Chemical Equilibrium

Equilibrium is a concept that permeates all of science. Thermal equilibrium is a simple example found in physics. Biologists, such as my roommate, Sapphire, think of homeostasis. Yet I, even before reaching college, when hearing the term, always thought of chemical equilibrium, a concept that can be both extremely simple and absurdly complicated.

Chemical equilibrium is an interesting phenomenon that describes the state of a reaction involving solutions or gases at its “end.” (Pure solids and liquids don’t apply here because reactions between these simply proceed until one reactant is completely used up. This reactant, by the way, is called the limiting reactant.) Equilibrium is the reason that, when reactions involve gas-phase or aqueous reactants, reactions don’t proceed to completely convert reactants to products.

The reason behind this has to do with concentrations and rates of reaction. With aqueous or gaseous reactants, as a reaction takes place, the concentrations of the reactants decrease (at a rate proportional to the rate of the reaction). As concentrations lower, the rate at which the “forward” (or reactants-to-products) reaction takes place decreases. Meanwhile, as concentrations of products increase, the rate at which the “reverse” (or products-to-reactants) reaction takes place increases. Eventually these two reactions will reach at a point where they are both proceeding at the same rate, and thus will balance each other out exactly. At this point, concentrations of reactants and products stop changing, and the reaction appears to have stopped. This is where the reaction is said to have reached equilibrium.

Why is this important, you ask? My chemistry professor would push his sleeves back, gesture at one side of the room, and say, “Half of you would be dead if we didn’t understand equilibrium.” I can safely say that when he pulled this on us, we all payed closer attention to what he said for the next few minutes. He was, naturally, correct.

Different reactions reach equilibrium with different concentrations of reactants and products. Some reactions produce mostly products—these are said to be product-favored. Reactant-favored reactions, on the other hand, shift only a little toward products, with mostly reactants remaining after equilibrium is reached. However, different conditions can be changed to shift equilibrium closer to one side or another. This is important to note from a commercial standpoint, as, in industry, producing the maximum amount of products is always favorable.

All right, but why would manipulating a reaction into producing more products contribute to many of us still being alive? Put quite simply, knowing how to do this helps us make fertilizer. College textbooks are extremely fond of pointing this out. Look at the following reaction:

N2 (g) + 3H2 (g) ⇌ 2NH3 (g)

This is the process by which nitrogen and hydrogen gas, two abundant gases, are combined to make ammonia. Ammonia is important in fertilizer because it is a usable form of nitrogen for plants. (Plants need nitrogen, but with the exception of a few that can “fix” it, they can’t pull it from our atmosphere.) Fritz Haber received a Nobel Prize for his groundbreaking work with this reaction. Haber determined the conditions that must be met for the reaction to produce the maximum amount of products, which is, of course, important when you’re dealing with fertilizer for food.

Changing different conditions such as temperature, pressure and concentrations of reactants and products can shift the equilibrium of a reaction one way or another. The ways each of these things can change equilibrium in a number of cases are summarized with Le Chatelier’s Principle, a wibbly-wobbly principal that causes joy and sorrow.


Did I make things needlessly confusing? Feel free to comment and ask a question. I’ll do my best.

Correct me if you think I’ve missed something important, but remember, I’m sensitive, so be nice.

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