Enthalpy Change

Specification Reference Physical Chemistry: Chemical energetics 5.1

Quick Notes

Enthalpy change (ΔH) is the heat energy change of a reaction or process measured at constant pressure.
Exothermic reactions release energy (ΔH is negative). Examples: Combustion, neutralisation.
Endothermic reactions absorb energy (ΔH is positive). Example: Thermal decomposition.
Standard conditions: 298 K, 101 kPa, 1 mol dm⁻³ solution. Represented by ⦵.
Types of enthalpy change:
- ΔH_r (Reaction) – Enthalpy change for a reaction in molar quantities of balanced reaction equation.
- ΔH_f (Formation) – Enthalpy change when 1 mole of a compound formed from its elements (all in standard states).
- ΔH_c (Combustion) – Enthalpy change when 1 mole of a substance is burned completely in oxygen.
- ΔH_neut (Neutralisation) – Enthalpy change when acid and base react to form 1 mole of water
We can use bond energies to estimate ΔH
- (ΔH = bonds broken – bonds made).
Enthalpy change can be calculated using experimental data and q = mcΔT and ΔH = –q / n.

Full Notes

What Is Enthalpy Change (ΔH)?

Enthalpy (H) is the total heat energy of a chemical system at constant pressure. ΔH is the change in enthalpy when a chemical reaction occurs.

If ΔH is negative, the reaction is described as exothermic.

Heat energy is released to the surroundings. The temperature of the surroundings increases.

Products have less energy than reactants.

Examples: Combustion, Neutralisation.

If ΔH is positive, the reaction is described as endothermic.

Heat energy is absorbed from the surroundings. The temperature of the surroundings decreases.

Products have more energy than reactants.

Example: Thermal decomposition of calcium carbonate.

Reaction Pathway Diagram

Reaction pathway diagrams show how the energy of reactants change in a reaction as they turn into new products. The energy difference between reactants and the profile peak represents activation energy (E_a).

Exothermic: reactants start high, products end lower, ΔH is negative.

Endothermic: reactants start low, products end higher, ΔH is positive.

Standard Conditions and Types of Enthalpy Change

Standard conditions are used to enable enthalpy changes to be compared for similar reactions.

Standard conditions (⦵): 298 K, 101 kPa, 1 mol dm⁻³ solution.

There are several types of enthalpy change you need to know for CIE - make sure you know the definitions fully.

ΔH_r (Reaction)

Example: H₂ + ½O₂ → H₂O, ΔH_r = –286 kJ mol⁻¹

ΔH_f (Formation)

Example: C(s) + 2H₂(g) → CH₄(g)

ΔH_c (Combustion)

Example: CH₄ + 2O₂ → CO₂ + 2H₂O

ΔH_neut (Neutralisation)

Example: H⁺ + OH⁻ → H₂O, typically ΔH = –57 kJ mol⁻¹

Bond Enthalpies and ΔH Calculations

The enthalpy change of a reaction can be estimated using:

Where:

Bonds broken (reactants) → Energy absorbed (+ΔH)
Bonds formed (products) → Energy released (–ΔH)

Bond enthalpy values are often averages meaning experimental values for ΔH will often be slightly different to calculated values.

Matt’s exam tip

Bond enthalpies are for substances in gaseous states. Sometimes you need to use enthalpy of vaporisation first in calculations.

Worked Example

Combustion of Methane (CH₄)

Reaction:
Given bond enthalpies:
C–H = +412 kJ mol⁻¹
O=O = +498 kJ mol⁻¹
C=O = +805 kJ mol⁻¹
O–H = +463 kJ mol⁻¹

Bonds Broken (Reactants – Energy Absorbed)
CH₄: 4 × C–H = 4 × 412 = 1648 kJ
O₂: 2 × O=O = 2 × 498 = 996 kJ
Total energy to break bonds = 1648 + 996 = 2644 kJ
Bonds Formed (Products – Energy Released)
CO₂: 2 × C=O = 2 × 805 = 1610 kJ
H₂O: 4 × O–H = 4 × 463 = 1852 kJ
Total energy released = 1610 + 1852 = 3462 kJ
Calculate Enthalpy Change
ΔH = Bonds broken – Bonds formed
ΔH = 2644 – 3462
ΔH = –818 kJ mol⁻¹ (exothermic reaction)

Calculating ΔH from Experimental Data

Calorimetry is an experimental technique used to measure enthalpy changes. See this page for more detailed information and practical procedures.

The key equation used is: q = mcΔT

q = heat energy change (J)
m = mass of substance heated (g)
c = specific heat capacity (J g⁻¹ K⁻¹) (for water, c = 4.18 J g⁻¹ K⁻¹)
ΔT = temperature change (K)

The enthalpy change per mole of reactant is found using: ΔH = –q / n
where n = moles of the limiting reactant. Convert ΔH to kJ mol⁻¹ by dividing q by 1000.

Matt’s exam tip

Remember the m in q = mcΔT is the mass of the surroundings or substance being heated. It isn’t the mass of reactants! You will often be given the mass of reactant used however this is for calculating moles with ΔH = –q / n.

Summary

Enthalpy change (ΔH) is the heat energy change of a reaction at constant pressure.
Exothermic: ΔH negative, heat released. Endothermic: ΔH positive, heat absorbed.
Standard conditions: 298 K, 101 kPa, 1 mol dm⁻³.
Bond enthalpies: ΔH = bonds broken – bonds formed.
Experimentally: q = mcΔT and ΔH = –q / n.