1. 1.2.4 Hess’s Law

2. Hess’s Law
Many chemical reactions occur in a series of steps rather than a single step.
Ex. the following reaction describes the burning (combustion) of carbon:
(1)     C (s) + O2(g) → CO2(g)    ΔH = kJ

3. But this is not the only way the reaction can take place, it may take more than one step:
if not enough oxygen is present, CO is produced instead:
(2)    C (s) + ½ O2(g) → CO (g)    ΔH = kJ
If more oxygen is added, CO will undergo combustion with oxygen:
(3)    CO (g) + ½ O2(g) → CO2 (g)
 ΔH = kJ

4. Hess’s Law
Watch what happens if we add together the second and third reaction:
(2) C (s) + ½ O2(g) → CO (g) (3) CO (g) + ½ O2(g) → CO2 (g) (1) C (s) + O2(g) → CO2(g)
If reactions 2 and 3 occur, they will produce the same result as if just reaction 1 occurred

5. Hess’s Law Be sure you see how these equations can be added together
Add things that are on the same side of the equation:
½O2+ ½O2= 1 O2
and cross things out on opposite side of the equation (the CO)
(2) C (s) + ½ O2(g) → CO (g)
(3) CO (g) + ½ O2(g) → CO2 (g)
(1) C (s) + O2(g) → CO2(g)

6. Hess’s Law
Now compare the energy in the second and third reactions with the amount of energy released in the original reaction:
Reaction 2= kJ
Reaction 3= kJ
Reaction 1= kJ
The energy for reaction is equal to the amount of energy for reaction 1

7. Hess’s Law
It doesn't matter if the reaction proceeds all at once or in series of steps; the net energy change is the same. This is Hess's Law.

8. Hess's Law
Definition:
The enthalpy change for any reaction depends only on the energy states of the final products and initial reactants not the pathway or the number of steps between the reactant and product

9. Example
Given the intermediate steps in the production of tetraphosphorus decaoxide, P4O10, calculate ΔHf for P4O10
Given the following reactions:
(1) 4 P + 3 O2 → P4O6   ΔH = kJ (2) P4O6 + 2 O2 → P4O10   ΔH = kJ
Recall the heat of formation reaction involves the production of one mole of the compound from it's elements. Thus, we want to calculate ΔH for:
(3) 4 P + 5 O2  → P4O10

10. Tips for for Using Hess's Law:
Consider the intermediate steps involved in a reaction and ΔH values for each.
Reverse intermediate reactions and change the sign of ΔH if necessary.
Multiply intermediate reactions as necessary to balance the overall equation and multiply ΔH values as required.
Determine the final ΔH from the algebraic sum of ΔH values for the intermediate reactions.

11. Another Example:
You are given the following two reactions (Reactions 1 and 2):
(1)C2H2(g) + 5/2 O2(g) → 2 CO2(g) + H2O(l) ΔH = kJ
(2)C6H6(l) + 15/2 O2(g) → 6 CO2(g) + 3 H2O(l) ΔH = kJ
Find ΔH for the following reaction and tell whether it is exothermic or endothermic: (3)3 C2H2(g) → C6H6(g) 

12. Practice problem- 1. Given the following equations:
4 NH3(g) + 3 O2 (g) → 2 N2 (g) + 6 H2O(l) 
ΔH° = kJ
4 NH3(g) + 5 O2 (g) → 4 NO (g) + 6 H2O (l)  
ΔH° = kJ
Using these two equations, determine the heat of formation, ΔHf, for nitrogen monoxide, NO.
Try on your own then we’ll go over it together