Many reactions do not go to completion but reach a balance in which reactants and products coexist. This part explains reversible reactions and dynamic equilibrium, uses Le Chatelier's principle to predict how changing conditions shifts the balance, and applies these ideas to choosing conditions in important industrial processes.
Reversible reactions
A reversible reaction can proceed in both the forward and the reverse direction, shown by a double arrow such as in N2 + 3 H2 <=> 2 NH3. As reactants form products, the products can re-form the reactants, so the reaction does not go fully to completion in a closed system. A familiar example is the hydration and dehydration of copper(II) sulfate: blue hydrated crystals lose water on heating to form white anhydrous powder, which turns blue again when water is added. Because both directions occur, the composition of the mixture depends on the conditions. Reversibility is the foundation of equilibrium, and only in a closed system, where nothing escapes, can a true balance between the two directions be reached.
Dynamic equilibrium
When a reversible reaction is carried out in a closed system, it eventually reaches dynamic equilibrium. At this point the forward and reverse reactions occur at exactly the same rate, so the concentrations of all reactants and products remain constant even though both reactions continue. The word dynamic stresses that the reactions have not stopped; particles are still changing in both directions, but the changes cancel out. Equilibrium can be approached from either side and gives the same final composition under the same conditions. The position of equilibrium describes the relative amounts of reactants and products present: it lies to the right if products predominate and to the left if reactants predominate.
Le Chatelier's principle
Le Chatelier's principle predicts how a system at equilibrium responds to a change: if a change is made to a system at equilibrium, the position of equilibrium shifts in the direction that tends to oppose, or partly cancel, that change. This single rule lets you predict the effect of altering concentration, pressure or temperature. The system never fully undoes the change but shifts so as to reduce its impact. Applying the principle requires you first to identify what has changed, then to decide which direction of shift would oppose it. The principle is qualitative: it tells you the direction of the shift but not the exact new amounts, which would need the equilibrium constant.
Effects of concentration and pressure
Increasing the concentration of a reactant disturbs the balance, so the equilibrium shifts toward the products to use up the added reactant; removing a product likewise pulls the equilibrium forward. Decreasing a reactant concentration shifts it back toward the reactants. For reactions involving gases, increasing the total pressure (by reducing the volume) shifts the equilibrium toward the side with fewer moles of gas, since that reduces the pressure. If both sides have equal moles of gas, pressure has no effect on the position. For example, in N2 + 3 H2 <=> 2 NH3 there are four moles of gas on the left and two on the right, so higher pressure favours ammonia formation.
Effect of temperature and catalysts
Temperature is the only factor that changes the value of the equilibrium constant. Raising the temperature shifts the equilibrium in the endothermic direction, which absorbs the added heat, while lowering it favours the exothermic direction. For an exothermic forward reaction such as the formation of ammonia, raising the temperature shifts the equilibrium back toward reactants, lowering the yield, while cooling increases the yield. A catalyst, by contrast, does not change the position of equilibrium at all: it speeds up the forward and reverse reactions equally, so equilibrium is reached faster but the final amounts of reactants and products are unchanged. This distinction between rate and position is frequently tested.
Industrial conditions and compromise
Industrial processes must balance yield against rate and cost, leading to compromise conditions. In the Haber process, N2 + 3 H2 <=> 2 NH3 is exothermic with fewer gas moles on the product side, so a high yield needs high pressure and low temperature. However, a low temperature makes the reaction too slow, so a moderate temperature of around 450 degrees Celsius is chosen as a compromise that gives a reasonable yield at an acceptable rate, with an iron catalyst to speed equilibrium. A high pressure of about 200 atmospheres improves yield but is limited by the cost and danger of high-pressure equipment. Unreacted gases are recycled to improve overall efficiency. The Contact process for sulfuric acid uses similar reasoning.
Key terms
Reversible reaction
A reaction that can proceed in both forward and reverse directions, shown by a double arrow.
Closed system
A system from which no reactants or products can escape, allowing equilibrium to be reached.
Dynamic equilibrium
The state where forward and reverse reactions occur at equal rates and concentrations stay constant.
Position of equilibrium
A description of the relative amounts of reactants and products present at equilibrium.
Le Chatelier's principle
A change to a system at equilibrium shifts the position so as to oppose that change.
Equilibrium constant
A value, fixed at a given temperature, that expresses the ratio of products to reactants at equilibrium.
Yield
The amount of product obtained, which industrial conditions aim to maximise economically.
Compromise conditions
Conditions chosen to balance a reasonable yield with an acceptable rate and cost.
Haber process
The industrial synthesis of ammonia from nitrogen and hydrogen using an iron catalyst.
Contact process
The industrial manufacture of sulfuric acid involving the oxidation of sulfur dioxide.
Catalyst
A substance that speeds attainment of equilibrium without changing its position.
Endothermic direction
The direction of a reversible reaction that absorbs heat, favoured by raising temperature.
Exam technique
State that equilibrium is dynamic: the forward and reverse rates are equal and reactions continue.
When using Le Chatelier's principle, name the change first, then the direction of shift that opposes it.
For pressure changes, compare the number of moles of gas on each side; equal moles means no effect.
Remember only temperature changes the equilibrium constant; raising it favours the endothermic direction.
Explain a catalyst as speeding up equilibrium without altering the yield or position.
In industrial questions, justify conditions as a compromise between yield, rate and cost.
Quick check
For the exothermic equilibrium N2 + 3 H2 <=> 2 NH3, which change increases the yield of ammonia?
Raising the temperature, because the forward reaction is favoured by heat
Increasing the pressure, because the equilibrium shifts to the side with fewer gas moles
Adding a catalyst, because it shifts the equilibrium toward ammonia
Decreasing the pressure, because the gases spread out and react more
Show answer
Answer: B. There are four moles of gas on the left and two on the right, so higher pressure shifts the equilibrium toward ammonia to reduce the pressure. Raising temperature favours the endothermic reverse direction, and a catalyst only changes rate, not position.