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

Section 1 of 26Collision theory

A reaction occurs only when particles collide with at least the activation energy and the correct orientation

• These are called successful (effective) collisions.
• Collisions with the wrong orientation or too little energy bounce apart without reaction.

Collision frequency

• Collision frequency the number of collisions per unit time, increased by:
• Higher concentration, so more particles in the same volume.
• Higher pressure for gases, so the same particles are in a smaller volume.
• Higher temperature, so particles move faster.
• Smaller particle size (larger surface area), so more particles are exposed.

Activation energy, Ea

Activation energy, Ea the minimum energy colliding particles need for a reaction

• It is relatively smaller for exothermic reactions than for endothermic reactions.

Section 2 of 26Measuring rates of reaction

Rate of reaction

Rate of reaction change in concentration of a reactant or product ÷ time, in mol dm^-3 s^-1

• A property that changes during the reaction is followed, such as mass, gas volume, colour or time to form a precipitate.

Mass loss

• Mass loss a balance records mass lost as gas escapes
• The gas must be dense enough, so CO₂ is suitable but H₂ is not.

Gas volume

• Gas volume gas is collected in a gas syringe or an inverted measuring cylinder over water
• The volume is read every few seconds and gives the amount of product.

Disappearing cross

• Disappearing cross times how long a cross takes to be hidden by sulfur forming from Na₂S₂O₃ + 2HCl
• It gives only one data point per run.

Concentration changes

• Concentration changes sampling and titration affect the rate unless the reaction is stopped by quenching

Section 3 of 26Rates from graphs

Rate at any instant

Rate at any instant the gradient of a concentration–time (or volume–time) graph

• Draw a tangent at that time and complete a large triangle to read the changes in y and x.
• For the initial rate, draw the tangent at t = 0.

Shape

• Shape the rate is fastest at the start and slows as reactants are used up, and a horizontal curve shows the reaction has finished

Section 4 of 26Maxwell–Boltzmann distributions and temperature

Maxwell–Boltzmann distributions and temperature

Maxwell–Boltzmann distribution

• Maxwell–Boltzmann distribution a graph of the number of particles against energy at a given temperature
• A few particles have very low energy, a few have very high energy and most are in between.
• The curve starts at the origin and never touches the energy axis at high energy.
• The area under the curve equals the total number of particles and does not change with temperature.

Effect of raising the temperature

• Effect of raising the temperature the curve flattens and its peak moves to the right
• A much larger proportion of particles have energy at least equal to E_a.
• Particles also collide more often, but the greater fraction above E_a is the more important factor.

Rule of thumb

• Rule of thumb a rise of 10 K roughly doubles the rate of many reactions

Section 5 of 26Catalysts and energy

Catalysts and energy

Catalyst

Catalyst a substance that increases the rate by providing an alternative pathway of lower activation energy and is unchanged at the end

• On a Maxwell–Boltzmann curve, E_a moves left so more particles can react, and the distribution does not change.
• On an enthalpy profile there is a lower hump with the same start and end levels.

No change to ΔH or position

• No change to ΔH or position catalysts do not alter the enthalpy change or the position of equilibrium

Homogeneous catalyst

• Homogeneous catalyst in the same phase as the reactants
• For example, reactants and catalyst all in solution.

Heterogeneous catalyst

• Heterogeneous catalyst in a different phase from the reactants
• Usually a solid with gaseous reactants, as with iron in the Haber process.
• It is the most common type in industry because it is easy to separate.

Advantages

• Advantages faster rate, lower energy costs from lower temperatures and pressures, and improved atom economy

Section 6 of 26Dynamic equilibrium

Dynamic equilibrium

Reversible reaction

• Reversible reaction products can react to reform reactants, shown with two opposing half arrows

Dynamic equilibrium

Dynamic equilibrium a state in a closed system in which the forward and reverse reactions occur at equal rates, so concentrations stay constant

• The reactions do not stop and the concentrations are not necessarily equal.
• It can be approached from either side.

Closed system

• Closed system no matter is exchanged with the surroundings
• Equilibrium involving gases can only be reached in a closed system.

