UNIT 2 CHEMISTRY · PEARSON AAQ · D1.2.2
Collision Theory and Activation Energy
FREE CHEMISTRY REVISION · TOPIC 34
Collision Theory and Activation Energy
Pearson BTEC Level 3 AAQ Applied Science
Unit 2: Principles and Applications of Chemistry · D1.2.2
What you will learn
Use collision theory to explain reaction-rate changes. Distinguish a collision from a successful collision, interpret activation energy, and explain the effects of temperature, concentration, gas pressure, solid surface-area-to-volume ratio and catalysis. Detailed Maxwell–Boltzmann curves belong to D1.2.4 and are not developed here.
1. From contact to a successful collision
Reacting particles must encounter one another for a collision-based reaction to occur. A collision is successful only when the encounter has enough energy for the reaction pathway and, where molecular geometry matters, a suitable relative orientation. Many encounters simply separate again without a chemical change.
A successful encounter can lead to rearrangement of electrons and bonds and the formation of products. Do not imagine that every collision breaks every bond, or that particles must come to a complete stop before reacting. Collision theory is an introductory model; detailed reactions may involve several elementary steps, surfaces or intermediates.
2. What activation energy means
Activation energy, Eₐ, is the energy barrier associated with a reaction pathway. In the introductory collision model, colliding particles need sufficient energy to reach a high-energy arrangement from which products can form. Activation energies are commonly given as molar quantities in kJ mol⁻¹. A single-particle collision energy and a molar activation energy must use compatible units if compared quantitatively.
On a simple energy profile, the forward activation energy is the difference between the reactant level and the pathway’s highest point. It is not the energy difference between reactants and products. The latter gives the overall energy change for the reaction in the chosen description. Even an exothermic reaction can have a substantial activation barrier and be slow at room temperature.
A high peak does not mean that the reaction takes longer along a time axis: the horizontal axis of an energy profile is reaction progress, not time. Real multistep catalysed pathways may have several peaks and intermediates; a single lower peak is a simplified comparison.
3. Concentration and gas pressure: encounter frequency
At fixed temperature, a higher concentration means more reacting particles in each unit volume. Appropriate particles encounter one another more often per unit volume, so the frequency of potentially successful collisions usually increases. The mean kinetic energy and the fraction exceeding the same energy threshold do not increase merely because concentration has risen.
For gaseous reactants compressed at the same temperature, the concentration and reactant partial pressure rise. This usually increases collision frequency per unit volume. Adding inert gas at unchanged volume and temperature increases total pressure but leaves reactant concentrations unchanged, so the simple encounter-frequency argument does not predict a faster reaction.
Keep the wording precise: a collision frequency is a number of encounters per time, often also per unit volume. A reaction rate expressed in mol dm⁻³ s⁻¹ is an amount change per volume per time. The two are related by the mechanism and the fraction of productive encounters; they are not interchangeable numbers.
4. Temperature: frequency and collision energy
Higher temperature increases average kinetic energy and average speed. Collisions generally occur more frequently, and a larger fraction of encounters has enough energy to overcome the activation barrier. The increase in the energetic fraction is often the major reason for a substantial rate increase.
It is insufficient to say only “the particles collide more”. Explain both motion and the chance of meeting the energy requirement. At a fixed temperature particles have a range of energies: some collisions can be successful even when most are not. Raising temperature does not give every particle the same energy and does not lower the activation energy for an unchanged pathway.
Use supplied measurements to establish a numerical rate change. There is no universal exact factor for a given temperature rise. For gases in a fixed-volume vessel, heating changes temperature and pressure together; concentration initially remains fixed because the amount and volume are unchanged. That case is different from compression at constant temperature.
5. A reacting solid: available surface sites
When a solid reacts with a solution or gas, encounters must occur at accessible solid surfaces. Smaller fragments of the same total solid amount usually expose more surface, increasing the number of sites available at once. This can increase the total number of productive surface encounters per second without changing the average energy of incoming particles.
For similar particles, smaller size also increases surface-area-to-volume ratio. For a cube, SA:V = 6/a, where a is its side length. This geometrical relationship explains why a large fraction of material can be close to an accessible surface in small particles. It does not mean that the solid gains extra atoms when crushed.
Particles buried inside a solid do not collide directly with solution particles until a surface becomes exposed. Clumping, protective layers, blocked catalyst sites, stirring and diffusion can affect accessibility. These qualifications explain why simply doubling a calculated area need not produce an exact twofold observed rate increase.
6. Catalysts change the pathway
A catalyst provides an alternative reaction pathway with a lower activation energy. At the same temperature, a larger fraction of encounters can satisfy the lower energy requirement. This increases the frequency of successful reaction events without requiring hotter or faster particles.
The catalyst is regenerated overall: it is chemically unchanged and unchanged in mass at the end, and the specified reaction products are not altered. It may take part in intermediate steps. A lower activation barrier does not mean the catalyst supplies all the reaction energy or makes the products have a different energy.
