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Chapter 5: Energetics 

5.1 Exothermic and endothermic reactions

  • Energetics deals with heat changes occurring during a chemical reactions.
  • Enthalpy is the amount of heat energy contained in a substance. It is stored in the chemical bonds as potential energy. When substances react, the difference in the enthalpy between the reactants and products (at constant pressure) results in a heat change which can be measured.
  • The reaction mixture is called the system and anything around the system is called the surroundings.
  • Thermochemical equations give the balanced equation with the enthalpy change.

e.g. H2 (g) + ½O2 (g) ® H2O (l); DH q = –286 kJ mol–1

H2 (g) + ½O2 (g) ® H2O (g); DH q = –242 kJ mol–1

State symbols must be shown as DH q depends on the state of the reactants or products.

  • In exothermic reactions heat is released to the surroundings.
  • In endothermic reactions heat is absorbed from the surroundings.
  • The standard enthalpy change (DH q) is the heat energy transferred under standard conditions (pressure 101.3 kPa, temperature 298 K). Only DH q can be measured, not H for the initial or final state of a system.
  • The standard enthalpy change of combustion is the enthalpy change for the complete combustion of one mole of a substance in its standard state in excess oxygen under standard conditions. All combustion reactions are exothermic.
  • The enthalpy of neutralization is the enthalpy change when one mol of H + (aq) reacts with one mol of OH (aq) ions. The reaction is exothermic as bond formation takes place: H + (aq) + OH (aq) ¯ H2O (l).
  • Exothermic reactions have negative ΔH The temperature of the reaction mixture rises as the chemicals give out heat.
  • Endothermic reactions have positive ΔH The temperature of the reaction mixture falls as the chemicals absorb heat.









An exothermic reaction: The products are more stable than the reactants as they have a lower enthalpy.



An endothermic reaction: The products are less stable than the reactants as they have a higher enthalpy.

5.2 Calculation of enthalpy changes

  • Calorimetry is the technique of measuring heat changes in physical processes and chemical reactions.
  • Heat changes can be calculated from the temperature changes:

heat change (q) = mass (m) ´ specific heat capacity (c) ´ temperature change (ΔT).

  • The specific heat capacity is the amount of heat energy required to raise the temperature of unit mass (e.g. 1 kg or 1 g) of a substance, by 1°C or 1 K.
  • and for reactions in aqueous solutions can be calculated if it is assumed that all the heat goes into the water.
 = –mH2O ´ cH2O ´ ΔTH2O/nfuelThe experiment is performed with a calorimeter which is a good conductor. This allows heat from the flame to pass to the water.  = –mH2O ´ cH2O ´ ΔTH2O/nlimiting reagentThe experiment is performed with a calorimeter which is an insulator of heat, which reduces heat losses from the system.

If a calorimeter absorbs heat: Q = (mH2O ´ cH2O ´ ΔTH2O) + (mcalor ´ ccalor ´ ΔTcalor).

Heat loss and incomplete combustion can lead to systematic errors in experimental results.

5.3 Hess’s law

  • Hess’s law states that the total enthalpy change for a reaction is independent of the route taken. It is a special case of the law of conservation of energy.



Hess’s law:

ΔH3 = ΔH1 + ΔH2

5.4 Bond enthalpies

  • Average bond energy is the energy required to break one mole of the same type of bonds in the gaseous state averaged over a variety of similar compounds.
  • Bond breaking absorbs energy and is endothermic. Bond making releases energy and is exothermic.

= Σ Ebonds broken – Σ Ebonds formed

When Σ Ebonds broken  > Σ Ebonds formed : the reaction is endothermic.

When Σ Ebonds formed  > Σ Ebonds broken : the reaction is exothermic.

15.1 Standard enthalpy changes of reaction

  • The standard state of an element or compound is its most stable state under the standard conditions (pressure 101.3 kPa, temperature 298 K).
  • The standard enthalpy change of combustion is the enthalpy change for the complete combustion of one mole of a substance in its standard state in excess oxygen under standard conditions.
  • The standard enthalpy change of formation is the enthalpy change when one mole of a substance is formed from its elements in their standard states under standard conditions.
  • The enthalpy of formation of any element in its stable state is zero, as there is no enthalpy change when an element is formed from itself.



Using  to find





Using  to find






calculated from  or are more accurate than  values based on bond enthalpies, which refer only to the gaseous state and are average values.

15.2 Born–Haber cycles

  • The first electron affinity is the enthalpy change when one mole of gaseous atoms attracts one mole of electrons: X (g) + e (g) ® X (g) .
  • The lattice enthalpy is the enthalpy change that occurs when one mole of a solid ionic compound is separated into gaseous ions under standard conditions. For example, for alkali metal halides: MX (s) ® M + (g) + X (g) .
  • depends on the attraction between the ions:
  • an increase in the ionic radius of the ions decreases .
  • an increase in ionic charge increases .
  • The Born–Haber cycle is a special case of Hess’s law for the formation of ionic compounds. It allows the experimental lattice enthalpy to be calculated from other enthalpy changes.
  • Theoretical lattice enthalpies can be calculated using a (purely) ionic model from the ionic charges and radii.
  • Differences between the theoretical and experimental lattice enthalpies give an indication of the covalent character of the compound; the greater the difference the more covalent the compound.

Born–Haber cycle for NaCl










= 411 + 107 + ½( + 243) + 496 – 349 = + 786.5 kJ mol1

15.3 Entropy

  • Entropy (S) is a property which quantifies the degree of disorder or randomness in a system.
  • Ordered states have low S, disordered states have high S: S (s)<. S (l)< S (g).
  • Generally matter and energy become more disordered, and Suniverse
  • = ΣS q (products) – ΣS q (reactants).

15.4 Spontaneity

  • Gibbs’ free energy (G) is the criterion for predicting the spontaneity of a reaction or process: it is related to . It gives the energy available to do useful work and is related to the enthalpy and entropy changes of the system: .
  • ∆Gsys <0 for a spontaneous process. ∆Gsys = 0 at equilibrium.
Calculating  (when T = 298 K) Calculating  (for all T)

T is in K. As the units of S are J mol–1 K–1 and H are kJ mol–1 they need to be changed to be consistent.

  • ∆Gsys and thus the direction of change varies with temperature.

At low temp: : exothermic reactions are spontaneous.

At high temp: : this allows some endothermic reactions to occur if .