightly soluble in cold water, but more soluble in hot water. Water solubility involves disruption of the ionic lattice and hydration of the lead(II) and chloride ions to give Pb2+(aq) and Cl-(aq).
Contributors and Attributions
Jim Clark (Chemguide.co.uk) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14%3A_The_Carbon_Family/1Group_14%3A_General_Chemistry/Chlorides_of_Gro.txt |
This page explores the oxidation states (oxidation numbers) adopted by the Group 4 elements (carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb)). It examines the increasing tendency of the elements to form compounds with +2 oxidation states, particularly for tin and lead.
Some examples of the trends in oxidation states
The typical oxidation state adopted by elements in Group 4 is +4, as in CCl4, SiCl4 and SnO2.
CH4, however, is not an example of carbon with an oxidation state of +4. Because carbon is more electronegative than hydrogen, its oxidation state is -4.
However, down the group, there are more examples of +2 oxidation states, such as SnCl2, PbO, and Pb2+. Tin's +4 state of is still more stable than its +2 state, but for lead and heavier elements, the +2 state is the more stable; it dominates the chemistry of lead.
An example from carbon chemistry
The only common example of carbon in a +2 oxidation state is carbon monoxide, CO. Carbon monoxide is a strong reducing agent because it is easily oxidized to carbon dioxide, which has a more thermodynamically stable oxidation state of +4. For example, carbon monoxide reduces many hot metal oxides to elemental metals; this reaction has many useful applications, one of which is the extraction of iron in a blast furnace.
Examples from tin chemistry
For tin and below, the +2 state is increasingly common, and there is a variety of both tin(II) and tin(IV) compounds. However, tin(IV) is the more stable oxidation state; it is therefore fairly easy to convert tin(II) compounds into tin(IV) compounds. This is best illustrated in that Sn2+ ions in solution are strong reducing agents.
A solution containing tin(II) ions (solvated tin(II) chloride, for example) reduces iodine to iodide ions. In the process, the tin(II) ions are oxidized to tin(IV) ions.
Tin(II) ions also reduce iron(III) ions to iron(II) ions: tin(II) chloride reduces iron(III) chloride to iron(II) chloride in solution. In the process, the tin(II) ions are oxidized to the more stable tin(IV) ions.
In addition, tin(II) ions are easily oxidized by powerful oxidizing agents like acidified potassium manganate(VII) (potassium permanganate). This reaction is used in a titration determination of the concentration of tin(II) ions in solution.
As a final example, in organic chemistry, tin and concentrated hydrochloric acid are traditionally used to reduce nitrobenzene to phenylamine (aniline). Tin is first oxidized to tin(II) ions and then further to preferred tin(IV) ions.
Examples from lead chemistry
With lead, the situation is reversed. The lead(II) oxidation state is the more stable; there is a strong tendency for lead(IV) compounds to react, forming lead(II) compounds. Lead(IV) chloride, for example, decomposes at room temperature to give lead(II) chloride and chlorine gas:
Lead(IV) oxide decomposes on heating to give lead(II) oxide and oxygen:
Lead(IV) oxide also reacts with concentrated hydrochloric acid, oxidizing chloride ions in the acid to chlorine gas. Once again, lead is reduced from the +4 to the more stable +2 state.
An explanation for the trends in oxidation states
There is nothing unusual about the stability of the +4 oxidation state in Group 4. Each of the elements in the group has the outer electronic structure ns2npx1npy1, where n is the period number, varying from 2 (for carbon) to 6 (for lead). In an oxidation state of +4, all valence electrons are directly involved in bonding.
Closer to the bottom of the group, there is an increasing tendency for the s2 pair to be uninvolved in bonding. This is often known as the inert pair effect, and is dominant in lead chemistry. There are two different explanations for this, depending on whether the formation of ionic or covalent bonds is in question.
The inert pair effect in the formation of ionic bonds
If the elements in Group 4 form 2+ ions, they lose their p electrons, leaving the s2 pair unused. For example, to form a lead(II) ion, lead loses its two 6p electrons, but the 6s electrons are left unchanged, an "inert pair".
Ionization energies usually decrease down a group as electrons get further from the nucleus. This is not the case in Group 4. This first chart shows how the total ionization energy needed to form a 2+ ion varies down the group. Values are given in kJ mol-1.
Notice the slight increase between tin and lead. This indicates that it is more difficult to remove the p electrons from lead than from tin.
However, examining the pattern for the loss of all four electrons in the chart below, this discrepancy between tin and lead is much more apparent. The relatively large increase between tin and lead is due to the greater difficulty in removing the 6s2 pair in lead than the corresponding 5s2 pair in tin.
(Again, the values are all in kJ mol-1, and the two charts are on approximately the same scale.)
These effects are due to the Theory of Relativity. Heavier elements such as lead experience a relativistic contraction of the electrons that draws the electrons closer to the nucleus than expected. Because they are closer, they are more difficult to remove. The heavier the element, the greater this effect becomes. This affects s electrons to a greater degree than p electrons.
In lead, the relativistic contraction makes it energetically more difficult to remove the 6s electrons than expected. The energy releasing terms when ions are formed (like lattice enthalpy or hydration enthalpy) cannot compensate for this extra energy. Therefore, it makes no energetic sense for lead to form 4+ ions.
The inert pair effect in the formation of covalent bonds
Carbon normally forms four covalent bonds rather than two. Using the electrons-in-boxes notation, the outer electronic structure of carbon looks like this:
There are only two unpaired electrons. Before carbon forms bonds, however, it normally promotes an s electron to the empty p orbital.
This leaves 4 unpaired electrons which (after hybridization) can go on to form 4 covalent bonds.
It is worth supplying the energy to promote the s electron, because the carbon can then form twice as many covalent bonds. Each covalent bond formed releases energy, and this is more than enough to supply the energy needed for the promotion.
One possible explanation for the reluctance of lead to do the same lies in decreasing bond energies down the group. Bond energies decrease as atoms get bigger and the bonding pair is further from the two nuclei and better screened from them.
For example, the energy released when two extra Pb-X bonds (where X is H or Cl or whatever) are formed may no longer be enough to compensate for the extra energy needed to promote a 6s electron into the empty 6p orbital. This would effect is amplified if the energy gap between the 6s and 6p orbitals is increased by the relativistic contraction of the 6s orbital.
Contributors and Attributions
Jim Clark (Chemguide.co.uk) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14%3A_The_Carbon_Family/1Group_14%3A_General_Chemistry/Oxidation_State_.txt |
This page briefly examines the oxides of carbon, silicon, germanium, tin and lead. It concentrates on the structural differences between carbon dioxide and silicon dioxide, and on the trends in acid-base behavior of the oxides down Group 4.
The structures of carbon dioxide and silicon dioxide
The physical properties of carbon dioxide differ significantly from those of silicon dioxide (also known as silicon(IV) oxide or silica). Carbon dioxide is a gas whereas silicon dioxide is a hard, high-melting solid. The other dioxides in Group 4 are also solids, making the structure of carbon dioxide the anomaly.
The structure of carbon dioxide
The fact that carbon dioxide is a gas indicates that it consists of small, simple molecules. Carbon can form these molecules because it can form double bonds with oxygen.
None of the other elements in Group 4 form double bonds with oxygen, so their oxides adopt completely different structures. When carbon forms bonds with oxygen, it promotes one of its 2s electrons into the empty 2p level. This produces 4 unpaired electrons.
These electrons are rearranged by hybridizing the 2s electron and one of the 2p electrons to make two sp1 hybrid orbitals of equal energy. The other 2p electrons are unaffected during this process.
The figure below illustrates this:
Notice that the two green lobes are two different hybrid orbitals, arranged as far from each other as possible (although the two hybrid orbitals have an arrangement similar to a p orbital, it is important not to confuse the two).
Oxygen's electronic structure is 1s22s22px22py12pz1, and its orbitals must also hybridize. In this case, sp2 hybrids are formed from the s orbital and two of the p orbitals, rearranging to form 3 orbitals of equal energy, leaving a temporarily unaffected p orbital.
As shown below, two of the sp2 hybrid orbitals contain lone pairs of electrons.
In the figure below, the carbon and oxygen atoms are arranged in pre-bonding position:
The green hybrid orbitals overlap end-to-end, forming covalent bonds. These are called sigma bonds, and are shown as orange in the next diagram. The sigma bonding brings the p orbitals close enough to overlap.
This overlap between the two sets of p orbitals produces two $pi$ bonds, similar to the $pi$ bond found in ethene. These $pi$ bonds are twisted at 90° to each other in the final molecule.
To form a carbon-oxygen double bond, it is necessary for the lobes of the p orbitals on the carbon and the oxygen to overlap correctly.
The structure of silicon dioxide
Silicon does not double bond with oxygen. Because silicon atoms are larger than carbon atoms, silicon-oxygen bonds are longer than carbon-oxygen bonds. Consider a hypothetical silicon-oxygen double bond, analogous to the carbon-oxygen double bond discussed above. Because silicon-oxygen bonds are longer than carbon-oxygen equivalents, the p orbitals on silicon and oxygen cannot overlap enough to form a stable pi bond. Therefore, only single bonds are formed.
There are several structures for silicon dioxide. One of the simplest is shown below:
This is similar to the structure of diamond, with each of the silicon atoms bridged to its four neighbors via an oxygen atom, forming a large network covalent structure. Strong bonds in three dimensions make silicon dioxide a hard, high melting point solid.
The acid-base behavior of the Group 4 oxides
The oxides of the elements at the top of Group 4 are acidic, but this acidity decreases down the group. Toward the bottom of the group the oxides are more basic, but do not lose their acidic character completely. A compound with both acidic and basic properties is called amphoteric. The trend, therefore, ranges from acidic oxides at the top of the group toward amphoteric oxides at the bottom.
Carbon and silicon oxides
Carbon monoxide
Carbon monoxide is usually treated as a neutral oxide, but it is slightly acidic. It does not react with water, but it can react with hot concentrated sodium hydroxide solution to give a solution of sodium methanoate.
$NaOH + CO \rightarrow HCOONa \nonumber$
The reaction of carbon monoxide with basic hydroxide ions displays its acidic character.
Dioxides of Carbon and Silicon
These compounds are both weakly acidic.
With water
• Silicon dioxide does not react with water, due to the energetic difficulty of breaking up the giant covalent structure.
• Carbon dioxide reacts slightly with water to produce hydrogen ions (strictly hydronium ions) and bicarbonate ions.
The chemical equation for CO2 is given below:
$H_2O (l) + CO_2(aq) \rightleftharpoons H^+(aq) + HCO_3^- (aq) \nonumber$
The solution of carbon dioxide in water is sometimes known as carbonic acid, but in fact only about 0.1% of the carbon dioxide actually reacts. The equilibrium lies well to the left.
With bases
At low temperatures, carbon dioxide reacts with sodium hydroxide to form solutions of either sodium carbonate or sodium bicarbonate, depending on the proportions used. Equations for these reactions are given below:
$2NaOH + CO_2 \rightarrow Na_2CO_3 + H_2O \nonumber$
$NaOH + CO_2 \rightarrow NaHCO_3 \nonumber$
At high temperatures and concentrations, silicon dioxide also reacts with sodium hydroxide, forming sodium silicate as shown:
$2NaOH + SiO_2 \rightarrow Na_2SiO_3 + H_2O \nonumber$
Another familiar reaction occurs in the Blast Furnace extraction of iron; in this process, calcium oxide (from limestone, one of the raw materials) reacts with silicon dioxide to produce a liquid slag of calcium silicate. This is another example of acidic silicon dioxide reacting with a base.
$CaO(s) + SiO_2(s) \rightarrow CaSiO_3 (l) \nonumber$
Germanium, tin and lead oxides
The monoxides
These elements form amphoteric oxides.
The basic nature of the oxides
These oxides react with acids to form salts. An example of this reaction, using hydrochloric acid, is given below:
$XO(s) + 2HCl(aq) \rightarrow XCl_2(aq) + H_2O (l) \nonumber$
where \ (X\) represents Ge or Sn; a modification is required for lead.
When lead(II) oxide is mixed with hydrochloric acid, an insoluble layer of lead(II) chloride forms over the oxide; this stops the reaction from proceeding. A more accurate chemical equation from that given above is the following:
$PbO(s) + 2HCl(aq) \rightarrow PbCl_2 (s) + H_2O \nonumber$
However, if a high enough concentration of HCl is used, the large excess of chloride ions can react with lead(II) chloride to produce soluble complexes such as PbCl42-. These ionic complexes are soluble in water, so the reaction continues. The equation for this process is given below:
$PbCl_2 (s) + 2Cl^-(aq) \rightarrow PbCl_4^{2-} (aq) \nonumber$
The acidic nature of the oxides
These oxides also react with bases like sodium hydroxide, without exception, as follows:
$XO(s) + 2OH^-(aq) \rightarrow XO_2^{2-}(aq) + H_2O(l) \nonumber$
Lead(II) oxide, for example, reacts to form PbO22- (plumbate(II)) ions.
The dioxides
These dioxides are amphoteric, as mentioned before.
The basic nature of the dioxides
These dioxides react first with concentrated hydrochloric acid to form compounds of the type XCl4:
$XO_2 + 4HCl \rightarrow XCl_4 + 2H_2O \nonumber$
These XCl4 compounds react with excess chloride ions in the hydrochloric acid to form complexes such as XCl62-.
$XCl_4 + 2Cl^- \rightarrow XCl_6^{2-} \nonumber$
In the case of lead(IV) oxide, the reaction requires ice-cold hydrochloric acid. If the reaction is carried out at warmer temperatures, the lead(IV) chloride decomposes to give lead(II) chloride and chlorine gas. This is because the preferred oxidation state of lead is +2 rather than +4.
The acidic nature of the dioxides
The dioxides react with hot concentrated sodium hydroxide, forming soluble complexes of the form [X(OH)6]2-.
$XO_2(s) + 2OH^-(aq) + 2H_2O (l) \rightarrow [X(OH)_6]^{2-} (aq) \nonumber$
Some sources suggest that lead(IV) oxide requires molten sodium hydroxide to react. In that case, the equation is modified as follows:
$PbO_2 (s) + 2NaOH (l) \rightarrow Na_2PbO_3(s) + H_2O (g) \nonumber$ | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14%3A_The_Carbon_Family/1Group_14%3A_General_Chemistry/Oxides_of_Group_.txt |
This page explores the trend from non-metallic to metallic behavior in the Group 4 elements (carbon (C), silicon (Si), germanium (Ge), tin (Sn) and lead (Pb)). It describes how this trend is evident in the structures and physical properties of the elements, and attempts to explain the trend.
Structures of the elements
The trend from non-metal to metal down the group is evident in the structures of the elements themselves. Carbon, at the top of the group, forms large network covalent structures in its two most familiar allotropes: diamond and graphite. Diamond has a three-dimensional structure of carbon atoms each bonded covalently to 4 other atoms. This diagram shows a representative portion of that structure:
This structure is also found in silicon and germanium and in one of the allotropes of tin, "grey tin" or "alpha-tin". The more common allotrope of tin ("white tin" or "beta-tin") is metallic, with its atoms held together by metallic bonds. The structure is a distorted close-packed arrangement. In a close-packed structure, each atom is surrounded by 12 neighboring atoms.
In lead and the heavier elements, the atoms are arranged in a 12-coordinated metallic structure.
From this information, it is clear that there is a trend from the typical covalency found in non-metals to the metallic bonding in metals, with an obvious inflection point between the two common tin allotropes.
Physical properties of the elements
Melting points and boiling points
If the trends in melting and boiling points down Group 4 are examined, it is difficult to comment on the shift from covalent to metallic bonding. The trends reflect the increasing weakness of the covalent or metallic bonds as the atoms get bigger and the bonds get longer. This trend is shown below:
The low value for tin's melting point compared with that of lead is presumably due to the distortion in tin's 12-coordinated structure. The tin values in the chart refer to metallic white tin.
Brittleness
A much clearer distinction between nonmetals and metals is shown when the brittleness of the elements is considered.
• Carbon in its diamond allotrope is very hard, reflecting the strength of the covalent bonds. However, if a diamond is hit with a hammer, it shatters.
• Silicon, germanium and grey tin (all with the same structure as diamond) are also brittle solids.
• However, white tin and lead have metallic structures. The atoms can move around without any permanent disruption of the metallic bonds; this leads to typical metallic properties like malleability and ductility. Lead in particular is fairly soft.
Electrical conductivity
• Diamond does not conduct electricity. In diamond the electrons are all tightly bound and not free to move.
• Silicon, germanium and grey tin are semiconductors.
• White tin and lead are metallic conductors.
This information shows clear trend between the typically non-metallic conductivity behavior of diamond, and the typically metallic behavior of white tin and lead.
Explaining the trends
One important characteristic of metals is that they form positive ions. This section examines factors which increase the likelihood of positive ions being formed down Group 4.
Electronegativity
Electronegativity measures the tendency of an atom to attract a bonding pair of electrons. It is usually measured on the Pauling scale, in which the most electronegative element (fluorine) is assigned an electronegativity of 4. The lower the electronegativity of an atom, the less strongly the atom attracts a bonding pair of electrons. That means that this atom will tend to lose the electron pair towards whatever else it is attached to. The atom we are interested in will therefore tend to carry either a partial positive charge or form a positive ion.
Metallic behavior is usually associated with low electronegativity. The trend in electronegativity in Group 4, and its implications for metallic behavior, can be examined using the figure below:
Electronegativity clearly decreases between carbon and silicon, but beyond silicon there is no definite trend. There therefore seems to be no relationship between the non-metal to metal trend and electronegativity values.
Ionization energies
When considering the formation of positive ions, a good start includes describing how ionization energies change down Group 4. Ionization energy is defined as the energy required to carry out each of the following changes (reported in kJ mol-1):
First ionization energy:
$X(g) \rightarrow X^+(g) + e^- \nonumber$
Second ionization energy:
$X^+(g) \rightarrow X^{2+}(g) + e^- \nonumber$
and so on for subsequent ionizations.
None of the Group 4 elements form 1+ ions, so looking at the first ionization energy alone is not helpful. Some of the elements do, however, form 2+ and (to some extent) 4+ ions. The first chart shows how the total ionization energy needed to form the 2+ ions varies down the group. The values are all reported in kJ mol-1.
The ionization energies decrease down the group, although there is a slight increase at lead. The trend exists because:
• The atoms are getting bigger because of the extra layers of electrons. The farther the outer electrons are from the nucleus, the less they are attracted; therefore, they are easier to remove.
• The outer electrons are screened from the full effect of the nucleus by the increasing number of inner electrons.
• These two effects outweigh the effect of increasing nuclear charge.
Examining the ionization energy required to form 4+ ions, the pattern is similar, but not as simple, as shown below (values again reported in kJ mol-1):
Large amounts of ionization energy are required to form 2+ ions, and even more energy is required for 4+ ions. However, in each case there is a decrease in ionization energy down the group; this implies that tin and lead could form positive ions. However, there is no indication from these figures that this is likely.
Carbon's ionization energies are so large that there is essentially no possibility of it forming simple positive ions.
Contributors and Attributions
Jim Clark (Chemguide.co.uk) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14%3A_The_Carbon_Family/1Group_14%3A_General_Chemistry/The_Trend_from_N.txt |
Carbon is one of the most common elements on earth, and greatly influences everyday life. Common molecules containing carbon include carbon dioxide (CO2) and methane (CH4). Many scientists in a variety of fields study of carbon: biologists investigating the origins of life; oceanographers measuring the acidification of the oceans; and engineers developing diamond film tools. This article details the periodic properties of the carbon family and briefly discusses of the individual properties of carbon, silicon, germanium, tin, lead, and flerovium.
Introduction
The carbon family, Group 14 in the p-block, contains carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb), and flerovium (Fl). Each of these elements has only two electrons in its outermost p orbital: each has the electron configuration ns2np2. The Group 14 elements tend to adopt oxidation states of +4 and, for the heavier elements, +2 due to the inert pair effect.
Members of this group conform well to general periodic trends. The atomic radii increase down the group, and ionization energies decrease. Metallic properties increase down the group. Carbon is a non-metal, silicon and germanium are metalloids, and tin and lead are poor metals (they conduct heat and electricity less effectively than other metals such as copper).
Despite their adherence to periodic trends, the properties of the carbon family vary greatly. For example, carbon is a non-metal and behaves as such, whereas tin and lead behave entirely as metals. In their elemental solid states, the Group 14 metalloids silicon and germanium act as electrical semiconductors, although silicon is largely non-metallic; their electrical conductivity can be affected in various degrees by doping, or adding of Group 13 or Group 15 elements in varying concentrations to the Group 14 solid matrix. These semiconductor properties have wide application for circuitry components in the electronics industry, such as diodes, transistors, and integrated circuit (IC) chips.
Element Symbol Atomic # Atomic Mass Classification Electron Configuration
Carbon C 6 12.011 Non-metal [He]2s22p2
Silicon Si 14 28.0855 Metalloid [Ne]3s23p2
Germanium Ge 32 72.61 Metalloid [Ar]3d104s24p2
Tin Sn 50 118.710 Metal [Kr]4d105s25p2
Lead Pb 82 207.2 Metal [Xe]4f145d106s26p2
Flerovium Fl 114 287 Metal [Rn]5f146d107s27p2
Carbon
Carbon is the fourth most abundant element on earth. It is of particular interest in organic chemistry, as it is the distinguishing feature of an organic compound. It is also considered the "backbone" of biology, as all life forms on earth are carbon-based. This is due to two important qualities of carbon: its small size and its unique electron configuration. Because carbon atoms are small, their p-orbital electrons overlap considerably and enable π bonds to form. Compare the molecular structures of CO2 and SiO2 below:
$\ce{CO_2}$ has double bonds between carbon and oxygen atoms, whereas $\ce{SiO2}$ has single bonds. The $\ce{CO_2}$ molecule exists freely in the gas phase. The $\ce{SiO_2}$ molecule, by contrast, always exists within a network of covalent bonds.
Carbon's electron configuration of allows it to form very stable bonds with oxygen and hydrogen. These bonds store an enormous amount of energy. The formation (fixation) and breakage (combustion) of these bonds in the carbon cycle facilitate earthly life:
• Carbon fixation: In photosynthesis, plants use energy from the sun and chlorophyll molecules to turn gaseous carbon dioxide from the atmosphere into simple carbohydrates like glucose:
$\ce{6CO2 + 6H2O + energy → C6H12O6 + 6O2} \nonumber$
• Carbon combustion: In aerobic respiration, plants and animals break carbohydrates down into carbon dioxide and water (as shown in the equation below) and use the energy released to fuel biological activities—growth, movement, etc. In addition, the combustion of carbohydrates found in fossil fuels provides energy needed for modern activities.
$\ce{C6H12O6 + 6O2 → 6CO2 + 6H2O + energy} \nonumber$
Next to sulfur, carbon is the element with the most allotropes. Carbon has three main solid state allotropes: graphite, diamond, and fullerenes (the most commonly known of which, buckminsterfullerene, is also known as a "buckyball"). These allotropes differ greatly in structure but are widely used in modern production.
Graphite and a diamond
Graphite has lubricating properties that make it extremely suitable for use in pencils. Because it is made up of planes of six-membered rings that can easily slide past one another, graphite glides easily and is hence used in combination with clay to form pencil "lead." Graphite is also used in a fibrous form for various plastics.
Carbon has very high melting and boiling points. Graphite is the most thermodynamically stable allotrope of carbon under ordinary conditions. In diamond, the more stable allotrope at extreme pressures (105 atm and up), each carbon atom is bonded to four others in a tetrahedral arrangement, resulting in the hardest naturally-occurring substance known. This hardness, combined with a good ability to dissipate heat, makes diamond and diamond film excellent materials in drill bits and other machine parts; however, the highest-quality natural diamonds are used mainly for jewelry, whereas lower-grade diamond or even synthetic diamond is used for industrial purposes.
Fullerenes (named after R. Buckminster Fuller) and nanotubes are a series of carbon allotropes in which carbon rings form more complex forms, including soccerball-like molecules (C60) and tubes resembling cylinders made of chicken wire. Graphene, a single carbon sheet with intriguing electronic properties, is the basis for these allotropes. Fullerenes occur when a certain percentage of hexagonal rings are assembled to form pentagonal rings, causing the sheet to contort into a roughly spherical "Buckyball." A carbon nanotube is simply graphene bent into a cylinder. Some of these allotropes are formed in the decomposition of graphite. Combustion can also yield alternate carbon forms. Heated coal without air forms coke. Similarly heated wood becomes charcoal as more volatile integrands are forced away.
There is a nearly innumerable amount of different carbon compounds, but several inorganic compounds are particularly important. Carbon monoxide (CO) is used for synthesizing other carbon compounds, reducing metal compounds to usable products, and in combination with other gases for fuel. Carbides, compounds of carbon and metals, are used in many industrial processes, often to stabilize other metal structures; calcium carbide is used to fabricate industrial chemical compounds, for example. Carbon disulfide and carbon tetrachloride are powerful solvents, (although since its classification as a carcinogen, CCl4 use has declined). Cyanide behaves similarly to halide ions, forming both a salt and an acid. Hydrocyanic acid (HCN) is a weak acid with an extremely low boiling point (room temperature in fact), and is used in plastic production. A cyanide dimer is called a cyanogen, and it is used in organic syntheses, fumigants, and rocket propellant.
Silicon
Although silicon plays a much smaller role in biology, it is the second most common element in the earth's crust (after oxygen) and is the backbone of the mineral world. It is classified as neither a metal or nonmetal, but a metalloid. Silicon is inert, primarily reacting with halogens. It may have functioned as a catalyst in the formation of the earliest organic molecules (Sadava 62). Plants depend on silicates (such as [SiO4]4-) to hold nutrients in the soil, where their roots can absorb them (Sadava 787). Silicon (primarily in the silica, SiO2, molecule) has been used for millennia in the creation of ceramics and glass. In more recent history, the name "Silicon Valley" attests to the element's importance in the computing industry— if carbon is the backbone of human intelligence, silicon is the backbone of artificial intelligence. Silicon is found in beach sand, and is a major component of concrete and brick.
Germanium
Germanium is a rare element used in the manufacture of semi-conductor devices. The physical and chemical properties of germanium are very similar to those of silicon. The semi-metal is found in coal, ore, and germanite. Germanium is gray-white in color and forms crystal structures.
Tin
Tin is a soft, malleable metal with a low melting point. It has two solid-state allotropes at regular temperatures and pressures, denoted α and β. At higher temperatures (above 13°C), tin exists as white tin, or β-tin, and is often used in alloys. At lower temperatures, tin can transform into gray tin, (α-tin); it loses its metallic properties and turns powdery. This causes the disintegration of items made from white tin alloys that have been exposed to the cold for long periods of time. The pipes in Europe's great pipe organs are a classic victim of this "tin pest." When a crystalline structure is broken, a "tin cry" is heard; this happens when a bar is bent. Gray tin is used to plate iron food cans to prevent them from rusting. Tin is malleable, ductile, and crystalline. It has 27 isotopes, 9 stable and 18 unstable. It is a superconductor at low temperatures. Tin reacts with bases, acid salts, and strong acids. Tin chlorides are good reducing agents and often used to reduce iron ores. Tin fluoride is often the anticavity "fluoride" additive in toothpastes.
Lead
Lead, (also known as plumbate), is similar to tin in that it is a soft, malleable metal with a low melting point. It was formerly widely used in water and sewage pipes, lending its Latin name (plumbum ) to the terms "plumber" and "plumbing." Lead is toxic to humans, especially children. Even low levels of exposure can cause nervous system damage and can prevent proper production of hemoglobin (the molecule in red blood cells responsible for carrying oxygen through the body). Because of this, there has been a concerted effort to reduce public exposure to lead, including an emphasis on using unleaded gasoline and unleaded paint. Lead is typically stable in an oxidation state of +2 or +4. Its oxides have many industrial uses as oxidizing agents, such as cathodes in lead-acid storage cells.
Flerovium
Flerovium (Fl) is also known as Element 114. It was discovered in 1998 by scientists in Dubna. It is radioactive and very short-lived.
Problems
Here are some questions to test your understanding of this material.
1. Recall the metallic properties. What makes tin and lead "poor" metals?
2. What makes graphite such a good material for pencil lead?
3. What makes diamonds so hard?
4. Why is tin used to plate iron cans?
5. Why are +2 and +4 the most common oxidation states of metals in this group?
Solution
1. They do not conduct heat or electricity very well.
2. It is composed of flat sheets, which are weakly bonded to one another, so they easily slide past each other and rub off on paper.
3. Each carbon atom forms bonds with four other carbon atoms in a tetrahedral crystal. This arrangement is extremely strong.
4. The tin plating prevents the iron can from oxidizing (rusting).
5. Because the valence electron configuration is $ns^2np^2$, the atoms tend to lose either all four outer shell electrons (resulting in a charge of +4) or, because of the inert pair effect, they may lose only the s electrons (resulting in a charge of +2).
Contributors and Attributions
• Elizabeth Sproat, Jessica Lin, Vancy Wong
1Group 14: General Properties and Reactions
Summary of Carbon Group (Group IVA) Trends:
1. Stabilization of (+2) oxidation state relative to (+4) of the elements down the group.
The halides, oxides, and sulfides of the M2+ ions become more stable on descending the group. For example SiCl4, SiBr4, and SiI4 are all stable. PbCl4 decomposes at 105 °C and PbI4 does not exist. Similarly the ease of oxidation of the M2+ halides decreases down the column.
PbCl2 may be converted to PbCl4 by heating in a stream of chlorine.
Similarly PbO2 is an oxidizing agent, whereas SnO2, GeO2, and SiO2are not.
The stabilities of the MII and MIV organometallic derivatives of the elements behave differently. Pb(C2H5)4 can be readily stored and is more stable than Pb(C2H5)2, which is not isolable as a solid.
The compounds in the lower oxidation state are in general more ionic, less likely to form molecular structures, the halides are less readily hydrolyzed and the oxides are less acidic.
2. Hydrides and alkyls become less stable down the column.
The M-H and M-C mean bond enthalpies decrease down the column and consequently the hydrides and alkyls become thermodynamically less stable and kinetically more reactive. Carbon, of course, forms a very wide range of hydrides, silicon forms primarily SiH4 and Si2H6 which are spontaneously inflammable. Higher silanes decompose readily to Si2H6. The Si-H bond polarities are opposite to C-H.
Silanes are strong reducing agents. The germanes GeH4, Ge2H6, and Ge3H8 are less flammable than SiH4 and are resistant to hydrolysis. SnH4 decomposes at 0°C to Sn and PbH4 is extremely unstable.
The organometallic derivatives of silicon and germaniu are very similar. They are more reactive than the carbon analogues because the M-C bonds are more polar and the central atom can exand its coordination number more easily. The rates of hydrolysis are in order:
Pb >> Sn >> Ge >Si
Organotin compounds more readily expand their coordination geometries and more readily form cationic species. Organolead compounds decompose readily at 100-200 °C by free radical processes.
3. Catenation
The element-element mean bond enthalpies decrease in the order:
C-C > Si-Si > Ge-Ge > Sn-Sn > Pb-Pb
and therefore the range of ring and polyhedral molecules diminish down the group. Carbon not only forms an extensive range of chain and ring compounds, but also polyhedral molecules such as prismanes, C6H6, and cubane (C8H8). Analogous compounds are known for Si, Ge, and Sn if the hydrogens are replaced by bulky organic substituents.
However, few examples exist for Pb, which form compounds containing the anionic Zintl polyhedral anoin Pb94-analogous to Sn52-.
4. Multiply bonded compounds
The ability of the elements to form multiple bonds diminishes in the series:
C-C > Si-Si > Ge-Ge > Sn-Sn > Pb-Pb
Because the pπ-pπ overlaps become less favorable. This has the following manifestations:
1. The elements below C the allotropes which would structurally resemble graphite which has a delocalized two dimensional π-system are not observed.
2. There are no simple analogues of ethane (C2H4) and ethyne (C2H2) and compounds of SI, Ge, and Sn with multiple bonds may only be isolated when there are bulky organic substituents on the group IV atoms. Furthermore, the Ge and Sn compounds do not have planar geometries.
3. The analogues of CO2 and CS2 have polymeric structures rather than triatomic molecular geometries.
4. The heavier elements do not form analogues of carbides with C22- and C32- multiply bonded ions.
5. Elements become progressively more metallic down the column.
Carbon is especially in its diamond and polyhedral forms is a typical non-metal, silicon is a semiconductor, and tin & lead are typical metals. Although tin has one modification (grey tin) which is isostructural with Ge, Si, and diamond. Lead only occurs in close packed structural forms.
6. The oxides become more basic down the column.
CO2 and SiO2 are acidic oxides, SnO2 is amphoteric, and GeO2 is mainly acidic with slight amphoteric character. The Si-O mean bond enthalpies are particularly large and this leads to a wide range of silicates. In general for the elements below carbon the M-O bonds are sufficiently strong that the oxides are susceptible to hydrolysis.
7. The typical coordination numbers increase down the group.
For carbon the tetradedral geometry predominates unless multiple bonds formed. For the havier elements the tetrahedral geometry is also widespread bu the larger size of the central atoms leads to the formation of compounds with higher coordination numbers, e.g.
SiF5-, Trigonal bipyramidal; SiF62-, Octahedral; SnCl5-, Trigonal bipyramidal; Sn(C6H5)(NO3)2(OP(C6H5)3, Pentgonal bipyramidal (7 coord); Sn(NO3)4, Dodecahedral (8 coord)
The increased facility to achieve the higher coordination numbers is also reflected in the transition from molecular to ppolymeric, e.g. CF4, SiF4, and GeF4 ar e molecular, where as SnF4 amnd PbF4 have infinite lattices based on octahedral metal centers.
Of course all of these compounds with coordination numbers is also reflected in the geometries of the oxo-anions:
trig. planar: CO32- tetrahedral: SiO42- octahedral: Ge(OH)62- Sn(OH)62- Pb(OH)62-
The chlorides of Ge, Sn, and Pb react with aqueous HCl to form the [MCl6]2- anions, where as SiCl4 hydrolyses and CCl4 is unreactive. However, SiF4 does form [SiF6]2- with HF.
8. The heavier elements form a wider range of complexes and more cationic complexes.
Si, Ge, Sn, and Pb all form oxalato-complexes [M(ox)3]2- and cationic complexes [M(acac)3}+.
9. Ease of reduction of halides.
Although C-Cl and Ge-Cl bonds in four valent compounds are reduced to the corresponding hydrides by Zn and HCl, Si-Cl and Sn-Cl bonds are not.
10. SiH4 hydrolyses in the presence of trace amounts of base more readily than CH4, GeH4, and SnH4. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14%3A_The_Carbon_Family/1Group_14%3A_General_Properties_and_Reactions/C.txt |
Carbon is the fourth most abundant element in the known universe but not nearly as common on the earth, despite the fact that living organisms contain significant amounts of the element. Common carbon compounds in the environment include the gases carbon dioxide ($CO_2$) and methane ($CH_4$).
Inorganic Chemistry of Carbon
Inorganic carbon is carbon extracted from ores and minerals, as opposed to organic carbon found in nature through plants and living things. Some examples of inorganic carbon are carbon oxides such as carbon monoxide and carbon dioxide; polyatomic ions, cyanide, cyanate, thiocyanate, carbonate and carbide. Carbon is an element that is unique to itself. Carbon forms strong single, double and triple bonds; therefore, it would take more energy to break these bonds than if carbon were to bond to another element.
For carbon monoxide the reaction is as follows:
$2C_{(s)} + O_2 \rightarrow 2CO_{(g)} \;\;\;\; \Delta H = -110.52\; kJ/mol \;CO \nonumber$
$C_{(s)} + O_2 \rightarrow CO_{2(g)} \;\;\;\; \Delta H = -393.51\; kJ/mol\; CO_2 \nonumber$
CO and CO2 are both gases. CO has no odor or taste and can be fatal to living organisms if exposed at even very small amounts (about a thousandth of a gram). This is because CO will bind to the hemoglobin that carries oxygen in the blood. CO2 will not become fatal unless living organisms are exposed to larger amounts of it, about 15%. CO2 influences the atmosphere and affects the temperature through the greenhouse gas effect. As heat is trapped in the atmosphere by CO2 gases, the Earth's temperature increases. The main source for CO2 in our atmosphere, amongst many, is volcanoes.
Allotropes
Carbon exists in several forms called allotropes. Diamond is one form with a very strong crystal lattice, known as a precious gem from the most ancient records. Another allotrope is graphite, in which the carbon atoms are arranged in planes which are loosely attracted to one another (hence its use as a lubricant). The recently discovered fullerenes are yet another form of carbon.
• Inorganic carbon may come in the form of diamond as a transparent, isotropic crystal. It is the hardest naturally occurring material on this earth. Diamond has four valence electrons, and when each electron bonds with another carbon it creates an sp3-hybridized atom. The boiling point of diamond is 4827°C.
• Unlike diamond, graphite is opaque, soft, dull and hexagonal. Graphite can be used as a conductor (electrodes) or even in pencils. Graphite consists of planes of sp2 hybridized carbon atoms in which each carbon is attached to three other carbons.
• Fullerenes are carbon cages with the formula $C_{2}n$ where $n > 13$. The most abundant fullerene is the spherical C60. Fullerene may contain atoms or molecules inside the cage (endohedral fullerenes) or covalently attached outside (exhedral or adduct fullerenes). The discovery of fullerenes is accredited to Richard Smalley and his team in 1985 at Rice University, who found it after photoablating the surface of graphite with a laser.
Applications
Carbon has a very high melting and boiling point and rapidly combines with oxygen at elevated temperatures. In small amounts it is an excellent hardener for iron, yielding the various steel alloys upon which so much of modern construction depends. An important (but rare) radioactive isotope of carbon, C-14, is used to date ancient objects of organic origin. It has a half-life of 5730 years, but there is only 1 atom of C-14 for every 1012 atoms of C-12 (the usual isotope of carbon). | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14%3A_The_Carbon_Family/Z006_Chemistry_of_Carbon_%28Z6%29.txt |
Silicon, the second most abundant element on earth, is an essential part of the mineral world. Its stable tetrahedral configuration makes it incredibly versatile, and it is used in various ways in our everyday lives. Found in everything from spaceships to synthetic body parts, silicon can be found all around us, and sometimes even in us.
Introduction
The name for silicon is taken from the Latin silex which means "flint". The element is second only to oxygen in abundance in the earth's crust and was discovered by Berzelius in 1824. The most common compound of silicon, \(SiO_2\), is the most abundant chemical compound in the earth's crust; we know it better as common beach sand.
Properties
Silicon is a crystalline semi-metal or metalloid. One of its forms is shiny, grey, and very brittle (it will shatter when struck with a hammer). It is a group 14 element in the same periodic group as carbon, but chemically behaves distinctly from all of its group counterparts. Silicon shares the bonding versatility of carbon, with its four valence electrons, but is otherwise a relatively inert element. However, under special conditions, silicon can be made to be a good deal more reactive. Silicon exhibits metalloid properties, is able to expand its valence shell, and is able to be transformed into a semiconductor, distinguishing it from its periodic group members.
Table 1: Properties of Silicon
Symbol Si
Atomic Number 14
Group 14 (Carbon Family)
Electron Configuration [Ne]3s23p2
Atomic Weight 28.0855 g
Density 2.57 g/mL
Melting Point 1414oC
Boiling Point 3265oC
Oxidation States 4, 3, 2, 1, -1, -2, -3, -4
Electronegativity 1.90
Stable Isotopes 28Si, 29Si, 30Si
Where Silicon is Found
27.6% of the Earth's crust is made up of silicon. Although it is so abundant, it is not usually found in its pure state, but rather its dioxide and hydrates. \(SiO_2\) is silicon's only stable oxide and is found in many crystalline varieties. Its purest form is quartz, but it is also found as jasper and opal. Silicon can also be found in feldspar, micas, olivines, pyroxenes and even in water (Figure 1). In another allotropic form, silicon is a brown amorphous powder most familiar in "dirty" beach sand. The crystalline form of silicon is the foundation of the semiconductor age.
Silicates
Silicon is most commonly found in silicate compounds. Silica is the one stable oxide of silicon, and has the empirical formula SiO2. Silica is not a silicon atom with two double bonds to two oxygen atoms. Silica is composed of one silicon atom with four single bonds to four oxygen molecules (Figure 2).
Silica, i.e. silicon dioxide, takes on this molecular form, instead of carbon dioxide's characteristic shape, because silicon's 3p orbitals make it more energetically favorable to create four single bonds with each oxygen rather than make two double bonds with each oxygen atom. This leads to silicates linking together in -Si-O-Si-O- networks called silicates. The empirical form of silica is SiO2 because, with respect to the net average of the silicate, each silicon atom has two oxygen atoms.
The tetrahedral SiO44- complex (see Figure 3), the core unit of silicates, can bind together in a variety of ways, creating a wide array of minerals. Silicon is an integral component in minerals, just as Carbon is an essential component of organic compounds.
Nesosilicates
In nesosilicates, the silicate tetrahedral does not share any oxygen molecules with other silicate tetrahedrals, and instead balances out its charge with other metals. The core structure of nesosilicate is simply a lone tetrahedral silica unit (Figure 4). The empirical formula for the core structure of a nesosilicate is SiO44-.
Nesosilicates make up the outer fringes of a group of minerals known as olivines.
Sorosilicates
In sorosilicates, two silicate tetrahedrals join together by sharing an oxygen atom at one of their corners. The core structure of a sorosilicate is a pair of silica tetrahedra. (see Figure 5)
The mineral thortveitie is an example of a sorosilicate complex.
Cyclosilicates
In cyclosilicates, three or more silica tetrahedrals share two corners of an oxygen atom. The core structure of a cyclosilicate is a closed ring of silica tetrahedra. (see Figure 6)
The mineral beryl is an example of a cyclosilicate complex.
Inosilicates
Inosilicates are complexes where each tetrahedral shares two corners with another silica tetrahedral to create a single chain (see Figure 7) or three corners to create a double chain (Figure 8). The core structure of an inosilicate is either an infinite single or double chain of silica tetrahedrals.
The mineral group pyroxenes are examples of single chain inosilicates.
Figure 8: The core of a double chain inosilicate
The mineral amphibole is an example of a double chain inosilicate.
Phyllosilicates
Phyllosilicates are silica complexes where each tetradedral shares three corners and creates a sheet of silicon and oxygen. (see Figure 9) The core complex of a phyllosilicate is an infinite sheet of connected silica tetrahedrals.
The mineral talc is an example of a phyllosilicate complex.
Tectosilicates
Tectosilicates are three-dimensional silicate structures. The core structure of a tectosilicate is an infinite network of connected silica tetrahedrals. (see Figure 10)
The mineral quartz is an example of a tectosilicate complex.
Although many silica complexes form network covalent solids, quartz is a particularly good example of a network solid. Silicates in general share the properties of covalent solids, and this affiliated array of properties makes them very useful in modern day industry.
Silanes
Silanes are silicon-hydrogen compounds. Carbon-hydrogen compounds form the backbone of the living world with seemingly endless chains of hydrocarbons. With the same valence configuration, and thus the same chemical versatility, silicon could conceivably play a role of similar organic importance. But silicon does not play an integral role in our day to day biology. One principal reason underlies this.
Like hydrocarbons, silanes progressively grow in size as additional silicon atoms are added. But there is a very quick end to this trend. The largest silane has a maximum of six silicon atoms (see Figure 11).
Hexasilane is the largest possible silane because Si-Si bonds are not particularly strong. In fact, silanes are rather prone to decomposition. Silanes are particularly prone to decomposition via reaction with oxygen. Silanes also have a tendency to swap out their hydrogens for other elements and become organosilanes. (see Figure 12)
Silanes have a variety of industrial and medical uses. Among other things, silanes are used as water repellents and sealants.
Silicones
Silicones are a synthetic silicon compound; they are not found in nature. When specific silanes are made to undergo a specific reaction, they are turned into silicone, a very special silicon complex. Silicone is a polymer and is prized for its versatility, temperature durability, low volatility, general chemical resistance and thermal stability. Silicone has a unique chemical structure, but it shares some core structural elements with both silicates and silanes. (See Figure 13)
Silicone polymers are used for a huge array of things. Among numerous other things, breast implants are made out of silicone.
Silicon Halides
Silicon has a tendency to react readily with halogens. The general formula depicting this is SiX4, where X represents any halogen. Silicon can also expand its valence shell, and the laboratory preparation of [SiF6]2- is a definitive example of this. However, it is unlikely that silicon could create such a complex with any halogen other than fluorine, because six of the larger halogen ions cannot physically fit around the central silicon atom.
Silicon halides are synthesized to purify silicon complexes. Silicon halides can easily be made to give up their silicon via specific chemical reactions that result in the formation of pure silicon.
Applications
Silicon is a vital component of modern day industry. Its abundance makes it all the more useful. Silicon can be found in products ranging from concrete to computer chips.
Electronics
The high tech sector's adoption of the title Silicon Valley underscores the importance of silicon in modern day technology. Pure silicon, that is, essentially pure silicon, has the unique ability of being able to discretely control the number and charge of the current that passes through it. This makes silicon play a role of utmost importance in devices such as transistors, solar cells, integrated circuits, microprocessors, and semiconductor devices, where such current control is a necessity for proper performance. Semiconductors exemplify silicon's use in contemporary technology.
Semiconductors
Semiconductors are unique materials that have neither the electrical conductivity of a conductor nor that of an insulator. Semiconductors lie somewhere in between these two classes, giving them a very useful property. Semiconductors are able to manipulate electric current. They are used to rectify, amplify, and switch electrical signals and are thus integral components of modern day electronics.
Semiconductors can be made out of a variety of materials, but the majority of semiconductors are made out of silicon. But semiconductors are not made out of silicates, or silanes, or silicones; they are made out of pure silicon, that is, essentially pure silicon crystal. Like carbon, silicon can make a diamond-like crystal. This structure is called a silicon lattice. (see Figure 15) Silicon is perfect for making this lattice structure because its four valence electrons allow it to perfectly bond to four of its silicon neighbors.
However, this silicon lattice is essentially an insulator, as there are no free electrons for any charge movement, and is therefore not a semiconductor. This crystalline structure is turned into a semiconductor when it is doped. Doping refers to a process by which impurities are introduced into ultra-pure silicon, thereby changing its electrical properties and turning it into a semiconductor. Doping turns pure silicon into a semiconductor by adding or removing a very, very small number of electrons, thereby making it neither an insulator nor a conductor, but a semiconductor with limited charge conduction. Subtle manipulation of pure silicon lattices via doping generates the wide variety of semiconductors that modern day electrical technology requires.
Semiconductors are made out of silicon for two fundamental reasons. Silicon has the properties needed to make semiconductors, and silicon is the second most abundant element on earth.
Glasses
Glass is another silicon derivate that is widely utilized by modern day society. If sand, a silica deposit, is mixed with sodium and calcium carbonate at temperatures near 1500 degrees Celsius, when the resulting product cools, glass forms. Glass is a particularly interesting state of silicon. Glass is unique because it represents a solid non-crystalline form of silicon. The tetrahedral silica elements bind together, but in no fundamental pattern behind the bonding. (see Figure 16)
The end result of this unique chemical structure is the often brittle, typically optically transparent material known as glass. This silica complex can be found virtually anywhere human civilization is found.
Glass can be tainted by adding chemical impurities to the basal silica structure. (see Figure 17) The addition of even a little Fe2O3 to pure silica glass gives the resultant mixed glass a distinctive green color.
Figure 17: Non-crystalline silica with unknown impurities
Fiber Optics
Modern fiber optic cables must relay data via undistorted light signals over vast distances. To undertake this task, fiber optic cables must be made of special ultra-high purity glass. The secret behind this ultra-high purity glass is ultra pure silica. To make fiber optic cables meet operational standards, the impurity levels in the silica of these fiber optic cables has been reduced to parts per billion. This level of purity allows for the vast communications network that our society has come to take for granted.
Ceramics
Silicon plays an integral role in the construction industry. Silicon, specifically silica, is a primary ingredient in building components such as bricks, cement, ceramics, and tiles.
Additionally, silicates, especially quartz, are very thermodynamically stable. This translates to silicon ceramics having high heat tolerance. This property makes silicon ceramics particularily useful for things ranging from space ship hulls to engine components. (see Figure 18)
Polymers
Silicone polymers represent another facet of silicon's usefulness. They are generally characterized by their flexibility, resistance to chemical attack, impermeability to water, and their ability to retain their properties at both high and low temperatures. This array of properties makes silicone polymers very useful. They are used in insulation, cookware, high temperature lubricants, medical equipment, sealants, adhesives, and even as an alternative to plastic in toys.
Production
As silicon is not normally found in its pure state, it must be chemically extracted from its naturally occurring compounds. The most prevalent form of naturally occurring silicon is silica. Silica is a strongly bonded compound, and it requires a good deal of energy to extract the silicon out of the silica complex. The principal means of extraction is via a chemical reaction at a very high temperature.
The synthesis of silicon is fundamentally a two-step process. First, use a powerful furnace to heat up the silica to temperatures over 1900 degrees celsius, and second, add carbon. At temperatures over 1900 degrees celsius, carbon will reduce the silica compound to pure silicon.
Purification
For some silicon applications, the purity of freshly produced silicon is not satisfactory. To meet the demand for high purity silicon, techniques have been devised to further refine the purity of extracted silicon.
Purification of silicon essentially involves taking synthesized silicon, turning it into a silicon compound that can be easily distilled, and then breaking up this newly formed silicon compound to yield an ultra pure silicon product. There are several distinct purification methods available, but most chemical forms of purification involve both silane and silicon halide complexes.
Trivia
• Silicon is the eighth most abundant element in the universe.
• Silicon was first identified in 1787 but first discovered as an element in 1824.
• Silicon is an important element in the metabolism of plants, but not very important in the metabolism of animals.
• Silicon is harmless to ingest and inject into the body but it is harmful to inhale.
• Silicosis is the name of the disease associated with inhaling too much of the silicon compound silica. It primarily afflicts construction workers.
• Silica is a major chemical component of asbestos.
Problems
Highlight area next to "Ans" to see answer
How many oxides does silicon have, and what are they?
Ans: 1 oxide O2
How does a silicate tetrahedral balance its charge if not bonded with another silicate?
Ans: By bonding to positively charged metals.
Carbon is to organic compounds as silica is to:
Ans: minerals
How big is the largest silicon-hydrogen compound?
Ans: The largest silane is hexasilane, with six silicon atoms and fourteen hydrogens.
Why is silicon important to computers?
Ans: It is used to make semiconductors.
Contributors and Attributions
• Thomas Bottyan (2010), Christina Rabago (2008)
Z014 Chemistry of Silicon (Z14)
Silicates are some of the most abundant minerals on Earth. They are some of the most common of the raw material that takes over 75% of the Earth's crust. A majority of igneous rocks and sedimentary rocks are made of silicate minerals. The most common type of silicate is (SiO4)4-.
There are many different types of silicates. Most of them have a general chemical formula of XxYy(ZzOo)Ww.
• X = +1 or +2 cations
• Y = +2, +3, or +4 cations
• Z = + 3 or +4 cations
• O = oxygen
• W = usually OH-, F-, or Cl-
• x, y, z, o, w = subscript numbers
Some of the subcategories of silicates are the following:
• Nesosilicates
• Sorosilicates
• Cyclosilicates
• Inosilicates
• Phyllosilicates
• Tectosilicates
Nesosilicates
Nesosilicates are made up of units of independent tetrahedrals. Some of the minerals that contain nesosilicates are olivine, garnet, zircon, kyanite, topaz, and staurolite. Olivine is important in the processes of igneous rock forming. It has a general formula of (Mg, Fe)2SiO4. Garnet belongs to the isomorphic group, where it often occurs as dodecahedron crystals such as pyrope, almandine, and grossularite. It is usually found in metamorphic rocks, and is known for being the January birthstone. Zircon, on the other hand, is marketed as gemstone and is oxidized to produce gemstones that are similar to diamonds known as cubic zirconia. Kyanite is a part of a polymorphic group (Al2OSiO4).
Ex. (SiO4)-4 or (Si3O12)-12
Sorosilicates
Sorosilicate is made up of two tetrahedrals shared by an oxygen. Some of the minerals that are classified as sorosilicates are hemimorphite, epidote, and allanite. Hemimorphite is usually found as bladed crystals. Epidote belongs to the isomorphic group, which is important in forming minerals. Lastly, allanite has a metamict structure that is usually black with no cleavage.
ex. (Si2O7)-6
Cyclosilicates
Cyclosilicates are made up of closed ring units of tetrahedrals sharing two oxygen atoms. They are known for their hardness and consist of a variety of gemstones. They also have poor cleavage. Some minerals that are classified as cyclosilicates are beryl, cordierite, and tourmaline. The gemstones that are classified as beryl include emerald (deep green), aquamarine (greenish-blue), and morganite (red). Tourmalines also have a variety of gemstones, which include rubellite (red-pink) and indicolite (dark blue). As for cordierite, it often show dichroism, meaning that it shows different colors at different concentrations.
Ex. (Si6O18)-12
Inosilicates
Inosilicates are made up of continuous double chain units of tetrahedrals, each sharing 2 and 3 oxygens. They include the pyroxene group, which are single chain minerals without hydroxides, and the amphibole group, which are double chains with hydroxides. The pyroxene group has two directional 90 degree cleavages. Some examples are enstatite-ferrosilite, diopside-hedenbergite, augite, and spodumene. As for the amphibole group, it has two directional cleavages at 124-56 degrees. Some examples are tremolite-actinolite and hornblende. Both of these groups are rock-forming minerals.
Ex. (SiO3)-2 or (Si2O6)-4
Phyllosilicates
Phyllosilicates comprise continuous sheet units of tetrahedrals, each sharing 3 oxygen atoms. They include the clay and mica minerals, which are rock-forming minerals. The clay group is made of hydrous aluminum layered silicates. Some examples are kaolinite and talc. On the other hand, the mica group consists of thin sheets and a multitude of ionic substitutions of Al3+ and Si4+. Some examples are muscovite (light color), biotite (black or dark colored), and lepidolite (pink colored and a source of lithium). The serpentine group also belongs to the phyllosilicates. Some examples are serpentine and crysotile.
