OCR GCSE Chemistry B (J258) — Making useful chemicals
Every time a chemical plant produces fertiliser, engineers are playing a high-stakes game of tug-of-war between speed, amount, and money.
In industrial processes like the Haber process, manufacturers want a high reaction rate to produce chemicals quickly and a high equilibrium yield to minimise wasted reactants. However, achieving both simultaneously is often impossible in a closed system at dynamic equilibrium.
According to collision theory, increasing the temperature increases the kinetic energy of particles. This means they move faster, resulting in more frequent collisions and a higher proportion of collisions exceeding the activation energy (). This creates a much faster rate of reaction.
However, we must also consider the equilibrium position. In many industrial processes, such as the Haber process, the forward reaction is an exothermic reaction. According to Le Chatelier’s principle, increasing the temperature will shift the equilibrium to the left to absorb the extra heat, significantly reducing the yield.
This creates a direct conflict: high temperatures give a fast rate but a low yield, whereas low temperatures give a high yield but a rate that is too slow to be profitable.
Increasing pressure forces gas particles closer together into a smaller volume. This leads to more frequent collisions per second, increasing the reaction rate.
Pressure also affects the equilibrium position if there is a different number of gas moles on each side of the equation.
In the Haber process equation above, there are 4 moles of reactant gas and 2 moles of product gas. High pressure shifts the equilibrium to the right (towards the side with fewer moles), which increases the yield.
Chemically, high pressure is ideal for both rate and yield. However, the trade-off is economic. Generating extremely high pressure requires vast amounts of energy for compressors and extremely expensive reinforced equipment to prevent explosion risks.
Because of these chemical and economic conflicts, industry relies on compromise conditions. These are conditions chosen to balance a high rate of production against a reasonable percentage yield, while remaining safe and financially viable.
For the Haber process, a compromise temperature of is used. This is high enough to achieve a fast rate, but low enough to maintain a yield of about 15–30% per pass. A pressure of is chosen because going higher would not generate enough extra profit to justify the massive equipment costs. To maximise efficiency, unreacted gases are recycled back into the reactor.
Adding a catalyst is a crucial part of chemical economics. A catalyst provides an alternative reaction pathway with a lower activation energy, increasing the rate of both the forward and reverse reactions equally.
Crucially, a catalyst does not change the final yield. Instead, it allows the system to reach equilibrium much faster. This means industrial plants can operate at lower temperatures, saving enormous amounts of energy and reducing environmental impact while still maintaining a profitable rate of production.
Students often state that adding a catalyst increases the equilibrium yield, but it actually only increases the speed at which the equilibrium yield is reached.
In 6-mark questions asking you to justify industrial conditions, you must explicitly use the word 'compromise' and link your points to both chemical factors (rate/yield) and economic factors (cost/safety).
When explaining the effect of temperature on rate using collision theory, always specify that collisions become 'more frequent' or happen 'more often per second'—do not just say 'more collisions'.
Always count the total moles of gas on each side of the balanced chemical equation before predicting how pressure will affect the equilibrium position.
Reaction rate
The speed at which reactants are converted into products, measured by the change in concentration of reactants or products over time.
Equilibrium yield
The amount of product present in the reaction mixture when a reversible reaction reaches dynamic equilibrium.
Dynamic equilibrium
A state in a closed system where the forward and reverse reactions occur at the same rate, resulting in no overall change in the concentrations of reactants and products.
Collision theory
The theory that for a reaction to occur, particles must collide with sufficient energy (activation energy) and the correct orientation.
Activation energy
The minimum amount of energy that colliding particles must possess for a successful reaction to occur.
Equilibrium position
The relative concentrations of reactants and products in a reversible reaction at dynamic equilibrium.
Exothermic reaction
A reaction that releases energy to the surroundings, meaning an increase in temperature will shift its equilibrium position to the reactant side.
Le Chatelier’s principle
If a system at equilibrium is subjected to a change in conditions, the system will shift its equilibrium position to counteract the change.
Compromise conditions
Reaction conditions (such as temperature and pressure) chosen to balance a high rate of production against a high percentage yield, while remaining safe and economically viable.
Catalyst
A substance that increases the rate of a chemical reaction without being used up by the overall reaction, by providing an alternative pathway with lower activation energy.
Put your knowledge into practice — try past paper questions for Chemistry B
Reaction rate
The speed at which reactants are converted into products, measured by the change in concentration of reactants or products over time.
Equilibrium yield
The amount of product present in the reaction mixture when a reversible reaction reaches dynamic equilibrium.
Dynamic equilibrium
A state in a closed system where the forward and reverse reactions occur at the same rate, resulting in no overall change in the concentrations of reactants and products.
Collision theory
The theory that for a reaction to occur, particles must collide with sufficient energy (activation energy) and the correct orientation.
Activation energy
The minimum amount of energy that colliding particles must possess for a successful reaction to occur.
Equilibrium position
The relative concentrations of reactants and products in a reversible reaction at dynamic equilibrium.
Exothermic reaction
A reaction that releases energy to the surroundings, meaning an increase in temperature will shift its equilibrium position to the reactant side.
Le Chatelier’s principle
If a system at equilibrium is subjected to a change in conditions, the system will shift its equilibrium position to counteract the change.
Compromise conditions
Reaction conditions (such as temperature and pressure) chosen to balance a high rate of production against a high percentage yield, while remaining safe and economically viable.
Catalyst
A substance that increases the rate of a chemical reaction without being used up by the overall reaction, by providing an alternative pathway with lower activation energy.
