text:
Alkali metal
|Alkali metals
|↓ Period
Lithium (Li)
Sodium (Na)
Potassium (K)
Rubidium (Rb)
Caesium (Cs)
|Francium (Fr)
The alkali metals are a group in the periodic table consisting of the chemical elements lithium (Li), sodium (Na),[note 1] potassium (K),[note 2] rubidium (Rb), caesium (Cs),[note 3] and francium (Fr).[4] This group lies in the s-block of the periodic table[5] as all alkali metals have their outermost electron in an s-orbital.[6][7][8] The alkali metals provide the best example of group trends in properties in the periodic table,[6] with elements exhibiting well-characterized homologous behaviour.[6]
The alkali metals have very similar properties: they are all shiny, soft, highly reactive metals at standard temperature and pressure[6] and readily lose their outermost electron to form cations with charge +1.[9]:28 They can all be cut easily with a knife due to their softness, exposing a shiny surface that tarnishes rapidly in air due to oxidation.[6] Because of their high reactivity, they must be stored under oil to prevent reaction with air,[10] and are found naturally only in salts and never as the free element.[10] In the modern IUPAC nomenclature, the alkali metals comprise the group 1 elements,[note 4] excluding hydrogen (H), which is nominally a group 1 element[4][12] but not normally considered to be an alkali metal[13][14] as it rarely exhibits behaviour comparable to that of the alkali metals.[15] All the alkali metals react with water, with the heavier alkali metals reacting more vigorously than the lighter ones.[6][16]
All the discovered alkali metals occur in nature: in order of abundance, sodium is the most abundant, followed by potassium, lithium, rubidium, caesium, and finally francium, which is very rare due to its extremely high radioactivity and thus occurs only in traces due to its presence in natural decay chains.[17][18] Experiments have been conducted to attempt the synthesis of ununennium (Uue), which is likely to be the next member of the group, but they have all met with failure.[19] However, ununennium may not be an alkali metal due to relativistic effects, which are predicted to have a large influence on the chemical properties of superheavy elements;[20] even if it does turn out to be an alkali metal, it is predicted to have some differences in physical and chemical properties from its lighter homologues.[21]:1729–1733
Most alkali metals have many different applications. Two of the most well-known applications of the pure elements are rubidium and caesium atomic clocks,[22] of which caesium atomic clocks are the most accurate and precise representation of time.[23][24] A common application of the compounds of sodium is the sodium-vapour lamp, which emits very efficient light.[25][26] Table salt, or sodium chloride, has been used since antiquity. Sodium and potassium are also essential elements, having major biological roles as electrolytes,[27][28] and although the other alkali metals are not essential, they also have various effects on the body, both beneficial and harmful.[29][30][31][32]
- 1 Characteristics
- 1.1 Chemical
- 1.1.1 Compounds and reactions
- 126.96.36.199 Reaction with water (alkali metal hydroxides)
- 184.108.40.206 Reaction with the group 14 elements
- 22.214.171.124 Reaction with the pnictogens (alkali metal pnictides)
- 18.104.22.168 Reaction with the chalcogens (alkali metal chalcogenides)
- 220.127.116.11 Reaction with hydrogen and the halogens (alkali metal hydrides and halides)
- 184.108.40.206 Coordination complexes
- 220.127.116.11 Ammonia solutions
- 22.214.171.124 Organometallic chemistry
- 1.1.1 Compounds and reactions
- 1.2 Physical
- 1.3 Nuclear
- 1.1 Chemical
- 2 Extensions
- 3 Other similar substances
- 4 History
- 5 Occurrence
- 6 Production and isolation
- 7 Applications
- 8 Biological role and precautions
- 9 Notes
- 10 References
- 11 Further reading
- 12 External links
Like other groups, the known members of this family show patterns in electronic configuration, especially the outermost shells, resulting in trends in chemical behavior:
|No. of electrons/shell
configuration[note 5]
|2, 1
|[He] 2s1
|2, 8, 1
|[Ne] 3s1
|2, 8, 8, 1
|[Ar] 4s1
|2, 8, 18, 8, 1
|[Kr] 5s1
|2, 8, 18, 18, 8, 1
|[Xe] 6s1
|2, 8, 18, 32, 18, 8, 1
|[Rn] 7s1
Most of the chemistry has been observed only for the first five members of the group. The chemistry of francium is not well established due to its extreme radioactivity;[6] thus, the presentation of its properties here is limited.
All the alkali metals are highly reactive and are never found in elemental forms in nature.[33] Because of this, they are usually stored in mineral oil or kerosene (paraffin oil).[10] They react aggressively with the halogens to form the alkali metal halides, which are white ionic crystalline compounds that are all soluble in water except lithium fluoride (LiF).[6] The alkali metals also react with water to form strongly alkaline hydroxides and thus should be handled with great care. The heavier alkali metals react more vigorously than the lighter ones; for example, when dropped into water, caesium produces a larger explosion than potassium.[6][16][23] The alkali metals have the lowest first ionisation energies in their respective periods of the periodic table[8] because of their low effective nuclear charge[6] and the ability to attain a noble gas configuration by losing just one electron. The second ionisation energy of all of the alkali metals is very high[6][8] as it is in a full shell that is also closer to the nucleus;[6] thus, they almost always lose a single electron, forming cations.[9]:28 The alkalides are an exception: they are unstable compounds which contain alkali metals in a −1 oxidation state, which is very unusual as before the discovery of the alkalides, the alkali metals were not expected to be able to form anions and were thought to be able to appear in salts only as cations. The alkalide anions have filled s-subshells, which gives them more stability and allows them to exist. All the stable alkali metals except lithium are known to be able to form alkalides,[34][35][36] and the alkalides have much theoretical interest due to their unusual stoichiometry and low ionisation potentials. Alkalides are chemically similar to the electrides, which are salts with trapped electrons acting as anions.[37] A particularly striking example of an alkalide is "inverse sodium hydride", H+Na−, as opposed to the usual sodium hydride, Na+H−:[38] it is unstable in isolation, due to its high energy resulting from the displacement of two electrons from hydrogen to sodium, although several derivatives are predicted to be metastable or stable.[38][39]
The chemistry of lithium shows several differences from that of the rest of the group as the small Li+ cation polarises anions and gives its compounds a more covalent character.[6] Lithium and magnesium have a diagonal relationship:[6] because of this, lithium has some similarities to magnesium. For example, lithium forms a stable nitride, a property common among all the alkaline earth metals (magnesium's group) but unique among the alkali metals.[40] In addition, among their respective groups, only lithium and magnesium form covalent organometallic compounds (e.g. LiMe and MgMe2).[41] Lithium fluoride is the only alkali metal halide that is not soluble in water,[6] and lithium hydroxide is the only alkali metal hydroxide that is not deliquescent.[6] Francium is also predicted show some differences due to its high atomic weight, causing its electrons to travel at considerable fractions of the speed of light and thus making relativistic effects more prominent. In contrast to the trend of decreasing electronegativities and ionisation energies of the alkali metals, francium's electronegativity and ionisation energy are predicted to be higher than caesium's due to the relativistic stabilisation of the 7s electrons; also, its atomic radius is expected to be abnormally low.[21]:1729[42]
Compounds and reactions[edit]
Reaction with water (alkali metal hydroxides)[edit]
|Reactions of the alkali metals with water, conducted by The Open University
All the alkali metals react vigorously or explosively with cold water, producing an aqueous solution of the strongly basic alkali metal hydroxide and releasing hydrogen gas.[43] This reaction becomes more vigorous going down the group: lithium reacts steadily with effervescence, but sodium and potassium can ignite and rubidium and caesium sink in water and generate hydrogen gas so rapidly that shock waves form in the water that may shatter glass containers.[6] When an alkali metal is dropped into water, it produces an explosion, of which there are two separate stages. The metal reacts with the water first, breaking the hydrogen bonds in the water and producing hydrogen gas; this takes place faster for the more reactive heavier alkali metals. Second, the heat generated by the first part of the reaction often ignites the hydrogen gas, causing it to burn explosively into the surrounding air. This secondary hydrogen gas explosion produces the visible flame above the bowl of water, lake or other body of water, not the initial reaction of the metal with water (which tends to happen mostly under water).[16]
Reaction with the group 14 elements[edit]
Lithium and sodium react with carbon to form acetylides, Li2C2 and Na2C2, which can also be obtained by reaction of the metal with acetylene. Potassium, rubidium, and caesium react with graphite; their atoms are intercalated between the hexagonal graphite layers, forming graphite intercalation compounds of formulae MC60 (dark grey, almost black), MC48 (dark grey, almost black), MC36 (blue), MC24 (steel blue), and MC8 (bronze) (M = K, Rb, or Cs). These compounds are over 200 times more electrically conductive than pure graphite, suggesting that the valence electron of the alkali metal is transferred to the graphite layers (e.g. M+
8).[44] Upon heating of KC8, the elimination of potassium atoms results in the conversion in sequence to KC24, KC36, KC48 and finally KC60. KC8 is a very strong reducing agent and is pyrophoric and explodes on contact with water.[45][46] While the large alkali metals (K, Rb, and Cs) initially form MC8, the smaller ones initially form MC6.[47]
When the alkali metals react with the heavier elements in the carbon group, ionic substances with cage-like structures are formed, such as the silicide M4Si4 (M = K, Rb, or Cs), which contains M+ and tetrahedral Si4−
4 ions.[44] The chemistry of alkali metal germanides, involving the germanide ion Ge4− and other cluster (Zintl) ions such as Ge2−
4, Ge4−
9, Ge2−
9, and [(Ge9)2]6−, is largely analogous to that of the corresponding silicides.[9] Alkali metal stannides are mostly ionic, sometimes with the stannide ion (Sn4−),[48] and sometimes with more complex Zintl ions such as Sn4−
9, which appears in tetrapotassium nonastannide (K4Sn9).[49] The monatomic plumbide ion (Pb4−) is unknown, and indeed its formation is predicted to be energetically unfavourable; alkali metal plumbides have complex Zintl ions, such as Pb4−
Reaction with the pnictogens (alkali metal pnictides)[edit]
Lithium, the lightest of the alkali metals, is the only alkali metal which reacts with nitrogen at standard conditions, and its nitride is the only stable alkali metal nitride. Nitrogen is an unreactive gas because breaking the strong triple bond in the dinitrogen molecule (N2) requires a lot of energy. The formation of an alkali metal nitride would consume the ionisation energy of the alkali metal (forming M+ ions), the energy required to break the triple bond in N2 and the formation of N3− ions, and all the energy released from the formation of an alkali metal nitride is from the lattice energy of the alkali metal nitride. The lattice energy is maximised with small, highly charged ions; the alkali metals do not form highly charged ions, only forming ions with a charge of +1, so only lithium, the smallest alkali metal, can release enough lattice energy to make the reaction with nitrogen exothermic, forming lithium nitride. The reactions of the other alkali metals with nitrogen would not release enough lattice energy and would thus be endothermic, so they do not form nitrides at standard conditions.[40] (Sodium nitride (Na3N) and potassium nitride (K3N), while existing, are extremely unstable, being prone to decomposing back into their constituent elements, and cannot be produced by reacting the elements with each other at standard conditions.)[51][52]
All the alkali metals react readily with phosphorus and arsenic to form phosphides and arsenides with the formula M3Pn (where M represents an alkali metal and Pn represents a pnictogen). This is due to the greater size of the P3− and As3− ions, so that less lattice energy needs to be released for the salts to form.[44] These are not the only phosphides and arsenides of the alkali metals: for example, potassium has nine different known phosphides, with formulae K3P, K4P3, K5P4, KP, K4P6, K3P7, K3P11, KP10.3, and KP15.[53] While most metals form arsenides, only the alkali and alkaline earth metals form mostly ionic arsenides. The structure of Na3As is complex with unusually short Na–Na distances of 328–330 pm which are shorter than in sodium metal, and this indicates that even with these electropositive metals the bonding cannot be straightforwardly ionic.[9] Other alkali metal arsenides not conforming to the formula M3As are known, such as LiAs, which has a metallic lustre and electrical conductivity indicating the presence of some metallic bonding.[9] The antimonides are unstable and reactive as the Sb3− ion is a strong reducing agent; reaction of them with acids form the toxic and unstable gas stibine (SbH3).[54] Bismuthides are not even wholly ionic; they are intermetallic compounds containing partially metallic and partially ionic bonds.[55]
Reaction with the chalcogens (alkali metal chalcogenides)[edit]
All the alkali metals react vigorously with oxygen at standard conditions. They form various types of oxides, such as simple oxides (containing the O2− ion), peroxides (containing the O2−
2 ion, where there is a single bond between the two oxygen atoms), superoxides (containing the O−
2 ion), and many others. Lithium burns in air to form lithium oxide, but sodium reacts with oxygen to form a mixture of sodium oxide and sodium peroxide. Potassium forms a mixture of potassium peroxide and potassium superoxide, while rubidium and caesium form the superoxide exclusively. Their reactivity increases going down the group: while lithium, sodium and potassium merely burn in air, rubidium and caesium are pyrophoric (spontaneously catch fire in air).[40]