Ex. (Si2O5)-2 or (Al Si3O10)-5
Tectosilicates
Tectosilicates consist of a continuous framework of tetrahedrals, each sharing all 4 oxygen atoms. Their structure has a great amount of Al-Si substitution. Some of the groups that are classified as tectosilicates are the SiO2 polymorphic group, the K-feldspar polymorphic group, the feldspathoid group, and the zeolite group. The SiO2 polymorphic group has a variety of quartz, such as smoky quartz, amethyst, and jasper. Some minerals in the K-felspar polymorphic group include orthoclase and microcline. Microcline has 1 Pb2+ ion replacing every 2 K1+ ions, showing an omission solid solution and causing a blue-green color in the mineral. The felspathoid group minerals are similar to feldspars but only have two-thirds of the amount of silica; they form a silica-deficient magma. Some examples are leucite and sodalite. Lastly, the zeolite group has hydrous silicates with ionic exchange and absorption properties that can act as water softeners by exchanging Na1+ ions for Ca2+ ions in solution. For example:
$Na_2Al_2Si_3O_{10}-2H_2O \rightarrow CaAl_2Si_3O_{10}-2H_2O. \nonumber$
Ex. (SiO2), (AlSiO4)-1, (Al2Si2O8)-2, or (Al2Si4O12)-2 | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14%3A_The_Carbon_Family/Z014_Chemistry_of_Silicon_%28Z14%29/Silicates.txt |
Learning Objectives
• Predict some chemical reactions for a set of conditions.
• Describe some chemical and physical properties of $\ce{C}$ and $\ce{Si}$.
Group 14 Elements C, Si, Ge, Sn, Pb
Group 14 elements play more important roles in our lives and our civilization than elements of any other group. Thus, every educated person should know something about them.
• Carbon - The element of organic chemistry and life.
• Silicon - The element of information technology.
• Germanium - The element of transistor bases.
• Tin and lead - Elements known to alchemists.
Chemistry of Carbon
Carbon exists as diamond, graphite, fullerenes, and charcoal. Their structures are interesting; so are their properties. You probably know a lot about diamond and graphite, but the fullerenes were discovered after 1970, and this discovery has opened a door for a lot of interesting research. Read about them in books, magazines and journals. You might find yourself working with them some day.
Among the fullerenes, one of the most common "molecules" has 60 $\ce C$ atoms, and is represented by $\ce{C60}$; a diagram is shown here. If you connect the 60 carbon atoms with bonds, the structure looks like a cage with 5- and 6- member rings. Synthesis, bonding, symmetry and stability of the cagelike fullerenes have already attracted a lot of attention, and their properties are even more fascinating.
Regarding carbon compounds, you already know something about $\ce{CO2}$ and $\ce{CO}$, including their roles in the environment. The hard carbides such as $\ce{Fe3C}$, $\ce{WC}$, and $\ce{TiC}$ are more interesting to material scientists and engineers for their application in cutting tools. The calcium carbide $\ce{CaC2}$ produced by reducing $\ce{CaO}$ by carbon was a valuable commodity at one time due to its reaction with water to give acetylene gas:
$\ce{CaC_2 + 2 H_2O \rightarrow Ca(OH)_2 + C_2H_{2\large{(g)}}}$
Acetylene is an important industrial gas for the manufacture of polymers.
Silicon
Do you know that:
• $\ce{Si}$ comprises 27.7 % of the Earth's crust?
• silicates, $\ce{SiO2}$ based minerals, are everywhere?
• $\ce{Si}$ is an essential element for bone growth?
• $\ce{Si}$ crystals are the bases of computer chips?
• the structure of $\ce{Si}$ is the same as that of diamond, and this feature is important for computer chips?
• How we convert $\ce{SiO2}$ into $\ce{Si}$ element?
Silicates
By weight, silicon is the most abundant element in the Earth's crust. It usually exists in the form of an oxide, $\ce{SiO2}$. This formula does not do justice to so many different materials we call silicates, but these substances are indeed $\ce{SiO2}$. Some of the minerals contain impurities.
In pure form, $\ce{SiO2}$ is quartz. Small particles of quartz are sand. They are hard. In the structure at the atomic level, every silicon atom is bonded to 4 oxygen atoms, and every oxygen is bonded to two $\ce{Si}$ atoms. The four $\ce{Si-O}$ bonds point towards the corners of a tetrahedron, as do the $\ce{C-C}$ bonds in the diamond structure. When an impurity is present, the quartz may be colored. Due to various arrangements of the $\ce{Si-O-Si}$ bonds, the same substance appears in many forms.
A basic unit of silicate structures is $\ce{SiO4^4-}$. The gemstone zircon has a formula $\ce{ZrSiO4}$, and olivine has a chemical formula of $\ce{(MgFe)2SiO4}$. Two $\ce{SiO4^4-}$ units combine to give the pyrosilicate unit $\ce{Si2O7^6-}$, and it appears in akermanite, $\ce{Ca2MgSi2O7}$. When the number of units increases, the tetrahedral units combine to form rings, chains, layers and 3-dimensional networks. Thus, the structure and classification of silicate is a major part of minerals. (This site has some interesting pictures.)
Silicon and Silane
Elemental silicon can be obtained from reduction of silicates. The reduction of sand, $\ce{SiO2}$ by carbon at 3300 K in the reaction,
$\ce{SiO_2 + 2 C \rightarrow Si_{\large{(l)}} + 2 CO} \;\;\; \textrm{at 3300 K}$
gives liquid silicon. The silicon so obtained is usually not pure, and for the computer industry, the element must be purified. Crystal growth and silicon fabrication dominated the industry in the 1980s and 1990s, and perhaps into the next century, and the production of the element is only the beginning of the process.
If a more reactive element, $\ce{Mg}$, is used in the reduction, $\ce{Mg2Si}$ is formed,
$\ce{SiO_2 + 4 Mg \rightarrow Mg_2Si + 2 MgO}$
$\ce{Mg2Si}$ is a compound, and it reacts with water to form silane.
Silane, $\ce{SiH4}$, can be produced by reacting $\ce{Mg2Si}$ with acids
$\ce{Mg_2Si + 4 H_2O \rightarrow 2 Mg(OH)_2 + SiH_4}$
and $\ce{SiH4}$ is ignited when it contacts air, being much more reactive than methane,
$\ce{SiH_4 + O_2 \rightarrow SiO_2 + H_2O}$
In a basic solution, $\ce{SiH4}$ reacts with water to give $\ce{SiO(OH)3-}$,
$\ce{SiH_4 + OH^- + 3H_2O \rightarrow SiO(OH)_3^- + 4 H_2}$
Silicon Halides
Silicon tetrafluoride is formed when glass ($\ce{SiO2}$) is exposed to $\ce{HF}$, and when $\ce{Si}$ reacts with $\ce{F2}$,
$\ce{SiO2 + 4 HF_{\large{(aq)}} \rightarrow 2 H2O + SiF_{4\large{(g)}}}$
$\ce{Si + 2 F2 \rightarrow SiF_{4\large{(g)}}}$
When chlorine passes through hot sand ($\ce{SiO2}$) and carbon, $\ce{SiCl4}$ is produced,
$\ce{SiO2 + C + 2 Cl_2 \rightarrow 2CO + SiCl_{4\large{(g)}}}$
$\ce{SiCl4}$ and $\ce{SiF4}$ react with water to give silicic acid,
$\ce{SiCl4 + 4H2O \rightarrow 4 HCl + Si(OH)_{4\large{(aq)}}}$,
$\ce{SiF4 + 4H2O \rightarrow 4 HF + Si(OH)_{4\large{(aq)}}}$.
Silicone Polymers
Silicones are polymers with the general formula $\mathrm{(R_2SiO_2)_{\large n}}$ or $\mathrm{(RSiO_3)_{\large n}}$, ($\ce{R}$ = $\ce{CH3}$, $\ce{C2H5}$, $\ce{C6H5}$, etc). The chain is held together by $\ce{Si-O-Si}$ bonds. A simple one is $\mathrm{((CH_3)_2SiO_2)_{\large n}}$,
Figure 2: Chemical structure of the silicone polydimethylsiloxane (PDMS).
Of course, the 4 bonds around the $\ce{Si}$ atoms point to the corners of a tetrahedron. These siloxane polymers are widely used as sealants, adhesives, additives, flame retardants, and lubricants. They have a wide application in industries. Depending on the organic group attached to silicon, the inorganic polymer silicones have been an important class of materials.
Questions
1. Which one lists the group 14 elements in order of increasing atomic weight?
1. $\textrm{B Al Ga In Tl}$
2. $\textrm{C Si Ge Sn Pb}$
3. $\textrm{N P As Sb Bi}$
4. $\textrm{O S Se Te Po}$
5. $\textrm{F Cl Br I At}$
2. Which allotrope of carbon is the hardest: diamond, graphite, or fullerenes?
3. What compound of carbon reacts with water to give acetylene gas?
4. When you want to extract silicon element, what do you use to reduce the sand: $\ce{SiO2}$, $\ce{C}$ or $\ce{Mg}$?
5. Which is stable towards air: methane or silane?
6. Give the name of polymers whose chains are held together by $\ce{Si-O-Si}$ bonds.
7. In the crystal structure of $\ce{Si}$, how many other $\ce{Si}$ atoms are connected to a particular $\ce{Si}$ atom?
Solutions
1. Answer b
Hint...
Knowing the groups of elements enables us to correlate their chemical properties. Each list of choices is a group of elements on the period table.
a = 3A, b = 14, c = 5A, d = 6A, e = 7A.
2. Answer diamond
Hint...
Diamond is the hardest thing in the world. Fullerenes are large molecules consisting of 40 to hundreds of carbon atoms, with $\mathrm{C_{60}}$ being the most common.
3. Answer . . .$\ce{CaC2}$
Hint...
The reaction to produce acetylene gas is
$\ce{CaC2 + H2O \rightarrow C2H2 + Ca(OH)2}$
Acetylene is still an important industrial gas, as raw material for polymers.
4. Answer ...$\ce{C}$
Hint...
Carbon or coke is used for silicon metal, because $\ce{Mg2Si}$ is formed if $\ce{Mg}$ is used.
5. Answer ... methane is stable
Hint...
Methane is the major component of natural gas, and it will not react with air until ignited, whereas silane ignites explosively as soon as it contacts air.
6. Answer ...silicones
Hint...
Silicon polymers are an important class of materials invented not too long ago.
7. Answer ... 4
Hint...
Silicon and diamond have the same crystal structure. The edge of unit cells of $\ce{Si}$ is larger than that of diamond. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14%3A_The_Carbon_Family/Z014_Chemistry_of_Silicon_%28Z14%29/Silicon_and.txt |
Germanium, categorized as a metalloid in group 14, the carbon family, has five naturally occurring isotopes. Germanium, abundant in the Earth's crust, has been said to improve the immune system of cancer patients. It is also used in transistors, but its most important use is in fiber-optic systems and infrared optics.
Introduction
The metalloid was one of the elements predicted by Mendeleev in 1871 (ekasilicon) to fill out his periodic table and was discovered in 1886 by Winkler. In a mine near Freiberg, Saxony, a new mineral was found in 1885. First the mineral was called argyrodite, but later, when Clemens Winkler examined this mineral he discovered that it was similar to antimony. At first he wanted to name it neptunium, but because this name was already taken he named it germanium in honor of his fatherland Germany.
The position of where germanium should be placed on the periodic table was under discussion during the time due to its similarities to arsenic and antimony. Due to Mendeleev's prediction of ekasilicon, germanium's place on the periodic table was confirmed because of the similar properties predicted and similar properties deduced from examining the mineral.
Characteristics
Like silicon, germanium is used in the manufacture of semi-conductor devices. Unlike silicon, it is rather rare (only about 1 part in 10 million parts in the earth's crust). The physical and chemical properties of germanium closely parallel those of silicon. Germanium (Ge) has an atomic number of 32. It is grayish-white, lustrous, hard, and has similar chemical properties to tin and silicon. Germanium is brittle and silvery-white under standard conditions. Under these conditions, germanium is known as α-germanium, which has a diamond cubic crystal structure. When germanium is under pressure above 120 kilobars, it has a different allotrope known as β-germanium. Germanium is one of the few substances like water that expands when it solidifies.
Atomic Number: 32
Atomic Symbol: Ge
Atomic Weight: 72.59
Atomic Radius: 122.5 pm
Series: Metalloid
Group: 14
Density: 5.353 g/cm3
Specific Heat: 0.32J/gK
Electronic Configuration: [Ar] 4s23d104p2
Melting Point: 938.25 C
Boiling Point: 2833 C
Common Oxidation States: 4,2
Germanium, a semiconductor, is the first metallic metal to become a superconductor in the presence of a strong electromagnetic field.
Isotopes
The five naturally occurring isotopes of germanium have atomic masses of 70, 72, 73, 74, and 76. Germanium 76 is slightly radioactive and is the least common. Germanium 74 is the most common isotope, having the greatest natural abundance of the five. When it is bombarded with alpha particles, Germanium 72 generates stable Se 77.
Chemical Properties
At a temperature of 250 °C, germanium slowly oxidizes to \(GeO_2\). Germanium dissolves slowly in concentrated sulfuric acid, and is insoluble in diluted acids and alkalis. It will react violently with molten alkalis to produce [GeO3]2-. The common oxidation states of germanium are +4 and +2. Under rare conditions, germanium also occurs in oxidation states of +3, +1, and -4.
There are two forms of germanium oxides: germanium dioxide (\(GeO_2\)) and germanium monoxide (\(GeO\)). By roasting germanium sulfide, germanium dioxide can be obtained. As for germanium monoxide, it can be obtained by the high temperature reaction of germanium dioxide and germanium metal. Germanium dioxide has the unusual property of a refractive index for light but transparency to infrared light.
Applications and Uses
Like silicon, germanium is used in the manufacture of semi-conductor devices. Unlike silicon, it is rather rare (only about 1 part in 10 million parts in the earth's crust). Although it forms a compound, germanium dioxide, just like silicon, it is generally extracted from the by-products of zinc refining.
Germanium is used mainly for fiber-optic systems and infrared optics. It is also used for polymerization catalysts, electronics, and as phosphors, for metallurgy, and for chemotherapy. Because germanium dioxide has a high index of refraction and low optical dispersion, it is useful for wide-angle camera lenses. As stated before, it is an important semi-conductor, so it is used in transistors. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14%3A_The_Carbon_Family/Z032_Chemistry_of_Germanium_%28Z32%29.txt |
Mentioned in the Hebrew scriptures, tin is of ancient origins. Tin is an element in Group 14 (the carbon family) and has mainly metallic properties. Tin has atomic number 50 and an atomic mass of 118.710 atomic mass units.
Introduction
Mentioned in the Hebrew scriptures, tin is of ancient origins. Early metal smiths were quick to learn that mixing copper with tin created a more durable metal (bronze), and it is principally for its alloys that tin is valued today. Named after the Etruscan god Tinia, the chemical symbol for tin is taken from the Latin stannum. The metal is silvery white and very soft when pure. It has the look of freshly cut aluminum, but the feel of lead.
Polished tin is slightly bluish. It has been used for many years in the coating of steel cans for food because it is more resistant to corrosion than iron. It forms a number of useful low-melting alloys (solders) which are used to connect electrical circuits. Bending a bar of tin produces a characteristic squealing sound called "tin cry". Tin shares chemical similarities with germanium and lead. Tin mining began in Australia in 1872, and today tin is used extensively in industry and commerce.
Table 1: Basic Properties of tin
Color white with blueish tinge
Hardness softer than gold, harder than lead
Atomic Radius 140 pm
Density
5.77g/cm3
Melting Point 232 degrees Celsius
Boiling Point 2623 degrees Celsius
Electrical Conductivity about 1/7th that of silver
Electrode Potential >0.192V
First Ionization Energy 709 kJ/mol
Ionic Radius 93 pm
Reactions of Tin
Hydrogen Tin not affected
Nitrogen Tin absorbs it instead of hydrogen in electric discharge
Argon No sign of a combination of tin with argon
Fluorine Does not react with tin at low temperatures, but at 100 degrees Celsius they form stannic fluoride. Perhaps one of the most familiar of tin compounds, $SnF_2$, tin(II) fluoride, goes by the trade name of fluoristan and is found in some fluoride toothpastes.
Chlorine Acts on tin at room temperature
Bromine Acts on tin at room temperature
Sulfur Unites directly with tin when heated
Selenium Reacts vigorously with tin
Tellurium Reacts vigorously with tin
Nitrogen Forms a compound by direct union with tin
Arsenic Reacts with tin under heat and light
Antimony Is dissolved by molten tin
Reaction of tin with oxygen
When heated in oxygen, tin produces stannic oxide:
$Sn_{(s)} + O_{2(g)} \rightarrow SnO_{2(s)} \nonumber$
Reaction of tin with water (steam)
$Sn_{(s)} + 2H_2O_{(g)} \rightarrow SnO_{2(s)} + 2H_{2(g)} \nonumber$
Isotopes
There are 10 known stable isotopes of tin, the most of any elements on the periodic table. This high number of stable isotopes could be attributed to the fact that the atomic number of $\ce{^{50}Sn}$ is a 'magic number' in nuclear physics.
Table 4: Isotopes of tin
Isotope % Natural Abundance
112 amu 0.95%
116 amu 14.24%
117 amu 7.57%
118 amu 24.01%
119 amu 8.58%
120 amu 32.97%
122 amu 4.71%
124 amu 5.98%
Allotropes of Tin
Tin has 3 allotropes: alpha, beta and gamma tin. Alpha tin is the most unstable form. Beta tin is the most commonly found allotrope of tin, and gamma tin only exists at very high temperatures.
Oxidation States of Tin
Tin, although it is found in Group 14 of the periodic table, is consistent with the trend found in Group 13, where the lower oxidation state is favored farther down a group. Tin can exist in two oxidation states, +2 and +4, but it displays a tendency to exist in the +4 oxidation state.
Common Compounds of Tin
Tin forms two main oxides, SnO and SnO2 (amphoteric).
Electron Configuration of Tin
Tin has a ground state electron configuration of 1s22s22p63s23p64s23d104p65s24d105p2 and can form covalent tin (II) compounds with its two unpaired p-electrons. In the three dimensional figure below, the first and most inner electron shell is represented by blue electrons, the second electron shell made up of eight electrons is represented by red electrons, the third shell containing eighteen electrons is represented with green electrons, and the next outer shell again contains eighteen electrons and is represented in purple.
Uses of Tin
Early metal smiths were quick to learn that mixing copper with tin created a more durable metal (bronze) and it is principally for its alloys that tin is valued today. Nearly half of the tin metal produced is used in solders, which are low melting point alloys used to join wires. Solders are important in electrician work and plumbing. Tin is also used as a coating for lead, zinc, and steel to prevent corrosion. Tin cans are widely used for storing foods; the first tin can was used in London in 1812.
Questions
Find the oxidation state of tin in the following compounds:
a. SnCl2 Answer:2
b. SnO2 Answer:4
Write an equation for the reaction of tin with water. Under what conditions does this reaction take place?
Answer: Sn(s) + 2H2O(g) → SnO2(s) + 2H2(g) Reaction takes place if water is heated to a high temperature to form steam.
Which of these reactions take place?
a. tin with oxygen Answer: YES
b. tin with hydrogen Answer: NO
c. tin with argon Answer: NO
d. tin with chlorine Answer: YES
Arrange the following in order of increasing atomic radius: Sn, K, Ag, C, Pb.
Answer: C<Sn<Pb<Ag<K
Arrange the following in order of decreasing ionization energy: Sn, Si, Pb, I, In.
Answer: Si> I > Sn > In > Pb
Contributors and Attributions
• Taylor Hughes, (UCSB) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14%3A_The_Carbon_Family/Z050_Chemistry_of_Tin_%28Z50%29.txt |
Known to the ancients, lead takes its name from the Anglo-Saxon word for the metal, and its symbol comes from the Latin plumbum (from which we get the modern word "plumber" since old plumbing was done with lead pipes).
Although lead is not very common in the earth's crust, what is there is readily available and easy to refine. Its chief use today is in lead-acid storage batteries such as those used in automobiles. In pure form it is too soft to be used for much else. Lead has a blue-white color when first cut but quickly dulls on exposure to air, forming Pb2O, one of the few lead(I) compounds. Most stable lead compounds contain lead in oxidation states of +2 or +4.
Various isotopes of lead come at the end of the natural decay series of elements like uranium, thorium and actinium. These are Pb-206, Pb-207 and Pb-208.
Contributors and Attributions
Stephen R. Marsden
Z082 Chemistry of Lead (Z82)
Lead plumbate, also called red lead, minium or Mennige (in German), is a mineral showing colors from light red to brown/yellow tints. As a pure chemical it shows a vivid red. Minium is rare and occurs in lead mineral deposits that have been subjected to severe oxidizing conditions. It also occurs as a result of mine fires. It is most often associated with galena, cerussite, massicot, litharge, native lead, wulfenite and mimetite.
Lead plumbate is obtained by heating lead monoxide ($PbO$) to 450-480°C in air:
$3 PbO + 1/2 O_2 \rightarrow Pb_3O_4 \nonumber$
or by oxidative annealing of lead white:
$3 Pb_2CO_3(OH)_2 + O_2 \rightarrow 2 Pb_3O_4 + 3 CO_2 + 3 H_2O \nonumber$
Lead plumbate decomposes into lead monoxide and oxygen above 550°C.
$Pb_3O_4$ can be seen formally as a lead(II)plumbate(IV), $Pb_2[PbO_4]$, or $2PbO\cdot PbO_2$. In nitric acid, the lead(II) oxide reacts forming lead nitrate, while the insoluble lead(IV) oxide is left unchanged:
$Pb_3O_4 + 4 HNO_3 \rightarrow 2 Pb(NO_3)_2 + PbO_2 + 2 H_2O \nonumber$
Lead plumbate is virtually insoluble in water. However, it dissolves in hydrochloric acid (which is present in the stomach), and is therefore toxic when ingested. Lead plumbate (in a mixture with linseed oil or other organic adhesives) has been used as an anti-corrosion paint for iron. It forms insoluble iron(II) and iron(III) plumbates when brought into contact with iron oxides and with elementary iron. However, its use as a protective undercoat paint is limited due to its toxicity.
Lead plumbate was used as a red pigment in ancient and medieval periods for paintings and the production of illuminated manuscripts (the term miniature is connected to the name of the substance).
Z114 Chemistry of Flerovium (Z114)
The synthesis of element 114 was reported in January of 1999 by scientists from the Joint Institute for Nuclear Research in Dubna (near Moscow) and Lawrence Livermore National Laboratory (in California). In an experiment lasting more than 40 days Russian scientists bombarded a film of Pu-244 supplied by Livermore scientists with a beam of Ca-48. One atom of element 114 was detected with a half-life of more than 30 seconds. This is about 100,000 times as long as the previously longest-lived isotope of 112 produced (element 113 has yet to be made) and bolsters the arguments of theorists who envision an "island of stability" in the super-heavy elements. In May of 2012 the IUPAC approved the name "Flerovium" (symbol Fl), in honor of the Flerov Laboratory of Nuclear Reactions where superheavy elements are synthesized.
Contributors and Attributions
Stephen R. Marsden | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_14%3A_The_Carbon_Family/Z082_Chemistry_of_Lead_%28Z82%29/Lead_Plumbate.txt |
The nitrogen family includes the following compounds: nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi). All Group 15 elements have the electron configuration ns2np3 in their outer shell, where n is the principal quantum number.
• Group 15: General Properties and Reactions
The nitrogen family includes the following compounds: nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi). All Group 15 elements have the electron configuration ns2np3 in their outer shell, where n is equal to the principal quantum number. The nitrogen family is located in the p-block in Group 15, as shown below.
• Chemistry of Nitrogen (Z=7)
Nitrogen is present in almost all proteins and plays important roles in both biochemical applications and industrial applications. Nitrogen forms strong bonds because of its ability to form a triple bond with itself and other elements. Thus, there is a lot of energy in the compounds of nitrogen. Before 100 years ago, little was known about nitrogen. Now, nitrogen is commonly used to preserve food and as a fertilizer.
• Chemistry of Phosphorus (Z=15)
Phosphorus (P) is an essential part of life as we know it. Without the phosphates in biological molecules such as ATP, ADP and DNA, we would not be alive. Phosphorus compounds can also be found in the minerals in our bones and teeth. It is a necessary part of our diet. In fact, we consume it in nearly all of the foods we eat. Phosphorus is quite reactive. This quality of the element makes it an ideal ingredient for matches because it is so flammable.
• Chemistry of Arsenic (Z=33)
Arsenic is situated in the 33rd spot on the periodic table, right next to Germanium and Selenium. Arsenic has been known for a very long time and the person who may have first isolated it is not known but credit generally is given to Albertus Magnus in about the year 1250. The element, which is classified as a metalloid, is named from the Latin arsenicum and Greek arsenikon which are both names for a pigment, yellow orpiment.
• Chemistry of Antimony (Z=51)
Antimony and its compounds have been known for centuries. Scientific study of the element began during the early 17th century, much of the important work being done by Nicolas Lemery. The name of the element comes from the Greek anti + monos for "not alone", while the modern symbol is rooted in the Latin-derived name of the common ore, stibnite.
• Chemistry of Bismuth (Z=83)
Bismuth, the heaviest non-radioactive naturally occurring element, was isolated by Basil Valentine in 1450. It is a hard, brittle metal with an unusually low melting point (271oC). Alloys of bismuth with other low-melting metals such as tin and lead have even lower melting points and are used in electrical solders, fuse elements and automatic fire sprinkler heads.
• Chemistry of Moscovium (Z=115)
In studies announced jointly by the Joint Institute for Nuclear Research in Dubna, Russia, and the Lawrence Livermore National Laboratory in the U.S., four atoms of element 113 were produced in 2004 via decay of element 115 after the fusion of Ca-48 and Am-243.
Thumbnail: White and red phosphorus. (CC-SA-BY 3.0; Peter Krimbacher).
Group 15: The Nitrogen Family
The nitrogen family includes the following compounds: nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi). All Group 15 elements have the electron configuration ns2np3 in their outer shell, where n is equal to the principal quantum number. The nitrogen family is located in the p-block in Group 15.
Periodic Trends
All Group 15 elements tend to follow the general periodic trends:
• Electronegativity (the atom's ability to attract electrons) decreases down the group.
• Ionization energy (the amount of energy required to remove an electron from the atom in its gas phase) decreases down the group.
• Atomic radii increase in size down the group.
• Electron affinity (the ability of the atom to accept an electron) decreases down the group.
• Melting point (amount of energy required to break bonds to change a solid phase substance to a liquid phase substance) increases down the group.
• Boiling point (amount of energy required to break bonds to change a liquid phase substance to a gas) increases down the group.
• Metallic character increases down the group.
Element/Symbol Atomic Number Mass Electron Configuration Covalent Radius (pm) Electronegativity First Ionizaton Energy (kJ/mol) Common Physical Form(s)
Properties of Group 15 Elements
Nitrogen (N) 7 14.01 1s2 2s2 2p3 75 3.0 1402 Colorless Gas
Phosphorus (P) 15 30.97 [Ne]3s2 3p3 110 2.1 1012 White Solid / Red Solid
Arsenic (As) 33 74.92 [Ar] 3d10 4s2 4p3 121 2.0 947 Yellow Solid / Gray Solid
Antimony (Sb) 51 121.76 [Kr] 4d10 5s2 5p3 140 1.9 834 Yellow Solid / Silver-White Metallic Solid
Bismuth (Bi) 83 208.98 [Xe] 4f14 5d10 6s2 6p3 155 1.9 703 Pink-White Metallic Solid
Nitrogen
Nitrogen was discovered in 1770 by Scheele and Priestley. This nonmetallic element has no color, taste, or odor, and is present in nature as a noncombustible gas. When compared with the rest of Group 15, nitrogen has the highest electronegativity, which makes it the most nonmetallic of the group. The common oxidation states of nitrogen are +5, +3, and -3. Nitrogen makes up about 0.002% of the earth's crust; however, it constitutes 78% of the earth’s atmosphere by volume. Nitrogen has also been discovered in the atmospheres of Venus and Mars. Venus has a 3.5% nitrogen volume in its atmosphere and Mars has a 2.7% nitrogen volume in its atmosphere. Nitrogen is found naturally in animal and plant proteins and in fossilized remains of ancient plant life. Important nitrogen-containing minerals are niter, KNO3, and soda niter, NaNO3, which are found in desert regions and are important components of fertilizers. Before the process of converting nitrogen into ammonia was discovered, sources of nitrogen were limited. One of the processes of converting nitrogen to ammonia, the Haber-Bosch process, is very important for the production of nitrogen. Nitrogen has very little solubility in liquids. N2 does not have any allotropes.
The unusually stable N2(g) nitrogen gas is the source from which all nitrogen compounds are ultimately derived. N2(g) is stable due to its electronic structure: the bond between the two nitrogen atoms of N2 is a triple covalent bond, which is strong and hard to break. The enthalpy change associated with breaking the bonds in N2 is highly endothermic: ΔH = +945.4 kJ. Nitrogen gas is used as a refrigerant, metal treatment, and pressurized gas for oil recovery. Additionally, the Gibbs energies of formation of nitrogen compounds show that their formations are nonspontaneous, and the following process does not occur at normal temperatures:
$1/2N_{2(g)} + 1/2O_2 \rightarrow NO_{(g)} \nonumber$
with
$ΔG_f= +86.55\;kJ\, mol^{-1} \nonumber$
The oxides and oxyacids of nitrogen include nitrous oxide (N2O), nitrogen oxide (NO), and nitrogen dioxide (NO2). Nitrous oxide, also called “laughing gas,” is used in dental work, child birth, and to increase the speed of cars. Nitrogen oxide is found in smog and neurotransmitters. Hydrazine, N2H4, is a poisonous, colorless liquid that explodes in air. Hydrazine is a good reducing agent, and methyl hydrazine is commonly used as a rocket fuel.
Phosphorus
Phosphorus is a nonmetallic element. The most common oxidation state of phosphorus is -3. Phosphorus is the eleventh most abundant element, making up 0.11% of the earth's crust. The main source of phosphorus compounds is phosphorus rocks. Phosphorous is not found pure in nature, but in the form of apatite ores. These include compounds such as fluorapatite (Ca5(PO4)3F), which in fluoridated water is used to strengthen teeth, and hydroxylapatite (Ca10(OH)2(PO4)6), a major component of tooth enamel and bone material. Phosphorus exhibits allotropic forms: the most common forms at room temperature are white phosphorus and red phosphorus. White phosphorus is a white, waxy solid that can be cut with a knife. It forms a tetrahedral molecule, P4. White phosphorus is toxic, while red phosphorus is nontoxic. Red phosphorous forms when white phosphorous is heated to 573 Kelvin and not exposed to air. Red phosphorus is less reactive than white phosphorus. Red phosphorus has a chain-like polymeric structure, and is more stable. Both white and red phosphorus are incendiary and have been used to make match tips, although the use of white phosphorus has been largely discontinued due to toxicity. Phosphorus has many applications: phosphorus trichloride (PCl3) is used in soaps, detergents, plastics, synthetic rubber nylon, motor oils, insecticides and herbicides; phosphoric acid, H3PO4, is used in fertilizers; phosphorus is also prevalent in the food industry, used in baking powders, instant cereals, cheese, the curing of ham, and in the tartness of soft drinks.
Arsenic
Arsenic is a highly poisonous metalloid. Because it is a metalloid, arsenic has a high density, moderate thermal conductivity, and a limited ability to conduct electricity. The oxidation states of arsenic are +5, +3, +2, +1, and -3. The three allotropic forms of arsenic are yellow, black, and gray; the gray allotrope is the most common. Compounds of arsenic are used in insecticides, weed killers, and alloys. The oxide of arsenic is amphoteric, meaning it can act as both an acid and a base. Arsenic is mainly obtained by heating arsenic-containing sulfides. The chemical formula for this process is given below:
$FeAsS_{(s)} \rightarrow FeS_{(s)} + As_{(g)} \nonumber$
The $As_{(g)}$ deposits as $As_{(s)}$, which can then be used to make other compounds. Arsenic can also be obtained by the reduction of arsenic(III) oxide with $CO_{(g)}$.
Antimony
Antimony is also a metalloid. The oxidation states of antimony are +3, -3, and +5. Antimony exhibits allotropy; the most stable allotrope is the metallic form, which is similar in properties to arsenic: high density, moderate thermal conductivity, and a limited ability to conduct electricity. The oxide of antimony is antimony (III) oxide, which is amphoteric, meaning it can act as both an acid and a base. Antimony is obtained mainly from its sulfide ores, and it vaporizes at low temperatures. Along with arsenic, antimony is commonly used in alloys. Arsenic, antimony, and lead form an alloy with desirable properties for electrodes in lead-acid batteries. Arsenic and antimony are also used to produce semiconductor materials such as GaAs, GaSb, and InSb in electronic devices.
Bismuth
Bismuth is a metallic element. The oxidation states of bismuth are +3 and +5. Bismuth is a poor metal (one with significant covalent character) that is similar to both arsenic and antimony. Bismuth is commonly used in cosmetic products and medicine. Out of the group, bismuth has the lowest electronegativity and ionization energy, which means that it is more likely to lose an electron than the rest of the Group 15 elements. This is why bismuth is the most metallic of Group 15. Bismuth is also a poor electrical conductor. The oxide of bismuth is bismuth(III) oxide; it acts as a base, as expected for a metal oxide. Bismuth is obtained as a by-product of the refining of other metals, allowing other metals to recycle their by-products into bismuth.
Problems
1. How much of the earth's crust is made up of nitrogen?
2. How much of earth's crust is not made up of phosphorus?
3. What kind of bond does N2 have?
4. How are red phosphorus and white phosphorus related to each other?
5. What is the electron configuration of arsenic?
6. Does bismuth have metallic properties or nonmetallic properties?
7. True or False: nitrogen and phosphorus are metals.
8. Which Group 15 element has the greatest atomic radius?
9. Which Group 15 element is the strongest reducing agent?
10. True or False: bismuth exhibits the most metallic character.
11. What is the most common physical form of nitrogen?
12. What is the process in which nitrogen can convert into ammonia?
13. Which element has the highest first ionization energy?
14. Complete and balance the following reaction: _N2(g) + _H2(g) -> ____
15. What is the common oxidation state of all Group 15 elements?
Answers
1. 0.002% of the earth's crust is made up of nitrogen.
2. The earth's crust is made up of 0.11% phosphorus, so 99.89% of the earth's crust is not made up of phosphorus.
3. N2 has a triple covalent bond that is strong and hard to break.
4. Red phosphorus and white phosphorus are both allotropes of phosphorus. Red phosphorus comes from white phosphorus when it is heated to about 573 Kelvin.
5. [Ar] 3d10 4s2 4p3
6. Bismuth has metallic properties.
7. False; both nitrogen and phosphorus are nonmetals.
8. According to periodic trends, bismuth has the greatest atomic radius.
9. According to periodic trends, bismuth is the strongest reducing agent since it has an electronegativity value of 1.9, which is the same as antimony, but it has a lower ionization energy of 703 kJ/mol, which means it is more likely to get oxidized.
10. True; bismuth is the only metallic element of Group 15.
11. The most common physical form of nitrogen is a colorless gas.
12. The process of converting nitrogen into ammonia is known as the Haber-Bosch process.
13. According to periodic trends, nitrogen would have the highest first ionization energy, which means that it does not want to lose an electron the most.
14. N2(g) + 3H2(g) -> 2NH3(g)
15. The common oxidation state for all Group 15 elements is -3.
Contributors and Attributions
• Kirenjot Grewal (UCD), Connie Sou (UCD) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_15%3A_The_Nitrogen_Family/1Group_15%3A_General_Properties_and_Reactions.txt |
Nitrogen is present in almost all proteins and plays important roles in both biochemical applications and industrial applications. Nitrogen forms strong bonds because of its ability to form a triple bond with itself and other elements. Thus, there is a lot of energy in the compounds of nitrogen. Before 100 years ago, little was known about nitrogen. Now, nitrogen is commonly used to preserve food and as a fertilizer.
Introduction
Nitrogen is found to have either 3 or 5 valence electrons and lies at the top of Group 15 on the periodic table. It can have either 3 or 5 valence electrons because it can bond in the outer 2p and 2s orbitals. Molecular nitrogen ($N_2$) is not reactive at standard temperature and pressure and is a colorless and odorless gas.
Nitrogen is a non-metal element that occurs most abundantly in the atmosphere; nitrogen gas (N2) comprises 78.1% of the volume of the Earth’s air. It only appears in 0.002% of the earth's crust by mass. Compounds of nitrogen are found in foods, explosives, poisons, and fertilizers. Nitrogen is found in DNA in the form of nitrogenous bases as well as in neurotransmitters. It is one of the most produced industrial gases, and is produced commercially as a gas and a liquid.
Table 1: General Properties of Nitrogen
Name and Symbol Nitrogen, N
Category non-metal
Atomic Weight 14.0067
Group 15
Electron Configuration 1s2 2s2 2p3
Valence Electrons 2, 5
Phase Gas
History
Nitrogen, which makes up about 78% of our atmosphere, is a colorless, odorless, tasteless and chemically unreactive gas at room temperature. Its name is derived from the Greek nitron + genes for soda forming. For many years during the 1500's and 1600's, scientists hinted that there was another gas in the atmosphere besides carbon dioxide and oxygen. It was not until the 1700's that scientists could prove there was in fact another gas that took up mass in the atmosphere of the Earth.
Nitrogen was discovered in 1772 by Daniel Rutherford (and independently by others such as Priestly and Cavendish); Rutherford was able to remove oxygen and carbon dioxide from a contained tube full of air. He showed that there was residual gas that did not support combustion as did oxygen or carbon dioxide. While his experiment was the one that proved that nitrogen existed, other experiments were also going on in London, where they called the substance "burnt" or "dephlogisticated air".
Nitrogen is the fourth most abundant element in humans, and it is more abundant in the known universe than carbon or silicon. Most commercially produced nitrogen gas is recovered from liquefied air. Of that amount, the majority is used to manufacture ammonia ($NH_3$) via the Haber process. Much is also converted to nitric acid ($HNO_3$).
Isotopes
Nitrogen has two naturally occurring isotopes, nitrogen-14 and nitrogen-15, which can be separated with chemical exchanges or thermal diffusion. Nitrogen also has isotopes with masses of 12, 13, 16, and 17, but they are radioactive.
• Nitrogen 14 is the most abundant form of nitrogen and makes up more than 99% of all nitrogen found on Earth. It is a stable compound and is non-radioactive. Nitrogen-14 has the most practical uses, and is found in agricultural practices, food preservation, biochemicals, and biomedical research. It is found in abundance in the atmosphere and among many living organisms. It has 5 valence electrons and is not a good electrical conductor.
• Nitrogen-15 is the other stable form of nitrogen. It is often used in medical research and preservation. The element is non-radioactive and therefore can also be sometimes used in agricultural practices. Nitrogen-15 is also used in brain research, specifically nuclear magnetic resonance spectroscopy (NMR), because unlike nitrogen-14 (nuclear spin of 1), it has a nuclear spin of 1/2, which has benefits when it comes to MRI research and NMR observations. Lastly, nitrogen-15 can be used as a label or in some proteins in biology. Scientists mainly use this compound for research purposes and have not yet seen its full potential for uses in brain research.
Compounds
The two most common compounds of nitrogen are potassium nitrate (KNO3) and sodium nitrate (NaNO3). These two compounds are formed by decomposing organic matter that has potassium or sodium present and are often found in fertilizers and byproducts of industrial waste. Most nitrogen compounds have a positive Gibbs free energy (i.e., reactions are not spontaneous).
The dinitrogen molecule ($N_2$) is an "unusually stable" compound, particularly because nitrogen forms a triple bond with itself. This triple bond is difficult to break. For dinitrogen to follow the octet rule, it must have a triple bond. Nitrogen has a total of 5 valence electrons, so doubling that, we would have a total of 10 valence electrons with two nitrogen atoms. The octet requires an atom to have 8 total electrons in order to have a full valence shell, therefore it needs to have a triple bond. The compound is also very inert, since it has a triple bond. Triple bonds are very hard to break, so they keep their full valence shell instead of reacting with other compounds or atoms. Think of it this way: each triple bond is like a rubber band; with three rubber bands, the nitrogen atoms are very attracted to each other.
Nitrides
Nitrides are compounds of nitrogen with a less electronegative atom; in other words they are compounds with atoms that have a less full valence shell. These compounds form with lithium and Group 2 metals. Nitrides usually have an oxidation state of -3.
$3Mg + N_2 \rightarrow Mg_3N_2 \label{1}$
When mixed with water, nitride will form ammonia; the nitride ion acts as a very strong base.
$N^{3-} + 3H_2O_{(l)} \rightarrow NH_3 + 3OH^-_{(aq)} \label{2}$
When nitrogen forms compounds with other atoms, it primarily forms covalent bonds. These are normally formed with other metals and look like: MN, M3N, and M4N. These compounds are typically hard, inert, and have high melting points because of nitrogen's ability to form triple covalent bonds.
Ammonium Ions
Nitrogen goes through fixation by reaction with hydrogen gas over a catalyst. This process is used to produce ammonia. As mentioned earlier, this process allows us to use nitrogen as a fertilizer because it breaks down the strong triple bond held by N2. The famous Haber-Bosch process for synthesis of ammonia looks like this:
$N_2 + 3H_2 \rightarrow 2NH_3 \label{3}$
Ammonia is a base and is also used in typical acid-base reactions.
$2NH_{3(aq)} + H_2SO_4 \rightarrow (NH_4)_2SO_{4(aq)} \label{4}$
Nitride ions are very strong bases, especially in aqueous solutions.
Oxides of Nitrogen
Nitrogen uses a variety of different oxidation numbers from +1 to +5 for oxide compounds. Almost all the oxides that form are gasses, and exist at 25 degrees Celsius. Oxides of nitrogen are acidic and easily attach protons.
$N_2O_5 + H_2O \rightarrow 2HNO_{3 (aq)} \label{5}$
The oxides play a large role in living organisms. They can be useful, yet dangerous.
• Dinitrogen monoxide (N2O) is an anesthetic used by dentists, also known as laughing gas.
• Nitrogen dioxide (NO2) is harmful. It binds to hemoglobin molecules, not allowing the molecule to release oxygen throughout the body. It is released from cars and is very toxic.
• Nitrate (NO3-) is a polyatomic ion.
• The more unstable nitrogen oxides allow for space travel.
Hydrides
Hydrides of nitrogen include ammonia (NH3) and hydrazine (N2H4).
• In aqueous solution, ammonia forms the ammonium ion, which we described above, and it has special amphiprotic properties.
• Hydrazine is commonly used as rocket fuel.
Applications of Nitrogen
• Nitrogen provides a blanketing for our atmosphere and is used for the production of chemicals and electronic compartments.
• Nitrogen is used as fertilizer in agriculture to promote growth.
• Pressurized gas for oil production.
• Refrigerant (such as freezing food fast).
• Explosives.
• Metal treatment/protectant via exposure to nitrogen instead of oxygen.
Problems
• Complete and balance the following equations
N2 + ___H2→ ___NH_
H2N2O2 → ?
2NH3 + CO2 → ?
__Mg + N2 → Mg_N_
N2H5 + H2O → ?
• What are the different isotopes of nitrogen?
• List the oxidation states of various nitrogen oxides: N2O, NO, N2O3, N2O4, N2O5
• List the different elements that nitrogen will react with to make it basic or acidic....
• List some uses of nitrogen
Answers
• Complete and balance the following equations
N2 + 3H2→ 2NH3(Haber process)
H2N2O2 → 2HNO
2NH3 + CO2 → (NH2)2CO + H2O
2Mg + 3N2 → Mg3N2
N2H5 + H2O → N2+ H+ + H2O
• What are the different isotopes of nitrogen?
Stable forms include nitrogen-14 and nitrogen-15.
• List the oxidation states of various nitrogen oxides: +1, +2, +3, +4, +5 respectively.
• List the different elements that nitrogen will react with to make it basic or acidic :Nitride ion is a strong base when reacted with water; ammonia is generally a weak acid.
• Uses of nitrogen include anesthetic, refrigerant, and metal protector.
Contributors
• Adam Wandell (UC Davis) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_15%3A_The_Nitrogen_Family/Z007_Chemistry_of_Nitrogen_%28Z7%29.txt |
Learning Objectives
• Compare properties of Group 15 elements.
• Explain the major application of phosphate.
• Describe the equilibria of the ionization of phosphoric acid.
Phosphorus (P) is an essential part of life as we know it. Without the phosphates in biological molecules such as ATP, ADP and DNA, we would not be alive. Phosphorus compounds can also be found in the minerals in our bones and teeth. It is a necessary part of our diet. In fact, we consume it in nearly all of the foods we eat. Phosphorus is quite reactive. This quality of the element makes it an ideal ingredient for matches because it is so flammable. Phosphorus is a vital element for plants and that is why we put phosphates in our fertilizer to help them maximize their growth.
Introduction
Phosphorus plays a big role in our existence but it can also be dangerous. When fertilizers containing phosphorus enter the water, it produces rapid algae growth. This can lead to eutrophication of lakes and rivers; i.e., the ecosystem has an increase of chemical nutrients and this can led to negative environmental effects. With all the excess phosphorus, plants grow rapidly then die, causing a lack of oxygen in the water and an overall reduction of water quality. It is thus necessary to remove excess phosphorus from our wastewater. The process of removing the phosphorus is done chemically by reacting the phosphorus with compounds such as ferric chloride, ferric sulfate, and aluminum sulfate or aluminum chlorohydrate. Phosphorus, when combined with aluminum or iron, becomes an insoluble salt. The solubility equilibrium constants of $FePO_4$ and $AlPO_4$ are 1.3x10-22 and 5.8x10-19, respectively. With solubilitys this low, the resulting precipitates can then be filtered out.
Another example of the dangers of phosphorus is in the production of matches. The flammable nature and cheap manufacturing of white phosphorus made it possible to easily make matches around the turn of the 20th century. However, white phosphorus is highly toxic. Many workers in match factories developed brain damage and a disease called "phosphorus necrosis of the jaw" from exposure to toxic phosphorus vapors. Excess phosphorus accumulation caused their bone tissue to die and rot away. For this reason, we now use red phosphorus or phosphorus sesquisulfide in "safety" matches.
Discovery of Phosphorus
Named from the Greek word phosphoros ("bringer of light"), elemental phosphorus is not found in its elemental form because this form is quite reactive. Because of this factor it took a long period of time for it to be "discovered". The first recorded isolation of phosphorus was by alchemist Hennig Brand in 1669, involving about 60 pails of urine. After letting a large amount of urine putrefy for a long time, Brand distilled the liquid to a paste, heated the paste, discarded the salt formed, and put the remaining substance under cold water to form solid white phosphorus. Brand's process was not very efficient; the salt he discarded actually contained most of the phosphorus. Nevertheless, he obtained some pure, elemental phosphorus for his efforts. Others of the time improved the efficiency of the process by adding sand, but still continued to discard the salt. Later, phosphorus was manufactured from bone ash. Currently, the process for manufacturing phosphorus does not involve large amounts of putrefied urine or bone ash. Instead, manufacturers use calcium phosphate and coke (Emsley).
Allotropes of Phosphorus
Phosphorus is a nonmetal, solid at room temperature, and a poor conductor of heat and electricity. Phosphorus occurs in at least 10 allotropic forms, the most common (and reactive) of which is so-called white (or yellow) phosphorus, which looks like a waxy solid or plastic. It is very reactive and will spontaneously inflame in air, so it is stored under water. The other common form of phosphorus is red phosphorus, which is much less reactive and is one of the components on the striking surface of a match book. Red phosphorus can be converted to white phosphorus by careful heating.
White phosphorus consists of $\ce{P4}$ molecules, whereas the crystal structure of red phosphorus has a complicated network of bonding. White phosphorus has to be stored in water to prevent natural combustion, but red phosphorus is stable in air.
When burned, red phosphorus also forms the same oxides as those obtained in the burning of white phosphosrus, $\ce{P4O6}$ when air supply is limited, and $\ce{P4O10}$ when sufficient air is present.
Diphosphorus (P2)
Diphosphorus ($P_2$) is the gaseous form of phosphorus that is thermodynamically stable above 1200 °C and until 2000 °C. It can be generated by heating white phosphorus (see below) to 1100 K and is very reactive with a bond-dissociation energy (117 kcal/mol or 490 kJ/mol) half that of dinitrogen ($N_2$).
White Phosphorus (P4)
White phosphorus (P4) has a tetrahedral structure. It is soft and waxy, but insoluble in water. Its glow occurs as a result of its vapors slowly being oxidized by the air. It is so thermodynamically unstable that it combusts in air. It was once used in fireworks and the U.S. military still uses it in incendiary bombs.
This Youtube video link shows various experiments with white phosphorus, which help show the physical and chemical properties of it. It also shows white phosphorus combusting with air.
Red Phosphorus and Violet Phosphorus (Polymeric)
Red Phosphorus has more atoms linked together in a network than white phosphorus does, which makes it much more stable. It is not quite as flammable, but given enough energy it still reacts with air. For this reason, we now use red phosphorus in matches.
Violet phosphorus is obtained from heating and crystallizing red phosphorus in a certain way. The phosphorus forms pentagonal "tubes".
Black Phosphorus (Polymeric)
Black phosphorus is the most stable form; the atoms are linked together in puckered sheets, like graphite. Because of these structural similarities black phosphorus is also flaky like graphite and possesses other similar properties.
Isotopes of Phosphorus
There are many isotopes of phosphorus, only one of which is stable (31P). The rest of the isotopes are radioactive with generally very short half-lives, which vary from a few nanoseconds to a few seconds. Two of the radioactive phosphorus isotopes have longer half-lives: 32P has a half-life of 14 days and 33P has a half-life of 25 days. These half-lives are long enough to be useful for analysis, and for this reason the isotopes can be used to mark DNA.
32P played an important role in the 1952 Hershey-Chase Experiment. In this experiment, Alfred Hershey and Martha Chase used radioactive isotopes of phosphorus and sulfur to determine that DNA was genetic material and not proteins. Sulfur can be found in proteins but not DNA, and phosphorus can be found in DNA but not proteins. This made phosphorus and sulfur effective markers of DNA and protein, respectively. The experiment was set up as follows: Hershey and Chase grew one sample of a virus in the presence of radioactive 35S and another sample of a virus in the presence of 32P. Then, they allowed both samples to infect bacteria. They blended the 35S and the 32P samples separately and centrifuged the two samples. Centrifuging separated the genetic material from the non-genetic material. The genetic material penetrated the solid that contained the bacterial cells at the bottom of the tube while the non-genetic material remained in the liquid. By analyzing their radioactive markers, Hershey and Chase found that the 32P remained with the bacteria, and the 35S remained in the supernatant liquid. These results were confirmed by further tests involving the radioactive phosphorus..
Phosphorus and Life
We get most elements from nature in the form of minerals. In nature, phosphorus exists in the form of phosphates. Rocks containing phosphate are fluoroapatite ($\ce{3Ca3(PO4)2.CaF2}$), chloroapaptite, ($\ce{3Ca3(PO4)2. CaCl2}$), and hydroxyapatite ($\ce{3Ca3(PO4)2. Ca(OH)2}$). These minerals are very similar to the bones and teeth. The arrangements of atoms and ions of bones and teeth are similar to those of the phosphate-containing rocks. In fact, when the $\ce{OH-}$ ions of the teeth are replaced by $\ce{F-}$, the teeth resist decay. This discovery led to a series of social and economical issues.
Nitrogen, phosphorus and potassium are key ingredients for plants, and their contents are key in all forms of fertilizers. From an industrial and economical viewpoint, phosphorus-containing compounds are important commodities. Thus, the chemistry of phosphorus has academic, commercial and industrial interests.
Chemistry of Phosphorus
As a member of the Nitrogen Family, Group 15 on the Periodic Table, phosphorus has 5 valence shell electrons available for bonding. Its valence shell configuration is 3s23p3. Phosphorus forms mostly covalent bonds. Any phosphorus rock can be used for the production of elemental phosphorus. Crushed phosphate rocks and sand ($\ce{SiO2}$) react at 1700 K to give phosphorus oxide, $\ce{P4O10}$:
$\ce{2 Ca3(PO4)2 + 6 SiO2 \rightarrow P4O10 + 6 CaSiO3} \label{1}$
$\ce{P4O10}$ can be reduced by carbon:
$\ce{P4O10 + 10 C \rightarrow P4 + 10 CO}. \label{2}$
Waxy solids of white phosphorus are molecular crystals consisting of $\ce{P4}$ molecules. They have an interesting property in that they undergo spontaneous combustion in air:
$\ce{P4 + 5 O2 \rightarrow P4O10} \label{3}$
The structure of $\ce{P4}$ can be understood by thinking of the electronic configuration (s2 p3) of $\ce{P}$ in bond formation. Sharing three electrons with other $\ce{P}$ atoms gives rise to the 6 $\ce{P-P}$ bonds, leaving a lone pair occupying the 4th position in a distorted tetrahedron.
When burned with insufficient oxygen, $\ce{P4O6}$ is formed:
$\ce{P4 + 3 O2 \rightarrow P4O6} \label{4}$
Into each of the $\ce{P-P}$ bonds, an $\ce{O}$ atom is inserted.
Burning phosphorus with excess oxygen results in the formation of $\ce{P4O10}$. An additional $\ce{O}$ atom is attached to the $\ce{P}$ directly:
$\ce{P4 + 5 O2 \rightarrow P4O10} \label{5}$
Thus, the oxides $\ce{P4O6}$ and $\ce{P4O10}$ share interesting features. Oxides of phosphorus, $\ce{P4O10}$, dissolve in water to give phosphoric acid,
$\ce{P4O10 + 6 H2O \rightarrow 4 H3PO4} \label{6}$
Phosphoric acid is a polyprotic acid, and it ionizes in three stages:
$\ce{H3PO4 \rightleftharpoons H+ + H2PO4-} \label{7a}$
$\ce{H2PO4- \rightleftharpoons H+ + HPO4^2-} \label{7b}$
$\ce{HPO4^2- \rightarrow H+ + PO4^3-} \label{7c}$
Phosphoric Acid
Phosphoric acid is a polyprotic acid, which makes it an ideal buffer. It gets harder and harder to separate the hydrogen from the phosphate, making the pKa values increase in basicity: 2.12, 7.21, and 12.67. The conjugate bases H2PO4-, HPO42-, and PO43- can be mixed to form buffer solutions.