Do not confuse Eₐ with the overall enthalpy change. A catalyst does not change the energy levels of the same reactants and products, the equilibrium position or the equilibrium constant at a fixed temperature. For a reversible reaction it accelerates approach to equilibrium by providing pathways for both directions; it does not guarantee complete conversion.
7. Compare the explanations
| Change | Collision frequency / sites | Energy requirement or distribution |
|---|---|---|
| Increase solution concentration at fixed T | Usually more reactant encounters per unit volume | Same mean kinetic energy and pathway Eₐ |
| Compress reacting gas at fixed T | More reactant encounters per unit volume | Same mean kinetic energy and pathway Eₐ |
| Increase temperature | Faster motion; more frequent encounters | Larger fraction has sufficient energy; Eₐ unchanged |
| Increase accessible solid area | More surface encounters possible at once | Incoming-particle energy and pathway Eₐ unchanged |
| Add appropriate catalyst at fixed T | More successful events; may introduce surface sites | Alternative pathway with lower Eₐ; no increase in mean kinetic energy |
An explanation should connect the actual change to the controlling particle behaviour. Avoid adding every rate-factor phrase to every answer. More surface area, for example, is not explained by hotter particles unless the temperature was also changed.
8. Six worked explanations
Example 1 — distinguish two kinds of collision
Two correctly oriented particles approach but have insufficient energy for the uncatalysed pathway. They collide and separate unchanged. Contact alone is not enough: the energetic requirement is not met. If a question states energy is sufficient but orientation is unsuitable, explain the failure using geometry rather than claiming that the energy disappeared.
Example 2 — doubled concentration at the same temperature
A reactant concentration rises from 0.100 to 0.200 mol dm⁻³ at 25 °C. The amount of that reactant per unit volume doubles. This creates more encounter opportunities, but the mean kinetic energy and Eₐ are unchanged. Predict an increase in rate where that reactant affects the rate; do not assert an exact doubling without a rate relationship or experimental evidence.
Example 3 — heating and successful encounters
In an illustrative model, a sample has 10 000 encounters s⁻¹, with 2.00% meeting an energy criterion. After heating it has 12 000 encounters s⁻¹, with 8.00% meeting the criterion. If the productive-orientation fraction is 0.500 in both cases, predicted successful encounters are 10 000 × 0.0200 × 0.500 = 100 s⁻¹ and 12 000 × 0.0800 × 0.500 = 480 s⁻¹. The factor is 4.80, larger than the frequency-only factor of 1.20. These are supplied model fractions, not a general rate law.
Example 4 — calculate two activation energies
For a simplified energy-level example, reactants are at 20 kJ mol⁻¹, products at −10 kJ mol⁻¹, the uncatalysed peak at 100 kJ mol⁻¹ and a catalysed peak at 65 kJ mol⁻¹, all relative to one common reference. Forward Eₐ values are 100 − 20 = 80 kJ mol⁻¹ and 65 − 20 = 45 kJ mol⁻¹. The barrier is reduced by 35 kJ mol⁻¹. The overall change is −10 − 20 = −30 kJ mol⁻¹ for both paths. These numbers illustrate the calculation and are not readings from the qualitative diagram.
Example 5 — explain crushed carbonate
Equal masses of calcium carbonate chips and powder react with equal excess-acid samples at the same temperature. Powder exposes more sites, allowing more acid–solid encounters per second initially. Its particles do not need to have more energy and its Eₐ has not automatically fallen. The same limiting carbonate amount can ultimately give the same CO₂ amount.
Example 6 — distinguish compression from inert-gas addition
A reacting gas mixture is first compressed isothermally to half its volume: reactant concentrations rise and more encounters per unit volume are expected. In a different test, inert gas is added at fixed volume and temperature: total pressure rises, but the initial reactant concentrations do not. Only the first change supports the stated collision-frequency explanation for a rate increase.
Common misconceptions
- Every collision makes products. Energy and suitable approach matter.
- Activation energy is the energy released overall. It is the pathway barrier; the net energy change is a different difference.
- Heating lowers Eₐ. It changes particle energies for the same pathway.
- Higher concentration makes every collision more energetic. At fixed temperature it changes number per volume.
- A catalyst raises particle energy. It supplies a lower-barrier pathway.
- Crushing weakens all chemical bonds in the reacting substance. The usual rate argument concerns accessible surface sites.
- The energy-profile x-axis measures seconds. It represents reaction progress.
Exam tips
- State the changed condition before applying collision theory.
- Distinguish total encounters from successful encounters.
- For temperature include a larger energetic fraction, not only faster motion.
- For solid reactants identify accessible surface and surface-area-to-volume ratio.
- Read activation energy from reactants to the relevant peak using a common reference.
- For catalysis state lower Eₐ, unchanged products and overall regeneration.
- Do not infer numerical rate laws from a qualitative particle drawing.
Check your understanding
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