Reactivity 1.2.1 – Bond Enthalpy

Energy Changes in Chemical Reactions

Chemical reactions involve changes in the bonding between atoms. During a chemical reaction:

• Bonds in the reactants must be broken to allow atoms to rearrange.
• New bonds must form between atoms to produce the products.

Bond Breaking is Endothermic

Breaking chemical bonds always requires an input of energy because it involves overcoming the electrostatic attraction between the bonded atoms. This energy is needed to separate atoms and pull them apart against the forces holding them together.

As a result, bond breaking is classified as an endothermic process, meaning energy is absorbed from the surroundings. The energy required to break a particular type of bond is called its bond enthalpy.

Bond Forming is Exothermic

Forming chemical bonds releases energy. When atoms bond, they move to a lower energy state and become more stable. The energy released corresponds to the strength of the new bond formed. This process is exothermic, as energy is given out to the surroundings.

The greater the bond strength, the more energy is released during bond formation.

Energy Profile Diagram

Energy profile diagrams visually represent the energy changes that occur during a chemical reaction:

• For an exothermic reaction, the products have a lower energy than the reactants. Energy is released.
• For an endothermic reaction, the products have a higher energy than the reactants. Energy is absorbed.

Both types of reactions require an initial input of energy known as the activation energy, which is the minimum energy required to initiate bond breaking in the reactants.

What Is Average Bond Enthalpy?

The average bond enthalpy is the average amount of energy required to break one mole of a specific type of covalent bond in the gaseous state. It is measured in kilojoules per mole (kJ/mol).

Because the same type of bond (e.g., O–H or C–H) can exist in different molecules with slightly different environments, the bond enthalpy value is averaged across many compounds to give a standardized value suitable for general use in calculations.

Key Formula for Enthalpy Change

The enthalpy change of a chemical reaction depends on the balance between energy absorbed to break bonds and energy released in forming new bonds. The overall enthalpy change of a reaction, \( \Delta H \), can be estimated using the following formula:

\( \Delta H = \sum \text{(Bond enthalpies of bonds broken)} – \sum \text{(Bond enthalpies of bonds formed)} \)

• Bonds broken: Energy absorbed (positive enthalpy change, endothermic)(\( \Delta H > 0 \))
• Bonds formed: Energy released (negative contribution to ΔH, exothermic)(\( \Delta H < 0 \))

Step-by-Step Calculation Method

1. Write the balanced chemical equation for the reaction.
2. List all bonds present in the reactants and products.
3. Use average bond enthalpy values (from the data booklet) to calculate:
   • Total energy required to break all bonds in the reactants
   • Total energy released when new bonds are formed in the products
4. Apply the formula:
   \( \Delta H = \text{Total energy in bonds broken} – \text{Total energy in bonds formed} \)

Why Are These Values Averages?

The bond enthalpies are averaged values for the following reasons:

• The energy required to break a particular bond (e.g., C–H) can vary slightly depending on the surrounding atoms and molecular structure.
• For example, the O–H bond in water is slightly different from the O–H bond in ethanol or methanol.
• To provide a single usable value, data booklets report mean bond enthalpies averaged over a wide range of compounds.

Limitations of Using Average Bond Enthalpy Values

• They are valid only for substances in the gas phase. If the actual state is liquid or solid (e.g., liquid water), the enthalpy change will differ.
• They do not account for resonance structures (e.g., benzene), which distribute bonding over multiple atoms.
• They ignore intermolecular forces and energy changes associated with phase changes (e.g., vaporization or condensation).
• Because of these limitations, the ΔH calculated using bond enthalpies is only an approximation. Experimental values (e.g., from calorimetry) are usually more accurate.