Reaction Dissociation Constant
Table 1: Ionization constants for the successive deprotonation of phosphoric acid states
$H_3PO_4 + H_2O \rightarrow H_3O^+ + H_2PO^{4-}$ Ka1=7.5x10-3
$H_2PO_4^{-} + H_2O \rightarrow H_3O^+ + HPO_4^{2-}$ Ka2=6.2x10-8
$H_2PO_4^{-} + H_2O \rightarrow H_3O^+ + PO_4^{3-}$ Ka3=2.14x10-13
Overall: $H_3PO_4 + 3H_2O \rightarrow 3 H_3O^+ + PO_4^{3-}$
Past and Present Uses of Phosphorus
Commercially, phosphorus compounds are used in the manufacture of phosphoric acid ($H_3PO_4$) (found in soft drinks and used in fertilizer compounding). Other compounds find applications in fireworks and, of course, phosphorescent compounds which glow in the dark. Phosphorus compounds are currently used in foods, toothpaste, baking soda, matches, pesticides, nerve gases, and fertilizers. Phosphoric acid is not only used in buffer solutions; it is also a key ingredient of Coca Cola and other sodas! Phosphorus compounds were once used in detergents as a water softener until they raised concerns about pollution and eutrophication. Pure phosphorus was once prescribed as a medicine and an aphrodisiac until doctors realized it was poisonous (Emsley).
Questions
1. About 85% of the total industrial output of phosphoric acid is used
a. in the detergent industry
b. to produce buffer solutions
c. in the paint industry
d. to produce superphosphate fertilizers
e. in the manufacture of plastics
2. What is the product when phosphorus pentoxide $\ce{P4O10}$ reacts with water? Give the formula of the product.
3. What is the phosphorus-containing product when $\ce{PCl3}$ reacts with water? Give the formula.
Solutions
1. Answer... d
The middle number, (for example, 6-5-8) specifies the percentage of phosphorus compound in a fertilizer. Phosphorus is an important element for plant life.
2. Answer $\ce{H3PO4}$
$\ce{P4O10 + 6 H2O \rightarrow 4 H3PO4} \nonumber$
3. Answer $\ce{H3PO3}$
$\ce{PCl3 + 3 H2O \rightarrow H3PO3 + 3 HCl} \nonumber$
This is a weaker acid than $\ce{H3PO4}$.
Contributors
• Aimee Kindel (UCD), Kirenjot Grewal (UCD), Tiffany Lui (UCD)
• Chung (Peter) Chieh (Professor Emeritus, Chemistry @ University of Waterloo)
Z033 Chemistry of Arsenic (Z33)
Contributors and Attributions
• Freddie Chak, Jonathan Molina, Tiffany Lui (University of California, Davis)
Stephen R. Marsden
Z051 Chemistry of Antimony (Z51)
Antimony and its compounds have been known for centuries. Scientific study of the element began during the early 17th century, much of the important work being done by Nicolas Lemery. The name of the element comes from the Greek anti + monos for "not alone", while the modern symbol is rooted in the Latin-derived name of the common ore, stibnite. Antimony is a hard, brittle metalloid which is alloyed with other metals to increase hardness. It is also used in some semi-conductor devices. The recovery of elemental antimony parallels that of arsenic: the sulfide ore (stibnite) is roasted in air and then heated with carbon.
Antimony trisulfide
Antimony trisulfide, $Sb_2S_3$, is a sulfide mineral commonly called stibnite or antimonite. Antimony trisulfide exists as a gray/black crystalline solid (orthorhombic crystals) and an amorphous red-orange powder. It turns black due to oxidation by air. Antimony trisulfide is the most important source for antimony. It is insoluble in water and melts at 550°C. The chemical symbol of antimony (Sb) is derived from stibnite.
Amorphous (red to yellow-orange) antimony trisulfide can be prepared by treating an antimony trichloride solution with hydrogen sulfide:
$2 SbCl_3 + 3 H_2S \rightarrow Sb_2S_3 + 6 HCl \nonumber$
When antimony trisulfide is melted with iron at approximately 600°C, the following reaction yields elementary antimony:
$Sb_2S_3 + 3 Fe \rightarrow 2 Sb + 3 FeS \nonumber$
$Sb_2S_3$ is used as a pigment, in pyrotechnics (glitter and fountain mixtures) and on safety matches. In combination with antimony oxides it is also used as a yellow pigment in glass and porcelain. Antimony trisulfide photoconductors are used in vidicons for CCTV.
Contributors and Attributions
Stephen R. Marsden
Z083 Chemistry of Bismuth (Z83)
Bismuth, the heaviest non-radioactive naturally occurring element, was isolated by Basil Valentine in 1450. It is a hard, brittle metal with an unusually low melting point (271oC). Alloys of bismuth with other low-melting metals such as tin and lead have even lower melting points and are used in electrical solders, fuse elements and automatic fire sprinkler heads.
The metal can be found in nature, often combined with copper or lead ores, but can also be extracted from bismuth(III) oxide by roasting with carbon. Compounds of bismuth are used in pigments for oil paintings and one is in a popular pink preparation for the treatment of common stomach upset.
Contributors and Attributions
Stephen R. Marsden
Z115 Chemistry of Moscovium (Z115)
In studies announced jointly by the Joint Institute for Nuclear Research in Dubna, Russia, and the Lawrence Livermore National Laboratory in the U.S., four atoms of element 113 were produced in 2004 via decay of element 115 after the fusion of Ca-48 and Am-243.
Contributors and Attributions
• Stephen R. Marsden | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_15%3A_The_Nitrogen_Family/Z015_Chemistry_of_Phosphorous.txt |
The oxygen family, also called the chalcogens, consists of the elements found in Group 16 of the periodic table and is considered among the main group elements. It consists of the elements oxygen, sulfur, selenium, tellurium and polonium. These can be found in nature in both free and combined states.
• Group 16: General Properties and Reactions
The oxygen family, also called the chalcogens, consists of the elements found in Group 16 of the periodic table and is considered among the main group elements. It consists of the elements oxygen, sulfur, selenium, tellurium and polonium. These can be found in nature in both free and combined states. The group 16 elements are intimately related to life.
• Chemistry of Oxygen (Z=8)
Oxygen is an element that is widely known by the general public because of the large role it plays in sustaining life. Without oxygen, animals would be unable to breathe and would consequently die. Oxygen is not only important to supporting life, but also plays an important role in many other chemical reactions. Oxygen is the most common element in the earth's crust and makes up about 20% of the air we breathe.
• Chemistry of Sulfur (Z=16)
Sulfur is a chemical element that is represented with the chemical symbol "S" and the atomic number 16 on the periodic table. Because it is 0.0384% of the Earth's crust, sulfur is the seventeenth most abundant element following strontium. Sulfur also takes on many forms, which include elemental sulfur, organo-sulfur compounds in oil and coal, H2S(g) in natural gas, and mineral sulfides and sulfates.
• Chemistry of Selenium (Z=34)
Element number 34, selenium, was discovered by Swedish chemist Jons Jacob Berzelius in 1817. Selenium is a non-metal and can be compared chemically to its other non-metal counterparts found in Group 16: The Oxygen Family, such as sulfur and tellurium.
• Chemistry of Tellurium (Z=52)
Discovered by von Reichenstein in 1782, tellurium is a brittle metalloid that is relatively rare. It is named from the Latin tellus for "earth". Tellurium can be alloyed with some metals to increase their machinability and is a basic ingredient in the manufacture of blasting caps. Elemental tellurium is occasionally found in nature but is more often recovered from various gold ores.
• Chemistry of Polonium (Z=84)
Polonium was discovered in 1898 by Marie Curie and named for her native country of Poland. The discovery was made by extraction of the remaining radioactive components of pitchblende following the removal of uranium. There is only about 10-6 g per ton of ore! Current production for research purposes involves the synthesis of the element in the lab rather than its recovery from minerals. This is accomplished by producing Bi-210 from the abundant Bi-209.
• Chemistry of Livermorium (Z=116)
In May of 2012 the IUPAC approved the name "Livermorium" (symbol Lv) for element 116. The new name honors the Lawrence Livermore National Laboratory (1952). A group of researchers of this Laboratory with the heavy element research group of the Flerov Laboratory of Nuclear Reactions took part in the work carried out in Dubna on the synthesis of superheavy elements including element 116.
Thumbnail: A sample of sulfur a member of the oxygen group of elements. (Public Domain; Ben Mills).
Group 16: The Oxygen Family (The Chalcogens)
The oxygen family, also called the chalcogens, consists of the elements found in Group 16 of the periodic table and is considered among the main group elements. It consists of the elements oxygen, sulfur, selenium, tellurium and polonium. These can be found in nature in both free and combined states. The group 16 elements are intimately related to life. We need oxygen all the time throughout our lives. Did you know that sulfur is also one of the essential elements of life? It is responsible for some of the protein structures in all living organisms. Many industries utilize sulfur, but emission of sulfur compounds is often seen more as a problem than the natural phenomenon. The metallic properties of Group 16 elements increase as the atomic number increases. The element polonium has no stable isotopes, and the isotope with mass number 209 has the longest half life of 103 years.
Properties and Periodic Trends
Properties of oxygen are very different from those of other elements of the group, but they all have 2 elections in the outer s orbital, and 4 electrons in the p orbitals, usually written as s2p4.
The electron configurations for each element are given below:
• Oxygen: 1s2 2s2 2p4
• Sulfur: 1s2 2s2p6 3s2p4
• Selenium: 1s2 2s2p6 3s2p6d10 4s2p4
• Tellurium: 1s2 2s2p6 3s2p6d10 4s2p6d10 5s2p4
• Polonium: 1s2 2s2p6 3s2p6d10 4s2p6d10f14 5s2p6d10 6s2p4
Example \(1\): Polonium
Polonium can be written as [Xe] 6s2 4f14 5d10 6p4
The trends of the properties in this group are interesting. Knowing the trend allows us to predict their reactions with other elements. Most trends are true for all groups of elements, and the group trends are due mostly to the size of the atoms and number of electrons per atom. The trends are described below:
1. The metallic properties increase in the order oxygen, sulfur, selenium, tellurium, or polonium. Polonium is essentially a metal. It was discovered by M. Curie, who named it after her native country Poland.
2. Electronegativity, ionization energy (or ionization potential IP), and electron affinity decrease for the group as atomic weight increases.
3. The atomic radii and melting points increase.
4. Oxygen differs from sulfur in chemical properties due to its small size. The differences between \(\ce{O}\) and \(\ce{S}\) are more than the differences between other members.
Metallic character increases down the group, with tellurium classified as a metalloid and polonium as a metal. Melting point, boiling point, density, atomic radius, and ionic radius all increase down the group. Ionization energy decreases down the group. The most common oxidation state is -2; however, sulfur can also exist at a +4 and +6 state, and +2, +4, and +6 oxidation states are possible for Se, Te, and Po.
Table \(1\): Select properties of Group 16 elements
Oxygen Sulfur Selenium Tellurium Polonium
Boiling Pt (°C) -182.962 444.674 685 989.9 962
Ionization Energy (kJ/mol) 1314 1000 941 869 812
Ionic Radius (pm) 140 184 198 221
Oxygen
Oxygen is a gas at room temperature and 1 atm, and is colorless, odorless, and tasteless. It is the most abundant element by mass in both the Earth's crust and the human body. It is second to nitrogen as the most abundant element in the atmosphere. There are many commercial uses for oxygen gas, which is typically obtained through fractional distillation. It is used in the manufacture of iron, steel, and other chemicals. It is also used in water treatment, as an oxidizer in rocket fuel, for medicinal purposes, and in petroleum refining.
Oxygen has two allotropes, O2 and O3. In general, O2 (or dioxygen) is the form referred to when talking about the elemental or molecular form because it is the most common form of the element. The O2 bond is very strong, and oxygen can also form strong bonds with other elements. However, compounds that contain oxygen are considered to be more thermodynamically stable than O2.
The latter allotrope, ozone, is a pale-blue poisonous gas with a strong odor. It is a very good oxidizing agent, stronger than dioxygen, and can be used as a substitute for chlorine in purifying drinking water without giving the water an odd taste. However, because of its unstable nature it disappears and leaves the water unprotected from bacteria. Ozone at very high altitudes in the atmosphere is responsible for protecting the Earth's surface from ultraviolet radiation; however, at lower altitudes it becomes a major component of smog.
Oxygen's primary oxidation states are -2, -1, 0, and -1/2 (in O2-), but -2 is the most common. Typically, compounds with oxygen in this oxidation state are called oxides. When oxygen reacts with metals, it forms oxides that are mostly ionic in nature. These can dissolve in water and react to form hydroxides; they are therefore called basic anhydrides or basic oxides. Nonmetal oxides, which form covalent bonds, are simple molecules with low melting and boiling points.
Compounds with oxygen in an oxidation state of -1 are referred to as peroxides. Examples of this type of compound include \(Na_2O_2\) and \(BaO_2\). When oxygen has an oxidation state of -1/2, as in \(O_2^-\), the compound is called a superoxide.
Oxygen is rarely the central atom in a structure and can never bond with more than 4 elements due to its small size and its inability to create an expanded valence shell. Oxygen reacts with hydrogen to form water, which is extensively hydrogen-bonded, has a large dipole moment, and is considered a universal solvent.
There are a wide variety of oxygen-containing compounds, both organic and inorganic: oxides, peroxides and superoxides, alcohols, phenols, ethers, and carbonyl-containing compounds such as aldehydes, ketones, esters, amides, carbonates, carbamates, carboxylic acids and anhydrides.
Sulfur
Sulfur is a solid at room temperature and 1 atm pressure. It is usually yellow, tasteless, and nearly odorless. It is the sixteenth most abundant element in Earth's crust. It exists naturally in a variety of forms, including elemental sulfur, sulfides, sulfates, and organosulfur compounds. Since the 1890s, sulfur has been mined using the Frasch process, which is useful for recovering sulfur from deposits that are under water or quicksand. Sulfur produced from this process is used in a variety of ways including in vulcanizing rubber and as fungicide to protect grapes and strawberries.
Sulfur is unique in its ability to form a wide range of allotropes, more than any other element in the periodic table. The most common state is the solid S8 ring, as this is the most thermodynamically stable form at room temperature. Sulfur exists in the gaseous form in five different forms (S, S2, S4, S6, and S8). In order for sulfur to convert between these compounds, sufficient heat must be supplied.
Two common oxides of sulfur are sulfur dioxide (SO2) and sulfur trioxide (SO3). Sulfur dioxide is formed when sulfur is combusted in air, producing a toxic gas with a strong odor. These two compounds are used in the production of sulfuric acid, which is used in a variety of reactions. Sulfuric acid is one of the top manufactured chemicals in the United States, and is primarily used in the manufacture of fertilizers.
Sulfur also exhibits a wide range of oxidation states, with values ranging from -2 to +6. It is often the central ion in a compound and can easily bond with up to 6 atoms. In the presence of hydrogen it forms the compound hydrogen sulfide, H2S, a poisonous gas incapable of forming hydrogen bonds and with a very small dipole moment. Hydrogen sulfide can easily be recognized by its strong odor that is similar to that of rotten eggs, but this smell can only be detected at low, nontoxic concentrations. This reaction with hydrogen epitomizes how differently oxygen and sulfur act despite their common valence electron configuration and common nonmetallic properties.
A variety of sulfur-containing compounds exist, many of them organic. The prefix thio- in front of the name of an oxygen-containing compound means that the oxygen atom has been substituted with a sulfur atom. General categories of sulfur-containing compounds include thiols (mercaptans), thiophenols, organic sulfides (thioethers), disulfides, thiocarbonyls, thioesters, sulfoxides, sulfonyls, sulfamides, sulfonic acids, sulfonates, and sulfates.
Selenium
Selenium appears as a red or black amorphous solid, or a red or grey crystalline structure; the latter is the most stable. Selenium has properties very similar to those of sulfur; however, it is more metallic, though it is still classified as a nonmetal. It acts as a semiconductor and therefore is often used in the manufacture of rectifiers, which are devices that convert alternating currents to direct currents. Selenium is also photoconductive, which means that in the presence of light the electrical conductivity of selenium increases. It is also used in the drums of laser printers and copiers. In addition, it has found increased use now that lead has been removed from plumbing brasses.
It is rare to find selenium in its elemental form in nature; it must typically be removed through a refining process, usually involving copper. It is often found in soils and in plant tissues that have bioaccumulated the element. In large doses, the element is toxic; however, many animals require it as an essential micronutrient. Selenium atoms are found in the enzyme glutathione peroxidase, which destroys lipid-damaging peroxides. In the human body it is an essential cofactor in maintaining the function of the thyroid gland. In addition, some research has shown a correlation between selenium-deficient soils and an increased risk of contracting the HIV/AIDS virus.
Tellurium
Tellurium is the metalloid of the oxygen family, with a silvery white color and a metallic luster similar to that of tin at room temperature. Like selenium, it is also displays photoconductivity. Tellurium is an extremely rare element, and is most commonly found as a telluride of gold. It is often used in metallurgy in combination with copper, lead, and iron. In addition, it is used in solar panels and memory chips for computers. It is not toxic or carcinogenic; however, when humans are exposed to too much of it they develop a garlic-like smell on their breaths.
Polonium
Polonium is a very rare, radioactive metal. There are 33 different isotopes of the element and all of the isotopes are radioactive. It exists in a variety of states, and has two metallic allotropes. It dissolves easily into dilute acids. Polonium does not exist in nature in compounds, but it can form synthetic compounds in the laboratory. It is used as an alloy with beryllium to act as a neutron source for nuclear weapons.
Polonium is a highly toxic element. The radiation it emits makes it very dangerous to handle. It can be immediately lethal when applied at the correct dosage, or cause cancer if chronic exposure to the radiation occurs. Methods to treat humans who have been contaminated with polonium are still being researched, and it has been shown that chelation agents could possibly be used to decontaminate humans.
Problems
1. What properties increase down the oxygen family?
2. What element can form the most allotropes in the periodic table?
3. What is photoconductivity and which elements display this property?
4. Ozone (\(O_3\)) is a contributor to smog: True or False
5. How many electrons do elements of the oxygen family have in their outermost shell?
6. What does the term "peroxide" refer to?
7. How many elements in the oxygen family are metals, and which one(s)?
8. What is the most common oxidation state for elements in the oxygen family?
9. What is the most abundant element by mass in the Earth's crust and in the human body?
Solutions
1. Melting point, boiling point, density, atomic radius, and ionic radius.
2. Sulfur.
3. Photoconductivity is when the electrical conductivity of an element increases in the presence of light. Both selenium and tellurium display this property.
4. True.
5. Six.
6. A compound that contains oxygen in the oxidation state of -1.
7. 1; Polonium.
8. -2.
9. Oxygen. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_16%3A_The_Oxygen_Family_(The_Chalcogens)/1Group_16%3A_General_Propertie.txt |
Oxygen is an element that is widely known by the general public because of the large role it plays in sustaining life. Without oxygen, animals would be unable to breathe and would consequently die. Oxygen not only is important to supporting life, but also plays an important role in many other chemical reactions. Oxygen is the most common element in the earth's crust and makes up about 20% of the air we breathe. Historically the discovery of oxygen as an element essential for combustion stands at the heart of the phlogiston controversy (see below).
The Origin and History
Oxygen is found in the group 16 elements and is considered a chalcogen. Named from the Greek oxys + genes, "acid-former", oxygen was discovered in 1772 by Scheele and independently by Priestly in 1774. Oxygen was given its name by the French scientist, Antoine Lavoisier.
Scheele discovered oxygen through an experiment which involved burning manganese oxide. Scheele came to find that the hot manganese oxide produced a gas which he called "fire air". He also came to find that when this gas was able to come into contact with charcoal, it produced beautiful bright sparks. All of the other elements produced the same gas. Although Scheele discovered oxygen, he did not publish his work until three years after another chemist, Joseph Priestly, discovered oxygen. Joseph Priestly, an English chemist, repeated Scheele's experiment in 1774 using a slightly different setup. Priestly used a 12 in burning glass and aimed the sunlight directly towards the compound that he was testing, mercuric oxide. As a result, he was able to "discover better air" that was shown to expand a mouse's lifetime to four times as long and caused a flame to burn with higher intensity. Despite these experiments, both chemists were not able to pinpoint exactly what this element was. It was not until 1775 that Antoine Lavoisier, a French chemist, was able to recognize this unknown gas as an element.
Our atmosphere currently contains about 21% of free oxygen. Oxygen is produced in various ways. The process of photochemical dissociation in which water molecules are broken up by ultraviolet rays produces about 1-2% of our oxygen. Another process that produces oxygen is photosynthesis which is performed by plants and photosynthetic bacteria. Photosynthesis occurs through the following general reaction:
$\ce{CO2 + H2O + h\nu \rightarrow} \text{organic compounds} \ce{+ O2} \nonumber$
The Dangers of Phlogiston
Phlogiston theory is the outdated belief that a fire-like element called phlogiston is contained within combustible bodies and released during combustion. The name comes from the Ancient Greek φλογιστόν phlogistón (burning up), from φλόξ phlóx (flame). It was first stated in 1667 by Johann Joachim Becher, and then put together more formally by Georg Ernst Stahl. The theory attempted to explain burning processes such as combustion and rusting, which are now collectively known as oxidation.
Properties
• Element number: 8
• Atomic weight 15.9994
• Color: gas form- colorless, liquid- pale blue
• Melting point: 54.36K
• Boiling point: 90.2 K
• Density: .001429
• 21% of earth's atmosphere
• Third most abundant element in the universe
• Most abundant element in Earth's crust at 45.4%
• 3 Stable isotopes
• Ionization energy: 13.618 eV
• Oxygen is easily reduced and is a great oxidizing agent making it readily reactive with other elements
Magnetic Properties of Oxygen
Oxygen (O2) is paramagnetic. An oxygen molecule has six valence electrons, so the O2 molecule has 12 valence electrons with the electron configuration shown below:
As shown, there are two unpaired electrons, which causes O2 to be paramagnetic. There are also eight valence electrons in the bonding orbitals and four in antibonding orbitals, which makes the bond order 2. This accounts for the double covalent bond that is present in O2.
Video $1$: A chemical demonstration of the paramagnetism of molecular oxygen, as shown by the attraction of liquid oxygen to magnets.
As shown in Video $1$, since molecular oxygen ($O_2$) has unpaired electrons, it is paramagnetic and is attracted to the magnet. In contrast, molecular nitrogen ($N_2$) has no unpaired electrons and is not attracted to the magnet.
General Chemistry of Oxygen
Oxygen normally has an oxidation state of -2, but is capable of having oxidation states of -2, -1, -1/2, 0, +1, and +2. The oxidation states of oxides, peroxides and superoxides are as follows:
• Oxides: O-2 ,
• peroxides: O2-2 ,
• superoxide: O2-1.
Oxygen does not react with itself, nitrogen, or water under normal conditions. Oxygen does, however, dissolve in water at 20 degrees Celsius and 1 atmosphere. Oxygen also does not normally react with bases or acids. Group 1 metals (alkaline metals) are very reactive with oxygen and must be stored away from oxygen in order to prevent them from becoming oxidized. The metals at the bottom of the group are more reactive than those at the top. The reactions of a few of these metals are explored in more detail below.
Lithium: Reacts with oxygen to form white lithium oxide in the reaction below.
$\ce{4Li + O_2 \rightarrow 2Li_2O} \label{1}$
Sodium: Reacts with oxygen to form a white mixture of sodium oxide and sodium peroxide. The reactions are shown below.
• Sodium oxide: $\ce{4Na + O_2 \rightarrow 2Na_2O} \label{2}$
• Sodium peroxide: $\ce{2Na + O_2 \rightarrow Na_2O_2} \label{3}$
Potassium: Reacts with oxygen to form a mixture of potassium peroxide and potassium superoxide. The reactions are shown below.
• Potassium peroxide: $\ce{2K + O_2 \rightarrow 2K_2O_2} \label{4}$
• Potassium superoxide: $\ce{K + O_2 \rightarrow KO_2} \label{5}$
Rubidium and Cesium: Both metals react to produce superoxides through the same process as that of the potassium superoxide reaction.
The oxides of these metals form metal hydroxides when they react with water. These metal hydroxides make the solution basic or alkaline, hence the name alkaline metals.
Group 2 metals (alkaline earth metals) react with oxygen through the process of burning to form metal oxides but there are a few exceptions.
Beryllium is very difficult to burn because it has a layer of beryllium oxide on its surface which prevents further interaction with oxygen. Strontium and barium react with oxygen to form peroxides. The reaction of barium and oxygen is shown below, and the reaction with strontium would be the same.
$\ce{Ba(s) + O2 (g) \rightarrow BaO2 (s) }\label{6}$
Group 13 reacts with oxygen in order to form oxides and hydroxides that are of the form $X_2O_3$ and $X(OH)_3$. The variable X represents the various group 13 elements. As you go down the group, the oxides and hydroxides get increasingly basic.
Group 14 elements react with oxygen to form oxides. The oxides formed at the top of the group are more acidic than those at the bottom of the group. Oxygen reacts with silicon and carbon to form silicon dioxide and carbon dioxide. Carbon is also able to react with oxygen to form carbon monoxide, which is slightly acidic. Germanium, tin, and lead react with oxygen to form monoxides and dioxides that are amphoteric, which means that they react with both acids and bases.
Group 15 elements react with oxygen to form oxides. The most important are listed below.
• Nitrogen: N2O, NO, N2O3, N2O4, N2O5
• Phosphorus: P4O6, P4O8, P2O5
• Arsenic: As2O3, As2O5
• Antimony: Sb2O3, Sb2O5
• Bismuth: Bi2O3, Bi2O5
Group 16 elements react with oxygen to form various oxides. Some of the oxides are listed below.
• Sulfur: SO, SO2, SO3, S2O7
• Selenium: SeO2, SeO3
• Tellurium: TeO, TeO2, TeO3
• Polonium: PoO, PoO2, PoO3
Group 17 elements (halogens) fluorine, chlorine, bromine, and iodine react with oxygen to form oxides. Fluorine forms two oxides with oxygen: F2O and F2O2. Both fluorine oxides are called oxygen fluorides because fluorine is the more electronegative element. One of the fluorine reactions is shown below.
$\ce{O2 (g) + F2 (g) \rightarrow F2O2 (g)} \label{7}$
Group 18: Some would assume that the Noble Gases would not react with oxygen. However, xenon does react with oxygen to form $\ce{XeO_3}$ and $\ce{XeO_4}$. The ionization energy of xenon is low enough for the electronegative oxygen atom to "steal away" electrons. Unfortunately, $\ce{XeO_3}$ is HIGHLY unstable, and it has been known to spontaneously detonate in a clean, dry environment.
Transition metals react with oxygen to form metal oxides. However, gold, silver, and platinum do not react with oxygen. A few reactions involving transition metals are shown below:
$2Sn_{(s)} + O_{2(g)} \rightarrow 2SnO_{(s)} \label{8}$
$4Fe_{(s)} + 3O_{2(g)} \rightarrow 2Fe_2O_{3(s)} \label{9A}$
$4Al_{(s)} + 3O_{2(g)} \rightarrow 2Al_2O_{3(s)} \label{9B}$
Reaction of Oxides
We will be discussing metal oxides of the form $X_2O$. The variable $X$ represents any metal that is able to bond to oxygen to form an oxide.
• Reaction with water: The oxides react with water to form a metal hydroxide.
$X_2O + H_2O \rightarrow 2XOH \nonumber$
• Reaction with dilute acids: The oxides react with dilute acids to form a salt and water.
$X_2O + 2HCl \rightarrow 2XCl + H_2O \nonumber$
Reactions of Peroxides
The peroxides we will be discussing are of the form $X_2O_2$. The variable $X$ represents any metal that can form a peroxide with oxygen.
Reaction with water: If the temperature of the reaction is kept constant despite the fact that the reaction is exothermic, then the reaction proceeds as follows:
$X_2O_2+ 2H_2O \rightarrow 2XOH + H_2O_2 \nonumber$
If the reaction is not carried out at a constant temperature, then the reaction of the peroxide and water will result in decomposition of the hydrogen peroxide that is produced into water and oxygen.
Reaction with dilute acid: This reaction is more exothermic than that with water. The heat produced causes the hydrogen peroxide to decompose to water and oxygen. The reaction is shown below.
$X_2O_2 + 2HCl \rightarrow 2XCl + H_2O_2 \nonumber$
$2H_2O_2 \rightarrow 2H_2O + O_2 \nonumber$
Reaction of Superoxides
The superoxides we will be talking about are of the form $XO_2$, with $X$ representing any metal that forms a superoxide when reacting with oxygen.
Reaction with water: The superoxide and water react in a very exothermic reaction that is shown below. The heat that is produced in forming the hydrogen peroxide will cause the hydrogen peroxide to decompose to water and oxygen.
$2XO_2 + 2H_2O \rightarrow 2XOH + H_2O_2 + O_2 \nonumber$
Reaction with dilute acids: The superoxide and dilute acid react in a very exothermic reaction that is shown below. The heat produced will cause the hydrogen peroxide to decompose to water and oxygen.
$2XO_2 + 2HCl \rightarrow 2XCl + H_2O_2 + O_2 \nonumber$
Allotropes of Oxygen
There are two allotropes of oxygen; dioxygen (O2) and trioxygen (O3) which is called ozone. The reaction of converting dioxygen into ozone is very endothermic, causing it to occur rarely and only in the lower atmosphere. The reaction is shown below:
$3O_{2 (g)} \rightarrow 2O_{3 (g)} \;\;\; ΔH^o= +285 \;kJ \nonumber$
Ozone is unstable and quickly decomposes back to oxygen but is a great oxidizing agent.
Miscellaneous Reactions
Reaction with Alkanes: The most common reactions that involve alkanes occur with oxygen. Alkanes are able to burn and it is the process of oxidizing the hydrocarbons that makes them important as fuels. An example of an alkane reaction is the reaction of octane with oxygen as shown below.
C8H18(l) + 25/2 O2(g) → 8CO2(g) + 9H2O(l) ∆Ho = -5.48 X 103 kJ
Reaction with ammonia: Oxygen is able to react with ammonia to produce dinitrogen (N2) and water (H2O) through the reaction shown below.
$4NH_3 + 3O_2 \rightarrow 2N_2 + 6H_2O \nonumber$
Reaction with Nitrogen Oxide: Oxygen is able to react with nitrogen oxide in order to produce nitrogen dioxide through the reaction shown below.
$NO + O_2 \rightarrow NO_2 \nonumber$
Problems
1. Is oxygen reactive with noble gases?
2. Which transition metals does oxygen not react with?
3. What is produced when an oxide reacts with water?
4. Is oxygen reactive with alkali metals? Why are the alkali metals named that way?
5. If oxygen is reactive with alkali metals, are oxides, peroxides or superoxides produced?
Solutions
1. No, noble gases are unreactive with oxygen.
2. Oxygen is mostly unreactive with gold and platinum.
3. When an oxide reacts with water, a metal hydroxide is produced.
4. Oxygen is very reactive with alkali metals. Alkali metals are given the name alkali because the oxides of these metals react with water to form a metal hydroxide that is basic or alkaline.
5. Lithium produces an oxide, sodium produces a peroxide, and potassium, cesium, and rubidium produce superoxides.
Contributors and Attributions
• Phillip Ball (UCD), Katharine Williams (UCD) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_16%3A_The_Oxygen_Family_(The_Chalcogens)/Z008_Chemistry_of_Oxygen_%28Z8.txt |
Learning Objectives
• Describe the chemistry of the oxygen group.
• Give the trend of various properties.
• Remember the names of Group 16 elements.
• Explain the Frasch process.
• Describe properties and applications of $\mathrm{H_2SO_4}$.
• Explain properties and applications of $\mathrm{H_2S}$.
Sulfur is a chemical element that is represented with the chemical symbol "S" and the atomic number 16 on the periodic table. Because it is 0.0384% of the Earth's crust, sulfur is the seventeenth most abundant element following strontium. Sulfur also takes on many forms, which include elemental sulfur, organo-sulfur compounds in oil and coal, H2S(g) in natural gas, and mineral sulfides and sulfates. This element is extracted by using the Frasch process (discussed below), a method where superheated water and compressed air are used to draw liquid sulfur to the surface. Offshore sites, Texas, and Louisiana are the primary sites that yield extensive amounts of elemental sulfur. However, elemental sulfur can also be produced by reducing H2S, commonly found in oil and natural gas. For the most part, though, sulfur is used to produce SO2(g) and H2SO4.
Known from ancient times (mentioned in the Hebrew scriptures as brimstone) sulfur was classified as an element in 1777 by Lavoisier. Pure sulfur is tasteless and odorless with a light yellow color. Samples of sulfur often encountered in the lab have a noticeable odor. Sulfur is the tenth most abundant element in the known universe.
Sulfur at a Glance
Atomic Number 16
Atomic Symbol S
Atomic Weight 32.07 grams per mole
Structure orthorhombic
Phase at room temperature solid
Classification nonmetal
Physical Properties of Sulfur
Sulfur has an atomic weight of 32.066 grams per mole and is part of group 16, the oxygen family. It is a nonmetal and has a specific heat of 0.706 J g-1 oC-1. The electron affinity is 200 kJ mol-1 and the electronegativity is 2.58 (unitless). Sulfur is typically found as a light-yellow, opaque, and brittle solid in large amounts of small orthorhombic crystals. Not only does sulfur have twice the density of water, but it is also insoluble in water. On the other hand, sulfur is highly soluble in carbon disulfide and slightly soluble in many common solvents. Sulfur can also vary in color and blackens upon boiling due to carbonaceous impurities. Even as little as 0.05% of carbonaceous matter darkens sulfur significantly.
Most sulfur is recovered directly as the element from underground deposits by injecting super-heated water and piping out molten sulfur (sulfur melts at 112o C). Compared to other elements, sulfur has the most allotropes. While the S8 ring is the most common allotrope, there are 6 other structures with up to 20 sulfur atoms per ring.
• Under appropriate conditions, sulfur vapor can contain $S$, $S_2$, $S_4$, $S_6$, and $S_8$.
• At room temperature, rhombic sulfur (Sα) is a stable solid comprising cyclic $S_8$ molecules.
• At 95.5 °C, rhombic sulfur becomes monoclinic sulfur (Sβ). The crystal structure found in monoclinic sulfur differs from that of rhombic sulfur. Monoclinic sulfur is also made up of $S_8$molecules.
• Monoclinic sulfur becomes liquid sulfur (Sλ) at 119 °C. Liquid sulfur is a straw-colored liquid made up of $S_8$ molecules and other cyclic molecules containing a range of six to twenty atoms.
• At 160 oC, this becomes a dark, viscous liquid called liquid sulfur (Sμ). The molecules are still made up of eight sulfur atoms, but the molecule opens up and transforms from a circle into a long spiral-chain molecule.
• At 180 °C, the chain length and viscosity reach their maximum. Chains break and viscosity decreases at temperatures that exceed 180 °C.
• Sulfur vapor is produced when liquid boils at 445 °C. In the vapor that is produced, $S_8$ molecules dominate, but as the vapor continues to heat up, the molecules break up into smaller groups of sulfur atoms.
• To produce plastic sulfur, S is poured into cold water. Plastic sulfur is rubberlike and is made up of long, spiral-chain molecules. If plastic sulfur sits for long, it will reconvert to rhombic sulfur.
While oxygen has fewer allotropes than sulfur, including $\ce{O}$, $\ce{O_2}$, $\ce{O_3}$, $\ce{O_4}$, $\ce{O_8}$, metallic $\ce{O}$ (and four other solid phases), many of these actually have a corresponding sulfur variant. However, sulfur has more tendency to catenate (the linkage of atoms of the same element into longer chains). Here are the values of the single and double bond enthalpies:
$\begin{array}{c|r} \ce {O-O} & \ce{142\ kJ/mol} \ \ce {S–S} & \ce{268\ kJ/mol} \ \ce {O=O} & \ce{499\ kJ/mol} \ \ce {S=S} & \ce{352\ kJ/mol} \ \end{array} \nonumber$
This means that $\ce{O=O}$ is stronger than $\ce{S=S}$, while $\ce{O–O}$ is weaker than $\ce{S–S}$. So, in sulfur, single bonds are favored and catenation is easier than in oxygen compounds. It seems that the reason for the weaker $\ce{S=S}$ double bonds has its roots in the size of the atom: it's harder for the two atoms to come to a small enough distance, so that the $p$ orbital overlap is small and the $\pi$ bond is weak. This is attested by looking down the periodic table: $\ce{Se=Se}$ has an even weaker bond enthalpy of $\ce{272 kJ/mol}$.
What happens when the solid sulfur melts? The $\ce{S8}$ molecules break up. When suddenly cooled, long chain molecules are formed in the plastic sulfur which behave like rubber. Plastic sulfur transforms into rhombic sulfur over time.
Compounds
Reading the following reactions, figure out and notice the change of the oxidation state of $\ce{S}$ in the reactants and products. Common oxidation states of sulfur are -2, 0, +4, and +6. Sulfur (brimstone, stone that burns) reacts with $\ce{O2}$ giving a blue flame (Figure $1$):
$\ce{S + O_2 \rightarrow SO_2} \nonumber$
$\ce{SO2}$ is produced whenever a metal sulfide is oxidized. It is recovered and oxidized further to give $\mathrm{SO_3}$, for production of $\mathrm{H_2SO_4}$. $\mathrm{SO_2}$ reacts with $\mathrm{H_2S}$ to form $\mathrm{H_2O}$ and $\ce{S}$.
$\mathrm{2 SO_2 + O_2 \rightleftharpoons 2 SO_3} \nonumber$
$\mathrm{SO_3 + H_2O \rightleftharpoons H_2SO_4} \;\;(\text{a valuable commodity}) \nonumber$
$\mathrm{SO_3 + H_2SO_4 \rightleftharpoons H_2S_2O_7} \;\;\; (\text{pyrosulfuric acid}) \nonumber$
Sulfur reacts with sulfite ions in solution to form thiosulfate,
$\ce{S + SO_3^{2-} -> S_2O_3^{2-}} \nonumber$
but the reaction is reversed in an acidic solution.
Oxides
There are many different stable sulfur oxides, but the two that are commonly found are sulfur dioxide and sulfur trioxide. Sulfur dioxide is a commonly found oxide of sulfur. It is a colorless, pungent, and nonflammable gas. It has a density of 2.8 kg/m3 and a melting point of -72.5 oC. Because organic materials are more soluble in $SO_2$ than in water, the liquid form is a good solvent. $SO_2$ is primarily used to produce $SO_3$. The direct combustion of sulfur and the roasting of metal sulfides yield $SO_2$ via the contact process:
$\underbrace{S(s) + O_2(g) \rightarrow SO_2(g)}_{\text{Direct combustion}} \nonumber$
$\underbrace{2 ZnS(s) + 3 O_2(g) \rightarrow 2 ZnO(s) + 2 SO_2(g)}_{\text{Roasting of metal sulfides}} \nonumber$
Sulfur trioxide is another one of the commonly found oxides of sulfur. It is a colorless liquid with a melting point of 16.9 oC and a density of kg/m3. $SO_3$ is used to produce sulfuric acid. $SO_2$ is used in the synthesis of $SO_3$:
$\underbrace{2 SO_2 (g) + O_2(g) \rightleftharpoons 2 SO_3(g)}_{\text{Exothermic, reversible reaction}} \nonumber$
This reaction needs a catalyst to be completed in a reasonable amount of time with $V_2O_5$ being the catalyst most commonly used.
Hydrogen Sulfide H2S
• Hydrogen sulfide, $\ce{H2S}$, is a diprotic acid. The equilibria below, $\mathrm{H_2S \rightleftharpoons HS^- + H^+} \nonumber$ $\mathrm{HS^- \rightleftharpoons S^{2-} + H^+} \nonumber$ have been discussed in connection with Polyprotic Acids.
Other Sulfur-containing Compounds
Perhaps the most significant compound of sulfur used in modern industrialized societies is sulfuric acid ($H_2SO_4$). Sulfur dioxide ($SO_2$) finds practical applications in bleaching and refrigeration but it is also a nuisance gas resulting from the burning of sulfurous coals. Sulfur dioxide gas then reacts with the water vapor in the air to produce a weak acid, sulfurous acid ($H_2SO_3$), which contributes to the acid rain problem.
• Sulfuric acid, H2SO4, is produced by reacting $SO_3$ with water. However, this often leads to pollution problems. SO3(g) is reacted with 98% H2SO4 in towers full of ceramic material to produce H2S2O7 or oleum. Water is circulated in the tower to maintain the correct concentration and the acid is diluted with water at the end in order to produce the correct concentration. Pure sulfuric acid has no color and odor, and it is an oily, hygroscopic liquid. However, sulfuric acid vapor produces heavy, white smoke and a suffocating odor.
• Dilute sulfuric acid, H2SO4(aq), reacts with metals and acts as a strong acid in common chemical reactions. It is used to produce H2(g) and liberate CO2(g) and can neutralize strong bases.
• Concentrated sulfuric acid, H2SO4 (conc.), has a strong affinity for water. In some cases, it removes H and O atoms. Concentrated sulfuric acid is also a good oxidizing agent and reacts with some metals.
$C_{12}H_{22}O_{11}(s) \rightarrow 12 C(s) + 11 H_2O(l) \nonumber$
(Concentrated sulfuric acid used in forward reaction to remove H and O atoms.)
Applications of Sulfuric Acid
• as a strong acid for making $\ce{HCl}$ and $\mathrm{HNO_3}$.
• as an oxidizing agent for metals.
• as a dehydrating agent.
• for manufacture of fertilizer and other commodities.
• Sulfurous acid (H2SO3) is produced when $SO_2$(g) reacts with water. It cannot be isolated in its pure form; however, it forms salts as sulfites. Sulfites can act as both reducing agents and oxidizing agents.
O2(g) + 2 SO32-(aq) $\rightarrow$ 2 SO42- (aq) (Reducing agent)
2 H2S(g) + 2 H+(aq) + SO32-(aq) $\rightarrow$ 3 H2O(l) + 3 S(s) (Oxidizing agent)
H2SO3 is a diprotic acid that acts as a weak acid in both steps, and H2SO4 is also a diprotic acid but acts as a strong acid in the first step and a weak acid in the second step. Acids like NaHSO3 and NaHSO4 are called acid salts because they are the product of the first step of these diprotic acids.
Boiling elemental sulfur in a solution of sodium sulfite yields thiosulfate. Not only are thiosulfates important in photographic processing, but they are also common analytical reagents used with iodine (like in the following two reactions).
$2 Cu^{2+}_{(aq)} + 5 I^-_{(aq)} \rightarrow 2 CuI_{(s)} + I^-_{3(aq)} \nonumber$
$I^-_{3(aq)} + 2 S_2O^{2-}_{3(aq)} \rightarrow 3 I^-_{(aq)} + S_4O^{2-}_{6(aq)} \nonumber$
with excess triiodide ion titrated with Na2S2O3(aq).
Other than sulfuric acid, perhaps the most familiar compound of sulfur in the chemistry lab is the foul-smelling hydrogen sulfide gas, $H_2S$, which smells like rotten eggs.
• Sulfur halides are compounds formed between sulfur and the halogens. Common compounds include SF2, S2F2, SF4, and SF6. While SF4 is a powerful fluorinating agent, SF6 is a colorless, odorless, unreactive gas. Compounds formed by sulfur and chloride include S2Cl2, SCl4, and SCl2. SCl2 is a red, bad-smelling liquid that is utilized to produce mustard gas ($S(CH_2CH_2Cl)_2$).
$SCl_2 + 2CH_2CH_2 \rightarrow S(CH_2CH_2Cl)_2 \nonumber$
Production -The Frasch Process
Sulfur can be mined by the Frasch process. This process has made sulfur a high purity (up to 99.9 percent pure) chemical commodity in large quantities. Most sulfur-containing minerals are metal sulfides, and the best known is perhaps pyrite ($\mathrm{FeS_2}$, known as fool's gold because of its golden color). The most common sulfate-containing mineral is gypsum, $\mathrm{CaSO_4 \cdot 2H_2O}$, also known as plaster of Paris.
The Frasch process is based on the fact that sulfur has a comparatively low melting point. The process forces (99.5% pure) sulfur out by using hot water and air. In this process, superheated water is forced down the outermost of three concentric pipes. Compressed air is pumped down the center tube, and a mixture of elemental sulfur, hot water, and air comes up the middle pipe. Sulfur is melted with superheated water (at 170 °C under high pressure) and forced to the surface of the earth as a slurry.
Sulfur is mostly used for the production of sulfuric acid, $\ce{H2SO4}$. Most sulfur mined by the Frasch process is used in industry for the manufacture of sulfuric acid. Sulfuric acid, the most abundantly produced chemical in the United States, is manufactured by the contact process. Most (about 70%) of the sulfuric acid produced in the world is used in the fertilizer industry. Sulfuric acid can act as a strong acid, a dehydrating agent, and an oxidizing agent. Its applications use these properties. Sulfur is an essential element of life in sulfur-containing proteins.
Applications
Sulfur has many practical applications. As a fungicide, sulfur is used to counteract apple scab in organically farmed apple production. Other crops that utilize sulfur fungicides include grapes, strawberries, and many vegetables. In general, sulfur is effective against mildew diseases and black spot. Sulfur can also be used as an organic insecticide. Sulfites are frequently used to bleach paper and preserve dried fruit.
The vulcanization of rubber includes the use of sulfur as well. Cellophane and rayon are produced with carbon disulfide, a product of sulfur and methane. Sulfur compounds can also be found in detergents, acne treatments, and agrichemicals. Magnesium sulfate (epsom salt) has many uses, ranging from bath additives to exfoliants. Sulfur is being increasingly used as a fertilizer as well. Because standard sulfur is hydrophobic, it is covered with a surfactant by bacteria before oxidation can occur. Sulfur is therefore a slow-release fertilizer. Lastly, sulfur functions as a light-generating medium in sulfur lamps.
Concentrated sulfuric acid was once one of the most produced chemicals in the United States; the majority of the H2SO4 that is now produced is used in fertilizer. It is also used in oil refining, production of titanium dioxide, and in emergency power supplies and car batteries. The mineral gypsum, or calcium sulfate dihydrate, is used in making plaster of Paris. Over one million tons of aluminum sulfate is produced each year in the United States by reacting H2SO4 and Al2O3. This compound is important in water purification. Copper sulfate is used in electroplating. Sulfites are used in the paper making industry because they produce a substance that coats the cellulose in the wood and frees the fibers of the wood for treatment.
Emissions and the Environment
Particles, SO2(g), and H2SO4 mist are the components of industrial smog. Because power plants burn coal or high-sulfur fuel oils, SO2(g) is released into the air. When catalyzed on the surfaces of airborne particles, SO2 can be oxidized to SO3. A reaction with NO2 works as well as shown in the following reaction:
$SO_{2(g)} + NO_{2(g)} \rightarrow SO_{3(g)} + NO_{(g)} \nonumber$
H2SO4 mist is then produced after SO3 reacts with water vapor in the air. If H2SO4 reacts with airborne NH3, (NH4)2SO4 is produced. When SO2(g) and H2SO4 reach levels that exceed 0.10 ppm, they are potentially harmful. By removing sulfur from fuels and controlling emissions, acid rain and industrial smog can be kept under control. Processes like fluidized bed combustion have been presented to remove SO2 from smokestack gases.
Outside Links
• Dhawale, S.W. "Thiosulfate: An interesting sulfur oxoanion that is useful in both medicine and industry--but is implicated in corrosion." J. Chem. Educ. 1993, 70, 12.
• Lebowitz, Samuel H. "A demonstration working model of the Frasch process for mining sulfur." J. Chem. Educ. 1931, 8, 1630.
• Nagel, Miriam C. "Herman Frasch, sulfur king (PROFILES)." J. Chem. Educ. 1981, 58, 60.
• Riethmiller, Steven. "Charles H. Winston and Confederate Sulfuric Acid." J. Chem. Educ. 1995 72 575.
• Sharma, B. D. "Allotropes and polymorphs." J. Chem. Educ. 1987, 64, 404.
• Silverstein, Todd P.; Zhang, Yi. "Sugar Dehydration without Sulfuric Acid: No More Choking Fumes in the Classroom!" J. Chem. Educ. 1998 75 748.
• Tykodi, R. J. "In praise of thiosulfate." J. Chem. Educ. 1990, 67, 146.
• Thomas Jefferson National Accelerator Facility - Office of Science Education."It's Elemental-The Element Sulfur." Jefferson Lab.
• Sulfur's Electron Shell
Problems
1. Draw a diagram that summarizes the allotropy of sulfur. Use symbols, arrows, and numbers.
2. Direct combustion of sulfur is the only method for producing SO2(g). True or False.
3. Sulfites are not oxidizing agents. They are good reducing agents. True or False.
4. Give the reaction for the production of sulfur trioxide.
5. Choose the incorrect statement.
1. Sulfur produces cellophane and rayon.
2. Standard sulfur is hydrophobic.
3. SO2 can oxidize to SO3
4. Sulfur influences the development of acid rain and industrial smog.
5. All of the above are correct.
6. Which reaction is responsible for the destruction of limestone and marble statues and buildings?
1. $\ce{CaCO3 \rightarrow CaO + CO2}$
2. $\ce{SO2 + H2O \rightarrow H2SO3}$
3. $\ce{BaO + CO2 \rightarrow BaCO3 \rightarrow BaSO3}$ upon reaction with $\ce{SO2}$
4. $\ce{CaCO3 + H2O \rightarrow Ca(OH)2 + CO2}$
5. $\ce{CaCO3 + SO2 \rightarrow CaSO3 + CO2 \rightarrow CaSO4}$ upon oxidation
7. Give the formula of thiosulfate ion.
8. What is the oxidation state of $\ce{S}$ in $\ce{SF6}$, $\ce{H2SO4}$, $\ce{NaHSO4}$, $\ce{SO4^2-}$, and $\ce{SO3}$?
9. What is the phase of sulfur at 298 K? Enter the type of crystals.
10. Give the name of the process by which sulfur is forced out of the ground using hot water and air.
Solutions
1. The diagram may be drawn in any way. However, the symbols (S2), (S4), (S6), (S?), and (S8(g)) must be included. The temperatures should be written next to the arrows.
2. False
3. False
4. $2 SO_{2(g)} + O_{2(g)} \rightarrow 2 SO_{3(g)}$
5. A
6. e.
Consider...
$\ce{SO2}$ in $\ce{H2SO3}$ is the acid in acid rain, which attacks $\ce{CaCO3}$, marble. $\ce{SO2}$ reduces pigments in organic matter.
7. $\ce{S2O3^2-}$
Consider...
Sulfate is $\ce{SO4^2-}$; replacement of an $\ce{O}$ by an $\ce{S}$ gives thiosulfate $\ce{S2O3^2-}$. The two $\ce{S}$ in $\ce{S2O3^2-}$ have different oxidation states: one is +6, the other is (-2), average +2.
8. 6
Consider...
Oxidation state for $\ce{S}$ in $\ce{H2SO3}$, $\ce{SO3^2-}$, $\ce{SO2}$, etc. is 4. The oxidation state of $\ce{S}$ is the same for all in the list.
9. rhombic sulfur
Consider...
The term rhombic describes a type of crystal. Monoclinic sulfur is meta stable at 298 K.
10. Frasch process
Consider...
The Frasch process is used to mine elemental sulfur. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_16%3A_The_Oxygen_Family_(The_Chalcogens)/Z016_Chemistry_of_Sulfur_%28Z1.txt |
Element number 34, selenium, was discovered by Swedish chemist Jons Jacob Berzelius in 1817. Selenium is a non-metal and can be compared chemically to its other non-metal counterparts found in Group 16: The Oxygen Family, such as sulfur and tellurium.
Properties
Chemical Symbol: Se
Atomic Number: 34
Atomic Weight: 78.96
Electron Configuration: [Ar] 4s23d104p4
Melting Point: 493.65 K
Boiling Point: 958 K
Electronegativity: 2.55 (Pauling)
Oxidation States: Se-2, Se+6, Se+4
Ionization Energies: First: 941 kJ/mol
Second: 2045 kJ/mol
Third: 2973.7 kJ/mol
History
Selenium was discovered by Berzelius in 1818. It is named for the Greek word for "moon", selene. The discovery of selenium was an important finding, but at the same time seemingly accidental. Fellow scientist Martin Klaproth discovered a contamination of sulfuric acid creating a red colored product which he believed to be due to the element tellurium. However, Berzelius went on to further analyze the impurity and came to the conclusion that it was an unknown element that shared properties similar to those of tellurium. Based on the Greek word “selene,” meaning moon, Jons Berzelius decided to call the newly found element selenium.
Allotropes and Physical Properties
Selenium can exist in multiple allotropes that are essentially different molecular forms of an element with varying physical properties. For example, one allotrope of selenium can be seen as an amphorous (“without crystalline shape”) red powder. Selenium also takes a crystalline hexagonal structure, forming a metallic gray allotrope which is known to be stable. The most thermodynamically stable allotrope of selenium is trigonal selenium, which also appears as a gray solid. Most selenium is recovered from the electrolytic copper refining process. This is usually in the form of the red allotrope.
Selenium is mostly noted for its important chemical properties, especially those dealing with electricity. Unlike sulfur, selenium is a semiconductor, meaning that it conducts some electricity, but not as well as conductors. Selenium is a photoconductor, which means it has the ability to change light energy into electrical energy. Not only is selenium able to convert light energy into electrical energy, but it also displays the property of photoconductivity. Photoconductivity is the idea that the electrical conductivity of selenium increases due to the presence of light -- or, in other words, it becomes a better photoconductor as light intensity increases.
Isotopes
Isotopes of an element are atoms that have the same atomic numbers but a different number of neutrons (different mass numbers) in their nuclei. Selenium is known to have over 20 different isotopes; however, only 5 of them are stable. The five stable isotopes of selenium are 74Se, 76Se, 77Se, 78Se, 80Se.
Uses
Due to selenium’s property of photoconductivity, it is known to be used in photocells, exposure meters in photography, and also in solar cells. Selenium can also be seen in the products of plain-paper photocopiers, laser printers and photographic toners. Besides its uses in the electronic industry, selenium is also popular in the glass-making industry. When selenium is added to glass, it is able to negate the color of other elements found in the glass and essentially decolorizes it. Selenium is also able to create a ruby-red colored glass when added. The element can also be used in the production of alloys and is an additive to stainless steel.
Health Hazards
Selenium, a trace element, is important in the diet and health of both plants and animals, but can be only taken in very small amounts. Exposure to an excess amount of selenium is known to be toxic and causes health problems. With an upper intake level of 400 micrograms per day that can be tolerated, too much selenium can lead to selenosis and may result in health problems and even death. Compounds of selenium are also known to be carcinogenic.
Chemical Reactivity
Reaction with hydrogen
Selenium forms hydrogen selenide, H2Se, a colorless flammable gas, when reacted with hydrogen.
