Now that you've gone through the ordeal of rates of reaction, we'll move on to something substantially easier: Hess' Law. This law simply states that when two reactions are added, other properties (such as ΔH) can be added as well. For example, the following two example reactions can be added as shown:
Reaction 1: A (s) + B (g) 2 C (g) ΔH°1 = -14.8 kJ
Reaction 2: 2 C (g) + D (g) 2 E (g) ΔH°2 = 3.4 kJ
Overall: A (s) + B (g) + D (g) 2 E (g) ΔH° = -11.4 kJ
Simply add the two ΔH° values together for the overall heat change. Equilibrium constants can also be combined for two reactions, but they cannot be added; the following example shows why:
Reaction 1: A (g) + 2 B (s) C (g) K1 = [C]/[A]
Reaction 2: 2 D (g) + E (g) F (g) + G (s) K2 = [F]/([D]2[E])
Overall: A (g) + 2 B (s) + 2 D (g) + E (g) C (g) + F (g) + G (s) Knet = ([C][F][G])/([A][D]2[E])
While the individual species are added to come up with the final reaction, the equilibrium constant multiplies the concentrations of these reactions together. Therefore, to find the equilibrium constant of two or more added reactions, you must multiply the K values of the component reactions together.
As explained in the previous page, the rate equation for two separate reactions being added together is dependent on the slower, rate-determining reaction. Since you're probably very tired of rate laws, we won't go over this material again!
That's all there is to Hess' Law! Next, we'll move on to the final lesson of our chapter, Catalysts. These are substances that accelerate a reaction without being consumed, and they are vital not only to industrial chemistry, but for the survival of life on earth as well.