Chemical equilibria and Le Chatelier's principle — AQA A-Level Chemistry
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Chemical equilibria and Le Chatelier's principle explained
In a reversible reaction the products can re-form the reactants, and the equation is written with a double arrow.
Read the full explanation
Start with pure reactants in a closed container and the forward reaction is initially fast because their concentrations are high; as products build up the reverse reaction accelerates, until the two rates match and the composition stops changing. The consequence is that the reaction never goes to completion: some reactant always remains, and the yield falls short of 100 per cent. Esterification is a familiar case, where ethanol and ethanoic acid reach equilibrium with ethyl ethanoate and water, and the yield can be improved by removing the ester as it forms or by using an excess of one reagent rather than by waiting longer.
In a reversible reaction at equilibrium: forward and reverse reactions proceed at equal rates the concentrations of reactants and products remain constant
Equilibrium is dynamic, not static. Both reactions are still happening, but at exactly equal rates, so every molecule of product formed is matched by one converted back and the macroscopic composition is fixed. The two features examiners want are precisely these: the forward and reverse reactions proceed at equal rates, and the concentrations of reactants and products remain constant. The second point needs care. Constant does not mean equal: an equilibrium mixture can contain a great deal of product and very little reactant, and writing that the concentrations of reactants and products are the same is refused. A closed system is needed, since anything that escapes cannot come back, but the closed system is a condition rather than one of the two features.
Le Chatelier’s principle.
Le Chatelier's principle states that when a system at equilibrium is subjected to a change, the position of equilibrium shifts in the direction that opposes, and so partially counteracts, that change. It is a rule of thumb for predicting direction, not an explanation of mechanism: underneath, a change in conditions alters the forward and reverse rates by different amounts, and the composition drifts until they are equal again. The word partially matters, since the shift never fully cancels the change imposed. The principle applies only to a system already at equilibrium in a closed container, and it predicts direction, never how far or how fast. Rate and position are separate questions, which is why a catalyst, which changes rates, moves nothing.
Le Chatelier's principle can be used to predict the effects of changes in temperature, pressure and concentration on the position of equilibrium in homogeneous reactions.
Apply the principle one variable at a time. Raising the temperature shifts the equilibrium in the endothermic direction, absorbing energy to oppose the rise; cooling shifts it the exothermic way. So if the forward reaction is exothermic, a lower temperature raises the yield of product, and a question that reports a falling yield of a product on cooling is telling you the forward reaction is endothermic. Raising the pressure shifts the equilibrium towards the side with fewer moles of gas; if both sides have the same number, pressure has no effect. Increasing the concentration of a reactant shifts the equilibrium to the right, removing a product does the same. In a homogeneous system every species shares one phase, so all of them count.
A catalyst does not affect the position of equilibrium.
A catalyst lowers the activation energy of the forward and the reverse reactions by the same amount, so both are speeded up equally. Equilibrium is therefore reached sooner, but it is the same equilibrium: the position is unmoved, the composition at equilibrium is identical, the equilibrium constant is unchanged and the yield is exactly what it would have been. This follows from the definition of equilibrium as equal forward and reverse rates, since multiplying both rates by the same factor leaves them equal at the same composition. It is also why catalysts matter industrially. They let a plant work at a lower temperature, where the yield of an exothermic reaction is better, without the reaction becoming uneconomically slow.
Students should be able to: use Le Chatelier’s principle to predict qualitatively the effect of changes in temperature, pressure and concentration on the position of equilibrium explain why, for a reversible reaction used in an industrial process, a compromise temperature and pressure may be used.
Industrial conditions are settled by three competing demands: yield, rate and cost. In an exothermic synthesis such as ammonia, a low temperature gives the best equilibrium yield but the reaction would be far too slow, so a compromise temperature of about 700 K is used with an iron catalyst, accepting a lower yield in return for reaching it quickly. Pressure works the other way. Fewer moles of gas on the product side means high pressure improves the yield and also the rate, but the thicker vessels, the pumping costs and the safety risks set a practical ceiling, so a compromise pressure around 200 atmospheres is chosen. Unreacted gases are recycled, which recovers much of the yield lost to the compromise.
