OCR A-Level Chemistry: Lattice Enthalpy, Entropy and Electrode Potentials — Complete Revision Guide (H432)
OCR A-Level Chemistry: Lattice Enthalpy, Entropy and Electrode Potentials
Lattice enthalpy, entropy and electrode potentials is the A2 extension of the energetics framework introduced in Enthalpy, Rates and Equilibrium. At AS, enthalpy was treated through Hess cycles and bond enthalpies, and "spontaneity" was a loose qualitative notion. At A2, the second law of thermodynamics enters explicitly: entropy as the system's measure of disorder, Gibbs free energy ΔG = ΔH - TΔS as the rigorous criterion for spontaneity, and electrode potentials as the experimental quantification of the driving force behind redox reactions. The Born-Haber cycle and the standard electrochemical-cell construction are routine items on Paper 1 (Periodic Table, Elements and Physical Chemistry) at A2.
H432 examiners weight this module heavily because it is the first place in the spec where candidates can quantify feasibility — and the first place where the limits of feasibility ("ΔG predicts whether, kinetics predicts when") become explicit. A single Paper 1 question on Module 5.2 can demand the construction of a Born-Haber cycle for an unfamiliar binary salt, the calculation of ΔG at a non-standard temperature, the EMF prediction for a redox pairing taken from a supplied data sheet, and the qualitative justification of why the predicted reaction is observed to be slow at room temperature. Candidates who internalise the thermodynamic framework as a hierarchy — lattice enthalpy explains ionic-solid energetics, Gibbs explains direction, electrode potentials explain quantitative drive — find these questions short. Candidates who memorise each calculation type separately struggle on the integrated items where two or three sub-skills are chained.
Course 9 of the H432 Chemistry learning path on LearningBro, Lattice Enthalpy, Entropy and Electrode Potentials, develops the thermodynamic and electrochemical machinery that underwrites every quantitative argument in A2 inorganic chemistry. It builds in three phases: lattice enthalpy with the Born-Haber cycle and the hydration/solution thermochemistry of dissolution; entropy and Gibbs free energy as the spontaneity framework; and standard electrode potentials, electrochemical cells, the EMF calculation and redox titrations. It sits adjacent to Acids, Bases and Buffers and feeds into Transition Elements and Aromatic on the OCR A-Level Chemistry learning path.
Guide Overview
The Lattice Enthalpy, Entropy and Electrode Potentials course is built as a sequence of lessons that move from lattice enthalpy through entropy and Gibbs into electrochemistry.
- Lattice Enthalpy Definition
- Born-Haber Cycle Construction
- Factors Affecting Lattice Enthalpy
- Hydration Enthalpy and Enthalpy of Solution
- Entropy and the Second Law
- Gibbs Free Energy and Spontaneity
- Standard Electrode Potential
- Electrochemical Cells and EMF
- Predicting Redox Feasibility
- Redox Titrations: Manganate and Iodine-Thiosulfate
OCR H432 Specification Coverage
This course addresses OCR H432 Module 5.2.1 (lattice enthalpy), Module 5.2.2 (entropy and Gibbs free energy) and Module 5.2.3 (redox and electrode potentials). The specification organises the topic into the thermochemistry of ionic crystals, the spontaneity criterion combining enthalpy and entropy, and the standard-electrode-potential framework for predicting and quantifying redox reactions (refer to the official OCR specification document for exact wording).
| Sub-topic | Spec area | Primary lesson(s) |
|---|---|---|
| Lattice enthalpy; Born-Haber cycle | OCR H432 Module 5.2.1 | Lattice Enthalpy Definition; Born-Haber Cycle Construction; Factors Affecting Lattice Enthalpy |
| Enthalpy of hydration and solution | OCR H432 Module 5.2.1 | Hydration Enthalpy and Enthalpy of Solution |
| Entropy; second law; ΔS calculations | OCR H432 Module 5.2.2 | Entropy and the Second Law |
| Gibbs free energy; ΔG = ΔH - TΔS | OCR H432 Module 5.2.2 | Gibbs Free Energy and Spontaneity |
| Standard electrode potential; SHE reference | OCR H432 Module 5.2.3 | Standard Electrode Potential |
| Cell EMF; cell notation | OCR H432 Module 5.2.3 | Electrochemical Cells and EMF |
| Predicting reaction feasibility from E° values | OCR H432 Module 5.2.3 | Predicting Redox Feasibility |
| Redox titrations (manganate, iodine-thiosulfate) | OCR H432 Module 5.2.3 | Redox Titrations: Manganate and Iodine-Thiosulfate |
Module 5.2 is examined on Paper 1 and recurs on Paper 3, particularly Born-Haber-cycle calculations, ΔG-based feasibility judgements at specified temperatures, and EMF-based feasibility predictions in unfamiliar redox systems.
