A-level Chemistry/WJEC/Module 2/Thermochemistry

Energetics is all about the energy transfer involved in the making and breaking of bonds. The breaking of a bond involves the injection of energy into the system so that the bond can no longer hold the parts together; the creation of bonds involves the removal of energy from the system so that components are captured by the binding forces. In a typical reaction some bonds are broken and some are created so some energy is injected and some is removed. Whether the total(net) of these amounts to an injection or removal determines whether the reaction is endothermic (involving an net injection of energy) or exothermic (involving a net removal of energy.)

The simple idea described above is complicated by the interaction of the system (reaction) being studied and its environment. This interaction amounts to the changes in volume of the system brought about by the reaction. To deal with this complexity we introduce the idea of enthalpy (given the symbol H). Enthalpy is not measured directly: we concern ourselves instead with CHANGES in enthalpy (ΔH) resulting from a reaction.

For a given reaction, change in enthalpy is equal to the heat energy change under conditions of constant pressure.

Enthalpy change

The heat energy transferred under constant pressure.

ΔH = H(products) – H(reactants) = mcΔT ÷ n

• negative ΔH

Such reactions include the combustion of a fuel like methane CH₄(g) + 2 O₂(g) → CO₂(g) + 2 H₂O(l) = -890 kJ mol⁻¹ Also including the oxidation of carbohydrates such as glucose C₆H₁₂O₆

• positive ΔH

Such reactions include the thermal decomposition of calcium carbonate CaCO₃(s) → CaO(s) + CO₂(g) = +178 kJ mol⁻¹

Standard conditions

We often deal with enthalpy change under standard conditions. i.e.

101 kPa, 298 K (WJEC use 1 atm = 101 kPa as standard pressure and 25 °C = 298 K for standard temperature).

Standard enthalpy of formation

Δ_fH^⦵

The enthalpy change when 1 mole of compound in its standard state is formed from its elements in their standard states (under standard conditions).

Standard enthalpy of combustion

Δ_cH^⦵

The enthalpy change when 1 mole of substance is completely burned in oxygen under standard conditions.

• q = mcΔT
  • heat energy = mass of substance x specific heat capacity x temperature change
• ΔH = q ÷ n
  • enthalpy change = heat energy ÷ amount

Hess's Law

The overall enthalpy change is independent of the route the chemical reaction takes place.

The enthalpy of a reaction which breaks 1 mole of a certain type of covalent bond by homolytic fission, under standard conditions.

Example:

Breaking 1 mole of Cl-Cl bonds:

Cl₂(g) → 2 Cl(g) E_d = 243 kJ mol^-1

Except for bonds which form diatomic molecules, multiple molecules can contain the same type of bond. For example, only Cl₂ has a Cl-Cl bond, but the C-Cl bond is present in a very wide range of compounds (CH₃Cl, CCl₄, CH₃CCl₃, etc). For most covalent bonds, we can look up the average bond enthalpy to find an estimate of the strength of the bond in a specific compound.

E_d is always positive because breaking bonds requires energy. Some biologist like to say that "breaking the phosphate bond in ATP releases energy" but this is wrong. Breaking the phosphate bond in ATP requires energy, like any bond-breaking process. When ADP and phosphate form from ATP and water, the new bonds formed in ADP and phosphate release more energy than was required to break the bonds in ATP and water.

We can estimate the energy of a reaction from the bond energies involved:

Δ_rH^⦵ = Σ (E_d(bonds broken)) - Σ (E_d(bonds made))

The Δ_rH^⦵ value calculated is not an exact figure, because the bond energies are average values and are not exactly the bond energies in the specific reaction.