What you'll learn
Physical Chemistry forms the foundation of your IGCSE Chemistry course, covering how matter behaves, changes state, and reacts. This section explores the particle model, energy transfers in reactions, reaction rates, and equilibrium. You'll develop the quantitative and conceptual skills needed to excel in both Paper 1 (multiple choice) and Paper 2 (structured questions).
Key terms and definitions
Endothermic reaction — a reaction that takes in energy from the surroundings, causing a temperature decrease; ΔH is positive
Exothermic reaction — a reaction that transfers energy to the surroundings, causing a temperature increase; ΔH is negative
Activation energy — the minimum energy that colliding particles must possess for a reaction to occur
Catalyst — a substance that increases the rate of a reaction without being used up, by providing an alternative pathway with lower activation energy
Dynamic equilibrium — a state in a reversible reaction where the forward and backward reactions occur at equal rates, and concentrations of reactants and products remain constant
Le Chatelier's Principle — when a change is made to a system at equilibrium, the position of equilibrium shifts to oppose that change
Rate of reaction — the speed at which reactants are converted into products, measured as change in concentration per unit time
Enthalpy change (ΔH) — the heat energy change in a reaction at constant pressure, measured in kJ/mol
Core concepts
States of matter and the particle model
Matter exists in three principal states: solid, liquid, and gas. The kinetic particle theory explains the properties of each state through particle arrangement, movement, and energy.
Solids have particles arranged in a fixed, regular pattern held together by strong forces of attraction. Particles vibrate about fixed positions but cannot move from place to place. Solids have a definite shape and volume.
Liquids have particles that are close together but arranged randomly. Forces of attraction are weaker than in solids, allowing particles to move past each other. Liquids have a definite volume but take the shape of their container.
Gases have particles that are far apart and move randomly at high speeds. Forces of attraction between particles are negligible. Gases fill their container completely and are easily compressed.
Changes of state occur when energy is transferred to or from a substance:
- Melting (solid → liquid) and boiling (liquid → gas) require energy input
- Freezing (liquid → solid) and condensing (gas → liquid) release energy
- Sublimation (solid → gas) and deposition (gas → solid) also occur in certain substances
During a change of state, temperature remains constant because energy is used to overcome forces between particles rather than increase their kinetic energy.
Energy changes in reactions
All chemical reactions involve energy transfers. The energy change is usually measured as heat energy gained or lost to the surroundings.
Exothermic reactions release energy to the surroundings:
- Combustion reactions (burning fuels)
- Neutralisation reactions (acid + alkali)
- Many oxidation reactions
- Most displacement reactions
- The formation of products from elements in their standard states when ΔH is negative
In exothermic reactions, the energy required to break bonds in reactants is less than the energy released when new bonds form in products. Temperature of the surroundings increases.
Endothermic reactions absorb energy from the surroundings:
- Thermal decomposition reactions
- Photosynthesis
- Dissolving certain salts (e.g., ammonium nitrate)
- Electrolysis
In endothermic reactions, more energy is needed to break bonds in reactants than is released when products form. Temperature of the surroundings decreases.
Reaction profiles (energy level diagrams) show the energy changes during reactions:
- The y-axis shows energy content
- Reactants and products are shown as horizontal lines
- The peak represents the transition state (activation energy barrier)
- For exothermic reactions, products are lower in energy than reactants
- For endothermic reactions, products are higher in energy than reactants
Rates of reaction
The rate of a chemical reaction can be increased by factors that increase the frequency and/or energy of particle collisions.
Factors affecting rate:
Concentration (for solutions) or pressure (for gases): Increasing concentration/pressure increases the number of particles in a given volume, leading to more frequent collisions and a faster rate.
Temperature: Increasing temperature gives particles more kinetic energy. They move faster, collide more frequently, and more particles possess energy equal to or greater than the activation energy.
Surface area: Increasing the surface area of solid reactants (e.g., using powder instead of lumps) provides more exposed particles available for collision, increasing rate.
Catalysts: These provide an alternative reaction pathway with lower activation energy, meaning more colliding particles have sufficient energy to react.
Measuring reaction rate:
- Volume of gas produced over time (using gas syringe or inverted burette)
- Mass loss over time (for reactions producing gas)
- Time taken for a precipitate to obscure a mark
- Change in colour/pH over time
Calculate mean rate using: rate = change in quantity / time taken
Collision theory states that for a reaction to occur:
- Particles must collide
- Collisions must have energy ≥ activation energy
- Particles must collide with correct orientation
Reversible reactions and equilibrium
Some reactions are reversible — products can react to form reactants again. This is shown using the ⇌ symbol.
For example: ammonium chloride ⇌ ammonia + hydrogen chloride
When a reversible reaction occurs in a closed system, it reaches dynamic equilibrium where:
- The forward and backward reactions continue to occur
- The rates of forward and backward reactions are equal
- Concentrations of all reactants and products remain constant
- There is no observable change in macroscopic properties
Le Chatelier's Principle predicts how the position of equilibrium shifts when conditions change:
Change in concentration:
- Adding more reactant shifts equilibrium to the right (more products formed)
- Removing a product shifts equilibrium to the right
- The system acts to oppose the change
Change in temperature:
- If forward reaction is exothermic, increasing temperature shifts equilibrium to the left (favouring the endothermic backward reaction)
- Decreasing temperature shifts equilibrium to the right
- The system moves to oppose the temperature change
Change in pressure (for gaseous systems):
- Increasing pressure shifts equilibrium toward the side with fewer gas molecules
- Decreasing pressure shifts equilibrium toward the side with more gas molecules
- If equal numbers of gas molecules on both sides, pressure has no effect on position
Catalysts have no effect on the position of equilibrium, but they allow equilibrium to be reached faster by speeding up both forward and backward reactions equally.