The smaller alkali metals tend to polarise the more complex anions (the peroxide and superoxide) due to their small size. This attracts the electrons in the more complex anions towards one of its constituent oxygen atoms, forming an oxide ion and an oxygen atom. This causes lithium to form the oxide exclusively on reaction with oxygen at room temperature. This effect becomes drastically weaker for the larger sodium and potassium, allowing them to form the less stable peroxides. Rubidium and caesium, at the bottom of the group, are so large that even the least stable superoxides can form. Because the superoxide releases the most energy when formed, the superoxide is preferentially formed for the larger alkali metals where the more complex anions are not polarised. (The oxides and peroxides for these alkali metals do exist, but do not form upon direct reaction of the metal with oxygen at standard conditions.)[40] In addition, the small size of the Li+ and O2− ions contributes to their forming a stable ionic lattice structure. Under controlled conditions, however, all the alkali metals, with the exception of francium, are known to form their oxides, peroxides, and superoxides. The alkali metal peroxides and superoxides are powerful oxidizing agents. Sodium peroxide and potassium superoxide react with carbon dioxide to form the alkali metal carbonate and oxygen gas, which allows them to be used in submarine air purifiers; the presence of water vapour, naturally present in breath, makes the removal of carbon dioxide by potassium superoxide even more efficient.[44][56]
Rubidium and caesium can form even more complicated oxides than the superoxides. Rubidium can form Rb6O and Rb9O2 upon oxidation in air, while caesium forms an immense variety of oxides, such as the ozonide CsO3[57][58] and several brightly coloured suboxides,[59] such as Cs
7O, Cs
4O, Cs
3, Cs
3O (dark-green[60]), CsO, Cs
2,[61] as well as Cs
2.[62][63] The latter may be heated under vacuum to generate Cs
The alkali metals can also react analogously with the heavier chalcogens (sulfur, selenium, tellurium, and polonium), and all the alkali metal chalcogenides are known (with the exception of francium's). Reaction with an excess of the chalcogen can similarly result in lower chalcogenides, with chalcogen ions containing chains of the chalcogen atoms in question. For example, sodium can react with sulfur to form the sulfide (Na2S) and various polysulfides with the formula Na2Sx (x from 2 to 6), containing the S2−
x ions.[44] Due to the basicity of the Se2− and Te2− ions, the alkali metal selenides and tellurides are alkaline in solution; when reacted directly with selenium and tellurium, alkali metal polyselenides and polytellurides are formed along with the selenides and tellurides with the Se2−
x and Te2−
x ions.[65] The alkali metal polonides are all ionic compounds containing the Po2− ion; they are very chemically stable and can be produced by direct reaction of the elements at around 300–400 °C.[9][66][67]
Reaction with hydrogen and the halogens (alkali metal hydrides and halides)[edit]
The alkali metals are among the most electropositive elements on the periodic table and thus tend to bond ionically to the most electronegative elements on the periodic table, the halogens, forming salts known as the alkali metal halides. This includes sodium chloride, otherwise known as common salt. The reactivity becomes higher from lithium to caesium and drops from fluorine to iodine. All of the alkali metal halides have the formula MX where M is an alkali metal and X is a halogen. They are all white ionic crystalline solids.[6][40] All the alkali metal halides are soluble in water except for lithium fluoride (LiF), which is insoluble in water due to its very high lattice enthalpy. The high lattice enthalpy of lithium fluoride is due to the small sizes of the Li+ and F− ions, causing the electrostatic interactions between them to be strong.[6] The alkali metals also react similarly with hydrogen to form ionic alkali metal hydrides.[44]
Coordination complexes[edit]
Alkali metal cations do not usually form coordination complexes with simple Lewis bases due to their low charge of just +1 and their relatively large size; thus the Li+ ion forms most complexes and the heavier alkali metal ions form less and less. In aqueous solution, the alkali metal ions exist as octahedral hexahydrate complexes ([M(H2O)6)]+), with the exception of the lithium ion, which due to its small size forms tetrahedral tetrahydrate complexes ([Li(H2O)4)]+); the alkali metals form these complexes because their ions are attracted by electrostatic forces of attraction to the polar water molecules. Because of this, anhydrous salts containing alkali metal cations are often used as desiccants.[44] Alkali metals also readily form complexes with crown ethers (e.g. 12-crown-4 for Li+, 15-crown-5 for Na+, and 18-crown-6 for K+) and cryptands due to electrostatic attraction.[44]
Ammonia solutions[edit]
Unlike most metals, the alkali metals dissolve slowly in liquid ammonia, forming hydrogen gas and the alkali metal amide (MNH2, where M represents an alkali metal). The process may be speeded up by a catalyst. The amide salt is quite insoluble and readily precipitates out of solution, leaving intensely coloured ammonia solutions of the alkali metals. The colour is due to the presence of solvated electrons, which contribute to the high electrical conductivity of these solutions. At low concentrations (below 3 M), the solution is dark blue and has ten times the conductivity of aqueous sodium chloride; at higher concentrations (above 3 M), the solution is copper-coloured and has approximately the conductivity of liquid metals like mercury.[9][44][69] In addition to the alkali metal amide salt and solvated electrons, such ammonia solutions also contain the alkali metal cation (M+), the neutral alkali metal atom (M), diatomic alkali metal molecules (M2) and alkali metal anions (M−). These are unstable and eventually become the more thermodynamically stable alkali metal amide and hydrogen gas. Solvated electrons are powerful reducing agents and are often used in chemical synthesis.[44]
Organometallic chemistry[edit]
Being the smallest alkali metal, lithium forms the widest variety of and most stable organometallic compounds, which are bonded covalently. Organolithium compounds are electrically non-conducting volatile solids or liquids that melt at low temperatures, and tend to form oligomers with the structure (RLi)x where R is the organic group. As the electropositive nature of lithium puts most of the charge density of the bond on the carbon atom, effectively creating a carbanion, organolithium compounds are extremely powerful bases and nucleophiles. For use as bases, butyllithiums are often used and are commercially available. An example of an organolithium compound is methyllithium ((CH3Li)x), which exists in tetrameric (x = 4) and hexameric (x = 6) forms.[44][70]
The application of organosodium compounds in chemistry is limited in part due to competition from organolithium compounds, which are commercially available and exhibit more convenient reactivity. The principal organosodium compound of commercial importance is sodium cyclopentadienide. Sodium tetraphenylborate can also be classified as an organosodium compound since in the solid state sodium is bound to the aryl groups. Organometallic compounds of the higher alkali metals are even more reactive than organosodium compounds and of limited utility. A notable reagent is Schlosser's base, a mixture of n-butyllithium and potassium tert-butoxide. This reagent reacts with propene to form the compound allylpotassium (KCH2CHCH2). cis-2-Butene and trans-2-butene equilibrate when in contact with alkali metals. Whereas isomerization is fast with lithium and sodium, it is slow with the higher alkali metals. The higher alkali metals also favor the sterically congested conformation.[71] Several crystal structures of organopotassium compounds have been reported, establishing that they, like the sodium compounds, are polymeric.[72] Organosodium, organopotassium, organorubidium and organocaesium compounds are all mostly ionic and are insoluble (or nearly so) in nonpolar solvents.[44]
The alkali metals are all silver-coloured except for caesium, which has a golden tint.[73] All are soft and have low densities,[6] melting points,[6] and boiling points.[6]
The table below is a summary of the key physical and atomic properties of the alkali metals. Data marked with question marks are either uncertain or are estimations partially based on periodic trends rather than observations.
|Alkali metal
|Standard atomic weight
(u)[note 6][75][76]
|Melting point
|Boiling point[8]
|First ionisation energy
|Atomic radius
|Flame test colour
|6.94(1)[note 7]
|Strong persistent orange or yellow[6][77]
|Lilac or pink[6][77]
|Red or reddish-violet[6][77]
|Blue or violet[6][77]
|[223][note 8]
|? 1.87
|? 0.7
Periodic trends[edit]
The alkali metals are more similar to each other than the elements in any other group are to each other.[6] For instance, when moving down the table, all known alkali metals show increasing atomic radius,[79] decreasing electronegativity,[79] increasing reactivity,[6] and decreasing melting and boiling points.[79] In general, their densities increase when moving down the table, with the exception that potassium is less dense than sodium.[79]
Atomic and ionic radii[edit]
The atomic radii of the alkali metals increase going down the group.[79] Because of the shielding effect, when an atom has more than one electron shell, each electron feels electric repulsion from the other electrons as well as electric attraction from the nucleus.[80] In the alkali metals, the outermost electron only feels a net charge of +1, as some of the nuclear charge (which is equal to the atomic number) is cancelled by the inner electrons; the number of inner electrons of an alkali metal is always one less than the nuclear charge. Therefore, the only factor which affects the atomic radius of the alkali metals is the number of electron shells. Since this number increases down the group, the atomic radius must also increase down the group.[79]
The ionic radii of the alkali metals are much smaller than their atomic radii. This is because the outermost electron of the alkali metals is in a different electron shell than the inner electrons, and thus when it is removed the resulting atom has one fewer electron shell and is smaller. Additionally, the effective nuclear charge has increased, and thus the electrons are attracted more strongly towards the nucleus and the ionic radius decreases.[6]
First ionisation energy[edit]
The first ionisation energy of an element or molecule is the energy required to move the most loosely held electron from one mole of gaseous atoms of the element or molecules to form one mole of gaseous ions with electric charge +1. The factors affecting the first ionisation energy are the nuclear charge, the amount of shielding by the inner electrons and the distance from the most loosely held electron from the nucleus, which is always an outer electron in main group elements. The first two factors change the effective nuclear charge the most loosely held electron feels. Since the outermost electron of alkali metals always feels the same effective nuclear charge (+1), the only factor which affects the first ionisation energy is the distance from the outermost electron to the nucleus. Since this distance increases down the group, the outermost electron feels less attraction from the nucleus and thus the first ionisation energy decreases.[79] (This trend is broken in francium due to the relativistic stabilization and contraction of the 7s orbital, bringing francium's valence electron closer to the nucleus than would be expected from non-relativistic calculations. This makes francium's outermost electron feel more attraction from the nucleus, increasing its first ionisation energy slightly beyond that of caesium.)[21]:1729
The second ionisation energy of the alkali metals is much higher than the first as the second-most loosely held electron is part of a fully filled electron shell and is thus difficult to remove.[6]
The reactivities of the alkali metals increase going down the group. This is the result of a combination of two factors: the first ionisation energies and atomisation energies of the alkali metals. Because the first ionisation energy of the alkali metals decreases down the group, it is easier for the outermost electron to be removed from the atom and participate in chemical reactions, thus increasing reactivity down the group. The atomisation energy measures the strength of the metallic bond of an element, which falls down the group as the atoms increase in radius and thus the metallic bond must increase in length, making the delocalised electrons further away from the attraction of the nuclei of the heavier alkali metals. Adding the atomisation and first ionisation energies gives a quantity closely related to (but not equal to) the activation energy of the reaction of an alkali metal with another substance. This quantity decreases going down the group, and so does the activation energy; thus, chemical reactions can occur faster and the reactivity increases down the group.[43]
Electronegativity is a chemical property that describes the tendency of an atom or a functional group to attract electrons (or electron density) towards itself.[85] If the bond between sodium and chlorine in sodium chloride were covalent, the pair of shared electrons would be attracted to the chlorine because the effective nuclear charge on the outer electrons is +7 in chlorine but is only +1 in sodium. The electron pair is attracted so close to the chlorine atom that they are practically transferred to the chlorine atom (an ionic bond). However, if the sodium atom was replaced by a lithium atom, the electrons will not be attracted as close to the chlorine atom as before because the lithium atom is smaller, making the electron pair more strongly attracted to the closer effective nuclear charge from lithium. Hence, the larger alkali metal atoms (further down the group) will be less electronegative as the bonding pair is less strongly attracted towards them.[79]
Because of the higher electronegativity of lithium, some of its compounds have a more covalent character. For example, lithium iodide (LiI) will dissolve in organic solvents, a property of most covalent compounds.[79] Lithium fluoride (LiF) is the only alkali halide that is not soluble in water,[6] and lithium hydroxide (LiOH) is the only alkali metal hydroxide that is not deliquescent.[6]
Melting and boiling points[edit]
Melting and boiling points of the alkali metals[8][note 14] Alkali metal Melting point Boiling point[8] Lithium 453.69 K (180.54 °C) 1615 K (1342 °C) Sodium 370.87 K (97.72 °C) 1156 K (883 °C) Potassium 336.53 K (63.38 °C) 1032 K (759 °C) Rubidium 312.46 K (39.31 °C) 961 K (688 °C) Caesium 301.59 K (28.44 °C) 944 K (671 °C) Francium ? 300 K (? 27 °C)[note 15] ? 950 K (? 677 °C)[78][note 15]
The melting point of a substance is the point where it changes state from solid to liquid while the boiling point of a substance (in liquid state) is the point where the vapor pressure of the liquid equals the environmental pressure surrounding the liquid[87][88] and all the liquid changes state to gas. As a metal is heated to its melting point, the metallic bonds keeping the atoms in place weaken so that the atoms can move around, and the metallic bonds eventually break completely at the metal's boiling point.[79][89] Therefore, the falling melting and boiling points of the alkali metals indicate that the strength of the metallic bonds of the alkali metals decreases down the group.[79] This is because metal atoms are held together by the electromagnetic attraction from the positive ions to the delocalised electrons.[79][89] As the atoms increase in size going down the group (because their atomic radius increases), the nuclei of the ions move further away from the delocalised electrons and hence the metallic bond becomes weaker so that the metal can more easily melt and boil, thus lowering the melting and boiling points.[79] (The increased nuclear charge is not a relevant factor due to the shielding effect.)[79]