Reaction with oxygen
Selenium burns in air, displaying a blue flame, and forms solid selenium dioxide.
$Se_{8(s)} + 8O_{2(g)} \rightarrow 8SeO_{2(s)} \nonumber$
Selenium is also known to form selenium trioxide, SeO3.
Reaction with halides
Selenium reacts with fluorine, F2, and burns to form selenium hexafluoride.
$Se_{8(s)} + 24F_{2(g)} \rightarrow 8SeF_{6(l)} \nonumber$
Selenium also reacts with chlorine and bromine to form diselenium dichloride, $Se_2Cl_2$ and diselenium dibromide, $Se_2Br_2$.
$Se_8 + 4Cl_2 \rightarrow 4Se_2Cl_{2(l)} \nonumber$
$Se_8 + 4Br_2 \rightarrow 4Se_2Br_{2(l)} \nonumber$
Selenium also forms $SeF_4$, $SeCl_2$ and $SeCl_4$.
Selenides
Selenium reacts with metals to form selenides. Example: Aluminum selenide
$3 Se_8 + 16 Al \rightarrow 8 Al_2Se_3 \nonumber$
Selenites
Selenium reacts to form salts called selenites, e.g., silver selenite (Ag2SeO3) and sodium selenite (Na2SeO3).
Problems
1. Describe selenium’s property of photoconductivity.
2. Does selenium react with hydrogen? If so, what compound is produced?
3. Describe selenium’s purpose as a trace element.
4. What are some common uses for selenium?
5. Does selenium react with oxygen?
Solutions
1. Selenium’s ability to change light energy into electrical energy increases as light intensity increases.
2. Yes, selenium reacts with hydrogen and forms hydrogen selenide, H2Se.
3. Selenium is important to the health of plants and animals, but is only safe in small amounts. Too much selenium can be toxic and cause serious health problems.
4. Selenium is used in the glass-making industry and also in electronics. It is used in photo cells, solar cells, photocopiers, laser printers and also photographic toners.
5. Selenium burns in air and forms selenium dioxide. It is also able to form selenium trioxide.
Reference
1. Minaev, V. S., S. P. Timoshenkov, and V. V. Kalugin. "Structural and Phase Transformations in Condensed Selenium." Journal of Optoelectronics and Advanced Materials, volume 7, number 4, 2005, pp. 1717–1741.
2. Mary Elvira Weeks and Henry M. Leicester. Discovery of the Elements, 7th edition. Easton, PA: Journal of Chemical Education, 1968.
3. Petrucci, Ralph H. General Chemistry. 9th ed. Upper Saddle River: Prentice Hall, 2007
Contributors and Attributions
• David Jin (UCD) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_16%3A_The_Oxygen_Family_(The_Chalcogens)/Z034_Chemistry_of_Selenium_%28.txt |
Discovered by von Reichenstein in 1782, tellurium is a brittle metalloid that is relatively rare. It is named from the Latin tellus for "earth". Tellurium can be alloyed with some metals to increase their machinability and is a basic ingredient in the manufacture of blasting caps. Elemental tellurium is occasionally found in nature but is more often recovered from various gold ores, all containing $AuTe_2$.
History
Tellurium was discovered in a gold ore from the mines in Zlatna, near present day Sibiu, Transylvania. The ore was known as "Faczebajer weißes blättriges Golderz" (white leafy gold ore from Faczebaja) or antimonalischer Goldkies (antimonic gold pyrite). In 1782, while serving as the Austrian chief inspector of mines in Transylvania, Franz-Joseph Müller von Reichenstein concluded that a certain ore did not contain antimony, but that it contained bismuth sulfide. However, the following year, he reported that this was erroneous and that the ore contained mostly gold and an unknown metal very similar to antimony. After 3 years of testing, Müller determined the specific gravity of the mineral and noted the radish-like odor of the white smoke, which passed off when the new metal was heated. In 1789, another Hungarian scientist, Pál Kitaibel, also discovered the element independently in an ore from Deutsch-Pilsen which had been regarded as argentiferous molybdenite, but later he gave the credit to Müller. In 1798, it was named by Martin Heinrich Klaproth, who earlier isolated it from the mineral calaverite.
Properties
Tellurium is a semimetallic, lustrous, crystalline, brittle, silver-white element. It is usually available as a dark grey powder and has metal and non-metal properties. Te forms many compounds corresponding to those of sulfur and selenium. When burned in the air, tellurium has a greenish-blue flame and forms tellurium dioxide as a result. Tellurium is unaffected by water or hydrochloric acid, but dissolves in nitric acid. It has an atomic mass of 127.6 g/mol-1 and a density of 6.24 g-cm-3. Its boiling point is 450 degrees Celsius and its melting point is 1390 °C.
Source and Abundance
There are eight naturally occurring isotopes of tellurium, of which three are radioactive. Tellurium is among the rarest stable solid elements in the Earth's crust. At 0.005 ppm, it is comparable to platinum in abundance. However, tellurium is far more abundant in the wider universe. Tellurium was originally and is most commonly found in gold tellurides. However, the largest source for modern production of tellurium is as a byproduct of blister copper refinement. The treatment of 500 tons of copper ore results in 0.45 kg of tellurium. Tellurium can also be found in lead deposits. Other tellurium sources, known as subeconomic deposits because the cost of abstraction outweighs the yield in tellurium, are lower-grade copper and some coal.
Originally, the copper tellurium ore is treated with sodium bicarbonate and elemental oxygen to produce a tellurium oxide salt, copper oxide, and carbon dioxide:
$Cu_2Te + Na_2CO_3 + 2O_2 \rightarrow 2CuO + Na_2TeO_3 + CO_2 \nonumber$
Then, the sodium tellurium oxide is treated with sulfuric acid to precipitate out tellurium dioxide, which can be treated with aqueous sodium hydroxide to reduce to pure tellurium and oxygen gas:
$TeO_2 + 2NaOH \rightarrow Na_2TeO_3 + H_2O \rightarrow Te + 2NaOH + O_2 \nonumber$
Industrial and Commercial Use
Tellurium has many unique industrial and commercial uses that improve product quality and quality-of-life. Many of the technologies that utilize tellurium have important uses for the energy industry, the military, and health industries. Tellurium is used to color glass and ceramics and can improve the machining quality of metal products. When added to copper alloys, tellurium makes the alloy more ductile, whereas it can prevent corrosion in lead products. Tellurium is an important component of infrared detectors used by the military as well as x-ray detectors used by a variety of fields including medicine, science, and security. In addition, tellurium-based catalysts are used to produce higher-quality rubber. CdTe films are one of the highest efficiency photovoltaics, metals that convert sunlight directly into electrical power, at 11-13% efficiency and are, therefore, widely used in solar panels. CdTe is a thin-film semiconductor that absorbs sunlight.
Tellurium can be replaced by other elements in some of its uses. For many metallurgical uses, selenium, bismuth, or lead are effective substitutes. Both selenium and sulfur can replace tellurium in rubber production. Technologies based on tellurium have global impacts. As a photovoltaic, CdTe is the second most utilized solar cell in the world, soon said to surpass crystalline silicon and become the first. According to the US military, tellurium-based infrared detectors are the reason that the military has such an advantage at night, an advantage which, in turn, has an effect on global and domestic politics.
Environmental Impacts
Tellurium extraction, as a byproduct of copper refinement, shares environmental impacts associated with copper mining and extraction. While a generally safe process, the removal of copper from other impurities in the ore can lead to leaching of various hazardous sediments. In addition, the mining of copper tends to lead to reduced water flow and quality, disruption of soils and erosion of riverbanks, and reduction of air quality.
Resource Limitations v. Demand
About 215-220 tons of tellurium are mined across the globe every year. In 2006, the US produced 40% of the global production, Peru produced 30%, Japan produced 20%, and Canada produced 10% of the world's tellurium supply (since the chart can't be any bigger). The leading countries in production are the United States with 50 tons per year, Japan with 40 tons per year, Canada with 16 tons per year, and Peru with 7 tons per year (year 2009). When pure, tellurium costs \$24 per 100 grams. Because tellurium is about as rare as platinum on earth, the United States Department of Energy expects a supply shortfall by the year 2025, despite the always improving extraction methods. As demand increases to provide the tellurium needed for solar panels and other such things, supply will continue to decrease and thus the price will skyrocket. This will cause waves in the sustainable energy movement as well as military practices and modern medicine.
Z084 Chemistry of Polonium (
Polonium was discovered in 1898 by Marie Curie and named for her native country of Poland. The discovery was made by extraction of the remaining radioactive components of pitchblende following the removal of uranium. There is only about 10-6 g per ton of ore! Current production for research purposes involves the synthesis of the element in the lab rather than its recovery from minerals. This is accomplished by producing Bi-210 from the abundant Bi-209. The new isotope of bismuth is then allowed to decay naturally into Po-210. The sample pictured above is actually a thin film of polonium on stainless steel.
Although radioactive, polonium has a few commercial uses. You can buy your own sample of polonium at a photography store. It is part of the special anti-static brushes for dusting off negatives and prints.
Contributors and Attributions
Stephen R. Marsden
Z116 Chemistry of Livermorium
In May of 2012 the IUPAC approved the name "Livermorium" (symbol Lv) for element 116. The new name honors the Lawrence Livermore National Laboratory (1952). A group of researchers of this laboratory with the heavy element research group of the Flerov Laboratory of Nuclear Reactions took part in the work carried out in Dubna on the synthesis of superheavy elements including element 116. Atoms of Lv were formed by the reaction of Cm-248 with Ca-48. The resulting Lv-292 decays by alpha emission to Fl-288.
Contributors and Attributions
Stephen R. Marsden | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_16%3A_The_Oxygen_Family_(The_Chalcogens)/Z052_Chemistry_of_Tellurium_%2.txt |
The halogens are located on the left of the noble gases on the periodic table. These five toxic, non-metallic elements make up Group 17 of the periodic table and consist of: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). Although astatine is radioactive and only has short-lived isotopes, it behaves similar to iodine and is often included in the halogen group. Because the halogen elements have seven valence electrons, they only require one additional electron to form a full octet. This characteristic makes them more reactive than other non-metal groups.
• Group 17: Physical Properties of the Halogens
It can be seen that there is a regular increase in many of the properties of the halogens proceeding down group 17 from fluorine to iodine. This includes their melting points, boiling points, intensity of their color, the radius of the corresponding halide ion, and the density of the element. On the other hand, there is a regular decrease in the first ionization energy as we go down this group. As a result, there is a regular increase in the ability to form high oxidation states.
• Group 17: Chemical Properties of the Halogens
Covers the halogens in Group 17: fluorine (F), chlorine (Cl), bromine (Br) and iodine (I). Includes trends in atomic and physical properties, the redox properties of the halogens and their ions, the acidity of the hydrogen halides, and the tests for the halide ions.
• Chemistry of Fluorine (Z=9)
Fluorine (F) is the first element in the Halogen group (group 17) in the periodic table. Its atomic number is 9 and its atomic weight is 19, and it's a gas at room temperature. It is the most electronegative element, given that it is the top element in the Halogen Group, and therefore is very reactive. It is a nonmetal, and is one of the few elements that can form diatomic molecules (F2).
• Chemistry of Chlorine (Z=17)
Chlorine is a halogen in group 17 and period 3. It is very reactive and is widely used for many purposes, such as as a disinfectant. Due to its high reactivity, it is commonly found in nature bonded to many different elements.
• Chemistry of Bromine (Z=35)
Bromine is a reddish-brown fuming liquid at room temperature with a very disagreeable chlorine-like smell. In fact its name is derived from the Greek bromos or "stench". It was first isolated in pure form by Balard in 1826. It is the only non-metal that is a liquid at normal room conditions. Bromine on the skin causes painful burns that heal very slowly. It is an element to be treated with the utmost respect in the laboratory.
• Chemistry of Iodine (Z=53)
Elemental iodine is a dark grey solid with a faint metallic luster. When heated at ordinary air pressures it sublimes to a violet gas. The name iodine is taken from the Greek ioeides which means "violet colored". It was discovered in 1811 by Courtois.
• Chemistry of Astatine (Z=85)
Astatine is the last of the known halogens and was synthesized in 1940 by Corson and others at the University of California. It is radioactive and its name, from the Greek astatos, means "unstable". The element can be produced by bombarding targets made of bismuth-209 with high energy alpha particles (helium nuclei). Astatine 211 is the product and has a half-life of 7.2 hours. The most stable isotope of astatine is 210 which has a half-life of 8.1 hours.
Thumbnail: Chlorine gas in an ampoule. (CC-BY-SA; W. Oelen (http://woelen.homescience.net/science/index.html)).
Group 17: The Halogens
Some chemical and physical properties of the halogens are summarized in Table $1$. It can be seen that there is a regular increase in many of the properties of the halogens proceeding down group 17 from fluorine to iodine. This includes their melting points, boiling points, the intensity of their color, the radius of the corresponding halide ion, and the density of the element. On the other hand, there is a regular decrease in the first ionization energy as we go down this group. As a result, there is a regular increase in the ability to form high oxidation states and a decrease in the oxidizing strength of the halogens from fluorine to iodine.
Table $1$: Properties of Group 17 (The Halogens)
Property F Cl Br I
Atomic number, Z 9 17 35 53
Ground state electronic configuration [He]2s2 2p5 [Ne]3s2 3p5 [Ar]3d10 4s2 4p5 [Kr]4d10 5s2 5p5
color pale yellow gas yellow-green gas red-brown liquid blue-black solid
Density of liquids at various temperatures, /kg m-3 1.51 (85 °K) 1.66 (203 °K) 3.19 (273 °K) 3.96 (393 °K)
Melting point, /K 53.53 171.6 265.8 386.85
Boiling point, /K 85.01 239.18 331.93 457.5
Enthalpy of atomization, ΔaH° (298K) / kJ mol-1 79.08 121.8 111.7 106.7
Standard enthalpy of fusion of X2, ΔfusH°(mp) / kJ mol-1 0.51 6.4 10.57 15.52
Standard enthalpy of vaporization of X2, ΔvapH°(bp) / kJ mol-1 6.62 20.41 29.96 41.57
First ionization energy, IE1 / kJ mol-1 1681 1251.1 1139.9 1008.4
ΔEAH1°(298K) / kJ mol-1 -333 -348 -324 -295
ΔhydH°(X-,g) / kJ mol-1 -504 -361 -330 -285
ΔhydS°(X-,g) / JK-1 mol-1 -150 -90 -70 -50
ΔhydG°(X-,g) / kJ mol-1 -459 -334 -309 -270
Standard redox potential, E°(X2 /2X-) /V 2.87 1.36 1.09 0.54
Covalent radius, rcov = ½ X-X bond length /pm 72 100 114.2 133.3
Ionic radius, rion for X- /pm 133 181 196 220
van der Waals radius, rv /pm 135 180 195 215
X-X(g)bond energy /kJ mol-1 159 243 193 151
H-X(g)bond energy /kJ mol-1 562 431 366 299
C-X(g)bond energy /kJ mol-1 484 338 276 238
Pauling electronegativity, χP 3.98 3.16 2.96 2.66
Color
The origin of the color of the halogens stems from the excitation between the highest occupied π* molecular orbital and the lowest unoccupied σ* molecular orbital. The energy gap between the HOMO and LUMO decreases according to F2 > Cl2 > Br2 > I2. The amount of energy required for excitation depends upon the size of the atom. Fluorine is the smallest element in the group and the force of attraction between the nucleus and the outer electrons is very large. As a result, it requires a large excitation energy and absorbs violet light (high energy) and so appears pale yellow. On the other hand, iodine needs significantly less excitation energy and absorbs yellow light of low energy. Thus it appears dark violet. Using similar arguments, it is possible to explain the greenish yellow color of chlorine and the reddish brown color of bromine.
Figure $1$: Molecular orbital diagram for fluorine.
The halogens show a variety of colors when dissolved in different solvents. Solutions of iodine can be bright violet in CCl4, pink or reddish brown in aromatic hydrocarbons, and deep brown in alcohols, for example. This variety can be explained by weak donor-acceptor interaction and complex formation. The presence of charge-transfer bands further supports this since they are thought to be derived from interaction with the HOMO σu* orbital.
The X-ray structure of some of these complexes have been obtained, and often the intense color can be used for characterization and determination such as the bright blue color of iodine in the presence of starch. In the case of the solid formed between dibromine and benzene, the structure is shown below and a new charge transfer band occurs at 292 nm. The Br-Br bond length is essentially unchanged from that of dibromine (228 pm).
Figure $2$: Structure of dibromine and benzene complex
In a study of the reaction of dibromine with substituted phosphines in diethyl ether, all but one showed a tetrahedral arrangement where one bromine was linked to the phosphorus.[3]
$R_3P + Br_2 (Et_2O, N_2/r.t.) \rightarrow R_3PBr_2 \label{1}$
The X-ray study of the triethylphosphine was interpreted as [Et3PBr]Br, where the Br-Br separation was 330 pm. This is considerably longer than the 228 pm found above and was taken to mean that the compound was ionic. In the case of the tri(perfluorophenyl)phosphine, however, the structure showed both bromines linked to give a trigonal bipyramid arrangement with D3 symmetry. Why (C6F5)3PBr2 was the only R3PBr2 compound that adopted trigonal bipyramidal geometry was reasoned to be due to the very low basicity of the parent tertiary phosphine.
Melting and Boiling Points
Intermolecular forces are the attractive forces between molecules without which all substances would be gases. The various types of these interactions span large differences in energy and for the halogens and interhalogens are generally quite small. The dispersion forces involved in these cases are called London forces (after Fritz Wolfgang London, 1900-1954). They are derived from momentary oscillations of electron charge in atoms and hence are present between all particles (atoms, ions and molecules).
The ease with which the electron cloud of an atom can be distorted to become asymmetric is termed the molecule's polarizability. The greater the number of electrons an atom has, the farther the outer electrons will be from the nucleus, and the greater the chance for them to shift positions within the molecule. This means that larger nonpolar molecules tend to have stronger London dispersion forces. This is evident when considering the diatomic elements in group 17, the Halogens. All of these diatomic elements are nonpolar, covalently bonded molecules. Descending the group, fluorine and chlorine are gases, bromine is a liquid, and iodine is a solid. For nonpolar molecules, the farther you go down the group, the stronger the London dispersion forces.
To picture how this occurs, compare the situation 1) where the electrons are evenly distributed and then consider 2) an instantaneous dipole that would arise from an uneven distribution of electrons on one side of the nucleus. When two molecules are close together, the instantaneous dipole of one molecule can induce a dipole in the second molecule. This results in synchronized motion of the electrons and an attraction between them. 3) Multiply this effect over numerous molecules and the overall result is that the attraction keeps these molecules together, and for diiodine is sufficient to make this a solid.
Figure $3$: On average the electron cloud for molecules can be considered to be spherical in shape. When two non-polar molecules approach, attractions or repulsions between the electrons and nuclei can lead to distortions in their electron clouds (i.e., dipoles are induced). When more molecules interact, these induced dipoles lead to intermolecular attraction.
The changes seen in the variation of MP and BP for the dihalogens and binary interhalogens can be attributed to the increase in the London dispersion forces of attraction between the molecules. In general they increase with increasing atomic number.
Figure 4:
Redox Properties
The most characteristic chemical feature of the halogens is their oxidizing strength. Fluorine has the strongest oxidizing ability, so that a simple chemical preparation is almost impossible and it must be prepared by electrolysis. Note that since fluorine reacts explosively with water, oxidizing it to dioxygen, finding reaction conditions for any reaction can be difficult. When fluorine is combined with other elements they generally exhibit high oxidation states. Chlorine is the next strongest oxidizing agent, but it can be prepared by chemical oxidation. Most elements react directly with chlorine, bromine and iodine, with decreasing reactivity going down the group, but often the reaction must be activated by heat or UV light. [2] The energy changes in redox process are:
1. Enthalpy of atomization,
2. ΔEAH1,
3. ΔhydH°(X-,g)
The redox potential, E°, X2/2X-, measures a free-energy change, usually dominated by the ΔH term. The values in the Table below show that there is a decrease in oxidizing strength proceeding down the group (2.87, 1.36, 1.09, 0.54 V). This can be explained by comparing the steps shown above.
• 1) atomization of the dihalide is the energy required to break the molecule into atoms:
$½ X_{2(g)} \rightarrow X_{(g)} \label{2}$
Note that only F2 and Cl2 are gases in their natural states, so the energies associated with atomization of Br2 and I2 require converting the liquid or solid to gas first.
• 2) ΔEAH1 is the energy liberated when the atom is converted into a negative ion and is related to the Electron Affinity
$X_{(g)} + e^- \rightarrow X^-_{(g)} \label{3}$
Addition of an electron to the small F atom is accompanied by larger e-/e- repulsion than is found for the larger Cl, Br or I atoms. This would suggest that the process for F should be less exothermic than for Cl and not fit the trend that shows a general decrease going down the group.
• 3) ΔhydH°(X-,g) is the energy liberated upon the hydration of the ion, the Hydration energy.
$X^-_{(g)} + H_2O \rightarrow X^-_{(aq)} \label{4}$
The overall reaction is then:
$½ X_{2(g)} \rightarrow X^-_{(aq)} \label{5}$
Table $2$
Halogen atomization energy
(kJ mol-1)
ΔEAH1
(kJ mol-1)
hydration enthalpy
(kJ mol-1)
overall
(kJ mol-1)
F +79.08 -333 -504 -758
Cl +121.8 -348 -361 -587
Br +111.7 -324 -330 -542
I +106.7 -295 -285 -473
This shows a very negative energy change for the fluoride compared to the others in the group. This comes about because of two main factors: the high hydration energy and the low atomization energy. For F2 2) is less than for Cl2, but since the energy needed to break the F-F bond is also less and the hydration more, the total energy drop is much greater. In spite of their lower atomization energies, Br2 and I2 are weaker oxidizing agents than Cl2 and this is due to their smaller ΔEAH1 and smaller ΔhydH°.
It can be seen that the ΔEAH1 value for fluorine is in between those for chlorine and bromine and so this value alone does not provide a good explanation for the observed variation.
Each of the halogens is able to oxidize any of the heavier halogens situated below it in the group. They can oxidize hydrogen and nonmetals such as:
$X_2 + H_{2(g)} \rightarrow 2HX_{(g)} \label{6}$
In water, the halogens disproportionate according to:
$X_2 + H_2O_{(l)} \rightarrow HX_{(aq)} + HXO_{(aq)} \label{7}$
where $X=Cl, Br, I$. When base is added then the reaction goes to completion forming hypohalites, or at higher temperatures, halates; for example, heating dichlorine:
$3Cl_{2(g)} + 6OH^-_{(aq)} \rightarrow ClO^-_{3(aq)} + 5Cl^-_{(aq)} + 3H_2O(l) \label{8}$
First Ionization Energies
The trend seen for the complete removal of an electron from the gaseous halogen atoms is that fluorine has the highest IE1 and iodine the lowest. To overcome the attractive force of the nucleus means that energy is required, so the Ionization Energies are all positive. The variation with size can be explained, since as the size increases it take less energy to remove an electron. This inverse relationship is seen for all the groups, not just group 17. As the distance from the nucleus to the outermost electrons increases, the attraction decreases so that those electrons are easier to remove. The high value of IE1 for fluorine is such that it does not exhibit any positive oxidation states, whereas Cl, Br and I can exist in oxidation states as high as 7.
Oxidation states
Fluorine is the most electronegative element in the periodic table and exists in all its compounds in either the -1 or 0 oxidation state. Chlorine, bromine, and iodine, however, can be found in a range of oxidation states including: +1, +3, +5, and +7, as shown below.
Table $3$: Common Oxidation States for the Halogens
Oxidation States Examples
-1 CaF2, HCl, NaBr, AgI
0 F2, Cl2, Br2, I2
1 HClO, ClF
3 HClO2, ClF3
5 HClO3, BrF5, [BrF6]-, IF5
7 HClO4, BrF6+, IF7, [IF8]-
In general, odd-numbered groups (like group 17) form odd-numbered oxidation states, and this can be explained since all stable molecules contain paired electrons (free radicals are obviously much more reactive). When covalent bonds are formed or broken, two electrons are involved, so the oxidation state changes by 2.
When difluorine reacts with diiodine, initially iodine monofluoride is formed.
$I_2 + F_2 \rightarrow 2IF \nonumber$
Adding a second difluorine uses two more iodine valence electrons to form two more bonds:
$2IF + F_2 \rightarrow IF_3 \nonumber$
Contributors and Attributions
• The Department of Chemistry, University of the West Indies)
0Group 17: Physical Properties of the Halogens
This page discusses the trends in the atomic and physical properties of the Group 7 elements (the halogens): fluorine, chlorine, bromine and iodine. Sections below cover the trends in atomic radius, electronegativity, electron affinity, melting and boiling points, and solubility, including a discussion of the bond enthalpies of halogen-halogen and hydrogen-halogen bonds.
Trends in Atomic Radius
The figure above shows the increase in atomic radius down the group.
Explaining the increase in atomic radius
The radius of an atom is determined by:
• the number of layers of electrons around the nucleus
• the pull the outer electrons feel from the nucleus.
Compare the numbers of electrons in each layer of fluorine and chlorine:
F 2,7
Cl 2,8,7
In each case, the outer electrons feel a net +7 charge from the nucleus. The positive charge on the nucleus is partially neutralized by the negative inner electrons.
This is true for all the atoms in Group 7: the outer electrons experience a net charge of +7..
The only factor affecting the size of the atom is therefore the number of layers of inner electrons surrounding the atom. More layers take up more space due to electron repulsion, so atoms increase in size down the group.
Trends in Electronegativity
Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. It is usually measured on the Pauling scale, on which the most electronegative element (fluorine) is assigned an electronegativity of 4.0. The figure below shows electronegativities for each halogen:
Notice that electronegativity decreases down the group. The atoms become less effective at attracting bonding pairs of electrons. This effect is illustrated below using simple dots-and-crosses diagrams for hydrogen fluoride and hydrogen chloride:
The bonding pair of electrons between the hydrogen and the halogen experiences the same net pull of +7 from both the fluorine and the chlorine. However, in the chlorine case, the nucleus is farther away from the bonding electrons, which are therefore not as strongly attracted as in the fluorine case.
The stronger attraction from the closer fluorine nucleus makes fluorine more electronegative than chlorine.
Summarizing the trend down the Group
As the halogen atoms increase in size, any bonding pair gets farther away from the halogen nucleus, and so is less strongly attracted toward it. Hence, down the group, the elements become less electronegative.
Trends in First Electron Affinity
The first electron affinity is the energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1- ions. In other words, it is the energy released in the following process:
$X(g) + e^- \rightarrow X^- (g) \nonumber$
First electron affinities have negative values by convention. For example, the first electron affinity of chlorine is -349 kJ mol-1. The negative sign indicates a release of energy.
The first electron affinities of the Group 7 elements
The electron affinity is a measure of the attraction between the incoming electron and the nucleus. There is a positive correlation between attraction and electron affinity. The trend down the group is illustrated below:
Notice that the trend down the group is inconsistent. The electron affinities generally decrease (meaning less heat is emitted), but the fluorine value deviates from this trend.
In the larger atom, the attraction from the more positive nucleus is offset by the additional screening electrons, so each incoming electron feels the effect of a net +7 charge from the center.
As the atom increases in size, the incoming electron is farther from the nucleus and so feels less attraction. The electron affinity therefore decreases down the group. However fluorine is a very small atom, with the incoming electron relatively close to the nucleus, and yet the electron affinity is smaller than expected.
Another effect must be considered in the case of fluorine. As the new electron comes approaches the atom, it enters a region of space already very negatively charged because of the existing electrons. The resulting repulsion from these electrons offsets some of the attraction from the nucleus.
Because the fluorine atom is very small, its existing electron density is very high. Therefore, the extra repulsion is particularly great and diminishes the attraction from the nucleus enough to lower the electron affinity below that of chlorine.
Trends in Melting Point and Boiling Point
Melting and boiling points increase down the group. As indicated by the graph above, fluorine and chlorine are gases at room temperature, bromine is a liquid and iodine a solid.
Explaining the trends in melting point and boiling point
All the halogens exist as diatomic molecules—F2, Cl2, and so on. van der Waals dispersion forces are the primary intermolecular attractions between one molecule and its neighbors. Larger molecules farther down the group have more electrons which can move around and form the temporary dipoles that create these forces.
The stronger intermolecular attractions down the group require more heat energy for melting or vaporizing, increasing their melting or boiling points.
Solubilities
Solubility in water
Fluorine reacts violently with water to produce aqueous or gaseous hydrogen fluoride and a mixture of oxygen and ozone; its solubility is meaningless. Chlorine, bromine, and iodine all dissolve in water to some extent, but there is again no discernible pattern. The following table shows the solubility of the three elements in water at 25°C:
solubility
(mol dm-3)
chlorine 0.091
bromine 0.21
iodine 0.0013
Chlorine dissolved in water produces a pale green solution. Bromine solution adopts a range of colors from yellow to dark orange-red depending on the concentration. Iodine solution in water is very pale brown. Chlorine reacts with water to some extent, producing a mixture of hydrochloric acid and chloric(I) acid (also known as hypochlorous acid). The reaction is reversible, and at any time only about a third of the chlorine molecules have reacted.
$Cl_2 + H_2O \rightleftharpoons HCl + HClO \nonumber$
Chloric(I) acid is sometimes symbolized as HOCl, indicating the actual bonding pattern. Bromine and iodine form similar compounds, but to a lesser extent. In both cases, about 99.5% of the halogen remains unreacted.
The solubility of iodine in potassium iodide solution
Although iodine is only slightly soluble in water, it dissolves freely in potassium iodide solution, forming a dark red-brown solution. A reversible reaction between iodine molecules and iodide ions gives I3- ions. These are responsible for the color. In the laboratory, iodine is often produced through oxidation of iodide ions. As long as there are any excess iodide ions present, the iodine reacts to form I3-. Once the iodide ions have all reacted, the iodine is precipitated as a dark gray solid.
Solubility in hexane
The halogens are much more soluble in organic solvents such as hexane than they are in water. Both hexane and the halogens are non-polar molecules, so the only intermolecular forces between them are van der Waals dispersion forces. Because of this, the attractions broken (between hexane molecules and between halogen molecules) are similar to the new attractions made when the two substances mix. Organic solutions of iodine are pink-purple in color.
Bond enthalpies (bond energies or bond strengths)
Bond enthalpy is the heat required to break one mole of covalent bonds to produce individual atoms, starting from the original substance in the gas state, and ending with gaseous atoms. For chlorine, Cl2(g), it is the heat energy required for the following reaction, per mole:
$Cl-Cl (g) \rightarrow 2Cl(g) \nonumber$
Although bromine is a liquid, the bond enthalpy is defined in terms of gaseous bromine molecules and atoms, as shown below:
Bond enthalpy in the halogens, X2(g)
Covalent bonding is effective because the bonding pair is attracted to both the nuclei at either side of it. It is that attraction which holds the molecule together. The extent of the attraction depends in part on the distances between the bonding pair and the two nuclei. The figure below illustrates such a covalent bond:
In all halogens, the bonding pair experiences a net +7 charge from either end of the bond, because the charge on the nucleus is offset by the inner electrons. As the atoms get larger down the group, the bonding pair is further from the nuclei and the strength of the bond should, in theory, decrease, as indicated in the figure below. The question is whether experimental data matches this prediction.
As is clear from the figure above, the bond enthalpies of the Cl-Cl, Br-Br and I-I bonds decreases as predicted, but the F-F bond enthalpy deviates.
Because fluorine atoms are so small, a strong bond is expected—in fact, it is remarkably weak. There must be another factor for consideration.
In addition to the bonding pair of electrons between the two atoms, each atom has 3 lone pairs of electrons in the outer shell. If the bond is very short,as in F-F, the lone pairs on the two atoms are close enough to cause significant repulsion, illustrated below:
In the case of fluorine, this repulsion is great enough to counteract much of the attraction between the bonding pair and the two nuclei. This weakens the bond.
Bond enthalpies in the hydrogen halides, HX(g)
If the halogen atom is attached to a hydrogen atom, this does not occur; there are no lone pairs on a hydrogen atom. Bond enthalpies for halogen-hydrogen bonds are given below:
As larger halogens are involved, the bonding pair is more distant from the nucleus. The attraction is lessened, and the bond should be weaker; this is supported by the data, without exception. This fact has significant implications for the thermal stability of the hydrogen halides— they are easily broken into hydrogen and the halogen on heating.
Hydrogen fluoride and hydrogen chloride are thermally very stable under typical laboratory conditions. Hydrogen bromide breaks down to some extent into hydrogen and bromine on heating, and hydrogen iodide is even less stable when heated. Weaker bonds are more easily broken.
Contributors and Attributions
Jim Clark (Chemguide.co.uk) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_17%3A_The_Halogens/0Group_17%3A_Physical_Properties_of_the_Halogens/Ato.txt |
The halogens are located on the left of the noble gases on the periodic table. These five toxic, non-metallic elements make up Group 17 of the periodic table and consist of: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). Although astatine is radioactive and only has short-lived isotopes, it behaves similarly to iodine and is often included in the halogen group. Because the halogen elements have seven valence electrons, they only require one additional electron to form a full octet. This characteristic makes them more reactive than other non-metal groups.
Introduction
Halogens form diatomic molecules (of the form X2, where X denotes a halogen atom) in their elemental states. The bonds in these diatomic molecules are non-polar covalent single bonds. However, halogens readily combine with most elements and are never seen uncombined in nature. As a general rule, fluorine is the most reactive halogen and astatine is the least reactive. All halogens form Group 1 salts with similar properties. In these compounds, halogens are present as halide anions with charge of -1 (e.g., Cl-, Br-, etc.). Replacing the -ine ending with an -ide ending indicates the presence of halide anions; for example, Cl- is named "chloride." In addition, halogens act as oxidizing agents—they exhibit the property to oxidize metals. Therefore, most of the chemical reactions that involve halogens are oxidation-reduction reactions in aqueous solution. When in the -1 oxidation state, with carbon or nitrogen in organic compounds, the halogens often form single bonds. When a halogen atom is substituted for a covalently-bonded hydrogen atom in an organic compound, the prefix halo- can be used in a general sense, or the prefixes fluoro-, chloro-, bromo-, or iodo- can be used for specific halogen substitutions. Halogen elements can cross-link to form diatomic molecules with polar covalent single bonds.
Chlorine (Cl2) was the first halogen to be discovered in 1774, followed by iodine (I2), bromine (Br2), fluorine (F2), and astatine (At, discovered last in 1940). The name "halogen" is derived from the Greek roots hal- ("salt") and -gen ("to form"). Together these words combine to mean "salt former", referencing the fact that halogens form salts when they react with metals. Halite is the mineral name for rock salt, a natural mineral consisting essentially of sodium chloride (NaCl). Lastly, the halogens are also relevant in daily life, whether it be the fluoride that goes into toothpaste, the chlorine that disinfects drinking water, or the iodine that facilitates the production of thyroid hormones in one's body.
Elements
Fluorine - Fluorine has an atomic number of 9 and is denoted by the symbol F. Elemental fluorine was first discovered in 1886 by isolating it from hydrofluoric acid. Fluorine exists as a diatomic molecule in its free state (F2) and is the most abundant halogen found in the Earth's crust. Fluorine is the most electronegative element in the periodic table. It appears as a pale yellow gas at room temperature. Fluorine also has a relatively small atomic radius. Its oxidation state is always -1 except in its elemental, diatomic state (in which its oxidation state is zero). Fluorine is extremely reactive and reacts directly with all elements except helium (He), neon (Ne) and argon (Ar). In H2O solution, hydrofluoric acid (HF) is a weak acid. Although fluorine is highly electronegative, its electronegativity does not determine its acidity; HF is a weak acid due to the fact that the fluoride ion is basic (pH>7). In addition, fluorine produces very powerful oxidants. For example, fluorine can react with the noble gas xenon and form the strong oxidizing agent Xenon Difluoride (XeF2). There are many uses for fluorine, which will be discussed later.
Chlorine - Chlorine has the atomic number 17 and the chemical symbol Cl. Chlorine was discovered in 1774 by extracting it from hydrochloric acid. In its elemental state, it forms the diatomic molecule Cl2. Chlorine exhibits multiple oxidation states, such as -1, +1, +3, +5, and +7. At room temperature it appears as a light green gas. Since the bond that forms between the two chlorine atoms is weak, the Cl2 molecule is very reactive. Chlorine reacts with metals to produce salts called chlorides. Chloride ions are the most abundant ions that dissolve in the ocean. Chlorine also has two isotopes: 35Cl and 37Cl. Sodium chloride is the most prevalent compound of the chlorides.
Bromine - Bromine has an atomic number of 35 with a symbol of Br. It was first discovered in 1826. In its elemental form, it is the diatomic molecule Br2. At room temperature, bromine is a reddish- brown liquid. Its oxidation states vary from -1, +1, 3, 4 and 5. Bromine is more reactive than iodine, but not as reactive as chlorine. Also, bromine has two isotopes: 79Br and 81Br. Bromine consists of bromide salts, which have been found in the sea. The world production of bromide has increased significantly over the years, due to its accessibility and longer existence. Like all of the other halogens, bromine is an oxidizing agent, and is very toxic.
Iodine - Iodine has the atomic number 53 and symbol I. Iodine has oxidation states -1, +1, +5 and +7. Iodine exists as a diatomic molecule, I2, in its elemental state. At room temperature, it appears as a violet solid. Iodine has one stable isotope: 127I. It was first discovered in 1811 through the use of seaweed and sulfuric acid. Currently, iodide ions can be isolated in seawater. Although iodine is not very soluble in water, the solubility may increase if particular iodides are mixed into the solution. Iodine has many important roles in life, including thyroid hormone production. This will be discussed in Part VI of the text.
Astatine - Astatine is a radioactive element with an atomic number of 85 and symbol At. Its possible oxidation states include: -1, +1, +3, +5, and +7. It is the only halogen that is not a diatomic molecule, and it appears as a black, metallic solid at room temperature. Astatine is a very rare element, so there is not that much known about this element. In addition, astatine has a very short radioactive half-life, no longer than a couple of hours. It was discovered in 1940 by synthesis. Also, it is thought that astatine is similar to iodine. However, these two elements are assumed to differ in their metallic character.
Table 1.1: Electron configurations of the halogens.
Halogen Electronic Configuration
Fluorine 1s2 2s2 2p5
Chlorine [Ne]3s2 3p5
Bromine [Ar]3d10 4s2 4p5
Iodine [Kr]4d10 5s2 5p5
Astatine [Xe]4f14 5d10 6s2 6p5
Periodic Trends
The periodic trends observed in the halogen group:
Melting and Boiling Points (increase down the group)
The melting and boiling points increase down the group because of the van der Waals forces. The size of the molecules increases down the group. This increase in size means an increase in the strength of the van der Waals forces.
$F < Cl < Br < I < At \nonumber$
Table 1.2: Melting and Boiling Points of Halogens
Halogen Melting Point (˚C) Boiling Point (˚C)
Fluorine -220 -188
Chlorine -101 -35
Bromine -7.2 58.8
Iodine 114 184
Astatine 302 337
Atomic Radius (increases down the group)
The size of the nucleus increases down a group (F < Cl < Br < I < At) because the numbers of protons and neutrons increase. In addition, more energy levels are added with each period. This results in a larger orbital, and therefore a longer atomic radius.
Table 1.3: Atomic Radii of Halogens
Halogen Covalent Radius (pm) Ionic (X-) radius (pm)
Fluorine 71 133
Chlorine 99 181
Bromine 114 196
Iodine 133 220
Astatine 150
Ionization Energy (decreases down the group)
If the outer valence electrons are not near the nucleus, it does not take as much energy to remove them. Therefore, the energy required to pull off the outermost electron is not as high for the elements at the bottom of the group since there are more energy levels. Also, the high ionization energy makes the element appear non-metallic. Iodine and astatine display metallic properties, so ionization energy decreases down the group (At < I < Br < Cl < F).
Table 1.4 Ionization Energy of Halogens
Halogen First Ionization Energy (kJ/mol)
Fluorine 1681
Chlorine 1251
Bromine 1140
Iodine 1008
Astatine 890±40
Electronegativity (decreases down the group)
The number of valence electrons in an atom increases down the group due to the increase in energy levels at progressively lower levels. The electrons are progressively further from the nucleus; therefore, the nucleus and the electrons are not as attracted to each other. An increase in shielding is observed. Electronegativity therefore decreases down the group (At < I < Br < Cl < F).
Table 1.5: Electronegativity of Halogens
Halogen Electronegativity
Fluorine 4.0
Chlorine 3.0
Bromine 2.8
Iodine 2.5
Astatine 2.2
Electron Affinity (decreases down the group)
Since the atomic size increases down the group, electron affinity generally decreases (At < I < Br < F < Cl). An electron will not be as attracted to the nucleus, resulting in a low electron affinity. However, fluorine has a lower electron affinity than chlorine. This can be explained by the small size of fluorine, compared to chlorine.
Table 1.6: Electron Affinity of Halogens
Halogen Electron Affinity (kJ/mol)
Fluorine -328.0
Chlorine -349.0
Bromine -324.6
Iodine -295.2
Astatine -270.1
Reactivity of Elements (decreases down the group)
The reactivities of the halogens decrease down the group ( At < I < Br < Cl < F). This is due to the fact that the atomic radius increases in size with an increase of electronic energy levels. This lessens the attraction for valence electrons of other atoms, decreasing reactivity. This decrease also occurs because electronegativity decreases down a group; therefore, there is less electron "pulling." In addition, there is a decrease in oxidizing ability down the group.
Hydrogen Halides and Halogen Oxoacids
Hydrogen Halides
A halide is formed when a halogen reacts with another, less electronegative element to form a binary compound. Hydrogen, for example, reacts with halogens to form halides of the form HX:
• Hydrogen Fluoride: HF
• Hydrogen Chloride: HCl
• Hydrogen Bromide: HBr
• Hydrogen Iodide: HI
Hydrogen halides readily dissolve in water to form hydrohalic (hydrofluoric, hydrochloric, hydrobromic, hydroiodic) acids. The properties of these acids are given below:
• The acids are formed by the following reaction: HX (aq) + H2O (l) X- (aq) + H3O+ (aq)
• All hydrogen halides form strong acids, except HF
• The acidity of the hydrohalic acids increases as follows: HF < HCl < HBr < HI
Hydrofluoric acid can etch glass and certain inorganic fluorides over a long period of time.
It may seem counterintuitive to say that HF is the weakest hydrohalic acid because fluorine has the highest electronegativity. However, the H-F bond is very strong; if the H-X bond is strong, the resulting acid is weak. A strong bond is determined by a short bond length and a large bond dissociation energy. Of all the hydrogen halides, HF has the shortest bond length and largest bond dissociation energy.
Halogen Oxoacids
A halogen oxoacid is an acid with hydrogen, oxygen, and halogen atoms. The acidity of an oxoacid can be determined through analysis of the compound's structure. The halogen oxoacids are given below:
• Hypochlorous Acid: HOCl
• Chlorous Acid: HClO2
• Chloric Acid: HClO3
• Perchloric Acid: HClO4
• Hypobromous Acid: HOBr
• Bromic Acid: HBrO3
• Perbromic Acid: HBrO4
• Hypoiodous Acid: HOI
• Iodic Acid: HIO3
• Metaperiodic Acid: HIO4; H5IO6
In each of these acids, the proton is bonded to an oxygen atom; therefore, comparing proton bond lengths is not useful in this case. Instead, electronegativity is the dominant factor in the oxoacid's acidity. Acidic strength increases with more oxygen atoms bound to the central atom.
States of Matter at Room Temperature
Table 1.7: States of Matter and Appearance of Halogens
States of Matter (at Room Temperature) Halogen Appearance
Solid Iodine Violet
Astatine Black/Metallic [Assumed]
Liquid Bromine Reddish-Brown
Gas Fluorine Pale Yellow-Brown
Chlorine Pale Green
Explanation for Appearance
The halogens' colors are results of the absorption of visible light by the molecules, which causes electronic excitation. Fluorine absorbs violet light, and therefore appears light yellow. Iodine, on the other hand, absorbs yellow light and appears violet (yellow and violet are complementary colors, which can be determined using a color wheel). The colors of the halogens grow darker down the group:
• Fluorine pale yellow/brown
• Chlorine pale green
• Bromine red-brown
• www.crscientific.com/brominecell4.jpg
• Iodine violet
• genchem.chem.wisc.edu/lab/PTL...ments/I/I.jpeg
• Astatine* black/metallic
• www4.msu.ac.th/satit/studentP...t/astatine.jpg
In closed containers, liquid bromine and solid iodine are in equilibrium with their vapors, which can often be seen as colored gases. Although the color for astatine is unknown, it is assumed that astatine must be darker than iodine's violet (i.e., black) based on the preceding trend.
Oxidation States of Halogens in Compounds
As a general rule, halogens usually have an oxidation state of -1. However, if the halogen is bonded to oxygen or to another halogen, it can adopt different states: the -2 rule for oxygen takes precedence over this rule; in the case of two different halogens bonded together, the more electronegative atom takes precedence and adopts the -1 oxidation state.
Example 1.1: Iodine Chloride (ICl)
Chlorine has an oxidation state of -1, and iodine will have an oxidation of +1. Chlorine is more electronegative than iodine, therefore giving it the -1 oxidation state.
Example 1.2: Perbromic acid (HBrO4)
Oxygen has a total oxidation state of -8 (-2 charge x 4 atoms= -8 total charge). Hydrogen has a total oxidation state of +1. Adding both of these values together, the total oxidation state of the compound so far is -7. Since the final oxidation state of the compound must be 0, bromine's oxidation state is +7.
A third exception to the rule is this: if a halogen exists in its elemental form (X2), its oxidation state is zero.
Table 1.8: Oxidation States of Halogens
Halogen Oxidation States in Compounds
Fluorine (always) -1*
Chlorine -1, +1, +3, +5, +7
Bromine -1, +1, +3, +4, +5
Iodine -1, +1,+5, +7
Astatine -1, +1, +3, +5, +7
Example 1.3: Fluorine
Why does fluorine always have an oxidation state of -1 in its compounds?
Solution
Electronegativity increases across a period, and decreases down a group. Therefore, fluorine has the highest electronegativity of all of the elements, indicated by its position on the periodic table. Its electron configuration is 1s2 2s2 2p5. If fluorine gains one more electron, the outermost p orbitals are completely filled (resulting in a full octet). Because fluorine has a high electronegativity, it can easily remove the desired electron from a nearby atom. Fluorine is then isoelectronic with a noble gas (with eight valence electrons); all its outermost orbitals are filled. Fluorine is much more stable in this state.
Applications of Halogens
Fluorine: Although fluorine is very reactive, it serves many industrial purposes. For example, it is a key component of the plastic polytetrafluoroethylene (called Teflon-TFE by the DuPont company) and certain other polymers, often referred to as fluoropolymers. Chlorofluorocarbons (CFCs) are organic chemicals that were used as refrigerants and propellants in aerosols before growing concerns about their possible environmental impact led to their discontinued use. Hydrochlorofluorocarbons (HFCs) are now used instead. Fluoride is also added to toothpaste and drinking water to help reduce tooth decay. Fluorine also exists in the clay used in some ceramics. Fluorine is associated with generating nuclear power as well. In addition, it is used to produce fluoroquinolones, which are antibiotics. Below is a list of some of fluorine's important inorganic compounds.
Table 1.9: Important Inorganic Compounds of Fluorine
Compound Uses
Na3AlF6 Manufacture of aluminum
BF3 Catalyst
CaF2 Optical components, manufacture of HF, metallurgical flux
ClF3 Fluorinating agent, reprocessing nuclear fuels
HF Manufacture of F2, AlF3, Na3AlF6, and fluorocarbons
LiF Ceramics manufacture, welding, and soldering
NaF Fluoridating water, dental prophylaxis, insecticide
SF6 Insulating gas for high-voltage electrical equipment
SnF2 Manufacture of toothpaste
UF6 Manufacture of uranium fuel for nuclear reactors
Chlorine: Chlorine has many industrial uses. It is used to disinfect drinking water and swimming pools. Sodium hypochlorite (NaClO) is the main component of bleach. Hydrochloric acid, sometimes called muriatic acid, is a commonly used acid in industry and laboratories. Chlorine is also present in polyvinyl chloride (PVC) and several other polymers. PVC is used in wire insulation, pipes, and electronics. In addition, chlorine is very useful in the pharmaceutical industry. Medicinal products containing chlorine are used to treat infections, allergies, and diabetes. The neutralized form of hydrochloride is a component of many medications. Chlorine is also used to sterilize hospital machinery and limit infection growth. In agriculture, chlorine is a component of many commercial pesticides: DDT (dichlorodiphenyltrichloroethane) was used as an agricultural insecticide, but its use was discontinued.
Bromine: Bromine is used in flame retardants because of its fire-resistant properties. It also found in the pesticide methyl bromide, which facilitates the storage of crops and eliminates the spread of bacteria. However, the excessive use of methyl bromide has been discontinued due to its impact on the ozone layer. Bromine is involved in gasoline production as well. Other uses of bromine include the production of photography film, the content in fire extinguishers, and drugs treating pneumonia and Alzheimer's disease.
Iodine: Iodine is important in the proper functioning of the thyroid gland of the body. If the body does not receive adequate iodine, a goiter (enlarged thyroid gland) will form. Table salt now contains iodine to help promote proper functioning of the thyroid hormones. Iodine is also used as an antiseptic. Solutions used to clean open wounds likely contain iodine, and it is commonly found in disinfectant sprays. In addition, silver iodide is important for photography development.
Astatine: Because astatine is radioactive and rare, there are no proven uses for this halogen element. However, there is speculation that this element could aid iodine in regulating the thyroid hormones. Also, 211At has been used in mice to aid the study of cancer.
VII. Outside Links
• Grube, Karl; Leffler, Amos J. "Synthesis of metal halides (ML)." J. Chem. Educ. 1993, 70, A204.
• This video provides information about some of the physical properties of chlorine, bromine, and iodine: http://www.youtube.com/watch?v=yP0U5rGWqdg
• The following video compares four halogens: fluorine, chlorine, bromine and iodine in terms of chemical reactions and physical properties. http://www.youtube.com/watch?v=u2ogMUDBaf4
• Color wheel referred to in the text: http://www.wou.edu/las/physci/ch462/c-wheel.gif
• Elson, Jesse. "A bonding parameter. III, Water solubilities and melting points of the alkali halogens." J. Chem. Educ.1969, 46, 86.
• Fessenden, Elizabeth. "Structural chemistry of the interhalogen compounds." J. Chem. Educ. 1951, 28, 619.
• Holbrook, Jack B.; Sabry-Grant, Ralph; Smith, Barry C.; Tandel, Thakor V. "Lattice enthalpies of ionic halides, hydrides, oxides, and sulfides: Second-electron affinities of atomic oxygen and sulfur." J. Chem. Educ. 1990, 67, 304.
• Kildahl, Nicholas K. "A procedure for determining formulas for the simple p-block oxoacids." J. Chem. Educ. 1991, 68, 1001.
• Liprandi, Domingo A.; Reinheimer, Orlando R.; Paredes, José F.; L'Argentière, Pablo C. "A Simple, Safe Way To Prepare Halogens and Study Their Visual Properties at a Technical Secondary School." J. Chem. Educ. 1999 76.
• Meek, Terry L. "Acidities of oxoacids: Correlation with charge distribution."J. Chem. Educ. 1992, 69, 270.
Practice Problems
1. Why does fluorine always have an oxidation state of -1 in its compounds?
2. Find the oxidation state of the halogen in each problem:
1. HOCl
2. KIO3
3. F2
3. What are three uses of chlorine?
4. Which element(s) exist(s) as a solid at room temperature?
5. Do the following increase or decrease down the group of halogens?
1. boiling point and melting point
2. electronegativity
3. ionization energy
Answers
1. Electronegativity increases across a period, and decreases down a group. Therefore, fluorine has the highest electronegativity out of all of the elements. Because fluorine has seven valence electrons, it only needs one more electron to acheive a noble gas configuration (eight valence electrons). Therefore, it will be more likely to pull off an electron from a nearby atom.
2. disinfecting water, pesticides, and medicinal products
1. +1 (Hydrogen has an oxidation state of +1, and oxygen has an oxidation state of -2. Therefore, chlorine must have an oxidation state of +1 so that the total charge can be zero)
2. +5 (Potassium's oxidation state is +1. Oxygen has an oxidation state of -2, so for this compound it is -6 (-2 charge x 3 atoms= -6). Since the total oxidation state has to be zero, iodine's oxidation state must be +5).
3. 0 (Elemental forms always have an oxidation state of 0.)
3. iodine and astatine
1. increases
2. decreases
3. decreases | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_17%3A_The_Halogens/0Group_17%3A_Physical_Properties_of_the_Halogens/Gro.txt |
1. All elements are diatomic and molecular and the boiling and melting points increase as a result of the increasing van der Waals interactions between diatomic molecules for the heavier elements.
2. The elements are typical non-metals in their physical and chemical properties. They form anionic compounds based on X- (X = halogen), which is associated with a complete octet.
The ionic compounds MX become progressively less ionic as the relative atomic mass of X increases, because of the decreasing electronegativity of the halogens. Iodine has the greatest tendency to form cationic species, e.g., I2+, I5+, because it has the lowest ionization energy. The cation Br2+ is known in Br2+Sb3F16-, and Br5+ has been reported.
3. The atoms also form strong covalent bonds with other non-metals. The mean bond enthalpies for E-X bonds are particularly large for fluorine and therefore a wide range of molecular fluorides are known; fluorine is particularly effective at bringing out the highest valencies of the non-metals and highest oxidation states of the metals.
4. The oxidizing ability of the halogens decreases markedly down the group: F2 > Cl2 > Br2 > I2, and only iodine is oxidized by nitric acid.
5. The stabilities of the hydrogen halides decrease down the group, but their acid strengths increase.
6. Only H-F forms strong hydrogen bonds and this is reflected in the boiling and melting points of the hydrogen halides.
7. The halogens form many interhalogen compounds with the less electronegative halogen surrounded by the more electronegative halogens. Neutral, anionic, and cationic interhalogen compounds are known. ICl and IBr are widely used in organic synthesis and are commercially available.
The most extensive series of compounds exists for iodine, e.g., IF7, IF5, ICl4-, ICl2-.Fluorine does not form any interhalogen compounds where it occupies the central position within the molecule.
