Edexcel International A Level Chemistry

Topic 12: Entropy and Energetics

Two questions about energy that enthalpy alone cannot answer. Why do endothermic changes happen, and how strong is the bonding inside an ionic solid? The first is answered by entropy, the second by lattice energy.

The one idea

Enthalpy tells you how much heat a reaction gives out or takes in. It does not tell you whether the reaction will happen. Some endothermic changes occur spontaneously (a cold pack, or ammonium nitrate dissolving), and some exothermic ones never start.

The missing quantity is entropy, a measure of how widely particles and energy are dispersed. A change is feasible when the total entropy of the system and its surroundings increases. Heat given out raises the entropy of the surroundings, which is how enthalpy re-enters the rule.

The second half of the topic turns the same energy thinking to ionic solids. A Born-Haber cycle is Hess's law applied to forming an ionic lattice, and it gives a lattice energy you cannot measure directly. Comparing that with theory shows covalent character. Combining the lattice energy with hydration gives an enthalpy of solution, and with entropy the rule for solubility.

Route through the topic. The total-entropy rule from block 12A is used again for solubility at the end of block 12B.

The five jobs

Job	Pages
Explain what entropy is, and say how it changes (solid, liquid, gas; moles of gas; dissolving)	Enthalpy is not enough to Why entropy changes
Calculate $\Delta S_\text{system}$ , $\Delta S_\text{surroundings}$ and $\Delta S_\text{total}$	Entropy change of the system and Surroundings and total entropy
Decide whether a reaction is feasible, find the temperature at which it becomes feasible, and separate this from speed	Balance and temperature to Thermodynamic and kinetic stability
Define the enthalpy changes for ionic solids, then build and use a Born-Haber cycle	Three definitions to Born-Haber calculations
Interpret lattice energies, calculate solution enthalpies, and predict solubility	Covalent character to Group 2 solubility

What you need first

Enthalpy changes and Hess's law from the energetics topic: you will use $\Delta_\text{f}H$ in cycles, and the idea that any route between the same two states gives the same total.
Ionisation energy and the shapes of ions from the atomic structure and bonding topics.
Group 2 solubility trends from Unit 2, which this topic explains.
Activation energy from kinetics, for the difference between feasible and fast.

Values used in these notes

Quantity	Value
gas constant, R	8.31 J mol⁻¹ K⁻¹ (only needed if you meet $\Delta G$ )
temperature	K = °C + 273
entropy	J K⁻¹ mol⁻¹
enthalpy	kJ mol⁻¹ (convert to J before dividing by T)

The standard entropy values used in the worked examples are given in the question, as they would be in an exam.

How to use these notes

Read the pages in order the first time. Later pages assume the earlier ones.
Each page opens with its one idea in a sentence. If you can explain that sentence to someone else, you know the page.
Coloured boxes mean the same thing on every page: definitions you may be asked to quote, key ideas, common mistakes, exam technique, and beyond the spec (interesting, but you will not be examined on it).
Each page ends with questions and fully worked answers. Cover the answers, try the questions, then mark yourself.
Use the topic on a page the night before an exam.

Spec checklist

Every point you can be examined on in this topic, and where it is explained. Tick them off as you go.

Spec	You should be able to	Where
12.1	understand that, since endothermic reactions can occur spontaneously at room temperature, enthalpy changes alone do not control whether reactions occur	Endothermic, yet it happens
12.2	understand entropy as a measure of disorder of a system in terms of the random dispersal of molecules and of energy quanta between molecules	Energy spreads out too Entropy The natural direction of change
12.3	understand that the entropy of a substance increases with temperature, that entropy increases as solid → liquid → gas and that perfect crystals at zero kelvin have zero entropy	The zero point Entropy rises with temperature Solid, liquid, gas
12.4	be able to interpret the natural direction of change as being in the direction of increasing total entropy (positive entropy change), including gases spread spontaneously through a room	A gas spreads out The natural direction of change Total entropy decides the direction
12.5(i)	understand why entropy changes occur during: changes of state	Changes of state
12.5(ii)	understand why entropy changes occur during: dissolving of a solid ionic lattice	Dissolving an ionic solid
12.5(iii)	understand why entropy changes occur during: reactions in which there is a change in the number of moles from reactants to products	Reactions with a change in moles
12.6	understand that the total entropy change of any reaction is the sum of the entropy change of the system and the entropy change of the surroundings, summarised by the expression: ΔS_total = ΔS_system + ΔS_surroundings	System and surroundings Total entropy decides the direction
12.7	be able to calculate the entropy change of the system for a reaction, ΔS_system, given the entropies of the reactants and products	The calculation
12.8	be able to calculate the entropy change in the surroundings, and hence ΔS_total, using the expression ΔS_surroundings = −ΔH / T	Heat and the surroundings
12.9(i)	understand that the feasibility of a reaction depends on: the balance between ΔS_system and ΔS_surroundings, so that even endothermic reactions can occur spontaneously at room temperature	The balance
12.9(ii)	understand that the feasibility of a reaction depends on: temperature, as higher temperatures decrease the magnitude of ΔS_surroundings so its contribution to ΔS_total is less. Students should be able to calculate the temperature at which a reaction is feasible. Students may also use ΔG = ΔH − TΔS_system in answers, although this approach is not a requirement of the specification.	Temperature changes the balance Where the total is zero Using ΔG
12.10	understand that reactions can occur as long as ΔS_total is positive even if one of the other entropy changes is negative	One negative term is fine
12.11	understand and distinguish between the concepts of thermodynamic stability and kinetic stability	Two kinds of stable
12.12(i)	be able to define the terms: standard enthalpy change of atomisation, Δ_atH	Enthalpy change of atomisation
12.12(ii)	be able to define the terms: electron affinity	Electron affinity
12.12(iii)	be able to define the terms: lattice energy (as the exothermic process for the formation of one mole of an ionic solid from its gaseous ions)	Lattice energy
12.13	be able to construct Born-Haber cycles and carry out related calculations	Two routes, one total The steps, in order Lattice energy from the cycle Magnesium chloride: multiples and a second step Magnesium oxide: an endothermic second electron affinity Working backwards
12.14	understand that a comparison of the experimental lattice energy value (from a Born-Haber cycle) with the theoretical value (obtained from electrostatic theory) in a particular compound indicates the degree of covalent bonding	Experimental and theoretical values
12.15	understand that polarisation of anions by cations leads to some covalency in an ionic bond, based on evidence from the Born-Haber cycle	Polarisation of the anion
12.16	be able to define the terms 'enthalpy change of solution, Δ_solH' and 'enthalpy change of hydration, Δ_hydH of an ion'	Two definitions
12.17	be able to use energy cycles and energy level diagrams to calculate the enthalpy change of solution of an ionic compound, using enthalpy change of hydration and lattice energy	The solution cycle
12.18	understand the effect of ionic charge and ionic radius on the values of enthalpy change of hydration and the lattice energy of an ionic compound	The data The explanation
12.19	be able to use entropy and enthalpy changes of solution values to predict the solubility of ionic compounds and discuss trends in the solubility of ionic compounds covered in Unit 2	The same test as before Temperature and the four cases The trends from Unit 2 Which term falls faster?