Chemical Kinetics — IB Chemistry SL Study Guide

For: IB Chemistry SL candidates sitting IB Chemistry SL.

Covers: IB Topic 6 — definition and measurement of rate of reaction, collision theory, factors affecting rate (concentration, surface area, temperature, catalyst), Maxwell-Boltzmann distribution.

A note on the practice questions: All worked questions in the "Practice Questions" section below are original problems written by us in the IB Chemistry SL style for educational use. They are not reproductions of past IBO papers.

1. Why Kinetics Matters

Topic 6 answers "how fast does this reaction go?" Topic 5 (Energetics) answers "how much energy?" and Topic 7 (Equilibrium) answers "how far?" Together they describe a chemical reaction completely.

About 7-9% of IB Chemistry SL covers Topic 6 directly, but kinetic ideas underpin Topic 7 (rate vs equilibrium) and Topic 16 HL (rate laws and mechanisms).

2. Rate of reaction

Rate = change in concentration per unit time: $$\text{Rate}=-\frac{\Delta [\text{reactant}]}{\Delta t}=+\frac{\Delta [\text{product}]}{\Delta t}$$

Units: mol·dm⁻³·s⁻¹.

For a reaction $aA+bB\to cC+dD$, the relationship is: $$\text{Rate}=-\frac{1}{a}\frac{\Delta [A]}{\Delta t}=-\frac{1}{b}\frac{\Delta [B]}{\Delta t}=+\frac{1}{c}\frac{\Delta [C]}{\Delta t}$$

Common ways to measure rate experimentally:

• Gas evolution: collect gas, measure volume vs time.
• Mass loss: mass on a balance vs time.
• Colorimetry: absorbance vs time (for coloured species).
• pH change: pH meter for acid-base reactions.
• Conductivity: ionic species change conductivity.

3. Collision theory

For a reaction to occur:

1. Particles must collide.
2. Collision must have enough energy ($E\ge E_a$, the activation energy).
3. Particles must collide in the correct orientation.

Increasing the frequency of collisions OR the fraction of effective collisions (those meeting both energy and orientation requirements) increases rate.

4. Factors affecting rate

Concentration / pressure: more particles per unit volume → more collisions per unit time → faster rate. Doubling [A] roughly doubles collisions involving A.

Surface area (heterogeneous reactions): more exposed reactant means more collision sites. Powdered marble + acid is much faster than a marble chip in the same acid.

Temperature: hotter particles move faster (more collisions), AND a higher fraction have $E\ge E_a$ (Maxwell-Boltzmann shift). Effect is exponential — a 10 °C rise typically doubles rate.

Catalyst: provides an alternative reaction pathway with lower activation energy. Increases the fraction of collisions with $E\ge E_a$ without being consumed.

5. Maxwell-Boltzmann distribution

Plots fraction of particles vs energy at a given temperature. Key features:

• Peak at the most probable energy.
• Long tail to the right (high-energy particles).
• Total area under curve = 1 (or total particle count).
• Particles with $E\ge E_a$ are those right of the activation-energy line.

When temperature increases:

• Peak shifts right and lower (broader spread).
• Tail extends further right.
• Fraction with $E\ge E_a$ increases substantially — this is the dominant kinetic effect of $T$.

When a catalyst is added:

• Distribution itself unchanged.
• Activation-energy line moves left (lower $E_a$).
• Fraction with $E\ge E_a$ increases.

6. Catalysts

A catalyst:

• Lowers $E_a$ by providing an alternative reaction pathway.
• Is not consumed in the overall reaction (may be regenerated step by step).
• Speeds up forward and reverse reactions equally — does not shift equilibrium position.
• Often involves a transition metal (Pt, Pd, Ni, Fe, V, etc.) or specific organic species (enzymes).

Homogeneous catalyst: same phase as reactants. Heterogeneous: different phase (e.g. solid catalyst with gas or liquid reactants — surface adsorption).

7. Worked Example

A solution of 0.50 M sodium thiosulfate is mixed with HCl. The reaction ${\text{S}}_2{\text{O}}_3^{2-}+2{\text{H}}^{+}\to \text{S}(s)+{\text{SO}}_2+{\text{H}}_2\text{O}$ produces a sulfur precipitate that obscures a cross drawn under the beaker. Time taken for the cross to disappear is 60 s.

(a) The experiment is repeated with the same volume but 1.0 M thiosulfate. Predict the time taken. (b) The original experiment is repeated at 35 °C instead of 25 °C. Predict the time taken qualitatively. (c) A drop of CuSO₄ is added (catalyst). Predict the effect on time and on equilibrium amount of sulfur produced.

Solution.

(a) Doubling [S₂O₃²⁻] roughly doubles rate → time roughly halves to ~30 s.

(b) A 10 °C rise typically doubles or roughly doubles rate. Time roughly halves to ~30 s, or even less. (Exact answer depends on activation energy.)

(c) Catalyst lowers $E_a$ → faster reaction → shorter time. Equilibrium amount of S unchanged (catalyst does not affect position of equilibrium, only how fast equilibrium is reached).

8. Common Pitfalls

• Stoichiometric coefficient in rate: rate of reaction is the same regardless of which species you measure — but use $\frac{1}{\text{coeff}}\frac{\Delta [X]}{\Delta t}$ to relate to species rates.
• Concentration vs amount: increasing volume of a fixed amount of reactant decreases concentration → slows rate. Common trap.
• Catalyst and equilibrium: catalysts speed both directions equally → no shift in equilibrium position. Don't say "catalyst increases yield".
• Maxwell-Boltzmann: temperature flattens AND broadens the curve. Total area stays constant, but spread changes.

9. Practice Questions

1. In a reaction $A+B\to C$, the initial rate doubles when [A] doubles, and increases 4× when [B] doubles. Determine the rate law.
2. The half-life of a first-order reaction is 30 s. How long until 87.5% of reactant has been consumed?
3. Sketch a Maxwell-Boltzmann distribution at temperatures $T_1<T_2$, with the activation-energy line marked. Indicate which fraction of particles can react in each case.

10. Quick Reference Cheatsheet

• Rate = $-\Delta [\text{reactant}]/\Delta t=+\Delta [\text{product}]/\Delta t$.
• Collision theory: collisions × correct orientation × enough energy.
• Increasing rate: ↑ concentration, ↑ surface area, ↑ temperature, add catalyst.
• $T$ ↑ 10 °C ≈ rate ×2 (rule of thumb).
• Catalyst: lowers $E_a$, not consumed, no equilibrium shift.
• Maxwell-Boltzmann: at higher $T$, peak right + lower; tail extends; fraction $\ge E_a$ ↑.

11. What's Next

Kinetics feeds directly into Topic 7 (Equilibrium) — a system at equilibrium has equal forward and reverse rates. HL Topic 16 extends kinetics with rate laws ($r=k[A{]}^m[B{]}^n$) and reaction mechanisms. Use Ollie for any specific mechanism or rate-law derivation.