Reversible reactions and equilibria

Eventually, the forward and reverse reactions occur at equal rates. The system has reached dynamic equilibrium. ‘Dynamic’ means that both reactions are still happening: particles continually change from reactants into products and back again.

At equilibrium, each substance is being produced as quickly as it is being used up. Its concentration therefore remains constant, provided the conditions do not change. Constant concentrations do not mean equal concentrations: an equilibrium mixture can contain much more product than reactant, or the reverse.

Equilibrium requires a closed system, so substances cannot enter or leave. For example, if a gaseous product escapes from an open container, it is no longer available for the reverse reaction. Keeping the substances together allows the two reaction rates to become equal.

Making ammonia in the Haber process

The Haber process manufactures ammonia from nitrogen and hydrogen. The nitrogen is extracted from the air, and the hydrogen is obtained from natural gas.

The balanced equation is:

$$\mathrm{N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)}$$

One molecule of nitrogen reacts with three molecules of hydrogen to form two molecules of ammonia. The state symbol $(g)$ shows that all three substances are gases.

The reaction is reversible: ammonia can decompose back into nitrogen and hydrogen. In a closed system, the reaction can reach dynamic equilibrium. Ammonia is then being formed and broken down at equal rates, leaving a mixture of all three gases.

The conditions to recall are:

• 450 °C
• 200 atmospheres pressure
• an iron catalyst.

These are compromise conditions: they balance ammonia yield, reaction rate and production cost rather than giving the greatest possible yield. The iron catalyst speeds up both reaction directions, allowing equilibrium to be reached more quickly without changing the equilibrium position.

Predicting changes in equilibrium position — Higher tier

The position of equilibrium describes the relative concentrations of reactants and products at equilibrium. A shift to the right favours products; a shift to the left favours reactants.

Le Chatelier’s principle gives a way to predict a shift: when conditions affecting an equilibrium change, the position shifts to counteract that change. After the shift, a new dynamic equilibrium can be established. The forward and reverse rates are equal again, but the mixture has different concentrations.

Temperature: favouring the direction that opposes heating or cooling — Higher tier

An exothermic reaction transfers energy to the surroundings; an endothermic reaction takes energy in. In a reversible reaction, if one direction is exothermic, the opposite direction is endothermic.

• Increasing temperature favours the endothermic direction, which takes in energy.
• Decreasing temperature favours the exothermic direction, which releases energy.

For the Haber reaction, ammonia formation is exothermic. Increasing temperature therefore shifts equilibrium to the left and reduces the equilibrium yield of ammonia. Decreasing temperature shifts it to the right and increases the yield.

However, a lower temperature also makes the reaction slower. The temperature of 450 °C is a compromise between obtaining a useful ammonia yield and producing it at a useful rate. This illustrates why reaction rate and equilibrium yield are different considerations.

Pressure: counting gaseous particles — Higher tier

For gaseous equilibria, increasing pressure favours the side with fewer gaseous particles. Decreasing pressure favours the side with more gaseous particles.

Use the coefficients in the balanced equation to compare the two sides. In the Haber equation, the left has four gaseous molecules in total: one nitrogen plus three hydrogen. The right has two ammonia molecules.

Increasing pressure therefore shifts equilibrium to the right, increasing the ammonia yield. Decreasing pressure shifts it to the left. If both sides of an equation have the same number of gaseous particles, a pressure change does not favour either side. Count gases only, not solids or liquids.

Concentration: using up what is added — Higher tier

Changing a concentration shifts equilibrium towards using up an added substance or replacing a removed substance.

For the Haber reaction:

• Adding nitrogen or hydrogen shifts equilibrium to the right, using up some of the added reactant and forming more ammonia.
• Adding ammonia shifts equilibrium to the left, using up some ammonia and forming nitrogen and hydrogen.
• Removing ammonia shifts equilibrium to the right, replacing some of the ammonia removed.

The shift counteracts the change; it does not necessarily restore the original concentrations.

• Nitrogen: extracted from air.
• Hydrogen: obtained from natural gas.
• Conditions: 450 °C, 200 atmospheres, iron catalyst.
• These balance yield, rate and cost.
• Catalyst: faster attainment of equilibrium, unchanged equilibrium position.

Equilibrium shifts — Higher tier

Le Chatelier’s principle: the position shifts to counteract a change.

Equal numbers of gaseous particles on both sides: no pressure-induced shift.

Haber application: the forward reaction is exothermic; four gaseous molecules become two. Lower temperature, higher pressure or adding a reactant favours ammonia formation.