Changing Equilibrium Conditions Higher
The relative amounts of reactants and products at equilibrium — the position of equilibrium — depend on the conditions. Change the conditions and the position moves. Higher Tier students must predict which way, using one beautifully simple rule:
If a system is at equilibrium and a change is made to any of the conditions, the system responds to counteract the change.
Whatever you do to an equilibrium, it shifts in the direction that partly undoes what you did. Add more of something and the equilibrium uses it up; heat it and the equilibrium absorbs the heat; increase the pressure and the equilibrium reduces the number of gas molecules. You are only asked for qualitative predictions, and the question will always give you the information you need — the equation, and which direction is exothermic.
The three changes
1 · Concentration. If the concentration of one substance is changed, the system is no longer at equilibrium, and the concentrations of all the substances change until equilibrium is reached again:
- Increase a reactant’s concentration → the equilibrium shifts to use it up → more products form until equilibrium is re-established.
- Decrease a product’s concentration (e.g. keep removing it) → the equilibrium shifts to replace it → more reactants react. This is how industry drags an equilibrium toward the product it wants.
2 · Temperature. Heat is “counteracted” by the direction that absorbs it — the endothermic direction; cooling favours the exothermic direction:
| Change | Favoured direction | Result |
|---|---|---|
| Increase temperature | Endothermic direction (absorbs the extra heat) | If the forward reaction is endothermic → more products. If it’s exothermic → fewer products. |
| Decrease temperature | Exothermic direction (releases heat to replace what was lost) | If the forward reaction is exothermic → more products. If it’s endothermic → fewer products. |
3 · Pressure (gaseous equilibria only). Before the rule, it is worth remembering what pressure actually is.
Pressure = force ÷ area. For a gas, the force is its molecules constantly striking the walls of the container, and the area is the inside surface of that container — set by its volume. Pack more gas molecules into the space (or squeeze the same molecules into a smaller volume) and the walls are struck more often, so the pressure rises.
So the number of gas molecules is what sets the force on the walls — and that is the one lever the equilibrium can pull. It cannot change the size of the container you have chosen, but it can shift to the side with fewer gas molecules: fewer molecules → fewer collisions with the walls → lower pressure.
So the system counteracts an increase in pressure by shifting toward the side with fewer gas molecules:
- Increase the pressure → equilibrium shifts toward the side with the smaller number of gas molecules, as shown by the symbol equation.
- Decrease the pressure → equilibrium shifts toward the side with the larger number of gas molecules.
- Count the molecules from the balanced equation’s coefficients — e.g. in N2 + 3H2 ⇌ 2NH3 there are 4 molecules on the left and 2 on the right.
Le Chatelier questions are marked on the complete chain of reasoning, not the conclusion alone. Build every answer in three links:
- State what the change is in the system’s terms (“temperature increased”, “4 gas molecules on the left, 2 on the right”).
- Apply the principle: the equilibrium shifts in the direction that counteracts it — naming that direction and why (“…shifts in the endothermic direction, which here is the reverse reaction, to absorb the extra heat”).
- Finish with the effect the question actually asked about: “…so the amount/yield of NH₃ at equilibrium decreases”. Stopping one step short — naming the shift but never the consequence — is one of the most common ways to lose the final mark.
Now try it. The simulator below builds the ammonia equilibrium from scratch — press Start to watch the forward and reverse reactions reach a dynamic balance, then change one condition at a time and read the examiner-style reasoning for every shift.
⚖️ The ammonia equilibrium, live
Remember: the forward reaction is exothermic. Press Start to reach equilibrium, then change one condition at a time and read what happens.
Only reactants so far — 6 N₂ and 18 H₂ bouncing around. Press Start to let them react and watch the mixture settle into equilibrium.
- Mixing up rate and position. A catalyst speeds up both directions equally: equilibrium is reached sooner, but the position (and yield) is unchanged. Don’t claim a catalyst increases yield.
- Counting molecules wrong for pressure. Count gas molecules only, using the balanced equation’s coefficients. If both sides have the same number of gas molecules, changing the pressure does not shift the position at all.
- Forgetting which direction is exothermic. The question always tells you (or gives data to work it out) — quote it in your answer: “the forward reaction is exothermic, so the endothermic direction is the reverse…”
Reading an equilibrium-yield graph
Higher Tier questions often give a graph of the % yield of a product against pressure, with a separate curve for each temperature. Read it with the same Le Chatelier logic, in two directions:
- Along a curve (pressure rising) — for N2 + 3H2 ⇌ 2NH3 the yield rises with pressure, because higher pressure favours the side with fewer gas molecules (2 on the right, not 4 on the left).
- Between curves (temperature) — the lower the temperature, the higher the yield, because the forward reaction is exothermic and cooling favours it.
