What controls how fast a chemical reaction takes place?
Reaction rate and how it is followed, collision theory, the effect of concentration, particle size, temperature and catalysts on rate, the activation energy, the activated complex and the potential energy diagram.
An SQA Higher Chemistry answer on controlling the rate of reaction, covering how rate is measured, collision theory, the effects of concentration, particle size, temperature and catalysts, the activation energy and the activated complex on a potential energy diagram, with worked rate calculations.
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What this key area is asking
The SQA wants you to explain reaction rate using collision theory, describe and explain the effects of concentration, particle size, temperature and catalysts, calculate average rate from experimental data, and interpret a potential energy diagram showing the activation energy and the activated complex. This is one of the most heavily examined key areas, and it rewards precise wording over vague statements.
Measuring rate
The average rate of a reaction is the change in a measurable quantity divided by the time taken:
You can follow a reaction by measuring the volume of gas produced, the loss of mass as a gas escapes, a change in concentration, or how long a precipitate takes to obscure a mark drawn under a flask. A rate curve (quantity against time) is steepest at the start, where reactant concentration is highest, and levels off as reactants run out.
Collision theory
Factors affecting rate
Worked example: average rate over an interval
Activation energy and the activated complex
On a potential energy diagram, the activation energy is the height of the barrier from the reactants up to the peak. The enthalpy change () is the difference between the potential energy of the products and that of the reactants: negative for an exothermic reaction, positive for endothermic. A catalyst provides an alternative pathway with a lower , so a greater proportion of collisions succeed, but it does not change .
Examples in context
Industrial nitric acid production at plants such as those once run by ICI relied on a platinum-rhodium gauze catalyst to oxidise ammonia. The catalyst lowers the activation energy so the reaction runs fast at a workable temperature, exactly the collision-theory logic above: a lower barrier means more successful collisions per second without raising the temperature, which would be more expensive. In a school lab, the same principle appears when manganese dioxide catalyses the decomposition of hydrogen peroxide: the gas-collection rate curve is far steeper with the catalyst present, yet the total volume of oxygen ( and amount of product) is unchanged.
Try this
Q1. Explain, using collision theory, why increasing the concentration of a reactant increases the rate. [2 marks]
- Cue. There are more reactant particles per unit volume, so collisions are more frequent and more successful collisions occur per second.
Q2. A reaction produces of gas in . Calculate the average rate. [1 mark]
- Cue. .
Q3. State the effect of a catalyst on the activation energy and on the enthalpy change. [2 marks]
- Cue. It lowers the activation energy by providing an alternative pathway; it does not change the enthalpy change.
Exam-style practice questions
Practice questions written in the style of SQA exam questions on this dot point, with worked answer explainers. The year tag is the paper they imitate, not the source.
SQA Higher 20193 marksA student followed the reaction between excess marble chips and dilute hydrochloric acid by measuring the volume of carbon dioxide gas collected. In the first 20 seconds, 48 cm cubed of gas was produced. Calculate the average rate of reaction, in cm cubed per second, over this time interval and state what happens to the rate as the reaction proceeds.Show worked answer →
This is a standard average rate calculation. Markers reward the correct formula, the correct substitution, a value with the right unit, and the qualitative comment.
Average rate is the change in the measured quantity divided by the change in time:
As the reaction proceeds the rate decreases. The concentration of hydrochloric acid falls as it is used up, so there are fewer acid particles per unit volume, fewer successful collisions per second, and the rate drops until the acid runs out and the curve levels off.
A common mark lost here is quoting the rate with no unit, or saying the rate is constant.
SQA Higher 20214 marksExplain, in terms of collision theory and the potential energy diagram, why a small increase in temperature produces a large increase in reaction rate, and state the effect of a catalyst on the activation energy.Show worked answer →
A 4 mark answer needs the kinetic-energy argument, the activation-energy threshold, the distribution idea, and the catalyst point.
Raising the temperature increases the average kinetic energy of the particles, so they collide more often and, more importantly, a much greater proportion of particles now have kinetic energy equal to or greater than the activation energy. Because the energy distribution rises steeply near the threshold, even a small temperature rise of about 10 degrees C can roughly double the number of particles above the barrier, so the number of successful collisions per second rises sharply.
On the potential energy diagram the activation energy is the height of the barrier from the reactants up to the activated complex at the peak. A catalyst lowers the activation energy by providing an alternative reaction pathway, so a greater proportion of collisions are successful at the same temperature and the rate increases.
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Sources & how we know this
- SQA Higher Chemistry Course Specification — SQA (2018)