Revise Chemical equilibria: reversible reactions, dynamic equilibrium for Cambridge 9701 (7.1)

A reversible reaction is one where the products can react to re-form the original reactants. We use a special double arrow symbol ($⇌$) to show this. For example, the famous Haber process for making ammonia: $N_2(g) + 3H_2(g) ⇌ 2NH_3(g)$. Initially, nitrogen and hydrogen react to form ammonia (the forward reaction). As ammonia molecules form, they start to decompose back into nitrogen and hydrogen (the reverse reaction). Eventually, the rates of these two opposing reactions become equal. At this point, the system has reached dynamic equilibrium.

If we have a system at equilibrium, what happens if we change the conditions? Le Chatelier's principle provides the answer. It states that if a change (a 'stress') is applied to a system at equilibrium, the position of equilibrium will shift to counteract that change. It's like the system 'fights back' to restore a balance. We can use this principle to predict how changes in concentration, pressure, and temperature will affect the yield of products.

Let's apply Le Chatelier's principle to the three main factors you need to know for your exam.

The principles of equilibrium are crucial for industrial chemists who want to maximise the yield of a desired product. The Haber process is a classic example. To maximise ammonia yield, Le Chatelier's principle suggests using high pressure (shifts equilibrium right) and low temperature (forward reaction is exothermic). However, there's a compromise. Low temperatures make the reaction rate incredibly slow. Therefore, a moderately high temperature (e.g., 400–450 °C) and a catalyst are used to achieve an acceptable rate, even though this slightly reduces the maximum possible yield. This balance between yield, rate, and cost is a key consideration in all industrial processes.