Rate of reaction (C5.2)
A perennial Gateway A 6-marker. Examiners want collision theory used as the explanation for every observation, not as a memorised paragraph at the start.
What is "rate"?
Rate of reaction = how fast the reactants are used up OR how fast the products are formed.
mean rate = amount of reactant used (or product made) ÷ time taken
Units: g/s, mol/s, cm³/s — match whatever you're measuring.
To find rate at a specific time on a graph: draw a tangent and calculate gradient.
Collision theory — the master explanation
For a reaction to happen, particles must:
- Collide with each other (so frequency of collisions matters).
- Have enough energy at the moment of collision (more than the activation energy, Eₐ).
Anything that increases the frequency of successful collisions speeds up the reaction.
The four factors
1. Concentration / pressure
- Higher concentration → more particles per unit volume → more frequent collisions → faster rate.
- For gases, increasing pressure is equivalent to increasing concentration.
2. Temperature
Two effects, both speed it up:
- Particles have more kinetic energy → move faster → more frequent collisions.
- More particles have energy ≥ Eₐ → higher proportion of collisions are successful.
A rough rule of thumb: a 10°C rise often doubles the rate.
3. Surface area (of solid reactant)
- Greater surface area → more particles exposed → more frequent collisions with the other reactant → faster rate.
- Powder reacts faster than chunks. Examiners love this with marble (CaCO₃) chips vs powder + acid.
4. Catalyst
A catalyst is a substance that increases the rate of reaction without being used up itself. It works by providing an alternative reaction pathway with a lower activation energy, so a higher proportion of collisions are successful.
Important examples:
- Iron in the Haber process (N₂ + H₂ → NH₃).
- Vanadium oxide in the Contact process.
- Enzymes are biological catalysts.
⚠ Common error: writing that catalysts "speed up the reactant particles". They don't — they lower the activation energy.
Reaction profiles
A reaction profile diagram plots energy on the y-axis vs progress of reaction on the x-axis.
- Reactants on the left at one energy level.
- Products on the right at another (lower for exothermic, higher for endothermic).
- A hump in the middle = activation energy.
- A catalyst lowers the hump but doesn't change reactant or product energies.
Measuring rate — three standard methods
| Method | Reaction example | What you measure |
|---|---|---|
| Gas syringe | CaCO₃ + HCl → CO₂ | Volume of gas vs time |
| Mass loss on a balance | CaCO₃ + HCl (with cotton-wool plug) | Mass loss vs time |
| Disappearing cross | Na₂S₂O₃ + HCl (cloudy precipitate) | Time for cross to disappear |
The gradient of a "volume of gas vs time" graph at any point IS the rate at that moment. Steepest gradient = fastest rate.
Required practical: marble chips and HCl
Set-up:
- Mass of marble chips fixed.
- Volume of HCl fixed.
- Vary one variable (e.g. surface area: chips vs powder; or concentration: 0.5, 1.0, 2.0 mol/dm³).
- Measure mass loss every 10 s OR volume of CO₂ in a gas syringe every 10 s.
- Plot a graph; calculate initial rate from the gradient at t = 0.
Control variables: temperature, total volume of acid, total mass of CaCO₃.
Common Gateway-paper mistakes
- Writing "more reactions happen" instead of "more frequent successful collisions".
- Forgetting that temperature affects both collision frequency and the proportion of successful collisions.
- Saying a catalyst is "used up" — they aren't.
- Confusing surface area with mass.
- Drawing a reaction profile where the catalyst lowers the products' energy (it doesn't — only the activation energy hump).
➜Try this— Quick check
A reaction between Mg + HCl produces 60 cm³ of gas in 30 s.
- Mean rate = 60 / 30 = 2 cm³/s.
- After repeating with HCl twice as concentrated, the same volume is produced in 15 s.
- New rate = 60 / 15 = 4 cm³/s — doubled.
AI-generated · claude-opus-4-7 · v3-ocr-combined-science