AQA A-Level Chemistry: Physical Chemistry Revision Guide -- Energetics, Kinetics, Equilibria, and Thermodynamics

Physical Chemistry is the mathematical and theoretical backbone of A-Level Chemistry. For many students, it is simultaneously the most rewarding and the most demanding part of the course -- rewarding because calculations produce definite, checkable answers, and demanding because the concepts build on one another in layers that require careful understanding.

On the AQA specification, Physical Chemistry spans energetics, kinetics, chemical equilibria, thermodynamics, and electrode potentials. Together, these topics answer fundamental questions -- how much energy does a reaction release, how fast does it go, how far does it go, and can we predict whether it will happen at all?

This guide works through each topic, covering the key specification content, essential calculations, and the conceptual understanding that examiners reward.

Energetics: Enthalpy Changes and Hess's Law

Energetics at AS level introduces the idea that chemical reactions involve energy changes, measured as enthalpy changes under constant pressure. At A-Level, this foundation is extended significantly.

Key Enthalpy Definitions

You must know precise definitions for a range of standard enthalpy changes. Examiners are strict about wording -- an imprecise definition will lose marks even if your understanding is sound.

• Standard enthalpy of formation -- the enthalpy change when one mole of a compound is formed from its elements in their standard states, under standard conditions (298 K, 100 kPa).
• Standard enthalpy of combustion -- the enthalpy change when one mole of a substance undergoes complete combustion under standard conditions.
• Standard enthalpy of atomisation -- the enthalpy change when one mole of gaseous atoms is formed from an element in its standard state.
• First ionisation energy -- the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous unipositive ions.
• First electron affinity -- the enthalpy change when one mole of gaseous atoms each gain one electron to form one mole of gaseous uninegative ions.
• Lattice enthalpy of formation -- the enthalpy change when one mole of an ionic compound is formed from its gaseous ions.

Revision tip: Write out each definition from memory, then compare it with the specification wording. Even small omissions -- forgetting to say "gaseous" or "one mole" -- can cost marks.

Hess's Law

Hess's law states that the total enthalpy change for a reaction is independent of the route taken, provided the initial and final conditions are the same. In practice, it allows you to calculate enthalpy changes that cannot be measured directly by constructing enthalpy cycles using enthalpies of formation or combustion.

For calculations using enthalpies of formation:

Delta H(reaction) = sum of Delta H(f) products -- sum of Delta H(f) reactants

For calculations using enthalpies of combustion:

Delta H(reaction) = sum of Delta H(c) reactants -- sum of Delta H(c) products

Note the reversal of the subtraction order. This is a common source of sign errors that costs many students marks.

Born-Haber Cycles

Born-Haber cycles are extended Hess's law cycles applied to ionic compounds, allowing you to calculate lattice enthalpies indirectly. A Born-Haber cycle for sodium chloride includes the following steps:

1. Atomisation of sodium (solid to gaseous atoms)
2. First ionisation of sodium (gaseous atoms to gaseous ions)
3. Atomisation of chlorine (half a mole of Cl2 gas to one mole of gaseous Cl atoms)
4. First electron affinity of chlorine (gaseous atoms gaining an electron)
5. Lattice enthalpy of formation (gaseous ions forming the solid lattice)

The cycle closes because the enthalpy of formation of NaCl equals the sum of all these steps. By rearranging, you can calculate whichever value is unknown -- most commonly the lattice enthalpy.

Exam technique: Draw the cycle clearly with all species labelled, including state symbols. Show each step with an arrow and label it with the correct enthalpy term. Examiners frequently ask you to construct a Born-Haber cycle from scratch, not just calculate from one.

Comparing Theoretical and Experimental Lattice Enthalpies

The lattice enthalpy from a Born-Haber cycle (experimental) can be compared with the value from a purely ionic model (theoretical). If the experimental value is significantly more exothermic, this indicates covalent character -- the ions are polarising each other. This is especially relevant for small, highly charged cations paired with large, easily polarised anions, and links energetics to Fajans' rules.

Kinetics: Rate Equations and the Arrhenius Equation

Kinetics is the study of reaction rates -- how fast reactions occur and what factors influence their speed. At A-Level, kinetics becomes quantitative, moving beyond the qualitative ideas of GCSE.

Rate Equations

The rate equation for a reaction expresses the rate in terms of the concentrations of reactants raised to various powers:

Rate = k[A]^m[B]^n

Here, k is the rate constant, [A] and [B] are concentrations of reactants, and m and n are the orders of reaction with respect to each reactant. The overall order is the sum m + n.

Critically, orders of reaction cannot be deduced from the balanced equation. They must be determined experimentally. This is one of the most important conceptual points in the topic. The rate equation reflects the mechanism of the reaction, not the stoichiometry.

