Some reactions run both ways
Many reactions are reversible: the products can react to reform the reactants. We write them with a double arrow, ⇌. In an open container the products often escape and the reaction just goes forward, but in a closed system — nothing in or out — something more interesting happens.
The forward reaction starts fast and slows as reactants are used up; the backward reaction starts at zero and speeds up as products build. Eventually the two rates become equal — dynamic equilibrium. The concentrations stop changing, but both reactions are still going, just at the same rate in each direction. "Dynamic" is the key word: it is balance, not stillness.
Le Chatelier's principle
The whole topic is one rule: if you change the conditions of a system at equilibrium, it shifts to oppose the change. That is Le Chatelier's principle, and it predicts which way the balance moves.
Add more of a reactant and the system shifts forward to use it up. Remove a product and it shifts forward to replace it. The system always moves in the direction that partly cancels whatever you did.
Temperature and pressure
Temperature: raising it favours the endothermic direction (the one that absorbs the extra heat); lowering it favours the exothermic direction. So heating a reaction whose forward direction is exothermic pushes the equilibrium back, giving less product.
Pressure (for gases): raising it shifts the equilibrium toward the side with fewer gas molecules, easing the squeeze. If both sides have the same number of gas molecules, pressure has no effect. A catalyst changes neither position — it only makes equilibrium arrive faster.
A catalyst does not shift equilibrium
A catalyst speeds up the forward and backward reactions equally, so it reaches equilibrium sooner but does not change how much product you end up with. Only concentration, temperature and pressure move the position.
The Haber process
The classic industrial equilibrium is the Haber process for ammonia: N₂ + 3H₂ ⇌ 2NH₃, forward exothermic. The conditions are a compromise. High pressure (~200 atm) favours ammonia (4 gas molecules → 2). Low temperature would favour it too (exothermic), but too low and the reaction is far too slow — so a moderate ~450 °C is used with an iron catalyst, trading some yield for a workable rate. Unreacted gases are recycled. The ammonia goes mostly into fertilisers.
The Contact process
Sulfuric acid — the most-produced industrial chemical — is made by the Contact process, and its key step is another equilibrium governed by exactly the same rules:
Sulfur is first burned to sulfur dioxide (S + O₂ → SO₂). The crucial reversible step then oxidises it further: 2SO₂ + O₂ ⇌ 2SO₃, forward exothermic. By Le Chatelier, a low temperature and high pressure would favour SO₃ — but in practice a moderate ~450 °C, a vanadium(V) oxide (V₂O₅) catalyst, and only a modest ~1–2 atm already give about a 96% conversion, so extremes are not worth the cost. Finally the SO₃ is absorbed into concentrated sulfuric acid to form oleum, then diluted to give more H₂SO₄.
| Haber (ammonia) | Contact (sulfuric acid) | |
|---|---|---|
| Key equilibrium | N₂ + 3H₂ ⇌ 2NH₃ | 2SO₂ + O₂ ⇌ 2SO₃ |
| Catalyst | iron | vanadium(V) oxide, V₂O₅ |
| Temperature | ~450 °C | ~450 °C |
| Pressure | ~200 atm | ~1–2 atm |
| Both forward reactions | exothermic | exothermic |
Same principle, different numbers
Both processes run an exothermic gas equilibrium, so both would prefer low temperature (yield) but use a moderate ~450 °C (rate) with a catalyst. The Contact process needs little extra pressure because its conversion is already very high — a neat reminder that industrial conditions balance yield, rate and cost.
A worked example
For N₂ + 3H₂ ⇌ 2NH₃ (forward exothermic), which way does the equilibrium shift if you (a) increase the pressure, (b) increase the temperature?
- (a) Left side has 4 gas molecules, right has 2. Higher pressure favours the side with fewer molecules.The system shifts to reduce the pressure you applied.
- So higher pressure shifts it forward → more ammonia.Fewer gas molecules on the product side.
- (b) The forward reaction is exothermic, so raising temperature favours the endothermic (backward) direction.The system absorbs the added heat by going the endothermic way.
- So higher temperature shifts it back → less ammonia (but faster).This is why the Haber process compromises on a moderate temperature.
(a) Higher pressure → more ammonia; (b) higher temperature → less ammonia (though the reaction is faster) — hence the compromise conditions.