Rates and Collision Theory

Rate of reaction questions appear on every paper in every series. The key to full marks is always linking back to collision theory: every factor that changes the rate does so by changing either the frequency of collisions or the energy of collisions (or both).

Collision theory basics

For a reaction to occur, reactant particles must:

1. Collide — particles must come into contact
2. With sufficient energy — the collision energy must be at least equal to the activation energy (Ea)

Collisions that meet both conditions are called successful collisions (or effective collisions). The rate of reaction depends on the frequency of successful collisions.

Activation energy is the minimum energy that colliding particles must have for a reaction to occur. It is the energy barrier that must be overcome. Different reactions have different activation energies.

The five factors and how they work

1. Concentration (solutions) / Pressure (gases)

Effect: Increasing concentration or pressure increases the rate.

Collision theory explanation: More particles in the same volume means particles are closer together, leading to more frequent collisions. More frequent collisions means more successful collisions per second, so the rate increases.

2. Temperature

Effect: Increasing temperature increases the rate.

Collision theory explanation: Temperature has a double effect:

• Particles move faster, so they collide more frequently
• Particles have more kinetic energy, so a greater proportion of collisions have energy equal to or exceeding the activation energy

The second effect is the dominant one. A 10 C rise in temperature can roughly double the rate because the proportion of particles exceeding the activation energy increases significantly.

3. Surface area (particle size)

Effect: Smaller particles (powder vs lumps) react faster.

Collision theory explanation: Smaller particles have a larger total surface area exposed to the other reactant. More surface is available for collisions, so collisions are more frequent and the rate increases.

Note: only the surface particles can collide. Particles inside a lump cannot react until the outer layer has reacted away.

4. Catalyst

Effect: A catalyst increases the rate without being used up.

Collision theory explanation: A catalyst provides an alternative reaction pathway with a lower activation energy. With a lower energy barrier, a greater proportion of collisions have sufficient energy to react, increasing the frequency of successful collisions.

Important: a catalyst is not used up in the reaction. It is chemically unchanged at the end and can be reused. It does not increase the amount of product formed, only the speed at which it forms.

5. Light (photochemical reactions)

Some reactions are started or accelerated by light (e.g. the reaction of methane with bromine). Light provides energy to break bonds and initiate the reaction. This is a minor point in the syllabus.

Summary table

Factor	Change	Effect on rate	Collision theory reason
Concentration	Increase	Increases	More particles per volume, more frequent collisions
Temperature	Increase	Increases	Faster particles, more frequent collisions AND more particles exceed Ea
Surface area	Increase (smaller pieces)	Increases	More surface exposed, more frequent collisions
Catalyst	Add	Increases	Lower Ea, greater proportion of successful collisions
Pressure (gases)	Increase	Increases	Particles closer together, more frequent collisions

Interpreting rate graphs

Volume of gas collected vs time

• Steeper initial gradient = faster rate
• Curve levels off when reaction is complete (limiting reagent used up)
• Same final volume = same amount of product formed
• Higher temperature or concentration: steeper curve, reaches the same final level sooner

Mass loss vs time

• Steeper initial drop = faster rate
• Graph levels off when reaction is complete
• Same final mass loss = same total gas escaped

Key graph comparisons

Higher concentration vs lower concentration: Steeper initial curve, reaches the same final amount sooner. Same final volume/mass because the same total moles of limiting reagent are present (if only concentration is changed, not amount).

Higher temperature vs lower temperature: Steeper initial curve, same final amount. Reaches completion sooner.

Catalyst vs no catalyst: Steeper initial curve, same final amount. The catalyst does not change how much product is made, only how fast.

Smaller particles vs larger particles: Steeper initial curve, same final amount (same mass of solid used, just different particle size).

Energy profile diagrams

An energy level diagram for a catalysed vs uncatalysed reaction shows:

• Same reactant energy level
• Same product energy level
• Same overall energy change (Delta H unchanged)
• The catalysed pathway has a lower activation energy hump

The catalyst does not change the enthalpy change of the reaction, only the activation energy.

Measuring rate experimentally

Common methods examiners test:

Method	What you measure	Suitable for
Gas syringe	Volume of gas produced over time	Reactions producing gas (e.g. Mg + HCl)
Mass loss (balance)	Mass decrease over time	Gas escaping (e.g. CaCO3 + HCl)
Disappearing cross	Time for cross to disappear	Na2S2O3 + HCl (sulfur precipitate)
Colour change / timing	Time for colour to appear/disappear	Clock reactions

For the disappearing cross experiment: rate is proportional to 1/time. A shorter time means a faster rate.

Common exam mistakes

1. “Particles move faster so they collide harder” — vague. Say “a greater proportion of particles have energy exceeding the activation energy.”
2. “A catalyst makes the reaction faster by giving particles more energy” — incorrect. A catalyst lowers the activation energy; it does not add energy.
3. “Increasing surface area increases the energy of collisions” — no. Surface area affects frequency of collisions only, not their energy.
4. “The catalyst is used up” — a catalyst is chemically unchanged at the end of the reaction.
5. Confusing rate with amount: A catalyst increases rate but does not increase the total amount of product.

Worked exam questions

Zinc reacts with dilute hydrochloric acid. The reaction is faster when powdered zinc is used instead of zinc granules. Explain why, using collision theory. [3 marks]

• Powdered zinc has a larger surface area than granules [1]
• More zinc particles are exposed / available to collide with acid particles [1]
• Collisions are more frequent, so there are more successful collisions per second and the rate increases [1]

Explain why increasing the temperature increases the rate of reaction between magnesium and hydrochloric acid. [3 marks]

• At higher temperature, particles have more kinetic energy and move faster [1]
• Collisions are more frequent [1]
• A greater proportion of collisions have energy equal to or greater than the activation energy, so more collisions are successful [1]

Examiner note: The third point about activation energy is essential for full marks. Just saying “particles move faster and collide more” is worth only 2 marks.

A student investigates the reaction of calcium carbonate with hydrochloric acid. She uses a catalyst. Sketch a graph showing how the volume of gas collected changes with time, with and without a catalyst. [3 marks]

• Both curves start at the origin (0,0) [1]
• The catalysed curve is steeper initially (faster rate) [1]
• Both curves reach the same final volume (same total amount of product) [1]

The catalysed curve levels off sooner because the reaction finishes faster.