1. Section 11.4 Calculating Heat Changes
OBJECTIVES:
Apply Hess’s law of heat summation to find heat changes for chemical and physical processes.

2. Section 11.4 Calculating Heat Changes
OBJECTIVES:
Calculate heat changes using standard heats of formation.

3. Hess’s Law
If you add two or more thermochemical equations to give a final equation, then you can also add the heats of reaction to give the final heat of reaction.
Called Hess’s law of heat summation
Example shown on page 314 for graphite and diamonds

4. Why Does It Work? If you turn an equation around, you change the sign:
If H2(g) + 1/2 O2(g)→ H2O(g) ΔH= kJ
then, H2O(g) → H2(g) + 1/2 O2(g) ΔH = kJ
also,
If you multiply the equation by a number, you multiply the heat by that number:
2 H2O(g) → 2 H2(g) + O2(g) ΔH = kJ

5. Why does it work?
You make the products, so you need their heats of formation
You “unmake” the products so you have to subtract their heats.
How do you get good at this?

6. Standard Heats of Formation
The ΔH for a reaction that produces 1 mol of a compound from its elements at standard conditions
Standard conditions: 25°C and 1 atm.
Symbol is
The standard heat of formation of an element = 0
This includes the diatomics

7. What good are they? ΔHo = ( Products) - ( Reactants)
Table 11.6, page 316 has standard heats of formation
The heat of a reaction can be calculated by:
subtracting the heats of formation of the reactants from the products
ΔHo
= (
Products) - (
Reactants)

8. Examples CH4 (g) = - 74.86 kJ/mol O2(g) = 0 kJ/mol
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
CH4 (g) = kJ/mol
O2(g) = 0 kJ/mol
CO2(g) = kJ/mol
H2O(g) = kJ/mol
ΔH= [ (-241.8)] - [ (0)]
ΔH= kJ

9. Examples
Page 318, #32
Page 318, #33