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 enthalpies (4th 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 pp. 382-383 #69, 71-73

pp. 742-743 #19, 27