The alkali metals all have the same crystal structure (body-centred cubic)[9] and thus the only relevant factors are the number of atoms that can fit into a certain volume and the mass of one of the atoms, since density is defined as mass per unit volume. The first factor depends on the volume of the atom and thus the atomic radius, which increases going down the group; thus, the volume of an alkali metal atom increases going down the group. The mass of an alkali metal atom also increases going down the group. Thus, the trend for the densities of the alkali metals depends on their atomic weights and atomic radii; if figures for these two factors are known, the ratios between the densities of the alkali metals can then be calculated. The resultant trend is that the densities of the alkali metals increase down the table, with an exception at potassium. Due to having the lowest atomic weight of all the elements in their period and having the largest atomic radius for their periods, the alkali metals are the least dense metals in the periodic table.[79] Lithium, sodium, and potassium are the only three metals in the periodic table that are less dense than water.[6]
|Alkali metal
|unstable: italics
odd-odd isotopes coloured pink
|No primordial isotopes
All the alkali metals have odd atomic numbers; hence, their isotopes must be either odd-odd (both proton and neutron number are odd) or odd-even (proton number is odd, but neutron number is even). Odd-odd nuclei have even mass numbers, while odd-even nuclei have odd mass numbers. Odd-odd primordial nuclides are rare because most odd-odd nuclei are highly unstable with respect to beta decay, because the decay products are even-even, and are therefore more strongly bound, due to nuclear pairing effects.[90]
Due to the great rarity of odd-odd nuclei, almost all the primordial isotopes of the alkali metals are odd-even (the exceptions being the light stable isotope lithium-6 and the long-lived radioisotope potassium-40). For a given odd mass number, there can be only a single beta-stable nuclide, since there is not a difference in binding energy between even-odd and odd-even comparable to that between even-even and odd-odd, leaving other nuclides of the same mass number (isobars) free to beta decay toward the lowest-mass nuclide. An effect of the instability of an odd number of either type of nucleons is that odd-numbered elements, such as the alkali metals, tend to have fewer stable isotopes than even-numbered elements. Of the 26 monoisotopic elements that have only a single stable isotope, all but one have an odd atomic number and all but one also have an even number of neutrons. Beryllium is the single exception to both rules, due to its low atomic number.[90]
All of the alkali metals except lithium and caesium have at least one naturally occurring radioisotope: sodium-22 and sodium-24 are trace radioisotopes produced cosmogenically,[91] potassium-40 and rubidium-87 have very long half-lives and thus occur naturally,[92] and all isotopes of francium are radioactive.[92] Caesium was also thought to be radioactive in the early 20th century,[93][94] although it has no naturally occurring radioisotopes.[92] (Francium had not been discovered yet at that time.) The natural radioisotope of potassium, potassium-40, makes up about 0.012% of natural potassium,[95] and thus natural potassium is weakly radioactive. This natural radioactivity became a basis for a mistaken claim of the discovery for element 87 (the next alkali metal after caesium) in 1925.[96][97]
Caesium-137, with a half-life of 30.17 years, is one of the two principal medium-lived fission products, along with strontium-90, which are responsible for most of the radioactivity of spent nuclear fuel after several years of cooling, up to several hundred years after use. It constitutes most of the radioactivity still left from the Chernobyl accident. 137Cs undergoes high-energy beta decay and eventually becomes stable barium-137. It is a strong emitter of gamma radiation. 137Cs has a very low rate of neutron capture and cannot be feasibly disposed of in this way, but must be allowed to decay.[98] 137Cs has been used as a tracer in hydrologic studies, analogous to the use of tritium.[99] Small amounts of caesium-134 and caesium-137 were released into the environment during nearly all nuclear weapon tests and some nuclear accidents, most notably the Goiânia accident and the Chernobyl disaster. As of 2005, caesium-137 is the principal source of radiation in the zone of alienation around the Chernobyl nuclear power plant.[100]
Although francium is the heaviest alkali metal that has been discovered, there has been some theoretical work predicting the physical and chemical characteristics of the hypothetical heavier alkali metals. Being the first period 8 element, the undiscovered element ununennium (element 119) is predicted to be the next alkali metal after francium and behave much like their lighter congeners; however, it is also predicted to differ from the lighter alkali metals in some properties.[21]:1729–1730 Its chemistry is predicted to be closer to that of potassium[102] or rubidium[21]:1729–1730 instead of caesium or francium. This is unusual as periodic trends, ignoring relativistic effects would predict ununennium to be even more reactive than caesium and francium. This lowered reactivity is due to the relativistic stabilisation of ununennium's valence electron, increasing ununennium's first ionisation energy and decreasing the metallic and ionic radii;[102] this effect is already seen for francium.[21]:1729–1730 This assumes that ununennium will behave chemically as an alkali metal, which, although likely, may not be true due to relativistic effects.[20] The relativistic stabilisation of the 8s orbital also increases ununennium's electron affinity far beyond that of caesium and francium; indeed, ununennium is expected to have an electron affinity higher than all the alkali metals lighter than it. Relativistic effects also cause a very large drop in the polarisability of ununennium.[21]:1729–1730 On the other hand, ununennium is predicted to continue the trend of melting points decreasing going down the group, being expected to have a melting point between 0 °C and 30 °C.[21]:1724
The stabilisation of ununennium's valence electron and thus the contraction of the 8s orbital cause its atomic radius to be lowered to 240 pm,[21]:1729–1730 very close to that of rubidium (247 pm),[6] so that the chemistry of ununennium in the +1 oxidation state should be more similar to the chemistry of rubidium than to that of francium. On the other hand, the ionic radius of the Uue+ ion is predicted to be larger than that of Rb+, because the 7p orbitals are destabilised and are thus larger than the p-orbitals of the lower shells. Ununennium may also show the +3 oxidation state,[21]:1729–1730 which is not seen in any other alkali metal,[9]:28 in addition to the +1 oxidation state that is characteristic of the other alkali metals and is also the main oxidation state of all the known alkali metals: this is because of the destabilisation and expansion of the 7p3/2 spinor, causing its outermost electrons to have a lower ionisation energy than what would otherwise be expected.[9]:28[21]:1729–1730 Indeed, many ununennium compounds are expected to have a large covalent character, due to the involvement of the 7p3/2 electrons in the bonding.[103]
Not as much work has been done predicting the properties of the alkali metals beyond ununennium. Although a simple extrapolation of the periodic table would put element 169, unhexennium, under ununennium, Dirac-Fock calculations predict that the next alkali metal after ununennium may actually be element 165, unhexpentium, which is predicted to have the electron configuration [Uuo] 5g18 6f14 7d10 8s2 8p1/22 9s1.[21]:1729–1730[101] Further calculations show that unhexpentium would follow the trend of increasing ionisation energy beyond caesium, having an ionisation energy comparable to that of sodium, and that it should also continue the trend of decreasing atomic radii beyond caesium, having an atomic radius comparable to that of potassium.[21]:1729–1730 However, the 7d electrons of unhexpentium may also be able to participate in chemical reactions along with the 9s electron, possibly allowing oxidation states beyond +1 and perhaps even making unhexpentium behave more like a boron group element than an alkali metal.[21]:1732–1733
The probable properties of the alkali metals beyond unhexpentium have not been explored yet as of 2012. In periods 8 and above of the periodic table, relativistic and shell-structure effects become so strong that extrapolations from lighter congeners become completely inaccurate. In addition, the relativistic and shell-structure effects (which stabilise the s-orbitals and destabilise and expand the d-, f-, and g-orbitals of higher shells) have opposite effects, causing even larger difference between relativistic and non-relativistic calculations of the properties of elements with such high atomic numbers.[21]:1732–1733 Due to the alkali and alkaline earth metals both being s-block elements, these predictions for the trends and properties of ununennium and unhexpentium also mostly hold quite similarly to the corresponding alkaline earth metals unbinilium (Ubn) and unhexhexium (Uhh).[21]:1729–1733
Other similar substances[edit]
The element hydrogen, with one electron per neutral atom, is usually placed at the top of Group 1 of the periodic table for convenience, but hydrogen is not normally considered to be an alkali metal;[13] when it is considered to be an alkali metal, it is because of its atomic properties and not its chemical properties.[14] Under typical conditions, pure hydrogen exists as a diatomic gas consisting of two atoms per molecule (H2);[104] however, the alkali metals only form diatomic molecules (such as dilithium, Li2) at high temperatures, when they are in the gaseous state.[105]
Hydrogen, like the alkali metals, has one valence electron[15] and reacts easily with the halogens,[15] but the similarities end there.[15] Its placement above lithium is primarily due to its electron configuration and not its chemical properties.[13][15] It is sometimes placed above carbon due to their similar electronegativities[106] or fluorine due to their similar chemical properties.[15][106]
The first ionisation energy of hydrogen (1312.0 kJ/mol) is much higher than that of the alkali metals.[81][82] As only one additional electron is required to fill in the outermost shell of the hydrogen atom, hydrogen often behaves like a halogen, forming the negative hydride ion, and is sometimes considered to be a halogen.[15] (The alkali metals can also form negative ions, known as alkalides, but these are little more than laboratory curiosities, being unstable.)[38][39] Under extremely high pressures, such as those found at the cores of Jupiter and Saturn, hydrogen does become metallic and behaves like an alkali metal; in this phase, it is known as metallic hydrogen.[107]
The ammonium ion (NH+
4) has very similar properties to the heavier alkali metals, acting as an alkali metal intermediate between potassium and rubidium,[108] and is often considered a close relative.[109][110][111] For example, most alkali metal salts are soluble in water, a property which ammonium salts share.[112] Ammonium is expected to behave stably as a metal (NH+
4 ions in a sea of electrons) at very high pressures (though less than the typical pressure where transitions from insulating to metallic behaviour occur around, 100 GPa), and could possibly occur inside the ice giants Uranus and Neptune, which may have significant impacts on their interior magnetic fields.[110][111] It has been estimated that the transition from a mixture of ammonia and dihydrogen molecules to metallic ammonium may occur at pressures just below 25 GPa.[110]
Thallium displays the +1 oxidation state[9]:28 that all the known alkali metals display,[9]:28 and thallium compounds with thallium in its +1 oxidation state closely resemble the corresponding potassium or silver compounds due to the similar ionic radii of the Tl+ (164 pm), K+ (152 pm) and Ag+ (129 pm) ions.[113][114] It was sometimes considered an alkali metal in continental Europe (but not in England) in the years immediately following its discovery,[114]:126 and was placed just after caesium as the sixth alkali metal in Dmitri Mendeleev's 1869 periodic table and Julius Lothar Meyer's 1868 periodic table.[115] (Mendeleev's 1871 periodic table and Meyer's 1870 periodic table put thallium in its current position in the boron group and leave the space below caesium blank.)[115] However, thallium also displays the oxidation state +3,[9]:28 which no known alkali metal displays[9]:28 (although ununennium, the undiscovered seventh alkali metal, is predicted to possibly display the +3 oxidation state).[21]:1729–1730 The sixth alkali metal is now considered to be francium.[4]
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The alkali metals are so called because their hydroxides are all strong alkalis when dissolved in water.[6]
Petalite (LiAlSi4O10) was discovered in 1800 by the Brazilian chemist José Bonifácio de Andrada in a mine on the island of Utö, Sweden.[116][117][118] However, it was not until 1817 that Johan August Arfwedson, then working in the laboratory of the chemist Jöns Jacob Berzelius, detected the presence of a new element while analyzing petalite ore.[119][120] This new element formed compounds similar to those of sodium and potassium, though its carbonate and hydroxide were less soluble in water and more alkaline than the other alkali metals.[121] Berzelius gave the unknown material the name "lithion/lithina", from the Greek word λιθoς (transliterated as lithos, meaning "stone"), to reflect its discovery in a solid mineral, as opposed to potassium, which had been discovered in plant ashes, and sodium, which was known partly for its high abundance in animal blood. He named the metal inside the material "lithium".[33][117][120]
Sodium compounds have been known since ancient times; salt (sodium chloride) has been an important commodity in human activities, as testified by the English word salary, referring to salarium, the wafers of salt sometimes given to Roman soldiers along with their other wages.[citation needed] In medieval Europe a compound of sodium[clarification needed] with the Latin name of sodanum was used as a headache remedy.[citation needed] Pure sodium was not isolated until 1807 by Humphry Davy through the electrolysis of caustic soda (now called sodium hydroxide),[122] a very similar method to the one used to isolate potassium earlier that year.
