Reaction Rates and Collision Theory

Learning Objectives (OCR Spec 3.2.2)

By the end of this lesson you should be able to:

• Define rate of reaction and state its units
• Describe how reaction rates can be measured experimentally
• Explain collision theory: effective collisions require sufficient energy and correct orientation
• Define activation energy Ea and describe its role in the rate of a reaction
• Sketch energy-profile diagrams and identify Ea

Rate of Reaction

The rate of reaction is the change in the concentration (or amount) of a reactant or product per unit time:

rate = Δ[concentration] / Δt

Units: mol dm⁻³ s⁻¹ (or mol dm⁻³ min⁻¹ for slow reactions). Rate can also be expressed in terms of mass or volume change per unit time.

Because the concentration of reactants falls and products rise during a reaction, a rate vs time graph is a curve that is steepest at t = 0 (when concentrations are highest) and levels off as the reaction proceeds.

Measuring Rate Experimentally

Any measurable property that changes during the reaction can be tracked against time. Common methods include:

Method	Suitable for	Comments
Volume of gas collected	Reactions producing gas (CaCO3 + HCl, Mg + HCl)	Gas syringe or water displacement
Mass loss	Reactions producing gas that escapes	Top-pan balance with cotton wool plug
pH	Reactions involving H⁺ or OH⁻	pH meter or data logger
Colour change	Reactions with coloured species (MnO4⁻, I2)	Colorimeter
Conductivity	Change in ion concentration	Conductivity probe
Titration of quenched samples	Slow reactions where small aliquots can be withdrawn	Sudden dilution + titration

The initial rate (slope at t = 0) can then be extracted from concentration-time graphs for each experiment.

Collision Theory

Reactions occur because particles collide. However, not every collision leads to a reaction. Collision theory states that for a reaction to occur the colliding particles must have:

1. Sufficient energy - at least the activation energy Ea
2. The correct orientation - reactive parts of the molecules must hit each other

A collision that meets both criteria is called an effective collision or a successful collision. Only a small fraction of all collisions are effective.

graph TD
  A[Two molecules collide] --> B{Enough energy?}
  B -->|No| C[Bounce apart - no reaction]
  B -->|Yes| D{Correct orientation?}
  D -->|No| C
  D -->|Yes| E[Reaction occurs]

Activation Energy Ea

Activation energy is the minimum energy that colliding particles must possess for a reaction to occur. Think of it as the energy barrier that must be overcome for reactant bonds to start breaking and product bonds to form. Ea is a property of the reaction and does not depend on the initial concentrations.

A high Ea means very few collisions are effective, so the rate is slow. A low Ea means a large fraction of collisions succeed, and the rate is fast. Catalysts reduce Ea (more on this in Lesson 9).

Energy Profile Diagram

An enthalpy (or energy) profile diagram shows how the potential energy of the system changes during the reaction. It identifies:

• The reactants (left) and products (right) levels
• The transition state peak
• The activation energy Ea (distance from reactants to peak)
• The overall enthalpy change ΔH (vertical distance from reactants to products)

For an exothermic reaction:

graph LR
  A[Reactants] -->|Ea forward| B[Transition state]
  B -->|Ea reverse > Ea forward| C[Products lower]

The activation energy for the reverse reaction is Ea(forward) + |ΔH| because the products must climb back over the same peak.

For an endothermic reaction the products sit above the reactants, and Ea(forward) > Ea(reverse).

Factors That Affect the Rate