A-level Chemistry/AQA/Module 2/Equilibria

Introduction

An equilibrium reaction is one in which the reactants, say A + B, collide successfully and react together to form the products, C + D. The 'reverse' also occurs simultaneously, in which C + D react to form A + B once again. When this system reaches a point in which the concentrations of A, B, C & D are constant, the reaction is said to be in Dynamic Equilibrium. At this point, the reaction appears to have stopped, but both the forward and backward reactions are still occurring at equal rates. However, it is very rare for this equilibrium position to be at the half-way point.

The above equilibrium reaction can be represented by the equation:

$A + B \leftrightarrow C + D$

Le Chatelier's Principle

Le Chatelier's principle can be used to predict the changes that will occur if the conditions are changed in a chemical equilibrium reaction. His principle states that:

"If a system at equilibrium is subject to a change in either pressure, temperature or concentration, then the system will move to oppose that change."

Using this principle is relatively straightforward; whatever you change in the reaction conditions, the system will move in the opposite direction to try and restore the equilibrium condition.

For example, let us consider the reaction in which Nitrogen Dioxide forms Dinitrogen Tetroxide

2NO₂ ↔ N₂O₄ ΔG = 45.53 kJ/mol

Temperature: It can be seen that the reaction here is endothermic (takes in heat from the surroundings). If the reaction conditions are changed, such that the temperature is increased, the equilibrium will attempt to oppose the change by decreasing the temperature and thereby shifting in the endothermic direction, to the RIGHT. Conversely, if the temperature is decreased, then the equilibrium will oppose the change by moving in the exothermic direction to raise the temperature once again, and hence the system will move to the LEFT.

Pressure: To determine the way an equilibrium reaction will shift when the overall pressure is changed, one must look at the number of moles of each species present in the reaction. In the above reaction concerning Nitrogen Dioxide it can be seen from the equation that the side with the greatest number of moles is the left. Therefore, increasing the pressure on this reaction will cause the equilibrium to move right, as this is the side with fewer moles of gas, which thereby reduces the pressure once again.

Concentration: If the concentration of one of the chemicals in an equilibrium reaction is changed, then the system will oppose the change by moving to either increase or decrease the concentration, depending on what change was made. Consider the following reaction:

H₂ + I₂ ↔ 2HI

If you were to add lots of Hydrogen Iodide to the above equilibrium, the system will move left to remove the excess HI concentration added. Likewise, if a large excess of Iodine was added, then the system would move to the right to reduce the concentration.

Equilibria & Industry

For many reactions, a compromise in reaction conditions must be met in order to achieve the best yield of product. Although, for example, a reaction may yield the greatest amount of product at extremely high temperatures, it is not feasible to generate such high temperatures in industry.

Consider the Haber Process, in which Hydrogen & Nitrogen react to form Ammonia:

3H₂ + N₂ ↔ 2NH₃ ΔG = -92.4 kJ/mol

The forward reaction is exothermic, and therefore the reaction is favoured by low temperatures. Even though this is the case, having temperatures that are too low will leave the reactant molecules with very little kinetic energy, and therefore collisions between the Hydrogen & Nitrogen will be very feeble and reaction will not occur, thus resulting in no Ammonia formation. It is therefore necessary for a compromise temperature of approx 450°C to be made, resulting in a mere 10-20% yield of Ammonia.

The Haber Process is also favoured by high pressures, since there are less moles of gas on the right of the equilibrium equation. Generating pressures that are extremely high is a difficult and costly process, since the energy required to generate such pressures can be high, as can the cost of the vessel used to withstand such high pressures. The compromise pressure here is 250atm.

Catalysts: Some reactions require the use of a catalyst in order to reach equilibrium in a feasible, efficient time. A catalyst does not affect the position of an equilibrium; it speeds up both the forward and backward reactions equally and therefore only increases the rate at which equilibrium is reached!

Catalysts work by providing an alternate reaction pathway of lower activation energy for both the reactants and products in an equilibrium reaction, and the activation energy remains unchanged.