While potash has been used since ancient times, it was not understood for most of its history to be a fundamentally different substance from sodium mineral salts. Georg Ernst Stahl obtained experimental evidence which led him to suggest the fundamental difference of sodium and potassium salts in 1702,[123] and Henri Louis Duhamel du Monceau was able to prove this difference in 1736.[124] The exact chemical composition of potassium and sodium compounds, and the status as chemical element of potassium and sodium, was not known then, and thus Antoine Lavoisier did include the alkali in his list of chemical elements in 1789.[125][126] Pure potassium was first isolated in 1807 in England by Sir Humphry Davy, who derived it from caustic potash (KOH, potassium hydroxide) by the use of electrolysis of the molten salt with the newly invented voltaic pile. Potassium was the first metal that was isolated by electrolysis.[127] Later that same year, Davy reported extraction of sodium from the similar substance caustic soda (NaOH, lye) by a similar technique, demonstrating the elements, and thus the salts, to be different.[122][125][126][128]
Rubidium was discovered in 1861 in Heidelberg, Germany by Robert Bunsen and Gustav Kirchhoff, the first people to suggest finding new elements by spectrum analysis, in the mineral lepidolite through the use of a spectroscope. Because of the bright red lines in its emission spectrum, they chose a name derived from the Latin word rubidus, meaning dark red or bright red.[129][130] Rubidium's discovery succeeded that of caesium, also discovered by Bunsen and Kirchhoff through spectroscopy.[131]
In 1860, Robert Bunsen and Gustav Kirchhoff discovered caesium in the mineral water from Dürkheim, Germany. Due to the bright-blue lines in its emission spectrum, they chose a name derived from the Latin word caesius, meaning sky-blue.[129][note 17][132] Caesium was the first element to be discovered spectroscopically, only one year after the invention of the spectroscope by Bunsen and Kirchhoff.[131]
There were at least four erroneous and incomplete discoveries[96][97][133][134] before Marguerite Perey of the Curie Institute in Paris, France discovered francium in 1939 by purifying a sample of actinium-227, which had been reported to have a decay energy of 220 keV. However, Perey noticed decay particles with an energy level below 80 keV. Perey thought this decay activity might have been caused by a previously unidentified decay product, one that was separated during purification, but emerged again out of the pure actinium-227. Various tests eliminated the possibility of the unknown element being thorium, radium, lead, bismuth, or thallium. The new product exhibited chemical properties of an alkali metal (such as coprecipitating with caesium salts), which led Perey to believe that it was element 87, caused by the alpha decay of actinium-227.[18] Perey then attempted to determine the proportion of beta decay to alpha decay in actinium-227. Her first test put the alpha branching at 0.6%, a figure that she later revised to 1%.[135] It was the last element discovered in nature, rather than by synthesis.[note 18]
The next element below francium (eka-francium) is very likely to be ununennium (Uue), element 119,[21]:1729–1730 although this is not completely certain due to relativistic effects.[20] The synthesis of ununennium was first attempted in 1985 by bombarding a target of einsteinium-254 with calcium-48 ions at the superHILAC accelerator at Berkeley, California. No atoms were identified, leading to a limiting yield of 300 nb.[19][136]
It is highly unlikely[19] that this reaction will be able to create any atoms of ununennium in the near future, given the extremely difficult task of making sufficient amounts of 254Es, which is favoured for production of ultraheavy elements because of its large mass, relatively long half-life of 270 days, and availability in significant amounts of several micrograms,[137] to make a large enough target to increase the sensitivity of the experiment to the required level; einsteinium has not been found in nature and has only been produced in laboratories. However, given that ununennium is only the first period 8 element on the extended periodic table, it may well be discovered in the near future through other reactions; indeed, another attempt to synthesise ununennium by bombarding a berkelium target with titanium ions is under way at the GSI Helmholtz Centre for Heavy Ion Research in Darmstadt, Germany.[138] Currently, none of the period 8 elements have been discovered yet, and it is also possible, due to drip instabilities, that only the lower period 8 elements, up to around element 128, are physically possible.[102][139] No attempts at synthesis have been made for any heavier alkali metals, such as unhexpentium, due to their extremely high atomic number.[21]:1737–1739
In the Solar System[edit]
The Oddo-Harkins rule holds that elements with even atomic numbers are more common that those with odd atomic numbers, with the exception of hydrogen. This rule argues that elements with odd atomic numbers have one unpaired proton and are more likely to capture another, thus increasing their atomic number. In elements with even atomic numbers, protons are paired, with each member of the pair offsetting the spin of the other, enhancing stability.[141][142][143] All the alkali metals have odd atomic numbers and they are not as common as the elements with even atomic numbers adjacent to them (the noble gases and the alkaline earth metals) in the Solar System. The heavier alkali metals are also less abundant than the lighter ones as the alkali metals from rubidium onward can only be synthesized in supernovae and not in stellar nucleosynthesis. Lithium is also much less abundant than sodium and potassium as it is poorly synthesized in both Big Bang nucleosynthesis and in stars: the Big Bang could only produce trace quantities of lithium, beryllium and boron due to the absence of a stable nucleus with 5 or 8 nucleons, and stellar nucleosynthesis could only pass this bottleneck by the triple-alpha process, fusing three helium nuclei to form carbon, and skipping over those three elements.[140]
On Earth[edit]
The Earth formed from the same cloud of matter that formed the Sun, but the planets acquired different compositions during the formation and evolution of the solar system. In turn, the natural history of the Earth caused parts of this planet to have differing concentrations of the elements. The mass of the Earth is approximately 5.98×1024 kg. It is composed mostly of iron (32.1%), oxygen (30.1%), silicon (15.1%), magnesium (13.9%), sulfur (2.9%), nickel (1.8%), calcium (1.5%), and aluminium (1.4%); with the remaining 1.2% consisting of trace amounts of other elements. Due to mass segregation, the core region is believed to be primarily composed of iron (88.8%), with smaller amounts of nickel (5.8%), sulfur (4.5%), and less than 1% trace elements.[144]
The alkali metals, due to their high reactivity, do not occur naturally in pure form in nature. They are lithophiles and therefore remain close to the Earth's surface because they combine readily with oxygen and so associate strongly with silica, forming relatively low-density minerals that do not sink down into the Earth's core. Potassium, rubidium and caesium are also incompatible elements due to their low ionic radii.[145]
Sodium and potassium are very abundant in earth, both being among the ten most common elements in Earth's crust;[17][146] sodium makes up approximately 2.6% of the Earth's crust measured by weight, making it the sixth most abundant element overall[147] and the most abundant alkali metal. Potassium makes up approximately 1.5% of the Earth's crust and is the seventh most abundant element.[147] Sodium is found in many different minerals, of which the most common is ordinary salt (sodium chloride), which occurs in vast quantities dissolved in seawater. Other solid deposits include halite, amphibole, cryolite, nitratine, and zeolite.[147]
Lithium, due to its relatively low reactivity, can be found in seawater in large amounts; it is estimated that seawater is approximately 0.14 to 0.25 parts per million (ppm)[148][149] or 25 micromolar.[150]
Rubidium is approximately as abundant as zinc and more abundant than copper. It occurs naturally in the minerals leucite, pollucite, carnallite, zinnwaldite, and lepidolite.[151] Caesium is more abundant than some commonly known elements, such as antimony, cadmium, tin, and tungsten, but is much less abundant than rubidium.[23]
Francium-223, the only naturally occurring isotope of francium,[75][76] is the product of the alpha decay of actinium-227 and can be found in trace amounts in uranium and thorium minerals.[152] In a given sample of uranium, there is estimated to be only one francium atom for every 1018 uranium atoms.[153][154] It has been calculated that there is at most 30 g of francium in the earth's crust at any time, due to its extremely short half-life of 22 minutes.[155][156]
Production and isolation[edit]
The production of pure alkali metals is difficult due to their extreme reactivity with commonly used substances, such as water. The alkali metals are so reactive that they cannot be displaced by other elements and must be isolated through high-energy methods such as electrolysis.[6][44]
Lithium salts have to be extracted from the water of mineral springs, brine pools, and brine deposits. The metal is produced electrolytically from a mixture of fused lithium chloride and potassium chloride.[157]
Potassium occurs in many minerals, such as sylvite (potassium chloride).[6] It is occasionally produced through separating the potassium from the chlorine in potassium chloride, but is more often produced through electrolysis of potassium hydroxide,[158] found extensively in places such as Canada, Russia, Belarus, Germany, Israel, United States, and Jordan, in a method similar to how sodium was produced in the late 1800s and early 1900s.[159] It can also be produced from seawater. Sodium occurs mostly in seawater and dried seabed,[6] but is now produced through electrolysis of sodium chloride by lowering the melting point of the substance to below 700 °C through the use of a Downs cell.[160][161] Extremely pure sodium can be produced through the thermal decomposition of sodium azide.[162]
For several years in the 1950s and 1960s, a by-product of the potassium production called Alkarb was a main source for rubidium. Alkarb contained 21% rubidium while the rest was potassium and a small fraction of caesium.[163] Today the largest producers of caesium, for example the Tanco Mine, Manitoba, Canada, produce rubidium as by-product from pollucite.[164] Today, a common method for separating rubidium from potassium and caesium is the fractional crystallization of a rubidium and caesium alum (Cs, Rb)Al(SO4)2·12H2O, which yields pure rubidium alum after approximately 30 different reactions.[164][165] The limited applications and the lack of a mineral rich in rubidium limits the production of rubidium compounds to 2 to 4 tonnes per year.[164] Caesium, however, is not produced from the above reaction. Instead, the mining of pollucite ore is the main method of obtaining pure caesium, extracted from the ore mainly by three methods: acid digestion, alkaline decomposition, and direct reduction.[164][166]
Francium-223, the only naturally occurring isotope of francium,[75][76] is produced naturally as the product of the alpha decay of actinium-227. Francium can be found in trace amounts in uranium and thorium minerals;[152] it has been calculated that at most there are 30 g of francium in the earth's crust at any given time.[155] As a result of its extreme rarity in nature, most francium is synthesized in the nuclear reaction 197Au + 18O → 210Fr + 5 n, yielding francium-209, francium-210, and francium-211.[167] The greatest quantity of francium ever assembled to date is about 300,000 neutral atoms,[168] which were synthesized using the nuclear reaction given above.[168]
From their silicate ores, all the alkali metals may be obtained the same way: sulfuric acid is first used to dissolve the desired alkali metal ion and aluminium(III) ions from the ore (leaching), whereupon basic precipitation removes aluminium ions from the mixture by precipitating it as the hydroxide. The remaining insoluble alkali metal carbonate is then precipitated selectively; the salt is then dissolved in hydrochloric acid. The result is then left to evaporate and the alkali metal can then be isolated through electrolysis.[44]
Lithium and sodium are typically isolated through electrolysis from their liquid chlorides, with calcium chloride typically added to lower the melting point of the mixture. The heavier alkali metals, however, is more typically isolated in a different way, where a reducing agent (typically sodium for potassium and magnesium or calcium for the heaviest alkali metals) is used to reduce the alkali metal chloride. The liquid or gaseous product (the alkali metal) then undergoes fractional distillation for purification.[44]
All of the discovered alkali metals excluding francium have many applications. Lithium is often used in batteries, and lithium oxide can help process silica. Lithium can also be used to make lubricating greases, air treatment, and aluminium production.[169]