Determining Orders of Reaction

Orders are determined from experimental data, typically using initial rate experiments:

• Zero order -- changing the concentration has no effect on the rate.
• First order -- doubling the concentration doubles the rate.
• Second order -- doubling the concentration quadruples the rate.

You may also be given concentration-time graphs. A first-order reaction produces an exponential decay curve with a constant half-life -- this constant half-life is a distinctive feature that examiners test regularly.

The Rate-Determining Step

The rate-determining step is the slowest step in a multi-step reaction mechanism. Only species that appear in or before the rate-determining step appear in the rate equation. This provides a powerful link between kinetics and mechanism -- if the rate equation is Rate = k[A] only, then B is not involved until after the slow step.

The Arrhenius Equation

The Arrhenius equation quantifies how the rate constant k changes with temperature:

k = Ae^(-Ea/RT)

Where A is the pre-exponential factor, Ea is the activation energy, R is the gas constant (8.314 J mol-1 K-1), and T is the temperature in Kelvin. Taking natural logarithms gives the linear form:

ln k = ln A -- Ea/RT

Plotting ln k against 1/T gives a straight line with gradient --Ea/R and y-intercept ln A. This is how activation energies are determined experimentally.

Calculation tip: Always convert temperatures to Kelvin and ensure energy units are consistent. If Ea is in kJ mol-1, convert to J mol-1 before dividing by R.

Chemical Equilibria: Kc, Kp, and Le Chatelier's Principle

Equilibrium is one of the most conceptually rich topics in Physical Chemistry. It brings together ideas of rate, energy, and concentration into a unified framework.

The Equilibrium Constant Kc

For a reversible reaction at equilibrium, the equilibrium constant Kc is defined as the product of the equilibrium concentrations of the products, each raised to the power of their stoichiometric coefficients, divided by the equivalent expression for the reactants.

For the general reaction aA + bB reversible cC + dD:

Kc = [C]^c[D]^d / [A]^a[B]^b

A large Kc indicates that the equilibrium position lies to the right (favouring products). A small Kc indicates it lies to the left (favouring reactants). Kc is constant at a given temperature -- it changes only when the temperature changes.

The Equilibrium Constant Kp

For reactions involving gases, the equilibrium constant can be expressed in terms of partial pressures rather than concentrations:

Kp = (pC)^c(pD)^d / (pA)^a(pB)^b

The partial pressure of a gas is calculated as the mole fraction of that gas multiplied by the total pressure. You need to be confident converting between moles, mole fractions, and partial pressures -- these calculations appear frequently and carry significant marks.

Calculation tip: Lay out your working in a clear table with columns for moles at equilibrium, mole fractions, and partial pressures. This structured approach reduces errors and makes your method visible to examiners.

Le Chatelier's Principle

Le Chatelier's principle states that if a system at equilibrium is subjected to a change in conditions, the position of equilibrium shifts to counteract the change.

• Concentration changes -- adding more reactant shifts equilibrium to the right. The value of Kc does not change.
• Pressure changes -- increasing pressure favours the side with fewer moles of gas. Kc does not change.
• Temperature changes -- increasing temperature shifts equilibrium in the endothermic direction. This is the only change that alters the value of K.
• Catalysts -- a catalyst does not change the position of equilibrium or the value of K. It only helps the system reach equilibrium faster.

Exam technique: Always state the direction of the shift and explain the consequence for specific species. Simply writing "the equilibrium shifts right" without elaboration rarely earns full marks.

Thermodynamics: Entropy and Gibbs Free Energy

Thermodynamics at A-Level extends your understanding of energy changes beyond enthalpy. It introduces entropy and Gibbs free energy, providing a more complete picture of what drives chemical reactions.

Entropy

Entropy (S) is a measure of the disorder or dispersal of energy in a system. The second law of thermodynamics states that any spontaneous process increases the total entropy of the universe (system plus surroundings).

Key principles:

• Gases have higher entropy than liquids, which have higher entropy than solids.
• Entropy increases when the number of moles of gas increases during a reaction.
• Dissolving a solid in a solvent generally increases entropy.
• Entropy increases with temperature.

The standard entropy change of a reaction is calculated as:

Delta S(system) = sum of S(products) -- sum of S(reactants)

Note that unlike enthalpy, absolute entropy values can be determined (the entropy of a perfect crystal at 0 K is zero -- this is the third law of thermodynamics). Standard entropy values are given in data tables in J K-1 mol-1.

