Energetics and thermodynamics

Study H2 Chemistry enthalpy, calorimetry and thermochemical cycles, then use entropy and Gibbs energy to reason about reaction feasibility.

GCE A-Level H2 Chemistry 2027
9 lessons

Start topic

Before you begin

This topic is about the energy changes in reactions and whether reactions are spontaneous. It defines enthalpy changes and measures them by calorimetry, calculates them with Hess’s law, bond energies and Born–Haber cycles, and then uses entropy and Gibbs free energy to predict whether a reaction can happen.

Be comfortable with: the mole and reacting quantities, ionisation energy, and converting between J and kJ, and between °C and K.

Learning goals

Enthalpy Changes and Energy Profiles
Calorimetry (q = mcΔT)
Hess’ Law and Cycles
Bond Enthalpy Calculations
Lattice Energy and Born–Haber Cycles
Entropy and Gibbs Free Energy

Syllabus statements covered

explain that most chemical reactions are accompanied by energy changes, principally in the form of heat usually associated with the breaking and forming of chemical bonds; the reaction can be exothermic (∆H negative) or endothermic (∆H positive)
construct and interpret an energy profile diagram, in terms of the enthalpy change of the reaction and of the activation energy (see also Section 8)
explain and use the terms: — enthalpy change of reaction and standard conditions, with particular reference to: formation; combustion; hydration; solution; neutralisation; atomisation
calculate enthalpy changes from appropriate experimental results, including the use of the relationship: heat change = mc∆T
apply Hess’ Law to construct simple energy cycles, e.g. Born-Haber cycle, and carry out calculations involving such cycles and relevant energy terms (including ionisation energy and electron affinity), with particular reference to: — determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion
explain and use the terms: — bond energy (∆H positive, i.e. bond breaking) (see also Section 2)
apply Hess’ Law to construct simple energy cycles, e.g. Born-Haber cycle, and carry out calculations involving such cycles and relevant energy terms (including ionisation energy and electron affinity), with particular reference to: — average bond energies
explain and use the terms: — lattice energy (∆H negative, i.e. gaseous ions to solid lattice)
explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy
apply Hess’ Law to construct simple energy cycles, e.g. Born-Haber cycle, and carry out calculations involving such cycles and relevant energy terms (including ionisation energy and electron affinity), with particular reference to: — the formation of a simple ionic solid and of its aqueous solution
explain and use the term entropy
discuss the effects on the entropy of a chemical system by the following: — change in temperature
discuss the effects on the entropy of a chemical system by the following: — change in phase
discuss the effects on the entropy of a chemical system by the following: — change in the number of particles (especially for gaseous systems) [quantitative treatment is not required]
predict whether the entropy change for a given process or reaction is positive or negative
state and use the equation involving standard Gibbs free energy change of reaction, ∆G⦵: ∆G⦵ = ∆H⦵ − T∆S⦵ [the calculation of standard entropy change, ∆S⦵, for a reaction using standard entropies, S⦵, is not required]
state whether a reaction or process will be spontaneous by using the sign of ∆G⦵
understand the limitations in the use of ∆G⦵ to predict the spontaneity of a reaction
predict the effect of temperature change on the spontaneity of a reaction, given standard enthalpy and entropy changes

Lessons

Work through them in order.

Practise and check

Or choose

Topic reference

Quick reference

To find	Use	Check
heat from a temperature change	q = mcΔ T	m of the solution or water heated, not of the reactant
enthalpy change from q	Δ H = -q/n	n matches the “per mole” in the definition
Δ H from formation data	∑ ν Δ Hf⦵(products) - ∑ ν Δ Hf⦵(reactants)	multiply by coefficients
Δ H from combustion data	∑ ν Δ Hc⦵(reactants) - ∑ ν Δ Hc⦵(products)	reverse order to formation data
Δ H from bond energies	∑ E(bonds broken) - ∑ E(bonds formed)	gases only; an estimate
lattice energy	Δ Hf⦵ = Δ Hₐₜ + IE + EA + Δ Hₗₐₜₜ, summed over every step	include IE₂ and EA₂ for 2+ and 2− ions
spontaneity	Δ G⦵ = Δ H⦵ - TΔ S⦵	Δ S in kJ mol⁻¹ K⁻¹; T in K
boundary temperature	T = Δ H⦵/Δ S⦵	only when both have the same sign

Core knowledge to remember

Δ H is negative for an exothermic reaction and positive for an endothermic one.
Standard conditions are 298 K and 1 bar, with solutions at 1 mol dm⁻³ and substances in their standard states.
Hess’s law: the enthalpy change of a reaction is independent of the route.
Enthalpy changes of formation and combustion are per mole of the substance formed or burned. An element in its standard state has Δ Hf⦵ = 0.
Lattice energy is the enthalpy change when one mole of an ionic solid forms from its gaseous ions. It is negative, and more negative for higher ionic charges and smaller ions.
A catalyst lowers the activation energy but does not change Δ H or Δ G⦵.
Entropy increases from solid to liquid to gas, with temperature, and when a reaction makes more moles of gas.
A reaction is spontaneous when Δ G⦵ is negative, but it may still be too slow to observe.

Common mistakes

Using the mass of the reactant instead of the solution in q = mcΔ T.
Copying the sign of the heat gained by the solution into Δ H.
Forgetting to multiply by coefficients in a Hess cycle, or to reverse the sign when reversing an equation.
Using “products minus reactants” with combustion data.
Leaving out a second ionisation energy or electron affinity in a Born–Haber cycle.
Mixing J and kJ in Δ G⦵ = Δ H⦵ - TΔ S⦵, or using °C.
Saying a spontaneous reaction must be fast.