Pure sodium has many applications, including use in sodium-vapour lamps, which produce very efficient light compared to other types of lighting,[25][26] and can help smooth the surface of other metals.[170][171] Sodium compounds have many applications as well, the most well-known compound being table salt.[citation needed] Sodium is also used in soap as salts of fatty acids.[citation needed]
Potassium compounds are often used as fertilisers[9]:73[172] as potassium is an important element for plant nutrition. Other potassium ions are often used to hold anions.[citation needed][clarification needed] Potassium hydroxide is a very strong base, and is used to control the pH of various substances.[173][174]
Rubidium and caesium are often used in atomic clocks.[22] Caesium atomic clocks are extraordinarily accurate; if a clock had been made at the time of the dinosaurs, it would be off by less than four seconds (after 80 million years).[23] For that reason, caesium atoms are used as the definition of the second.[24] Rubidium ions are often used in purple fireworks,[175] and caesium is often used in drilling fluids in the petroleum industry.[23][176]
Francium has no commercial applications,[153][154][177] but because of francium's relatively simple atomic structure, among other things, it has been used in spectroscopy experiments, leading to more information regarding energy levels and the coupling constants between subatomic particles.[178] Studies on the light emitted by laser-trapped francium-210 ions have provided accurate data on transitions between atomic energy levels, similar to those predicted by quantum theory.[179]
Biological role and precautions[edit]
Lithium naturally only occurs in traces in biological systems and has no known biological role, but does have effects on the body when ingested.[29] Lithium carbonate is used as a mood stabiliser in psychiatry to treat bipolar disorder (manic-depression) in daily doses of about 0.5 to 2 grams, although there are side-effects.[29] Excessive ingestion of lithium causes drowsiness, slurred speech and vomiting, among other symptoms,[29] and poisons the central nervous system,[29] which is dangerous as the required dosage of lithium to treat bipolar disorder is only slightly lower than the toxic dosage.[29][180] Its biochemistry, the way it is handled by the human body and studies using rats and goats suggest that it is an essential trace element, although the natural biological function of lithium in humans has yet to be identified.[181][182]
Sodium and potassium occur in all known biological systems, generally functioning as electrolytes inside and outside cells.[27][28] Sodium is an essential nutrient that regulates blood volume, blood pressure, osmotic equilibrium and pH; the minimum physiological requirement for sodium is 500 milligrams per day.[183] Sodium chloride (also known as common salt) is the principal source of sodium in the diet, and is used as seasoning and preservative, such as for pickling and jerky; most of it comes from processed foods.[184] The DRI for sodium is 1.5 grams per day,[185] but most people in the United States consume more than 2.3 grams per day,[186] the minimum amount that promotes hypertension;[187] this in turn causes 7.6 million premature deaths worldwide.[188]
Potassium is the major cation (positive ion) inside animal cells,[27] while sodium is the major cation outside animal cells.[27][28] The concentration differences of these charged particles causes a difference in electric potential between the inside and outside of cells, known as the membrane potential. The balance between potassium and sodium is maintained by ion pumps in the cell membrane.[189] The cell membrane potential created by potassium and sodium ions allows the cell to generate an action potential—a "spike" of electrical discharge. The ability of cells to produce electrical discharge is critical for body functions such as neurotransmission, muscle contraction, and heart function.[189]
Rubidium has no known biological role, but may help stimulate metabolism,[30][190][191] and, similarly to caesium,[30][31] replace potassium in the body causing potassium deficiency.[30][191] Caesium compounds are rarely encountered by most people, but most caesium compounds are mildly toxic because of chemical similarity of caesium to potassium, allowing the caesium to replace the potassium in the body, causing potassium deficiency.[31] Exposure to large amounts of caesium compounds can cause hyperirritability and spasms, but as such amounts would not ordinarily be encountered in natural sources, caesium is not a major chemical environmental pollutant.[192] The median lethal dose (LD50) value for caesium chloride in mice is 2.3 g per kilogram, which is comparable to the LD50 values of potassium chloride and sodium chloride.[193] Caesium chloride has been promoted as an alternative cancer therapy,[194] but has been linked to the deaths of over 50 patients, on whom it was used as part of a scientifically unvalidated cancer treatment.[195] Radioisotopes of caesium require special precautions: the improper handling of caesium-137 gamma ray sources can lead to release of this radioisotope and radiation injuries. Perhaps the best-known case is the Goiânia accident of 1987, in which an improperly-disposed-of radiation therapy system from an abandoned clinic in the city of Goiânia, Brazil, was scavenged from a junkyard, and the glowing caesium salt sold to curious, uneducated buyers. This led to four deaths and serious injuries from radiation exposure. Together with caesium-134, iodine-131, and strontium-90, caesium-137 was among the isotopes distributed by the Chernobyl disaster which constitute the greatest risk to health.[100]
Francium has no biological role[196] and is most likely to be toxic due to its extreme radioactivity, causing radiation poisoning,[32] but since the greatest quantity of francium ever assembled to date is about 300,000 neutral atoms,[168] it is unlikely that most people will ever encounter francium.
- ^ The symbol Na for sodium is derived from its Latin name, natrium; this is still the name for the element in some languages, such as German and Russian. In early English texts, the symbol So for the English name sodium is sometimes seen; this is wholly obsolete.
- ^ The symbol K for potassium is derived from its Latin name, kalium; this is still the name for the element in some languages, such as German and Russian. In early English texts, the symbol Po for the English name potassium is sometimes seen; this is wholly obsolete, and presently would collide with the symbol for polonium (also Po).
- ^ Caesium is the spelling recommended by the International Union of Pure and Applied Chemistry (IUPAC).[1] The American Chemical Society (ACS) has used the spelling cesium since 1921,[2][3] following Webster's Third New International Dictionary.
- ^ In both the old IUPAC and the CAS systems for group numbering, this group is known as group IA (pronounced as "group one A", as the "I" is a Roman numeral).[11]
- ^ Noble gas notation is used for conciseness; the nearest noble gas that precedes the element in question is written first, and then the electron configuration is continued from that point forward.
- ^ The number given in parentheses refers to the measurement uncertainty. This uncertainty applies to the least significant figure(s) of the number prior to the parenthesized value (ie. counting from rightmost digit to left). For instance, 1.00794(7) stands for 1.00794±0.00007, while 1.00794(72) stands for 1.00794±0.00072.[74]
- ^ The value listed is the conventional value suitable for trade and commerce; the actual value may range from 6.938 to 6.997 depending on the isotopic composition of the sample.[76]
- ^ The element does not have any stable nuclides, and a value in brackets indicates the mass number of the longest-lived isotope of the element.[75][76]
- ^ The values are in picometres (pm). The shade of the box ranges from red to yellow as the radius increases. The atomic and ionic radii are displayed on the same scale of colour.
- ^ The shade of the box ranges from red to yellow as the ionisation energy decreases.
- ^ A different source gives 4.0712 ± 0.00004 eV (392.811(4) kJ/mol).[42]
- ^ The shade of the box ranges from red to yellow as the electronegativity decreases.
- ^ Linus Pauling estimated the electronegativity of francium at 0.7 on the Pauling scale, the same as caesium;[83] the value for caesium has since been refined to 0.79, although there are no experimental data to allow a refinement of the value for francium.[84] Francium has a slightly higher ionization energy than caesium,[42] 392.811(4) kJ/mol as opposed to 375.7041(2) kJ/mol for caesium, as would be expected from relativistic effects, and this would imply that caesium is the less electronegative of the two.
- ^ The shade of the box ranges from red to yellow as the melting and boiling points decrease. The melting and boiling points are displayed on different scales of colour.
- ^ a b Francium's melting point was claimed to have been calculated to be around 27 °C (80 °F, 300 K).[86] However, the melting point is uncertain because of the element's extreme rarity and radioactivity. Thus, the estimated boiling point value of 677 °C (1250 °F, 950 K) is also uncertain. Because radioactive elements give off heat, francium would almost certainly be a liquid at standard conditions[vague] if enough were to be produced.
- ^ The shade of the box ranges from red to yellow as the density increases.
- ^ Bunsen quotes Aulus Gellius Noctes Atticae II, 26 by Nigidius Figulus: Nostris autem veteribus "caesia" dicta est, quae a Graecis glaukopis, ut Nigidius ait, "de colore caeli quasi caelia.
- ^ Some synthetic elements, like technetium and plutonium, have later been found in nature.
- ^ The asterisk denotes an excited state.[citation needed]
- ^ International Union of Pure and Applied Chemistry (2005). Nomenclature of Inorganic Chemistry (IUPAC Recommendations 2005). Cambridge (UK): RSC–IUPAC. ISBN 0-85404-438-8. pp. 248–49. Electronic version..
- ^ Coghill, Anne M.; Garson, Lorrin R., eds. (2006). The ACS Style Guide: Effective Communication of Scientific Information (3rd ed.). Washington, D.C.: American Chemical Society. p. 127. ISBN 0-8412-3999-1..
- ^ Coplen, T. B.; Peiser, H. S. (1998). "History of the recommended atomic-weight values from 1882 to 1997: a comparison of differences from current values to the estimated uncertainties of earlier values". Pure Appl. Chem. 70 (1): 237–257. doi:10.1351/pac199870010237..
- ^ a b c International Union of Pure and Applied Chemistry (2005). Nomenclature of Inorganic Chemistry (IUPAC Recommendations 2005). Cambridge (UK): RSC–IUPAC. ISBN 0-85404-438-8. pp. 51. Electronic version..
- ^ Leach, Mark R. (1999–2012). "The Internet Database of Periodic Tables". meta-synthesis.com. Retrieved 20 May 2012.
- ^ a b c d e f g h i j k l m n o p q r s t u v w x y z aa ab ac ad ae af ag ah ai aj ak al am an Royal Society of Chemistry. "Visual Elements: Group 1 – The Alkali Metals". Visual Elements. Royal Society of Chemistry. Retrieved 13 January 2012.
- ^ "Periodic Table: Atomic Properties of the Elements". nist.gov. National Institute of Standards and Technology. September 2010. Retrieved 17 February 2012.
- ^ a b c d e f g h Lide, D. R., ed. (2003). CRC Handbook of Chemistry and Physics (84th ed.). Boca Raton, FL: CRC Press.
- ^ a b c d e f g h i j k l m n o p Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN 0080379419.
- ^ a b c The OpenLearn team (2012). "Alkali metals". OpenLearn. The Open University. Retrieved 9 July 2012.
- ^ Fluck, E. (1988). "New Notations in the Periodic Table". Pure Appl. Chem. (IUPAC) 60 (3): 431–436. doi:10.1351/pac198860030431. Retrieved 24 March 2012.
- ^ "IUPAC Periodic Table of the Elements". iupac.org. International Union of Pure and Applied Chemistry. 21 January 2011. Retrieved 22 February 2012.
- ^ a b c "International Union of Pure and Applied Chemistry > Periodic Table of the Elements". IUPAC. Retrieved 1 May 2011.
- ^ a b Folden, Cody (31 January 2009). "The Heaviest Elements in the Universe". Saturday Morning Physics at Texas A&M. Retrieved 9 March 2012.
- ^ a b c d e f g Vinson, Greg (2008). "Hydrogen is a Halogen". HydrogenTwo.com. Retrieved 14 January 2012.
- ^ a b c Gray, Theodore. "Alkali Metal Bangs". Theodore Gray. Retrieved 13 May 2012.
- ^ a b "Abundance in Earth's Crust". WebElements.com. Retrieved 14 April 2007.
- ^ a b Adloff, Jean-Pierre; Kaufman, George B. (2005-09-25). Francium (Atomic Number 87), the Last Discovered Natural Element. The Chemical Educator 10 (5). Retrieved on 26 March 2007.
- ^ a b c Lougheed, R.W.; Landrum, J.H.; Hulet, E.K.; Wild, J.F.; Dougan, R.J.; Dougan, A.D.; Gäggeler, H.; Schädel, M; Moody, K.J.; Gregorich, K.E.; Seaborg, G. (1985). "Search for superheavy elements using 48Ca + 254Esg reaction". Physical Review C 32 (5): 1760–1763. Bibcode:1985PhRvC..32.1760L. doi:10.1103/PhysRevC.32.1760.
- ^ a b c Gäggeler, Heinz W. (5–7 November 2007). "Gas Phase Chemistry of Superheavy Elements". Lecture Course Texas A&M. Retrieved 26 February 2012.