Gibbs Free Energy

Gibbs free energy (Delta G) combines enthalpy and entropy into a single quantity that determines whether a reaction is thermodynamically feasible:

Delta G = Delta H -- T Delta S

A reaction is thermodynamically feasible when Delta G is negative. This means:

• If Delta H is negative and Delta S is positive, Delta G is always negative -- the reaction is feasible at all temperatures.
• If Delta H is positive and Delta S is negative, Delta G is always positive -- the reaction is never feasible.
• If both Delta H and Delta S are positive, the reaction becomes feasible at high temperatures (when the T Delta S term outweighs Delta H).
• If both Delta H and Delta S are negative, the reaction is feasible at low temperatures (when Delta H outweighs the T Delta S term).

Important caveat: Thermodynamic feasibility does not guarantee that a reaction will actually occur. A reaction may have a negative Delta G but be kinetically hindered by a high activation energy barrier. Diamond converting to graphite is the classic example -- thermodynamically feasible but kinetically so slow that diamonds are effectively permanent.

Calculating the Temperature at Which a Reaction Becomes Feasible

Setting Delta G = 0 gives the temperature at which a reaction transitions between feasible and non-feasible:

T = Delta H / Delta S

This is a common calculation question. Remember to use consistent units -- if Delta H is in kJ mol-1 and Delta S is in J K-1 mol-1, you must convert one of them before dividing.

Electrode Potentials

Electrode potentials bring together ideas from energetics, equilibria, and redox chemistry. This topic explains how electrochemical cells work and introduces a quantitative method for predicting the feasibility of redox reactions.

Standard Electrode Potentials

A standard electrode potential (E-standard) is the voltage measured when a half-cell is connected to a standard hydrogen electrode under standard conditions (298 K, 100 kPa, 1.00 mol dm-3 solutions). The standard hydrogen electrode is assigned a potential of 0.00 V by convention.

Half-cells are arranged in an electrochemical series, from the most negative E-standard (strongest reducing agents) to the most positive (strongest oxidising agents). The more negative the electrode potential, the greater the tendency of the species to lose electrons (be oxidised). The more positive, the greater the tendency to gain electrons (be reduced).

Calculating Cell EMF

The electromotive force (EMF) of an electrochemical cell is calculated as:

E(cell) = E(positive electrode) -- E(negative electrode)

Or equivalently:

E(cell) = E(cathode) -- E(anode)

Where the cathode is the electrode where reduction occurs (more positive E-standard) and the anode is where oxidation occurs (more negative E-standard). A positive cell EMF indicates that the reaction is thermodynamically feasible under standard conditions.

Predicting Feasibility of Redox Reactions

To determine whether a redox reaction is feasible, identify the two half-equations, look up their standard electrode potentials, and calculate the cell EMF. A positive EMF indicates feasibility under standard conditions. However, standard electrode potentials apply only to standard conditions, and a positive EMF indicates thermodynamic feasibility -- kinetic factors may still prevent the reaction from occurring at a measurable rate.

Electrochemical Cells and Fuel Cells

You should also understand rechargeable cells (such as lithium-ion batteries) and fuel cells (such as hydrogen-oxygen fuel cells). The key idea is that chemical energy is converted to electrical energy through redox reactions at separate electrodes, with an electrolyte allowing ion transfer. Hydrogen fuel cells produce only water as a chemical product, but challenges remain around hydrogen production and storage.

Linking the Topics Together

These Physical Chemistry topics are not isolated -- they form an interconnected framework. Enthalpy is half the feasibility story; entropy completes it through Gibbs free energy. Kinetics tells you how fast a reaction reaches equilibrium; K values tell you where the balance lies. Electrode potentials connect to Gibbs free energy through Delta G = --nFE. When answering synoptic questions, demonstrating these connections is what separates good answers from excellent ones.

Exam Strategy for Physical Chemistry Questions

Physical Chemistry questions on AQA papers fall into predictable categories. For calculation questions, show every step -- formula, substitution, arithmetic, and units. Method marks are available even if you make an arithmetic error, but only if your working is visible. For definition questions, learn the specification wording precisely -- paraphrasing often costs marks. For explain and justify questions, use correct chemical terminology and always consider both thermodynamic and kinetic factors when discussing feasibility. For graph interpretation, know what the gradient and intercept represent on plots such as ln k against 1/T or rate against concentration.

Prepare with LearningBro

Physical Chemistry rewards consistent practice and a solid conceptual foundation. The following LearningBro courses are designed to help you build both:

• AQA A-Level Chemistry: Physical Chemistry in Depth -- focused, in-depth coverage of every Physical Chemistry topic on the AQA specification
• AQA A-Level Chemistry: Kinetics and Equilibria -- targeted practice on rate equations, the Arrhenius equation, Kc, Kp, and Le Chatelier's principle
• AQA A-Level Chemistry -- comprehensive coverage of the full AQA A-Level Chemistry course, including organic, inorganic, and physical topics