A-level Chemistry/AQA/Module 2/Equilibria

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

Ammonia

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.

Methanol

Methanol is used mainly as a chemical feedstock, that is, as a starting material for making other chemicals particularly used in manufacturing methanal. Methanal is used in turn to make plastics e.g. Bakelite.

Indycars in the US are fueled by methanol which have an advantage over petrol fueled cars as methanol fires can be put out with water.

Every year 33 million tonnes of methanol are made worldwide mostly from the following reaction.

CO + 3H₂ ↔ CH₃OH ΔH = -91kJ mol⁻¹

The starting mixture is called systhesis gas and is made by reacting methane or propane with steam. Le Chatelier's principle tells us that the methanol sysnthesis reaction will give us the highest yield at a low temperature and high pressure. But compromise conditions are used in this case a temperature of 500K and a pressure of 10,000 kPa produces around 5-10% yield.

Ethanol

This is the alcohol in alcoholic drinks and therefore is unsurprising that it has been made for thousands of years by fermentation of sugars such as glucose while using the enzymes in yeast as a catalyst.

C₆H₁₂O₆ → 2C₂H₅OH + 2CO₂

Ethanol is also used in drugs, cosmetics, detergents and inks and at present the main source of ethanol is ethene made from crude oil obtained by fractional ditilation then cracking.

Ethanol is made by the reversible reaction of hydration (adding of water) to ethene and is speeded up by the catalyst of phosphoric acid absorbed on silica.

H₂C=CH₂ + H₂O ↔ CH₃CH₂OH ΔH = -46kJ mol⁻¹

The products and reactants are all gaseous at the temperature used. Applying Le Chatelier's principle to this equilibrium predicts that the maximum yield will be produced with a high pressure forcing the equilibrium to move to the right with fewer gas molecules, a low temperature forcing the equilibrium to the right to give out heat and exces steam to force equilibrium to the right again to reduce the steam concentration. However the practical problems include; the low temperature will reduce the reaction rate and thus how quickly the equilibrium is reached, the high pressure tends to cause ethene to polymerise (to polyethene) also the high pressure increases costs of building the plant and energy cost to run it and too much steam will simply dilute the catalyst.

In practice conditions of about 570K and 6500 kPa pressure are used obtaining a yield of ethanol at around 5%. The unreacted ethene is separated from the reaction mixture and recycled over the catalyst again until a 95% conversion is obtained.

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.