- ^ a b c d e f g h i j k l m n o p q r s t u v Hoffman, Darleane C.; Lee, Diana M.; Pershina, Valeria (2006). "Transactinides and the future elements". In Morss; Edelstein, Norman M.; Fuger, Jean. The Chemistry of the Actinide and Transactinide Elements (3rd ed.). Dordrecht, The Netherlands: Springer Science+Business Media. ISBN 1-4020-3555-1.
- ^ a b "Cesium Atoms at Work". Time Service Department—U.S. Naval Observatory—Department of the Navy. Retrieved 20 December 2009.
- ^ a b c d e Butterman, William C.; Brooks, William E.; Reese, Jr., Robert G. (2004). "Mineral Commodity Profile: Cesium" (PDF). United States Geological Survey. Archived from the original on 22 November 2009. Retrieved 27 December 2009.
- ^ a b "The NIST reference on Constants, Units, and Uncertainty". National Institute of Standards and Technology.
- ^ a b Lindsey, Jack L (1997). Applied illumination engineering. p. 112. ISBN 978-0-88173-212-2.
- ^ a b Kane, Raymond; Sell, Heinz (2001). Revolution in lamps: A chronicle of 50 years of progress. p. 241. ISBN 978-0-88173-351-8.
- ^ a b c d Winter, Mark. "WebElements Periodic Table of the Elements | Potassium | biological information". WebElements. Retrieved 13 January 2012.
- ^ a b c Winter, Mark. "WebElements Periodic Table of the Elements | Sodium | biological information". WebElements. Retrieved 13 January 2012.
- ^ a b c d e f Winter, Mark. "WebElements Periodic Table of the Elements | Lithium | biological information". Webelements. Retrieved 15 February 2011.
- ^ a b c d Winter, Mark. "WebElements Periodic Table of the Elements | Rubidium | biological information". Webelements. Retrieved 15 February 2011.
- ^ a b c Winter, Mark. "WebElements Periodic Table of the Elements | Caesium | biological information". WebElements. Retrieved 13 January 2012.
- ^ a b "Francium – Element information, properties and uses | Periodic Table". Visual Elements Periodic Table. Royal Society of Chemistry. 2012. Retrieved 27 June 2012.
- ^ a b Krebs, Robert E. (2006). The History and Use of Our Earth's Chemical Elements: A Reference Guide. Westport, Conn.: Greenwood Press. ISBN 0-313-33438-2.
- ^ J. L. Dye, J. M. Ceraso, Mei Lok Tak, B. L. Barnett, F. J. Tehan (1974). "Crystalline salt of the sodium anion (Na−)". J. Am. Chem. Soc. 96 (2): 608–609. doi:10.1021/ja00809a060.
- ^ F. J. Tehan, B. L. Barnett, J. L. Dye (1974). "Alkali anions. Preparation and crystal structure of a compound which contains the cryptated sodium cation and the sodium anion". J. Am. Chem. Soc. 96 (23): 7203–7208. doi:10.1021/ja00830a005.
- ^ J. L. Dye (1979). "Compounds of Alkali Metal Anions". Angew. Chem. Int. Ed. Engl. 18 (8): 587–598. doi:10.1002/anie.197905871.
- ^ M. Y. Redko, R. H. Huang, J. E. Jackson, J. F. Harrison, J. L. Dye (2003). "Barium azacryptand sodide, the first alkalide with an alkaline Earth cation, also contains a novel dimer, (Na2)2−". J. Am. Chem. Soc. 125 (8): 2259–2263. doi:10.1021/ja027241m. PMID 12590555.
- ^ a b c M. Y. Redko, M. Vlassa, J. E. Jackson, A. W. Misiolek, R. H. Huang RH, J. L. Dye (2002). ""Inverse sodium hydride": a crystalline salt that contains H+ and Na−". J. Am. Chem. Soc. 124 (21): 5928–5929. doi:10.1021/ja025655.
- ^ a b Agnieszka Sawicka, Piotr Skurski, and Jack Simons (2003). "Inverse Sodium Hydride: A Theoretical Study". J. Am. Chem. Soc. 125 (13): 3954–3958. doi:10.1021/ja021136v. PMID 12656631.
- ^ a b c d e Clark, Jim (2005). "Reaction of the Group 1 Elements with Oxygen and Chlorine". chemguide. Retrieved 27 June 2012.
- ^ Shriver, Duward; Atkins, Peter (2006). Inorganic Chemistry. W. H. Freeman. p. 259. ISBN 978-0716748786. Retrieved 10 November 2012.
- ^ a b c Andreev, S.V.; Letokhov, V.S.; Mishin, V.I., (1987). "Laser resonance photoionization spectroscopy of Rydberg levels in Fr". Phys. Rev. Lett. 59 (12): 1274–76. Bibcode:1987PhRvL..59.1274A. doi:10.1103/PhysRevLett.59.1274. PMID 10035190.
- ^ a b Clark, Jim (2005). "Reaction of the Group 1 Elements with Water". chemguide. Retrieved 18 June 2012.
- ^ a b c d e f g h i j k l m n o Averill, Bruce A.; Eldredge, Patricia (2007). "21.3: The Alkali Metals". Chemistry: Principles, Patterns, and Applications with Student Access Kit for Mastering General Chemistry (1st ed.). Prentice Hall. ISBN 9780805337990. Retrieved 24 June 2013.
- ^ Catherine E. Housecroft; Alan G. Sharpe (2008). "Chapter 14: The group 14 elements". Inorganic Chemistry, 3rd Edition. Pearson. p. 386. ISBN 978-0-13-175553-6.
- ^ NIST Ionizing Radiation Division 2001 - Technical Highlights
- ^ N. Emery et al. (2008). "Review: Synthesis and superconducting properties of CaC6". Sci. Technol. Adv. Mater. 9 (4): 044102. Bibcode:2008STAdM...9d4102E. doi:10.1088/1468-6996/9/4/044102.
- ^ S.M. Kauzlarich,(1994), Zintl Compounds, Encyclopedia of Inorganic Chemistry, John Wiley & sons, ISBN 0-471-93620-0
- ^ Hoch, Constantin; Wendorff, Marco; Röhr, Caroline (2002). "Tetrapotassium nonastannide, K4Sn9". Acta Crystallographica Section C Crystal Structure Communications 58 (4): i45. doi:10.1107/S0108270102002032.
- ^ Duncan H. Gregory, Paul M. O'Meara, Alexandra G. Gordon, Jason P. Hodges, Simine Short, and James D. Jorgensen (2002). "Structure of Lithium Nitride and Transition-Metal-Doped Derivatives, Li3−x−yMxN (M = Ni, Cu): A Powder Neutron Diffraction Study". Chem. Mater. 14 (5): 2063–2070. doi:10.1021/cm010718t.
- ^ Fischer, D., Jansen, M. (2002). "Synthesis and structure of Na3N". Angew Chem 41 (10): 1755. doi:10.1002/1521-3773(20020517)41:10<1755::AID-ANIE1755>3.0.CO;2-C.
- ^ Fischer, D.; Cancarevic, Z.; Schön, J. C.; Jansen, M. Z. (2004). "Synthesis and structure of K3N". Z. anorg allgem Chemie 630 (1): 156. doi:10.1002/zaac.200300280.. 'Elusive Binary Compound Prepared' Chemical & Engineering News 80 No. 20 (20 May 2002)
- ^ H.G. Von Schnering, W. Hönle Phosphides - Solid-state Chemistry Encyclopedia of Inorganic Chemistry Ed. R. Bruce King (1994) John Wiley & Sons ISBN 0-471-93620-0
- ^ Kahlenberg, Louis (2008). Outlines of Chemistry – A Textbook for College Students. READ BOOKS. pp. 324–325. ISBN 1-4097-6995-X.
- ^ "Welcome to Arthur Mar's Research Group". University of Alberta. University of Alberta. 1999–2013. Retrieved 24 June 2013.
- ^ Lindsay, D. M.; Garland, D. A. (1987). "ESR spectra of matrix-isolated lithium superoxide". The Journal of Physical Chemistry 91 (24): 6158–61. doi:10.1021/j100308a020.
- ^ Vol'nov, I. I.; Matveev, V. V. (1963). "Synthesis of cesium ozonide through cesium superoxide". Bulletin of the Academy of Sciences, USSR Division of Chemical Science 12 (6): 1040–1043. doi:10.1007/BF00845494.
- ^ Tokareva, S. A. (1971). "Alkali and Alkaline Earth Metal Ozonides". Russian Chemical Reviews 40 (2): 165–174. Bibcode:1971RuCRv..40..165T. doi:10.1070/RC1971v040n02ABEH001903.
- ^ Simon, A. (1997). "Group 1 and 2 Suboxides and Subnitrides — Metals with Atomic Size Holes and Tunnels". Coordination Chemistry Reviews 163: 253–270. doi:10.1016/S0010-8545(97)00013-1.
- ^ Tsai, Khi-Ruey; Harris, P. M.; Lassettre, E. N. (1956). "The Crystal Structure of Tricesium Monoxide". Journal of Physical Chemistry 60 (3): 345–347. doi:10.1021/j150537a023.
- ^ Okamoto, H. (2009). "Cs-O (Cesium-Oxygen)". Journal of Phase Equilibria and Diffusion 31: 86. doi:10.1007/s11669-009-9636-5.
- ^ Band, A.; Albu-Yaron, A.; Livneh, T.; Cohen, H.; Feldman, Y.; Shimon, L.; Popovitz-Biro, R.; Lyahovitskaya, V.; Tenne, R. (2004). "Characterization of Oxides of Cesium". The Journal of Physical Chemistry B 108 (33): 12360–12367. doi:10.1021/jp036432o.
- ^ Brauer, G. (1947). "Untersuchungen ber das System Cäsium-Sauerstoff". Zeitschrift für anorganische Chemie 255: 101. doi:10.1002/zaac.19472550110.
- ^ Butterman, William C.; Brooks, William E.; Reese, Jr., Robert G. (2004). "Mineral Commodity Profile: Cesium" (PDF). United States Geological Survey. Retrieved 27 December 2009.
- ^ House, James E. (2008). Inorganic chemistry. Academic Press. p. 524. ISBN 0-12-356786-6.
- ^ Moyer, Harvey V. (1956). "Chemical Properties of Polonium". In Moyer, Harvey V. Polonium. Oak Ridge, Tenn.: United States Atomic Energy Commission. pp. 33–96. doi:10.2172/4367751. TID-5221..
- ^ Bagnall, K. W. (1962). "The Chemistry of Polonium". Adv. Inorg. Chem. Radiochem. Advances in Inorganic Chemistry and Radiochemistry 4: 197–229. doi:10.1016/S0065-2792(08)60268-X. ISBN 978-0-12-023604-6..
- ^ Alberto, R.; Ortner, K.; Wheatley, N.; Schibli, R.; Schubiger, A. P. (2001). "Synthesis and properties of boranocarbonate: a convenient in situ CO source for the aqueous preparation of [99mTc(OH2)3(CO)3]+". J. Am. Chem. Soc. 121 (13): 3135–3136. doi:10.1021/ja003932b.
- ^ Cotton, F.A.; Wilkinson, G. (1972). Advanced Inorganic Chemistry. John Wiley and Sons Inc. ISBN 0-471-17560-9.
- ^ Brown, T. L.; Rogers, M. T. (1957). "The Preparation and Properties of Crystalline Lithium Alkyls". Journal of the American Chemical Society 79 (8): 1859–1861. doi:10.1021/ja01565a024.
- ^ Manfred Schlosser (1988). "Superbases for organic synthesis". Pure and Appl. Chem. 60 (11): 1627–1634. doi:10.1351/pac198860111627.
- ^ Clegg, William; Conway, Ben; Kennedy, Alan R.; Klett, Jan; Mulvey, Robert E.; Russo, Luca (2011). "Synthesis and Structures of \(Trimethylsilyl)methyl]sodium and -potassium with Bi- and Tridentate N-Donor Ligands". European Journal of Inorganic Chemistry 2011 (5): 721. doi:10.1002/ejic.201000983.
- ^ Gray, Theodore. "Facts, pictures, stories about the element Cesium in the Periodic Table". The Wooden Periodic Table Table. Retrieved 13 January 2012.
- ^ "Standard Uncertainty and Relative Standard Uncertainty". CODATA reference. National Institute of Standards and Technology. Retrieved 26 September 2011.
- ^ a b c d Wieser, Michael E.; Berglund, Michael (2009). "Atomic weights of the elements 2007 (IUPAC Technical Report)". Pure Appl. Chem. (IUPAC) 81 (11): 2131–2156. doi:10.1351/PAC-REP-09-08-03. Retrieved 7 February 2012.