8. Oxygen fluorides are extremely strong and reactive oxidants and have been explored as potential rocket fuels; the oxides become less reactive down the column and more numerous. Iodine forms a particularly wide range of oxides.
9. The perhalates, EO4-, are only known for Cl, Br, and I. They exhibit an alternation in their oxidizing abilities, and the perbromates are particularly strong oxidizing agents.
10. In the highest oxidation state (+7) the relative oxidizing ability is:
Br> I > Cl
and results in the formation of the corresponding +5 oxoanions,
ClO4- + 2e- = ClO3- E° = 1.20 V
BrO4- + 2e- = BrO3- E° = 1.85 V
IO4- + 2e- = IO3- E° = 1.63 V
The hypohalite ions disproportionate according to the equation:
2XO- = 2X- + XO3-
The equilibrium constants are 1027 for ClO-/Cl- :(the reaction is slow at room temperature), 1015 for BrO-/Br-, and 1020for IO-/I-. HOF has been prepared from ice + F2 but is very reactive, decomposing to HF + O2.
Phy
This page discusses the trends in some atomic and physical properties of the Group 17 elements (the halogens): fluorine, chlorine, bromine and iodine. Sections below describe the trends in atomic radius, electronegativity, electron affinity, melting and boiling points, and solubility. There is also a section on the bond enthalpies (and strengths) of halogen-halogen bonds (for example, the Cl-Cl bond) and of hydrogen-halogen bonds (e.g., the H-Cl bond).
Trends in Atomic Radius
You can see that the atomic radius increases as you go down the group.
The radius of an atom is governed by
• Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. It is usually measured on the Pauling scale, on which the most electronegative element (fluorine) is given an electronegativity of 4.0.
As shown in the figure above, electronegativity decreases from fluorine to iodine; the atoms become less effective at attracting bonding pairs of electrons as they grow larger. This can be visualized using dots-and-crosses diagrams for hydrogen fluoride and hydrogen chloride.
The bonding electrons between the hydrogen and the halogen experience the same net charge of +7 from either the fluorine or the chlorine. However, in the chlorine case, the nucleus is further away from the bonding pair. Therefore, electrons are not as strongly attracted to the chlorine nucleus as they are to the fluorine nucleus.
The stronger attraction to the closer fluorine nucleus makes fluorine more electronegative.
Summarizing the trend down the group
As the halogen atoms get larger, any bonding pair is farther and farther away from the halogen nucleus, and so is less strongly attracted towards it. Hence, the elements become less electronegative as you go down the group. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_17%3A_The_Halogens/0Group_17%3A_Physical_Properties_of_the_Halogens/Hal.txt |
Covers the halogens in Group 17: fluorine (F), chlorine (Cl), bromine (Br) and iodine (I). Includes trends in atomic and physical properties, the redox properties of the halogens and their ions, the acidity of the hydrogen halides, and the tests for the halide ions.
1Group 17: General Reactions
This page examines the redox reactions involving halide ions and concentrated sulfuric acid, using these reactions to discuss the trend in reducing ability of the ions from fluoride to iodide. Two types of reactions might occur when concentrated sulfuric acid is added to a solid ionic halide like sodium fluoride, chloride, bromide or iodide. The concentrated sulfuric acid can act as both an acid and an oxidizing agent.
Concentrated sulfuric acid acting as an acid
Concentrated sulfuric acid transfers a proton to the halide ion to produce a gaseous hydrogen halide, which immediately escapes from the system. If the hydrogen halide is exposed to moist air, steam fumes are formed. For example, concentrated sulfuric acid reacts with solid sodium chloride at low temperatures to produce hydrogen chloride and sodium bisulfate, as in the following equation:
$NaCl + H_2SO_4 \rightarrow HCl + NaHSO_4 \nonumber$
All the halide ions behave similarly.
Concentrated sulfuric acid acting as an oxidizing agent
With fluoride or chloride
Concentrated sulfuric acid is not a strong enough oxidizing agent to oxidize fluoride or chloride. In those cases, only the steamy fumes of the hydrogen halide—hydrogen fluoride or hydrogen chloride—are produced. In terms of the halide ions, fluoride and chloride are not strong enough reducing agents to reduce the sulfuric acid. This is not the case for bromides and iodides.
With bromide
Bromide is a strong enough reducing agent to reduce sulfuric acid. Bromide is oxidized to bromine in the process, as in the half-equation below:
$2Br^- \rightarrow Br_2 + 2e^- \nonumber$
Bromide reduces sulfuric acid to sulfur dioxide gas, decreasing the oxidation state of sulfur from +6 to +4. The half-equation for this transition is as follows:
$H_2SO_4 + 2H^+ + 2e^- \rightarrow SO_2 + 2H_2O \nonumber$
These two half-equations can be combined into the overall ionic equation for the reaction:
$H_2SO_4 + 2H^+ + 2Br^- \rightarrow Br_2 +SO_2 +2H_2O \nonumber$
In practice, this reaction is confirmed by the steamy fumes of hydrogen bromide contaminated with the brown color of bromine vapor. The sulfur dioxide is a colorless gas, so its presence cannot be directly observed.
With Iodide
Iodide is a stronger reducing agent than bromide, and it is oxidized to iodine by the sulfuric acid:
$2I^- \rightarrow I_2 +2e^- \nonumber$
The reduction of the sulfuric acid is more complicated than with bromide. Iodide is powerful enough to reduce it in three steps:
• sulfuric acid to sulfur dioxide (sulfur oxidation state = +4)
• sulfur dioxide to elemental sulfur (oxidation state = 0)
• sulfur to hydrogen sulfide (sulfur oxidation state = -2).
The most abundant product is hydrogen sulfide. The half-equation for its formation is as follows:
$H_2SO_4 + 8H^+ + 8e^- \rightarrow H_2S + 4H_2O \nonumber$
Combining these two half-equations gives the following net ionic equation:
$H_2SO_4 + 8H^+ + 8I^- \rightarrow 4I_2 + H_2S + 4H_2O \nonumber$
This is confirmed by a trace of steamy fumes of hydrogen iodide and a large amount of iodine. The reaction is exothermic: purple iodine vapor is formed, with dark gray solid iodine condensing around the top of the reaction vessel. There is also a red color where the iodine comes into contact with solid iodide salts. The red color is due to the I3- ion formed by reaction between I2 molecules and I- ions. Hydrogen sulfide gas can be detected by its "rotten egg" smell, but this gas is intensely poisonous.
Summary of the trend in reducing ability
• Fluoride and chloride cannot reduce concentrated sulfuric acid.
• Bromide reduces sulfuric acid to sulfur dioxide. In the process, bromide ions are oxidized to bromine.
• Iodide reduces sulfuric acid to a mixture of products including hydrogen sulfide. Iodide ions are oxidized to iodine.
• The reducing ability of halide ions increases down the group.
Explaining the trend
An over-simplified explanation
The following explanation is only (partially) accurate if fluoride is neglected. When a halide ion acts as a reducing agent, it transfers electrons to something else. That means that the halide ion itself loses electrons. The larger the halide ion, the farther the outer electrons are from the nucleus, and the more they are shielded by inner electrons. It therefore gets easier for the halide ions to lose electrons down the group because there is less attraction between the outer electrons and the nucleus. This argument seems valid, but it is flawed. The energetics of the change must be examined.
A more detailed explanation
Enthalpy change variation between halogens
The amount of heat evolved or absorbed when a solid halide (like sodium chloride) is converted into an elemental halogen must be considered. Taking sodium chloride as an example, the following energetic quantities are important:
• The energy required to break the attractions between the ions in the sodium chloride (the lattice enthalpy).
• The energy required to remove the electron from the chloride ion. This is the reverse of the electron affinity of the chlorine (the electron affinity can be acquired from a data table and negated).
• The energy recovered when the chlorine atoms convert to diatomic chlorine. Energy is released when these bonds are formed. Chlorine is simple because it is a gas. In bromine and iodine, heat is also released during condensation to a liquid or a solid, respectively. To account for this, it is simpler to think in terms of atomization energy rather than bond energy. The value of interest is the reverse of atomization energy.
Atomization energy is the energy needed to produce 1 mole of isolated gaseous atoms starting from an element in its standard state (gas for chlorine, and liquid for bromine, for example - both of them as X2). The figure below shows how this information fits together:
The enthalpy change shown by the green arrow in the diagram for each of the halogens must be compared. The diagram shows that the overall change involving the halide ions is endothermic (the green arrow is pointing up toward a higher energy).
This is not the total enthalpy change for the whole reaction. Heat is emitted when the changes involving the sulfuric acid occur. That is the same irrespective of the halogen in question. The total enthalpy change is the sum of the enthalpy changes for the halide ion half-reaction and the sulfuric acid half-reaction.
The table below shows the energy changes that vary from halogen to halogen. The process is assumed to start from the solid sodium halide. The values for the lattice enthalpies for other solid halides would be different, but the pattern would be the same.
heat needed to break up NaX lattice
(kJ mol-1)
heat needed to remove electron from halide ion
(kJ mol-1)
heat released in forming halogen molecules
(kJ mol-1)
sum of these
(kJ mol-1)
F +902 +328 -79 +1151
Cl +771 +349 -121 +999
Br +733 +324 -112 +945
I +684 +295 -107 +872
The overall enthalpy change for the halide half-reaction:
The sum of the enthalpy changes, in the final column, is decreasingly endothermic down the group. The total change in enthalpy (including the sulfuric acid) is also less positive.
The amount of heat produced in the half-reaction involving the sulfuric acid must be great enough to make the reactions with the bromide or iodide feasible, but not enough to compensate for the more positive values produced by the fluoride and chloride half-reactions.
Exploring the changes in the various energy terms
In this section, the individual energy terms in the table that are most important in making the halogen half-reaction less endothermic down the group are determined.
Chlorine to iodine
From chlorine to iodine, the lattice enthalpy changes most, decreasing by 87 kJ mol-1. By contrast, the energy required to remove the electron decreases by only 54 kJ mol-1. Both of these terms matter, but the decrease in lattice enthalpy is the more significant. This quantity decreases because the ions are getting larger. That means that they are farther away from each other, and so the attractions between positive and negative ions in the solid lattice are lessened.
The simplified explanation mentioned earlier is misleading because it concentrates on the less-important decrease in the amount of energy needed to remove the electron from the ion.
Fluorine
Fluoride ions are very difficult to oxidize to fluorine. The table above shows that this has nothing to do with the amount of energy required to remove an electron from a fluoride ion. It actually takes less energy to remove an electron from a fluoride ion than from a chloride ion. The generalization that an electron becomes easier to remove as the ion becomes larger does not apply here.
Fluoride ions are so small that the electrons experience strong repulsion from each other. This outweighs the effect of their closeness to the nucleus and makes them easier to remove than the simplified argument predicts.
There are two important reasons why fluoride ions are so difficult to oxidize. The first is the comparatively high lattice enthalpy of the solid fluoride. This is due to the small size of the fluoride ion, which means that the positive and negative ions are very close together and therefore strongly attracted to each other. The other factor is the small amount of heat that is released when the fluorine atoms combine to make fluorine molecules (see the table above). This is due to the low bond enthalpy of the F-F bond. The reason for this low bond enthalpy is discussed on a separate page.
What if the halide ions were in solution rather than in a solid?
This discussion has focused on the energetics of the process starting from solid halide ions because that is the standard procedure when using concentrated sulfuric acid. Halides could also be oxidized in solution with another oxidizing agent.
The trend is exactly the same. Fluoride is difficult to oxidize and it becomes easier down the group toward iodide; in other words, fluoride ions are not good reducing agents, but iodide ions are.
The explanation starts from the hydrated ions in solution rather than solid ions. In a sense, this has already been done on another page. Fluorine is a very powerful oxidizing agent because it very readily forms its negative ion in solution. It is therefore energetically difficult to reverse the process. By contrast, for energetic reasons, iodine is relatively reluctant to form its negative ion in solution. Therefore, it is relatively easy for it to revert back to iodine. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_17%3A_The_Halogens/1Group_17%3A_General_Reactions/Halide_Ions_as_Reduci.txt |
This page examines the trend in oxidizing ability of the Group 17 elements (the halogens): fluorine, chlorine, bromine and iodine. It considers the ability of one halogen to oxidize the ions of another, and how this changes down the group.
Basic facts
Consider a situation in which one halogen (chlorine, for example) is reacted with the ions of another (iodide, perhaps) from a salt solution. In the chlorine and iodide ion case, the reaction is as follows:
$\ce{Cl_2 + 2I^- \rightarrow 2Cl^- + I_2} \nonumber$
• The iodide ions lose electrons to form iodine molecules. In other words, they are oxidized.
• The chlorine molecules gain electrons to form chloride ions— they are reduced.
This is therefore a redox reaction in which chlorine acts as an oxidizing agent.
Fluorine
Fluorine must be excluded from this discussion because its oxidizing abilities are too strong. Fluorine oxidizes water to oxygen, as in the equation below, and so it is impossible to carry out reactions with it in aqueous solution.
$\ce{2F_2 + 2H_2O \rightarrow 4HF + O_2} \nonumber$
Chlorine, Bromine and Iodine
In each case, a halogen higher in the group can oxidize the ions of one lower down. For example, chlorine can oxidize bromide ions to bromine:
$\ce{Cl_2 + 2Br^- \rightarrow 2Cl^- + Br_2} \nonumber$
The bromine forms an orange solution. As shown below, chlorine can also oxidize iodide ions to iodine:
$\ce{Cl_2 +2I^- \rightarrow 2Cl^- + I_2} \nonumber$
The iodine appears either as a red solution if little chlorine is used, or as a dark gray precipitate if the chlorine is in excess.
Bromine can only oxidize iodide ions, and is not a strong enough oxidizing agent to convert chloride ions into chlorine. A red solution of iodine is formed (see the note above) until the bromine is in excess. Then a dark gray precipitate is formed.
$\ce{Br_2 + 2I^- \rightarrow 2Br^- + I_2} \nonumber$
Iodine won't oxidize any of the other halide ions, except possibly the extremely radioactive and rare astatide ions.
To summarize
• Oxidation is the loss of electrons. Each of the elements (for example, chlorine) could potentially take electrons from something else and are subsequently ionized (e.g., Cl-). This means that they are all potential oxidizing agents.
• Fluorine is such a powerful oxidizing agent that solution reactions are unfeasible.
• Chlorine has the ability to take electrons from both bromide ions and iodide ions. Bromine and iodine cannot reclaim those electrons from the chloride ions formed.
• This indicates that chlorine is a more powerful oxidizing agent than either bromine or iodine.
• Similarly, bromine is a more powerful oxidizing agent than iodine. Bromine can remove electrons from iodide ions, producing iodine; iodine cannot reclaim those electrons from the resulting bromide ions.
In short, oxidizing ability decreases down the group.
Explaining the trend
Whenever one of the halogens is involved in oxidizing a species in solution, the halogen end is reduced to a halide ion associated with water molecules. The following figure illustrates this process:
Down the group, the ease with which these hydrated ions are formed decreases; the halogens become less effective as oxidizing agents, taking electrons from something else less readily. The reason that the hydrated ions form less readily down the group is due to several complicated factors. Unfortunately, this explanation is often over-simplified, giving a faulty and misleading explanation. The wrong explanation is dealt with here before a proper explanation is given.
The Incorrect Explanation
The following explanation is normally given for the trend in oxidizing ability of chlorine, bromine and iodine. The ease of ionization depends on how strongly the new electrons are attracted. As the atoms get larger, the new electrons are further from the nucleus and increasingly shielded by the inner electrons (offsetting the effect of the greater nuclear charge). The larger atoms are therefore less effective at attracting new electrons and forming ions. This is equivalent to saying electron affinity decreases down the group. Electron affinity is described in detail on another page.
The problem with this argument is that it does not include fluorine. Fluorine's tendency to form a hydrated ion is much higher than that of chlorine. However, fluorine's electron affinity is less than that of chlorine. This contradicts the above argument. This problem stems from examining a single part of a very complicated process. The argument about atoms accepting electrons applies only to isolated atoms in the gas state picking up electrons to form isolated ions, also in the gas state. The argument must be generalized.
In reality:
• The halogen starts as a diatomic molecule, X2. This may be a gas, liquid or solid at room temperature, depending on the halogen.
• The diatomic molecule must split into individual atoms (atomization)
• Each atom gains an electron (electron affinity; this is the element of the process of interest in the faulty explanation.)
• The isolated ions are surrounded by water molecules; hydrated ions are formed (hydration).
The Correct Explanation
The table below shows the energy involved in each of these changes for atomization energy, electron affinity, and hydration enthalpy (hydration energy):
atomization energy
(kJ mol-1)
electron affinity
(kJ mol-1)
hydration enthalpy
(kJ mol-1)
overall
(kJ mol-1)
F +79 -328 -506 -755
Cl +121 -349 -364 -592
Br +112 -324 -335 -547
I +107 -295 -293 -481
Consider first the fifth column, which shows the overall heat evolved, the sum of the energies in the previous three columns.
The amount of heat evolved decreases quite dramatically from the top to the bottom of the group, with the biggest decrease between fluorine and chlorine. Fluorine generates a large amount of heat when it forms its hydrated ion, chlorine a lesser amount, and so on down the group.
The first electron affinity is defined as the energy released when 1 mole of gaseous atoms each acquire an electron to form 1 mole of gaseous 1- ions, as in the following equation: In symbol terms:
$X(g) + e^- \rightarrow X^-(g) \nonumber$
The fifth column measures the energy released when 1 mole of gaseous ions dissolves in water to produce hydrated ions, as in the following equation, which is not equivalent to that above:
$X^-(g) \rightarrow X^- (aq) \nonumber$
Why is fluorine a stronger oxidizing agent than chlorine?
There are two main factors. First, the atomization energy of fluorine is abnormally low. This reflects the low bond enthalpy of fluorine.
The main reason, however, is the very high hydration enthalpy of the fluoride ion. That is because fluoride is very small. There is a very strong attraction between fluoride ions and water molecules. The stronger the attraction, the more heat is evolved when the hydrated ions are formed.
Why does oxidizing ability decrease from chlorine to bromine to iodine?
The decrease in atomization energy between these three elements is relatively small, and would tend to make the overall change more negative down the group. It is helpful to look at the changes in electron affinity and hydration enthalpy down the group. Using the figures from the previous table:
change in electron affinity
(kJ mol-1)
change in hydration enthalpy
(kJ mol-1)
Cl to Br +25 +29
Br to I +29 +42
Both of these effects contribute, but the more important factor—the one that changes the most—is the change in the hydration enthalpy. Down the group, the ions become less attractive to water molecules as they get larger. Although the ease with which an atom attracts an electron matters, it is not as important as the hydration enthalpy of the negative ion formed.
The faulty explanation is incorrect even if restricted to chlorine, bromine and iodine:
• This is the energy needed to produce 1 mole of isolated gaseous atoms starting from an element in its standard state (gas for chlorine, and liquid for bromine, for example, both of the form X2).
• For a gas like chlorine, this is simply half of the bond enthalpy (because breaking a Cl-Cl bond produces 2 chlorine atoms, not 1). For a liquid like bromine or a solid like iodine, it also includes the energy that is needed to convert them into gases. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_17%3A_The_Halogens/1Group_17%3A_General_Reactions/Halogens_as_Oxidizing.txt |
The halogens react with each other to form interhalogen compounds. The general formula of most interhalogen compounds is XYn, where n = 1, 3, 5 or 7, and X is the less electronegative of the two halogens. The compounds which are formed by the union of two different halogens are called interhalogen compounds. There are never more than two types of halogen atoms in an interhalogen molecule. These are of four general types:
1. AX- type : ClF, BrF, BrCl, ICl, IBr,
2. AX3-type: ClF3, BrF3, (ICl3)2,
3. AX5-type: ClF5, BrF5, IF5,
4. AX7-type: IF7.
The interhalogen compounds of type AX and AX3 are formed between halogens having a very low electronegativity difference (e.g., ClF, ClF3). The interhalogen compounds of type AX5 and AX7 are formed by larger atoms having a low electronegativity with the smaller atoms having a high electronegativity. This is because it is possible to fit a greater number of smaller atoms around a larger one (e.g., BrF5, IF7).
Interhalogens are all prone to hydrolysis and ionize to give rise to polyatomic ions. The interhalogens are generally more reactive than halogens except for F. This is because A-X bonds in interhalogens are weaker than the X-X bonds in dihalogen molecules. Reactions of interhalogens are similar to those of halogens. Hydrolysis of interhalogen compounds gives a halogen acid and an oxy-acid.
Nomenclature
To name an interhalogen compound, the less electronegative element is placed on to the left in formulas and naming is straightforward.
Properties
Some properties of interhalogen compounds are listed below. They are all prepared by direct combination of the elements, although since in some cases more than one product is possible, the conditions may be varied by altering the temperature and relative proportions. For example, under the same conditions, difluorine reacts with dichlorine to give ClF, with dibromine to give BrF3, but with diiodine to give IF5.
Compound ClF BrF BrCl ICl IBr ClF3 BrF3 IF3 I2Cl6 ClF5 BrF5 IF5 IF7
Appearance at 298K Colorless gas Pale brown gas impure Red solid Black solid Colorless gas Yellow liquid Yellow solid Orange solid Colorless gas Colorless liquid Colorless liquid Colorless gas
Stereochemistry linear linear linear linear linear T-shaped T-shaped T-shaped planar square-based pyramid square-based pyramid square-based pyramid pentagonal bipyramid
Melting point /K 117 ~240 dissoc. 300(a) 313 197 282 245 (dec) 337 (sub) 170 212.5 282.5 278 (sub)
Boiling point /K 173 ~293 ~278 ~373 389 285 399 - - 260 314 373 -
ΔfH°(298 K) /kJ mol-1 -50.3 -58.5 14.6 -23.8 -10.5 -163.2 -300.8 ~-500 -89.3 -255 -458.6 -864.8 -962
Dipole moment for gas-phase molecule /D 0.89 1.42 0.52 1.24 0.73 0.6 1.19 - 0 - 1.51 2.18 0
Bond distances in gas-phase molecules except for IF3 and I2Cl6 / pm 163 176 214 232 248.5 160 (eq), 170 (ax) 172 (eq), 181 (ax) 187 (eq), 198 (ax) 238 (terminal) 268 (bridge) 172 (basal), 162 (apical) 178 (basal), 168 (apical) 187 (basal), 185 (apical) 186 (eq), 179 (ax)
Structures
The structures found for the various interhalogens conform to what would be expected based on the VSEPR model. For XY3 the shape can be described as T-shaped with 2 lone pairs sitting in two of the equatorial positions of a trigonal bipyramid. For XY5 the shape is a square pyramid with the unpaired electrons sitting in an axial position of an octahedron, and XY7 is a pentagonal bipyramid.
XY diatomic interhalogens
The interhalogens with formula XY have physical properties intermediate between those of the two parent halogens. The covalent bond between the two atoms has some ionic character, with the larger element, X, becoming oxidized and having a partial positive charge. Most combinations of F, Cl, Br and I are known, but not all are stable.
• Chlorine monofluoride (ClF), the lightest interhalogen, is a colorless gas with a boiling point of 173 °K.
• Bromine monofluoride (BrF) has not been obtained in the pure form - it dissociates into the trifluoride and free bromine. Similarly, iodine monofluoride is unstable - iodine reacts with fluorine to form a pentafluoride.
• Iodine monofluoride (IF) is unstable and disproportionates rapidly and irreversibly at room temperature: $\ce{5IF → 2I2 + IF5} \nonumber$. However, its molecular properties have been determined by spectroscopy: the iodine-fluorine distance is 190.9 pm and the I-F bond dissociation energy is around 277 kJ mol-1. ΔHf° = -95.4 kJ mol-1 and ΔGf° = -117.6 kJ mol-1, both at 298 K.
$\ce{IF}$ can be generated by the following reactions:
I2 + F2 → 2IF at -45 °C in CCl3F;
I2 + IF3 → 3IF at -78 °C in CCl3F;
I2 + AgF → IF + AgI at 0 °C.
• Bromine monochloride (BrCl) is an unstable red-brown gas with a boiling point of 5 °C.
• Iodine monochloride (ICl) consists of red transparent crystals which melt at 27.2 °C to form a choking brownish liquid (similar in appearance and weight to bromine). It reacts with HCl to form the strong acid HICl2. The crystal structure of iodine monochloride consists of puckered zig-zag chains, with strong interactions between the chains.
• Iodine monobromide (IBr) is made by direct combination of the elements to form a dark red crystalline solid. It melts at 42 °C and boils at 116 °C to form a partially dissociated vapor.
XY3 interhalogens
• Chlorine trifluoride (ClF3) is a colorless gas that condenses to a green liquid and freezes to a white solid. It is made by reacting chlorine with an excess of fluorine at 250° C in a nickel tube. It reacts more violently than fluorine, often explosively. The molecule is planar and T-shaped.
• Bromine trifluoride (BrF3) is a yellow-green liquid that conducts electricity - it ionizes to form [BrF2+] + [BrF4-]. It reacts with many metals and metal oxides to form similar ionized entities; with some others it forms the metal fluoride plus free bromine and oxygen. It is used in organic chemistry as a fluorinating agent. It has the same molecular shape as chlorine trifluoride.
• Iodine trifluoride (IF3) is a yellow solid which decomposes above -28 °C. It can be synthesized from the elements, but care must be taken to avoid the formation of IF5. F2 attacks I2 to yield IF3 at -45 °C in CCl3F. Alternatively, at low temperatures, the fluorination reaction I2 + 3XeF2 → 2IF3 + 3Xe can be used. Not much is known about iodine trifluoride as it is so unstable.
• Iodine trichloride (ICl3) forms lemon yellow crystals which can be melted under pressure to a brown liquid. It can be made from the elements at low temperature, or from iodine pentoxide and hydrogen chloride. It reacts with many metal chlorides to form tetrachloriodides, and hydrolyses in water. The molecule is a planar dimer, with each iodine atom surrounded by four chlorine atoms. In the melt it is conductive, which may indicate dissociation: $\ce{I_2Cl_6 → ICl_2^{+} + ICl_4^{-}} \nonumber$
Chlorine trifluoride, ClF3, was first reported in 1931; it is primarily used for the manufacture of uranium hexafluoride, UF6, as part of nuclear fuel processing and reprocessing by the reaction:
$\ce{U + 3 ClF_3 → UF_6 + 3 ClF} \nonumber$
U isotope separation is difficult because the two isotopes have very nearly identical chemical properties, and can only be separated gradually using small mass differences. (235U is only 1.26% lighter than 238U.) A cascade of identical stages produces successively higher concentrations of 235U. Each stage passes a slightly more concentrated product to the next stage and returns a slightly less concentrated residue to the previous stage.
There are currently two generic commercial methods employed internationally for enrichment: gaseous diffusion (referred to as first generation) and gas centrifuge (second generation), which consumes only 6% as much energy as gaseous diffusion. These both make use of the volatility of UF6.
ClF3 has been investigated as a high-performance storable oxidizer in rocket propellant systems. Handling concerns, however, prevented its use.
$ClF_3$ is Hypergolic
Hypergolic means explodes on contact with no need for any activator. One observer made the following comment about $ClF_3$:
"It is, of course, extremely toxic, but that's the least of the problem. It is hypergolic* with every known fuel, and so rapidly hypergolic that no ignition delay has ever been measured. It is also hypergolic with such things as cloth, wood, and test engineers, not to mention asbestos, sand, and water with which it reacts explosively. It can be kept in some of the ordinary structural metals-steel, copper, aluminium, etc.-because of the formation of a thin film of insoluble metal fluoride which protects the bulk of the metal, just as the invisible coat of oxide on aluminium keeps it from burning up in the atmosphere. If, however, this coat is melted or scrubbed off, and has no chance to reform, the operator is confronted with the problem of coping with a metal-fluorine fire. For dealing with this situation, I have always recommended a good pair of running shoes."
It is believed that prior to and during World War II, ClF3 code named N-stoff ("substance N") was being stockpiled in Germany for use as a potential incendiary weapon and poison gas. The plant was captured by the Russians in 1944, but there is no evidence that the gas was actually ever used during the war.
XY5 interhalogens
• Chlorine pentafluoride (ClF5) is a colorless gas, made by reacting chlorine trifluoride with fluorine at high temperatures and high pressures. It reacts violently with water and most metals and nonmetals.
• Bromine pentafluoride (BrF5) is a colorless fuming liquid, made by reacting bromine trifluoride with fluorine at 200° C. It is physically stable, but reacts violently with water and most metals and nonmetals.
• Iodine pentafluoride (IF5) is a colorless liquid, made by reacting iodine pentoxide with fluorine, or iodine with silver fluoride. It is highly reactive, even slowly with glass. It reacts with elements, oxides and carbon halides. The molecule has the form of a tetragonal pyramid.
• Primary amines react with iodine pentafluoride to form nitriles after hydrolysis with water. $R-CH_2-NH_2 → R-CN \nonumber$
XY7 interhalogens
• Iodine heptafluoride (IF7) is a colorless gas. It is made by reacting the pentafluoride with fluorine. IF7 is chemically inert, having no lone pair of electrons in the valency shell; in this it resembles sulfur hexafluoride. The molecule is a pentagonal bipyramid. This compound is the only interhalogen compound possible where the larger atom is carrying seven of the smaller atoms.
• All attempts to form bromine heptafluoride have met with failure; instead, bromine pentafluoride and fluorine gas are produced.
Diatomic Interhalogens (AX)
The interhalogens of form XY have physical properties intermediate between those of the two parent halogens. The covalent bond between the two atoms has some ionic character, the less electronegative element, X, being oxidized and having a partial positive charge. Most combinations of F, Cl, Br and I are known, but not all are stable.
• Chlorine monofluoride (ClF): The lightest interhalogen compound, ClF is a colorless gas with a normal boiling point of -100 °C.
• Bromine monofluoride (BrF): BrF has not been obtained in a pure form; it dissociates into the trifluoride and free bromine.
• Iodine monofluoride (IF): IF is unstable and decomposes at 0 C, disproportionating into elemental iodine and iodine pentafluoride.
• Bromine monochloride (BrCl): A red-brown gas with a boiling point of 5 °C.
• Iodine monochloride (ICl): Red transparent crystals which melt at 27.2 °C to form a choking brownish liquid (similar in appearance and weight to bromine). It reacts with HCl to form the strong acid HICl2. The crystal structure of iodine monochloride consists of puckered zig-zag chains, with strong interactions between the chains.
• Iodine monobromide (IBr): Made by direct combination of the elements to form a dark red crystalline solid. It melts at 42 °C and boils at 116 °C to form a partially dissociated vapor.
Tetra-atomic Interhalogens (AX3)
• Chlorine trifluoride (ClF3) is a colorless gas which condenses to a green liquid and freezes to a white solid. It is made by reacting chlorine with an excess of fluorine at 250 °C in a nickel tube. It reacts more violently than fluorine, often explosively. The molecule is planar and T-shaped. It is used in the manufacture of uranium hexafluoride.
• Bromine trifluoride (BrF3) is a yellow-green liquid which conducts electricity and ionizes to form [BrF2+] + [BrF4-]. It reacts with many metals and metal oxides to form similar ionized entities; with some others it forms the metal fluoride plus free bromine and oxygen. It is used in organic chemistry as a fluorinating agent. It has the same molecular shape as chlorine trifluoride.
• Iodine trifluoride (IF3) is a yellow solid which decomposes above -28 °C. It can be synthesized from the elements, but care must be taken to avoid the formation of IF5. F2 attacks I2 to yield IF3 at -45 °C in CCl3F. Alternatively, at low temperatures, the fluorination reaction I2+ 3XeF2 --> 2IF3 + 3Xe can be used. Not much is known about iodine trifluoride as it is so unstable.
• Iodine trichloride (ICl3) forms lemon yellow crystals which can be melted under pressure to a brown liquid. It can be made from the elements at low temperature, or from iodine pentoxide and hydrogen chloride. It reacts with many metal chlorides to form tetrachloriodides, and hydrolyses in water. The molecule is a planar dimer, with each iodine atom surrounded by four chlorine atoms.
Hexa-atomic Interhalogens (AX5)
• Chlorine pentafluoride (ClF5) is a colorless gas, made by reacting chlorine trifluoride with fluorine at high temperatures and high pressures. It reacts violently with water and most metals and nonmetals.
• Bromine pentafluoride (BrF5) is a colorless fuming liquid, made by reacting bromine trifluoride with fluorine at 200Å C. It is physically stable, but reacts violently with water and most metals and nonmetals.
• Iodine pentafluoride (IF5) is a colorless liquid, made by reacting iodine pentoxide with fluorine, or iodine with silver fluoride. It is highly reactive, even slowly with glass. It reacts with elements, oxides and carbon halides.The molecule has the form of a tetragonal pyramid.
Octa-atomic interhalogens (AX7)
• Iodine heptafluoride (IF7) is a colourless gas. It is made by reacting the pentafluoride with fluorine. IF7 is chemically inert, having no lone pair of electrons in the valency shell; in this it resembles sulfur hexafluoride. The molecule is a pentagonal bipyramid. This compound is the only interhalogen compound possible where the larger atom is carrying seven of the smaller atoms
• All attempts to form bromine heptafluoride (BrF7) have failed and instead produce bromine pentafluoride (BrF5) gas.
Summary
All interhalogens are volatile at room temperature. All are polar due to differences in their electronegativity. These are usually covalent liquids or gases due to small electronegativity differences among them. Some compounds partially ionize in solution. For example: $\ce{2 ICl \rightarrow I^{+} + ICl_2^{-}} \nonumber$ Interhalogen compounds are more reactive than normal halogens, except for fluorine. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_17%3A_The_Halogens/1Group_17%3A_General_Reactions/Interhalogens.txt |
This page describes reactions of the halogens that do not fall under the other categories in other pages in this section. All the reactions described here are redox reactions.
Reactions with hydrogen
The following examples illustrate the decrease in reactivity of the halogens down Group 7.
Fluorine combines explosively with hydrogen even under cold, dark conditions, evolving hydrogen fluoride gas.
A mixture of chlorine and hydrogen explodes if exposed to sunlight or a flame, producing hydrogen chloride gas. This reaction can be controlled by lighting a jet of hydrogen and then lowering it into a gas jar of chlorine. The hydrogen burns at a slower, constant rate, and hydrogen chloride gas is formed as before.
Bromine vapor and hydrogen combine with a mild explosion when ignited. Hydrogen bromide gas is formed.
Iodine and hydrogen combine only partially even on constant heating. An equilibrium exists between the hydrogen and the iodine and hydrogen iodide gas.
Each of these reactions has an equation of the form:
$H_2 + X_2 \rightarrow 2HX \nonumber$
A minor exception is made for iodine: the single arrow is replaced with a reversible sign.
Reactions with phosphorus
Care must be taken when analyzing the rates of these reactions; analogous reactions must be compared. For example, it is nonsensical to compare the rate at which phosphorus reacts with gaseous chlorine with the rate at which it reacts with liquid bromine. There is more contact between phosphorus and liquid bromine than between phosphorus and gaseous chlorine.
The formation of trihalides, PX3
All halogens react with phosphorus to form, in the first instance, phosphorus(III) halides of the form PX3.
There are two common forms of phosphorus: white phosphorus (sometimes called yellow phosphorus) and red phosphorus. White phosphorus is more reactive than red phosphorus. This video on YouTube shows the reaction between red phosphorus and bromine. This is a violent reaction under cold conditions, and white phosphorus behaves even more dramatically.
When writing the equations for these reactions, it is important to remember that white phosphorus is molecular, consisting of P4 molecules, whereas red phosphorus is polymeric, indicated by the symbol P. The reaction for white phosphorus and bromine is as follows:
$P_4 + 6Br_2 \rightarrow 4PBr_3 \nonumber$
The red phosphorus equation is shown below:
$2P + 3Br_2 \rightarrow 2PBr_3 \nonumber$
The formation of pentahalides, PX5
In excess chlorine or bromine, phosphorus reacts to form phosphorus(V) chloride or bromide. Most simply, using white phosphorus:
$P_4 + 10Cl_2 \rightarrow 4PCl_5 \nonumber$
The reaction between phosphorus(III) chloride and phosphorus(V) chloride is reversible:
$PCl_3 + Cl_2 \rightleftharpoons PCl_5 \nonumber$
An excess of chlorine pushes this equilibrium to the right. Phosphorus does not form a pentaiodide, in contrast; this is likely because five large iodine atoms cannot physically fit around the central phosphorus atom.
Reactions with sodium
All halogens react with sodium to produce sodium halides. A common reaction between hot sodium and chlorine gas produces a bright orange flame and white sodium chloride.
$2Na + Cl_2 \rightarrow 2NaCl \nonumber$
Hot sodium will also burn in bromine or iodine vapor to produce sodium bromide or sodium iodide. Each of these reactions produces an orange flame and a white solid.
Reactions with iron
With the exception of iodine, iron burns in halogen vapor, forming iron(III) halides. Iodine is less reactive, and produces iron(II) iodide.
Fluorine
Cold iron wool burns in cold fluorine to give iron(III) fluoride. Anhydrous iron(III) fluoride is described as either white or pale green. A standard inorganic chemistry textbook by Cotton and Wilkinson describes it as white. The reaction is given below:
$2Fe + 3F_2 \rightarrow 2FeF_3 \nonumber$
This is a rapid reaction, in which the iron burns and is oxidized to an iron(III) compound—in other words, from an oxidation state of zero in the elemental metal to an oxidation state of +3 in the iron(III) compound.
Chlorine
Chlorine gas in contact with hot iron forms iron(III) chloride. Anhydrous iron(III) chloride forms black crystals; any trace of water present in the apparatus or in the chlorine reacts with the crystals, turning them reddish-brown. The equation for this reaction is given below:
$2Fe + 3Cl_2 \rightarrow 2FeCl_3 \nonumber$
The iron is again oxidized from a state of zero to +3.
Bromine
Bromine vapor passed over hot iron triggers a similar, slightly less vigorous reaction, shown below; iron(III) bromide is produced. Anhydrous iron(III) bromide usually appears as a reddish-brown solid.
$2Fe + 3Br_2 \rightarrow 2FeBr_3 \nonumber$
In this reaction the iron is again oxidized to a +3 state.
Iodine
The reaction between hot iron and iodine vapor produces gray iron(II) iodide, and is much less vigorous. This reaction, the equation for which is given below, is difficult to carry out because the product is always contaminated with iodine.
$Fe + 2I_2 \rightarrow FeI_2 \nonumber$
Iodine is only capable of oxidizing iron to the +2 oxidation state.
Reactions with solutions containing iron(II) ions
Only the reactions of chlorine, bromine, and iodine can be considered. Aqueous fluorine is very reactive with water. Chlorine and bromine are strong enough oxidizing agents to oxidize iron(II) ions to iron(III) ions. In the process, chlorine is reduced to chloride ions, bromine to bromide ions.
This process is easiest to visualize with ionic equations:
For the bromine equation, Br is substituted for Cl.
The pale green solution containing the iron(II) ions turns into a yellow or orange solution containing iron(III) ions. Iodine is not a strong enough oxidizing agent to oxidize iron(II) ions, so there is no reaction. In fact, the reverse reaction proceeds. Iron(III) ions are strong enough oxidizing agents to oxidize iodide ions to iodine as shown:
$2Fe^{3+} + 2I^- \rightarrow 2Fe^{2+} + I_2 \nonumber$
Reactions with sodium hydroxide solution
Once again, only chlorine, bromine, and iodine are considered.
The reaction of chlorine with cold sodium hydroxide solution
Chlorine and cold, dilute sodium hydroxide react as follows:
$2NaOH + Cl_2 \rightarrow NaCl + NaClO + H_2O \nonumber$
NaClO (sometimes written as NaOCl) symbolizes sodium chlorate(I). The traditional name for this compound is sodium hypochlorite; the solution on the product side of the equation is commonly sold as bleach.
Consider this reaction in terms of oxidation states. Chlorine displays an obvious state change from its elemental form to ionic compounds. The oxidation numbers for each element are shown below:
Chlorine is the only element that changes oxidation state—it is both oxidized and reduced. One atom is reduced because its oxidation state has decreased; the other is oxidized. This is a good example of a disproportionation reaction, a reaction in which a single substance is both oxidized and reduced.
The reaction of chlorine with hot sodium hydroxide solution
Chlorine reacts with hot, concentrated sodium hydroxide as follows:
$6NaOH + 3Cl_2 \rightarrow 5NaCl + NaClO_3 + 3H_2O \nonumber$
The product formed is sodium chlorate(V) - NaClO3. As before, the oxidation states of each element are calculated. Once again, the only change is in chlorine, from 0 in the chlorine molecules on the reactant side to -1 (in the NaCl) and +5 (in the NaClO3). This is another example of a disproportionation reaction.
Balancing equations for these reactions
The first equation is simple to balance. The second one is more difficult; oxidation states are used to derive it.
The two main products of the reaction are NaCl and NaCIO3, so the reaction can be tentatively written as follows:
$NaOH + Cl_2 \rightarrow NaCl + NaClO_3 + ? \nonumber$
In its conversion to NaCl, the oxidation state of the chlorine decreases from 0 to -1. When converted to NaClO3, it increases from 0 to +5. The positive and negative oxidation state changes must cancel out, so for every NaClO3 formed, there must be 5 NaCl:
$NaOH + Cl_2 \rightarrow 5NaCl + NaClO_3 + ? \nonumber$
Now it is a simple task to balance the sodium and the chlorine atoms, after which there are enough hydrogen and oxygen atoms to make 3H2O.
The reactions involving bromine and iodine
These are essentially similar to that of chlorine; the difference lies in the reaction temperatures. The tendency to form the ion with the halogen in the +5 oxidation state increases rapidly down the group.
Bromine and sodium hydroxide solution
For bromine, the formation of the sodium bromate(V) happens at around room temperature. Sodium bromate(I) must be formed at about 0°C.
Iodine and sodium hydroxide solution
In this case, sodium iodate(V) is formed at any temperature. Cotton and Wilkinson (Advanced Inorganic Chemistry 3rd edition page 477) say that the iodate(I) ion is unknown in solution.
Contributors and Attributions
Jim Clark (Chemguide.co.uk) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_17%3A_The_Halogens/1Group_17%3A_General_Reactions/More_Reactions_of_Hal.txt |
Consider a reaction between one halogen—chlorine, for example—and the ions of another—iodide, in this case. The iodide ions are dissolved from a salt such as sodium iodide or potassium iodide. The sodium or potassium ions are spectator ions and therefore irrelevant to the reaction, which proceeds as follows:
$Cl_2 + 2I^- \rightarrow 2Cl^- + I_2 \nonumber$
• The iodide ions lose electrons to form iodine molecules; they are oxidized.
• The chlorine molecules gain electrons to form chloride ions. They are reduced.
This is a redox reaction in which chlorine is acting as an oxidizing agent. The driving force force this reaction is straightforward to identify from the table of Standard Reduction Potentials (Table P2).
Contributors and Attributions
• Jim Clark (ChemGuide)
Testing for Halide Io
This page discusses the tests for halide ions (fluoride, chloride, bromide and iodide) using silver nitrate and ammonia.
Using silver nitrate solution
This test is carried out in a solution of halide ions. The solution is acidified by adding dilute nitric acid. The nitric acid reacts with, and removes, other ions that might also form precipitates with silver nitrate. Silver nitrate solution is then added, and the halide can be identified from the following products:
ion present observation
F- no precipitate
Cl- white precipitate
Br- very pale cream precipitate
I- very pale yellow precipitate
The chloride, bromide and iodide precipitates are shown in the photograph:
The chloride precipitate is easily identified, but the other two are quite similar to each other. They can only be differentiated in a side-by-side comparison. All the precipitates change color if they are exposed to light, taking on gray or purple tints. The absence of a precipitate with fluoride ions is unhelpful unless it is known that a halogen is present; otherwise, it indicates that there is no chloride, bromide, or iodide.
The chemistry of the test
The precipitates are insoluble silver halides: silver chloride, silver bromide or silver iodide. The formation of these is illustrated in the following equations:
$Ag^+_{aq} + Cl^-_{(aq)} \rightarrow AgCl_{(s)} \nonumber$
$Ag^+_{aq} + Br^-_{(aq)} \rightarrow AgBr_{(s)} \nonumber$
$Ag^+_{aq} + I^-_{(aq)} \rightarrow AgI_{(s)} \nonumber$
Silver fluoride is soluble, so no precipitate is formed.
$Ag^+_{aq} + F^-_{(aq)} \rightarrow Ag^+_{aq} + F^-_{(aq)} \nonumber$
Confirming the precipitate using ammonia solution
Ammonia solution is added to the precipitates.
original precipitate Observation
AgCl precipitate dissolves to give a colorless solution
AgBr precipitate is almost unchanged using dilute ammonia solution, but dissolves in concentrated ammonia solution to give a colorless solution
AgI precipitate is insoluble in ammonia solution of any concentration
There are no absolutely insoluble ionic compounds. A precipitate forms if the concentrations of the ions in solution in water exceed a certain value, unique to every compound. This value is known as the solubility product. For the silver halides, the solubility product is given by the expression:
$K_{sp} = [Ag^+][X^-] \nonumber$
The square brackets indicate molar concentrations, with units of mol L-1.
• If the product of the concentrations of ions is less than the solubility product, no precipitate is formed.
• If the product of the concentrations exceeds this value, a precipitate is formed.
Essentially, the product of the ionic concentrations is never greater than the solubility product value. Enough solid is always precipitated to lower the ionic product to the solubility product. The table below lists solubility products from silver chloride to silver iodide (a solubility product for silver fluoride cannot be reported because it is too soluble).
Ksp (mol2dm-6)
AgCl 1.8 x 10-10
AgBr 7.7 x 10-13
AgI 8.3 x 10-17
The compounds are all quite insoluble, but become even less so down the group.
The purpose of ammonia
The ammonia combines with silver ions to produce a complex ion called the diamminesilver(I) ion, [Ag(NH3)2]+. This is a reversible reaction, but the complex is very stable, and the position of equilibrium lies well to the right. The equation for this reaction is given below:
A solution in contact with one of the silver halide precipitates contains a very small concentration of dissolved silver ions. The effect of adding the ammonia is to lower this concentration still further. If the adjusted silver ion concentration multiplied by the halide ion concentration is less than the solubility product, some precipitate dissolves to restore equilibrium.
This occurs with silver chloride, and with silver bromide if the ammonia is concentrated. The more concentrated ammonia pushes the equilibrium even further to the right, lowering the silver ion concentration even more.
The silver iodide is so insoluble that ammonia cannot lower the silver ion concentration enough for the precipitate to dissolve.
An alternative test using concentrated sulfuric acid
Adding concentrated sulfuric acid to a solid sample of one of the halides gives the following results:
ion present observation
F- steamy acidic fumes (of HF)
Cl- steamy acidic fumes (of HCl)
Br- steamy acidic fumes (of HBr) contaminated with brown bromine vapor
I- some HI fumes with large amounts of purple iodine vapor and a red compound in the reaction vessel
The only possible confusion is between a fluoride and a chloride—they behave identically under these conditions. They can be distinguished by dissolving the original solid in water and then testing with silver nitrate solution. The chloride gives a white precipitate; the fluoride produces none. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_17%3A_The_Halogens/1Group_17%3A_General_Reactions/Oxidizing_Ability_of_.txt |
This page discusses the acidity of the hydrogen halides: hydrogen fluoride, hydrogen chloride, hydrogen bromide and hydrogen iodide. It begins by describing their physical properties and synthesis and then explains what happens when they react with water to make acids such as hydrofluoric acid and hydrochloric acid.
Physical properties
The hydrogen halides are colorless gases at room temperature, producing steamy fumes in moist air. The boiling points of these compounds are shown in the figure below:
Hydrogen fluoride has an abnormally high boiling point for a molecule of its size (293 K or 20°C), and can condense under cool conditions. This is due to the fact that hydrogen fluoride can form hydrogen bonds. Because fluorine is the most electronegative of all the elements, the fluorine-hydrogen bond is highly polarized. The hydrogen atom carries a high partial positive charge (δ+); the fluorine is fairly negatively charged (δ-).
In addition, each fluorine atom has 3 lone pairs of electrons. Fluorine's outer electrons are at the n=2 level, and the lone pairs represent small, highly charged regions of space. Hydrogen bonds form between the δ+ hydrogen on one HF molecule and a lone pair on the fluorine of another one.The figure below illustrates this association:
The other hydrogen halides do not form hydrogen bonds because the larger halogens are not as electronegative as fluorine; therefore, the bonds are less polar. In addition, their lone pairs are at higher energy levels, so the halogen does not carry such an intensely concentrated negative charge; therefore, other hydrogen atoms are not attracted as strongly.
Making hydrogen halides
There are several ways of synthesizing hydrogen halides; the method considered here is the reaction between an ionic halide, like sodium chloride, and an acid like concentrated phosphoric(V) acid, H3PO4, or concentrated sulfuric acid.
Making hydrogen chloride
When concentrated sulfuric acid is added to sodium chloride under cold conditions, the acid donates a proton to a chloride ion, forming hydrogen chloride. In the gas phase, it immediately escapes from the system.
$Cl^- + H_2SO_4 \rightarrow HCl + HSO_4^- \nonumber$
The full equation for the reaction is:
$NaCl + H_2SO_4 \rightarrow HCl + NaHSO_4 \nonumber$
Sodium bisulfate is also formed in the reaction. Concentrated phosphoric(V) acid reacts similarly, according to the following equation:
$Cl^- + H_3PO_4 \rightarrow HCl + H_2PO_4^- \nonumber$
The full ionic equation showing the formation of the salt, sodium biphosphate(V), is given below:
$NaCl + H_3PO_4 \rightarrow HCl + NaH_2PO_4 \nonumber$
Making other hydrogen halides
All hydrogen halides can be formed by the same method, using concentrated phosphoric(V) acid.
Concentrated sulfuric acid, however, behaves differently. Hydrogen fluoride can be made with sulfuric acid, but hydrogen bromide and hydrogen iodide cannot.
The problem is that concentrated sulfuric acid is an oxidizing agent, and as well as producing hydrogen bromide or hydrogen iodide, some of the halide ions are oxidized to bromine or iodine. Phosphoric acid does not have this ability because it is not an oxidant.
The acidity of the hydrogen halides
Hydrogen chloride as an acid
By the Bronsted-Lowry definition of an acid as a proton donor, hydrogen chloride is an acid because it transfers protons to other species. Consider its reaction with water.
Hydrogen chloride gas is soluble in water; its solvated form is hydrochloric acid. Hydrogen chloride fumes in moist air are caused by hydrogen chloride reacting with water vapor in the air to produce a cloud of concentrated hydrochloric acid.
A proton is donated from the hydrogen chloride to one of the lone pairs on a water molecule.
A coordinate (dative covalent) bond is formed between the oxygen and the transferred proton.
The equation for the reaction is the following:
$H_2O + HCl \rightarrow H_3O^+ + Cl^- \nonumber$
The H3O+ ion is the hydroxonium ion (also known as the hydronium ion or the oxonium ion). This is the normal form of protons in water; sometimes it is shortened to the proton form, H+(aq), for brevity.
When hydrogen chloride dissolves in water (to produce hydrochloric acid), almost all the hydrogen chloride molecules react in this way. Hydrochloric acid is therefore a strong acid. An acid is strong if it is fully ionized in solution.
Hydrobromic acid and hydriodic acid as strong acids
Hydrogen bromide and hydrogen iodide dissolve in (and react with) water in exactly the same way as hydrogen chloride does. Hydrogen bromide forms hydrobromic acid; hydrogen iodide gives hydriodic acid. Both of these are also strong acids.
Hydrofluoric acid as an exception
By contrast, although hydrogen fluoride dissolves freely in water, hydrofluoric acid is only a weak acid; it is similar in strength to organic acids like methanoic acid. The complicated reason for this is discussed below.
The bond enthalpy of the H-F bond
Because the fluorine atom is so small, the bond enthalpy (bond energy) of the hydrogen-fluorine bond is very high. The chart below gives values for all the hydrogen-halogen bond enthalpies:
bond enthalpy
(kJ mol-1)
H-F +562
H-Cl +431
H-Br +366
H-I +299
In order for ions to form when the hydrogen fluoride reacts with water, the H-F bond must be broken. It would seem reasonable to say that the relative reluctance of hydrogen fluoride to react with water is due to the large amount of energy needed to break that bond, but this explanation does not hold.
The energetics of the process from HX(g) to X-(aq)
The energetics of this sequence are of interest:
All of these terms are involved in the overall enthalpy change as you convert HX(g) into its ions in water.
However, the terms involving the hydrogen are the same for every hydrogen halide. Only the values for the red terms in the diagram need be considered. The values are tabulated below:
bond enthalpy of HX
(kJ mol-1)
electron affinity of X
(kJ mol-1)
hydration enthalpy of X-
(kJ mol-1)
sum of these
(kJ mol-1)
HF +562 -328 -506 -272
HCl +431 -349 -364 -282
HBr +366 -324 -335 -293
HI +299 -295 -293 -289
There is virtually no difference in the total HF and HCl values.
The large bond enthalpy of the H-F bond is offset by the large hydration enthalpy of the fluoride ion. There is a very strong attraction between the very small fluoride ion and the water molecules. This releases a lot of heat (the hydration enthalpy) when the fluoride ion becomes wrapped in water molecules.
Other attractions in the system
The energy terms considered previously have concerned HX molecules in the gas phase. To reach a more correct explanation, the molecules must first be considered as unreacted aqueous HX molecules. The equation for this is given below:
The equation is incorporated into an improved energy cycle as follows:
Unfortunately, values for the first step in the reaction are not readily available. However, in each case, the initial separation of the HX from water molecules is endothermic. Energy is required to break the intermolecular attractions between the HX molecules and water.
That energy is much greater for hydrogen fluoride because it forms hydrogen bonds with water. The other hydrogen halides experience only the weaker van der Waals dispersion forces or dipole-dipole attractions.
The overall enthalpy changes (including all the stages in the energy cycle) for the reactions are given in the table below:
$HX(aq) + H_2O (l) \rightarrow H_3O^+ (aq) + X^- (aq) \nonumber$
enthalpy change
(kJ mol-1)
HF -13
HCl -59
HBr -63
H-I -57
The enthalpy change for HF is much smaller in magnitude than that for the other three hydrogen halides, but it is still negative exothermic change. Therefore, more information is needed to explain why HF is a weak acid.
Entropy and free energy considerations
The free energy change, not the enthalpy change, determines the extent and direction of a reaction.
Free energy change is calculated from the enthalpy change, the temperature of the reaction and the entropy change during the reaction.
