Chemical reactions involve bond-breaking and bond-forming

In a chemical reaction, the atoms in the reactants are rearranged to produce new substances called products. This process involves:

• Breaking chemical bonds in the reactants - An endothermic process that requires energy input.
• Forming new chemical bonds to create the products - An exothermic process that releases energy.

The enthalpy change (ΔH) of a reaction depends on the relative amounts of energy involved in bond breaking and bond formation:

• If bond breaking requires more energy than is released during bond formation, he reaction has a positive ΔH.
• If bond breaking requires less energy than is released during bond formation, the reaction has a negative ΔH.

Bond enthalpy is the energy required to break a bond

Bond enthalpy (BE) is defined as the energy needed to break one mole of a particular bond by homolytic fission in one mole of gaseous covalent molecules under standard conditions.

Key points about bond enthalpy:

• It is an endothermic process, so bond enthalpy values are always positive.
• Homolytic fission distributes the bonding electrons equally between the two new species, forming radicals (denoted by the • symbol).
• Bond enthalpy values are average values derived from experimental data on breaking the same bond in various compounds.

For example, the bond enthalpy of the H–H bond in a hydrogen molecule (H₂) is equal to the enthalpy change of the following reaction:

H_2(g) ➔ 2H•_(g) ΔH = +436 kJ mol^-1

Bond enthalpies vary with chemical environment

The bond enthalpy for a specific bond can vary depending on its chemical environment within a molecule.

For instance, if water (H₂O) underwent a series of steps where one hydrogen atom was removed at a time, the bond enthalpy for each O–H bond would differ:

As a result, bond enthalpy values provided in data tables are averaged values. For example, the average bond enthalpy of the O–H bond is +464 kJ mol^-1.

Calculating enthalpy changes using bond enthalpies

You can estimate the enthalpy change of a reaction (ΔH) using bond enthalpy data:

ΔH = ΣBE of bonds broken $-$ ΣBE of bonds formed

Where:

• ΣBE of bonds broken = Sum of the bond enthalpies of all bonds broken in the reactants.
• ΣBE of bonds formed = Sum of the bond enthalpies of all bonds formed in the products.

However, since bond enthalpies are average values, the calculated ΔH will differ from the actual value. Additionally, bond enthalpy data does not account for intermolecular forces, which is especially important for solid or liquid reactants and products.

Worked example 1 - Calculating ΔH for the combustion of methane

Calculate the enthalpy change (ΔH) for the combustion of methane (CH₄) with oxygen to form carbon dioxide and water. The balanced chemical equation is:

CH₄ + 2O₂ ➔ CO₂ + 2H₂O

Given the average bond enthalpies:

Step 1: Identify and count the bonds broken and formed

Bonds broken: 4 x C-H and 2 x O=O

Bonds formed: 2 x C=O and 4 x O-H

Step 2: Calculate total energy absorbed in breaking bonds

Energy to break bonds = (4 x 414) + (2 x 498) = 1,656 + 996 = 2,652 kJ mol^-1^

Step 3: Calculate total energy released in forming bonds

Energy released in bond formation = (2 x 799) + (4 x 464) = 1,598 + 1,856 = 3,454 kJ mol^-1^

Step 4: Calculate ΔH for the reaction

ΔH = ΣBE of bonds broken $-$ ΣBE of bonds formed

ΔH = 2,652 $-$ 3,454 = $-$802 kJ mol^-1^

The negative sign indicates that the reaction is exothermic, releasing 802 kJ mol^-1^ of energy.

Worked example 2 - Calculating ΔH for the synthesis of ammonia

Calculate the enthalpy change (ΔH) for the synthesis of ammonia (NH₃) from nitrogen and hydrogen. The balanced chemical equation is:

N₂ + 3H₂ ➔ 2NH₃

Given the average bond enthalpies:

Step 1: Identify and count the bonds broken and formed

Bonds broken: 1 x N≡N and 3 x H-H

Bonds formed: 6 x N-H

Step 2: Calculate total energy absorbed in breaking bonds

Energy to break bonds = (1 x 945) + (3 x 436) = 945 + 1,308 = 2,253 kJ mol^-1^

Step 3: Calculate total energy released in forming bonds

Energy released in bond formation = (6 x 391) = 2,346 kJ mol^-1^

Step 4: Calculate ΔH for the reaction

ΔH = ΣBE of bonds broken $-$ ΣBE of bonds formed

ΔH = 2,253 $-$ 2,346 = $-$93 kJ mol^-1^

The negative sign indicates that the reaction is exothermic, releasing 93 kJ mol^-1^ of energy.