- So the best yield sits in the top-right corner: high pressure, low temperature. Why industry doesn’t simply operate there is the very next point.
Yield of ammonia against pressure at three temperatures.
This is where the two halves of C6 finally meet. Take a forward reaction that is exothermic and makes fewer gas molecules, like N2 + 3H2 ⇌ 2NH3. The highest equilibrium yield wants a low temperature (favouring the exothermic forward direction) and a high pressure (favouring the 2-molecule side). But a low temperature also means a slow rate — you could wait days for that yield.
So real processes settle on a compromise: a moderate temperature (some yield given up to keep the rate usable), a catalyst (which speeds the reaction up without moving the position, so it lifts the rate at no cost to the yield), and a high pressure where the equation — and the expense of the equipment — allow it. And because each pass through the reactor only converts some of the mixture, the unreacted reactants are separated off and recycled back through — so very little is wasted, and the overall yield of the process is high even when a single pass is modest. If a question asks why a chosen temperature isn’t the one that gives the best yield, the answer is always this rate–yield trade-off.
The industrial optimisation and compromise, and reading an equilibrium-yield graph is a Higher-Tier data skill for everyone. The full industrial detail — recycling the unreacted gases, and choosing the conditions for a named process such as the Haber process — belongs to the Using Resources topic (4.10), and is Triple Chemistry only. It’s shown here, kept generic, just to bring the rate–yield idea to life — Combined Science students don’t need the named-process specifics.
🧪 Exam-style questions Higher
In the equilibrium ICl(l) + Cl2(g) ⇌ ICl3(s), the forward reaction is exothermic. What happens to the amount of ICl3 if the temperature is increased? Tick (✓) one box.
For the equilibrium 2SO2(g) + O2(g) ⇌ 2SO3(g), what is the effect of increasing the pressure? Tick (✓) one box.
Ethanol is made industrially by the equilibrium C2H4(g) + H2O(g) ⇌ C2H5OH(g). The ethanol is condensed and removed as it forms. Why does this increase the yield? Tick (✓) one box.
Hydrogen for industry is made using the equilibrium CH4(g) + H2O(g) ⇌ CO(g) + 3H2(g). The forward reaction is endothermic. Predict and explain the effect on the equilibrium yield of hydrogen of (a) increasing the temperature, and (b) increasing the pressure.
Show a model answer
(a) Increasing the temperature:
- The equilibrium shifts in the endothermic direction to absorb the extra heat — here that is the forward direction. 1 mark
- So the equilibrium shifts right and the yield of hydrogen increases. 1 mark
(b) Increasing the pressure:
- There are 2 gas molecules on the left and 4 on the right; higher pressure favours the side with fewer gas molecules — the left. 1 mark
- So the equilibrium shifts left and the yield of hydrogen decreases. 1 mark
Examiner note — both parts use the same three links: the change, the counteracting direction with its reason, and the effect on yield. Notice the question gave the thermal information (“the forward reaction is endothermic”) — quote it, don’t guess it.
Ammonia is made industrially from the equilibrium N2(g) + 3H2(g) ⇌ 2NH3(g). The forward reaction is exothermic. Explain how the temperature, the pressure and a catalyst should be chosen so that ammonia is made economically.
This is a levels-of-response question, marked on how well your points are linked into a clear account.
Show a model answer
Levels of response:
- Level 3 (5–6): Relevant points are identified, given in detail and logically linked to form a clear account.
- Level 2 (3–4): Relevant points are identified, and there are attempts at logical linking; the account is not fully clear.
- Level 1 (1–2): Points are identified and stated simply, but their relevance is not clear and there is no attempt at logical linking.
Indicative content — a full answer weighs yield against rate and cost for each condition, not just “use a high pressure”.
Temperature
- A low temperature gives a higher yield, because it favours the exothermic forward reaction… 1 mark
- …but a low temperature also makes the rate slow, so a moderate (compromise) temperature is used — some yield is given up to keep the rate usable. 1 mark
Pressure
- A high pressure gives a higher yield, because it favours the side with fewer gas molecules (2 on the right, not 4), and it also speeds the reaction up… 1 mark
- …but very high pressures need expensive, strong equipment and are hazardous, so a high but not extreme pressure is the compromise. 1 mark
Catalyst and recycling
- A catalyst speeds the reaction up without moving the position of equilibrium — it raises the rate at no cost to the yield (and lets a lower temperature be used). 1 mark
- The unreacted nitrogen and hydrogen are recycled, so little is wasted and the overall process is economical. 1 mark
Conclusion (needed for Level 3) — a top-level answer pairs each condition with both its effect on the yield and the rate-or-cost reason it has to be a compromise, then draws the threads together: the chosen conditions are the ones that make ammonia fast enough and cheaply enough, even though none of them gives the highest possible yield on its own. Listing the conditions without the trade-off caps the marks.