Skip to main content
ExamExplained

Reaction rates and collision theory (SACE Stage 2 Chemistry)

Syllabus dot point

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

8 min read

Quick answer

Collision theory says a reaction occurs only when particles collide with enough energy (at least the activation energy, EaE_a) and the correct orientation. The reaction rate depends on the frequency of these successful collisions. Anything that increases that frequency, higher concentration, greater surface area, higher pressure (for gases), or higher temperature, increases the rate. Temperature also raises the proportion of particles exceeding EaE_a, which is why it has the largest effect.

Jump to a section
  1. What this dot point is asking
  2. Lead worked calculation
  3. The collision theory model
  4. Explaining each factor
  5. The Maxwell-Boltzmann distribution
  6. Measuring rate
  7. Why it matters for managing processes
  8. Exam-style questions

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 Zn+2HCl→ZnCl2+H2\text{Zn} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2, 48 mL48\ \text{mL} of hydrogen gas is collected in the first 30 s30\ \text{s}, and a total of 72 mL72\ \text{mL} after 90 s90\ \text{s}. Find the average rate over the first 30 s30\ \text{s} and over the interval 3030 to 90 s90\ \text{s}, then comment.

Rate over the first 30 s

rate=ΔVΔt=4830=1.6 mL s−1\text{rate} = \frac{\Delta V}{\Delta t} = \frac{48}{30} = 1.6\ \text{mL s}^{-1}

Rate over 30 to 90 s

rate=72−4890−30=2460=0.40 mL s−1\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

Successful collisions and activation energy

For a reaction to happen, colliding particles must have a combined energy at least equal to the activation energy (EaE_a), the minimum energy needed to break the existing bonds and start forming new ones, and they must collide with the correct orientation. Collisions that fail either requirement bounce apart unchanged. Reaction rate is proportional to the frequency of successful collisions per second.

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 EaE_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 EaE_a can react.

Why temperature has the biggest effect

Raising the temperature shifts the whole distribution to higher energy and flattens its peak. The area under the curve beyond EaE_a, the fraction of particles able to react, grows steeply. A modest temperature rise therefore increases the number of successful collisions far more than it increases the total number of collisions, which is why a 10 ∘C10\ ^\circ\text{C} rise can roughly double the rate.

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 changetime\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.

Saying temperature only increases collision frequency

A frequent SACE error is explaining temperature solely by "particles move faster and collide more often". That effect is real but small. The dominant reason temperature speeds reactions is that a much larger fraction of particles now exceed the activation energy, so a far greater proportion of collisions are successful. Always mention EaE_a and the Maxwell-Boltzmann distribution. A second slip is confusing surface area (a solid property) with concentration (a solution property); a single lump and a powder of the same mass have the same amount but very different rates.

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 questions

Questions in the style of SACE Board exam questions on this dot point, each with a worked answer. They are written by ExamExplained unless tagged "Past paper"; the year shows the paper a question is modelled on.

SACE 2021-style4 marks
Using collision theory, explain why increasing the temperature of a reaction by 10 ∘C10\ ^\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 (EaE_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 EaE_a (the area under the curve beyond EaE_a) increases sharply, far more than the modest rise in total collision frequency. (1 mark)

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

SACE 2019-style3 marks
In an experiment, 0.50 g0.50\ \text{g} of magnesium ribbon reacts completely with excess hydrochloric acid in 40 s40\ \text{s}, releasing hydrogen gas. (a) Calculate the average rate of consumption of magnesium in g s−1\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 =mass reactedtime=0.5040=1.25×10−2 g s−1= \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)

Practise this

Sources & how we know this

Based on the official SACE Board subject outline. Not yet individually reviewed by a person; always check SACE Board and your school's materials for current rules.

ExamExplained