2.6Unit 2. Energetics, Group Chemistry, Halogenoalkanes and Alcohols
Introduction to Kinetics & Equilibria — IAL Edexcel Linear (Chemistry) Revision Notes & Topical Questions
15 min read 24 parts Lesson 6 of 53 1 quiz
Lesson summary
2.6 Introduction to Kinetics & Equilibria — collision theory and rate of reaction — IAL EDEXCEL LINEAR (CHEMISTRY) notes: Use collision theory to explain how concentration, pressure, temperature, surface area and activation energy affect rate, and define and calculate the rate of…
- Successful (effective) collision
- Unsuccessful (ineffective) collision
- Activation energy, Ea
- Pressure (gases)
- Surface area (smaller particles)
- Maxwell–Boltzmann distribution
Part 1 of 7Collision theory
Collision theory
Reacting particles must collide, and some collisions break old bonds and form new bonds.
The rate depends on:
- collision frequency,
- collision energy,
- activation energy.
Successful (effective) collision
- Successful (effective) collision particles collide with the correct orientation and with at least the activation energy.
Unsuccessful (ineffective) collision
- Unsuccessful (ineffective) collision the particles have the wrong orientation or too little energy and simply bounce apart.
Rate of reaction depends on the number of successful collisions per unit time.
Part 2 of 7Activation energy
Activation energy, Ea
- Activation energy, Ea the minimum energy colliding particles need for a reaction to occur.
Even with the correct orientation, a collision without at least Ea does not react.
In an exothermic reaction the reactants are higher in energy than the products, so Ea is relatively small.
In an endothermic reaction the reactants are lower in energy than the products, so Ea is relatively large.
An effective collision needs both the correct orientation and energy of at least Ea.
Part 3 of 7Factors changing collision frequency
Factors changing collision frequency
Concentration
- Concentration more particles in the same volume give more frequent collisions.
Pressure (gases)
- Pressure (gases) the same number of particles in a smaller volume gives more frequent collisions.
Temperature
- Temperature particles have more kinetic energy and move faster, so collisions are more frequent.
Surface area (smaller particles)
- Surface area (smaller particles) more particles are exposed in the same volume, so more collisions occur.
Raising temperature also increases the proportion of collisions with E ≥ Ea,
- so its effect on rate is the largest; a rise of 10 K roughly doubles many reaction rates.
A catalyst lowers Ea, so more collisions are successful.
Part 4 of 7Summary of factors affecting rate
| Change | Effect on collisions | Effect on rate |
|---|---|---|
| Increase concentration / pressure | More particles per volume, more frequent collisions | Increases |
| Increase temperature | More frequent collisions and many more with E ≥ Ea | Large increase |
| Increase surface area | More exposed particles to collide | Increases |
| Add a catalyst | Lower Ea so more collisions are successful | Increases |
Part 5 of 7Rate of reaction
Rate of reaction
- Rate of reaction the change in concentration of a reactant or product per unit time, in mol dm−3 s−1.
Rate = change in amount of reactant or product (mol dm−3) ÷ time (s).
During a reaction reactant concentration falls and product concentration rises.
- The rate is fastest at the start and slows as reactants are used up.
Reaction progress can be followed by measuring the reactant lost, the product formed, the time for a colour change or the time for a precipitate to form.
Part 6 of 7Rate from a concentration–time graph
1
Plot concentration (or volume or mass) against time.
2
The rate at any instant is the gradient of a tangent to the curve.
- Draw the tangent at the required time.
- Complete a large triangle, ideally intersecting gridlines, and read off the change in y and x.
- Gradient = change in y ÷ change in x.
3
The initial rate is the gradient of the tangent at time zero.
4
Properties used to follow rate (colour, mass, volume) are taken as proportional to concentration.
Part 7 of 7More from the specification
More from the specification
Analytical chromatographic separation
- Analytical chromatographic separation differential partitioning between mobile and stationary phases provides an effective means to isolate reaction components.
Carbon dioxide
- A gas that turns limewater cloudy.
Maxwell-Boltzmann distribution
- A graph showing how particle energies are spread at a given temperature.
Reversible reaction
- A reaction in which the products can react to re-form the reactants.
Closed system
- A system in which nothing can enter or leave.
Maxwell–Boltzmann distribution
- Maxwell–Boltzmann distribution a curve of number of molecules against energy, where the area under it equals the total number of molecules.
Dynamic equilibrium
- Dynamic equilibrium forward and reverse reactions continue at equal rates, so concentrations stay constant, though not necessarily equal.
Position of equilibrium
- Position of equilibrium if a condition changes, the system shifts to oppose the change, for example removing a product shifts it right and raises yield.
Industrial conditions
- Industrial conditions chosen as a compromise between yield, rate, energy cost and safety.
Key terms
| Term | Meaning |
|---|---|
| Rate of reaction | The change in concentration of a reactant or product per unit time. |
| Collision theory | Reactions occur only when particles collide with the correct orientation and at least the activation energy. |
| Activation energy | The minimum energy colliding particles need to react. |
| Successful collision | A collision with the correct orientation and enough energy to react. |
| Collision frequency | The number of collisions between particles per unit time. |
| Initial rate | The rate at time zero, found from the tangent at the start of the curve. |
Avoid these mistakes
1
Saying a higher temperature only gives more frequent collisions — more particles also have energy above the activation energy.
2
Saying a catalyst raises the energy of the particles — it provides a lower-energy pathway.
3
Saying all collisions lead to reaction — only collisions with enough energy and correct orientation do.
4
Drawing a tangent that does not touch the curve at the point required.
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