Equilibrium position

In a reversible reaction, products can react to form reactants again. At dynamic equilibrium, the forward and reverse reactions continue at equal rates, so the amounts of substances remain constant. Those amounts need not be equal. The equilibrium position describes their relative amounts: an equilibrium lying towards the products gives a greater equilibrium yield of the desired product.

Changing conditions can disturb an equilibrium. The reaction then shifts in the direction that partly counteracts the change, until a new equilibrium is established. To predict whether a product yield rises or falls, first decide which direction is favoured, then find the product on the equation.

Changing concentration

If you add more of a reactant to an equilibrium mixture, the reaction shifts in the direction that uses some of it up. For a reaction written $A + B \rightleftharpoons C$, adding $A$ favours the forward reaction, so more $C$ forms. Removing $C$ also favours the forward reaction, because the mixture responds by replacing some of the product removed. Conversely, adding $C$ or removing a reactant favours the reverse reaction. These predictions concern substances whose concentrations can meaningfully be changed in the reacting mixture, such as dissolved substances or gases.

Changing temperature

The two directions of a reversible reaction have opposite energy changes: if the forward reaction is exothermic and releases heat, the reverse reaction is endothermic and takes in heat. Increasing temperature favours the endothermic direction, which takes in some of the added energy. Decreasing temperature favours the exothermic direction. You therefore need to know which direction is exothermic before predicting what happens to the yield of a particular product.

For instance, if the forward reaction makes the desired product and is exothermic, cooling shifts the equilibrium towards that product. Heating shifts it towards the reactants. Yet cooling usually slows the reaction. A temperature that gives the greatest equilibrium yield may therefore be too slow to make the product at a useful rate.

Changing pressure

Pressure matters to equilibrium position when gases react and the two sides of the balanced equation contain different total numbers of gaseous particles. Increasing pressure favours the side with fewer gaseous particles; decreasing pressure favours the side with more. Count the gaseous particles using the numbers in front of their formulae, not the number of different substances. If both sides have the same total, a pressure change does not favour either side.

For example, $2\mathrm{NO_2}(g) \rightleftharpoons \mathrm{N_2O_4}(g)$ has two gaseous particles on the left for every one on the right. Increasing pressure shifts the equilibrium right, increasing the equilibrium yield of $\mathrm{N_2O_4}$. Lowering pressure favours the left.

Choosing useful conditions

Chemists consider yield and rate together when choosing conditions to produce a particular product. For ammonia, $\mathrm{N_2}(g)+3\mathrm{H_2}(g)\rightleftharpoons2\mathrm{NH_3}(g)$, the forward reaction is exothermic. A lower temperature favours ammonia at equilibrium, while a higher pressure favours it because four gaseous particles on the left form two on the right. But a very low temperature makes production slow. Very high temperatures or pressures cost energy to maintain, and extreme conditions can put equipment under strain or cause it to fail. Chemical engineers therefore choose workable conditions rather than simply those giving the highest possible equilibrium yield.

A catalyst offers a different benefit. It speeds up both the forward and reverse reactions, so equilibrium is reached sooner. It does not shift the equilibrium position or increase the equilibrium yield. This distinction matters when suggesting conditions: a catalyst can improve the rate without sacrificing yield, whereas changing temperature or pressure may alter both the yield and the practical cost of production.