Position of equilibrium

• Position of equilibrium shifting right gives more products and shifting left gives more reactants

Section 7 of 26Le Chatelier's principle

• Le Chatelier's principle if a change is made to a system in dynamic equilibrium, the position of equilibrium shifts to oppose the change

Change	Shift
Increase concentration of a reactant	Towards products
Remove a product	Towards products
Increase pressure	Towards the side with fewer moles of gas
Increase temperature	In the endothermic direction
Add a catalyst	No shift; equilibrium is reached faster

• Gases only
  • only gaseous species count when comparing moles for pressure effects, and equal moles of gas on each side give no shift

Section 8 of 26Industrial compromises

1Haber process N₂ + 3H₂ ⇌ 2NH₃, ΔH

−92 kJ mol^-1

• Raising the temperature shifts left and lowers the yield, but gives a faster rate, so about 450 °C is used.
• Raising the pressure shifts right because 4 moles of gas become 2, but about 200 atm is used because higher pressure is expensive and hazardous.
• The iron catalyst speeds up the rate without changing the yield.
• Ammonia is liquefied and removed, and unreacted gases are recycled.

2Contact process 2SO₂ + O₂ ⇌ 2SO₃

• Contact process 2SO₂ + O₂ ⇌ 2SO₃ uses a V₂O₅ catalyst at about 450 °C and 1–2 atm
• The yield is already high, so high pressure is not worth the cost.

3Ethanol from ethene

• Ethanol from ethene uses H₃PO₄, 300 °C and 60 atm, following the same reasoning

4Factors considered

• Factors considered cost and profit, energy use and environmental impact, and health and safety

Section 9 of 26More from the specification

1Collision Theory and Activation Energy

• For a chemical change to take place, colliding particles must collide with correct orientation and sufficient energy equal to or greater than the activation energy (E_a), ensuring collisions result in successful chemical transformation.

2Maxwell-Boltzmann Distribution and Temperature

• A Maxwell-Boltzmann distribution curve shows the spread of molecular kinetic energies in a gas or liquid;
• increasing temperature shifts the peak to the right and flattens it, significantly increasing the fraction of particles possessing energy >= E_a.

3Catalysts and Activation Energy

• Adding a catalyst provides an alternative reaction pathway with a lower activation energy,
• meaning a greater area under the Maxwell-Boltzmann distribution curve exceeds E_a at existing temperature.

4Optimising Industrial Reaction Conditions

• Chemical engineers balance reaction conditions such as temperature, pressure,
• and catalyst efficiency to optimize yield and rate while minimizing fossil fuel consumption and carbon dioxide emissions.

5Carbon dioxide

• A gas that turns limewater cloudy.

6Maxwell-Boltzmann distribution

• A graph showing how particle energies are spread at a given temperature.

Section 10 of 26Collision theory

Successful collision

• Successful collision particles collide with the correct orientation and at least the activation energy, E_a

Rate depends on collision frequency and collision energy

• Collision frequency rises with concentration, total pressure for gases, temperature and smaller particle size (greater surface area).

Rate of reaction

Rate of reaction change in amount of a reactant or product ÷ time, with units mol dm^-3 s^-1 for concentration

Section 11 of 26Methods of measuring rates

Methods of measuring rates

Mass loss

• Mass loss a balance records the mass lost as gas escapes

Gas volume

• Gas volume gas is collected in a gas syringe or an inverted measuring cylinder over water, reading the volume every few seconds
• This is a measure of the amount of product.

Colour change or precipitate

• Colour change or precipitate the disappearing cross experiment times how long until a cross is obscured
• Na₂S₂O₃ + 2HCl → 2NaCl + SO₂ + H₂O + S.
• It generates only one data point per run.

Section 12 of 26Rates from graphs

Graph

• Graph plot the amount of product against time
  • (or reactant remaining)

Gradient

Gradient the rate

• Draw a tangent to the curve at the required time and complete a large triangle to read the change in y and x.
• Use the tangent at t = 0 for the initial rate.

Finished reaction

• Finished reaction a curve that becomes horizontal shows that the reaction has finished

Section 13 of 26Maxwell–Boltzmann distributions

1Maxwell–Boltzmann distribution

• Maxwell–Boltzmann distribution plots the number of particles against their energy at a certain temperature
• A few particles have very low energy, a few have very high energy and most are in between.
• The area under the curve equals the total number of particles.

2Drawing it

• Drawing it the curve starts at the origin, never touches the x-axis, and the area beyond E_a is shaded

Section 14 of 26Effect of temperature on rate

• At a higher temperature the curve flattens and its peak moves to the right
  • More particles have energy greater than or equal to E_a.

