Reaction Rates & Equilibrium
CollegeHow fast, and how far
Kinetics asks how fast a reaction goes; equilibrium asks how far. Collision theory states that a reaction occurs only when particles collide with enough energy (the activation energy, Ea) and the correct orientation. Anything that increases the frequency or energy of effective collisions increases the rate:
- Concentration / pressure ↑ — more particles per volume, more collisions.
- Temperature ↑ — faster particles, more collisions exceeding Ea.
- Surface area ↑ — more exposed particles to collide with.
- Catalyst — provides an alternative pathway with lower Ea, unchanged itself.
A reversible reaction reaches dynamic equilibrium when forward and reverse rates are equal — concentrations stop changing though both reactions continue. Its position is described by the equilibrium constant:
Le Chatelier's principle: if a system at equilibrium is disturbed, it shifts to oppose the change.
A tug-of-war that never stops
Dynamic equilibrium is easy to misread as "the reaction stopped". It has not. Molecules keep converting in both directions, but at equal rates, so the net concentrations hold steady — like a bridge where people cross both ways at the same rate, leaving the count on each side constant while nobody stands still.
Le Chatelier's principle lets you predict the response to a disturbance. Add more reactant, and the system consumes it by shifting forward. Remove product, and it shifts forward to replace it. Raise the temperature, and it favours the endothermic direction (which absorbs the added heat). Increase the pressure on a gaseous equilibrium, and it shifts toward the side with fewer gas molecules. This is the scientific basis of industrial optimisation — the Haber process for ammonia, for instance, uses high pressure and a carefully chosen temperature, plus an iron catalyst, precisely to push and speed the equilibrium favourably.
A catalyst speeds up both the forward and reverse reactions equally. It helps equilibrium arrive sooner but does not change the equilibrium position or the value of Kc.
Worked example — applying Le Chatelier to the Haber process
The reaction is N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = −92 kJ mol⁻¹ (exothermic). Predict how to maximise the yield of ammonia.
- Count gas moles each side: left = 1 + 3 = 4 mol gas; right = 2 mol gas.
- Apply pressure: increasing pressure shifts toward fewer gas molecules (the right), so high pressure favours NH₃.
- Apply temperature: the forward reaction is exothermic, so lowering temperature shifts right and favours NH₃.
- Weigh the kinetic cost: low temperature also makes the reaction very slow, so in practice a compromise temperature (~450 °C) balances yield against speed.
- Add a catalyst: iron speeds the approach to equilibrium without shifting it — the yield is reached faster.
- Conclusion: high pressure (~200 atm), moderate temperature, and an iron catalyst — exactly the real industrial conditions.
How does a catalyst speed up a reaction without being consumed?
It offers an alternative reaction pathway with a lower activation energy, so a larger fraction of collisions have enough energy to react. The catalyst participates in intermediate steps but is regenerated by the end, emerging chemically unchanged.
For the equilibrium in a sealed flask, what does a large Kc (≫ 1) tell you?
Kc is the ratio of product to reactant terms at equilibrium. A large Kc means the numerator (products) dominates — at equilibrium the mixture is mostly products, so the reaction lies far to the right ("goes nearly to completion").
For N₂ + 3H₂ ⇌ 2NH₃, which way does the equilibrium shift if some NH₃ is removed?
Removing product disturbs the balance; by Le Chatelier the system shifts to replace it, i.e. forward (to the right), making more ammonia. Continuously removing NH₃ is used industrially to keep driving the reaction.
Why does powdered zinc react with acid faster than a single lump of the same mass?
Powdering exposes far more surface area. Reaction happens only at the solid's surface, so more exposed atoms means more frequent collisions with acid particles per second, and therefore a faster rate — even though the total mass is the same.
How the ideas connect
Every key idea in this chapter, branching from the core concept — use it to see the whole picture at a glance.
The key facts, visualised
Worked problems, step by step
Follow each solution line by line, then try to reproduce it on paper before moving on.
Example 1In N2 + 3H2 <-> 2NH3 (exothermic), how does raising pressure shift equilibrium?
- Higher pressure favors the side with fewer gas moles.
- Left has 4 moles of gas, right has 2 moles.
- The system shifts right to reduce pressure, making more NH3.
Example 2For that same exothermic reaction, how does lowering temperature affect NH3 yield?
- Lowering temperature removes heat, a product of an exothermic reaction.
- The system shifts to replace the heat, favoring the exothermic (forward) direction.
- So the equilibrium moves right.
Now you try
Work each one out first, then tap to reveal the worked answer.