Your focus
- Explain why the percentage yield of a reversible reaction cannot reach 100 per cent even under ideal handling.
- Describe how the forward and reverse rates change from the moment of mixing until equilibrium is reached.
- Suggest two ways of increasing the yield of a named reversible reaction and justify each.
Show all 24 objectives
- Distinguish between yield lost to the position of equilibrium and yield lost during isolation and purification.
- Deduce from a graph of concentrations against time whether a mixture has reached equilibrium, and identify the time at which it did so.
- Explain why constant concentrations do not imply equal concentrations of reactants and products.
- Describe evidence that the reactions continue after equilibrium has been reached.
- Explain why a system open to the atmosphere cannot reach equilibrium when a gaseous product is formed.
- State Le Chatelier's principle in a form that includes opposing the change imposed on the system.
- Identify the imposed change in a described scenario before predicting the direction of any shift.
- Explain why the principle predicts direction only and says nothing about how quickly equilibrium is restored.
- Distinguish between a change that alters the position of equilibrium and one that alters only the rate.
- Predict the direction of shift and the effect on yield when the temperature of a named equilibrium is raised or lowered.
- Deduce the sign of ∆H for a forward reaction from reported changes in yield with temperature.
- Predict the effect of a pressure change by comparing the number of moles of gas on each side of the equation.
- Explain the effect on the position of equilibrium of removing a product as it is formed.
- Explain why speeding up both the forward and reverse reactions equally leaves the position of equilibrium unchanged.
- State the effect of a catalyst on the equilibrium yield, on the equilibrium constant and on the time taken to reach equilibrium.
- Identify the incorrect statement in a set of claims about catalysts in equilibrium systems and justify the choice.
- Explain why a catalyst allows an industrial equilibrium process to run at a lower temperature.
- Predict the temperature and pressure that would maximise the equilibrium yield of a named industrial reaction.
- Explain why the conditions actually used differ from those predicted by yield alone.
- Justify the use of a catalyst in an industrial equilibrium process in terms of rate and operating temperature.
- Explain how recycling unreacted gases compensates for the yield sacrificed by a compromise condition.
Chemical equilibria and Le Chatelier's principle exam tips
Marking Points
- one mark for stating that the reaction is reversible, so the products re-form the reactants
- one mark for explaining an incomplete yield by the reaction reaching an equilibrium rather than going to completion
- one mark for identifying a closed system as a requirement for equilibrium to be established
- one mark for suggesting a valid way to increase the yield, such as removing a product or using an excess of one reactant
- one mark for stating that the forward and reverse reactions proceed at equal rates
- one mark for stating that the concentrations of reactants and products remain constant
- one mark for identifying the equilibrium as dynamic, with both reactions still occurring
- one mark for recognising that a closed system is required so that no substance enters or leaves
- one mark for stating that the position of equilibrium moves to oppose the change imposed
- one mark for identifying the change imposed, in temperature, pressure or concentration, before predicting the direction
- one mark for expressing the shift as a direction, to the left or to the right, or towards a named side
- one mark for recognising that the principle predicts direction only, not the extent or the rate of change
- one mark for identifying whether the forward reaction is exothermic or endothermic when temperature is changed
- one mark for stating that the equilibrium shifts in the direction that opposes the change, quoting the direction
- one mark for counting the moles of gas on each side before predicting the effect of a pressure change
- one mark for the consequence for the yield of the named substance
- one mark for stating that a catalyst increases the rates of the forward and reverse reactions equally
- one mark for concluding that the position of equilibrium is unchanged
- one mark for stating that the equilibrium constant is unaffected
- one mark for identifying the benefit as equilibrium being reached more quickly, or a lower operating temperature being usable
- one mark for predicting the direction of shift and the effect on yield for the stated change
- one mark for identifying the conflict, that the conditions giving the best yield give an unacceptably slow rate
- one mark for naming the economic or safety limit on pressure, such as the cost of the plant or of compression
- one mark for stating that the compromise condition balances an acceptable yield against an acceptable rate and cost
Examiner Tips
- 💡If a percentage yield question mentions an ester or an ammonia synthesis, equilibrium is almost always one of the accepted reasons.