1

Particles gain kinetic energy and move faster.

2

There are more collisions per second.

3

A much larger fraction of the collisions have energy at least equal to E_a.

4

There are more effective collisions per second, so the rate increases.

• Importance
  • the greater proportion above E_a is the more important factor

Section 15 of 26Catalysts and energy

Catalyst

• Catalyst increases the rate by providing an alternative pathway with a lower activation energy
• E_a moves to the left on the Maxwell–Boltzmann curve, so a greater proportion of particles can react.
• The distribution itself does not change.

Unchanged

• Unchanged catalysts are not used up and do not change the enthalpy change or the position of equilibrium

Homogeneous catalyst

• Homogeneous catalyst in the same phase as the reactants, for example all in solution

Heterogeneous catalyst

• Heterogeneous catalyst in a different phase, usually a solid with gaseous reactants
• It is the most common type in industry because it is easy to separate.

Benefits

• Benefits lower energy costs from lower temperatures and pressures, and improved selectivity

Section 16 of 26More from the specification

More from the specification

Dynamic Equilibrium Characteristics

• In a closed system, dynamic chemical equilibrium is established when the rate of the forward reaction equals the rate of the reverse reaction,
• and macroscopic concentrations of reactants and products remain constant over time.

Le Chatelier's Principle and Temperature Shifts

• When temperature increases, the equilibrium position shifts in the direction of the endothermic reaction to absorb excess thermal energy, altering the equilibrium constant and changing the overall equilibrium yield.

Maxwell-Boltzmann Distribution and Temperature Effects

• Raising system temperature shifts the peak of the Maxwell-Boltzmann energy distribution curve to the right and flattens it, significantly increasing the proportion of particles with energy equal to or exceeding activation energy.

Industrial Compromise Conditions for Maximum Yield

• In industrial processes like the Haber process, operating conditions represent an economic compromise between high equilibrium yield (favored by lower temperatures) and acceptable reaction rates.
  • (temperature and pressure)

Closed system

• A system in which nothing can enter or leave.

Collision theory

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

Endothermic

• A reaction that takes in thermal energy from the surroundings.

Dynamic equilibrium

• The state in which the forward and reverse reactions occur at the same rate.

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.

Section 17 of 26Reversible reactions and dynamic equilibrium

Reversible reactions and dynamic equilibrium

Reversible reaction

• Reversible reaction can proceed in both directions, shown with two opposing half arrows

Dynamic equilibrium

• Dynamic equilibrium in a closed system the forward and reverse reactions occur at equal rates
• Concentrations stay constant although both reactions continue.
• The concentrations are not necessarily equal.

Open system

• Open system matter or energy can be lost, so equilibrium cannot be established
• A gas such as CO₂ escaping lets a decomposition go to completion.

Position of equilibrium

• Position of equilibrium the relative amounts of products and reactants
• A shift to the right means more products and a shift to the left means more reactants.

Section 18 of 26Le 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

Change	Shift	Reason
Increase concentration of a reactant	Towards products	Uses up some of the added reactant
Remove a product	Towards products	Replaces some of the product removed
Increase pressure	Towards the side with fewer moles of gas	Reduces the pressure
Increase temperature	In the endothermic direction	Absorbs the extra heat
Add a catalyst	No shift	Speeds both directions equally; equilibrium is reached sooner

Section 19 of 26Effect of concentration

Effect of concentration

Adding a reactant or removing a product

• Adding a reactant or removing a product shifts the position to the right

Adding a product or removing a reactant

• Adding a product or removing a reactant shifts the position to the left

Increasing yield

• Increasing yield increase the concentration of reactants or remove products as they form
• Adding water to an aqueous ionic equilibrium dilutes all ions equally, so there is no change when the moles of ions are equal on both sides.

Section 20 of 26Effect of pressure

Effect of pressure

Gaseous species only

• Gaseous species only pressure affects only reactions in which reactants or products are gases

Increase in pressure

• Increase in pressure shifts towards the side with fewer moles of gas

Decrease in pressure

• Decrease in pressure shifts towards the side with more moles of gas

Solids do not count

• For CaCO₃(s) ⇌ CaO(s) + CO₂(g), only the CO₂ counts, so a pressure increase shifts left.

