Summary — What This Chapter Covers
- Exothermic reactions have ΔH negative (heat given out); endothermic reactions have ΔH positive.
- A reaction pathway diagram shows ΔH and the activation energy, EA.
- Standard conditions (⦵): 298 K and 101 kPa.
- Know the definitions of ΔHr, ΔHf, ΔHc and ΔHneut.
- ΔHr = Σ(bonds broken) − Σ(bonds made). Some bond energies are exact; most are averages.
- From experiments: q = mcΔT and ΔH = −mcΔT/n.
- Hess's law: the total enthalpy change for a reaction is independent of the route taken.
1. Enthalpy Change, ΔH
Exothermic and endothermic
Sign conventions
Chemical reactions are accompanied by enthalpy changes. Enthalpy change, ΔH, is the heat energy transferred at constant pressure.
- Exothermic: energy is transferred to the surroundings; the temperature of the surroundings rises; ΔH is negative. Examples: combustion, neutralisation.
- Endothermic: energy is taken in from the surroundings; the temperature falls; ΔH is positive. Examples: thermal decomposition of CaCO₃, dissolving NH₄NO₃.
Reaction pathway diagrams
Enthalpy against progress of reaction
The diagram shows the enthalpy of reactants and products and the energy barrier between them. Activation energy, EA, is the minimum energy needed for a reaction to occur (for a collision to be effective). ΔH is the difference between products and reactants.
For an exothermic reaction the products are lower than the reactants (ΔH negative, arrow pointing down). For an endothermic reaction the products are higher (ΔH positive, arrow up). EA is always measured up from the reactants to the peak.
Standard enthalpy changes
Definitions to learn word for word
Standard conditions (shown by ⦵): 298 K and 101 kPa, with substances in their standard (normal physical) states and solutions at 1 mol dm⁻³.
| Term | Definition | Example |
|---|
| Enthalpy change of reaction, ΔHr | The enthalpy change when the amounts of reactants shown in the equation react to give products. | N₂(g) + 3H₂(g) → 2NH₃(g) ΔHr⦵ = −92 kJ mol⁻¹ |
| Enthalpy change of formation, ΔHf | The enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions. | 2C(s) + 3H₂(g) + ½O₂(g) → C₂H₅OH(l) |
| Enthalpy change of combustion, ΔHc | The enthalpy change when one mole of a substance is burnt completely in excess oxygen under standard conditions. | CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(l) |
| Enthalpy change of neutralisation, ΔHneut | The enthalpy change when one mole of water is formed by the reaction of an acid with an alkali under standard conditions. | H⁺(aq) + OH⁻(aq) → H₂O(l) about −57 kJ mol⁻¹ |
ΔHf of an element in its standard state is zero by definition.
Bond energies
Why reactions transfer energy
bonds are broken and made
needs
releases
From bond energies
ΔHr = Σ(bond energies of bonds broken) − Σ(bond energies of bonds made)
Exact and average bond energies
average
Measuring enthalpy changes