Topic-by-Topic Walkthrough
Lattice Enthalpy and the Born-Haber Cycle
The lattice enthalpy lesson defines lattice enthalpy as the enthalpy change when one mole of an ionic solid is formed from its constituent gaseous ions — exothermic by convention. The Born-Haber cycle lesson develops the standard five-step indirect route to lattice enthalpy: atomisation of the metal, atomisation of the non-metal, ionisation of the metal (sum of successive IEs), electron affinity of the non-metal (sum if multiple negative charges, with second EA endothermic because of electron-electron repulsion), and the enthalpy of formation of the ionic solid from its elements. Hess's law then closes the cycle. The classic worked example is NaCl: atomisation of Na (+108 kJ mol⁻¹) + ionisation Na (+496) + atomisation Cl (+122) + electron affinity Cl (-349) + lattice enthalpy NaCl(s) = formation ΔH NaCl (-411), giving lattice enthalpy = -788 kJ mol⁻¹.
The factors affecting lattice enthalpy lesson develops the inverse-distance and product-of-charges scaling: lattice enthalpy is more exothermic for smaller ions (which can approach closer) and for ions with higher charges. The progression LiF (-1037), NaF (-918), KF (-816) is the inverse-distance trend; the progression NaF (-918), MgO (-3791) is the charge-product trend (×4 for Mg²⁺O²⁻ versus ×1 for Na⁺F⁻, partly offset by larger radii). The polarisation-of-anion concept (a small highly charged cation distorts the electron cloud of a large anion, giving covalent character) explains why some compounds deviate from purely ionic models — and is the conceptual bridge to the d-block covalent character of transition elements.
Hydration, Solution and the Dissolution Cycle
The hydration enthalpy lesson develops the dissolution-cycle competition. Lattice enthalpy is the energy needed to break the ionic lattice into gaseous ions (endothermic when reversed); hydration enthalpy is the energy released when those ions are surrounded by water molecules. The sum gives the enthalpy of solution: ΔH_solution = -ΔH_lattice + ΔH_hydration(cation) + ΔH_hydration(anion). A salt with strongly exothermic hydration relative to lattice enthalpy dissolves with overall exothermic ΔH_solution; one with a strongly negative lattice enthalpy (small, highly charged ions) tends to be insoluble. The hydration enthalpy itself scales inversely with ionic radius and with charge — Li⁺ is the most strongly hydrated Group 1 cation, F⁻ the most strongly hydrated Group 17 anion.
Entropy and Gibbs Free Energy
The entropy lesson develops entropy as the system's measure of disorder, with the second law stating that the entropy of the universe always increases in a spontaneous process. Gaseous states have higher entropy than liquids than solids; more complex molecules have higher entropy than simpler ones; mixed states have higher entropy than separated. ΔS_system = Σ S°(products) - Σ S°(reactants); the canonical sign-prediction rules are an increase in gas moles → positive ΔS, dissolution of a solid → typically positive ΔS, condensation or freezing → negative ΔS.
The Gibbs lesson develops ΔG = ΔH - TΔS as the spontaneity criterion: ΔG < 0 means thermodynamically feasible, ΔG > 0 means not feasible, ΔG = 0 means at equilibrium. The four combinations of ΔH and ΔS give: exothermic with positive ΔS always feasible (e.g. combustion); exothermic with negative ΔS feasible only at low T; endothermic with positive ΔS feasible only at high T (e.g. NaHCO₃ thermal decomposition); endothermic with negative ΔS never feasible. The temperature at which ΔG changes sign is T = ΔH/ΔS. This explains why ice melts above 273 K (endothermic, ΔS > 0, so feasible at higher T) and why water freezes below 273 K (exothermic, ΔS < 0, so feasible at lower T). The exam-relevant subtlety is that ΔG predicts feasibility but says nothing about rate — a feasible reaction with high activation energy may not proceed at observable speed (diamond → graphite at room temperature is the canonical case).
