Bond energy: The amount of energy needed to break 1 mole of bonds in a gaseous
covalent compound to form products (gas) at a constant temperature and pressure
Example: H2(g) + 436 kJ → 2H(g)
The greater the bond energy, the stronger the bond.
Breaking bonds is an endothermic process – energy is positive
Forming bonds is exothermic and the energy is negative
Bond energies are based on average energies of bonds from several compounds: C-H
based on bond energy from CH4, C2H6 etc.
Triple bonds > Double bonds > Single bonds
The heat of a reaction may be determined from bond energies:
ΔHrxn = ∑ B.E. reactants - ∑ B.E. products
Only defined for gas state!
Example: Determine the heat of the reaction using bond energies:
N2(g) + 3H2(g) → 2NH3(g)
1. Draw Lewis Structures
2. Determine how many and what type of bonds are broken
3. Look up bond energies and calculate
945kJ+ 3(436kJ) – 6(391kJ) = -93kJ/mol
Another examples:
Find H for the following reaction given the following bond energies:
Draw the Lewis Structures:
Bond
H-H
O=O
O-H
Bond
Energy
(kJ/mol)
436
499
463
What bonds are involved?
2 H-H bonds are broken.
1 O=O bond is broken
2 O-H bonds are formed per water molecule, and there are 2 water molecules formed,
therefore 4 O-H bonds are formed
Substitute the values given into the equation:
ΔHrxn = 2(436) + 499 – 4(463) = -481kJ
Lattice Energy: The energy change that occurs with the formation of one mole of
formula units in crystalline state from particles in the gaseous state
Always exothermic
M+ (g) + X- (g) → MX(s) + lattice energy
The energy is based on
1. The amount of charge on the ions
2. The distance between ionic nuclei (size of ions)
The greater the ionic charge and the smaller the ionic radius, the stronger the ionic bond
and larger the lattice energy