Reaction Kinetics

1. Rate of Reaction

Definition: Rate of reaction is the concentration change per unit time of a reactant or product. (Basically it just tells us the speed of the reaction — but write the formal wording in the exam.)

• Δ means change. Use Δ[reactant] or Δ[product] — both give the rate.
• Measure whichever is easier: a solid reactant → its mass falls; a gas product → collect it and watch its pressure / volume rise.

① Getting the rate from a concentration–time graph

Gradient of a concentration–time graph = rate of reaction.

1. Go to the time asked (e.g. t₁) on the curve.
2. Draw a neat tangent to the curve at that point.
3. Work out the tangent’s gradient with (y₂−y₁)/(x₂−x₁) — that is the rate at t₁.

② Trend of the rate over time

Draw tangents at every point and see what the gradient does — exactly like reading gradients in kinematics:

• Very steep → less steep → even less → completely horizontal (m = 0). Horizontal line has gradient 0.
• So the rate is highest at the start (t = 0) and falls to zero at the end — same story whether you plot reactant or product.

③ Is the reaction reversible? (only from a reactant graph)

Reactant [conc]–time graph…	Meaning
falls all the way to 0	reactants fully converted to products ⇒ not a reversible (complete) reaction
falls, then goes constant above 0	reactants not fully converted ⇒ reversible; mixture holds reactants and products ⇒ equilibrium established (at t4 in the sketch)
a product graph	can’t tell! Product always rises then levels off, reversible or not

2. Collision Theory & Activation Energy

Collision theory: two species (atoms / molecules / particles) only react if they ① collide with ② a high enough energy — greater than or equal to the activation energy. Colliding weakly, or not colliding at all, ⇒ no reaction.

Activation energy (E_a): the minimum energy required to start the reaction.

Reactants must collide with this or a greater energy to break bonds and then form new ones. Remember: bond breaking is endothermic.

3. Boltzmann Distribution

Boltzmann plotted number of particles (y-axis) against energy (x-axis) for the particles in a sample (e.g. a gas). It looks like a normal bell curve, but it is not symmetrical: it starts at the origin and has a long tail to the right.

• Starts at the origin (0,0) — energy can’t be negative, so the curve stops at zero. Almost no particles have zero energy.
• Most particles have a moderate energy (the peak ≈ the average). Few have very low or very high energy.
• Never touches the x-axis on the right — a rare particle may have monstrously high energy.

Where is E_a on this graph?

Normally things don’t react at room temperature — you have to heat them / give them energy. So E_a sits on the right-hand side, in the region where few particles have that much energy.

Area beyond the E_a line (red) = particles with E ≥ E_a — these can react if they collide.

Area before it (green) = particles that cannot react — not enough energy.

4. Factors Affecting Rate of Reaction

① Concentration

Increasing the concentration of reactants usually increases the rate (not always!) — the “not always” is A2 stuff. For AS just assume it does. Doubling the concentration does not automatically double the rate (could be ×4, or no change).

② Surface area

Increasing the surface area (e.g. powdering a solid reactant) increases the rate. The other reactant can attack from more sides — more of the solid is exposed.

Try to put the words “collision frequency” into your answers.

③ Temperature

Increasing temperature increases the rate — in two ways:

• (1) collision frequency ↑ as particles travel faster (more collisions).
• (2) increases the number of particles with energy ≥ E_a (so each collision is more likely to be successful). ← the bigger effect

Think: fast particles ⇒ crowded market ⇒ more collisions.

	Peak	Area beyond Ea	Ea itself
T ↑	shifts right & lower	increases	unchanged
T ↓	shifts left & higher	decreases	unchanged

5. Catalysts

Using a catalyst increases the rate of reaction.

• It only lowers E_a — nothing else. Lowered “automatically” because the alternative route is easier.
• It must not be used up: if it is consumed it’s a reactant. If it’s regenerated (same amount at the end) → catalyst.

Example: A + B → C needs 200 J. With catalyst X: first A + X → AX (needs only 50 J), then AX + B → C + X. Same overall reaction, but the energy needed to complete it has dropped.

Homogeneous vs heterogeneous

Homogeneous catalyst: in the same phase (same state) as the reactants.

Heterogeneous catalyst: in a different phase (different state) to the reactants. (homo = same, hetero = different)

Reaction	Catalyst	Catalyst vs reactants	Type
1) Decomposition of H2O2 2H2O2 → 2H2O + O2	MnO2	solid vs liquid	Heterogeneous
2) Haber process N2 + 3H2 ⇌ 2NH3	Fe (iron)	solid vs gases	Heterogeneous
3) Contact process 2SO2 + O2 ⇌ 2SO3	V2O5 (vanadium(V) oxide / pentoxide)	solid vs gases	Heterogeneous
4) Hydrogenation of C=C bond	Ni (nickel) — or Pt	solid vs gas / liquid alkene	Heterogeneous
5) Esterification alcohol + carboxylic acid ⇌ ester + water	H2SO4	aqueous vs aqueous / liquid	Homogeneous

Catalyst on the Boltzmann graph & energy profile

• The E_a line moves left / back → more area beyond it → more successful collisions.
• Be reasonable: shift it a little, keep it near the right end — don’t drag it right to the peak.
• Temperature changes the curve’s shape; a catalyst does NOT — it only moves the E_a line.

Change	Boltzmann curve shape	Ea line	Area beyond Ea
Temperature ↑	changes (peak right & lower)	same place	↑
Catalyst	unchanged	moves LEFT	↑

6. Past-Paper Questions (all MCQs)

Six questions covered at the end of the class — this topic comes almost only in MCQs. Tip: if unsure, first cross out options you know are wrong.

7. Quick Recap — Cheat-Sheet