Chemical Energetics

The study of chemical energetics deals with the energy changes during a chemical reaction. During chemical reactions, energy is either released or absorbed. Chemical energetics tells us whether a chemical reaction can take place spontaneously or not and about the feasibility of a reaction.

Course Contents

  • Enthalpy changes: ∆H, of formation; combustion; hydration; solution; neutralisation; atomisation; bond energy; lattice energy; electron affinity
  • Hess’ Law, including Born-Haber cycles
  • Entropy and Free Energy

Learning Outcomes

Candidates should be able to:

  • (a) 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)
  • (b) 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)
  • (c) explain and use the terms:
  • (i) enthalpy change of reaction and standard conditions, with particular reference to: formation; combustion; hydration; solution; neutralisation; atomisation
  • (ii) bond energy (∆H positive, i.e. bond breaking) (see also Section 2)
  • (iii) lattice energy (∆H negative, i.e. gaseous ions to solid lattice)
  • (d) calculate enthalpy changes from appropriate experimental results, including the use of the relationship:
  • heat change = mc∆T
  • (e) explain, in qualitative terms, the effect of ionic charge and of ionic radius on the numerical magnitude of a lattice energy
  • (f) 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:
  • (i) determining enthalpy changes that cannot be found by direct experiment, e.g. an enthalpy change of formation from enthalpy changes of combustion
  • (ii) the formation of a simple ionic solid and of its aqueous solution
  • (iii) average bond energies
  • (g) explain and use the term entropy
  • (h) discuss the effects on the entropy of a chemical system by the following:
  • (i) change in temperature
  • (ii) change in phase
  • (iii) change in the number of particles (especially for gaseous systems)
  • (iv) mixing of particles (quantitative treatment is not required)
  • (i) predict whether the entropy change for a given process or reaction is positive or negative
  • (j) 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)
  • (k) state whether a reaction or process will be spontaneous by using the sign of ∆G⦵
  • (l) understand the limitations in the use of ∆G⦵ to predict the spontaneity of a reaction
  • (m) predict the effect of temperature change on the spontaneity of a reaction, given standard enthalpy and entropy changes