This mini-lesson walks you through the whole of Edexcel iGCSE 4CH1 Section 3 — Physical chemistry: energetics (exo/endothermic, calorimetry, Q = mcΔT and ΔH, bond energies), rates of reaction (collision theory, catalysts, the marble-chip & hydrogen peroxide practicals) and reversible reactions & dynamic equilibrium.
Make a guess: why does a match need a strike to start burning?
Mission 1Why does a burning match need a push?
Mission 2What decides how fast a reaction goes?
Mission 3Can a reaction run backwards?
Work through each screen, answer the questions as you go (some are wordy, some are calculations) and collect ⭐ stars. Press Start when you're ready.
Energetics
Exothermic and endothermic
Every reaction involves an energy change with the surroundings:
An exothermic reaction gives out heat energy — the surroundings get warmer (e.g. combustion, neutralisation, most displacement reactions).
An endothermic reaction takes in heat energy — the surroundings get colder (e.g. some salts dissolving, thermal decomposition).
Feel the test tube: warmer = exothermic, colder = endothermic.
Watch out: exothermic means energy is released to the surroundings, so a thermometer in the mixture reads a rise. "Exo gives, endo takes."
Quick check
Warmer or colder?
?A student adds ammonium nitrate to water and the temperature of the solution falls from 21 °C to 14 °C. What does this tell us?
Energetics · practical
Calorimetry — measuring the heat change
In a simple calorimetry experiment we measure the temperature change of a known mass of water (or solution) and calculate the heat energy change:
Q = m c ΔTQ = heat energy (J) = mass of water (g) × specific heat capacity (4.2 J/g°C) × temperature change (°C)
A polystyrene cup + lid reduces heat loss for combustion, displacement, dissolving and neutralisation reactions.
Once you have Q (in J), divide by the number of moles that reacted to get the molar enthalpy change ΔH (and add a sign: − for exothermic, + for endothermic).
Worked example
50 g of water rose by 12 °C in a neutralisation. c = 4.2 J/g°C.
Q = m c ΔT = 50 × 4.2 × 12 = 2520 J = 2.52 kJ
If 0.025 mol of acid reacted: ΔH = −2520 ÷ 0.025 = −100 800 J/mol ≈ −100.8 kJ/mol (exothermic).
Sign rule: exothermic ΔH is negative (energy leaves); endothermic ΔH is positive.
Calculate
Your turn — heat energy change
1In a calorimetry experiment, 100 g of water rose in temperature by 8.0 °C. Using c = 4.2 J/g°C, calculate the heat energy change Q.
J
Hint: Q = m c ΔT = 100 × 4.2 × 8.0.
Energetics · energy level diagrams
Reaction profile diagrams
An energy level (reaction profile) diagram shows the energy of reactants and products, the activation energy (Ea — the minimum energy to start the reaction) and the overall ΔH.
Exothermic: products below reactants, ΔH negative. Endothermic: products above, ΔH positive. The "hump" is the activation energy Ea.Quick check
Reading the profile
?On a reaction profile, the products sit higher than the reactants. What can you conclude?
Mission 1 check
Why does a burning match need a push?
?A match head burns in an exothermic reaction. Yet a match left in its box never lights by itself. Why must it be struck first?
Energetics · bond energies
Bond breaking vs bond making
Energy changes come from bonds:
Breaking bonds takes in energy → endothermic.
Making bonds gives out energy → exothermic.
ΔH = (bonds broken) − (bonds made)= total energy IN to break reactant bonds − total energy OUT when product bonds form
ΔH = 678 − 862 = −184 kJ/mol → exothermic (more energy out making bonds than in breaking them).
Watch out: never reverse it — it is broken minus made. If "made" is bigger, ΔH is negative (exothermic).
Calculate
Your turn — bond energy ΔH
2For H₂ + Br₂ → 2HBr, the bond energies (kJ/mol) are: H–H = 436, Br–Br = 193, H–Br = 366. Calculate ΔH for the reaction.
kJ/mol
Hint: broken = 436 + 193 = 629; made = 2 × 366 = 732; ΔH = 629 − 732. Don't forget the minus sign.
Mission 2 of 3
What decides how fast a reaction goes?
Make a guess: what decides how fast a reaction goes?
Rates of reaction
Collision theory
For particles to react they must collide with at least the activation energy Ea. The rate depends on how frequent and how energetic these collisions are.
The factors that speed up a reaction (4CH1):
Surface area of a solid ↑ → more particles exposed → more collisions.
Concentration of a solution ↑ → particles closer → more collisions.
Pressure of a gas ↑ → particles closer → more collisions.
