Matching rates of change in the Haber process

My favorite reaction is the Haber process, the synthesis of ammonia from nitrogen and hydrogen gases. After balancing the reaction (see the section “Balancing the Haber process,” earlier in this chapter), you end up with

N_2(g) + 3H_2(g) → 2NH_3(g)

Written this way, the reaction says that hydrogen and nitrogen react to form ammonia — and this keeps on happening until you use up one or both of the reactants. But this isn’t quite true.

If this reaction occurs in a closed container (which it has to, with everything being gases), then the nitrogen and hydrogen react and ammonia is formed — but some of the ammonia soon starts to decompose into nitrogen and hydrogen, like this:

2NH_3(g) → N_2(g) + 3H_2(g)

In the container, then, you actually have two exactly opposite reactions occurring — nitrogen and hydrogen combine to give ammonia, and ammonia decomposes to give nitrogen and hydrogen.

Instead of showing the two separate reactions, you can show one reaction and use a double arrow like this:

N_2(g) + 3H_2(g) ↔ 2NH_3(g) + heat

You put the nitrogen and hydrogen on the left because that’s what you initially put into the reaction container. The reaction is exothermic (it gives off heat), so I’m showing heat as a product of the reaction on the right.  Now these two reactions occur at different speeds, but sooner or later, the two speeds become the same, and the relative amounts of nitrogen, hydrogen, and ammonia become constant.

This is an example of a chemical equilibrium. At any given time in the Haber process, you have nitrogen and hydrogen reacting to form ammonia and ammonia decomposing to form nitrogen and hydrogen. When the system reaches equilibrium, the amounts of all chemical species become constant but not necessarily the same.