For simplicity, entropy can be thought of as a measure of the amount of disorder in a system. Entropy is given the symbol S. If a system becomes more disordered, then its entropy increases. If it becomes more ordered, its entropy decreases.
The key equation is given below:
In simple terms, for a reaction to happen, the free energy change must be negative. But more accurately, the free energy change can be used to calculate a value for the equilibrium constant for a reaction using the following expression:
The term Ka is the equilibrium constant for the reaction below:
$HX(aq) + H_2O(l) \rightleftharpoons H_3O^+(aq) + X^-(aq) \nonumber$
The values for TΔS (needed to calculate ΔG) for the four reactions at a temperature of 298 K are tabulated below:
TS
(kJ mol-1)
HF -29
HCl -13
HBr -4
H-I +4
Notice that at the top of the group, the systems become more ordered when the HX reacts with the water. The entropy of the system (the amount of disorder) decreases, particularly for the hydrogen fluoride.
The reason for this is that the very strong attraction between H3O+ and F-(aq) imposes a lot of order on the system, as does the attraction between the water molecules and the various ions present. These attractions are each greatest for the small fluoride ions.
The total effect on the free energy change, and therefore the value of the equilibrium constant, can now be considered. These values are calculated in the following table:
H
(kJ mol-1)
TS
(kJ mol-1)
G
(kJ mol-1)
Ka
(mol dm-3)
HF -13 -29 +16 1.6 x 10-3
HCl -59 -13 -46 1.2 x 108
HBr -63 -4 -59 2.2 x 1010
HI -57 +4 -61 5.0 x 1010
The values for these estimated equilibrium constants for HCl, HBr and HI are so high that the reaction can be considered "one-way". The ionization is virtually 100% complete. These are all strong acids, increasing in strength down the group.
By contrast, the estimated Ka for hydrofluoric acid is small. Hydrofluoric acid only ionizes to a limited extent in water. Therefore, it is a weak acid.
The estimated value for HF in the table can be compared to the experimental value:
• Experimental value: 5.6 x 10-4 mol dm-3
• Estimated value: 1.6 x 10-3 mol dm-3
These values differ by an order of magnitude, but because of the logarithmic relationship between the free energy and the equilibrium constant, a very small change in ΔG has a very large effect on Ka.
To have the values in close agreement, ΔG would have to increase from +16 to +18.5 kJ mol-1. Given the uncertainty in the values used to calculate ΔG, the difference between the calculated value and the experimental value could easily fall within this range.
Summary: Why is hydrofluoric acid a weak acid?
The two main factors are:
• Entropy decreases dramatically when the hydrogen fluoride reacts with water. This is particularly noticeable with hydrogen fluoride because the attraction of the small fluoride ions produced imposes significant order on the surrounding water molecules and nearby hydronium ions. The effect decreases with larger halide ions.
• Very strong hydrogen bonding exists between the hydrogen fluoride molecules and water molecules. This costs a large amount of energy to break. This effect does not occur in the other hydrogen halides.
Contributors and Attributions
Jim Clark (Chemguide.co.uk) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_17%3A_The_Halogens/1Group_17%3A_General_Reactions/The_Acidity_of_the_Hy.txt |
Fluorine (F) is the first element in the Halogen group (group 17) in the periodic table. Its atomic number is 9 and its atomic weight is 19, and it's a gas at room temperature. It is the most electronegative element, given that it is the top element in the Halogen Group, and therefore is very reactive. It is a nonmetal, and is one of the few elements that can form diatomic molecules (F2). It has 5 valence electrons in the 2p level. Its electron configuration is 1s22s22p5. It will usually form the anion F- since it is extremely electronegative and a strong oxidizing agent. Fluorine is a Lewis acid in weak acid, which means that it accepts electrons when reacting. Fluorine has many isotopes, but the only stable one found in nature is F-19.
Quick Reference Table
Symbol F
Atomic Number 9
Group 17 (Halogens)
Electron Configuration 1s22s22p5
Atomic Weight 18.998 g
Density 1.7 g/L
Melting Point -219.62oC
Boiling Point -188.12oC
Critical Point 144.13K, 5.172 MPa
Oxidation States -1
Electronegativity 3.98
Stable Isotopes F-19
Brief History
In the late 1600's minerals which we now know contain fluorine were used in etching glass. The discovery of the element was prompted by the search for the chemical substance which was able to attack glass (it is HF, a weak acid). The early history of the isolation and work with fluorine and hydrogen fluoride is filled with accidents since both are extremely dangerous. Eventually, electrolysis of a mixture of KF and HF (carefully ensuring that the resulting hydrogen and fluorine would not come in contact) in a platinum apparatus yielded the element.
Fluorine was discovered in 1530 by Georgius Agricola. He originally found it in the compound Fluorspar, which was used to promote the fusion of metals. It was used in this application until 1670, when Schwanhard discovered its usefulness in etching glass. Pure fluorine (from the Latin fluere, for "flow") was not isolated until 1886 by Henri Moissan, burning and even killing many scientists along the way. It has many uses today; a particular one was being used in the Manhattan project to help create the first nuclear bomb.
Electronegativity of Fluorine
Fluorine is the most electronegative element on the periodic table, which means that it is a very strong oxidizing agent and accepts other elements' electrons. Fluorine's atomic electron configuration is 1s22s22p5 (see Figure 2).
Fluorine is the most electronegative element because it has 5 electrons in its 2p shell. The optimal electron configuration of the 2p orbital contains 6 electrons, so since fluorine is so close to ideal electron configuration, the electrons are held very tightly to the nucleus. The high electronegativity of fluorine explains its small radius because the positive protons have a very strong attraction to the negative electrons, holding them closer to the nucleus than the bigger and less electronegative elements do.
Reactions of Fluorine
Because of its reactivity, elemental fluorine is never found in nature, and no other chemical element can displace fluorine from its compounds. Fluorine bonds with almost any element, both metals and nonmetals, because it is a very strong oxidizing agent. It is very unstable and reactive since it is so close to its ideal electron configuration. It forms covalent bonds with nonmetals, and since it is the most electronegative element, is always going to be the element that is reduced. It can also form a diatomic element with itself ($F_2$), or covalent bonds where it oxidizes other halogens ($ClF$, $ClF_3$, $ClF_5$). It will react explosively with many elements and compounds such as hydrogen and water. Elemental fluorine is slightly basic, which means that when it reacts with water it forms $OH^-$.
$2F_2+2H_2O \rightarrow O_2+4HF \tag{1}$
When combined with hydrogen, fluorine forms hydrofluoric acid ($HF$), which is a weak acid. This acid is very dangerous and when dissociated can cause severe damage to the body because while it may not be painful initially, it passes through tissues quickly and can cause deep burns that interfere with nerve function.
$HF+H_2O \rightarrow H_3O^++F^- \tag{2}$
There are also some organic compounds made of fluorine, ranging from nontoxic to highly toxic. Fluorine forms covalent bonds with carbon, which sometimes form into stable aromatic rings. When carbon reacts with fluorine, the reaction is complex and forms a mixture of $CF_4$, $C_2F_6$, and $C_5F_{12}$.
$C_{(s)} + F_{2(g)} \rightarrow CF_{4(g)} + C_2F_6 + C_5F_{12} \tag{3}$
Fluorine reacts with oxygen to form $OF_2$ because fluorine is more electronegative than oxygen. The reaction goes:
$2F_2 + O_2 \rightarrow 2OF_2 \tag{4}$
Fluorine is so electronegative that sometimes it will even form molecules with noble gases like Xenon, such as the molecule xenon difluoride, $XeF_2$.
$Xe + F_2 \rightarrow XeF_2 \tag{5}$
Fluorine also forms strong ionic compounds with metals. Some common ionic reactions of fluorine are:
$F_2 + 2NaOH \rightarrow O_2 + 2NaF +H_2 \tag{6}$
$4F_2 + HCl + H_2O \rightarrow 3HF + OF_2 + ClF_3 \tag{7}$
$F_2 + 2HNO_3 \rightarrow 2NO_3F + H_2 \tag{8}$
Applications of Fluorine
Compounds of fluorine are present in fluoridated toothpaste and in many municipal water systems, where they help to prevent tooth decay. And, of course, fluorocarbons such as Teflon have made a major impact on life in the 20th century. There are many applications of fluorine:
• Rocket fuels
• Polymer and plastics production
• Teflon and tefzel production
• When combined with oxygen, used as a refrigerator cooler
• Hydrofluoric acid used for glass etching
• Purify public water supplies
• Uranium production
• Air conditioning
Sources
Fluorine can either be found in nature or produced in a lab. To make it in a lab, compounds like potassium fluoride are put through electrolysis with hydrofluoric acid to create pure Fluorine and other compounds. It can be carried out with a variety of compounds, usually ionic ones involving fluorine and a metal. Fluorine can also be found in nature in various minerals and compounds. The two main compounds it can be found in are Fluorspar ($CaF_2$) and Cryolite ($Na_3AlF_6$).
Problems
(Highlight to view answers)
1. Q. What is the electron configuration of fluorine? of F-?
A. 1s22s22p5
1s22s22p6
2. Q. Is fluorine usually oxidized or reduced? explain.
A. Fluorine is usually reduced because it accepts an electron from other elements since it is so electronegative.
3. Q. What are some common uses of fluorine?
A. Toothpaste, plastics, rocket fuels, glass etching, etc.
4. Q. Does fluorine form compounds with nonmetals? if so, give two examples, one of them being of an oxide.
A. OF2, ClF
5. Q. What group is fluorine in? (include name of group and number)
A. 17, Halogens
Contributors and Attributions
• Rachel Feldman (University of California, Davis)
• Stephen R. Marsden | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_17%3A_The_Halogens/Z009_Chemistry_of_Fluorine_%28Z9%29.txt |
Chlorine is a halogen in group 17 and period 3. It is very reactive and is widely used for many purposes, such as as a disinfectant. Due to its high reactivity, it is commonly found in nature bonded to many different elements.
Chlorine, which is similar to fluorine but not as reactive, was prepared by Sheele in the late 1700's and shown to be an element by Davy in 1810. It is a greenish-yellow gas with a disagreeable odor (you can detect it near poorly balanced swimming pools). Its name comes from the Greek word chloros, meaning greenish-yellow. In high concentration it is quite toxic and was used in World War I as a poison gas.
Properties
Atomic Number 17
Atomic Weight 35.457
Electron Configuration [Na]3s23p5
1st Ionization Energy 1251 kJ/mol
Ionic Radius 181 pm
Density (Dry Gas) 3.2 g/L
Melting Point -101°C
Boiling Point -34.05°C
Specific Heat 0.23 g cal/g/°C
Heat of Vaporization 68 g cal/g
Heat of Fusion 22 g cal/g
Critical Temperature 114°C
Standard Electron Potential
$Cl_2 + 2e^- \rightarrow 2Cl^-$
1.358V
At room temperature, pure chlorine is a yellow-green gas. Chlorine is easily reduced, making it a good oxidation agent. By itself, it is not combustible, but many of its reactions with different compounds are exothermic and produce heat. Because chlorine is so highly reactive, it is found in nature in a combined state with other elements, such as NaCl (common salt) or KCl (sylvite). It forms strong ionic bonds with metal ions.
Like fluorine and the other members of the halogen family, chlorine is diatomic in nature, occurring as $Cl_2$ rather than Cl. It forms -1 ions in ionic compounds with most metals. Perhaps the best known compound of that type is sodium chloride, common table salt (NaCl).
Small amounts of chlorine can be produced in the lab by oxidizing $HCl$ with $MnO_2$. On an industrial scale, chlorine is produced by electrolysis of brines or even sea water. Sodium hydroxide (also in high demand) is a by-product of the process.
In addition to the ionic compounds that chlorine forms with metals, it also forms molecular compounds with non-metals such as sulfur and oxygen. There are four different oxides of the element. Hydrogen chloride gas (from which we get hydrochloric acid) is an important industrial product.
Reactions with Water
Usually, reactions of chlorine with water are for disinfection purposes. Chlorine is only slightly soluble in water, with its maximum solubility occurring at 49° F. After that, its solubility decreases until 212° F. At temperatures below that range, it forms crystalline hydrates (usually $Cl_2$) and becomes insoluble. Between that range, it usually forms hypochlorous acid ($HOCl$). This is the primary reaction used for water/wastewater disinfection and bleaching.
$Cl_2+H_2O \rightarrow HOCl + HCl \nonumber$
At the boiling temperature of water, chlorine decomposes water:
$2Cl_2+2H_2O \rightarrow 4HCl + O_2 \nonumber$
Reactions with Oxygen
Although chlorine usually has a -1 oxidation state, it can have oxidation states of +1, +3, +4, or +7 in certain compounds, such as when it forms oxoacids with the alkali metals.
Oxidation State Compound
+1 NaClO
+3 NaClO2
+5 NaClO3
+7 NaClO4
Reactions with Hydrogen
When H2 and Cl2 are exposed to sunlight or high temperatures, they react quickly and violently in a spontaneous reaction. Otherwise, the reaction proceeds slowly.
$H_2+Cl_2 \rightarrow 2HCl \nonumber$
HCl can also be produced by reacting chlorine with compounds containing hydrogen, such as hydrogen sulfide
Reactions with Halogens
Chlorine, like many of the other halogens, can form interhalogen compounds (examples include BrCl, ICl, ICl2). The heavier element in one of these compounds acts as the central atom. For chlorine, this occurs when it is bonded to fluorine in ClF, ClF3, and ClF5.
Reactions with Metals
Chlorine reacts with most metals and forms metal chlorides, with most of these compounds being soluble in water. Examples of insoluble compounds include $AgCl$ and $PbCl_2$. Gaseous or liquid chlorine usually does not have an effect on metals such as iron, copper, platinum, silver, and steel at temperatures below 230°F. At high temperatures, however, it reacts rapidly with many of the metals, especially if the metal is in a form that has a high surface area (such as when powdered or made into wires).
Example: Oxidizing Iron
Chlorine can oxidize iron:
$Cl_2+Fe \rightarrow FeCl_2 \nonumber$
Half Reactions:
$Fe \rightarrow Fe^{+2} +2e^- \nonumber$
$Cl_2+2e^- \rightarrow 2Cl^- \nonumber$
Isotopes
$\ce{^35}Cl$ and $\ce{^37}Cl$ are the two natural, stable isotopes of chlorine. $\ce{^36}Cl$, a radioactive isotope, occurs only in trace amounts as a result of cosmic rays in the atmosphere. Chlorine is usually a mixture of 75% $\ce{^35}Cl$ and 25% $\ce{^37}Cl$. Besides these isotopes, the other isotopes must be artificially produced. A table containing some common isotopes is found below:
Isotope Atomic Mass Half-Life
$\ce{^33}Cl$ 32.986 2.8 seconds
$\ce{^34}Cl$ 33.983 33 minutes
$\ce{^35}Cl$ 34.979 Stable ($\infty$)
$\ce{^36}Cl$ 35.978 400,000 years
$\ce{^37}Cl$ 35.976 Stable ($\infty$)
$\ce{^38}Cl$ 37.981 39 Minutes
Production and Uses
Chlorine is a widely used chemical with many applications.
Water Treatment
Chlorine is used in the disinfection (removal of harmful microorganisms) of water and wastewater. In the United States, it is almost exclusively used. Chlorine was first used to disinfect drinking water in 1908, using sodium hypochlorite (NaOCl):
$NaOCl+ H_2O \rightarrow HOCl+NaOH \nonumber$
Following widespread use of sodium hypochlorite to disinfect water, diseases caused by unclean water decreased greatly. Compared to other methods, it is effective at lower concentrations and is inexpensive.
Polyvinyl Chloride (PVC)
Polyvinyl chloride is a plastic which is widely manufactured throughout the globe, and is responsible for nearly a third of the world’s use of chlorine. It is usually manufactured by first taking EDC (ethylene dichloride) and then making it into a vinyl chloride, the basic unit for PVC. From then on, vinyl chloride monomers are linked together to form a polymer. PVC becomes malleable at high temperatures, making it flexible and ideal for many purposes from pipes to clothing. However, PVC is toxic. When in gaseous form and inhaled, it can cause damage to the lungs, the body’s blood circulation, and the nervous system. The production of PVC has many regulations surrounding it due to the many harmful effects that the plastic itself and the intermediates involved have on the environment and on human health.
Paper Bleaching
Paper is one of the most widely consumed products in the world. Before wood is made into a paper product, however, it must be turned into pulp (separated fibrous material). This pulp has a color that ranges from light to dark brown. Chlorine is used to bleach the pulp to turn it into a bright, white color, which makes it desirable for consumers. The process usually involves a number of steps, depending on the nature of the pulp.
Problems
1) Solve and balance the following equations
1. $H_2S + Cl_2 + H_2O \rightarrow$
2. $Sb + Cl_2 +H_2O \rightarrow$
2) Write the electron configuration for chlorine.
3) What is the molecular geometry of the following? (See Valence Bond Theory)
1. $ClO_2$
2. $ClF_5$
4) What are the naturally occurring chlorine isotopes?
5) When does chlorine have an oxidation state of +5?
Answers
1) Solve and balance the following equations:
1. H2S + 4Cl2 + 4H20 --> H2S04 + 8HCl
2. 2Sb + 3Cl2 --> 2SbCl3
2) The electron configuration of chlorine is: 1s22s22p63s23p5
3) What is the molecular geometry of the following?
1. $ClO_2$ -Bent or angular; the Cl is bonded to two ligands, has one lone pair and one unpaired electron.
2. $ClF_5$ -Square pyramid; the Cl is bonded to five ligands and has one lone pair.
4) The naturally occurring chlorine isotopes are chlorine-35 and chlorine-37. While chlorine-36 does occur naturally, it is radioactive and unstable.
5) Chlorine has an oxidation state of +5 when it reacts with oxoacids with the Alkali Metals.
References
1. Sconce, J.S. Chlorine: Its Manufacture, Properties, and Uses. Reinhold Corporation, 1962.
2. Stringer, Ruth, and Paul Johnston. Chlorine and the Enviroment. Norwell: Kluwer Academic, 2001.
3. Reynolds, Tom D. Unit Operations and Processes in Environmental Engineering. Brooks/Cole Engineering Division, a Division of Wadsworth Inc, 1982. 523-532
4. Davis, Stanley N., DeWayne Cecil, Marek Zreda, and Pankaj Sharma. "Chlorine-36 and the Initial Value Problem." Hydrogeology Journal 6.1 (1998): 104-14. SpringerLink. Web. 23 May 2010. <www.springerlink.com/content/3205uburlwx2x48g/>
5. Pettrucci, Ralph H. General Chemistry: Principles and Modern Applications. 9th. Upper Saddle River: Pearson Prentice Hall, 2007
Contributors and Attributions
• Judy Hsia (University of California, Davis)
Z017 Chemistry of Chlorine (Z17)
This page describes the manufacture of chlorine by the electrolysis of sodium chloride solution using a diaphragm cell and a membrane cell. Both cells rely on the same underlying chemistry, but differ in detail.
Background chemistry
Chlorine is manufactured by electrolyzing sodium chloride solution. This process generates three useful substances: chlorine, sodium hydroxide, and hydrogen.
The chemistry of the electrolysis process
Sodium chloride solution contains the following:
• sodium ions
• chloride ions
• protons (from the water)
• hydroxide ions (from the water)
The protons and hydroxide ions come from the equilibrium:
$H_2O (l) \rightleftharpoons H^+(aq) + OH^- (aq) \label{1}$
At any time, the concentration of protons or hydroxide ions is very small; the position of equilibrium lies well to the left.
At the anode
The negative ions, chloride and hydroxide, are attracted to the positively charged anode. It is easier to oxidize hydroxide ions to oxygen than to oxidize chloride ions to chlorine, but there are far more chloride ions arriving at the anode than hydroxide ions.
The major reaction at the anode is therefore:
$2Cl^-_{(aq)}\rightarrow Cl_{2(g)}+2e^- \label{2}$
Two chloride ions each give up an electron to the anode, and the atoms produced combine into chlorine gas. The chlorine is, however, contaminated with small amounts of oxygen because of a reaction involving hydroxide ions, which also give up electrons:
$4OH^-_{(aq)} \rightarrow 2H_2O_{(l)} + O_{2(g)} + 4e^- \label{3}$
The chlorine must be purified by removing this oxygen.
At the cathode
Sodium ions and protons (from the water) are attracted to the negative cathode. It is much easier for a proton to pick up an electron than for a sodium ion to do so. Therefore, the following reaction occurs:
$2H^+_{(aq)} + 2e^- \rightarrow H_{2(g)} \label{4}$
As protons convert into hydrogen gas, the equilibrium below shifts to the right to replace them:
The net effect of this process is a buildup of sodium ions and newly-produced hydroxide ions at the cathode. In other words, sodium hydroxide solution is formed.
The necessity of keeping all products separate
If chlorine comes into contact with hydrogen, it produces a mixture that explodes violently on exposure to sunlight or heat, producing hydrogen chloride gas. Clearly these gases must remain separated. However, chlorine also reacts with sodium hydroxide solution to produce a mixture of sodium chloride and sodium chlorate(I), also known as sodium hypochlorite; this mixture is commonly sold as bleach. In addition, when the desired products are chlorine and sodium hydroxide rather than bleach, chlorine and sodium hydroxide must also be kept apart. The diaphragm and membrane cells are designed to keep all the products separate.
The diaphragm
The diaphragm is made of a porous mixture of asbestos and polymers. The solution can seep through it from the anode compartment into the cathode compartment. Notice that there is a higher level of liquid on the anode side. This ensures that the liquid always flows from left to right, preventing any of the sodium hydroxide solution from coming into contact with chlorine products.
Production of chlorine at the anode
Chlorine is produced at the titanium anode according to the following equation:
$2Cl^- (aq) - 2e^- \rightarrow Cl_2(g) \nonumber$
The product is contaminated with some oxygen because of the reaction below:
$4OH^-(aq) - 4e^- \rightarrow 2H_2O(l) + O_2(g) \nonumber$
The chlorine is purified by liquefaction under pressure. The oxygen remains a gas when compressed at ordinary temperatures.
Production of hydrogen at the cathode
Hydrogen is produced at the steel cathode by the following process:
$2H^+ (aq) + 2e^- \rightarrow H_2 (g) \nonumber$
Production of the sodium hydroxide
A dilute solution of sodium hydroxide solution is also produced at the cathode (see above for the explanation of what happens at the cathode). It is highly contaminated with unreacted sodium chloride.
The sodium hydroxide solution leaving the cell is concentrated by evaporation. During this process, most of the sodium chloride crystallizes out as solid salt. The salt can be separated, dissolved in water, and passed through the cell again. Even after concentration, sodium hydroxide still contains a small percentage of sodium chloride.
The membrane
The membrane is made from a polymer that only allows the diffusion of positive ions. That means that only the sodium ions can pass through the membrane; the chloride ions are blocked. The advantage of this is that the sodium hydroxide formed in the right-hand compartment is never contaminated with sodium chloride. The sodium chloride solution must be pure. If it contains any other metal ions, these can also pass through the membrane and so contaminate the sodium hydroxide solution.
Production of chlorine
Chlorine is produced at the titanium anode according to the following equation:
$2Cl^- (aq) + 2e^- \rightarrow Cl_2(g) \nonumber$
It is contaminated with some oxygen because of the parallel reaction below:
$4OH^-(aq) + 4e^- \rightarrow 2H_2O(l) + O_2 \nonumber$
The chlorine is purified by liquefaction under pressure. The oxygen remains in the gas phase when compressed at normal temperatures.
Production of hydrogen
Hydrogen is produced at the nickel cathode as follows:
$2H^+ (aq) +2e^- \rightarrow H_2 (g) \nonumber$
Production of sodium hydroxide
An approximately 30% solution of sodium hydroxide solution is also produced at the cathode (see the background chemistry section for an explanation of what happens at the cathode).
Z035 Chemistry of Bromine (Z35)
Bromine is a reddish-brown fuming liquid at room temperature with a very disagreeable chlorine-like smell. In fact its name is derived from the Greek bromos or "stench". It was first isolated in pure form by Balard in 1826. It is the only non-metal that is a liquid at normal room conditions. Bromine on the skin causes painful burns that heal very slowly. It is an element to be treated with the utmost respect in the laboratory.
Most bromine is produced by displacement from ordinary sea water. Chlorine (which is more active) is generally used to dislodge the bromine from various compounds in the water. Before leaded gasolines were removed from the market, bromine was used in an additive to help prevent engine "knocking". Production now is chiefly devoted to dyes, disinfectants and photographic chemicals.
Contributors and Attributions
Stephen R. Marsden
Z053 Chemistry of Iodine (Z53)
Elemental iodine is a dark grey solid with a faint metallic luster. When heated at ordinary air pressures it sublimes to a violet gas. The name iodine is taken from the Greek ioeides which means "violet colored". It was discovered in 1811 by Courtois.
Commercially iodine is recovered from seaweed and brines. It is an important trace element in the human diet, required for proper function of the thyroid gland. Thus iodine is added to table salt ("iodized") to insure against iodine deficiencies. Radioactive isotopes of iodine are used in medical tracer work involving the thyroid and also to treat diseases of that gland.
Contributors and Attributions
Stephen R. Marsden
Z085 Chemistry of Astatine (Z85)
Astatine was formerly known as alabamine. It has no stable isotopes and was first synthetically produced (1940) at the University of California.
Name: Astatine
Symbol: At
Atomic Number: 85
Atomic Mass: (210.0) amu
Melting Point: 302.0 °C (575.15 K, 575.6 °F)
Boiling Point: 337.0 °C (610.15 K, 638.6 °F)
Number of Protons/Electrons: 85
Number of Neutrons: 125
Classification: Halogen
Crystal Structure: Unknown
Density @ 293 K: Unknown
Color: Unknown
Date of Discovery: 1940
Discoverer: D.R. Corson
Name Origin: From the Greek word astatos (unstable)
Uses: No uses known
Obtained From: Man-made
Oxidation Number: -1, +5
Astatine is the last of the known halogens and was synthesized in 1940 by Corson and others at the University of California. It is radioactive and its name, from the Greek astatos, means "unstable". The element can be produced by bombarding targets made of bismuth-209 with high energy alpha particles (helium nuclei). Astatine 211 is the product and has a half-life of 7.2 hours. The most stable isotope of astatine is 210, which has a half-life of 8.1 hours.
Not much is known about the chemical properties of astatine, but it is expected to react like the other halogens, although much less vigorously, and it should be more metallic than iodine. There should be tiny quantities of astatine in the earth's crust as products of other radioactive decays, but their existence would be short-lived.
Contributors and Attributions
Stephen R. Marsden | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_17%3A_The_Halogens/Z017_Chemistry_of_Chlorine_%28Z17%29/The_Manufacture.txt |
The noble gases (Group 18) are located in the far right of the periodic table and were previously referred to as the "inert gases" due to the fact that their filled valence shells (octets) make them extremely nonreactive. The noble gases were characterized relatively late compared to other element groups.
Thumbnail: Vial of glowing ultrapure neon. (CC SA; Jurii via http://images-of-elements.com/neon.php).
Group 18: The Noble Gases
The noble gases (Group 18) are located in the far right of the periodic table and were previously referred to as the "inert gases" due to the fact that their filled valence shells (octets) make them extremely nonreactive. The noble gases were characterized relatively late compared to other element groups.
The History
The first person to discover the noble gases was Henry Cavendish in the late 180th century. Cavendish distinguished these elements by chemically removing all oxygen and nitrogen from a container of air. The nitrogen was oxidized to $NO_2$ by electric discharges and absorbed by a sodium hydroxide solution. The remaining oxygen was then removed from the mixture with an absorber. The experiment revealed that 1/120 of the gas volume remained un-reacted in the receptacle. The second person to isolate, but not typify, them was William Francis (1855-1925). Francis noted the formation of gas while dissolving uranium minerals in acid.
Argon
In 1894, John William Strutt discovered that chemically-obtained pure nitrogen was less dense than the nitrogen isolated from air samples. From this breakthrough, he concluded that another, unknown gas was present in the air. With the aid of William Ramsay, Strutt managed to replicate and modify Cavendish's experiment to better understand the inert component of air in his original experiment. The researchers' procedure differed from the Cavendish procedure: they removed the oxygen by reacting it with copper, and removed the nitrogen in a reaction with magnesium. The remaining gas was properly characterized and the new element was named "argon," which originates from the Greek word for "inert."
Helium
Helium was first discovered in 1868, manifesting itself in the solar spectrum as a bright yellow line with a wavelength of 587.49 nanometers. This discovery was made by Pierre Jansen. Jansen initially assumed it was a sodium line. However, later studies by Sir William Ramsay (who isolated helium on Earth by treating a variety of rare elements with acids) confirmed that the bright yellow line from his experiment matched up with that in the spectrum of the sun. From this, British physicist William Crookes identified the element as helium.
Neon, Krypton, Xenon
These three noble gases were discovered by Morris W. Travers and Sir William Ramsay in 1898. Ramsay discovered neon by chilling a sample of the air to a liquid phase, warming the liquid, and capturing the gases as they boiled off. Krypton and xenon were also discovered through this process.
Radon
In 1900, while studying the decay chain of radium, Friedrich Earns Dorn discovered the last gas in Group 18: radon. In his experiments, Dorn noticed that radium compounds emanated radioactive gas. This gas was originally named niton after the Latin word for shining, "nitens". In 1923, the International Committee for Chemical Elements and International Union of Pure Applied Chemistry (IUPAC) decided to name the element radon. All isotopes of radon are radioactive. Radon-222 has the longest half-life at less than 4 days, and is an alpha-decay product of Radium-226 (part of the U-238 to Pb-206 radioactive decay chain).
The Electron Configurations for Noble Gases
• Helium 1s2
• Neon [He] 2s2 2p6
• Argon [Ne] 3s2 3p6
• Krypton [Ar] 3d10 4s2 4p6
• Xenon [Kr] 4d10 5s2 5p6
• Radon [Xe] 4f14 5d10 6s2 6p6
Table 1: Trends within Group 18
Atomic # Atomic mass Boiling point (K) Melting point (K) 1st Ionization (E/kJ mol-1) Density (g/dm3) Atomic radius (pm)
He 2 4.003 4.216 0.95 2372.3 0.1786 31
Ne 10 20.18 27.1 24.7 2080.6 0.9002 38
Ar 18 39.948 87.29 83.6 1520.4 1.7818 71
Kr 36 83.3 120.85 115.8 1350.7 3.708 88
Xe 54 131.29 166.1 161.7 1170.4 5.851 108
Rn 86 222.1 211.5 202.2 1037.1 9.97 120
The Atomic and Physical Properties
• Atomic mass, boiling point, and atomic radii INCREASE down a group in the periodic table.
• The first ionization energy DECREASES down a group in the periodic table.
• The noble gases have the largest ionization energies, reflecting their chemical inertness.
• Down Group 18, atomic radius and interatomic forces INCREASE resulting in an INCREASED melting point, boiling point, enthalpy of vaporization, and solubility.
• The INCREASE in density down the group is correlated with the INCREASE in atomic mass.
• Because the atoms INCREASE in atomic size down the group, the electron clouds of these non polar atoms become increasingly polarized, which leads to weak van Der Waals forces among the atoms. Thus, the formation of liquids and solids is more easily attainable for these heavier elements because of their melting and boiling points.
• Because noble gases’ outer shells are full, they are extremely stable, tending not to form chemical bonds and having a small tendency to gain or lose electrons.
• Under standard conditions all members of the noble gas group behave similarly.
• All are monotomic gases under standard conditions.
• Noble gas atoms, like the atoms in other groups, INCREASE steadily in atomic radius from one period to the next due to the INCREASING number of electrons.
• The size of the atom is positively correlated to several properties of noble gases. The ionization potential DECREASES with an INCREASING radius, because the valence electrons in the larger noble gases are further away from the nucleus; they are therefore held less tightly by the atom.
• The attractive force INCREASES with the size of the atom as a result of an INCREASE in polarizability and thus a DECREASE in ionization potential.
• Overall, noble gases have weak interatomic forces, and therefore very low boiling and melting points compared with elements of other groups.
For covalently-bonded diatomic and polyatomic gases, heat capacity arises from possible translational, rotational, and vibrational motions. Because monatomic gases have no bonds, they cannot absorb heat as bond vibrations. Because the center of mass of monatomic gases is at the nucleus of the atom, and the mass of the electrons is negligible compared to the nucleus, the kinetic energy due to rotation is negligible compared to the kinetic energy of translation (unlike in di- or polyatomic molecules where rotation of nuclei around the center of mass of the molecule contributes significantly to the heat capacity). Therefore, the internal energy per mole of a monatomic noble gas equals its translational contribution, $\frac{3}{2}RT$, where $R$ is the universal gas constant and $T$ is the absolute temperature.
For monatomic gases at a given temperature, the average kinetic energy due to translation is practically equal regardless of the element. Therefore at a given temperature, the heavier the atom, the more slowly its gaseous atoms move. The mean velocity of a monatomic gas decreases with increasing molecular mass, and given the simplified heat capacity situation, noble gaseous thermal conductivity decreases with increasing molecular mass.
Applications of Noble Gases
Helium
Helium is used as a component of breathing gases due to its low solubility in fluids or lipids. This is important because other gases are absorbed by the blood and body tissues when under pressure during scuba diving. Because of its reduced solubility, little helium is taken into cell membranes; when it replaces part of the breathing mixture, helium causes a decrease in the narcotic effect of the gas at far depths. The reduced amount of dissolved gas in the body means fewer gas bubbles form, decreasing the pressure of the ascent. Helium and Argon are used to shield welding arcs and the surrounding base metal from the atmosphere.
Helium is used in very low temperature cryogenics, particularly for maintaining superconductors (useful for creating strong magnetic fields) at a very low temperatures. Helium is also the most common carrier gas in gas chromatography.
Neon
Neon has many common and familiar applications: neon lights, fog lights, TV cine-scopes, lasers, voltage detectors, luminous warnings, and advertising signs. The most popular application of neon is the neon tubing used in advertising and elaborate decorations. These tubes are filled with neon and helium or argon under low pressure and submitted to electrical discharges. The color of emitted light is depends on the composition of the gaseous mixture and on the color of the glass of the tube. Pure Neon within a colorless tube absorbs red light and reflects blue light, as shown in the figure below. This reflected light is known as fluorescent light.
One of the many colors of neon lights.
Argon
Argon has a large number of applications in electronics, lighting, glass, and metal fabrications. Argon is used in electronics to provide a protective heat transfer medium for ultra-pure silicon crystal semiconductors and for growing germanium. Argon can also fill fluorescent and incandescent light bulbs, creating the blue light found in "neon lamps." By utilizing argon's low thermal conductivity, window manufacturers provide a gas barrier needed to produce double-pane insulated windows. This insulation barrier improves the windows' energy efficiencies. Argon also creates an inert gas shield during welding, flushes out melted metals to eliminate porosity in casting, and provides an oxygen- and nitrogen-free environment for annealing and rolling metals and alloys.
Argon plasma light bulb.
Krypton
Similarly to argon, krypton can be found in energy efficient windows. Because of its superior thermal efficiency, krypton is sometimes chosen over argon for insulation. It is estimated that 30% of energy efficient windows sold in Germany and England are filled with krypton; approximately 1.8 liters of krypton are used in these countries. Krypton is also found in fuel sources, lasers and headlights. In lasers, krypton functions as a control for a desired optic wavelength. It is usually mixed with a halogen (most likely fluorine) to produce excimer lasers. Halogen sealed beam headlights containing krypton produce up to double the light output of standard headlights. In addition, Krypton is used for high performance light bulbs, which have higher color temperatures and efficiency because the krypton reduces the rate of evaporation of the filament.
Krypton laser.
Xenon
Xenon has various applications in incandescent lighting, x-ray development, plasma display panels (PDPs), and more. Incandescent lighting uses xenon because less energy can be used to obtain the same light output as a normal incandescent lamp. Xenon has also made it possible to obtain better x-rays with reduced amounts of radiation. When mixed with oxygen, it can enhance the contrast in CT imaging. These applications have had great impact on the health care industries. Plasma display panels (PDPs) using xenon as one of the fill gases may one day replace the large picture tubes in television and computer screens.
Nuclear fission products may include several radioactive isotopes of xenon, which absorb neutrons in nuclear reactor cores. The formation and elimination of radioactive xenon decay products are factors in nuclear reactor control.
Radon
Radon is reported as the second most frequent cause of lung cancer, after cigarette smoking. However, it also has beneficial applications in radiotherapy, arthritis treatment, and bathing. In radiotherapy, radon has been used in implantable seeds, made of glass or gold, primarily used to treat cancers. It has been said that exposure to radon mitigates auto-immune diseases such as arthritis. Some arthritis sufferers have sought limited exposure to radioactive mine water and radon to relieve their pain. "Radon Spas" such as Bad Gastern in Austria and Onsen in Japan offer a therapy in which people sit for minutes to hours in a high-radon atmosphere, believing that low doses of radiation will boost up their energy. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_18%3A_The_Noble_Gases/1Group_18%3A_Properties_of_Nobel_Gases.txt |
The noble gases (Group 18) are located in the far right of the periodic table and were previously referred to as the "inert gases" due to the fact that their filled valence shells (octets) make them extremely nonreactive.
The Chemical Properties
Noble gases are odorless, colorless, nonflammable, and monotonic gases that have low chemical reactivity.
Atomic Number Element Number of Electrons/Shell
2 Helium 2
10 Neon 2,8
18 Argon 2,8,8
36 Krypton 2,8,18,8
54 Xenon 2,8,18,18,8
86 Radon 2,8,18,32,18,8
The full valence electron shells of these atoms make noble gases extremely stable and unlikely to form chemical bonds because they have little tendency to gain or lose electrons. Although noble gases do not normally react with other elements to form compounds, there are some exceptions. Xe may form compounds with fluoride and oxide.
Example 1: Xenon Fluorides
Xenon Difluoride (\(XeF_2\))
• Dense white crystallized solid
• Powerful fluorinating agent
• Covalent inorganic fluorides
• Stable xenon compound
• Decomposes on contact with light or water vapor
• Linear geometry
• Moisture sensitive
• Low vapor pressure
Xenon Tetrafluoride (\(XeF_4\))
• Colorless Crystals
• Square planar geometry
• Discovered in 1963
Xenon Hexafluoride (\(XeF_6\))
• Strongest fluorinating agent
• Colorless solid
• Highest coordination of the three binary fluorides of xenon (\(XeF_2\) and \(XeF_4\))
• Formation is exergonic, and the compound is stable at normal temperatures
• Readily sublimes into intense yellow vapors
• Structure lacks perfect octahedral symmetry
Example 2: Xenon Oxide
Xenon Tetroxide (XeO4)
• Yellow crystalline solid
• Relatively stable
• Oxygen is the only element that can bring xenon up to its highest oxidation state of +8
Two other short-lived xenon compounds with an oxidation state of +8, XeO3F2 and XeO2F4, are produced in the reaction of xenon tetroxide with xenon hexafluoride.
Example 3: Radon Compounds
Radon difluoride (RnF2) is one of the few reported compounds of radon. Radon reacts readily with fluorine to form a solid compound, but this decomposes on attempted vaporization and its exact composition is uncertain. The usefulness of radon compounds is limited because of the noble gas's radioactivity. The longest-lived isotope, 222Ra, has a half-life of only 3.82 days.
Z002 Chemistry of Helium (Z2)
Helium is at the top of the noble gas group (which also contains neon, argon, krypton, xenon, and radon) and is the least reactive element. Helium has many interesting characteristics, such as making balloons float and raising the pitch of one's voice; these applications are discussed below.
Introduction
Helium is the second most abundant element in the universe, next to hydrogen. Helium is colorless, odorless, and tasteless. It has a very low boiling point, and is monatomic. Helium is small and extremely light, and is the least reactive of all elements; it does not react with any other elements or ions, so there are no helium-bearing minerals in nature. Helium was first observed by studying the sun, and was named after the Greek word for the sun, Helios.
Physical Properties
Color Colorless
Phase at Room Temperature Gas
Density 0.0002 g/cm3
Boiling Point 4.2 K
Heat of Vaporization 0.1 kJ/mol
Thermal Conductivity 0.15 J/m sec K
Source Natural gas
Atomic Properties
Electron Configuration 1s2
Number of Isotopes 7 (2 liquid)
Electron Affinity 0 kJ/mol
First Ionization Energy 2372.3 kJ/mol
Second Ionization Energy 5250.3 kJ/mol
Polarizability 0.198 Å3
Atomic Weight 4.003
Atomic Volume 27.2 cm3/mol
Atomic Radius 31 pm
Abundance
In Earth's Crust 8x10-3
In Earth's Ocean 7×10-6
In Human Body 0%
Occurrence and production
Helium is one of the most abundant elements in the universe. Large quantities are produced in the energy-producing fusion reactions in stars. Previously, helium was rarely used, because only .0004% of Earth's atmosphere is helium—that equates to one helium molecule for every 200,000 air molecules, including oxygen, hydrogen, and nitrogen. However, the discovery of helium-rich wells in Texas, Russia, Poland, Algeria, China, and Canada has made helium more accessible.
Helium is produced in minerals through radioactive decay. Helium is extracted from natural gas deposits, which often contain as much as 10% helium. These natural gas reserves are the only industrially-available source of helium. The total world helium resources theoretically add up to 25.2 billion cubic meters; the United States contains 11.1 billion cubic meters. The extracted gas is subjected to chemical pre-purification, using an alkaline wash to remove carbon dioxide and hydrogen sulfide. The remaining gas is cooled to -200°C, where all materials, except helium gas, are liquefied.
History
Helium was first discovered in 1868 by the French astronomer P. J. C. Jenssen, who was studying the chromosphere of the Sun during a solar eclipse. He used a spectrometer to resolve the light into its spectrum, in which each color represents a different gaseous element. He observed a new yellow light, concluding that it indicated the presence of an element not previously known. In 1895, the existence of helium on Earth was proved by Sir William Ramsay. Heating cleveite (a radioactive mineral) released an inert gas, which was found to be helium; this helium is a by-product of the natural decay of radioactive elements. The chemists Norman Lockyer and Edward Frankland confirmed helium as an element and named it after helios, the Greek word for the Sun.
Applications and hazards
Helium has a number of applications due to its inert nature. Liquefied helium has cryogenic properties, and is used to freeze biological materials for long term storage and later use. Twenty percent of industrial helium use is in wielding and industrial applications. Helium protects the heated parts of metals such as aluminum and titanium from air. Mixtures of helium and oxygen are used in tanks for underwater breathing devices: due to its low density, helium gas allows oxygen to stream easily through the lungs. Because helium remains a gas, even at temperatures low enough to liquefy hydrogen, it is used as pressure gas to move liquid hydrogen into rocket engines. Its inert nature also makes helium useful for cooling nuclear power plants.
The most commonly known characteristic of helium is that it is lighter than air. It can levitate balloons during parties and fly blimps over sports stadiums. Helium has 92% of the lifting power of hydrogen; however, it is safer to use because it is noncombustible and has lower rate of diffusion than that of hydrogen gas. The famous Hindenburg disaster is an example of the hazards of using combustible gas like hydrogen. Because helium was previously very expensive only available from natural gas reserves in U.S., Nazi Germany had only hydrogen gas at its disposal. The consequences were devastating, as shown below:
Currently, helium is found in other natural gas reserves around the world. The cost of helium has decreased from \$2500/ft3 in 1915 to \$0.15/ft3 in 1989. Helium is what keeps the Goodyear blimps afloat over stadiums.
Helium is often inhaled from balloons to produce a high, squeaky voice. This practice can be very harmful. Inhaling helium can lead to loss of consciousness and cerebral arterial gas embolism, which can temporarily lead to complete blindness. This occurs when blood vessels in the lungs rupture, allowing the gas to gain access to the pulmonary vasculature and subsequently the brain.
Characteristics
Gas and plasma phases
Helium is naturally found in the gas state. Helium is the second least reactive element and noble gas (after neon). Its low atomic mass, thermal conductivity, specific heat, and sound speed are greatest after hydrogen. Due to the small size of helium atoms, the diffusion rate through solids is three times greater than that of air and 65% greater than that of hydrogen. The element is inert, monatomic in standard conditions, and the least water soluble gas.
At normal ambient temperatures, helium has a negative Joule Thomson coefficient. Thus, upon free expansion, helium naturally heats up. However, below its Joule Thomson inversion temperature (32-50 K at 1 atm), it cools when allowed to freely expand. Once cooled, helium can be liquefied through expansion cooling. Helium is commonly found throughout the universe as plasma, a state in which electrons are not bound to nuclei. Plasmas have high electrical conductivities and are highly influenced by magnetic and electric fields.
Solid and liquid phases
Helium is the only element that cannot be solidified by lowering the temperature at ordinary pressures; this must be accompanied by a pressure increase. The volume of solid helium, 3He and 4He, can be decreased by more than 30% by applying pressure. Solid helium has a projected density of 0.187 ± 0.009 g/mL at 0 K and 25 bar. Solid helium also has a sharp melting point and a crystalline structure. There are two forms of liquid helium: He4I and He4II.
Helium I
Helium I is formed when temperature falls below 4.22 K and above the lambda point of 2.1768 K. It is a clear liquid that boils when heat is applied and contracts when temperature is lowered. Below the lambda point, helium does not boil, but expands. Helium I has a gas-like index of refraction of 1.026 which makes its surface difficult to see. It has a very low viscosity and a density 1/8th that of water. This property can be explained with quantum mechanics. Both helium I and II are quantum fluids, displaying atomic properties on a macroscopic scale due to the fact that the boiling point of helium is so close to absolute zero.
Helium II
At 2.174 K, helium I forms into helium II. Its properties are very unusual, and the substance is described as superfluid. Superfluid is a quantum-mechanical state of matter; the two-fluid model for helium II explains why one portion of helium atoms exists in a ground state, flowing with zero viscosity, and another portion is in an excited state, behaving like an ordinary fluid. The viscosity of He4II is so low that there is no internal friction.
He4II can conduct heat 300 times more effectively than silver, making it the best heat conductor known. Its thermal conductivity is a million times that of helium I and several hundred times that of copper. The conductivity and viscosity of helium II do not obey classical rules, but are consistent with the rules of quantum mechanics. When temperature is lowered, helium II expands in volume. It cannot be boiled, but evaporates directly to gas when heated.
In this superfluid state, liquid helium can flow through thin capillaries or cracks much faster than helium gas. It also exhibits a creeping effect, moving along the surface seemingly against gravity. Helium II creeps along the sides of a open vessel until it reaches a warmer region where it evaporates. As a result of the creeping behavior and the ability to leak rapidly through tiny openings, helium II is very difficult to confine. Helium II also exhibits a fountain effect. Suppose a chamber allows a reservoir of helium II to filter superfluid and non-superfluid helium. When the interior of the container is heated, superfluid helium converts to non-superfluid helium to maintain equilibrium. This creates intense pressure on the superfluid helium, causing the liquid to fountain out of the container.
Isotopes
Helium has eight known isotopes but only two are stable: 3He and 4He. 3He is found in only very small quantities compared to 4He. It is produced in trace amounts by the beta decay of tritium. This form is found in abundance in stars, as a product of nuclear fusion. Extraplanetary materials have trace amounts of 3He from solar winds. 4He is produced by the alpha decay of heavier radioactive elements on Earth. It is an unusually stable isotope because its nucleons are arranged in complete shells.
Problems
1. What happens when a lit cigarette is thrown at a leaking, high-pressured helium cylinder?
a). nothing
b). the cigarette is incinerated before touching the cylinder
c). the cylinder explodes
d). the cylinder becomes a flame thrower
1. How many isotopes of helium are known? ______
2. (Helium gas, or helium II liquid) leaks faster than the other when stored in a opened cylinder (at STP).
3. What happens when a section of divided petri dish is filled with helium II at 2.173K?
a). It starts boiling.
b). It starts "creeping" over the divider, soon filling up the other sections of the dish.
c). It evaporates and soon leaves the dish.
d.) It solidifies and expands, breaking the dividers of the petri dish, and filling up the whole dish.
1. Helium was first discovered through ________.
Contributors and Attributions
• Jun-Hyun Hwang - University of California, Davis | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_18%3A_The_Noble_Gases/2Group_18%3A_Reactions_of_Nobel_Gases.txt |
Contributors and Attributions
• Jonathan Molina - (UCD)
Z018 Chemistry of Argon (Z18)
Argon is it is colorless, tasteless and odorless noble gas that is located in Group 18 on the Periodic Table. It was discovered by Henry Cavendish in 1785 and was named Argon, which is derived from the Greek word "argos" meaning inactive. Cavendish formed oxides of nitrogen by passing electric currents through air, then dissolved them in water to get nitric acid, but was unable to get all of the air to react. He suspected that there was a then unidentified gas component of air; Ramsay and Rayleigh went on to isolate this component in 1894, and the new found element was thus named Argon.
Contributors and Attributions
• Katherine Cubbon (UCD)
Stephen R. Marsden
Z036 Chemistry of Kryton
Krypton is one of the six Noble Gas elements (Group 18), which are widely known for their relative "inertness" and difficulty in forming chemical compounds with any other elements, due to these elements having full valence shells. Contrary to original thinking, however, Krypton has been made to react with the highly electronegative elements and is used in lighting and other commercial purposes.
Facts
• Element number: 36
• Electron configuration: [Ar]3d104s24p6
• Atomic weight: 83.798g/mol
• Color: colorless, odorless, tasteless
• Light: large number of spectral lines, strongest being green and yellow/ whitish emission
• Solidified: white and crystalline/ face-centered cubic crystal structure
• Melting point: 115.79K
• Boiling point: 119.92K
• Critical point: 209.41K
• Specific heat capacity: 20.786 J/mol K
• 0.000108-0.000114% of atmosphere
• 6 Stable isotopes
• Produced by breakdown of uranium and plutonium in the earth's crust at a very small %
The Origin and History
Krypton is found in the Group 18 elements, otherwise known as the Noble Gases. In 1785, Henry Cavendish suggested that air contained nonreactive gases after he was unsuccessful in getting a sample of air to react. A century later, British chemists John Rayleigh and William Ramsey began to isolate these inert gases (beginning with Argon) and seperated them in their own group on the periodic table since each of these elements had full electron valence shells. One of these gases, Krypton, was discovered along with Neon and Xenon by Rayleigh and fellow chemist Morris Travers in 1898 in a residue left from evaporating almost all components of liquid air. The name Krypton is derived from the Greek word "kryptos", meaning "hidden". However, the inert quality of these gases was disproved when Xenon compounds were created in 1962 and a Krypton compound (KrF2) was synthesized successfully a year later. This proved that this group of gases is not necessarily inert. Although both Kr and Xe have full valence shells, they are both the most easily ionized of the group. It simply took an element of high electronegativity, in this case Fluorine, to force Xe and Kr to react under high temperatures.
Isolation
It ranks sixth in abundance in the atmosphere. As with the other noble gases, krypton is isolated from the air by liquefaction.
Compounds & Isotopes
Although Krypton is naturally chemically nonreactive, krypton difluoride was synthesized in 1963.
$Kr_{(g)}+F_{2(g)} \rightarrow KrF_{2(g)} \nonumber$
It has also been discovered that Krypton can bond with other atoms besides Fluorine, however such compounds are much more unstable than krypton difluoride. For example, KrF2 can bond with nitrogen when it reacts with [HC≡NH]+[AsF6] under -50°C to form [HC≡N–Kr–F]+. There have been other reports of successfully synthesizing additional Krypton compounds, but none have been verified. Krypton has 6 stable isotopes: 78Kr, 80Kr, 82Kr, 83Kr, 84Kr, and 86Kr. There are a total of 31 isotopes of Krypton, and the only isotope besides the six given that occur naturally is 81Kr which is a product of atmospheric reactions between the other natural isotopes.
Applications
Krypton gas is used in various kinds of lights, from small bright flashlight bulbs to special strobe lights for airport runways. Due to Krypton's large number of spectral lines, it's ionized gas is white, which is why light bulbs that are krypton based are used in photography and studio lighting in the film industry. In neon lights, Krypton reacts with other gases to produce a bright yellow light as well. The isotope 85Kr can also be used in combination with phosphors to produce materials that shine in the dark due to the fact that this particular isotope of Krypton reflects off of phosphors. Krypton is also used in lasers as a control for a desired wavelength, especially in red lasers because Krypton has a much higher light density in the red spectral region than other gases such as Neon, which is why krypton-based lasers are used to produce red light in laser-light shows.
Perhaps one of the most significant uses of Krypton is in the krypton-fluoride laser which is used in nuclear fusion energy research.
other.nrl.navy.mil/LaserFusio...ercreation.htm
Isotopes of Krypton
Yet another important application of Krypton, specifically 83Kr, is in Magnetic Resonance Imaging (MRI), which is used instead of other gases because of its high spin and smaller/less polar electron cloud compared to other noble gases such as Xenon. It is used to distinguish hydrophobic and hydrophillic regions containing an airway.
An international agreement was made in 1960 to base the length of the meter on the wavelength of light emitted by 86Kr (605.78 nm). However, this was changed in 1983 when the International Bureau of Weights and Measures determined the meter to be the distance that light travels in a vacuum in 1/299,792,458 s.
Problems
1. How is the element Krypton isolated?
2. Why can fluorine react with Krypton to form a compound?
3. What is the electron configuration of Krypton?
4. What colors are most pronounced in the Krypton spectral emission?
5. How many stable isotopes of Krypton are there?
Solutions
1. By evaporating the componnts of liquid air.
2. Because fluorine is highly electronegative and Krypton is readily ionizable.
3. [Ar]3d104s24p6
4. Green and yellow
5. 6
Contributors and Attributions
• Megan Meadows (UCD)
Z086 Chemistry of Radon (Z86)
Contributors and Attributions
• Boundless
• Stephen R. Marsden
Z54 Chemistry of Xenon (Z54)
Xenon is an element under the Noble gases group and is on period 7 of the periodic table. This element is most notable for its bright luminescence in light bulbs. Xenon is unique for being the first noble gas element to be synthesized into a compound.
Contributors and Attributions
• Albert Young (UCD) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Block/2_p-Block_Elements/Group_18%3A_The_Noble_Gases/Z010_Chemistry_of_Neon_%28Z10%29.txt |
Alkali metals are the chemical elements found in Group 1 of the periodic table. The alkali metals include: lithium, sodium, potassium, rubidium, cesium, and francium. Although often listed in Group 1 due to its electronic configuration, hydrogen is not technically an alkali metal since it rarely exhibits similar behavior. The word "alkali" received its name from the Arabic word "al qali," meaning "from ashes", which since these elements react with water to form hydroxide ions, creating alkaline solutions (pH>7).
