8.1 Rate of Reactions
🎯 Learning outcomes
After completing this lesson, you should be able to:
- explain and use the terms rate of reaction, frequency of collisions, effective collisions and non-effective collisions
- explain qualitatively, in terms of the frequency of effective collisions, the effect of concentration and pressure changes on the rate of a reaction
- use experimental data to calculate the rate of a reaction
1⏱️What is rate of reaction?
Chemical kinetics is the branch of chemistry that studies the rates of reactions — how fast reactants turn into products — and the factors that affect them.
The rate of reaction is the change in the concentration of a reactant or product per unit time.
Consider a simple reaction where reactant A is converted into product B:
As the reaction proceeds, [A] falls while [B] rises. Because rate is always quoted as a positive quantity, a minus sign is needed when the rate is written in terms of a reactant (whose concentration is decreasing):
Concentration is measured in mol dm⁻³, calculated as moles ÷ volume (dm³). You’ll need this to interpret rate data later in this lesson.
2💥Collision theory
For a reaction to happen, particles (atoms, ions or molecules) must first collide with one another. But collisions alone are not enough — most collisions do not lead to a reaction at all.
Frequency of collisions — the number of collisions between particles occurring per unit time.
Effective collision — a collision that results in a reaction (bonds break and new bonds form).
Non-effective (ineffective) collision — a collision where particles simply bounce apart unchanged, with no reaction taking place.
For a collision to be effective, two conditions must both be satisfied:
- Proper orientation — the particles must collide with the correct parts of each molecule facing one another, so the right bonds are positioned to break and reform.
- Sufficient energy — the collision must supply at least the activation energy needed to start breaking the existing bonds.
A very common mistake is to say that increasing collision frequency always means a faster reaction. What actually matters is the frequency of effective collisions — always mention orientation and energy when explaining rate changes at A level.
In the reaction NO₂(g) + CO(g) → NO(g) + CO₂(g), two molecules may collide many times per second. If they collide side-on (wrong orientation) or too gently (insufficient energy), nothing happens — they simply separate again unchanged.
3🔥Activation energy & reaction profiles
The activation energy, EA, is the minimum energy required by colliding particles for a collision to be effective.
Even in reactions that release energy overall, particles must first climb over this energy “barrier” before products can form. Reaction profile diagrams (energy vs. progress of reaction) show this barrier clearly.
A catalyst works by providing an alternative reaction pathway with a lower activation energy, so a greater fraction of collisions become effective at the same temperature — increasing the rate without being used up itself. You’ll meet this in more detail in a later lesson.
4🧪Effect of concentration & pressure on rate
Increasing concentration
As concentration increases, there are more reactant particles in the same volume, so the frequency of collisions between them increases. This increases the frequency of effective collisions too, so the rate of reaction increases.
Increasing pressure (gaseous reactions)
Increasing the pressure of a gaseous reaction compresses the same number of particles into a smaller volume. This is equivalent to increasing their concentration, so once again the frequency of collisions increases, increasing the rate of the reaction.
Concentration and pressure changes affect only the frequency of collisions — they do not change the activation energy, and they do not make individual particles collide “harder” or with more energy. Never write that particles “gain energy” when concentration increases; the energy distribution of the particles is unchanged, only how often they meet.
If the concentration of a reactant is doubled, the frequency of collisions (and so the frequency of effective collisions) roughly doubles too — for a reaction that is first order in that reactant, this means the initial rate also roughly doubles.
5📊Calculating rate from experimental data
To find a rate experimentally, we monitor how the concentration of a reactant or product — or a measurable quantity proportional to it, such as the volume of gas evolved, mass lost, or colour intensity — changes over time, and plot it on a graph.
The rate at any moment equals the gradient (slope) of the tangent to the concentration–time curve at that point. The initial rate is the gradient of the tangent drawn at t = 0, where the curve is steepest.
The tangent drawn at the very start of the curve above passes through (0 s, 0.80 mol dm⁻³) and (20 s, 0 mol dm⁻³):
Notice the curve becomes less steep as time goes on — this is because the reactant concentration is falling, so collisions (and effective collisions) become less frequent, and the rate decreases.
Using raw data directly
Sometimes you are given a table of raw readings instead of a smooth curve — for example, the volume of gas collected over time in a reaction that produces a gas:
| Time / s | 0 | 10 | 20 | 30 | 40 |
|---|---|---|---|---|---|
| Volume of gas / cm³ | 0 | 18 | 30 | 37 | 40 |
Average rate between 0 s and 10 s: (18 − 0) cm³ ÷ (10 − 0) s = 1.8 cm³ s⁻¹.
Average rate between 20 s and 30 s: (37 − 30) cm³ ÷ (30 − 20) s = 0.7 cm³ s⁻¹ — noticeably slower, since the reactant is being used up.
📌 Exam tips & quick summary
✅ Do remember
- rate = change in concentration ÷ time (mol dm⁻³ s⁻¹)
- an effective collision needs correct orientation AND sufficient energy (≥ EA)
- ↑ concentration or ↑ pressure → ↑ frequency of (effective) collisions → ↑ rate
- instantaneous rate = gradient of the tangent to a concentration–time graph
- a catalyst lowers EA, opening an alternative pathway
🚫 Do not confuse
- collision frequency ≠ collision effectiveness — always link rate changes back to effective collisions
- ↑ concentration/pressure changes how often particles meet — it does not change each particle’s energy or EA
- rate is always positive, even though a reactant’s concentration is decreasing (hence the minus sign in −Δ[A]/Δt)
🔁 Quick recap
- Rate: Δ[conc.] / time (mol dm⁻³ s⁻¹)
- Activation energy, EA: minimum energy needed for an effective collision
- Effective collision: correct orientation + energy ≥ EA
- On a graph: rate at any point = gradient of the tangent at that point