- ^ a b c d e Wieser, Michael E.; Coplen, Tyler B. (2011). "Atomic weights of the elements 2009 (IUPAC Technical Report)". Pure Appl. Chem. (IUPAC) 83 (2): 359–396. doi:10.1351/PAC-REP-10-09-14. Retrieved 11 February 2012.
- ^ a b c d e Clark, Jim (2005). "Flame Tests". chemguide. Retrieved 29 January 2012.
- ^ a b Klehr, Wolfram (21 May 2007). "Francium". apsidium.com. Archived from the original on 9 May 2008. Retrieved 25 April 2012.
- ^ a b c d e f g h i j k l m n o Clark, Jim (2005). "Atomic and Physical Properties of the Group 1 Elements". chemguide. Retrieved 30 January 2012.
- ^ L. Brown, Theodore; LeMay, Jr., H. Eugene; Bursten, Bruce E.; Burdge, Julia R. (2003). Chemistry: The Central Science (8th ed.). US: Pearson Education. ISBN 0-13-061142-5.[dead link]
- ^ a b J.E. Huheey, E.A. Keiter, and R.L. Keiter in Inorganic Chemistry : Principles of Structure and Reactivity, 4th edition, HarperCollins, New York, USA, 1993.
- ^ a b A.M. James and M.P. Lord in Macmillan's Chemical and Physical Data, Macmillan, London, UK, 1992.
- ^ Pauling, Linus (1960). The Nature of the Chemical Bond (Third ed.). Cornell University Press. p. 93. ISBN 978-0-8014-0333-0.
- ^ Allred, A. L. (1961). "Electronegativity values from thermochemical data". J. Inorg. Nucl. Chem. 17 (3–4): 215–221. doi:10.1016/0022-1902(61)80142-5.
- ^ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "Electronegativity".
- ^ "Francium". Los Alamos National Laboratory. 15 December 2003. Retrieved 19 February 2012.
- ^ David. E. Goldberg (1988). 3,000 Solved Problems in Chemistry (1st ed.). McGraw-Hill. ISBN 0-07-023684-4. Section 17.43, page 321
- ^ Louis Theodore, R. Ryan Dupont and Kumar Ganesan (Editors) (1999). Pollution Prevention: The Waste Management Approach to the 21st Century. CRC Press. ISBN 1-56670-495-2. Section 27, p. 15
- ^ a b Clark, Jim (2000). "Metallic Bonding". chemguide. Retrieved 23 March 2012.
- ^ a b Various authors (2002). Lide, David R., ed. Handbook of Chemistry & Physics (88th ed.). CRC. ISBN 0-8493-0486-5. OCLC 179976746. Retrieved 2008-05-23.
- ^ "Universal Nuclide Chart". Nucleonica. Institute for Transuranium Elements. 2007–2012. Retrieved 2011-04-17.
- ^ a b c Sonzogni, Alejandro. "Interactive Chart of Nuclides". National Nuclear Data Center: Brookhaven National Laboratory. Retrieved 4 October 2012.
- ^ Patton, I. Jocelyn; Waldbauer, L. J. (1926). "The Radioactivity of the Alkali Metals". Chemical Reviews 3: 81. doi:10.1021/cr60009a003.
- ^ McLennan, J. C.; Kennedy, W. T. (1908). "On the radioactivity of potassium and other alkali metals". Philosophical Magazine. 6 16 (93): 377–395. doi:10.1080/14786440908636519.
- ^ "Potassium-40". Human Health Fact Sheet. Argonne National Laboratory, Environmental Science Division. August 2005. Retrieved 7 February 2012.
- ^ a b Fontani, Marco (10 September 2005). "The Twilight of the Naturally-Occurring Elements: Moldavium (Ml), Sequanium (Sq) and Dor (Do)". International Conference on the History of Chemistry. Lisbon. pp. 1–8. Archived from the original on 24 February 2006. Retrieved 8 April 2007.
- ^ a b Van der Krogt, Peter (10 January 2006). "Francium". Elementymology & Elements Multidict. Retrieved 8 April 2007.
- ^ National Institute of Standards and Technology. "Radionuclide Half-Life Measurements". Retrieved 2011-11-07.
- ^ http://www.bt.cdc.gov/radiation/isotopes/cesium.asp
- ^ a b The Radiological Accident in Goiânia. IAEA. 1988.
- ^ a b c d Pyykkö, Pekka (2011). "A suggested periodic table up to Z ≤ 172, based on Dirac–Fock calculations on atoms and ions". Physical Chemistry Chemical Physics 13 (1): 161–8. Bibcode:2011PCCP...13..161P. doi:10.1039/c0cp01575j. PMID 20967377.
- ^ a b c Seaborg, G. T. (c. 2006). "transuranium element (chemical element)". Encyclopædia Britannica. Retrieved 16 March 2010.
- ^ Thayer, John S. (2010). Chemistry of heavier main group elements. pp. 81, 84. doi:10.1007/9781402099755_2.
- ^ Emsley, J. (1989). The Elements. Oxford: Clarendon Press. pp. 22–23.
- ^ Chemical Bonding, Mark J. Winter, Oxford University Press, 1994, ISBN 0-19-855694-2
- ^ a b Cronyn, Marshall W. (August 2003). "The Proper Place for Hydrogen in the Periodic Table". Journal of Chemical Education 80 (8): 947–951. Bibcode:2003JChEd..80..947C. doi:10.1021/ed080p947.
- ^ Wigner, E.; Huntington, H.B. (1935). "On the possibility of a metallic modification of hydrogen". Journal of Chemical Physics 3 (12): 764. Bibcode:1935JChPh...3..764W. doi:10.1063/1.1749590.
- ^ Mark R. Leach. "2002 Inorganic Chemist's Periodic Table". Retrieved 16 October 2012.
- ^ Holleman, A. F.; Wiberg, E. (2001), Inorganic Chemistry, San Diego: Academic Press, ISBN 0-12-352651-5
- ^ a b c Stevenson, D. J. (20 November 1975). "Does metallic ammonium exist?". Nature (Nature Publishing Group) 258 (5532): 222–223. Bibcode:1975Natur.258..222S. doi:10.1038/258222a0. Retrieved 13 January 2012.
- ^ a b Bernal, M. J. M.; Massey, H. S. W. (3 February 1954). "Metallic Ammonium". Monthly Notices of the Royal Astronomical Society (Wiley-Blackwell for the Royal Astronomical Society) 114: 172–179. Bibcode:1954MNRAS.114..172B.
- ^ "Solubility Rules!". Retrieved 4 January 2014.
- ^ R. D. Shannon (1976). "Revised effective ionic radii and systematic studies of interatomic distances in halides and chalcogenides". Acta Cryst A32 (5): 751–767. Bibcode:1976AcCrA..32..751S. doi:10.1107/S0567739476001551.
- ^ a b Crookes, William (1864). "On Thallium". The Journal of the Chemical Society, London (Harrison & Sons) XVII: 112–152. doi:10.1039/js8641700112. Retrieved 13 January 2012.
- ^ a b Leach, Mark R. (1999–2012). "The Internet Database of Periodic Tables". meta-synthesis.com. Retrieved 6 April 2012.
- ^ Ralph, Jolyon; Chau, Ida (24 August 2011). "Petalite: Petalite mineral information and data". Retrieved 27 November 2011.
- ^ a b Winter, Mark. "WebElements Periodic Table of the Elements | Lithium | historical information". Retrieved 27 November 2011.
- ^ Weeks, Mary (2003). Discovery of the Elements. Whitefish, Montana, United States: Kessinger Publishing. p. 124. ISBN 0-7661-3872-0. Retrieved 10 August 2009.
- ^ "Johan Arfwedson". Archived from the original on 5 June 2008. Retrieved 10 August 2009.
- ^ a b van der Krogt, Peter. "Lithium". Elementymology & Elements Multidict. Retrieved 5 October 2010.
- ^ Clark, Jim (2005). "Compounds of the Group 1 Elements". chemguide. Retrieved 10 August 2009.
- ^ a b Davy, Humphry (1808). "On some new phenomena of chemical changes produced by electricity, in particular the decomposition of the fixed alkalies, and the exhibition of the new substances that constitute their bases; and on the general nature of alkaline bodies". Philosophical Transactions of the Royal Society of London 98: 1–44. doi:10.1098/rstl.1808.0001.
- ^ Marggraf, Andreas Siegmund (1761). Chymische Schriften. p. 167.
- ^ du Monceau, H. L. D. "Sur la Base de Sel Marine". Memoires de l'Academie royale des Sciences (in French): 65–68.
- ^ a b Weeks, Mary Elvira (1932). "The discovery of the elements. IX. Three alkali metals: Potassium, sodium, and lithium". Journal of Chemical Education 9 (6): 1035. Bibcode:1932JChEd...9.1035W. doi:10.1021/ed009p1035.
- ^ a b Siegfried, R. (1963). "The Discovery of Potassium and Sodium, and the Problem of the Chemical Elements". Isis 54 (2): 247–258. doi:10.1086/349704. JSTOR 228541.
- ^ Enghag, P. (2004). "11. Sodium and Potassium". Encyclopedia of the elements. Wiley-VCH Weinheim. ISBN 3-527-30666-8.
- ^ Shaposhnik, V. A. (2007). "History of the discovery of potassium and sodium (on the 200th anniversary of the discovery of potassium and sodium)". Journal of Analytical Chemistry 62 (11): 1100–1102. doi:10.1134/S1061934807110160.
- ^ a b Weeks, Mary Elvira (1932). "The discovery of the elements. XIII. Some spectroscopic discoveries". Journal of Chemical Education 9 (8): 1413–1434. Bibcode:1932JChEd...9.1413W. doi:10.1021/ed009p1413.
- ^ Kirchhoff, G.; Bunsen, R. (1861). "Chemische Analyse durch Spectralbeobachtungen". Annalen der Physik und Chemie 189 (7): 337–381. Bibcode:1861AnP...189..337K. doi:10.1002/andp.18611890702.
- ^ a b Kaner, Richard (2003). "C&EN: It's Elemental: The Periodic Table – Cesium". American Chemical Society. Retrieved 25 February 2010.
- ^ Oxford English Dictionary, 2nd Edition
- ^ "Alabamine & Virginium". TIME. 15 February 1932. Retrieved 1 April 2007.
- ^ MacPherson, H. G. (1934). "An Investigation of the Magneto-Optic Method of Chemical Analysis". Physical Review (American Physical Society) 47 (4): 310–315. Bibcode:1935PhRv...47..310M. doi:10.1103/PhysRev.47.310.
- ^ "Francium". McGraw-Hill Encyclopedia of Science & Technology 7. McGraw-Hill Professional. 2002. pp. 493–494. ISBN 0-07-913665-6.
- ^ van der Krogt, Peter. "Ununennium". Elementymology & Elements Multidict. Retrieved 14 February 2011.
- ^ Schadel, M.; Brüchle, W.; Brügger, M.; Gäggeler, H.; Moody, K.; Schardt, D.; Sümmerer, K.; Hulet, E.; Dougan, A. et al. (1986). "Heavy isotope production by multinucleon transfer reactions with 254Es". Journal of the Less Common Metals 122: 411. doi:10.1016/0022-5088(86)90435-2.
- ^ "Modern alchemy: Turning a line". The Economist. 12 May 2012. Retrieved 5 October 2012.
- ^ Emsley, John (2011). Nature's Building Blocks: An A-Z Guide to the Elements (New ed.). New York, NY: Oxford University Press. p. 593. ISBN 978-0-19-960563-7.
- ^ a b Lodders, Katharina (2003). "Solar System Abundances and Condensation Temperatures of the Elements". The Astrophysical Journal 591 (2): 1220–1247. Bibcode:2003ApJ...591.1220L. doi:10.1086/375492.
- ^ Oddo, Giuseppe (1914). "Die Molekularstruktur der radioaktiven Atome". Zeitschrift für anorganische Chemie 87: 253. doi:10.1002/zaac.19140870118.
- ^ Harkins, William D. (1917). "The Evolution of the Elements and the Stability of Complex Atoms. I. A New Periodic System Which Shows a Relation Between the Abundance of the Elements and the Structure of the Nuclei of Atoms". Journal of the American Chemical Society 39 (5): 856. doi:10.1021/ja02250a002.
- ^ North, John (2008). Cosmos an illustrated history of astronomy and cosmology (Rev. and updated ed.). Univ. of Chicago Press. p. 602. ISBN 978-0-226-59441-5.
