Help your learners grasp equilibrium constants

To properly understand the sustainable manufacture of many products in industry, students need to comprehend the concepts associated with chemical equilibrium. When reactions are reversible, for example, it’s important for manufacturers to optimise reaction conditions to produce as much of the desired product as possible. They must also minimise the quantity of starting materials, identify reactions with high atom economy and limit the amount of waste material. Manufacturers also need to consider the energy demands of high temperature and pressure reactions.

A well-known and commonly studied industrial reaction is the production of ammonia from hydrogen and nitrogen, the Haber–Bosch process. Industrial production of ammonia allows the mass manufacture of chemical fertilisers, which in turn increases crop yields and helps to meet the needs of a growing population. The process also provides a potential ethical dilemma because ammonia is a precursor for the synthesis of numerous explosives.

What students need to know

We can use Le Chatelier’s principle to make simple qualitative predictions about how equilibrium reactions will respond to changes in conditions. The principle is that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change and re-establish an equilibrium. To put this more simply, when changes are imposed on a system, the position of equilibrium alters to oppose this change. This principle allows us to predict the effects of changes in temperature, pressure and concentration on the yield of a reversible reaction.

At 14–16, students learn the qualitative interpretation of Le Chatelier’s principle, where they can easily develop misconceptions. For example, learners can wrongly think an exothermic reaction would be favourable after an increase in temperature because releasing energy would minimise the system change.

Use reversible chemical reactions with visible changes to show students how the yield is affected

Also, students are often taught rates of reaction directly before equilibrium. When describing situations, they confuse the rules for rate with equilibrium, causing further misunderstanding. You can separate these topics to minimise this problem.

Bridging the gap to post-16 study

Post-16 courses recap the principles studied at 14–16. Students then move on to a more quantitative understanding of chemical equilibrium. Central to this is the definition and calculation of equilibrium constants, and how these constants help describe positions of equilibrium.

A key point about the equilibrium constant is that it is constant at all concentrations and pressures, and only temperature affects this value. The equilibrium constant, in terms of concentration, for the reaction aA(g) + bB(g) ⇌ cC(g) + dD(g) is defined by the law of chemical equilibrium as:

Developing the equation

In order to use this equation, the reaction must be homogeneous (ie all substances must be in the same phase), there must be dynamic equilibrium, and the concentrations of each substance at equilibrium must be clear. The value of K_c determines the favoured side of the equation, and therefore the relative proportion of species in the mixture. So, if K_c is much greater than 1, the mixture contains mainly products; if K_c equals 1, the equilibrium will lie midway between reactants and products; and if K_c is much less than 1, the mixture contains mainly reactants.

We can use the Haber–Bosch process in a worked example to show this. The task could be to determine the equilibrium constant for a mixture of 0.030 mol dm^-3 N₂, 0.037 mol dm^-3 H₂ and 0.016 mol dm^-3 NH₃ at a constant temperature, given the equilibrium N₂(g) + 3H₂(g) ⇌ 2NH₃(g).

When conditions change in the reaction, students can then use these mathematical expressions to explain how the change in position of equilibrium occurs. For example, if ammonia was removed from the equilibrium, then the value of the numerator of the K_c expression decreases. This causes the value of the K_c to decrease. To restore K_c to its original constant value, the numerator must increase, and the denominator decrease. In other words, the concentrations of N₂ and H₂ must decrease (be used up) and increase for NH₃ (be produced). So, the forward reaction is favoured, the position of equilibrium shifts right and the dynamic equilibrium is re-established.

This change can also be described in terms of rates of reaction. When a closed system is in dynamic equilibrium, the rates of the forward and backward reaction are equal. However, when the concentration of NH₃ is decreased, the rate of the reverse reaction decreases. In that instant, however, the rate of the forward reaction remains the same. The forward reaction is then higher than the reverse, which increases the concentration of NH₃ and re-establishes the equilibrium.