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# Andrew Rosen

1 = 1x1010
m

## Bond Energy: Energy required to break a bond

Chemical Bond: Results from simultaneous attraction of electrons for two nuclei
Bond formed = energy released
Bond broken = energy absorbed
Unbonded atoms have greater PE while bonded atoms have lower PE
Ionic Bonds (Electrovalent): Transfer of electrons from one valence to another
Think of properties of salt
Valence Electron: Electrons in the highest-energy electron shell of an atom
Shown by Lewis Dot Diagram
Metallic Bond: Metal atoms bonded to several neighboring metal atoms
The electrons are free to move within the 3-d structure
Octet Rule: Atoms gain or lose electrons to achieve the electron configuration of the
closest noble gas (until they are surrounded by 8 electrons)
Loss of electrons from an atom = endothermic and form a cation
Addition of electrons to an atom = exothermic and form an anion
Lattice Energy: Stability of a solid ionic substance (attracted due to opposite
electrostatic charges)
Energy required to separate a mole of a solid ionic compound into gaseous ions
E=k

Q1 Q2
R

## Q = charge for ion R = Distance between ion centers

k=

2.31x10-19 J*nm
Large ion charges and small distances between oppositely charged ions yield most
stable ionic solids
Formation of ionic compounds is exothermic and breaking up of ionic compounds is
endothermic
Energy for Formation of an Ionic Solid from its Ions:
Example: LiF(s)
Li(s) Li(g)

## *Must remove an electron so it must be in a gas state*

161 KJ

Andrew Rosen
Li(g) Li(g)+ + 1e-

## F2(g) F(g) *Must obtain single atom*77KJ

F(g) + 1e- Fe-(g)

## Li(s)+ + F(g)- LiF(s) -1047KJ

Li(s) + F2(g) LiF(s) *Sum up energies* -617KJ
Comparing Magnitudes of Lattice Energy:
Example: MgO and NaF

## MgO Lattice Energy of MgO 3916 KJ

=
=
4
NaF Lattice Energy of NaF 923 KJ
MgO (+2)(2)
=
=4
NaF (+1)(1)
Greater charges means higher energy
*If the charges are the same for two compounds, look at the atomic radii and see which
is smaller to find the one with more lattice energy
Isoelectronic Series: Ions consisting of the same number of electrons
The size of an ion in an isoelectronic series decreases with increasing atomic
number (nuclear charge)
When electrons are removed, the radius is smaller and when an electron is added,
Electronegativity: Ability of an atom in a molecule to attract electrons to itself
Partial charges are written with a

The more electronegative atom has a partial negative charge and vice versa
Partial charges can also be written with an arrow above the bond in the direction
towards the flow of electrons (tip on the partially negative side and end on the
partially positive side)
The greater the difference in electronegativity between two atoms, the more polar
the bond
Ionic bonds are formed due to a high electronegativity difference
Covalent bonds are formed due to a low electronegativity difference

Andrew Rosen
Ionization Energy: How strongly an atom holds on to its electrons
Electron Affinity: Measure of how strongly an atom attracts additional electrons
Dipole: Formed when two electrical charges of equal magnitude but opposite sign are
separated by a distance
Dipole Moment ( ): Quantitative measures of two electrical charges with equal
magnitude separated by a distance (r)

= Qr

1D = 3.34x10-30C*m
charge

## 1 e = 1.60x10-19 C for an electron

Q is magnitude of

Dipole moment increases with an increase in the quantity of charge separating the
atoms and distance between the charges
Covalent Bond: Sharing of electrons between two atoms
Think of properties of butter
How to Make Lewis Dot Structures:
2) Ions: If it is a positive ion, then subtract an electron. If there is a negative ion, then
3) Draw in the single bonds and find the electrons remaining
4) Make outer atoms have 8 valence
5) Make sure the total amount of electrons is correct
6) If the central atom has less than 8 valence, then create double/triple bonds if
possible
7) If central atom has more than 8 valence, then its okay
8) Dont forget to put a bracket and an overall charge if its an ion
Atoms with more than 8 valence will put the extra valence electrons in a d orbital
Degenerate: Equal energies (hybridized and natural state)
Resonance Structure:
Equivalent Lewis Structures
Double sided arrows are used
Formal Charges:
Charge an atom in a molecule would have if all atoms had the same
electronegativity
Is used to decide which resonance structure is most reasonable

Andrew Rosen
Number of valence electrons in the isolated atom minus the number of electrons
assigned to the atom in the Lewis Structure
Each bond counts as 1 electron for each atom
Formal charges that are closest to 0 are most common
The more electronegative atom should have a negative formal charge if none of
them zero out
*Compounds that follow the octet rule are most common despite the formal charge
rule
How to Make Lewis Dot Structures:
2) Ions: If it is a positive ion, then subtract an electron. If there is a negative ion, then
3) Draw in the single bonds and find the electrons remaining
4) Make outer atoms have 8 valence
5) Make sure the total amount of electrons is correct
6) If the central atom has less than 8 valence, then create double/triple bonds if
possible
7) If central atom has more than 8 valence, then its okay
* Can only find more than an octet for elements in the third period and below since there
needs to be an empty d orbital to accommodate the additional electrons
* The larger the central atom, the larger the number of atoms that can surround it, which
means a greater chance for more than an octet
* Smaller outside atoms with high ENs create more chance for greater than octet
molecules
Bond Enthalpy:
Energy needed to break 1 mole of bonds in the gaseous state
Bond Enthalpy is always positive since energy is required to break bonds

H rx n =

## D ( bonds broken ) (bonds formed )

Bond strength and bond enthalpy increase from single to double to triple bonds
As bond enthalpy increases, the bond length decreases
*If unsure what to do, break up all the bonds and they will net out in the end
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