2.6Unit 2. Energetics, Group Chemistry, Halogenoalkanes and Alcohols

Introduction to Kinetics & Equilibria — IAL Edexcel Linear (Chemistry) Revision Notes & Topical Questions

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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

Section 1 of 24Collision 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.

Section 2 of 24Activation 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 E_a does not react.

In an exothermic reaction the reactants are higher in energy than the products, so E_a is relatively small.

In an endothermic reaction the reactants are lower in energy than the products, so E_a is relatively large.

An effective collision needs both the correct orientation and energy of at least E_a.

Section 3 of 24Factors 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 ≥ E_a,

so its effect on rate is the largest; a rise of 10 K roughly doubles many reaction rates.

A catalyst lowers E_a, so more collisions are successful.

Section 4 of 24Rate 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.

Section 5 of 24Rate from a concentration–time graph

Plot concentration (or volume or mass) against time.

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.

The initial rate is the gradient of the tangent at time zero.

Properties used to follow rate (colour, mass, volume) are taken as proportional to concentration.

Section 6 of 24More from the specification

Section 7 of 24Measuring rate by mass loss

A gas that escapes from the vessel reduces the mass, e.g. CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂.

Record the mass every few seconds and plot mass lost against time.

The gas must be dense enough for the change to be measured on a 2 or 3 d.p. balance.

CO₂ (M_r = 44) is suitable; H₂ (M_r = 2) is not.

Section 8 of 24Measuring rate by gas volume

Measuring rate by gas volume

Collect the gas in a gas syringe, or over water in an inverted measuring cylinder if the gas is not water soluble.

Read the volume every few seconds and plot volume against time, which shows the amount of product against time.

Example reaction: Mg + 2HCl → MgCl₂ + H₂.

Section 9 of 24Other rate methods

Disappearing cross

Disappearing cross time for a cross seen through the flask to disappear as sulfur forms.
Na₂S₂O₃ + 2HCl → 2NaCl + SO₂ + H₂O + S(s).
It gives only one piece of data per experiment and is subjective.

Sampling and titrating can change the rate, unless the reaction is stopped by quenching.

Timing the point when something visible happens, e.g. a piece of magnesium dissolving, is a convenient alternative.

Faster reactions are easier to measure by averaging over the whole reaction.

Section 10 of 24Initial rate from a tangent

Initial rate from a tangent

Plot the results and draw a tangent to the curve at t = 0.

Calculate the gradient of the tangent as change in y ÷ change in x, with units.

Draw the tangent with a ruler and make the triangle large to reduce error.

Section 11 of 24Maxwell-Boltzmann distribution and temperature

A Maxwell-Boltzmann curve shows the distribution of particle energies at a given temperature.

A few particles have very low energy, a few have very high energy and most are in between.

The area under the curve is the total number of particles and does not change with temperature.
The curve starts at the origin and never touches the energy axis at high energy.

Only a small proportion of particles have energy of at least E_a.

A higher temperature flattens the curve and moves the peak to the right.

More particles have E ≥ E_a, so a larger proportion of collisions is successful.
Particles also move faster, so collisions are more frequent.

Section 12 of 24Catalysts

Catalyst

Catalyst increases the rate by providing an alternative pathway with a lower activation energy and is unchanged at the end.

On a Maxwell-Boltzmann curve, E_a moves to the left.

A larger area (more particles) is above E_a.

On a reaction profile, the catalysed route has a lower peak while the reactant and product levels, and so ΔH, are unchanged.

Advantages of catalysts

They increase the rate so more product forms in a given time.
They reduce energy costs because reactions can run at lower temperature and pressure.
They can increase atom economy by reducing unwanted reactions, e.g. zeolites in the synthesis of phenol.

Homogeneous

Types of catalyst

Homogeneous
- in the same phase as the reactants, e.g. all in solution.
Heterogeneous
- in a different phase from the reactants, e.g. a solid catalyst with gaseous reactants;
  - this is most common in industry.

Section 13 of 24Reversible reactions and dynamic equilibrium

Some reactions go to completion and stop when a reactant is used up.

In a reversible reaction the products can react to re-form the reactants, shown with ⇌.

Dynamic equilibrium

Dynamic equilibrium the rate of the forward reaction equals the rate of the backward reaction in a closed system, so concentrations are constant.
Molecules keep reacting in both directions.
In 2HI ⇌ H₂ + I₂, HI breaks down at the same rate as it forms.
The same equilibrium is reached starting from reactants or from products.

Concentrations at equilibrium are constant, not equal.

Section 14 of 24Closed and open systems

Closed and open systems

Closed system

Closed system no reactants or products escape.

Open system

Open system matter and energy can be lost to the surroundings.

Reactions entirely in solution can reach equilibrium in open flasks because little material is lost.

Reactions involving gases reach equilibrium only in a closed system.

CaCO₃ ⇌ CaO + CO₂ reaches equilibrium when CO₂ cannot escape.
In an open vessel CO₂ is lost and the decomposition goes to completion.

Section 15 of 24More from the specification

1Dynamic chemical equilibrium

A reversible reaction in a closed system reaches dynamic equilibrium when the rate of the forward reaction equals the rate of the backward reaction,
which means the concentrations of reactants and products remain constant.

2Le Chatelier's principle and equilibrium shifts

Changing reaction conditions (temperature, pressure,
or concentration) causes the equilibrium position to shift in the direction that opposes the disturbance, directly influencing product yield.

3Pressure effects on gas equilibria

For reactions involving gases such as the decomposition of carbonates to form carbon dioxide, an increase in system pressure shifts the equilibrium towards the side with fewer moles of gas.

