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Lessons / Rates and equilibrium / GCSE and A-level

Why does a reaction push back when you change it?

The A-level picture: the equilibrium constant, and which changes alter it.

  1. A sealed tube of brown gas10 cm
  2. A pocket of gas1 µm
  3. A few hundred molecules50 nm
  4. Two molecules joining1 nm
2NO2(g) ⇌ N2O4(g) ΔH = −57 kJ mol−1
What the simulation shows, in words

The same box of NO₂ and N₂O₄ as the Core page, with sliders for temperature and volume and a button to add NO₂. After each change the counts drift and settle at new steady values. The graph shows both counts over time, so you can see the system respond and level off.

The equilibrium constant

For a reaction at equilibrium, one ratio of concentrations always comes out the same at a given temperature. For this reaction:

Kc = [N2O4] / [NO2]2

Square brackets mean the concentration at equilibrium, in mol dm⁻³. The power 2 comes from the 2 in front of NO₂ in the equation. The units follow from the expression: (mol dm⁻³) ÷ (mol dm⁻³)² = dm³ mol⁻¹.

A large Kc means the equilibrium lies well to the right, a small one well to the left.

Concentration and pressure: the position moves, K doesn’t

Add NO₂ and the bottom of the fraction gets bigger, so the fraction is now smaller than Kc. The system isn’t at equilibrium any more. NO₂ joins up, the top grows and the bottom shrinks, until the fraction equals Kc again. That is Le Chatelier’s principle, explained: the shift happens because the ratio has to get back to Kc.

Pressure works the same way. Halve the volume and every concentration doubles. The top of the fraction doubles but the bottom goes up four times, because it is squared. The fraction falls below Kc, so the equilibrium moves right until it is restored. This is why only reactions with different numbers of gas molecules on each side respond to pressure.

Temperature: K itself changes

Temperature is different, because it changes Kc. The forward reaction is exothermic (ΔH = −57 kJ mol⁻¹), so heating makes Kc smaller: less N₂O₄ at equilibrium, more NO₂, a browner gas. For an endothermic forward reaction, heating makes Kc larger.

A catalyst changes neither the position nor Kc. It lowers the activation energy of the forward and reverse reactions by the same amount, so both speed up equally.

Choosing conditions in industry

For making ammonia, N₂ + 3H₂ ⇌ 2NH₃, ΔH = −92 kJ mol⁻¹:

  • Low temperature would give a better yield (exothermic forward reaction) but a slow rate. A compromise of about 450 °C is used, with an iron catalyst to make up the speed.
  • High pressure helps both yield (4 gas molecules become 2) and rate, but costs money and needs strong equipment. About 200 atmospheres is the compromise usually taught.
  • Ammonia is removed as it forms and unreacted gases are recycled, so the overall conversion is high even though each pass is not.
Safety

Nitrogen dioxide is toxic and ammonia is corrosive and toxic. This is the theory only. Don’t try to make either.