Bond Breaking and Formation for the ESAT

The Energetics of Chemical Bonds

Every chemical reaction involves two distinct processes: the breaking of existing chemical bonds in the reactants and the formation of new chemical bonds in the products. These processes are energetically opposite.

Bond Breaking is Endothermic

To break a chemical bond, energy must be supplied to the system to overcome the attractive forces between atoms, such as the electrostatic attraction between nuclei and the shared pair of electrons in a covalent bond. Because energy is taken in from the surroundings, bond breaking is an endothermic process. In calculations, the energy associated with breaking bonds is always assigned a positive value.

Bond Formation is Exothermic

When atoms or ions join together to form new bonds, they move to a more stable, lower energy state. This transition results in the release of energy to the surroundings. Therefore, bond formation is an exothermic process. In calculations, the energy released when bonds are made is treated as a negative contribution to the overall enthalpy change.

Using Bond Energy Data

Bond energy (or bond enthalpy) is the amount of energy required to break one mole of a particular type of bond in a gaseous molecule. To determine the overall energy change ($\Delta H$) of a reaction, we compare the total energy required to break the reactant bonds with the total energy released when the product bonds are formed.

The Calculation Method

To calculate the enthalpy change of a reaction using bond energies, follow these three steps:

1. Calculate the total energy required to break all bonds in the reactants. Ensure you multiply the bond energy by the number of bonds of that type in the molecule and the stoichiometry of the balanced equation.
2. Calculate the total energy released when all bonds in the products are formed. Again, account for every bond present in the product molecules.
3. Subtract the energy released from the energy taken in: $\Delta H = \text{Total energy to break bonds} - \text{Total energy released by forming bonds}$

If the result is negative, the reaction is exothermic, as more energy was released than was required. If the result is positive, the reaction is endothermic.

Worked Example: The Combustion of Methane

Calculate the enthalpy change for the complete combustion of methane: $\text{CH}_4(\text{g}) + 2\text{O}_2(\text{g}) \rightarrow \text{CO}_2(\text{g}) + 2\text{H}_2\text{O}(\text{g})$

Bond Energy Data:

• $\text{C-H}$: $413 \text{ kJ mol}^{-1}$
• $\text{O=O}$: $495 \text{ kJ mol}^{-1}$
• $\text{C=O}$: $799 \text{ kJ mol}^{-1}$
• $\text{O-H}$: $463 \text{ kJ mol}^{-1}$

Step 1: Energy to break bonds (Reactants)

• One molecule of $\text{CH}_4$ has four $\text{C-H}$ bonds: $4 \times 413 = 1652 \text{ kJ}$
• Two molecules of $\text{O}_2$ have two $\text{O=O}$ bonds: $2 \times 495 = 990 \text{ kJ}$
• Total energy required $= 1652 + 990 = 2642 \text{ kJ}$

Step 2: Energy released by forming bonds (Products)

• One molecule of $\text{CO}_2$ has two $\text{C=O}$ bonds: $2 \times 799 = 1598 \text{ kJ}$
• Two molecules of $\text{H}_2\text{O}$ have four $\text{O-H}$ bonds ($2$ per molecule): $4 \times 463 = 1852 \text{ kJ}$
• Total energy released $= 1598 + 1852 = 3450 \text{ kJ}$

Step 3: Calculate $\Delta H$ $\Delta H = 2642 - 3450 = -808 \text{ kJ mol}^{-1}$

The reaction is exothermic because $\Delta H$ is negative.

Average Bond Enthalpies and the Gas Phase

Most bond energy values provided in exam tables are average bond enthalpies. This means they are the mean value for a specific bond type measured across a wide variety of different compounds. While these averages allow for useful estimations, the actual energy of a bond can vary slightly depending on the specific molecular environment. Furthermore, bond energy calculations assume that all reactants and products are in the gaseous state. If a substance is a liquid or solid, the calculation would also need to account for the energy involved in changes of state, such as the heat of vaporisation or fusion.