§-Syllabus dot point

SAChemistrySyllabus dot point

What factors control how fast a reaction proceeds, and why?

Use collision theory to explain the effect of concentration, surface area, temperature and pressure on reaction rate.

Collision theory, activation energy and the Maxwell-Boltzmann distribution explaining how concentration, surface area, temperature, pressure and catalysts change reaction rate, with worked SACE-style rate calculations.

Generated by Claude Opus 4.88 min answerUpdated 2026-06-02

Reviewed by: AI editorial process; not yet individually human-reviewed

Have a quick question? Jump to the Q&A page

Jump to a section

What this dot point is asking
Lead worked calculation
The collision theory model
Explaining each factor
The Maxwell-Boltzmann distribution
Measuring rate
Why it matters for managing processes

What this dot point is asking

SACE expects you to use collision theory to explain each factor, interpret the Maxwell-Boltzmann distribution, and calculate average reaction rates from experimental data.

Lead worked calculation

Average rate from gas volume data

In the reaction $\text{Zn} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2$ , $48\ \text{mL}$ of hydrogen gas is collected in the first $30\ \text{s}$ , and a total of $72\ \text{mL}$ after $90\ \text{s}$ . Find the average rate over the first $30\ \text{s}$ and over the interval $30$ to $90\ \text{s}$ , then comment.

Rate over the first 30 s

\text{rate} = \frac{\Delta V}{\Delta t} = \frac{48}{30} = 1.6\ \text{mL s}^{-1}

Rate over 30 to 90 s

\text{rate} = \frac{72 - 48}{90 - 30} = \frac{24}{60} = 0.40\ \text{mL s}^{-1}

Interpret

The rate drops fourfold as the reaction proceeds. As reactants are used up, their concentration falls, so collisions become less frequent and the rate decreases. This is why reaction rate is always highest at the start.

The collision theory model

Explaining each factor

Concentration: more reactant particles in a given volume means collisions happen more often per second, so the rate increases.
Surface area: breaking a solid into smaller pieces (or a powder) exposes more particles to collision, increasing the frequency of successful collisions and the rate.
Pressure (gases): raising the pressure squeezes the same particles into a smaller volume, increasing their concentration and so the collision frequency and rate.
Temperature: gives particles more kinetic energy, so they collide both more often and, crucially, with more of them exceeding $E_a$ .

The Maxwell-Boltzmann distribution

The Maxwell-Boltzmann distribution shows the spread of kinetic energies among particles at a given temperature. Only those to the right of $E_a$ can react.

Measuring rate

Rate is the change in a measurable quantity per unit time: volume of gas produced, mass lost, change in concentration, or time for a fixed change (such as a precipitate obscuring a cross). The average rate over an interval is $\dfrac{\text{change}}{\text{time}}$ ; the instantaneous rate at a point is the gradient of the tangent to a concentration-time graph. Because reactant concentration falls as the reaction proceeds, the rate is greatest at the start and decreases over time.

Why it matters for managing processes

Controlling rate is central to industrial chemistry: too slow and production is uneconomic, too fast and a reaction may become unsafe. Collision theory provides the levers, concentration, surface area, pressure, temperature and catalysts, that engineers adjust to manage real chemical processes such as the Haber process.

Exam-style practice questions

Practice questions written in the style of SACE Board exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.

SACE 20214 marksUsing collision theory, explain why increasing the temperature of a reaction by

10\ ^\circ\text{C}

can roughly double the rate, even though the number of collisions only increases by a few per cent. Refer to the Maxwell-Boltzmann distribution and activation energy.

Show worked answer →

Rate depends on the frequency of successful collisions, those with energy at least equal to the activation energy ( $E_a$ ) and the correct orientation. (1 mark)

Raising the temperature increases the average kinetic energy of the particles, so the Maxwell-Boltzmann distribution shifts and broadens to higher energies. (1 mark)

The fraction of particles with energy greater than or equal to $E_a$ (the area under the curve beyond $E_a$ ) increases sharply, far more than the modest rise in total collision frequency. (1 mark)

Because the proportion of collisions that exceed $E_a$ rises steeply, the number of successful collisions per second roughly doubles, even though total collisions rise only slightly. (1 mark)

SACE 20193 marksIn an experiment,

0.50\ \text{g}

of magnesium ribbon reacts completely with excess hydrochloric acid in

40\ \text{s}

, releasing hydrogen gas. (a) Calculate the average rate of consumption of magnesium in

\text{g s}^{-1}

. (b) State and explain one change to the magnesium that would increase the rate, using collision theory.

Show worked answer →

(a) Average rate $= \dfrac{\text{mass reacted}}{\text{time}} = \dfrac{0.50}{40} = 1.25 \times 10^{-2}\ \text{g s}^{-1}$ . (1 mark)

(b) Using powdered magnesium instead of ribbon increases the surface area. (1 mark) With more surface area exposed, there are more sites where acid particles can collide with magnesium per second, so the frequency of successful collisions and the rate both increase. (1 mark)