Chemical Energetics: Exothermic and Endothermic Reactions - Calculating Enthalpy Change using Bond Energies

This section explores the concept of enthalpy change in chemical reactions, focusing on how to calculate it using bond energies. We will differentiate between exothermic and endothermic reactions and learn a practical method for determining the enthalpy change based on the breaking and forming of chemical bonds.

Exothermic and Endothermic Reactions

Exothermic Reactions: Reactions that release heat into the surroundings. The enthalpy change (ΔH) is negative.
Example: Combustion of methane (natural gas) is exothermic.

Endothermic Reactions: Reactions that absorb heat from the surroundings. The enthalpy change (ΔH) is positive.
Example: Melting ice is endothermic.

Enthalpy Change (ΔH)

Enthalpy change is a measure of the heat absorbed or released in a chemical reaction at constant pressure. It's typically expressed in units of kJ/mol (kilojoules per mole).

Calculating ΔH using Bond Energies

The enthalpy change of a reaction can be estimated using the average bond energies of the reactants and products. The principle is that breaking bonds requires energy (endothermic), and forming bonds releases energy (exothermic). The overall enthalpy change is the difference between the energy required to break bonds in the reactants and the energy released when bonds are formed in the products.

The formula is: $$ \Delta H = \text{Σ (Bond energies of bonds broken)} - \text{Σ (Bond energies of bonds formed)} $$

Where:

• Σ (Bond energies of bonds broken) represents the sum of the bond energies of all the bonds that are broken in the reactants.
• Σ (Bond energies of bonds formed) represents the sum of the bond energies of all the bonds that are formed in the products.

Example Calculation

Consider the reaction: CH₄(g) + 2O₂(g) → CO₂(g) + 2H₂O(g)

Let's assume the following bond energies (These are approximate values):

• C-H bond: 413 kJ/mol
• O=O bond: 498 kJ/mol
• C=O bond: 799 kJ/mol
• O-H bond: 463 kJ/mol

Reactants:

• CH₄: 4 bonds x 413 kJ/mol = 1652 kJ/mol
• O₂: 2 bonds x 498 kJ/mol = 996 kJ/mol

Products:

• CO₂: 2 bonds x 799 kJ/mol = 1598 kJ/mol
• H₂O: 2 bonds x 463 kJ/mol = 926 kJ/mol

ΔH = (1652 kJ/mol + 996 kJ/mol) - (1598 kJ/mol + 926 kJ/mol) = 2648 kJ/mol - 2524 kJ/mol = 124 kJ/mol

Since ΔH is positive, the reaction is endothermic.

Table of Bond Energies (Approximate Values - Important for calculations)

Bond	Bond Energy (kJ/mol)
C-H	413
O=O	498
C=O	799
O-H	463
N-H	463
N≡N	945
C-C	347
S-S	266
S-H	595
Cl-Cl	242
Cl-H	431

Important Considerations

• These bond energies are average values and can vary slightly depending on the specific molecule.
• The calculation provides an *estimate* of the enthalpy change. It's not always perfectly accurate, especially for complex reactions.
• The calculation assumes that all bonds in the reactants and products are in their ground state.

Further Learning

For a deeper understanding, explore the concept of Hess's Law, which allows us to calculate enthalpy changes for reactions that cannot be directly measured.