Catalysts — AQA A-level Chemistry

SPEC 3.1.5.4–3.1.5.5 • KINETICS

Concentration, Pressure & Catalysts

How concentration, pressure and catalysts affect reaction rate through collision theory

Factors Affecting Rate

In this lesson, we explore three factors that affect reaction rate: concentration, pressure, and catalysts. Each works by changing either the frequency of collisions or the proportion of successful collisions.

Recap from Lesson 5.01

For a reaction to occur, particles must:

Collide with each other

Have energy ≥ activation energy (Eₐ)

Collide with the correct orientation

Concentration and pressure affect collision frequency — how often particles collide per second. Catalysts affect the activation energy — making it easier for collisions to be successful.

FactorWhat It ChangesEffect on Rate

ConcentrationCollision frequencyHigher conc. → faster rate

Pressure (gases)Collision frequencyHigher pressure → faster rate

CatalystActivation energyLower Eₐ → faster rate

Effect of Concentration

Key Principle

Increasing the concentration of a reactant increases the rate of reaction.

When concentration increases, there are more reactant particles per unit volume. This means particles are closer together and collide more frequently, leading to a higher frequency of effective collisions.

The diagram shows a higher concentration of particles on the right — more particles present in the same volume means the chances and frequency of collisions between reacting particles is increased, causing an increased rate of reaction.

Explanation Chain

Higher concentration → More particles per unit volume → Particles closer together → Higher collision frequency → Higher frequency of effective collisions → Faster reaction rate

M-B Diagram Note

Increasing concentration does not change the shape of the Maxwell-Boltzmann curve (the energy distribution stays the same). However, there are more particles overall, so the curve would be taller with a larger area underneath.

Effect of Pressure (Gases Only)

Key Principle

Increasing the pressure of a gaseous reaction increases the rate of reaction.

Increasing pressure has the same effect as increasing concentration. When pressure increases, gas particles are forced into a smaller volume, so there are more particles per unit volume.

The diagram shows that at higher pressure, the same number of particles occupy a smaller volume. When the pressure is increased, the molecules have less space in which they can move, resulting in an increased collision frequency and therefore increased rate of reaction.

Explanation Chain

Higher pressure → Particles in smaller volume → Particles closer together → Higher collision frequency → Higher frequency of effective collisions → Faster reaction rate

Important

Pressure only affects reactions involving gases. For reactions in solution, use concentration instead. Liquids and solids are essentially incompressible.

Worked Example01

Explain, using collision theory, why increasing the concentration of hydrochloric acid increases the rate of its reaction with magnesium ribbon.

1

Increasing the concentration of HCl means there are more H⁺ ions per unit volume.

2

The particles are closer together, so they collide more often.

3

This increases the frequency of collisions between H⁺ ions and magnesium atoms.

4

Higher collision frequency leads to a higher frequency of effective collisions (where particles have E ≥ Eₐ and correct orientation).

5

Therefore, the rate of reaction increases.

Key points: more particles per volume → closer together → more collisions → more effective collisions → faster rate

What is a Catalyst?

Definition

A catalyst is a substance that increases the rate of a chemical reaction without being changed in chemical composition or amount at the end of the reaction.

Catalysts work by providing an alternative reaction pathway that has a lower activation energy than the uncatalysed reaction. This means more particles have sufficient energy to react.

PropertyExplanation

Increases rateReaction happens faster with catalyst present

Not used upSame amount of catalyst at start and end

Chemically unchangedCatalyst has same composition after reaction

Alternative pathwayDifferent mechanism with lower Eₐ

Types of Catalyst

Homogeneous catalyst: In the same phase as the reactants (e.g., all in solution).

Heterogeneous catalyst: In a different phase to the reactants (e.g., solid catalyst with gaseous reactants).

Exam Tip

A catalyst is not a reactant. It participates in the mechanism but is regenerated. Always state that the catalyst provides an "alternative pathway with lower activation energy" — this exact phrase is expected in mark schemes.

How Catalysts Work

The catalyst provides an alternative reaction route with a lower activation energy. This doesn't change the energy of reactants or products — only the energy barrier between them.

The diagram shows that the catalyst allows the reaction to take place through a different mechanism, which has a lower activation energy than the original reaction. The energy of reactants and products stays the same — only the activation energy barrier is reduced.

Key Concept

The catalyst lowers the activation energy (Eₐ). With a lower Eₐ, a greater proportion of molecules have energy ≥ Eₐ, so more collisions are successful and the reaction rate increases.

Catalysts on M-B Diagrams

A catalyst does not change the Maxwell-Boltzmann distribution curve. The curve stays exactly the same shape and position. Instead, the catalyst lowers the activation energy line on the diagram.

The total shaded area (purple + red) under the curve shows the number of particles with energy greater than Eₐ when a catalyst is present. This area is much larger than the red shaded area alone, which shows the number of particles with energy greater than Eₐ without a catalyst. By lowering Eₐ, a greater proportion of molecules have sufficient energy for an effective collision.

