EntropyAQA A-Level Chemistry: Revision notes
Section 1
Why enthalpy change is not enough
Many exothermic changes happen readily, which suggests that energy is released when a change is favourable. But the enthalpy change, ΔH, is not sufficient to explain whether a change is feasible (able to happen without continuing input of energy).
Some endothermic changes occur readily at room temperature: ammonium nitrate dissolving in water, a solid melting, or a liquid evaporating. Some exothermic changes do not proceed at all. A second quantity is needed: entropy.
Section 2
Entropy as disorder
Entropy (S) is a measure of the disorder of a system: the number of ways the particles and their energy can be arranged. The greater the number of arrangements, the greater the entropy.
A positive entropy change (ΔS) means the system becomes more disordered. The units are J K⁻¹ mol⁻¹.
All substances have a positive absolute entropy value. A perfect crystal at 0 K has the lowest possible entropy, so absolute entropy values can be tabulated at 298 K as standard entropies (S°).
Entropy is measured in J K⁻¹ mol⁻¹ but enthalpy is in kJ mol⁻¹. Convert one to match the other before combining them.
Section 3
Entropy and physical changes
Entropy depends on the state of a substance. Particles in a solid are held in fixed positions and are the most ordered. A liquid is more disordered, and a gas, whose particles move randomly and are far apart, has by far the highest entropy: S(gas) ≫ S(liquid) > S(solid).
- Melting and boiling increase entropy, so ΔS is positive.
- Freezing and condensing decrease entropy, so ΔS is negative.
- Dissolving a solid usually increases entropy as the ordered lattice breaks into mobile ions.
- Raising the temperature raises entropy, because the particles have more ways of sharing out the energy.
Section 4
Entropy and chemical changes
In a chemical reaction the entropy change is mostly determined by changes in the number of moles of gas.
- More gas particles in the products than in the reactants: ΔS is positive, e.g. CaCO₃(s) → CaO(s) + CO₂(g).
- Fewer gas particles in the products: ΔS is negative, e.g. N₂(g) + 3H₂(g) → 2NH₃(g).
- No change in the number of gas particles: ΔS is small, positive or negative.
A gas produced from solids or solutions gives a large positive ΔS. More particles overall in the same state also gives a positive ΔS.
To predict the sign of ΔS, count moles of gas on each side of the equation first, then consider solids and liquids.
Section 5
Calculating entropy changes
The entropy change of a reaction is found from absolute entropy values:
ΔS = ΣS(products) − ΣS(reactants)
Multiply each value by its coefficient in the balanced equation.
Worked example. C(graphite) + O₂(g) → CO₂(g). S = 5.7, 205 and 214 J K⁻¹ mol⁻¹. ΔS = 214 − (5.7 + 205) = +3.3 J K⁻¹ mol⁻¹
The value is small because the number of moles of gas does not change. Always give a sign and the units J K⁻¹ mol⁻¹.
Forgetting the coefficients, or subtracting products from reactants. Write out ΣS(products) and ΣS(reactants) separately first.
That's the notes covered.
Carry on to the next subtopic.
Exam questions on Entropy
- A student discusses entropy using the three physical states of water. She considers how the disorder of the particles in ice, liquid water and steam changes as a sample of water is heated from below 0 °C to above 100 °C at constant pressure.Explain why the entropy of steam is greater than the entropy of liquid water at the same temperature.2 marks
- Ammonia is made industrially in the Haber process: N₂(g) + 3H₂(g) → 2NH₃(g). Standard entropies (J K⁻¹ mol⁻¹): N₂(g) = 192; H₂(g) = 131; NH₃(g) = 193.Predict, with a reason, the sign of the entropy change for the reaction Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g).2 marks
- Calcium carbonate decomposes on strong heating: CaCO₃(s) → CaO(s) + CO₂(g), ΔH = +178 kJ mol⁻¹. Standard entropies (J K⁻¹ mol⁻¹): CaCO₃(s) = 93; CaO(s) = 40; CO₂(g) = 214. Relative formula mass of CaCO₃ = 100.1.Calculate the entropy change for the decomposition of calcium carbonate. Include units.3 marks
Written by the Exaim team, led by Shaun Daswani (Head of Upper Secondary, Improve ME Institute; MSc Financial Mathematics, Imperial College London; BSc, UCL) and Jason Daswani (operational lead, Improve ME Institute; LSE).