4Equilibrium concentration-time graph

On a concentration-time graph, equilibrium is identified where the curves for reactants consumed and products formed become horizontal parallel lines, indicating no further change in composition.

5Carbon dioxide

A gas that turns limewater cloudy.

6Collision theory

Particles react only if they collide with at least the activation energy and the correct orientation.

Section 16 of 24Position of equilibrium

Position of equilibrium

Position of equilibrium the relative amounts of products and reactants in the equilibrium mixture.

A shift to the left increases reactant concentration; a shift to the right increases product concentration.

Le Chatelier's principle

Le Chatelier's principle if a change is made to a system in dynamic equilibrium, the position of equilibrium moves to counteract the change.

A catalyst does not change the position; it only helps equilibrium to be reached faster.

Section 17 of 24Effect of concentration

Increasing the concentration of a reactant shifts the position to the right to use it up.

Removing a product shifts the position to the right to replace it.

In CH₃COOH + C₂H₅OH ⇌ CH₃COOC₂H₅ + H₂O, adding ester or removing ethanol shifts the position to the left.

Adding water to Ce⁴⁺ + Fe²⁺ ⇌ Ce³⁺ + Fe³⁺ has no effect, because dilution changes all ions equally.

Section 18 of 24Effect of pressure

Pressure affects only equilibria involving gases.

Increasing pressure shifts the position towards the side with fewer moles of gas; decreasing pressure shifts it towards more moles.

Solids and liquids are not counted.

N₂O₄(g) ⇌ 2NO₂(g): higher pressure shifts left (1 mole against 2).

CaCO₃(s) ⇌ CaO(s) + CO₂(g): only CO₂ is a gas, so higher pressure shifts left.

2NO₂ ⇌ 2NO + O₂: lower pressure shifts right (2 moles against 3).

Section 19 of 24Effect of temperature and deducing ΔH

Effect of temperature and deducing ΔH

Increasing temperature shifts the position in the endothermic direction, absorbing the added energy.

Decreasing temperature shifts it in the exothermic direction.

If more product forms on heating, the forward reaction is endothermic.

H₂ + CO₂ ⇌ H₂O + CO has ΔH = +41.2 kJ mol⁻¹, so heating shifts it right.
For Ag₂CO₃(s) ⇌ Ag₂O(s) + CO₂(g), more CO₂ on heating shows the forward reaction is endothermic.

Section 20 of 24The Haber process compromise

N₂(g) + 3H₂(g) ⇌ 2NH₃(g) is exothermic in the forward direction.

Temperature

A higher temperature increases the rate but shifts the position left, lowering the yield.
A lower temperature gives a better yield but the rate is too slow.
About 450 °C gives an acceptable yield (roughly 35%) in an acceptable time.
Producing 100 g per hour (2400 g per day) can be better than 1000 g per day.

Pressure

Higher pressure increases the yield (4 moles of gas become 2) and the rate.
About 20 MPa (200 atm) is used; doubling it to about 40 MPa raises the yield only by about 7%.
High pressure costs more energy and equipment and is hazardous.

An iron catalyst speeds up the reaction without changing the yield, and unreacted gases are recycled.

Section 21 of 24Factors in industrial compromises

Financial

Financial does the extra yield pay for the cost of altered conditions?

Energy and environment

Energy and environment do the conditions need more energy or more fossil fuel?

Health and safety

Health and safety are the conditions safe for workers and the plant?

The same ideas apply to other processes, e.g. the Contact process 2SO₂ + O₂ ⇌ 2SO₃ (V₂O₅ catalyst, about 450 °C, 1–2 atm) and ethanol from ethene (H₃PO₄, 300 °C, 60 atm).

Section 22 of 24More from the specification

Section 23 of 24Key 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.

Term	Meaning
Initial rate	The rate at time zero, found from the gradient of a tangent.
Maxwell-Boltzmann distribution	A graph showing how particle energies are spread at a given temperature.
Catalyst	A substance that increases rate by an alternative pathway of lower activation energy and is unchanged at the end.
Heterogeneous catalyst	A catalyst in a different phase from the reactants.
Dynamic equilibrium	A state in a closed system where forward and reverse rates are equal and concentrations are constant.
Closed system	A system in which no matter enters or leaves.
Quenching	Stopping a reaction in a sample so it can be analysed.

Term	Meaning
Position of equilibrium	The relative amounts of reactants and products in the equilibrium mixture.
Le Chatelier's principle	A system at equilibrium shifts to oppose a change imposed on it.
Endothermic direction	The direction of reaction that absorbs energy.
Compromise conditions	Conditions chosen to balance rate, yield and cost.
Recycling	Returning unreacted reactants to the reactor.

Section 24 of 24Exam technique

Saying a higher temperature only gives more frequent collisions — more particles also have energy above the activation energy.

Saying a catalyst raises the energy of the particles — it provides a lower-energy pathway.

Saying all collisions lead to reaction — only collisions with enough energy and correct orientation do.

Drawing a tangent that does not touch the curve at the point required.

Saying concentrations are equal at equilibrium — they are constant.

Saying the reaction stops at equilibrium — both reactions continue at equal rates.

Saying a catalyst changes ΔH — it only lowers the activation energy.

Saying the peak of the curve rises when temperature rises — it moves right and becomes lower.

Using hydrogen with a mass-loss method — its mass is too small to measure.

Counting solids when comparing moles of gas — only gases count.

Saying a catalyst increases yield — it speeds up both directions equally.

Thinking a higher temperature always increases yield — it lowers yield for exothermic reactions.

Confusing rate and position — temperature and pressure change both, but a catalyst changes rate only.

Last updated: 18 October 2026