- 💡Use the double arrow consistently once you have described a system as reversible.
- 💡Distinguish practical losses, such as transfer and purification, from the chemical limit set by the equilibrium.
- 💡Learn the two features word for word; this question is asked in almost identical form year after year.
- 💡Write remain constant, never stay equal, when describing the concentrations.
- 💡If asked for evidence that an equilibrium is dynamic, refer to a labelled isotope appearing on both sides over time.
- 💡Use the verb shifts or moves, and name a direction; favours earns nothing.
- 💡State the change first, then the direction, then the consequence for the yield; that is the usual three mark structure.
- 💡Keep position of equilibrium and rate of reaction in separate sentences so the two ideas are not blurred.
- 💡Write the balanced equation and count the gas moles on each side before answering any pressure question.
- 💡For temperature, state the sign of ∆H for the forward reaction first, then the direction of the shift.
- 💡Give the effect and the explanation as two separate statements, since they are usually two separate marks.
- 💡Answer yield questions and rate questions separately; a catalyst belongs only in the rate answer.
- 💡If a multiple choice option says a catalyst has no effect on the equilibrium constant, that statement is correct.
- 💡Mention the industrial payoff, a lower temperature at an acceptable rate, when a question asks why a catalyst is used.
- 💡Structure a compromise answer in three moves: what the equilibrium wants, what the rate wants, what is chosen and why.
- 💡Name the cost or safety factor explicitly for pressure; a general statement about expense rarely scores.
- 💡Quote the sign of ∆H for the forward reaction before you argue about temperature.
Common Mistakes
- writing a single arrow in an equation for a reaction described as reversible
- explaining a low yield only as loss during transfer when the reaction itself is an equilibrium
- assuming a reversible reaction stops when equilibrium is reached
- claiming that leaving the reaction longer will push it to completion
- confusing reversibility with a reaction that can be undone by a different reagent
- writing that the concentrations of reactants and products are equal or are the same, which examiners refuse
- saying the reactions stop when equilibrium is reached
- offering a closed system as one of the two required features, where it is ignored as an answer
- stating equal amounts rather than constant concentrations
- claiming equilibrium means the reaction is exactly half complete
- saying the equilibrium favours a side, which examiners ignore, rather than shifts or moves in a direction
- writing that the system opposes the change completely, so the change is cancelled
- applying the principle to a system that has not yet reached equilibrium
- using it to predict a change in the rate of reaction rather than the position
- answering with the observed colour or yield without naming the direction of the shift
- shifting the equilibrium in the exothermic direction when the temperature is raised
- counting all moles, including liquids and solids, when deciding the effect of pressure
- predicting a pressure effect for a reaction with equal numbers of gaseous moles on each side
- saying the yield increases without naming the direction the equilibrium moved
- claiming that adding more of a reactant increases the equilibrium constant
- saying a catalyst increases the yield of product
- claiming it speeds the forward reaction more than the reverse
- asserting that it changes the value of the equilibrium constant
- stating that it makes the forward reaction more exothermic
- assuming a catalyst must be a solid in every equilibrium system
- arguing only from yield and concluding the plant should run cold, with no mention of rate
- saying a catalyst is used to improve the yield of the compromise process
- claiming high pressure is avoided because it reduces the yield
- forgetting that recycling unreacted gas raises the overall conversion
- treating the compromise temperature as the one giving the highest yield