• Group 1: Properties of Alkali Metals
This page discusses the trends in some atomic and physical properties of the Group 1 elements - lithium, sodium, potassium, rubidium and cesium. Sections below cover the trends in atomic radius, first ionization energy, electronegativity, melting and boiling points, and density.
• Group 1: Reactivity of Alkali Metals
Alkali metals are among the most reactive metals. This is due in part to their larger atomic radii and low ionization energies. They tend to donate their electrons in reactions and have an oxidation state of +1. These metals are characterized by their soft texture and silvery color. They also have low boiling and melting points and are less dense than most elements. All these characteristics can be attributed to these elements' large atomic radii and weak metallic bonding.
• Chemistry of Hydrogen (Z=1)
Hydrogen is one of the most important elements in the world. It is all around us. It is a component of water (H2O), fats, petroleum, table sugar (C6H12O6), ammonia (NH3), and hydrogen peroxide (H2O2)—things essential to life, as we know it. This module will explore several aspects of the element and how they apply to the world.
• Chemistry of Lithium (Z=3)
Chlorine is a halogen in Lithium is a rare element found primarily in molten rock and saltwater in very small amounts. It is understood to be non-vital in human biological processes, although it is used in many drug treatments due to its positive effects on the human brain. Because of its reactive properties, humans have utilized lithium in batteries, nuclear fusion reactions, and thermonuclear weapons.
• Chemistry of Sodium (Z=11)
Sodium is metallic element found in the first group of the periodic table. As the sixth most abundant element in the Earth's crust, sodium compounds are commonly found dissolved in the oceans, in minerals, and even in our bodies.
• Chemistry of Potassium (Z=19)
In its pure form, potassium has a white-sliver color but it quickly oxidizes upon exposure to air, tarnishing in minutes if it is not stored under oil or grease. Potassium is essential to several aspects of plant, animal, and human life and is thus mined, manufactured, and consumed in huge quantities around the world.
• Chemistry of Rubidium (Z=37)
Rubidium (Latin: rubidius = red) is similar in physical and chemical characteristics to potassium, but much more reactive. It is the seventeenth most abundant element and was discovered by its red spectral emission in 1861 by Bunsen and Kirchhoff. Its melting point is so low you could melt it in your hand if you had a fever (39°C). But that would not be a good idea because it would react violently with the moisture in your skin.
• Chemistry of Cesium (Z=55)
Cesium is so reactive that it will even explode on contact with ice! It has been used as a "getter" in the manufacture of vacuum tubes (i.e., it helps remove trace quantities of remaining gases). An isotope of cesium is used in the atomic clocks.
• Chemistry of Francium (Z=87)
Francium is the last of the known alkali metals and does not occur to any significant extent in nature. All known isotopes are radioactive and have short half-lives (22 minutes is the longest).
Group 01: Hydrogen and the Alkali Metals
This page discusses the trends in some atomic and physical properties of the Group 1 elements - lithium, sodium, potassium, rubidium and cesium. Sections below cover the trends in atomic radius, first ionization energy, electronegativity, melting and boiling points, and density.
The chart below shows the increase in atomic radius down the group.
The radius of an atom is governed by two factors:
1. The number of layers of electrons around the nucleus
2. The attraction the outer electrons feel from the nucleus
Compare the electronic configurations of lithium and sodium:
• Li: 1s22s1
• Na: 1s22s22p63s1
In each element, the outer electron experiences a net charge of +1 from the nucleus. The positive charge on the nucleus is canceled out by the negative charges of the inner electrons. This effect is illustrated in the figure below:
This is true for each of the other atoms in Group 1. The only factor affecting the size of the atom is the number of layers of inner electrons which surround the atom. More layers of electrons take up more space, due to electron-electron repulsion. Therefore, the atoms increase in size down the group.
The first ionization energy of an atom is defined as the energy required to remove the most loosely held electron from each of one mole of gaseous atoms, producing one mole of singly charged gaseous ions; in other words, it is the energy required for 1 mole of this process:
$X(g) \rightarrow X^+ (g) + e^- \nonumber$
A graph showing the first ionization energies of the Group 1 atoms is shown above. Notice that first ionization energy decreases down the group. Ionization energy is governed by three factors:
• the charge on the nucleus,
• the amount of screening by the inner electrons,
• the distance between the outer electrons and the nucleus.
Down the group, the increase in nuclear charge is exactly offset by the increase in the number of inner electrons. As mentioned before, in each of the elements Group 1, the outermost electrons experience a net charge of +1 from the center. However, the distance between the nucleus and the outer electrons increases down the group; electrons become easier to remove, and the ionization energy falls.
Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. It is usually measured on the Pauling scale, on which the most electronegative element (fluorine) is given an electronegativity of 4.0 (Table A2).
A graph showing the electronegativities of the Group 1 elements is shown above. Each of these elements has a very low electronegativity when compared with fluorine, and the electronegativities decrease from lithium to cesium.
Picture a bond between a sodium atom and a chlorine atom. The bond can be considered covalent, composed of a pair of shared electrons. The electron pair will be pulled toward the chlorine atom because the chlorine nucleus contains many more protons than the sodium nucleus. This is illustrated in the figure below:
The electron pair is so close to the chlorine that an effective electron transfer from the sodium atom to the chlorine atom occurs—the atoms are ionized. This strong attraction from the chlorine nucleus explains why chlorine is much more electronegative than sodium.
Now compare this with a lithium-chlorine bond. The net pull from each end of the bond is the same as before, but the lithium atom is smaller than the sodium atom. That means that the electron pair is going to be more strongly attracted to the net +1 charge on the lithium end, and thus closer to it.
In some lithium compounds there is often a degree of covalent bonding that is not present in the rest of the group. Lithium iodide, for example, will dissolve in organic solvents; this is a typical property of covalent compounds. The iodine atom is so large that the pull from the iodine nucleus on the pair of electrons is relatively weak, and a fully-ionic bond is not formed.
Summarizing the trend down the group
As the metal atoms increase in size, any bonding electron pair becomes farther from the metal nucleus, and so is less strongly attracted towards it. This corresponds with a decrease in electronegativity down Group 1. With the exception of some lithium compounds, the Group 1 elements each form compounds that can be considered ionic. Each is so weakly electronegative that in a Group 1-halogen bond, we assume that the electron pair on a more electronegative atom is pulled so close to that atom that ions are formed.
The figure above shows melting and boiling points of the Group 1 elements. Both the melting and boiling points decrease down the group.
When any of the Group 1 metals is melted, the metallic bond is weakened enough for the atoms to move more freely, and is broken completely when the boiling point is reached. The decrease in melting and boiling points reflects the decrease in the strength of each metallic bond.
The atoms in a metal are held together by the attraction of the nuclei to electrons which are delocalized over the whole metal mass. As the atoms increase in size, the distance between the nuclei and these delocalized electrons increases; therefore, attractions fall. The atoms are more easily pulled apart to form a liquid, and then a gas. As previously discussed, each atom exhibits a net pull from the nuclei of +1. The increased charge on the nucleus down the group is offset by additional levels of screening electrons. As before, the trend is determined by the distance between the nucleus and the bonding electrons.
The densities of the Group 1 elements increase down the group (except for a downward fluctuation at potassium). This trend is shown in the figure below:
The metals in this series are relatively light—lithium, sodium, and potassium are less dense than water (less than 1 g cm-3). It is difficult to develop a simple explanation for this trend because density depends on two factors, both of which change down the group. The atoms are packed in the same way, so the two factors considered are how many atoms can be packed in a given volume, and the mass of the individual atoms. The amount packed depends on the individual atoms' volumes; these volumes, in turn, depends on their atomic radius.
Atomic radius increases down a group, so the volume of the atoms also increases. Fewer sodium atoms than lithium atoms, therefore, can be packed into a given volume. However, as the atoms become larger, their masses increase. A given number of sodium atoms will weigh more than the same number of lithium atoms. Therefore, 1 cm3 of sodium contains fewer atoms than the same volume of lithium, but each atom weighs more. Mathematical calculations are required to determine the densities.
Contributors and Attributions
Jim Clark (Chemguide.co.uk)
2Group 1: Reactivity of Alkali Metals
Alkali metals are among the most reactive metals. This is due in part to their larger atomic radii and low ionization energies. They tend to donate their electrons in reactions and have an oxidation state of +1. These metals are characterized by their soft texture and silvery color. They also have low boiling and melting points and are less dense than most elements. Lithium, sodium, and potassium float on water because of their low density. All these characteristics can be attributed to these elements' large atomic radii and weak metallic bonding. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_01%3A_Hydrogen_and_the_Alkali_Metals/1Group_1%3A_Properties_of_Alkali_Metals.txt |
Hydrogen is a colorless, odorless and tasteless gas that is the most abundant element in the known universe. It is also the lightest (in terms of atomic mass) and the simplest, having only one proton and one electron (and no neutrons in its most common isotope). It is all around us. It is a component of water (H2O), fats, petroleum, table sugar (C6H12O6), ammonia (NH3), and hydrogen peroxide (H2O2)—things essential to life, as we know it.
Hydrogen Facts
• Atomic Number: 1
• Atomic Symbol: H
• Atomic Weight: 1.0079
• Electronic Configuration: 1s1
• Oxidation States: 1, -1
• Atomic Radius: 78 pm
• Melting Point: -259.34°C
• Boiling Point: -252.87° C
• Elemental Classification: Non-Metal
• At Room Temperature: Colorless & Odorless Diatomic Gas
History of Hydrogen
Hydrogen comes from Greek meaning “water producer” (“hydro” =water and “gennao”=to make). First isolated and identified as an element by Cavendish in 1766, hydrogen was believed to be many different things. Cavendish himself thought that it was "inflammable air from metals", owing to its production by the action of acids on metals. Before that, Robert Boyle and Paracelsus both used reactions of iron and acids to produce hydrogen gas and Antoine Lavoisier gave hydrogen its name because it produced water when ignited in air. Others thought it was pure phlogiston because of its flammability. Hydrogen is among the ten most abundant elements on the planet, but very little is found in elemental form due to its low density and reactivity. Much of the terrestrial hydrogen is locked up in water molecules and organic compounds like hydrocarbons.
Properties of Hydrogen
Hydrogen is a nonmetal and is placed above group in the periodic table because it has ns1 electron configuration like the alkali metals. However, it varies greatly from the alkali metals as it forms cations (H+) more reluctantly than the other alkali metals. Hydrogen‘s ionization energy is 1312 kJ/mol, while lithium (the alkali metal with the highest ionization energy) has an ionization energy of 520 kJ/mol.
Because hydrogen is a nonmetal and forms H- (hydride anions), it is sometimes placed above the halogens in the periodic table. Hydrogen also forms H2 dihydrogen like halogens. However, hydrogen is very different from the halogens. Hydrogen has a much smaller electron affinity than the halogens.
H2 dihydrogen or molecular hydrogen is non-polar with two electrons. There are weak attractive forces between H2 molecules, resulting in low boiling and melting points. However, H2 has very strong intramolecular forces; H2 reactions are generally slow at room temperature due to strong H—H bond. H2 is easily activated by heat, irradiation, or catalysis. Activated hydrogen gas reacts very quickly and exothermically with many substances.
Hydrogen also has an ability to form covalent bonds with a large variety of substances. Because it makes strong O—H bonds, it is a good reducing agent for metal oxides. Example: CuO(s) + H2(g) → Cu(s) + H2O(g) H2(g) passes over CuO(s) to reduce the Cu2+ to Cu(s), while getting oxidized itself.
Reactions of Hydrogen
Hydrogen's low ionization energy makes it act like an alkali metal:
$H_{(g)} \rightarrow H^+_{(g)} + e^- \nonumber$
However, it half-filled valence shell (with a $1s^1$ configuration) with one $e^-$ also causes hydrogen to act like a halogen non-metal to gain noble gas configuration by adding an additional electron
$H_{(g)} + e^- \rightarrow H^-_{(g)} \nonumber$
Reactions of Hydrogen with Active Metals
Hydrogen accepts e- from an active metal to form ionic hydrides like LiH. By forming an ion with -1 charge, the hydrogen behaves like a halogen.
Group 1 metals
$2M_{(s)}+H_{2(g)} \rightarrow 2MH_{(s)} \nonumber$
with $M$ representing Group 1 Alkali metals
Examples:
• $2K_{(s)}+H_{2(g)} \rightarrow 2KH_{(s)}$
• $2K_{(s)}+Cl_{2(g)} \rightarrow 2KCl_{(s)}$
Group 2 metals
$M_{(s)}+H_{2(g)} \rightarrow MH_{2(s)} \nonumber$
with $M$ representing Group 2 Alkaline Earth metals
Example:
• $Ca_{(s)}+H_{2(g)} \rightarrow CaH_{2(s)}$
• $Ca_{(s)}+Cl_{2(g)} \rightarrow CaCl_{2(s)}$
Reactions of Hydrogen with Nonmetals
Unlike metals forming ionic bonds with nonmetals, hydrogen forms polar covalent bonds. Despite being electropositive like the active metals that form ionic bonds with nonmetals, hydrogen is much less electropositive than the active metals, and forms covalent bonds.
Hydrogen + Halogen → Hydrogen Halide
$H_{2(g)}+ Cl_{2(g)} \rightarrow HCl_{(g)} \nonumber$
Hydrogen gas reacting with oxygen to produce water and a large amount of heat: Hydrogen + Oxygen → Water
$(H_{2(g)}+O_{2(g)} \rightarrow H_2O_{(g)} \nonumber$
Reactions with Transition Metals
Reactions of hydrogen with Transition metals (Group 3-12) form metallic hydrides. There is no fixed ratio of hydrogen atom to metal because the hydrogen atoms fill holes between metal atoms in the crystalline structure.
Uses & Application
The vast majority of hydrogen produced industrially today is made either from treatment of methane gas with steam or in the production of "water gas" from the reaction of coal with steam. Most of this hydrogen is used in the Haber process to manufacture ammonia.
Hydrogen is also used for hydrogenation vegetable oils, turning them into margarine and shortening, and some is used for liquid rocket fuel. Liquid hydrogen (combined with liquid oxygen) is a major component of rocket fuel (as mentioned above combination of hydrogen and oxygen relapses a huge amount of energy). Because hydrogen is a good reducing agent, it is used to produce metals like iron, copper, nickel, and cobalt from their ores.
Because one cubic feet of hydrogen can lift about 0.07 lbs, hydrogen lifted airships or Zeppelins became very common in the early 1900s.However, the use of hydrogen for this purpose was largely discontinued around World War II after the explosion of The Hindenburg; this prompted greater use of inert helium, rather than flammable hydrogen for air travel.
Video Showing the explosion of The Hindenburg. (Video from Youtube)
Recently, due to the fear of fossil fuels running out, extensive research is being done on hydrogen as a source of energy.Because of their moderately high energy densities liquid hydrogen and compressed hydrogen gas are possible fuels for the future.A huge advantage in using them is that their combustion only produces water (it burns “clean”). However, it is very costly, and not economically feasible with current technology.
Combustion of fuel produces energy that can be converted into electrical energy when energy in the steam turns a turbine to drive a generator. However, this is not very efficient because a great deal of energy is lost as heat. The production of electricity using voltaic cell can yield more electricity (a form of usable energy). Voltaic cells that transform chemical energy in fuels (like H2 and CH4) are called fuel cells. These are not self-contained and so are not considered batteries. The hydrogen cell is a type of fuel cell involving the reaction between H2(g) with O2(g) to form liquid water; this cell is twice as efficient as the best internal combustion engine. In the cell (in basic conditions), the oxygen is reduced at the cathode, while the hydrogen is oxidized at the anode.
Reduction: O2(g)+2H2O(l)+4e- → 4OH-(aq)
Oxidation: H2(g) + 2OH-(aq) → 2H2O(l) + 2e-
Overall: 2H2(g) + O2(g) → 2H2O(l)
E°cell= Reduction- Oxidation= E°O2/OH- - E°H2O/H2 = 0.401V – (-0.828V) = +1.23
However, this technology is far from being used in everyday life due to its great costs.
Image of A Hydrogen Fuel Cell. (Image made by Ridhi Sachdev)
Natural Occurrence & Other Sources
Naturally Occurring Hydrogen
Hydrogen is the fuel for reactions of the Sun and other stars (fusion reactions). Hydrogen is the lightest and most abundant element in the universe. About 70%- 75% of the universe is composed of hydrogen by mass. All stars are essentially large masses of hydrogen gas that produce enormous amounts of energy through the fusion of hydrogen atoms at their dense cores. In smaller stars, hydrogen atoms collided and fused to form helium and other light elements like nitrogen and carbon(essential for life). In the larger stars, fusion produces the lighter and heavier elements like calcium, oxygen, and silicon.
On Earth, hydrogen is mostly found in association with oxygen; its most abundant form being water (H2O). Hydrogen is only .9% by mass and 15% by volume abundant on the earth, despite water covering about 70% of the planet. Because hydrogen is so light, there is only 0.5 ppm (parts per million) in the atmosphere, which is a good thing considering it is EXTREMELY flammable.
Other Sources of Hydrogen
Hydrogen gas can be prepared by reacting a dilute strong acid like hydrochloric acids with an active metal. The metal becomes oxides, while the H+ (from the acid) is reduced to hydrogen gas. This method is only practical for producing small amounts of hydrogen in the lab, but is much too costly for industrial production:
$Zn_{(s)} + 2H^+_{(aq)} \rightarrow Zn^{2+}_{(aq)} + H_{2(g)} \nonumber$
The purest form of H2(g) can come from electrolysis of H2O(l), the most common hydrogen compound on this plant. This method is also not commercially viable because it requires a significant amount of energy ($\Delta H = 572 \;kJ$):
$2H_2O_{(l)} \rightarrow 2H_{2(g)} + O_{2(g)} \nonumber$
$H_2O$ is the most abundant form of hydrogen on the planet, so it seems logical to try to extract hydrogen from water without electrolysis of water. To do so, we must reduce the hydrogen with +1 oxidation state to hydrogen with 0 oxidation state (in hydrogen gas). Three commonly used reducing agents are carbon (in coke or coal), carbon monoxide, and methane. These react with water vapor form H2(g):
$C_{(s)} + 2H_2O_{(g)} \rightarrow CO(g) + H_{2(g)} \nonumber$
$CO_{(g)} + 2H_2O_{(g)} \rightarrow CO2 + H_{2(g)} \nonumber$
Reforming of Methane:
$CH_{4(g)} + H_2O_{(g)} \rightarrow CO(g) + 3H_{2(g)} \nonumber$
These three methods are most industrially feasible (cost effective) methods of producing H2(g).
Isotopes
There are two important isotopes of hydrogen. Deuterium (2H) has an abundance of 0.015% of terrestrial hydrogen and the nucleus of the isotope contains one neutron.
• Protium (1H) is the most common isotope, consisting of 99.98% of naturally occurring hydrogen. It is a nucleus containing a single proton.
• Deuterium (2H) is another an isotope containing a proton and neutron, consisting of only 0.0156% of the naturally occurring hydrogen. Commonly indicated with symbol D and sometimes called heavy hydrogen, deuterium is separated by the fractional distillation of liquid hydrogen but it can also be produced by the prolonged electrolysis of ordinary water. Approximately 100,000 gallons of water will produce a single gallon of D2O, "heavy water". This special kind of water has a higher density, melting point, and boiling point than regular water and used as a moderator in some fission power reactors. Deuterium fuel is used in experimental fusion reactors. Replacing protium with deuterium has important uses for exploring reaction mechanisms via the kinetic isotope effect.
• Tritium (3H) contains two neutrons in its nucleus and is radioactive with a 12.3-year half-life, which is continuously formed in the upper atmosphere due to cosmic rays. It is can also be made in a lab from Lithium-6 in a nuclear reactor. Tritium is also used in hydrogen bombs. It is very rare (about 1 in every 1,018 atoms) and is formed in the environment by cosmic ray bombardment. Most tritium is manufactured by bombarding Li with neutrons. Tritium is used in thermonuclear weapons and experimental fusion reactors.
Problems
1. Write the reaction of Na(s) with H2(g).
2. What is the name of the radioactive isotope of hydrogen?
3. What characteristics of alkali metals does hydrogen display?
4. What characteristics of halogens does hydrogen display?
5. How does the electronegativity of hydrogen compare to that of the halogens?
6. What is the electron configuration of a neutral hydrogen atom.
Answers
1. 2Na(s) + H2(g)→ 2NaH(s)
2. Tritium
3. Hydrogen is placed above group in the periodic table because it has ns1 electron configuration like the alkali metals. However, it varies greatly from the alkali metals as it forms cations (H+) more reluctantly than the other alkali metals. Hydrogen‘s ionization energy is 1312 kJ/mol, while lithium (the alkali metal with the highest ionization energy) has an ionization energy of 520 kJ/mol.
4. Because hydrogen is a nonmetal and forms H- (hydride anions), it is sometimes placed above the halogens in the periodic table. Hydrogen also forms H2 dihydrogen like halogens. However, hydrogen is very different from the halogens. Hydrogen has a much smaller electron affinity than the halogens.
5. Hydrogen is less electronegative than the halogens.
6. 1s1
Contributors and Attributions
• Ridhi Sachdev (UC Davis)
Stephen R. Marsden | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_01%3A_Hydrogen_and_the_Alkali_Metals/Z001_Chemistry_of_Hydrogen_%28Z1%29.txt |
Lithium is a rare element found primarily in molten rock and saltwater in very small amounts. It is understood to be non-vital in human biological processes, although it is used in many drug treatments due to its positive effects on the human brain. Because of its reactive properties, humans have utilized lithium in batteries, nuclear fusion reactions, and thermonuclear weapons.
Introduction
Lithium was first identified as a component of of the mineral petalite and was discovered in 1817 by Johan August Arfwedson, but not isolated until some time later by W.T. Brande and Sir Humphry Davy. In its mineral forms it accounts for only 0.0007% of the earth's crust. It compounds are used in certain kinds of glass and porcelain products. More recently lithium has become important in dry-cell batteries and nuclear reactors. Some compounds of lithium have been used to treat manic depressives.
Lithium is an alkali metal with the atomic number = 3 and an atomic mass of 6.941 g/mol. This means that lithium has 3 protons, 3 electrons and 4 neutrons (6.941 - 3 = ~4). Being an alkali metal, lithium is a soft, flammable, and highly reactive metal that tends to form hydroxides. It also has a pretty low density and under standard conditions, it is the least dense solid element.
Properties
Lithium is the lightest of all metals and is named from the Greek work for stone (lithos). It is the first member of the Alkali Metal family. It is less dense than water (with which it reacts) and forms a black oxide in contact with air.
Table 1. Properties of lithium.
Atomic Number 3
Atomic Mass 6.941 g/mol
Atomic Radius 152 pm
Density 0.534 g/cm3
Color light silver
Melting point 453.69 K
Boiling point 1615 K
Heat of fusion 3.00 kJ/mol
Heat of vaporization 147.1 kJ/mol
Specific heat capacity 24.860 kJ/mol
First ionization energy 520.2 kJ/mol
Oxidation states +1, -1
Electronegativity 0.98
Crystal structure body-centered cubic
Magnetism paramagnetic
2 stable isotopes 6Li (7.5%) and 7Li (92.5%)
Being on the upper left side of the Periodic Table, lithium has a fairly low electronegativity and electron affinity as compared to the rest of the elements. Also, lithium has high metallic character and subsequently lower nonmetallic character when compared with the other elements. Lithium has a higher atomic radius than most of the elements on the Periodic Table. In compounds lithium (like all the alkali metals) has a +1 charge. In its pure form it is soft and silvery white and has a relatively low melting point (181oC).
Reactivity
Lithium is part of the Group 1 Alkali Metals, which are highly reactive and are never found in their pure form in nature. This is due to their electron configuration, in that they have a single valence electron (Figure 1) which is very easily given up in order to create bonds and form compounds.
_↑ ↓_ _↑__
1s2 2s1
Reactions with Water
When placed in contact with water, pure lithium reacts to form lithium hydroxide and hydrogen gas.
$2Li (s) + 2H_2O (l) \rightarrow 2LiOH (aq) + H_2 (g) \nonumber$
Out of all the group 1 metals, lithium reacts the least violently, slowly releasing the hydrogen gas which may create a bright orange flame only if a substantial amount of lithium is used. This occurs because lithium has the highest activation energy of its group - that is, it takes more energy to remove lithium's one valence electron than with other group 1 elements, because lithium's electron is closer to its nucleus. Atoms with higher activation energies will react slower, although lithium will release more total heat through the entire process.
Reactions with Air
Pure lithium will form lithium hydroxide due to moisture in the air, as well as lithium nitride ($Li_3N$) from $N_2$ gas, and lithium carbonate $(Li_2CO_3$) from carbon dioxide. These compounds give the normally the silver-white metal a black tarnish. Additionally, it will combust with oxygen as a red flame to form lithium oxide.
$4Li (s) + O_2 (g) \rightarrow 2Li_2O \nonumber$
Applications
In its mineral forms it accounts for only 0.0007% of the earth's crust. It compounds are used in certain kinds of glass and porcelain products. More recently lithium has become important in dry-cell batteries and nuclear reactors. Some compounds of lithium have been used to treat manic depressives.
Batteries
Lithium is able to be used in the function of a Lithium battery in which the Lithium metal serves as the anode. Lithium ions serve in lithium ion batteries (chargeable) in which the lithium ions move from the negative to positive electrode when discharging, and vice versa when charging.
Heat Transfer
Lithium has the highest specific heat capacity of the solids, Lithium tends to be used as a cooler for heat transfer techniques and applications.
Sources and Extraction
Lithium is most commonly found combined with aluminum, silicon, and oxygen to form the minerals known as spodumene (LiAl(SiO3)2) or petalite/castorite (LiAlSi4O10). These have been found on each of the 6 inhabited continents, but they are mined primarily in Western Australia, China, and Chile. Mineral sources of lithium are becoming less essential, as methods have now been developed to make use of the lithium salts found in saltwater.
Extraction from minerals
The mineral forms of lithium are heated to a high enough temperature (1200 K - 1300 K) in order to crumble them and thus allow for subsequent reactions to more easily take place. After this process, one of three methods can be applied.
1. The use of sulfuric acid and sodium carbonate to allow the iron and aluminum to precipitate from the ore - from there, more sodium carbonate is applied to the remaining material allow the lithium to precipitate out, forming lithium carbonate. This is treated with hydrochloric acid to form lithium chloride.
2. The use of limestone to calcinate the ore, and then leaching with water, forming lithium hydroxide. Again, this is treated with hydrochloric acid to form lithium chloride.
3. The use of sulfuric acid, and then leaching with water, forming lithium sulfate monohydrate. This is treated with sodium carbonate to form lithium carbonate, and then hydrochloric acid to form lithium chloride.
The lithium chloride obtained from any of the three methods undergoes an oxidation-reduction reaction in an electrolytic cell, to separate the chloride ions from the lithium ions. The chloride ions are oxidized, and the lithium ions are reduced.
$2Cl^- - 2e^- \rightarrow Cl_2 \;\; \text{(oxidation)} \nonumber$
$Li^+ + e^- \rightarrow Li \;\; \text{(reduction)} \nonumber$
Extraction from Saltwater
Saltwater naturally contains lithium chloride, which must be extracted in the form of lithium carbonate, then it is re-treated, separated into its ions, and reduced in the same electrolytic process as in extraction from lithium ores. Only three saltwater lakes in the world are currently used for lithium extraction, in Nevada, Chile, and Argentina.
Saltwater is channeled into shallow ponds and over a period of a year or more, water evaporates out to leave behind various salts. Lime is used to remove the magnesium salt, so that the remaining solution contains a fairly concentrated amount of lithium chloride. The solution is then treated with sodium carbonate in order for usable lithium carbonate to precipitate out.
Problems
1. With which group of elements will lithium form compounds the most easily with?
2. What is the electron configuration of Li+?
3. What are some common uses of lithium?
4. For a lithium-ion battery containing LiCoO2, should the compound be placed in the anode or cathode?
5. Given that 7Li is 7.0160 amu and 6Li is 6.0151 amu, and their percent abundance is 92.58% and 7.42% respectively, what is the atomic mass of lithium?
Solutions
1. Group 17 Halogens (lithium forms strongly inic bonds with them, as halogens are highly electronegative and lithium has a free electron)
2. 1s2
3. Lithium-ion batteries, disposable lithium batteries, pyrotechnics, creation of strong metal alloys, etc.
4. Anode - lithium is oxidized (LiCoO2 → Li+ + CoO2)
5. 6.942 g/mol
Contributors and Attributions
• Katherine Szelong (UCD), Kevin Fan
Z011 Chemistry of Sodium (Z11)
Sodium is metallic element found in the first group of the periodic table. As the sixth most abundant element in the Earth's crust, sodium compounds are commonly found dissolved in the oceans, in minerals, and even in our bodies.
Contributors and Attributions
• Helen Min (University of California, Davis) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_01%3A_Hydrogen_and_the_Alkali_Metals/Z003_Chemistry_of_Lithium_%28Z3%29.txt |
Potassium is a group 1 metal, abbreviated as K on the periodic table. In its pure form, potassium has a white-sliver color, but quickly oxidizes upon exposure to air and tarnishing in minutes if it is not stored under oil or grease. Potassium is essential to several aspects of plant, animal, and human life and is thus mined, manufactured, and consumed in huge quantities around the world.
The seventh most abundant element, potassium was discovered and isolated in 1807 by Sir Humphry Davy. Important compounds of potassium include potassium hydroxide (used in some drain cleaners), potassium superoxide, $KO_2$, which is used in respiratory equipment and potassium nitrate, used in fertilizers and pyrotechnics. Potassium, like sodium, melts below the boiling point of water (63 °C) and is less dense than water also. Like most of the alkali metals, potassium compounds impart a characteristic color to flames. In the case of the 19th element, the color is pale lavender. Like sodium ions, the presence of potassium ions in the body is essential for the correct function of many cells.
Table $1$: Basic Chemical and Physical Properties
Atomic Number 19
Atomic Mass 39.098 g/mol
Electronegativity 0.8
Density 0.862 g/cm3 (at 0o C) (floats on water)
Melting Point 63.65o C
Boiling Point 773.9o C
Atomic Radius 227 pm
Ionic Radius 0.133 (+1)
Isotopes 5
Electronic Shell [Ar] 4s1
1st Ionization Energy 418.8 kJ/mol
Electrode Potential -2.924
Hardness 0.5
Crystal Lattice body-centered cubic
Specific Heat 0.741 J/gK
Heat of Fusion 59.591 J/g
Heat of Vaporization 2075 kJ/g
Electron Configuration 1s22s22p63s23p64s1
Notable Reactions with Phosphorus
Potassium reacts so violently with water that it bursts into flame. The silvery white metal is very soft and reacts rapidly with the oxygen in air. Its chemical symbol is derived from the Latin word kalium which means "alkali". Its English name is from potash which is the common name for a compound containing it.
Table $1$: Key reactions of Potassium
Reactant Reaction Product
H2 begins slowly at ca 200°C; rapid above 300°C KH
O2 begins slowly with solid; fairly rapid with liquid K2O, K2O2, KO2
H2O extremely vigorous and frequently results in hydrogen–air explosions KOH, H2
C(graphite) 150–400°C KC4, KC8, KC24
CO forms unstable carbonyls (KCO)
NH3 dissolves as K; iron, nickel, and other metals catalyze in gas and liquid phase KNH2
S molten state in liquid ammonia K2S, K2S2, K2S4
F2, Cl2, Br2 violent to explosive KF, KCl, KBr
I2 ignition KI
CO2 occurs readily, but is sometimes explosive CO, C, K2CO3
Potassium in the Environment
Potassium has a 2.6% abundance by mass in the earth's crust and is found mostly in mineral form as part of feldspars (groups of minerals) and clays. Potassium easily leaches out of these minerals over time and thus has a relatively high concentration in sea water as well (0.75g/L). Today, most of the world's potassium is mined in Canada, the U.S., and Chile but was originally monopolized by Germany.
Potassium and Living Organisms
Plants, animals and humans all depend on potassium for survival and good health. The element is part of many bodily fluids and assists related functions of the human body. Most notably, potassium aids nerve functions and is found in several cell types (including skeletal cells, smooth muscle cells, endocrine cells, cardiac cells, and central neurons). Plants depend on potassium for healthy growth. Potassium found in animal excretions and dead plants easily binds to clay in the soil they fall on and is thus utilized by plants. The element helps maintain osmotic pressure and cell size and plays a role in photosynthesis and energy production.
Applications
95% of manufactured potassium is used in fertilizers and the rest is used to produce specific compounds of potassium, such as potassium hydroxide ($KOH$), which can then be turned into potassium carbonate ($K_2CO_3$). Potassium carbonate is used in glass manufacturing and potassium hydroxide is found in liquid soaps and detergents. Potassium chloride is used in many pharmaceuticals and other salts of potassium are used in baking, photography, tanning leather, and iodized salt. In these cases, potassium is utilized for its negative anion.
Potassium can be obtained through various known reactions, all of which require heat treatment:
$K_2CO_3+2C \overset{\Delta}{\longrightarrow} 3CO+2K \label{1}$
$2KCl+CaC_2 \overset{\Delta}{\longrightarrow} CaCl_2+2C+K \label{2}$
$2KN_3 \overset{\Delta}{\longrightarrow} 3N_2+2K \label{3}$
Due to expenses, these processes are not commercially adaptable. Therefore the element is commonly obtained through reduction at elevated heats (i.e., pyrometallurgy). Sodium is often combined with $KCl$, $KOH$, or $K_2CO_3$ to produce potassium sodium alloys and in the 1950's the Mine Safety Appliances Company developed a reduction process that yields high purity potassium:
$KCl+Na \overset{\Delta}{\longrightarrow} K+NaCl \label{4}$
The reaction is heated in a special device equipped with a furnace, heat-exchanger tubes, a fractionating column, a $KCl$ feed, a waste removal system, and a vapor condensing system. Because the reaction attains equilibrium quickly, potassium can be removed continuously as a product in order to shift equilibrium to the right and produce even more potassium in its place.
Alloys of potassium include $NaK$ (Sodium) and $KLi$ (Lithium). Both of these alloys produce metals of low vapor pressure and melting points.
Problems
1. Why is Potassium never found pure in nature?
2. Write out the chemical reaction between potassium and water.
3. Name 3 uses of potassium.
4. Where is Potassium on the periodic table. What are a few things you can deduce just from this location?
5. Name a common alloy of Potassium. What are the beneficial properties of this alloy?
Answers
1. It is too reactive. Potassium is a very strong reducing agent because of its desire to achieve an inert gas electron configuration (like the other alkali metals). This means that it will easily give up electrons, giving it the ability to reduce numerous other elements.
2. $K+H_2O \rightarrow KOH + H_{2(g)}$: Like other group 1 metals, potassium reacts readily with water to generate hydrogen gas.
3. Potassium is used in glass making and is found in fertilizers and soaps.
4. Potassium is in group one, and is the 4th element down in it's column. This tells us that it is an alkali metal. It is very reactive, has a low density, and is a good reducing agent.
5. Potassium can form an alloy with $Na$ that has a low vapor pressure and melting point.
Z037 Chemistry of Rubidium (Z37)
Rubidium (Latin: rubidius = red) is similar in physical and chemical characteristics to potassium, but much more reactive. It is the seventeenth most abundant element and was discovered by its red spectral emission in 1861 by Bunsen and Kirchhoff. Its melting point is so low you could melt it in your hand if you had a fever (39°C). But that would not be a good idea because it would react violently with the moisture in your skin.
Rubidium was once thought to be quite rare but recent discoveries of large deposits indicate that there is plenty to use. However at present it finds only limited application in the manufacture of cathode ray tubes.
Contributors and Attributions
Stephen R. Marsden
Z055 Chemistry of Cesium (Z55)
Cesium is a bright silvery metal which is a liquid in a warm room (28oC). Its name is from the Latin caesius which is a description of a sky blue spectral emission by which it was discovered in 1860 by Bunsen and Kirchhoff.
Cesium is so reactive that it will even explode on contact with ice! It has been used as a "getter" in the manufacture of vacuum tubes (i.e., it helps remove trace quantities of remaining gases). An isotope of cesium is used in the atomic clocks.
Contributors and Attributions
• Stephen R. Marsden
Z087 Chemistry of Francium (Z87)
Francium is the last of the known alkali metals and does not occur to any significant extent in nature. All known isotopes are radioactive and have short half-lives (22 minutes is the longest).
The existence of Francium was predicted by Dmitri Mendeleev in the 1870's and he presumed it would have chemical and physical properties similar to cesium. That may well be, but not enough francium has been isolated to test.
Numerous historical claims to the discovery of element 87 were made resulting in the names russium, virginium, and moldavium. However, the confirmed discovery is credited to Marguerite Perey who was an assistant to Marie Curie at the Radium Institute in Paris. She named the element after her native country.
Contributors and Attributions
Stephen R. Marsden | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_01%3A_Hydrogen_and_the_Alkali_Metals/Z019_Chemistry_of_Potassium_%28Z19%29.txt |
The Group 2 alkaline earth metals include Beryllium, Magnesium, Calcium, Barium, Strontium and Radium and are soft, silver metals that are less metallic in character than the Group 1 Alkali Metals. Although many characteristics are common throughout the group, the heavier metals such as Ca, Sr, Ba, and Ra are almost as reactive as the Group 1 Alkali Metals. All the elements in Group 2 have two electrons in their valence shells, giving them an oxidation state of +2.
• Group 2: Chemical Properties of Alkali Earth Metals
Covers the elements beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr) and barium (Ba). Includes trends in atomic and physical properties, trends in reactivity, the solubility patterns in the hydroxides and sulfates, trends in the thermal decomposition of the nitrates and carbonates, and some of the atypical properties of beryllium.
• Group 2: Physical Properties of Alkali Earth Metals
This page explores the trends in some atomic and physical properties of the Group 2 elements: beryllium, magnesium, calcium, strontium and barium. Sections below cover the trends in atomic radius, first ionization energy, electronegativity, and physical properties.
• Chemistry of Beryllium (Z=4)
The name Beryllium comes from the Greek beryllos which is the name for the gemstone beryl. The element is a high-melting, silver-white metal which is the first member of the alkaline earth metals. It is not abundant in the environment and occurs mainly in the mineral beryl with aluminum and silicon.
• Chemistry of Magnesium (Z=12)
Magnesium is a group two element and is the eighth most common element in the earth's crust. Magnesium is light, silvery-white, and tough. Like aluminum, it forms a thin layer around itself to help prevent itself from rusting when exposed to air. Fine particles of magnesium can also catch on fire when exposed to air.
• Chemistry of Calcium (Z=20)
Calcium is the 20th element in the periodic table. It is a group 2 metal, also known as an alkaline-earth metal, and no populated d-orbital electrons. Calcium is the fifth most abundant element by mass (3.4%) in both the Earth's crust and in seawater. All living organisms (in fact, even dead ones) have and need calcium for survival.
• Chemistry of Strontium (Z=38)
Strontium is a group 2 element that does not occur as a free element due to its extreme reactivity with oxygen and water. It occurs naturally only in compounds with other elements such as strontianite. It is softer than calcium and decomposes water more vigorously. It has a silver appearance but then turns yellow with the formation of oxide. Strontium is named after the Scottish village on Strontian.
• Chemistry of Barium (Z=56)
Barium is a soft, silvery white metal, and has a melting point of 1000 K. Because of its reaction to air, barium cannot be found in nature in its pure form but can be extracted from the mineral barite.
• Chemistry of Radium (Z=88)
Radium takes its name from the Latin word radius or ray. All isotopes of radium are radioactive and many exhibit luminescence, reacting readily with oxygen and water. The metal was discovered and isolated in 1911 by Marie Curie. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_02_Elements%3A_The_Alkaline_Earth_Metals.txt |
The observed trends in the properties of the group 3 elements are similar to those of groups 1 and 2. Due to their ns2(n − 1)d1 valence electron configurations, the chemistry of all four elements is dominated by the +3 oxidation state formed by losing all three valence electrons. As expected based on periodic trends, these elements are highly electropositive metals and powerful reductants, with La (and Ac) being the most reactive.
Group 04: Transition Metals
Because the elements of group 4 have a high affinity for oxygen, all three metals occur naturally as oxide ores that contain the metal in the +4 oxidation state resulting from losing all four ns2(n − 1)d2 valence electrons.
• Group 4 Elemental Properties
Because the elements of group 4 have a high affinity for oxygen, all three metals occur naturally as oxide ores that contain the metal in the +4 oxidation state resulting from losing all four valence electrons.
• Chemistry of Hafnium
More abundant than better known metals such as silver and gold, hafnium (from the Latin Hafnia, a name for Copenhagen) was not discovered until 1923 by Coster and de Hevesy. The reason is the similarity of hafnium to zirconium. Mendeleev had predicted the existence of element 72 but had wrongly suggested it might be found along with titanium ores. Instead it lay hidden with "pure" samples of zirconium. Later Niels Bohr predicted the arrangement of outer electrons for element 72.
• Chemistry of Rutherfordium
In 1964 researchers in the Soviet Union at Dubna announced their discovery of element 104. A similar claim was made by researchers at the University of California at Berkeley. The Soviet scientists claimed to have bombarded a target of Pu-242 with Ne-22, resulting in a nucleus with 104 protons and a mass number of 260. The Berkeley team used a Cf-249 target and isotopes of carbon for projectiles, resulting in isotopes of 104 with mass numbers of 257 and 259. Several other isotopes were also prep
• Chemistry of Titanium
Discovered independently by William Gregor and Martin Klaproth in 1795, titanium (named for the mythological Greek Titans) was first isolated in 1910. Gregor, a Cornish vicar and amateur chemist isolated an impure oxide from ilmenite. This page deals with the uses of titanium and its extraction from the ore, rutile.
• Chemistry of Zirconium
Named for the mineral zircon in which it can be found, zirconium was discovered in 1789 by Klaproth and eventually isolated in 1824 by Berzelius. The metal reacts with oxygen and nitrogen in the atmosphere to form a protective coating that inhibits further corrosion. It is resistant to weak acids and even forms a low-temperature superconductor when alloyed with niobium.
Group 05: Transition Metals
All group 5 metals are normally found in nature as oxide ores that contain the metals in their highest oxidation state (+5). Because of the lanthanide contraction, the chemistry of Nb and Ta is so similar that these elements are usually found in the same ores.
• Chemistry of Dubnium
Dubnium (Db) is a transactinide element and is highly radioactive. The most stable known isotope, dubnium-268, has a half-life of just above a day, which greatly limits the extent of possible research on dubnium. Dubnium does not occur naturally on Earth and is produced artificially.
• Chemistry of Niobium
Pure niobium looks much like steel but resists corrosion better due to a thin coating of oxide that forms on all exposed surfaces. The only acid that attacks Nb at room temperature is HF. Above 200° the metal becomes more reactive.
• Chemistry of Tantalum
Tantalum is a heavy, gray metal that resembles the more expensive platinum in many respects and is sometimes used as an economical substitute for that element. The metal comprises only 0.0002% of the earth's crust. Tantalum alloys are corrosion and wear resistant and find use in dental and surgical tools. Tantalum oxide is used in some electronic components and a composite of tantalum carbide (TaC) and graphite is one of the hardest materials ever produced.
• Chemistry of Vanadium
Vanadium (V) takes its name from the Scandinavian goddess Vanadis and was discovered in 1801 by Andrés Manuel del Rio. It was isolated in 1867 by Henry Roscoe as a silvery-white metal that is somewhat heavier than aluminum but lighter than iron. It has excellent corrosion resistance at room temperature.
• Group 5 Elemental Properties
Group 06: Transition Metals
The group 6 metals are slightly less electropositive than those of the three preceding groups, and the two heaviest metals are essentially the same size because of the lanthanide contraction.
Group 07: Transition Metals
All three group 7 elements have seven valence electrons and can form compounds in the +7 oxidation state. The chemistry of the group 7 metals (Mn, Tc, and Re) is dominated by lower oxidation states. Compounds in the maximum possible oxidation state (+7) are readily reduced.
Group 08: Transition Metals
The chemistry of group 8 is dominated by iron, whose high abundance in Earth’s crust is due to the extremely high stability of its nucleus. Ruthenium and osmium, on the other hand, are extremely rare elements, with terrestrial abundances of only about 0.1 ppb and 5 ppb, respectively, and they were not discovered until the 19th century. Because of the high melting point of iron (1538°C), early humans could not use it for tools or weapons. The advanced techniques needed to work iron were first developed by the Hittite civilization in Asia Minor sometime before 2000 BC, and they remained a closely guarded secret that gave the Hittites military supremacy for almost a millennium. With the collapse of the Hittite civilization around 1200 BC, the technology became widely distributed, however, leading to the Iron Age. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_03%3A_Transition_Metals.txt |
Cobalt is one of the least abundant of the first-row transition metals. Its oxide ores, however, have been used in glass and pottery for thousands of years to produce the brilliant color known as “cobalt blue,” and its compounds are consumed in large quantities in the paint and ceramics industries. The heavier elements of group 9 are also rare, with terrestrial abundances of less than 1 ppb; they are generally found in combination with the heavier elements of groups 8 and 10 in Ni–Cu–S ores.
• Chemistry of Cobalt
Cobalt (Co) lies with the transition metals on the periodic table. Cobalt was first discovered in 1735 by George Brandt in Stockholm Sweden. It is used in many places today, such as, magnets materials, paint pigments, glasses, and even cancer therapy. The word cobalt is from the German word kobold, which means "goblin" or "evil spirit" this term was used by miners that was really difficult to mine and harmful to the miners health.
• Chemistry of Iridium
Iridium has the reputation of being the most corrosion resistant of all metals. It was discovered in 1803 by Smithson Tennant in 1803 and named for the Latin iris, or "rainbow" because it forms a large number of very colorful compounds. The pure metal is very difficult to machine into useful shapes because of its hardness and its principal use is as a hardening agent for platinum.
• Chemistry of Meitnerium
Element 109 was first synthesized by researchers at Darmstadt, (West) Germany in August of 1983. For 10 days they hurled a beam of iron-58 ions into a bismuth-209 target. They detected the formation of one nucleus of Mt-267 which rapidly "boiled off" a neutron, reverting to Mt-266. This decayed within milliseconds to give (element 107)-262, etc. A committee of the IUPAC suggested the name Meitnerium (Mt) after the German physicist Lise Meitner. Final approval of the name and symbol was given in
• Chemistry of Rhodium
World production of rhodium (from the Greek rhodon, "rose") is about 10 tons. While the metal itself has few applications, it is an important alloying agent used as a hardener for platinum and palladium.
Group 10: Transition Metals
Group 10 metals are white to light grey in color, and possess a high luster, a resistance to tarnish (oxidation), are highly ductile, and enter into oxidation states of +2 and +4, with +1 being seen in special conditions.
Group 11: Transition Metals
The “coinage metals”, copper, silver, and gold, have held great importance in societies throughout history, both symbolically and practically. For centuries, silver and gold have been worn by royalty to parade their wealth and power. On occasion, these metals were even used in art. Although the most important oxidation state for group 11 is +1, the elements are relatively unreactive, with reactivity decreasing from Cu to Au.
Group 12: Transition Metals
Group 12 elements have partially filled (n − 1)d subshells, and hence are not, strictly speaking, transition metals. Nonetheless, much of their chemistry is similar to that of the elements that immediately precede them in the d block. The group 12 metals are similar in abundance to those of group 11, and they are almost always found in combination with sulfur. Group 12 metals tend have low melting and boiling points (due to the weak metallic bonding of the ns2 electrons) and charges of +2 or +1. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_09%3A_Transition_Metals.txt |
The boron family contains elements in group 13 of the periodic talbe and include the semi-metal boron (B) and the metals aluminum (Al), gallium (Ga), indium (In), and thallium (Tl). Aluminum, gallium, indium, and thallium have three electrons in their outermost shell (a full s orbital and one electron in the p orbital) with the valence electron configuration ns2np1. The elments of the boron family adopts oxidation states +3 or +1. The +3 oxidation states are favorable except for the heavier elements, such as Tl, which prefer the +1 oxidation state due to its stability; this is known as the inert pair effect. The elements generally follow periodic trends except for certain Tl deviations:
• Group 13: Chemical Reactivity
The boron family contains the semi-metal boron (B) and metals aluminum (Al), gallium (Ga), indium (In), and thallium (Tl).
• Group 13: Physical Properties of Group 13
The boron family contains the semi-metal boron (B) and metals aluminum (Al), gallium (Ga), indium (In), and thallium (Tl).
• Chemistry of Boron (Z=5)
Boron is the fifth element of the periodic table (Z=5), located in Group 13. It is classified as a metalloid due it its properties that reflect a combination of both metals and nonmetals.
• Chemistry of Aluminum (Z=13)
Aluminum (also called Aluminium) is the third most abundant element in the earth's crust. It is commonly used in the household as aluminum foil, in crafts such as dyeing and pottery, and also in construction to make alloys. In its purest form the metal is bluish-white and very ductile. It is an excellent conductor of heat and electricity and finds use in some wiring. When pure it is too soft for construction purposes but addition of small amounts of silicon and iron hardens it significantly.
• Chemistry of Gallium (Z=31)
Gallium is the chemical element with the atomic number 31 and symbol Ga on the periodic table. It is in the Boron family (group 13) and in period 4. Gallium was discovered in 1875 by Paul Emile Lecoq de Boisbaudran. Boisbaudran named his newly discovered element after himself, deriving from the Latin word, “Gallia,” which means “Gaul.” Elemental Gallium does not exist in nature but gallium (III) salt can be extracted in small amounts from bauxite and zinc ores.
• Chemistry of Indium (Z=49)
Indium has the chemical symbol In and the atomic number 49. It has the electron configuration [Kr] 2s22p1 and may adopt the +1 or +3 oxidation state; however, the +3 state is more common. It is a soft, malleable metal that is similar to gallium. Indium forms InAs, which is found in photoconductors in optical instruments. The physical properties of indium include its silver-white color and the "tin cry" it makes when bent. Indium is soluble in acids, but does not react with oxygen at room tempera
• Chemistry of Thalium (Z=81)
Thallium has the chemical symbol Tl and atomic number 81. It has the electron configuration \([Xe] 2s^22p^1\) and has a +3 or +1 oxidation state. As stated above, because thallium is heavy, it has a greater stability in the +1 oxidation state (inert pair effect). Therefore, it is found more commonly in its +1 oxidation state. Thallium is soft and malleable.
• Chemistry of Nihonium (Z=113)
In studies announced jointly by the Joint Institute for Nuclear Research in Dubna, Russia, and the Lawrence Livermore National Laboratory in the U.S., four atoms of element 113 were produced in 2004 via decay of element 115 after the fusion of Ca-48 and Am-243.
Thumbnail: Crystals of 99.999% gallium. (CC-SA-BY 3.0; Foobar) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_13%3A_The_Boron_Family.txt |
Carbon is one of the most common elements on earth, and greatly influences everyday life. Common molecules containing carbon include carbon dioxide (CO2) and methane (CH4). Many scientists in a variety of fields study of carbon: biologists investigating the origins of life; oceanographers measuring the acidification of the oceans; and engineers developing diamond film tools. This article details the periodic properties of the carbon family and briefly discusses of the individual properties of carbon, silicon, germanium, tin, lead, and flerovium.
• Group 14: General Chemistry
Covers the Group 4 (IUPAC: Group 14) chemistry (carbon, silicon, germanium, tin and lead) and specifically the trend from non-metal to metal as you go down the group, and the increasing tendency towards an oxidation state of +2. Also a certain amount of chemistry of the chlorides and oxides.
• Group 14: General Properties and Reactions
Carbon is one of the most common elements on earth, and greatly influences everyday life. This article details the periodic properties of the carbon family and briefly discusses of the individual properties of carbon, silicon, germanium, tin, lead, and flerovium.
• Chemistry of Carbon (Z=6)
Organic chemistry involves structures and reactions of mainly carbon and hydrogen. Inorganic chemistry deal with interactions of all other pure elements besides carbon, amongst geo/biochemistry. So where does inorganic chemistry of carbon fit in? The inorganic chemistry of carbon also known as inorganic carbon chemistry, is the chemistry of carbon that does not fall within the organic chemistry zone.
• Chemistry of Silicon (Z=14)
Silicon, the second most abundant element on earth, is an essential part of the mineral world. Its stable tetrahedral configuration makes it incredibly versatile and is used in various way in our every day lives. Found in everything from spaceships to synthetic body parts, silicon can be found all around us, and sometimes even in us.
• Chemistry of Germanium (Z=32)
Germanium, categorized as a metalloid in group 14, the Carbon family, has five naturally occurring isotopes. Germanium, abundant in the Earth's crust has been said to improve the immune system of cancer patients. It is also used in transistors, but its most important use is in fiber-optic systems and infrared optics.
• Chemistry of Tin (Z=50)
Mentioned in the Hebrew scriptures, tin is of ancient origins. Tin is an element in Group 14 (The carbon family) and has mainly metallic properties. Tin has atomic number 50 and an atomic mass of 118.710 atomic mass units. Tin, or Sn (from the Latin name Stannum) has been known since ancient times, although it could only be obtained by extraction from its ore. Tin shares chemical similarities with germanium and lead. Tin mining began in Australia in 1872 and today Tin is used extensively.
• Chemistry of Lead (Z=82)
Although lead is not very common in the earth's crust, what is there is readily available and easy to refine. Its chief use today is in lead-acid storage batteries such as those used in automobiles. In pure form it is too soft to be used for much else. Lead has a blue-white color when first cut but quickly dulls on exposure to air, forming Pb2O, one of the few lead(I) compounds. Most stable lead compounds contain lead in oxidation states of +2 or +4.
• Chemistry of Flerovium (Z=114)
The synthesis of element 114 was reported in January of 1999 by scientists from the Joint Institute for Nuclear Research in Dubna (near Moscow) and Lawrence Livermore National Laboratory (in California). In an experiment lasting more than 40 days Russian scientists bombarded a film of Pu-244 supplied by Livermore scientists with a beam of Ca-48. One atom of element 114 was detected with a half-life of more than 30 seconds. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_14%3A_The_Carbon_Family.txt |
The nitrogen family includes the following compounds: nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi). All Group 15 elements have the electron configuration ns2np3 in their outer shell, where n is the principal quantum number.
• Group 15: General Properties and Reactions
The nitrogen family includes the following compounds: nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb), and bismuth (Bi). All Group 15 elements have the electron configuration ns2np3 in their outer shell, where n is equal to the principal quantum number. The nitrogen family is located in the p-block in Group 15, as shown below.
