Bond Enthalpies
How does a chemical reaction have energy?
Bond Energy
 Energy required to make/break a chemical bond
 Endothermic reactions
 Products have more energy than reactants
 More energy to BREAK bonds
 Exothermic reactions
 Reactants have more energy than products
 More energy to FORM bonds
Bond Enthalpy
 Focuses on the energy/heat between products and reactants as
it relates to chemical bonding
 Amount of energy absorbed to break a chemical bond--amount of energy released to form a bond.
 Multiple chemical bonds take more energy to break and release
more energy at formation
 Amount of energy absorbed = amount of energy released
to break chemical bond
to form a chemical bond
Calculating ΔHrxn. by bond
th
enthalpies (4 method)
 Least accurate method
 ΔH = ΣBE (bonds broken)
- ΣBE (bonds formed)
Example 1:
 Using average bond enthalpy data, calcaulate ΔH for the
following reaction.
 CH4 + 2O2  CO2 + 2H2O
ΔH = ?
Bond
Average Bond Enthalpy
C-H
413 kJ/mol
O=O
495 kJ/mol
C-O
358 kJ/mol
C=O
799 kJ/mol
O-H
467 kJ/mol
Entropy
Spontaneous vs.
Nonspontaneous
1) Spontaneous Process
 Occurs WITHOUT help outside of the system, natural
 Many are exothermic—favors energy release to create an
energy reduction after a chemical reaction

Ex. Rusting iron with O2 and H2O, cold coffee in a mug
 Some are endothermic

Ex. Evaporation of water/boiling, NaCl dissolving in water
Spontaneous vs.
Nonspontaneous
2) Nonspontaneous Process

REQUIRES help outside system to perform chemical
reaction, gets aid from environment
 Ex. Water cannot freeze at standard conditions (25°C, 1atm),
cannot boil at 25°C
**Chemical processes that are spontaneous
have a nonspontaneous process in reverse **
Entropy (S)
 Measure of a system’s disorder
 Disorder is more favorable than order
 ΔS = S(products) - S(reactants)
 ΔS is (+) with increased disorder
 State function
 Only dependent on initial and final states of a reaction
 Ex. Evaporation, dissolving, dirty house
Thermodynamic Laws
1st Law of Thermodynamics
 Energy cannot be created or destroyed
2nd Law of Thermodynamics
 The entropy of the universe is always increasing.
 Naturally favors a disordered state
When does a system become MORE
disordered from a chemical reaction?
(ΔS > 0)
1) Melting
2) Vaporization
3) More particles present in the products than the reactants
 4C3H5N3O9 (l)  6N2 (g) + 12CO2 (g) + 10H2O (g) + O2
(g)
4) Solution formation with liquids and solids
5) Addition of heat
When does a system become LESS
disordered from a chemical reaction?
(ΔS < 0)
1) Solution formation with liquids and gases
3rd Law of Thermodynamics
The entropy (ΔS) of a perfect crystal is 0 at a
temperature of absolute zero (0°K).
 No particle motion at all in crystal structure
 All motion stops
How do we determine if a chemical
reaction is spontaneous?
1) Change in entropy (ΔS)
2) Gibbs Free Energy (ΔG)
Change in entropy (ΔS)
 For a chemical reaction to be spontaneous (ΔST > 0), there
MUST be an increase in system’s entropy (Δssys> 0) and the
reaction MUST be exothermic (Δssurr > 0).
 Exothermic reactions are favored, NOT endothermic reactions.
 Exothermic (ΔH < 0, ΔS > 0)
 Endothermic (ΔH > 0, ΔS < 0)
 ΔST = Δssys + Δssurr
 If ΔST > 0, then the chemical reaction is spontaneous
Example 1:
Will entropy increase or decrease for the following?
a) N2 (g) + 3H2 (g)  2NH3 (g)
b) 2KClO3 (s)  2KCl (s) + 3O2 (g)
c) CO(g) + H2O(g) 
CO2 (g) + H2 (g)
d) C12H22O11 (s)  C12H22O11
How do we calculate the entropy change
(ΔS) in a chemical reaction?
 Same method as using the enthalpies of formation to
calculate ΔH and use the same table.
 aA + bB 
cC + dD
ΔS° =[c (ΔS°C) + d(ΔS°D)] - [a (ΔS°A) + b (ΔS°B)]
Example 2: Calculate ΔS° for the
following reaction at 25°C….
4HCl(g) +
O2 (g) 
2Cl2 (g) + 2H2O (g)
Homework
 Finish problems #16-19 on enthalpy worksheet.
 pp. 742-473 #19, 23, 26, 27