Industrial applications: The Haber Process
The Haber Process synthesises ammonia from nitrogen and hydrogen:
N₂(g) + 3H₂(g) ⇌ 2NH₃(g) ΔH = -92 kJ/mol (exothermic)
Industrial conditions:
- Temperature: 450°C (compromise temperature)
- Pressure: 200 atmospheres
- Catalyst: iron
Explaining the conditions:
Higher pressure favours ammonia formation (4 molecules on left, 2 on right), but very high pressures are expensive and require strong, costly equipment. 200 atmospheres is an economic compromise.
Lower temperature favours the exothermic forward reaction, producing more ammonia. However, low temperatures make the reaction too slow. 450°C is a compromise between yield and rate.
The iron catalyst speeds up both reactions, allowing equilibrium to be reached faster, but doesn't affect the position of equilibrium.
Unreacted nitrogen and hydrogen are recycled to improve efficiency and reduce waste.
Calculations involving energy changes
You may need to calculate energy changes from experimental data.
Energy change = mass × specific heat capacity × temperature change
q = m × c × ΔT
Where:
- q = energy change in joules (J)
- m = mass of substance (usually water) in grams (g)
- c = specific heat capacity (4.2 J/g°C for water)
- ΔT = temperature change in °C
To find enthalpy change per mole (ΔH), divide the energy change by the number of moles that reacted.
Example: When 0.5 mol of ethanol burns, it heats 200 g of water from 20°C to 65°C.
Energy transferred to water = 200 × 4.2 × (65-20) = 37,800 J = 37.8 kJ
Enthalpy change per mole = -37.8 / 0.5 = -75.6 kJ/mol (negative because exothermic)
Worked examples
Example 1: Calculating rate of reaction
Question: In an experiment, 48 cm³ of carbon dioxide gas is produced in 2 minutes when calcium carbonate reacts with hydrochloric acid. Calculate the mean rate of reaction in cm³/s.
Solution:
- Volume produced = 48 cm³
- Time = 2 minutes = 2 × 60 = 120 seconds
- Rate = change in volume / time
- Rate = 48 / 120 = 0.4 cm³/s
(1 mark for converting to seconds, 1 mark for correct calculation, 1 mark for units)
Example 2: Applying Le Chatelier's Principle
Question: The Contact Process produces sulfur trioxide: 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) ΔH = -197 kJ/mol
Explain how increasing the temperature affects the yield of sulfur trioxide. (3 marks)
Solution:
- The forward reaction is exothermic (ΔH is negative) (1)
- Increasing temperature favours the endothermic reaction / backward reaction (1)
- Therefore, the yield of sulfur trioxide decreases / equilibrium shifts to the left (1)
Example 3: Energy calculations
Question: A student adds 5.0 g of ammonium nitrate to 50.0 g of water in a polystyrene cup. The temperature decreases from 21.0°C to 16.5°C.
(a) Calculate the energy absorbed when the ammonium nitrate dissolves. (Specific heat capacity of water = 4.2 J/g°C) (2 marks)
(b) State whether this is an endothermic or exothermic change. (1 mark)
Solution:
(a) q = m × c × ΔT q = 50.0 × 4.2 × (21.0 - 16.5) (1) q = 50.0 × 4.2 × 4.5 = 945 J (or 0.945 kJ) (1)
(b) Endothermic (temperature decreased, so energy was absorbed) (1)
Common mistakes and how to avoid them
Confusing exothermic and endothermic: Remember EXothermic = EXit of energy. Temperature of surroundings rises in exothermic reactions and falls in endothermic reactions.
Thinking catalysts shift equilibrium position: Catalysts only speed up the rate at which equilibrium is reached; they don't change the equilibrium position or yield.
Incorrect units in rate calculations: Always check whether time should be in seconds or minutes. The question will specify the units required for the answer.
Assuming all reactions go to completion: In reversible reactions at equilibrium, significant amounts of both reactants and products are present. The position of equilibrium determines the ratio.
Forgetting to consider all gas molecules when applying pressure changes: Count the total number of gaseous molecules on each side of the equation. Solids and liquids don't count.
Not converting temperature to Kelvin when required: For IGCSE, temperature changes are usually in °C, but check the question carefully. ΔT in °C equals ΔT in K.
Exam technique for "Physical Chemistry"
Command words matter: "State" requires a brief answer with no explanation (1 mark). "Explain" requires reasoning or mechanism (usually 2-3 marks). "Calculate" requires working shown for method marks.
Show all working in calculations: Even if your final answer is wrong, you can gain method marks. Write the formula, substitute values, then calculate. Include units in your final answer.
Use data from the question: Many questions provide ΔH values, specific heat capacity, or other data. Underlining key numbers helps you spot what to use.
For equilibrium questions, always reference the specific equation given: Don't make general statements. Say "the forward reaction is exothermic" rather than just "it's exothermic."
Quick revision summary
Physical Chemistry covers the particle model, energy changes, reaction rates, and equilibrium. Exothermic reactions release energy (ΔH negative); endothermic reactions absorb energy (ΔH positive). Rate of reaction increases with higher temperature, concentration, pressure, surface area, or by adding a catalyst. Reversible reactions can reach dynamic equilibrium in closed systems. Le Chatelier's Principle predicts how equilibrium position shifts when conditions change. Industrial processes like the Haber Process use compromise conditions to balance yield, rate, and cost.