Critical Distinction

Temperature changes the shape of the M-B curve (shifts right, becomes broader).

Catalyst does NOT change the curve — it only moves the Eₐ line to the left.

Worked Example02

Explain, with reference to a Maxwell-Boltzmann distribution diagram, how a catalyst increases the rate of a reaction.

1

A catalyst provides an alternative reaction pathway with a lower activation energy.

2

On a Maxwell-Boltzmann diagram, the distribution curve stays the same (temperature hasn't changed).

3

The Eₐ line moves to the left (to a lower energy value).

4

The shaded area to the right of Eₐ increases, meaning a greater proportion of molecules have E ≥ Eₐ.

5

This increases the frequency of effective collisions, so the rate of reaction increases.

Key: Catalyst lowers Eₐ → more molecules above Eₐ → higher frequency of effective collisions → faster rate

Temperature vs Catalyst

Both temperature increase and adding a catalyst increase reaction rate, but they work in fundamentally different ways:

Temperature Increase

Changes the M-B distribution

Curve shifts right and becomes broader

Peak becomes lower

Eₐ stays the same

More molecules have E ≥ Eₐ because curve has shifted

Also slightly increases collision frequency

Adding a Catalyst

M-B distribution stays unchanged

Curve shape and position identical

Peak stays the same

Eₐ decreases (line moves left)

More molecules have E ≥ Eₐ because Eₐ is lower

No change to collision frequency

Summary

Temperature changes the distribution of molecular energies.

Catalyst changes the energy threshold required for reaction.

Worked Example03

A company wants to speed up an industrial reaction. Compare and contrast the effects of (a) increasing temperature and (b) using a catalyst on a Maxwell-Boltzmann diagram.

(a) Increasing temperature:

•

The M-B curve shifts to the right and becomes broader and lower.

•

The total area under the curve stays the same (same number of molecules).

•

The Eₐ line stays in the same position.

•

More molecules have E ≥ Eₐ because the distribution has shifted.

(b) Using a catalyst:

•

The M-B curve remains unchanged.

•

The Eₐ line moves to the left (lower activation energy).

•

More molecules have E ≥ Eₐ because the threshold has lowered.

Both result in more molecules exceeding Eₐ, but by different mechanisms.

Practice

Question 1: In a solution, the reaction between sodium thiosulfate (Na₂S₂O₃) and hydrochloric acid produces a yellow precipitate of sulfur. Explain why doubling the concentration of sodium thiosulfate approximately doubles the initial rate of this reaction.

Doubling the concentration of Na₂S₂O₃ means double the number of thiosulfate ions per unit volume.

The particles are closer together on average.

This doubles the frequency of collisions between S₂O₃²⁻ ions and H⁺ ions.

Double the collision frequency means double the frequency of effective collisions (where E ≥ Eₐ with correct orientation).

Therefore, the initial rate approximately doubles.

Practice

Question 2: A student claims: "A catalyst works by giving molecules more energy so they can overcome the activation energy barrier." Explain why this statement is incorrect, and give the correct explanation for how a catalyst works.

Why the statement is incorrect:

A catalyst does not change the energy of the molecules. The Maxwell-Boltzmann distribution remains unchanged when a catalyst is added.

The correct explanation:

A catalyst provides an alternative reaction pathway with a lower activation energy.

On a M-B diagram, the Eₐ line moves to the left (to a lower value).

With a lower Eₐ, a greater proportion of molecules already have E ≥ Eₐ, so the frequency of effective collisions increases and the rate increases.

Practice

Question 3: In the Haber process (N₂ + 3H₂ ⇌ 2NH₃), the gases are compressed to high pressure (around 200 atm). Explain how this high pressure affects the rate of the forward reaction.

At high pressure, the gas molecules (N₂ and H₂) are forced into a smaller volume.

This means there are more molecules per unit volume.

The collision frequency increases — molecules collide more often per second.

More collisions per second leads to more successful collisions (with E ≥ Eₐ).

Therefore, the rate of the forward reaction increases.

Summary

Key Concepts

FactorHow It WorksEffect on M-B Diagram

ConcentrationMore particles per volume → higher collision frequencyCurve same shape but TALLER (more particles)

PressureSame as concentration (for gases only)Same as concentration

CatalystProvides alternative pathway with lower EₐCurve unchanged; Eₐ line moves LEFT

TemperatureMore molecules have E ≥ Eₐ (from 5.02)Curve shifts RIGHT, becomes broader

Catalyst Definition

A catalyst increases the rate of reaction without being chemically changed. It provides an alternative pathway with lower activation energy.

Exam Checklist

Concentration/pressure → "more particles per unit volume" → "particles closer together" → "higher frequency of effective collisions"

Catalyst → "alternative pathway" → "lower activation energy" → "greater proportion with E ≥ Eₐ"

On M-B diagram: catalyst moves Eₐ left; temperature changes the curve shape; concentration makes curve taller