Electrode Potentials and Electrochemical Cells
The standard electrode potential lesson develops electrode potentials as the EMF measured when a half-cell is connected to a standard hydrogen electrode (SHE, defined as 0.00 V) under standard conditions: 1 mol dm⁻³ ion concentration, 100 kPa gas pressure, 298 K, platinum inert electrode. Each half-cell is written as a reduction (e.g. Zn²⁺ + 2e⁻ → Zn, E° = -0.76 V; Cu²⁺ + 2e⁻ → Cu, E° = +0.34 V). The more positive the E°, the greater the tendency to be reduced (i.e. the stronger oxidising agent).
The electrochemical cells and EMF lesson develops cell construction with two half-cells joined by a salt bridge and a voltmeter, and the standard cell notation Zn(s) | Zn²⁺(aq) || Cu²⁺(aq) | Cu(s) where || is the salt bridge. EMF = E°(cathode) - E°(anode) = E°(positive electrode) - E°(negative electrode), where the more-positive E° half-cell is the cathode and undergoes reduction. The Daniell cell gives EMF = 0.34 - (-0.76) = 1.10 V. The predicting redox feasibility lesson develops the rule: a redox reaction is feasible if the resulting cell EMF is positive, which corresponds to ΔG = -nFE° being negative.
Redox Titrations
The redox titrations lesson develops the two canonical titrations on the H432 spec. Manganate(VII)-iron(II) uses MnO₄⁻/H⁺ as the oxidant, self-indicating because the purple MnO₄⁻ decolourises in the conical flask until the iron(II) is exhausted, then the first excess drop tints the solution purple. The stoichiometry from the combined half-equations is MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺, i.e. 1:5. The iodine-thiosulfate titration uses I₂/S₂O₃²⁻ with starch indicator added near the endpoint (a blue-black iodine-starch complex disappears at endpoint), and the stoichiometry is I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻, i.e. 1:2. Both procedures are anchors for PAG 5 (Redox titration) and recur in transition elements.
A Typical H432 Paper 1 Question
A standard Paper 1 prompt gives candidates a Born-Haber cycle with one quantity missing — typically the second electron affinity of an oxide ion (endothermic because of electron-electron repulsion) or the lattice enthalpy of an unfamiliar fluoride or oxide. The route is fixed: complete the cycle by ensuring each arrow direction is labelled correctly, identify the missing arrow as the unknown, apply Hess's law (the sum of one direction equals the sum of the other), and solve for the unknown. The discriminator at the top band is the explicit sign treatment — second electron affinities are positive (endothermic), atomisation enthalpies are positive, ionisation enthalpies are positive, electron affinities for the first electron are negative, lattice formation enthalpies are negative — and the explicit acknowledgement that a "lattice enthalpy" given as a positive number in some textbooks may be the lattice dissociation enthalpy rather than the lattice formation enthalpy. Reading the question's definition carefully is itself a mark-bearing step.
Synoptic Links
Energetics and electrochemistry thread through the rest of A2 chemistry. The lattice and hydration enthalpy framework explains solubility trends across the Group 2 sulfates and hydroxides revisited from periodicity, Group 2 and halogens. The Gibbs free energy framework connects to the equilibrium constant via ΔG° = -RT ln K from quantitative rates and equilibrium. Electrode potentials underwrite the redox chemistry of transition elements and aromatic, where the manganate-iron(II) titration anchors the assessment of vanadium and chromium oxidation states. The thermodynamic-vs-kinetic distinction (feasible-not-fast) is revisited as the rationale for why catalysts speed reactions without changing K.
Paper 3 'Unified chemistry' items deploy this module in two characteristic ways. The first is the battery-or-fuel-cell scenario: candidates are given a primary, secondary or fuel-cell electrochemistry context (Li-ion, hydrogen fuel cell, alkaline cell, methanol fuel cell) and asked to write the two half-equations, the overall cell reaction, the standard EMF from a supplied data table, and the consequences of operating away from standard conditions. The second is the metal-extraction scenario: candidates are given a thermodynamic argument for why a particular metal cannot be extracted by reduction with carbon below a certain temperature (because ΔG of the carbon-oxygen reaction is not negative enough) and asked to identify the crossover temperature on an Ellingham-style diagram, which is a direct application of ΔG = ΔH - TΔS. The discriminating moves at the top band are explicit unit treatments (kJ vs J for ΔH, K vs °C for T, and the conversion from J K⁻¹ mol⁻¹ for ΔS) and the explicit reading of E° signs from the supplied data sheet.
