Explain how concentration, temperature, surface area and catalysts change reaction rate, and use Le Chatelier's principle to predict how an equilibrium shifts.
Why this matters
Two separate questions decide whether a reaction is useful: how fast does it go, and how far does it go? Food spoils faster on the counter than in the fridge (rate). An ammonia plant runs at high pressure to push more nitrogen and hydrogen into product (equilibrium). Engineers, chemists and lab staff deal with both every day, whether they are speeding up a production step or keeping a reagent stable on the shelf.
Part 1: reaction rate (kinetics)
The rate of a reaction is how quickly reactants are used up or products form, for example in mol/L per second.
Collision theory
Particles react only when they collide, with enough energy (at least the activation energy, Eₐ) and in the right orientation. Anything that makes useful collisions more frequent or more energetic speeds the reaction up.
Four factors that change the rate
| factor | effect | why |
|---|---|---|
| concentration (or gas pressure) | higher -> faster | more particles in the same space, more collisions |
| temperature | higher -> faster | particles move faster and more collisions exceed Eₐ |
| surface area | more surface -> faster | more particles of a solid are exposed |
| catalyst | faster | provides a lower-energy pathway (lower Eₐ) and is not used up |
Examples:
- Sugar dissolves and reacts faster as powder than as a cube (surface area). Fine powders can even be explosion hazards in industry for this reason.
- Refrigeration slows the reactions that spoil food (temperature). A common rough rule is that many reactions near room temperature go about twice as fast for every 10 °C rise; it is a rule of thumb, not a law.
- Enzymes are biological catalysts: catalase speeds the breakdown of hydrogen peroxide, 2 H₂O₂ → 2 H₂O + O₂, enormously.
- A car's catalytic converter speeds the conversion of carbon monoxide and unburned fuel to CO₂ and water.
Energy diagram, in words
Picture a hill between reactants and products. The height of the hill is the activation energy. A catalyst digs a lower pass through the hill; it does not change where the reactants start or where the products end up, so it does not change the overall energy released.
Part 2: chemical equilibrium
Many reactions are reversible: products can turn back into reactants.
N2(g) + 3 H2(g) <=> 2 NH3(g)
At first the forward reaction dominates. As products build up, the reverse reaction speeds up, until both run at the same rate. From then on the concentrations stay constant. This is dynamic equilibrium: the reactions have not stopped, they are balanced.
The equilibrium constant
For aA + bB <=> cC + dD:
K = [C]^c [D]^d / ( [A]^a [B]^b )
For the ammonia reaction, K = [NH₃]² / ([N₂][H₂]³).
- K much greater than 1: products are favoured at equilibrium.
- K much less than 1: reactants are favoured.
K changes only with temperature. Concentrations can change, but at a given temperature the system always settles back to the same K.
Le Chatelier's principle
If you disturb a system at equilibrium, it shifts to partly undo the disturbance.
Using the Haber process, which is exothermic (it releases heat, ΔH ≈ −92 kJ per mole of N₂ reacting):
| change | shift | reason |
|---|---|---|
| add N₂ or H₂ | right, more NH₃ | uses up the added reactant |
| remove NH₃ as it forms | right | replaces the removed product |
| increase pressure (smaller volume) | right | 4 moles of gas -> 2: fewer gas particles lowers pressure |
| increase temperature | left, less NH₃ | the reverse reaction absorbs heat |
| add a catalyst | no shift | speeds both directions equally; equilibrium is reached sooner |
This creates a real engineering trade-off. Low temperature gives more ammonia at equilibrium but makes the reaction slow. Plants therefore run at a compromise temperature (roughly 400–450 °C), high pressure (around 150–300 atm) and an iron catalyst, and they remove ammonia continuously. Rate and equilibrium are designed together.
Everyday equilibrium
- A sealed bottle of fizzy drink holds CO₂ in equilibrium with dissolved CO₂. Opening it lowers the pressure, so the equilibrium shifts and gas comes out of solution.
- Haemoglobin picks up oxygen in the lungs (high O₂) and releases it in tissues (low O₂), a shift driven by concentration.
Common mistakes
- Thinking equilibrium means equal amounts. It means equal forward and reverse rates; amounts can be very unequal.
- Saying a catalyst increases yield. It only gets you to equilibrium faster.
- Changing K by adding reactant. Concentrations shift, but K changes only with temperature.
- Confusing fast with favourable. A reaction can have a huge K and still be extremely slow (diamond turning to graphite).
Quick recap
- Rate rises with concentration, temperature, surface area and a catalyst.
- Catalysts lower the activation energy and are not consumed.
- Equilibrium is dynamic: forward and reverse rates are equal; K describes the balance.
- Le Chatelier: the system shifts to counteract a change; temperature is the only thing that changes K.
Try it
Open the Do tab and take the quiz: you will predict rate changes and equilibrium shifts for real situations.