- ^ Morgan, J. W.; Anders, E. (1980). "Chemical composition of Earth, Venus, and Mercury". Proceedings of the National Academy of Sciences 77 (12): 6973–6977. Bibcode:1980PNAS...77.6973M. doi:10.1073/pnas.77.12.6973. PMC 350422. PMID 16592930.
- ^ Albarède, Francis (2003). Geochemistry: an introduction. Cambridge University Press. ISBN 978-0-521-89148-6.
- ^ "List of Periodic Table Elements Sorted by Abundance in Earth's crust". Israel Science and Technology Homepage. Retrieved 15 April 2007.
- ^ a b c Lide, D. R., ed. (2005). CRC Handbook of Chemistry and Physics (86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5.
- ^ "Lithium Occurrence". Institute of Ocean Energy, Saga University, Japan. Retrieved 13 March 2009.[dead link]
- ^ "Some Facts about Lithium". ENC Labs. Retrieved 15 October 2010.
- ^ Schwochau, Klaus (1984). "Extraction of metals from sea water". Topics in Current Chemistry. Topics in Current Chemistry. 124/1984: 91–133. doi:10.1007/3-540-13534-0_3. ISBN 978-3-540-13534-0.
- ^ Wise, M. A. (1995). "Trace element chemistry of lithium-rich micas from rare-element granitic pegmatites". Mineralogy and Petrology 55 (13): 203–215. Bibcode:1995MinPe..55..203W. doi:10.1007/BF01162588.
- ^ a b CRC Handbook of Chemistry and Physics 4. CRC. 2006. p. 12. ISBN 0-8493-0474-1.
- ^ a b c Emsley, John (2001). Nature's Building Blocks. Oxford: Oxford University Press. pp. 151–153. ISBN 0-19-850341-5.
- ^ a b Gagnon, Steve. "Francium". Jefferson Science Associates, LLC. Archived from the original on 31 March 2007. Retrieved 1 April 2007.
- ^ a b Winter, Mark. "Geological information". Francium. The University of Sheffield. Retrieved 26 March 2007.
- ^ "It's Elemental — The Periodic Table of Elements". Jefferson Lab. Archived from the original on 29 April 2007. Retrieved 14 April 2007.
- ^ Ober, Joyce A. "Lithium" (PDF). United States Geological Survey. pp. 77–78. Archived from the original on 11 July 2007. Retrieved 19 August 2007.
- ^ Winter, Mark. "WebElements Periodic Table of the Elements | Potassium | Essential information". Webelements. Retrieved 27 November 2011.
- ^ Lemke, Charles H.; Markant, Vernon H. (2001). "Sodium and Sodium Alloys". Kirk-Othmer Encyclopedia of Chemical Technology. doi:10.1002/0471238961.1915040912051311.a01.pub2. ISBN 0471238961.
- ^ Pauling, Linus. General Chemistry (1970 ed.). Dover Publications.
- ^ "Los Alamos National Laboratory – Sodium". Retrieved 8 June 2007.
- ^ Merck Index, 9th ed., monograph 8325
- ^ "Cesium and Rubidium Hit Market". Chemical & Engineering News 37 (22): 50–56. 1959. doi:10.1021/cen-v037n022.p050.
- ^ a b c d Butterman, William C.; Brooks, William E.; Reese, Jr., Robert G. (2003). "Mineral Commodity Profile: Rubidium" (PDF). United States Geological Survey. Retrieved 4 December 2010.
- ^ bulletin 585. United States. Bureau of Mines. 1995.
- ^ Burt, R. O. (1993). "Caesium and cesium compounds". Kirk-Othmer encyclopedia of chemical technology 5 (4th ed.). New York: John Wiley & Sons, Inc. pp. 749–764. ISBN 978-0-471-48494-3.
- ^ Stancari, G.; Veronesi, S.; Corradi, L.; Atutov, S. N.; Calabrese, R.; Dainelli, A.; Mariotti, E.; Moi, L.; Sanguinetti, S.; Tomassetti, L. (2006). "Production of Radioactive Beams of Francium". Nuclear Instruments and Methods in Physics Research Section A: Accelerators, Spectrometers, Detectors and Associated Equipment 557 (2): 390–396. Bibcode:2006NIMPA.557..390S. doi:10.1016/j.nima.2005.11.193.
- ^ a b c Luis A. Orozco (2003). "Francium". Chemical and Engineering News.
- ^ USGS (2011). "Lithium" (PDF). Retrieved 4 December 2011.
- ^ Stampers, National Association of Drop Forgers and (1957). Metal treatment and drop forging.
- ^ Harris, Jay C (1949). Metal cleaning bibliographical abstracts. p. 76.
- ^ Cordel, Oskar (1868). Die Stassfurter Kalisalze in der Landwirthschalt: Eine Besprechung ... (in German). L. Schnock. Retrieved 29 May 2011.
- ^ Toedt, John; Koza, Darrell; Cleef-Toedt, Kathleen Van (2005). "Personal Cleansing Products: Bar Soap". Chemical composition of everyday products. Greenwood Publishing Group. ISBN 978-0-313-32579-3.
- ^ Schultz, H. et al. (2006). "Potassium compounds". Ullmann's Encyclopedia of Industrial Chemistry A22. p. 95. doi:10.1002/14356007.a22_031.pub2. ISBN 3-527-30673-0.
- ^ Koch, E.-C. (2002). "Special Materials in Pyrotechnics, Part II: Application of Caesium and Rubidium Compounds in Pyrotechnics". Journal Pyrotechnics 15: 9–24.
- ^ Heiserman, David L. (1992). Exploring Chemical Elements and their Compounds. McGraw-Hill. pp. 201–203. ISBN 0-8306-3015-5.
- ^ Winter, Mark. "Uses". Francium. The University of Sheffield. Archived from the original on 31 March 2007. Retrieved 25 March 2007.
- ^ Gomez, E; Orozco, L A; Sprouse, G D (7 November 2005). "Spectroscopy with trapped francium: advances and perspectives for weak interaction studies". Rep. Prog. Phys. 69 (1): 79–118. Bibcode:2006RPPh...69...79G. doi:10.1088/0034-4885/69/1/R02.
- ^ Peterson, I (11 May 1996). "Creating, cooling, trapping francium atoms". Science News 149 (19): 294. doi:10.2307/3979560. Retrieved 11 September 2009.
- ^ Gray, Theodore. "Facts, pictures, stories about the element Lithium in the Periodic Table". theodoregray.com. Retrieved 9 January 2012.
- ^ Howland, Robert H. (September 2007). "Lithium: Underappreciated and Underused?". Psychiatric Annals 37 (9). Retrieved 6 November 2012.
- ^ Zarse, Kim; Terao, Takeshi; Tian, Jing; Iwata, Noboru; Ishii, Nobuyoshi; Ristow, Michael (August 2011). "Low-dose lithium uptake promotes longevity in humans and metazoans". European Journal of Nutrition (Springer) 50 (5): 387–9. doi:10.1007/s00394-011-0171-x. PMC 3151375. PMID 21301855. Retrieved 6 November 2012.
- ^ "Sodium". Northewestern University. Retrieved 21 November 2011.[dead link]
- ^ "Sodium and Potassium Quick Health Facts". Retrieved 7 November 2011.
- ^ "Dietary Reference Intakes: Water, Potassium, Sodium, Chloride, and Sulfate". Food and Nutrition Board, Institute of Medicine, United States National Academies. 11 February 2004. Retrieved 23 November 2011.
- ^ U.S. Department of Agriculture; U.S. Department of Health and Human Services (December 2010). Dietary Guidelines for Americans, 2010 (PDF) (7th ed.). p. 22. ISBN 978-0-16-087941-8. OCLC 738512922. Retrieved 23 November 2011.
- ^ Geleijnse, J. M.; Kok, F. J.; Grobbee, D. E. (2004). "Impact of dietary and lifestyle factors on the prevalence of hypertension in Western populations". European Journal of Public Health 14 (3): 235–239. doi:10.1093/eurpub/14.3.235. PMID 15369026.
- ^ Lawes, C. M.; Vander Hoorn, S.; Rodgers, A.; International Society of Hypertension (2008). "Global burden of blood-pressure-related disease, 2001". Lancet 371 (9623): 1513–1518. doi:10.1016/S0140-6736(08)60655-8. PMID 18456100.[dead link]
- ^ a b Mikko Hellgren, Lars Sandberg, Olle Edholm (2006). "A comparison between two prokaryotic potassium channels (KirBac1.1 and KcsA) in a molecular dynamics (MD) simulation study". Biophys. Chem. 120 (1): 1–9. doi:10.1016/j.bpc.2005.10.002. PMID 16253415.
- ^ Relman, AS (1956). "The Physiological Behavior of Rubidium and Cesium in Relation to That of Potassium". The Yale journal of biology and medicine 29 (3): 248–62. PMC 2603856. PMID 13409924.
- ^ a b Meltzer, HL (1991). "A pharmacokinetic analysis of long-term administration of rubidium chloride". Journal of clinical pharmacology 31 (2): 179–84. doi:10.1002/j.1552-4604.1991.tb03704.x. PMID 2010564.
- ^ Pinsky, Carl; Bose, Ranjan; Taylor, J. R.; McKee, Jasper; Lapointe, Claude; Birchall, James (1981). "Cesium in mammals: Acute toxicity, organ changes and tissue accumulation". Journal of Environmental Science and Health, Part A 16 (5): 549– 567. doi:10.1080/10934528109375003.
- ^ Johnson, Garland T.; Lewis, Trent R.; Wagner, D. Wagner (1975). "Acute toxicity of cesium and rubidium compounds". Toxicology and Applied Pharmacology 32 (2): 239–245. doi:10.1016/0041-008X(75)90216-1. PMID 1154391.
- ^ Sartori H. E. (1984). "Cesium therapy in cancer patients". Pharmacol Biochem Behav 21 (Suppl 1): 11–13. doi:10.1016/0091-3057(84)90154-0. PMID 6522427.
- ^ Wood, Leonie. "'Cured' cancer patients died, court told". The Sydney Morning Herald. 20 November 2010.
- ^ Winter, Mark. "WebElements Periodic Table of the Elements | Francium | biological information". WebElements. Retrieved 15 February 2011.
Further reading[edit]
- Bauer, Brent A., Robert Houlihan, Michael J. Ackerman, Katya Johnson, and Himeshkumar Vyas (2006). "Acquired Long QT Syndrome Secondary to Cesium Chloride Supplement". Journal of Alternative and Complementary Medicine 12 (10): 1011–1014. doi:10.1089/acm.2006.12.1011. PMID 17212573.
- Campbell, Linda M., Aaron T. Fisk, Xianowa Wang, Gunter Kock, and Derek C. Muir (2005). "Evidence for Biomagnification of Rubidium in Freshwater and Marine Food Webs". Canadian Journal of Fisheries and Aquatic Sciences 62 (5): 1161–1167. doi:10.1139/f05-027.
- Chang, Cheng-Hung, and Tian Y. Tsong (2005). "Stochastic Resonance of Na, K-Ion Pumps on the Red Cell Membrane". AIP Conference Proceedings: 18th International Conference on Noise and Fluctuations 780. American Institute of Physics. pp. 587–590. doi:10.1063/1.2036821. ISBN 0-7354-0267-1.
- Erermis, Serpil, Muge Tamar, Hatice Karasoy, Tezan Bildik, Eyup S. Ercan, and Ahmet Gockay (1997). "Double-Blind Randomised Trial of Modest Salt Restriction in Older People". Lancet 350 (9081): 850–854. doi:10.1016/S0140-6736(97)02264-2. PMID 9310603.
- Joffe, Russell T., Stephen T. Sokolov and Anthony J. Levitt (2006). "Lithium and Triiodothyronine Augmentation of Antidepressants". Canadian Journal of Psychiatry 51 (12): 791–3. PMID 17168254.
- Krachler, M, and E Rossipal (1999). "Trace Elements Transfer From Mother to the Newborn – Investigations on Triplets of Colostrum, Maternal and Umbilical Sera". European Journal of Clinical Nutrition 53 (6): 486–494. doi:10.1038/sj.ejcn.1600781. PMID 10403586.
- Stein, Benjamin P., Stephen G. Benka, Phillip F. Schewe, and Bertram Schwarzhild (1996). "Physics Update". Physics Today 49 (6): 9. Bibcode:1996PhT....49f...9S. doi:10.1063/1.2807642.
External links[edit]
- "Group 1: The Alkali Metals". Visual Elements. Royal Society of Chemistry. Retrieved 8 December 2009.
- Atomic and Physical Properties of the Group 1 Elements An in-depth look at alkali metals
- Alkali Metal Bangs Filmed reactions of five-gram samples of the alkali metals with water
|Periodic table (Large version)