Equal moles of gas on each side

• Equal moles of gas on each side no change in the position

Section 21 of 26Effect of temperature

Increase in temperature

• Increase in temperature shifts in the endothermic direction

Decrease in temperature

• Decrease in temperature shifts in the exothermic direction

Deducing ΔH

Type	Description
Deducing ΔH	if more product forms at higher temperature, the forward reaction is endothermic
Example	CO2 yield increases from Ag2CO3 ⇌ Ag2O + CO2 on heating.

Section 22 of 26The yield–rate compromise

Haber process N₂ + 3H₂ ⇌ 2NH₃, exothermic in the forward direction

• Higher temperature raises the rate but lowers the equilibrium yield.
• Lower temperature raises the yield but the rate is too slow, so about 450 °C is used with an iron catalyst.
• Higher pressure raises the yield because 4 moles of gas become 2.
• About 200 atm is used because higher pressure needs expensive equipment and more energy and is more hazardous.
• Unreacted gases are recycled.

Rate versus yield

• Rate versus yield a faster rate with lower yield can give more product overall
• For example, 100 g per hour gives 2400 g per day, which is better than 1000 g per day.

Section 23 of 26Other industrial considerations

1Contact process 2SO₂ + O₂ ⇌ 2SO₃

• Contact process 2SO₂ + O₂ ⇌ 2SO₃ exothermic with 3 moles of gas becoming 2
• Low temperature and high pressure favour the product.
• In practice about 450 °C, 1–2 atm and a vanadium(V) oxide catalyst are used because the yield is already high and high pressure is costly.

2

Iron(III)–thiocyanate equilibrium Fe³⁺ + SCN⁻ ⇌ [FeSCN]²⁺

• Adding Fe³⁺ deepens the red colour.
• Adding NaOH removes Fe³⁺ as the hydroxide, so the colour fades.

3Factors weighed

• Factors weighed cost and profit, energy and environmental impact, and health and safety

Section 24 of 26More from the specification

More from the specification

Collision frequency

• Number of collisions per second

Effective collision

• Collision that leads to products.

Collision theory

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

Particle

• A tiny portion of matter, too small to see.

Volume

• The space taken up by an object.

Activation energy

• The minimum energy colliding particles need to react.

Successful collision

• A collision with the correct orientation and enough energy to react.

Carbon dioxide

• A gas that turns limewater cloudy.

Maxwell-Boltzmann distribution

• A graph showing how particle energies are spread at a given temperature.

Section 25 of 26Key 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 particles need for a collision to result in reaction.
Maxwell–Boltzmann distribution	A graph showing how particle energies are distributed at a given temperature.
Catalyst	A substance that speeds up a reaction by providing a lower-energy pathway and is not used up.
Dynamic equilibrium	A state in which forward and reverse reactions occur at equal rates in a closed system.
Le Chatelier's principle	A system at equilibrium responds to a change by opposing it.

Term	Meaning
Activation energy	The minimum energy particles need for a collision to result in a reaction.
Maxwell–Boltzmann distribution	A graph of the distribution of particle energies at a given temperature.
Initial rate	The rate at the start of the reaction (t = 0).
Catalyst	A substance that speeds up a reaction by providing a lower-energy pathway without being used up.
Collision frequency	The number of collisions between particles per unit time.

Term	Meaning
Dynamic equilibrium	A state in which forward and backward reactions continue at equal rates in a closed system.
Position of equilibrium	The relative amounts of reactants and products at equilibrium.
Le Chatelier's principle	A system at equilibrium responds to a change by opposing it.
Closed system	A system in which no matter is exchanged with the surroundings.
Reversible reaction	A reaction in which products can react to reform the reactants.

Section 26 of 26Exam technique

1

Claiming a catalyst increases the energy of the particles — it lowers the activation energy

2

Drawing the higher-temperature curve with the same peak height — it is lower and shifted right

3

Counting solids when comparing moles of gas

4

Saying equilibrium means the concentrations are equal or the reactions have stopped

5

Saying a catalyst increases the yield — it only helps equilibrium to be reached sooner

1

Claiming a catalyst increases the energy of the particles

2

Drawing the higher-temperature curve with the same peak height

3

Saying only collision frequency explains the effect of temperature

4

Using the average gradient in place of the tangent at the required time

1

Stating that a catalyst increases the yield — it only speeds up reaching equilibrium

2

Counting solids when comparing moles of gas

3

Saying equilibrium means the concentrations are equal

4

Believing the reactions stop at equilibrium — both continue at equal rates