What Examiners Reward
Top-band marks on this module cluster around sign discipline and explicit definitions. For Born-Haber cycles, examiners want every arrow labelled with its named enthalpy quantity (atomisation, first ionisation, second ionisation, first electron affinity, second electron affinity, lattice formation, formation), every arrow direction correct, and every sign explicit. For Gibbs calculations, they want ΔH in kJ mol⁻¹ converted to J mol⁻¹ before adding to TΔS (because ΔS is in J K⁻¹ mol⁻¹), and the explicit comment that ΔG = 0 corresponds to the crossover temperature T = ΔH/ΔS rather than to equilibrium concentration. For electrode-potential calculations, they want the half-equations written as reductions in both rows (the convention) and EMF = E°(cathode) - E°(anode) with explicit identification of which is which. For redox-titration arithmetic, they want the combined balanced equation drawn before any mole arithmetic, and the explicit mole-ratio reading off the balanced equation.
Common pitfalls cluster around six recurring mistakes. First, sign errors in the second electron affinity for oxide (O → O⁻ is exothermic, O⁻ → O²⁻ is endothermic because of electron-electron repulsion in the smaller ion). Second, computing TΔS in J mol⁻¹ but ΔH in kJ mol⁻¹ and forgetting to convert before subtracting. Third, predicting feasibility from EMF and then claiming the reaction will proceed at observable speed (feasibility says nothing about rate). Fourth, misreading the cell notation, taking the left-hand half-cell as the cathode when by convention the right-hand half-cell (the more-positive E°) is the cathode. Fifth, writing the manganate-iron(II) titration with a 1:1 mole ratio rather than the correct 1:5. Sixth, applying standard electrode potentials to non-standard conditions without flagging that the actual EMF will differ via the Nernst equation. Each is a one- or two-mark deduction that compounds quickly across a multi-step thermodynamic-electrochemical problem.
Practical Activity Groups (PAGs)
This course anchors PAG 5 (Redox titration) in full through the manganate-iron(II) and iodine-thiosulfate titrations. The course also anchors elements of PAG 11 (pH measurement) through the cell-EMF measurements that use a high-impedance voltmeter, and connects synoptically to the buffer and titration work of acids, bases and buffers. The Born-Haber cycle and ΔG calculations are written exercises rather than lab practicals but commonly appear as evaluative items applied to lab data.
Going Further
Undergraduate analogues of this material extend in several directions. First, the lattice enthalpy framework generalises into the Madelung constant and into the Kapustinskii equation for estimating lattice energies without a full crystal-structure calculation. Second, the Gibbs free energy framework generalises into the chemical potential and into the thermodynamics of solutions, electrolytes and reactions away from standard state. Third, electrode potentials generalise into the Nernst equation E = E° - (RT/nF) ln Q, which gives the cell potential at non-standard concentrations and is the foundation of pH electrodes, ion-selective electrodes and biological electrochemistry (nerve action potentials, ATP synthase). Oxbridge-style interview prompts on this material include: "Why does MgO have a lattice enthalpy approximately four times that of NaF?" "Use the Gibbs framework to predict whether NaHCO₃ thermal decomposition is feasible at room temperature, and at what temperature it becomes feasible." "Why is a feasible redox reaction (negative ΔG, positive EMF) not necessarily a fast one — and what does this say about how electrochemical cells are designed?"
Authorship and Sign-off
This guide was authored independently by John Haigh, paraphrasing OCR H432 Modules 5.2.1, 5.2.2 and 5.2.3 as descriptive use. No verbatim spec text, mark-scheme phrasing, examiner-report quotation, or past-paper question reference appears. The worked examples are original.
Start at the Lattice Enthalpy, Entropy and Electrode Potentials course and work through every lesson in sequence. Once the Born-Haber cycle, the Gibbs free energy criterion and the electrode-potential framework are automatic, the rest of A2 inorganic chemistry resolves into substitution-and-rearrangement against a known thermodynamic backbone — and the feasibility questions become recognition rather than recall.