Temperature ↑ → particles move faster AND with more energy → more frequent AND more successful collisions.
Catalyst → lowers Ea, so more collisions succeed.
Watch out: raising the temperature does two things — collisions become both more frequent and more energetic (a bigger fraction reach Ea). Don't only mention speed.
Sort it
Why does it speed up?
Tap the main reason each change increases the rate (collision theory).
Rates · catalysts
Catalysts lower the activation energy
A catalyst increases the rate of a reaction but is chemically unchanged at the end. It works by providing an alternative pathway with a lower activation energy, so more collisions succeed.
The catalyst lowers the energy "hump" (Ea) but leaves the start and end energies — and so ΔH — unchanged.
Practical (3.16): different solids (e.g. manganese(IV) oxide) catalyse the decomposition of hydrogen peroxide — the catalyst is recovered unchanged at the end.
Quick check
How a catalyst works
?Which statement about a catalyst is correct?
Rates · measuring & graphs
Measuring rate & reading graphs
Rate is measured by following how fast a product forms or a reactant is used up — e.g. volume of gas collected, mass lost, or the time for a mark to disappear.
mean rate = quantity ÷ timee.g. mean rate = volume of gas (cm³) ÷ time (s)
The steeper the curve, the faster the rate. Both level off (plateau) when the reaction finishes; the same amount of the limiting reactant (e.g. the same mass of marble chips with the acid in excess) gives the same final volume.
Practical (3.15): marble chips + dilute hydrochloric acid give off carbon dioxide. Measure the gas volume (or the mass lost) at regular times, changing the chip size or the acid concentration — a steeper curve means a faster rate. Smaller chips or more concentrated acid speed it up.
Calculate
Your turn — mean rate
3In the marble chips + hydrochloric acid practical, 48 cm³ of carbon dioxide is collected in 120 s. Calculate the mean rate of reaction.
cm³/s
Hint: mean rate = volume ÷ time = 48 ÷ 120.
Mission 2 check
What decides how fast a reaction goes?
?Zinc reacts with dilute sulfuric acid. A student swaps zinc lumps for the same mass of zinc powder, at the same temperature. Hydrogen now bubbles off much faster. What is the best explanation?
Mission 3 of 3
Can a reaction run backwards?
Make a guess: does a reversible reaction ever finish?
Reversible reactions & equilibria
Reversible reactions ⇌
Some reactions are reversible — products can react to re-form reactants. We show this with the special symbol ⇌.
Heating blue hydrated copper(II) sulfate drives off water to give white anhydrous CuSO₄; adding water reverses it (and warms — exothermic).
Heating ammonium chloride: NH₄Cl ⇌ NH₃ + HCl (it decomposes, then re-forms on cooling).
The two half-arrows of ⇌ show both directions happening.Dynamic equilibrium
Dynamic equilibrium
In a sealed (closed) container, a reversible reaction reaches dynamic equilibrium. At equilibrium:
the forward and reverse reactions occur at the same rate;
the concentrations of reactants and products stay constant.
Dynamic, not stopped — both reactions keep going, but at equal rates, so amounts stay constant.
Watch out: equilibrium does not mean the reaction has stopped, and it does not mean amounts of reactant and product are equal — only that they are no longer changing.
Match it
Equilibrium facts
Tap a term on the left, then its correct match on the right.
Changing the conditions
Temperature & pressure shift equilibrium
If you change the conditions, the position of equilibrium moves to oppose the change (this idea is sometimes called Le Chatelier's principle — not named in 4CH1, but it explains the rules):
Temperature: an increase (decrease) shifts equilibrium in the endothermic (exothermic) direction.
Pressure: an increase (decrease) shifts equilibrium toward the side with fewer (more) moles of gas.
Catalyst: does not change the position — it speeds forward and reverse equally, so equilibrium is just reached faster.
Worked example — the Haber process N₂ + 3H₂ ⇌ 2NH₃ (forward exothermic; 4 mol gas → 2 mol gas). High pressure (~200 atm) shifts right (fewer moles → more NH₃). A lower temperature would also shift right, but the compromise ~450 °C is used to keep the rate fast enough. An iron catalyst speeds it up without changing the position.
Predict
Your turn — shift the equilibrium
4For 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), the forward reaction is exothermic. The pressure is increased. Which way does the position of equilibrium move?
Mission 3 check
Can a reaction run backwards?
?Brown nitrogen dioxide is sealed in a tube: 2NO₂ ⇌ N₂O₄ (N₂O₄ is colourless). The brown colour fades a little, then stays the same for days. Why does it stop changing?