Equilibrium position

Step 1: Count the moles of gas on each side of the balanced equation.

• Left side (reactants): $1 + 2 = 3$ moles of gas
• Right side (products): $1$ mole of gas

Step 2: Apply Le Chatelier's Principle for pressure.

• Decreasing the pressure causes the system to shift to the side with more gas molecules to oppose the change.

Step 3: State the final conclusion.

• The equilibrium shifts to the left, meaning the amount of gas C produced decreases.

The Role of Catalysts

You can snap a piece of chalk easily, but try snapping a diamond — some reactions naturally need a lot more energy to get started than others. The minimum amount of energy that particles must have to react when they collide is called the activation energy ($E_a$).

• A catalyst is a substance that speeds up a reaction without being used up by providing an alternative pathway with a lower activation energy.
• Because the activation energy is lower, a higher proportion of particle collisions are successful.

In a reversible reaction, a catalyst increases the rate of the forward and reverse reactions equally. This means a catalyst has no effect on the position of equilibrium or the final yield (the amount of product obtained). However, it does allow the system to reach dynamic equilibrium much faster.

Choosing Compromise Conditions

Industrial chemical production is a careful balancing act between making as much product as possible and making it quickly enough to be profitable. Compromise conditions are chosen to provide the best balance between a high rate of product formation, a high percentage yield, low production costs, and safe operation.

• The Haber Process: Produces ammonia using the exothermic reaction $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$. A low temperature would give a high yield, but the rate would be far too slow. A temperature of 450°C is chosen as a compromise. A high pressure of 200–250 atm increases both the yield and the rate, balancing production with safety limits. An iron catalyst speeds up the process, and unreacted gases are recycled.
• Hydration of Ethene: Produces ethanol via the reaction $C_2H_4(g) + H_2O(g) \rightleftharpoons C_2H_5OH(g)$. It uses a temperature of 300°C, a pressure of 60–70 atm, and a phosphoric acid catalyst. The pressure is limited to manage costs and prevent the ethene from spontaneously polymerising.
• The Contact Process: In the stage producing sulfur trioxide ($2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)$), a vanadium(V) oxide catalyst and a temperature of 450°C are used. The pressure is kept very low (1–2 atm) because the yield is already 96–99%, and higher pressure would be an unnecessary expense that increases the risk of leaking corrosive gas.

Economic and Safety Drawbacks

Every time you boil a kettle, you are paying for the electrical energy used — scaling this up to industrial chemical plants involves massive costs. Operating at extreme temperatures and pressures comes with significant drawbacks:

• Economic Costs for Temperature: Maintaining high temperatures requires massive continuous energy input, which increases running costs. High temperatures can also decrease the yield of exothermic reactions.
• Economic Costs for Pressure: High pressures require specialised, thick-walled reaction vessels and reinforced stainless steel piping, which massively increases capital costs (the initial expense of building the plant). Running powerful compressors also adds high running costs.
• Safety Risks: High pressures carry a catastrophic risk of explosion if the equipment fails. High temperatures cause thermal stress and metal fatigue, leading to leaks and severe burn hazards. Extreme conditions can also make chemicals more dangerous, such as sulfur trioxide becoming more corrosive.