This chemistry grades 9-10 lesson on chemical equilibrium notes gives the key ideas first, then worked examples and diagrams. After that, practise with the worksheet on chemical equilibrium and check yourself with the quiz.
Quick answer: A reversible reaction can go in both directions, and in a closed system it reaches dynamic equilibrium when the forward and reverse rates are equal and concentrations stay constant. Le Chatelier’s principle says the equilibrium shifts to oppose any change. Also, the Haber process makes ammonia from nitrogen and hydrogen using compromise conditions.
Level: Chemistry grades 9-10 (GCSE · IGCSE · O Level · Class 9-10) · Curriculum: HS-PS1-6
Chemical equilibrium notes: the key ideas
What is a reversible reaction?
Many reactions go only one way: once a candle has burned, you cannot turn the gases back into wax. However, some reactions can run in both directions. Heat blue hydrated copper(II) sulfate and it turns white, giving off water. Next, add water to the white powder and it turns blue again. Chemists show this with a double arrow: CuSO₄·5H₂O ⇌ CuSO₄ + 5H₂O.
A reversible reaction is one in which the products can react to form the reactants again. In a closed system, a reversible reaction reaches dynamic equilibrium: the forward and reverse reactions both continue at the same rate, so the concentrations stop changing.
What does dynamic equilibrium mean?
At first only the forward reaction happens, and it is fast. Moreover, as products build up, the reverse reaction begins and speeds up while the forward rate slows. Eventually the two rates are equal. Nothing seems to change, but particles are still reacting in both directions, which is why the equilibrium is dynamic.
Equilibrium does not mean equal amounts of reactants and products, and it does not mean the reaction has stopped. The amounts stay constant because the forward and reverse rates are equal. The system must also be closed, or gases can escape and the system never reaches equilibrium.
What is Le Chatelier’s principle?
Le Chatelier’s principle: if you change a system at equilibrium, the position of equilibrium shifts to oppose the change.
| Change made | Equilibrium shifts | Reason |
|---|---|---|
| Add a reactant | Towards the products | Uses up the extra reactant |
| Remove a product | Towards the products | Makes more product to replace it |
| Increase pressure | To the side with fewer gas molecules | Reduces the pressure |
| Increase temperature | In the endothermic direction | Absorbs the extra energy |
| Add a catalyst | No shift | Speeds up both directions equally |
Chemical equilibrium: a worked example with NO₂
Predict the effect on the brown gas NO₂ in the reaction N₂O₄(g) ⇌ 2NO₂(g), where the forward reaction is endothermic (colourless N₂O₄, brown NO₂).
Heating favours the endothermic direction, so the equilibrium shifts right and the mixture turns darker brown. Raising the pressure favours the side with fewer gas molecules: 1 molecule on the left against 2 on the right, so it shifts left and the mixture turns paler.
How is ammonia made in the Haber process?
Industry makes ammonia, used for fertilisers that feed crops from the Punjab to Kenya, from nitrogen (from the air) and hydrogen (from natural gas): N₂ + 3H₂ ⇌ 2NH₃. In addition, the forward reaction is exothermic (ΔH = −92 kJ/mol) and reduces the number of gas molecules from 4 to 2.
Chemical equilibrium and industrial conditions
Le Chatelier suggests a low temperature and a high pressure would give the most ammonia. Also, in industry the conditions are a compromise. Cold conditions make the reaction far too slow, so a temperature of about 450 °C and an iron catalyst give a reasonable rate. A very high pressure needs thick pipes and costly compressors, so industry uses about 200 atm. Also, only a fraction of the gases turns into ammonia on each pass; we cool the ammonia until it turns liquid and recycle the unreacted nitrogen and hydrogen.
Chemical equilibrium: a worked example on ammonia
A reactor holds 100 mol of N₂ and 300 mol of H₂. Moreover, at equilibrium 25% of the nitrogen has reacted. Then, find the amounts present and the percentage of ammonia in the mixture.
N₂ reacted = 25 mol, so N₂ left = 75 mol. H₂ reacted = 3 × 25 = 75 mol, so H₂ left = 225 mol. NH₃ formed = 2 × 25 = 50 mol. Total = 75 + 225 + 50 = 350 mol, so NH₃ = 50 ÷ 350 × 100 = 14.3%.
Check three things in every prediction: count the gas molecules on each side for pressure, look at the sign of ΔH for temperature, and remember that a catalyst never moves the position, it only gets there sooner.
- Reversible reactions use ⇌ and reach equilibrium only in a closed system
- At equilibrium the forward and reverse rates are equal; concentrations are constant, not equal
- Le Chatelier: the equilibrium shifts to oppose the change made
- Higher pressure favours fewer gas molecules; heating favours the endothermic direction
- Haber process: N₂ + 3H₂ ⇌ 2NH₃, 450 °C, 200 atm, iron catalyst, recycle
Reactions that go both ways: frequently asked questions
What does dynamic equilibrium mean?
Both the forward and the reverse reactions keep happening at the same rate, so the amounts of reactants and products stay constant. In addition, the reaction has not stopped; the two rates balance.
Does a catalyst change the position of equilibrium?
No. A catalyst speeds up the forward and the reverse reaction equally, so the system reaches equilibrium sooner but the amounts at equilibrium are the same.
Why does the Haber process run at 450 °C instead of a low temperature?
A low temperature gives a better yield of ammonia but the reaction is too slow. Also, about 450 °C with an iron catalyst is a compromise between a reasonable yield and a fast rate.
How does pressure affect an equilibrium?
Raising the pressure shifts the equilibrium to the side with fewer gas molecules. In the Haber process there are 4 molecules on the left and 2 on the right, so higher pressure gives more ammonia.
Why does the Haber process recycle the unreacted gas?
Only a fraction of the nitrogen and hydrogen reacts on each pass through the reactor. Recycling the rest saves raw materials and increases the overall yield.