• Chemistry of Nitrogen (Z=7)
Nitrogen is present in almost all proteins and plays important roles in both biochemical applications and industrial applications. Nitrogen forms strong bonds because of its ability to form a triple bond with itself and other elements. Thus, there is a lot of energy in the compounds of nitrogen. Before 100 years ago, little was known about nitrogen. Now, nitrogen is commonly used to preserve food and as a fertilizer.
• Chemistry of Phosphorus (Z=15)
Phosphorus (P) is an essential part of life as we know it. Without the phosphates in biological molecules such as ATP, ADP and DNA, we would not be alive. Phosphorus compounds can also be found in the minerals in our bones and teeth. It is a necessary part of our diet. In fact, we consume it in nearly all of the foods we eat. Phosphorus is quite reactive. This quality of the element makes it an ideal ingredient for matches because it is so flammable.
• Chemistry of Arsenic (Z=33)
Arsenic is situated in the 33rd spot on the periodic table, right next to Germanium and Selenium. Arsenic has been known for a very long time and the person who may have first isolated it is not known but credit generally is given to Albertus Magnus in about the year 1250. The element, which is classified as a metalloid, is named from the Latin arsenicum and Greek arsenikon which are both names for a pigment, yellow orpiment.
• Chemistry of Antimony (Z=51)
Antimony and its compounds have been known for centuries. Scientific study of the element began during the early 17th century, much of the important work being done by Nicolas Lemery. The name of the element comes from the Greek anti + monos for "not alone", while the modern symbol is rooted in the Latin-derived name of the common ore, stibnite.
• Chemistry of Bismuth (Z=83)
Bismuth, the heaviest non-radioactive naturally occurring element, was isolated by Basil Valentine in 1450. It is a hard, brittle metal with an unusually low melting point (271oC). Alloys of bismuth with other low-melting metals such as tin and lead have even lower melting points and are used in electrical solders, fuse elements and automatic fire sprinkler heads.
• Chemistry of Moscovium (Z=115)
In studies announced jointly by the Joint Institute for Nuclear Research in Dubna, Russia, and the Lawrence Livermore National Laboratory in the U.S., four atoms of element 113 were produced in 2004 via decay of element 115 after the fusion of Ca-48 and Am-243.
Thumbnail: White and red phosphorus. (CC-SA-BY 3.0; Peter Krimbacher). | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_15%3A_The_Nitrogen_Family.txt |
The oxygen family, also called the chalcogens, consists of the elements found in Group 16 of the periodic table and is considered among the main group elements. It consists of the elements oxygen, sulfur, selenium, tellurium and polonium. These can be found in nature in both free and combined states.
• Group 16: General Properties and Reactions
The oxygen family, also called the chalcogens, consists of the elements found in Group 16 of the periodic table and is considered among the main group elements. It consists of the elements oxygen, sulfur, selenium, tellurium and polonium. These can be found in nature in both free and combined states. The group 16 elements are intimately related to life.
• Chemistry of Oxygen (Z=8)
Oxygen is an element that is widely known by the general public because of the large role it plays in sustaining life. Without oxygen, animals would be unable to breathe and would consequently die. Oxygen is not only important to supporting life, but also plays an important role in many other chemical reactions. Oxygen is the most common element in the earth's crust and makes up about 20% of the air we breathe.
• Chemistry of Sulfur (Z=16)
Sulfur is a chemical element that is represented with the chemical symbol "S" and the atomic number 16 on the periodic table. Because it is 0.0384% of the Earth's crust, sulfur is the seventeenth most abundant element following strontium. Sulfur also takes on many forms, which include elemental sulfur, organo-sulfur compounds in oil and coal, H2S(g) in natural gas, and mineral sulfides and sulfates.
• Chemistry of Selenium (Z=34)
Element number 34, selenium, was discovered by Swedish chemist Jons Jacob Berzelius in 1817. Selenium is a non-metal and can be compared chemically to its other non-metal counterparts found in Group 16: The Oxygen Family, such as sulfur and tellurium.
• Chemistry of Tellurium (Z=52)
Discovered by von Reichenstein in 1782, tellurium is a brittle metalloid that is relatively rare. It is named from the Latin tellus for "earth". Tellurium can be alloyed with some metals to increase their machinability and is a basic ingredient in the manufacture of blasting caps. Elemental tellurium is occasionally found in nature but is more often recovered from various gold ores.
• Chemistry of Polonium (Z=84)
Polonium was discovered in 1898 by Marie Curie and named for her native country of Poland. The discovery was made by extraction of the remaining radioactive components of pitchblende following the removal of uranium. There is only about 10-6 g per ton of ore! Current production for research purposes involves the synthesis of the element in the lab rather than its recovery from minerals. This is accomplished by producing Bi-210 from the abundant Bi-209.
• Chemistry of Livermorium (Z=116)
In May of 2012 the IUPAC approved the name "Livermorium" (symbol Lv) for element 116. The new name honors the Lawrence Livermore National Laboratory (1952). A group of researchers of this Laboratory with the heavy element research group of the Flerov Laboratory of Nuclear Reactions took part in the work carried out in Dubna on the synthesis of superheavy elements including element 116.
Thumbnail: A sample of sulfur a member of the oxygen group of elements. (Public Domain; Ben Mills).
Group 16: The Oxygen Family
The oxygen family, also called the chalcogens, consists of the elements found in Group 16 of the periodic table and is considered among the main group elements. It consists of the elements oxygen, sulfur, selenium, tellurium and polonium. These can be found in nature in both free and combined states. The group 16 elements are intimately related to life. We need oxygen all the time throughout our lives. Did you know that sulfur is also one of the essential elements of life? It is responsible for some of the protein structures in all living organisms. Many industries utilize sulfur, but emission of sulfur compounds is often seen more as a problem than the natural phenomenon. The metallic properties of Group 16 elements increase as the atomic number increases. The element polonium has no stable isotopes, and the isotope with mass number 209 has the longest half life of 103 years.
Properties of oxygen are very different from those of other elements of the group, but they all have 2 elections in the outer s orbital, and 4 electrons in the p orbitals, usually written as s2p4.
The electron configurations for each element are given below:
• Oxygen: 1s2 2s2 2p4
• Sulfur: 1s2 2s2p6 3s2p4
• Selenium: 1s2 2s2p6 3s2p6d10 4s2p4
• Tellurium: 1s2 2s2p6 3s2p6d10 4s2p6d10 5s2p4
• Polonium: 1s2 2s2p6 3s2p6d10 4s2p6d10f14 5s2p6d10 6s2p4
Example \(1\): Polonium
Polonium can be written as [Xe] 6s2 4f14 5d10 6p4
The trends of the properties in this group are interesting. Knowing the trend allows us to predict their reactions with other elements. Most trends are true for all groups of elements, and the group trends are due mostly to the size of the atoms and number of electrons per atom. The trends are described below:
1. The metallic properties increase in the order oxygen, sulfur, selenium, tellurium, or polonium. Polonium is essentially a metal. It was discovered by M. Curie, who named it after her native country Poland.
2. Electronegativity, ionization energy (or ionization potential IP), and electron affinity decrease for the group as atomic weight increases.
3. The atomic radii and melting points increase.
4. Oxygen differs from sulfur in chemical properties due to its small size. The differences between \(\ce{O}\) and \(\ce{S}\) are more than the differences between other members.
Metallic character increases down the group, with tellurium classified as a metalloid and polonium as a metal. Melting point, boiling point, density, atomic radius, and ionic radius all increase down the group. Ionization energy decreases down the group. The most common oxidation state is -2; however, sulfur can also exist at a +4 and +6 state, and +2, +4, and +6 oxidation states are possible for Se, Te, and Po.
Table \(1\): Select properties of Group 16 elements
Oxygen Sulfur Selenium Tellurium Polonium
Boiling Pt (°C) -182.962 444.674 685 989.9 962
Ionization Energy (kJ/mol) 1314 1000 941 869 812
Ionic Radius (pm) 140 184 198 221
Oxygen
Oxygen is a gas at room temperature and 1 atm, and is colorless, odorless, and tasteless. It is the most abundant element by mass in both the Earth's crust and the human body. It is second to nitrogen as the most abundant element in the atmosphere. There are many commercial uses for oxygen gas, which is typically obtained through fractional distillation. It is used in the manufacture of iron, steel, and other chemicals. It is also used in water treatment, as an oxidizer in rocket fuel, for medicinal purposes, and in petroleum refining.
Oxygen has two allotropes, O2 and O3. In general, O2 (or dioxygen) is the form referred to when talking about the elemental or molecular form because it is the most common form of the element. The O2 bond is very strong, and oxygen can also form strong bonds with other elements. However, compounds that contain oxygen are considered to be more thermodynamically stable than O2.
The latter allotrope, ozone, is a pale-blue poisonous gas with a strong odor. It is a very good oxidizing agent, stronger than dioxygen, and can be used as a substitute for chlorine in purifying drinking water without giving the water an odd taste. However, because of its unstable nature it disappears and leaves the water unprotected from bacteria. Ozone at very high altitudes in the atmosphere is responsible for protecting the Earth's surface from ultraviolet radiation; however, at lower altitudes it becomes a major component of smog.
Oxygen's primary oxidation states are -2, -1, 0, and -1/2 (in O2-), but -2 is the most common. Typically, compounds with oxygen in this oxidation state are called oxides. When oxygen reacts with metals, it forms oxides that are mostly ionic in nature. These can dissolve in water and react to form hydroxides; they are therefore called basic anhydrides or basic oxides. Nonmetal oxides, which form covalent bonds, are simple molecules with low melting and boiling points.
Compounds with oxygen in an oxidation state of -1 are referred to as peroxides. Examples of this type of compound include \(Na_2O_2\) and \(BaO_2\). When oxygen has an oxidation state of -1/2, as in \(O_2^-\), the compound is called a superoxide.
Oxygen is rarely the central atom in a structure and can never bond with more than 4 elements due to its small size and its inability to create an expanded valence shell. Oxygen reacts with hydrogen to form water, which is extensively hydrogen-bonded, has a large dipole moment, and is considered a universal solvent.
There are a wide variety of oxygen-containing compounds, both organic and inorganic: oxides, peroxides and superoxides, alcohols, phenols, ethers, and carbonyl-containing compounds such as aldehydes, ketones, esters, amides, carbonates, carbamates, carboxylic acids and anhydrides.
Sulfur
Sulfur is a solid at room temperature and 1 atm pressure. It is usually yellow, tasteless, and nearly odorless. It is the sixteenth most abundant element in Earth's crust. It exists naturally in a variety of forms, including elemental sulfur, sulfides, sulfates, and organosulfur compounds. Since the 1890s, sulfur has been mined using the Frasch process, which is useful for recovering sulfur from deposits that are under water or quicksand. Sulfur produced from this process is used in a variety of ways including in vulcanizing rubber and as fungicide to protect grapes and strawberries.
Sulfur is unique in its ability to form a wide range of allotropes, more than any other element in the periodic table. The most common state is the solid S8 ring, as this is the most thermodynamically stable form at room temperature. Sulfur exists in the gaseous form in five different forms (S, S2, S4, S6, and S8). In order for sulfur to convert between these compounds, sufficient heat must be supplied.
Two common oxides of sulfur are sulfur dioxide (SO2) and sulfur trioxide (SO3). Sulfur dioxide is formed when sulfur is combusted in air, producing a toxic gas with a strong odor. These two compounds are used in the production of sulfuric acid, which is used in a variety of reactions. Sulfuric acid is one of the top manufactured chemicals in the United States, and is primarily used in the manufacture of fertilizers.
Sulfur also exhibits a wide range of oxidation states, with values ranging from -2 to +6. It is often the central ion in a compound and can easily bond with up to 6 atoms. In the presence of hydrogen it forms the compound hydrogen sulfide, H2S, a poisonous gas incapable of forming hydrogen bonds and with a very small dipole moment. Hydrogen sulfide can easily be recognized by its strong odor that is similar to that of rotten eggs, but this smell can only be detected at low, nontoxic concentrations. This reaction with hydrogen epitomizes how differently oxygen and sulfur act despite their common valence electron configuration and common nonmetallic properties.
A variety of sulfur-containing compounds exist, many of them organic. The prefix thio- in front of the name of an oxygen-containing compound means that the oxygen atom has been substituted with a sulfur atom. General categories of sulfur-containing compounds include thiols (mercaptans), thiophenols, organic sulfides (thioethers), disulfides, thiocarbonyls, thioesters, sulfoxides, sulfonyls, sulfamides, sulfonic acids, sulfonates, and sulfates.
Selenium
Selenium appears as a red or black amorphous solid, or a red or grey crystalline structure; the latter is the most stable. Selenium has properties very similar to those of sulfur; however, it is more metallic, though it is still classified as a nonmetal. It acts as a semiconductor and therefore is often used in the manufacture of rectifiers, which are devices that convert alternating currents to direct currents. Selenium is also photoconductive, which means that in the presence of light the electrical conductivity of selenium increases. It is also used in the drums of laser printers and copiers. In addition, it has found increased use now that lead has been removed from plumbing brasses.
It is rare to find selenium in its elemental form in nature; it must typically be removed through a refining process, usually involving copper. It is often found in soils and in plant tissues that have bioaccumulated the element. In large doses, the element is toxic; however, many animals require it as an essential micronutrient. Selenium atoms are found in the enzyme glutathione peroxidase, which destroys lipid-damaging peroxides. In the human body it is an essential cofactor in maintaining the function of the thyroid gland. In addition, some research has shown a correlation between selenium-deficient soils and an increased risk of contracting the HIV/AIDS virus.
Tellurium
Tellurium is the metalloid of the oxygen family, with a silvery white color and a metallic luster similar to that of tin at room temperature. Like selenium, it is also displays photoconductivity. Tellurium is an extremely rare element, and is most commonly found as a telluride of gold. It is often used in metallurgy in combination with copper, lead, and iron. In addition, it is used in solar panels and memory chips for computers. It is not toxic or carcinogenic; however, when humans are exposed to too much of it they develop a garlic-like smell on their breaths.
Polonium
Polonium is a very rare, radioactive metal. There are 33 different isotopes of the element and all of the isotopes are radioactive. It exists in a variety of states, and has two metallic allotropes. It dissolves easily into dilute acids. Polonium does not exist in nature in compounds, but it can form synthetic compounds in the laboratory. It is used as an alloy with beryllium to act as a neutron source for nuclear weapons.
Polonium is a highly toxic element. The radiation it emits makes it very dangerous to handle. It can be immediately lethal when applied at the correct dosage, or cause cancer if chronic exposure to the radiation occurs. Methods to treat humans who have been contaminated with polonium are still being researched, and it has been shown that chelation agents could possibly be used to decontaminate humans.
Problems
1. What properties increase down the oxygen family?
2. What element can form the most allotropes in the periodic table?
3. What is photoconductivity and which elements display this property?
4. Ozone (\(O_3\)) is a contributor to smog: True or False
5. How many electrons do elements of the oxygen family have in their outermost shell?
6. What does the term "peroxide" refer to?
7. How many elements in the oxygen family are metals, and which one(s)?
8. What is the most common oxidation state for elements in the oxygen family?
9. What is the most abundant element by mass in the Earth's crust and in the human body?
Solutions
1. Melting point, boiling point, density, atomic radius, and ionic radius.
2. Sulfur.
3. Photoconductivity is when the electrical conductivity of an element increases in the presence of light. Both selenium and tellurium display this property.
4. True.
5. Six.
6. A compound that contains oxygen in the oxidation state of -1.
7. 1; Polonium.
8. -2.
9. Oxygen. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_16%3A_The_Oxygen_Family/1Group_16%3A_General_Properties_and_Reactions.txt |
Oxygen is an element that is widely known by the general public because of the large role it plays in sustaining life. Without oxygen, animals would be unable to breathe and would consequently die. Oxygen not only is important to supporting life, but also plays an important role in many other chemical reactions. Oxygen is the most common element in the earth's crust and makes up about 20% of the air we breathe. Historically the discovery of oxygen as an element essential for combustion stands at the heart of the phlogiston controversy (see below).
The Origin and History
Oxygen is found in the group 16 elements and is considered a chalcogen. Named from the Greek oxys + genes, "acid-former", oxygen was discovered in 1772 by Scheele and independently by Priestly in 1774. Oxygen was given its name by the French scientist, Antoine Lavoisier.
Scheele discovered oxygen through an experiment which involved burning manganese oxide. Scheele came to find that the hot manganese oxide produced a gas which he called "fire air". He also came to find that when this gas was able to come into contact with charcoal, it produced beautiful bright sparks. All of the other elements produced the same gas. Although Scheele discovered oxygen, he did not publish his work until three years after another chemist, Joseph Priestly, discovered oxygen. Joseph Priestly, an English chemist, repeated Scheele's experiment in 1774 using a slightly different setup. Priestly used a 12 in burning glass and aimed the sunlight directly towards the compound that he was testing, mercuric oxide. As a result, he was able to "discover better air" that was shown to expand a mouse's lifetime to four times as long and caused a flame to burn with higher intensity. Despite these experiments, both chemists were not able to pinpoint exactly what this element was. It was not until 1775 that Antoine Lavoisier, a French chemist, was able to recognize this unknown gas as an element.
Our atmosphere currently contains about 21% of free oxygen. Oxygen is produced in various ways. The process of photochemical dissociation in which water molecules are broken up by ultraviolet rays produces about 1-2% of our oxygen. Another process that produces oxygen is photosynthesis which is performed by plants and photosynthetic bacteria. Photosynthesis occurs through the following general reaction:
$\ce{CO2 + H2O + h\nu \rightarrow} \text{organic compounds} \ce{+ O2} \nonumber$
The Dangers of Phlogiston
Phlogiston theory is the outdated belief that a fire-like element called phlogiston is contained within combustible bodies and released during combustion. The name comes from the Ancient Greek φλογιστόν phlogistón (burning up), from φλόξ phlóx (flame). It was first stated in 1667 by Johann Joachim Becher, and then put together more formally by Georg Ernst Stahl. The theory attempted to explain burning processes such as combustion and rusting, which are now collectively known as oxidation.
Properties
• Element number: 8
• Atomic weight 15.9994
• Color: gas form- colorless, liquid- pale blue
• Melting point: 54.36K
• Boiling point: 90.2 K
• Density: .001429
• 21% of earth's atmosphere
• Third most abundant element in the universe
• Most abundant element in Earth's crust at 45.4%
• 3 Stable isotopes
• Ionization energy: 13.618 eV
• Oxygen is easily reduced and is a great oxidizing agent making it readily reactive with other elements
Magnetic Properties of Oxygen
Oxygen (O2) is paramagnetic. An oxygen molecule has six valence electrons, so the O2 molecule has 12 valence electrons with the electron configuration shown below:
As shown, there are two unpaired electrons, which causes O2 to be paramagnetic. There are also eight valence electrons in the bonding orbitals and four in antibonding orbitals, which makes the bond order 2. This accounts for the double covalent bond that is present in O2.
Video $1$: A chemical demonstration of the paramagnetism of molecular oxygen, as shown by the attraction of liquid oxygen to magnets.
As shown in Video $1$, since molecular oxygen ($O_2$) has unpaired electrons, it is paramagnetic and is attracted to the magnet. In contrast, molecular nitrogen ($N_2$) has no unpaired electrons and is not attracted to the magnet.
General Chemistry of Oxygen
Oxygen normally has an oxidation state of -2, but is capable of having oxidation states of -2, -1, -1/2, 0, +1, and +2. The oxidation states of oxides, peroxides and superoxides are as follows:
• Oxides: O-2 ,
• peroxides: O2-2 ,
• superoxide: O2-1.
Oxygen does not react with itself, nitrogen, or water under normal conditions. Oxygen does, however, dissolve in water at 20 degrees Celsius and 1 atmosphere. Oxygen also does not normally react with bases or acids. Group 1 metals (alkaline metals) are very reactive with oxygen and must be stored away from oxygen in order to prevent them from becoming oxidized. The metals at the bottom of the group are more reactive than those at the top. The reactions of a few of these metals are explored in more detail below.
Lithium: Reacts with oxygen to form white lithium oxide in the reaction below.
$\ce{4Li + O_2 \rightarrow 2Li_2O} \label{1}$
Sodium: Reacts with oxygen to form a white mixture of sodium oxide and sodium peroxide. The reactions are shown below.
• Sodium oxide: $\ce{4Na + O_2 \rightarrow 2Na_2O} \label{2}$
• Sodium peroxide: $\ce{2Na + O_2 \rightarrow Na_2O_2} \label{3}$
Potassium: Reacts with oxygen to form a mixture of potassium peroxide and potassium superoxide. The reactions are shown below.
• Potassium peroxide: $\ce{2K + O_2 \rightarrow 2K_2O_2} \label{4}$
• Potassium superoxide: $\ce{K + O_2 \rightarrow KO_2} \label{5}$
Rubidium and Cesium: Both metals react to produce superoxides through the same process as that of the potassium superoxide reaction.
The oxides of these metals form metal hydroxides when they react with water. These metal hydroxides make the solution basic or alkaline, hence the name alkaline metals.
Group 2 metals (alkaline earth metals) react with oxygen through the process of burning to form metal oxides but there are a few exceptions.
Beryllium is very difficult to burn because it has a layer of beryllium oxide on its surface which prevents further interaction with oxygen. Strontium and barium react with oxygen to form peroxides. The reaction of barium and oxygen is shown below, and the reaction with strontium would be the same.
$\ce{Ba(s) + O2 (g) \rightarrow BaO2 (s) }\label{6}$
Group 13 reacts with oxygen in order to form oxides and hydroxides that are of the form $X_2O_3$ and $X(OH)_3$. The variable X represents the various group 13 elements. As you go down the group, the oxides and hydroxides get increasingly basic.
Group 14 elements react with oxygen to form oxides. The oxides formed at the top of the group are more acidic than those at the bottom of the group. Oxygen reacts with silicon and carbon to form silicon dioxide and carbon dioxide. Carbon is also able to react with oxygen to form carbon monoxide, which is slightly acidic. Germanium, tin, and lead react with oxygen to form monoxides and dioxides that are amphoteric, which means that they react with both acids and bases.
Group 15 elements react with oxygen to form oxides. The most important are listed below.
• Nitrogen: N2O, NO, N2O3, N2O4, N2O5
• Phosphorus: P4O6, P4O8, P2O5
• Arsenic: As2O3, As2O5
• Antimony: Sb2O3, Sb2O5
• Bismuth: Bi2O3, Bi2O5
Group 16 elements react with oxygen to form various oxides. Some of the oxides are listed below.
• Sulfur: SO, SO2, SO3, S2O7
• Selenium: SeO2, SeO3
• Tellurium: TeO, TeO2, TeO3
• Polonium: PoO, PoO2, PoO3
Group 17 elements (halogens) fluorine, chlorine, bromine, and iodine react with oxygen to form oxides. Fluorine forms two oxides with oxygen: F2O and F2O2. Both fluorine oxides are called oxygen fluorides because fluorine is the more electronegative element. One of the fluorine reactions is shown below.
$\ce{O2 (g) + F2 (g) \rightarrow F2O2 (g)} \label{7}$
Group 18: Some would assume that the Noble Gases would not react with oxygen. However, xenon does react with oxygen to form $\ce{XeO_3}$ and $\ce{XeO_4}$. The ionization energy of xenon is low enough for the electronegative oxygen atom to "steal away" electrons. Unfortunately, $\ce{XeO_3}$ is HIGHLY unstable, and it has been known to spontaneously detonate in a clean, dry environment.
Transition metals react with oxygen to form metal oxides. However, gold, silver, and platinum do not react with oxygen. A few reactions involving transition metals are shown below:
$2Sn_{(s)} + O_{2(g)} \rightarrow 2SnO_{(s)} \label{8}$
$4Fe_{(s)} + 3O_{2(g)} \rightarrow 2Fe_2O_{3(s)} \label{9A}$
$4Al_{(s)} + 3O_{2(g)} \rightarrow 2Al_2O_{3(s)} \label{9B}$
Reaction of Oxides
We will be discussing metal oxides of the form $X_2O$. The variable $X$ represents any metal that is able to bond to oxygen to form an oxide.
• Reaction with water: The oxides react with water to form a metal hydroxide.
$X_2O + H_2O \rightarrow 2XOH \nonumber$
• Reaction with dilute acids: The oxides react with dilute acids to form a salt and water.
$X_2O + 2HCl \rightarrow 2XCl + H_2O \nonumber$
Reactions of Peroxides
The peroxides we will be discussing are of the form $X_2O_2$. The variable $X$ represents any metal that can form a peroxide with oxygen.
Reaction with water: If the temperature of the reaction is kept constant despite the fact that the reaction is exothermic, then the reaction proceeds as follows:
$X_2O_2+ 2H_2O \rightarrow 2XOH + H_2O_2 \nonumber$
If the reaction is not carried out at a constant temperature, then the reaction of the peroxide and water will result in decomposition of the hydrogen peroxide that is produced into water and oxygen.
Reaction with dilute acid: This reaction is more exothermic than that with water. The heat produced causes the hydrogen peroxide to decompose to water and oxygen. The reaction is shown below.
$X_2O_2 + 2HCl \rightarrow 2XCl + H_2O_2 \nonumber$
$2H_2O_2 \rightarrow 2H_2O + O_2 \nonumber$
Reaction of Superoxides
The superoxides we will be talking about are of the form $XO_2$, with $X$ representing any metal that forms a superoxide when reacting with oxygen.
Reaction with water: The superoxide and water react in a very exothermic reaction that is shown below. The heat that is produced in forming the hydrogen peroxide will cause the hydrogen peroxide to decompose to water and oxygen.
$2XO_2 + 2H_2O \rightarrow 2XOH + H_2O_2 + O_2 \nonumber$
Reaction with dilute acids: The superoxide and dilute acid react in a very exothermic reaction that is shown below. The heat produced will cause the hydrogen peroxide to decompose to water and oxygen.
$2XO_2 + 2HCl \rightarrow 2XCl + H_2O_2 + O_2 \nonumber$
Allotropes of Oxygen
There are two allotropes of oxygen; dioxygen (O2) and trioxygen (O3) which is called ozone. The reaction of converting dioxygen into ozone is very endothermic, causing it to occur rarely and only in the lower atmosphere. The reaction is shown below:
$3O_{2 (g)} \rightarrow 2O_{3 (g)} \;\;\; ΔH^o= +285 \;kJ \nonumber$
Ozone is unstable and quickly decomposes back to oxygen but is a great oxidizing agent.
Miscellaneous Reactions
Reaction with Alkanes: The most common reactions that involve alkanes occur with oxygen. Alkanes are able to burn and it is the process of oxidizing the hydrocarbons that makes them important as fuels. An example of an alkane reaction is the reaction of octane with oxygen as shown below.
C8H18(l) + 25/2 O2(g) → 8CO2(g) + 9H2O(l) ∆Ho = -5.48 X 103 kJ
Reaction with ammonia: Oxygen is able to react with ammonia to produce dinitrogen (N2) and water (H2O) through the reaction shown below.
$4NH_3 + 3O_2 \rightarrow 2N_2 + 6H_2O \nonumber$
Reaction with Nitrogen Oxide: Oxygen is able to react with nitrogen oxide in order to produce nitrogen dioxide through the reaction shown below.
$NO + O_2 \rightarrow NO_2 \nonumber$
Problems
1. Is oxygen reactive with noble gases?
2. Which transition metals does oxygen not react with?
3. What is produced when an oxide reacts with water?
4. Is oxygen reactive with alkali metals? Why are the alkali metals named that way?
5. If oxygen is reactive with alkali metals, are oxides, peroxides or superoxides produced?
Solutions
1. No, noble gases are unreactive with oxygen.
2. Oxygen is mostly unreactive with gold and platinum.
3. When an oxide reacts with water, a metal hydroxide is produced.
4. Oxygen is very reactive with alkali metals. Alkali metals are given the name alkali because the oxides of these metals react with water to form a metal hydroxide that is basic or alkaline.
5. Lithium produces an oxide, sodium produces a peroxide, and potassium, cesium, and rubidium produce superoxides.
Contributors and Attributions
• Phillip Ball (UCD), Katharine Williams (UCD) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_16%3A_The_Oxygen_Family/Z008_Chemistry_of_Oxygen_%28Z8%29.txt |
Learning Objectives
• Describe the chemistry of the oxygen group.
• Give the trend of various properties.
• Remember the names of Group 16 elements.
• Explain the Frasch process.
• Describe properties and applications of $\mathrm{H_2SO_4}$.
• Explain properties and applications of $\mathrm{H_2S}$.
Sulfur is a chemical element that is represented with the chemical symbol "S" and the atomic number 16 on the periodic table. Because it is 0.0384% of the Earth's crust, sulfur is the seventeenth most abundant element following strontium. Sulfur also takes on many forms, which include elemental sulfur, organo-sulfur compounds in oil and coal, H2S(g) in natural gas, and mineral sulfides and sulfates. This element is extracted by using the Frasch process (discussed below), a method where superheated water and compressed air are used to draw liquid sulfur to the surface. Offshore sites, Texas, and Louisiana are the primary sites that yield extensive amounts of elemental sulfur. However, elemental sulfur can also be produced by reducing H2S, commonly found in oil and natural gas. For the most part, though, sulfur is used to produce SO2(g) and H2SO4.
Known from ancient times (mentioned in the Hebrew scriptures as brimstone) sulfur was classified as an element in 1777 by Lavoisier. Pure sulfur is tasteless and odorless with a light yellow color. Samples of sulfur often encountered in the lab have a noticeable odor. Sulfur is the tenth most abundant element in the known universe.
Sulfur at a Glance
Atomic Number 16
Atomic Symbol S
Atomic Weight 32.07 grams per mole
Structure orthorhombic
Phase at room temperature solid
Classification nonmetal
Physical Properties of Sulfur
Sulfur has an atomic weight of 32.066 grams per mole and is part of group 16, the oxygen family. It is a nonmetal and has a specific heat of 0.706 J g-1 oC-1. The electron affinity is 200 kJ mol-1 and the electronegativity is 2.58 (unitless). Sulfur is typically found as a light-yellow, opaque, and brittle solid in large amounts of small orthorhombic crystals. Not only does sulfur have twice the density of water, but it is also insoluble in water. On the other hand, sulfur is highly soluble in carbon disulfide and slightly soluble in many common solvents. Sulfur can also vary in color and blackens upon boiling due to carbonaceous impurities. Even as little as 0.05% of carbonaceous matter darkens sulfur significantly.
Most sulfur is recovered directly as the element from underground deposits by injecting super-heated water and piping out molten sulfur (sulfur melts at 112o C). Compared to other elements, sulfur has the most allotropes. While the S8 ring is the most common allotrope, there are 6 other structures with up to 20 sulfur atoms per ring.
• Under appropriate conditions, sulfur vapor can contain $S$, $S_2$, $S_4$, $S_6$, and $S_8$.
• At room temperature, rhombic sulfur (Sα) is a stable solid comprising cyclic $S_8$ molecules.
• At 95.5 °C, rhombic sulfur becomes monoclinic sulfur (Sβ). The crystal structure found in monoclinic sulfur differs from that of rhombic sulfur. Monoclinic sulfur is also made up of $S_8$molecules.
• Monoclinic sulfur becomes liquid sulfur (Sλ) at 119 °C. Liquid sulfur is a straw-colored liquid made up of $S_8$ molecules and other cyclic molecules containing a range of six to twenty atoms.
• At 160 oC, this becomes a dark, viscous liquid called liquid sulfur (Sμ). The molecules are still made up of eight sulfur atoms, but the molecule opens up and transforms from a circle into a long spiral-chain molecule.
• At 180 °C, the chain length and viscosity reach their maximum. Chains break and viscosity decreases at temperatures that exceed 180 °C.
• Sulfur vapor is produced when liquid boils at 445 °C. In the vapor that is produced, $S_8$ molecules dominate, but as the vapor continues to heat up, the molecules break up into smaller groups of sulfur atoms.
• To produce plastic sulfur, S is poured into cold water. Plastic sulfur is rubberlike and is made up of long, spiral-chain molecules. If plastic sulfur sits for long, it will reconvert to rhombic sulfur.
While oxygen has fewer allotropes than sulfur, including $\ce{O}$, $\ce{O_2}$, $\ce{O_3}$, $\ce{O_4}$, $\ce{O_8}$, metallic $\ce{O}$ (and four other solid phases), many of these actually have a corresponding sulfur variant. However, sulfur has more tendency to catenate (the linkage of atoms of the same element into longer chains). Here are the values of the single and double bond enthalpies:
$\begin{array}{c|r} \ce {O-O} & \ce{142\ kJ/mol} \ \ce {S–S} & \ce{268\ kJ/mol} \ \ce {O=O} & \ce{499\ kJ/mol} \ \ce {S=S} & \ce{352\ kJ/mol} \ \end{array} \nonumber$
This means that $\ce{O=O}$ is stronger than $\ce{S=S}$, while $\ce{O–O}$ is weaker than $\ce{S–S}$. So, in sulfur, single bonds are favored and catenation is easier than in oxygen compounds. It seems that the reason for the weaker $\ce{S=S}$ double bonds has its roots in the size of the atom: it's harder for the two atoms to come to a small enough distance, so that the $p$ orbital overlap is small and the $\pi$ bond is weak. This is attested by looking down the periodic table: $\ce{Se=Se}$ has an even weaker bond enthalpy of $\ce{272 kJ/mol}$.
What happens when the solid sulfur melts? The $\ce{S8}$ molecules break up. When suddenly cooled, long chain molecules are formed in the plastic sulfur which behave like rubber. Plastic sulfur transforms into rhombic sulfur over time.
Compounds
Reading the following reactions, figure out and notice the change of the oxidation state of $\ce{S}$ in the reactants and products. Common oxidation states of sulfur are -2, 0, +4, and +6. Sulfur (brimstone, stone that burns) reacts with $\ce{O2}$ giving a blue flame (Figure $1$):
$\ce{S + O_2 \rightarrow SO_2} \nonumber$
$\ce{SO2}$ is produced whenever a metal sulfide is oxidized. It is recovered and oxidized further to give $\mathrm{SO_3}$, for production of $\mathrm{H_2SO_4}$. $\mathrm{SO_2}$ reacts with $\mathrm{H_2S}$ to form $\mathrm{H_2O}$ and $\ce{S}$.
$\mathrm{2 SO_2 + O_2 \rightleftharpoons 2 SO_3} \nonumber$
$\mathrm{SO_3 + H_2O \rightleftharpoons H_2SO_4} \;\;(\text{a valuable commodity}) \nonumber$
$\mathrm{SO_3 + H_2SO_4 \rightleftharpoons H_2S_2O_7} \;\;\; (\text{pyrosulfuric acid}) \nonumber$
Sulfur reacts with sulfite ions in solution to form thiosulfate,
$\ce{S + SO_3^{2-} -> S_2O_3^{2-}} \nonumber$
but the reaction is reversed in an acidic solution.
Oxides
There are many different stable sulfur oxides, but the two that are commonly found are sulfur dioxide and sulfur trioxide. Sulfur dioxide is a commonly found oxide of sulfur. It is a colorless, pungent, and nonflammable gas. It has a density of 2.8 kg/m3 and a melting point of -72.5 oC. Because organic materials are more soluble in $SO_2$ than in water, the liquid form is a good solvent. $SO_2$ is primarily used to produce $SO_3$. The direct combustion of sulfur and the roasting of metal sulfides yield $SO_2$ via the contact process:
$\underbrace{S(s) + O_2(g) \rightarrow SO_2(g)}_{\text{Direct combustion}} \nonumber$
$\underbrace{2 ZnS(s) + 3 O_2(g) \rightarrow 2 ZnO(s) + 2 SO_2(g)}_{\text{Roasting of metal sulfides}} \nonumber$
Sulfur trioxide is another one of the commonly found oxides of sulfur. It is a colorless liquid with a melting point of 16.9 oC and a density of kg/m3. $SO_3$ is used to produce sulfuric acid. $SO_2$ is used in the synthesis of $SO_3$:
$\underbrace{2 SO_2 (g) + O_2(g) \rightleftharpoons 2 SO_3(g)}_{\text{Exothermic, reversible reaction}} \nonumber$
This reaction needs a catalyst to be completed in a reasonable amount of time with $V_2O_5$ being the catalyst most commonly used.
Hydrogen Sulfide H2S
• Hydrogen sulfide, $\ce{H2S}$, is a diprotic acid. The equilibria below, $\mathrm{H_2S \rightleftharpoons HS^- + H^+} \nonumber$ $\mathrm{HS^- \rightleftharpoons S^{2-} + H^+} \nonumber$ have been discussed in connection with Polyprotic Acids.
Other Sulfur-containing Compounds
Perhaps the most significant compound of sulfur used in modern industrialized societies is sulfuric acid ($H_2SO_4$). Sulfur dioxide ($SO_2$) finds practical applications in bleaching and refrigeration but it is also a nuisance gas resulting from the burning of sulfurous coals. Sulfur dioxide gas then reacts with the water vapor in the air to produce a weak acid, sulfurous acid ($H_2SO_3$), which contributes to the acid rain problem.
• Sulfuric acid, H2SO4, is produced by reacting $SO_3$ with water. However, this often leads to pollution problems. SO3(g) is reacted with 98% H2SO4 in towers full of ceramic material to produce H2S2O7 or oleum. Water is circulated in the tower to maintain the correct concentration and the acid is diluted with water at the end in order to produce the correct concentration. Pure sulfuric acid has no color and odor, and it is an oily, hygroscopic liquid. However, sulfuric acid vapor produces heavy, white smoke and a suffocating odor.
• Dilute sulfuric acid, H2SO4(aq), reacts with metals and acts as a strong acid in common chemical reactions. It is used to produce H2(g) and liberate CO2(g) and can neutralize strong bases.
• Concentrated sulfuric acid, H2SO4 (conc.), has a strong affinity for water. In some cases, it removes H and O atoms. Concentrated sulfuric acid is also a good oxidizing agent and reacts with some metals.
$C_{12}H_{22}O_{11}(s) \rightarrow 12 C(s) + 11 H_2O(l) \nonumber$
(Concentrated sulfuric acid used in forward reaction to remove H and O atoms.)
Applications of Sulfuric Acid
• as a strong acid for making $\ce{HCl}$ and $\mathrm{HNO_3}$.
• as an oxidizing agent for metals.
• as a dehydrating agent.
• for manufacture of fertilizer and other commodities.
• Sulfurous acid (H2SO3) is produced when $SO_2$(g) reacts with water. It cannot be isolated in its pure form; however, it forms salts as sulfites. Sulfites can act as both reducing agents and oxidizing agents.
O2(g) + 2 SO32-(aq) $\rightarrow$ 2 SO42- (aq) (Reducing agent)
2 H2S(g) + 2 H+(aq) + SO32-(aq) $\rightarrow$ 3 H2O(l) + 3 S(s) (Oxidizing agent)
H2SO3 is a diprotic acid that acts as a weak acid in both steps, and H2SO4 is also a diprotic acid but acts as a strong acid in the first step and a weak acid in the second step. Acids like NaHSO3 and NaHSO4 are called acid salts because they are the product of the first step of these diprotic acids.
Boiling elemental sulfur in a solution of sodium sulfite yields thiosulfate. Not only are thiosulfates important in photographic processing, but they are also common analytical reagents used with iodine (like in the following two reactions).
$2 Cu^{2+}_{(aq)} + 5 I^-_{(aq)} \rightarrow 2 CuI_{(s)} + I^-_{3(aq)} \nonumber$
$I^-_{3(aq)} + 2 S_2O^{2-}_{3(aq)} \rightarrow 3 I^-_{(aq)} + S_4O^{2-}_{6(aq)} \nonumber$
with excess triiodide ion titrated with Na2S2O3(aq).
Other than sulfuric acid, perhaps the most familiar compound of sulfur in the chemistry lab is the foul-smelling hydrogen sulfide gas, $H_2S$, which smells like rotten eggs.
• Sulfur halides are compounds formed between sulfur and the halogens. Common compounds include SF2, S2F2, SF4, and SF6. While SF4 is a powerful fluorinating agent, SF6 is a colorless, odorless, unreactive gas. Compounds formed by sulfur and chloride include S2Cl2, SCl4, and SCl2. SCl2 is a red, bad-smelling liquid that is utilized to produce mustard gas ($S(CH_2CH_2Cl)_2$).
$SCl_2 + 2CH_2CH_2 \rightarrow S(CH_2CH_2Cl)_2 \nonumber$
Production -The Frasch Process
Sulfur can be mined by the Frasch process. This process has made sulfur a high purity (up to 99.9 percent pure) chemical commodity in large quantities. Most sulfur-containing minerals are metal sulfides, and the best known is perhaps pyrite ($\mathrm{FeS_2}$, known as fool's gold because of its golden color). The most common sulfate-containing mineral is gypsum, $\mathrm{CaSO_4 \cdot 2H_2O}$, also known as plaster of Paris.
The Frasch process is based on the fact that sulfur has a comparatively low melting point. The process forces (99.5% pure) sulfur out by using hot water and air. In this process, superheated water is forced down the outermost of three concentric pipes. Compressed air is pumped down the center tube, and a mixture of elemental sulfur, hot water, and air comes up the middle pipe. Sulfur is melted with superheated water (at 170 °C under high pressure) and forced to the surface of the earth as a slurry.
Sulfur is mostly used for the production of sulfuric acid, $\ce{H2SO4}$. Most sulfur mined by the Frasch process is used in industry for the manufacture of sulfuric acid. Sulfuric acid, the most abundantly produced chemical in the United States, is manufactured by the contact process. Most (about 70%) of the sulfuric acid produced in the world is used in the fertilizer industry. Sulfuric acid can act as a strong acid, a dehydrating agent, and an oxidizing agent. Its applications use these properties. Sulfur is an essential element of life in sulfur-containing proteins.
Applications
Sulfur has many practical applications. As a fungicide, sulfur is used to counteract apple scab in organically farmed apple production. Other crops that utilize sulfur fungicides include grapes, strawberries, and many vegetables. In general, sulfur is effective against mildew diseases and black spot. Sulfur can also be used as an organic insecticide. Sulfites are frequently used to bleach paper and preserve dried fruit.
The vulcanization of rubber includes the use of sulfur as well. Cellophane and rayon are produced with carbon disulfide, a product of sulfur and methane. Sulfur compounds can also be found in detergents, acne treatments, and agrichemicals. Magnesium sulfate (epsom salt) has many uses, ranging from bath additives to exfoliants. Sulfur is being increasingly used as a fertilizer as well. Because standard sulfur is hydrophobic, it is covered with a surfactant by bacteria before oxidation can occur. Sulfur is therefore a slow-release fertilizer. Lastly, sulfur functions as a light-generating medium in sulfur lamps.
Concentrated sulfuric acid was once one of the most produced chemicals in the United States; the majority of the H2SO4 that is now produced is used in fertilizer. It is also used in oil refining, production of titanium dioxide, and in emergency power supplies and car batteries. The mineral gypsum, or calcium sulfate dihydrate, is used in making plaster of Paris. Over one million tons of aluminum sulfate is produced each year in the United States by reacting H2SO4 and Al2O3. This compound is important in water purification. Copper sulfate is used in electroplating. Sulfites are used in the paper making industry because they produce a substance that coats the cellulose in the wood and frees the fibers of the wood for treatment.
Emissions and the Environment
Particles, SO2(g), and H2SO4 mist are the components of industrial smog. Because power plants burn coal or high-sulfur fuel oils, SO2(g) is released into the air. When catalyzed on the surfaces of airborne particles, SO2 can be oxidized to SO3. A reaction with NO2 works as well as shown in the following reaction:
$SO_{2(g)} + NO_{2(g)} \rightarrow SO_{3(g)} + NO_{(g)} \nonumber$
H2SO4 mist is then produced after SO3 reacts with water vapor in the air. If H2SO4 reacts with airborne NH3, (NH4)2SO4 is produced. When SO2(g) and H2SO4 reach levels that exceed 0.10 ppm, they are potentially harmful. By removing sulfur from fuels and controlling emissions, acid rain and industrial smog can be kept under control. Processes like fluidized bed combustion have been presented to remove SO2 from smokestack gases.
• Dhawale, S.W. "Thiosulfate: An interesting sulfur oxoanion that is useful in both medicine and industry--but is implicated in corrosion." J. Chem. Educ. 1993, 70, 12.
• Lebowitz, Samuel H. "A demonstration working model of the Frasch process for mining sulfur." J. Chem. Educ. 1931, 8, 1630.
• Nagel, Miriam C. "Herman Frasch, sulfur king (PROFILES)." J. Chem. Educ. 1981, 58, 60.
• Riethmiller, Steven. "Charles H. Winston and Confederate Sulfuric Acid." J. Chem. Educ. 1995 72 575.
• Sharma, B. D. "Allotropes and polymorphs." J. Chem. Educ. 1987, 64, 404.
• Silverstein, Todd P.; Zhang, Yi. "Sugar Dehydration without Sulfuric Acid: No More Choking Fumes in the Classroom!" J. Chem. Educ. 1998 75 748.
• Tykodi, R. J. "In praise of thiosulfate." J. Chem. Educ. 1990, 67, 146.
• Thomas Jefferson National Accelerator Facility - Office of Science Education."It's Elemental-The Element Sulfur." Jefferson Lab.
• Sulfur's Electron Shell
Problems
1. Draw a diagram that summarizes the allotropy of sulfur. Use symbols, arrows, and numbers.
2. Direct combustion of sulfur is the only method for producing SO2(g). True or False.
3. Sulfites are not oxidizing agents. They are good reducing agents. True or False.
4. Give the reaction for the production of sulfur trioxide.
5. Choose the incorrect statement.
1. Sulfur produces cellophane and rayon.
2. Standard sulfur is hydrophobic.
3. SO2 can oxidize to SO3
4. Sulfur influences the development of acid rain and industrial smog.
5. All of the above are correct.
6. Which reaction is responsible for the destruction of limestone and marble statues and buildings?
1. $\ce{CaCO3 \rightarrow CaO + CO2}$
2. $\ce{SO2 + H2O \rightarrow H2SO3}$
3. $\ce{BaO + CO2 \rightarrow BaCO3 \rightarrow BaSO3}$ upon reaction with $\ce{SO2}$
4. $\ce{CaCO3 + H2O \rightarrow Ca(OH)2 + CO2}$
5. $\ce{CaCO3 + SO2 \rightarrow CaSO3 + CO2 \rightarrow CaSO4}$ upon oxidation
7. Give the formula of thiosulfate ion.
8. What is the oxidation state of $\ce{S}$ in $\ce{SF6}$, $\ce{H2SO4}$, $\ce{NaHSO4}$, $\ce{SO4^2-}$, and $\ce{SO3}$?
9. What is the phase of sulfur at 298 K? Enter the type of crystals.
10. Give the name of the process by which sulfur is forced out of the ground using hot water and air.
Solutions
1. The diagram may be drawn in any way. However, the symbols (S2), (S4), (S6), (S?), and (S8(g)) must be included. The temperatures should be written next to the arrows.
2. False
3. False
4. $2 SO_{2(g)} + O_{2(g)} \rightarrow 2 SO_{3(g)}$
5. A
6. e.
Consider...
$\ce{SO2}$ in $\ce{H2SO3}$ is the acid in acid rain, which attacks $\ce{CaCO3}$, marble. $\ce{SO2}$ reduces pigments in organic matter.
7. $\ce{S2O3^2-}$
Consider...
Sulfate is $\ce{SO4^2-}$; replacement of an $\ce{O}$ by an $\ce{S}$ gives thiosulfate $\ce{S2O3^2-}$. The two $\ce{S}$ in $\ce{S2O3^2-}$ have different oxidation states: one is +6, the other is (-2), average +2.
8. 6
Consider...
Oxidation state for $\ce{S}$ in $\ce{H2SO3}$, $\ce{SO3^2-}$, $\ce{SO2}$, etc. is 4. The oxidation state of $\ce{S}$ is the same for all in the list.
9. rhombic sulfur
Consider...
The term rhombic describes a type of crystal. Monoclinic sulfur is meta stable at 298 K.
10. Frasch process
Consider...
The Frasch process is used to mine elemental sulfur. | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_16%3A_The_Oxygen_Family/Z016_Chemistry_of_Sulfur_%28Z16%29.txt |
Element number 34, selenium, was discovered by Swedish chemist Jons Jacob Berzelius in 1817. Selenium is a non-metal and can be compared chemically to its other non-metal counterparts found in Group 16: The Oxygen Family, such as sulfur and tellurium.
Properties
Chemical Symbol: Se
Atomic Number: 34
Atomic Weight: 78.96
Electron Configuration: [Ar] 4s23d104p4
Melting Point: 493.65 K
Boiling Point: 958 K
Electronegativity: 2.55 (Pauling)
Oxidation States: Se-2, Se+6, Se+4
Ionization Energies: First: 941 kJ/mol
Second: 2045 kJ/mol
Third: 2973.7 kJ/mol
History
Selenium was discovered by Berzelius in 1818. It is named for the Greek word for "moon", selene. The discovery of selenium was an important finding, but at the same time seemingly accidental. Fellow scientist Martin Klaproth discovered a contamination of sulfuric acid creating a red colored product which he believed to be due to the element tellurium. However, Berzelius went on to further analyze the impurity and came to the conclusion that it was an unknown element that shared properties similar to those of tellurium. Based on the Greek word “selene,” meaning moon, Jons Berzelius decided to call the newly found element selenium.
Allotropes and Physical Properties
Selenium can exist in multiple allotropes that are essentially different molecular forms of an element with varying physical properties. For example, one allotrope of selenium can be seen as an amphorous (“without crystalline shape”) red powder. Selenium also takes a crystalline hexagonal structure, forming a metallic gray allotrope which is known to be stable. The most thermodynamically stable allotrope of selenium is trigonal selenium, which also appears as a gray solid. Most selenium is recovered from the electrolytic copper refining process. This is usually in the form of the red allotrope.
Selenium is mostly noted for its important chemical properties, especially those dealing with electricity. Unlike sulfur, selenium is a semiconductor, meaning that it conducts some electricity, but not as well as conductors. Selenium is a photoconductor, which means it has the ability to change light energy into electrical energy. Not only is selenium able to convert light energy into electrical energy, but it also displays the property of photoconductivity. Photoconductivity is the idea that the electrical conductivity of selenium increases due to the presence of light -- or, in other words, it becomes a better photoconductor as light intensity increases.
Isotopes
Isotopes of an element are atoms that have the same atomic numbers but a different number of neutrons (different mass numbers) in their nuclei. Selenium is known to have over 20 different isotopes; however, only 5 of them are stable. The five stable isotopes of selenium are 74Se, 76Se, 77Se, 78Se, 80Se.
Uses
Due to selenium’s property of photoconductivity, it is known to be used in photocells, exposure meters in photography, and also in solar cells. Selenium can also be seen in the products of plain-paper photocopiers, laser printers and photographic toners. Besides its uses in the electronic industry, selenium is also popular in the glass-making industry. When selenium is added to glass, it is able to negate the color of other elements found in the glass and essentially decolorizes it. Selenium is also able to create a ruby-red colored glass when added. The element can also be used in the production of alloys and is an additive to stainless steel.
Health Hazards
Selenium, a trace element, is important in the diet and health of both plants and animals, but can be only taken in very small amounts. Exposure to an excess amount of selenium is known to be toxic and causes health problems. With an upper intake level of 400 micrograms per day that can be tolerated, too much selenium can lead to selenosis and may result in health problems and even death. Compounds of selenium are also known to be carcinogenic.
Chemical Reactivity
Reaction with hydrogen
Selenium forms hydrogen selenide, H2Se, a colorless flammable gas, when reacted with hydrogen.
Reaction with oxygen
Selenium burns in air, displaying a blue flame, and forms solid selenium dioxide.
$Se_{8(s)} + 8O_{2(g)} \rightarrow 8SeO_{2(s)} \nonumber$
Selenium is also known to form selenium trioxide, SeO3.
Reaction with halides
Selenium reacts with fluorine, F2, and burns to form selenium hexafluoride.
$Se_{8(s)} + 24F_{2(g)} \rightarrow 8SeF_{6(l)} \nonumber$
Selenium also reacts with chlorine and bromine to form diselenium dichloride, $Se_2Cl_2$ and diselenium dibromide, $Se_2Br_2$.
$Se_8 + 4Cl_2 \rightarrow 4Se_2Cl_{2(l)} \nonumber$
$Se_8 + 4Br_2 \rightarrow 4Se_2Br_{2(l)} \nonumber$
Selenium also forms $SeF_4$, $SeCl_2$ and $SeCl_4$.
Selenides
Selenium reacts with metals to form selenides. Example: Aluminum selenide
$3 Se_8 + 16 Al \rightarrow 8 Al_2Se_3 \nonumber$
Selenites
Selenium reacts to form salts called selenites, e.g., silver selenite (Ag2SeO3) and sodium selenite (Na2SeO3).
Problems
1. Describe selenium’s property of photoconductivity.
2. Does selenium react with hydrogen? If so, what compound is produced?
3. Describe selenium’s purpose as a trace element.
4. What are some common uses for selenium?
5. Does selenium react with oxygen?
Solutions
1. Selenium’s ability to change light energy into electrical energy increases as light intensity increases.
2. Yes, selenium reacts with hydrogen and forms hydrogen selenide, H2Se.
3. Selenium is important to the health of plants and animals, but is only safe in small amounts. Too much selenium can be toxic and cause serious health problems.
4. Selenium is used in the glass-making industry and also in electronics. It is used in photo cells, solar cells, photocopiers, laser printers and also photographic toners.
5. Selenium burns in air and forms selenium dioxide. It is also able to form selenium trioxide.
Reference
1. Minaev, V. S., S. P. Timoshenkov, and V. V. Kalugin. "Structural and Phase Transformations in Condensed Selenium." Journal of Optoelectronics and Advanced Materials, volume 7, number 4, 2005, pp. 1717–1741.
2. Mary Elvira Weeks and Henry M. Leicester. Discovery of the Elements, 7th edition. Easton, PA: Journal of Chemical Education, 1968.
3. Petrucci, Ralph H. General Chemistry. 9th ed. Upper Saddle River: Prentice Hall, 2007
Contributors and Attributions
• David Jin (UCD) | textbooks/chem/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/Descriptive_Chemistry/Elements_Organized_by_Group/Group_16%3A_The_Oxygen_Family/Z034_Chemistry_of_Selenium_%28Z34%29.txt |
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