5.5.c. Bond Enthalpy

Chemical reactions involve the breaking and making of chemical bonds. Energy is required to break a bond and energy is released when a bond is formed. It is possible to relate heat of reaction to changes in energy associated with breaking and making of chemical bonds.

Bond enthalpy refers to the energy required to break one mole of bonds in gaseous molecules. To understand this concept, remember that during a chemical reaction, bonds between atoms in molecules are broken, and new bonds are formed as the reaction proceeds. When we say energy is required to break bonds, this means that energy must be supplied to overcome the forces holding atoms together. Conversely, when new bonds form, energy is released into the environment.

Think of bond enthalpy like a rubber band. If you want to stretch (break) the rubber band, you need to pull hard on it until it snaps. This requires effort (energy). Once snapped, if you let the rubber band go, it snaps back to its original shape, releasing the energy stored in it. This is similar to how energy is released when bonds form in a chemical reaction.

Consider the following process in which the bonds in one mole of dihydrogen gas (H2) are broken: H2(g) → 2H(g); ∆H–HH = 435.0 kJ mol–1 The enthalpy change involved in this process is the bond dissociation enthalpy of H–H bond. The bond dissociation enthalpy is the change in enthalpy when one mole of covalent bonds of a gaseous covalent compound is broken to form products in the gas phase.

For diatomic molecules like hydrogen gas (H2), the bond enthalpy can be measured as the energy required to break the hydrogen-hydrogen bond to form two hydrogen atoms. This is an essential concept because it quantifies the strength of the bond between the two hydrogen atoms. The value of 435 kJ/mol indicates that 435 kilojoules of energy must be supplied to break one mole of H2 molecules into two moles of H atoms.

Imagine stretching a tightly-held string between two people. The effort it takes to pull the string apart can be likened to the energy needed to break the bond. The tighter (stronger) the string is held, the more effort you’ll need to pull it apart, just like with bond enthalpy.

In the case of polyatomic molecules, bond dissociation enthalpy is different for different bonds within the same molecule.

Polyatomic molecules, such as methane (CH4), have multiple types of bonds. In such cases, the bond enthalpy will vary between different types of bonds. For example, the energy required to break the first carbon-hydrogen bond in methane may differ from breaking the second or third. This variability arises due to the influence of other nearby atoms (their electronic effects) and the structural environment within the molecule.

If you think about a necklace made of different types of beads, each bead (or bond) might require a different amount of effort to remove based on how tightly it's strung on the string. The energy to remove each bead is akin to the different bond dissociation enthalpies for various bonds within a polyatomic molecule.

In such cases we use mean bond enthalpy of C – H bond. For example in CH4, ∆C–H H is calculated as: ∆C–H H = ¼ (∆a H) = ¼ (1665 kJ mol–1) = 416 kJ mol–1.

Mean bond enthalpy is an average value for bond strengths in molecules that contain multiple identical bonds. For methane, the four C-H bonds are identical, but the energy required to break them sequentially is different. To simplify calculations, we can average these values to find a mean value for bond strength, which helps in predicting reaction energies and understanding molecular stability.

Consider a classroom full of chairs arranged in rows. Each row can be viewed as a bond type. While the chairs in the front row might be difficult to move out for cleaning versus the last row that is easy to move, if you were to take an average of the effort required to move any chair in the classroom, you’d get a general idea of ‘average chair difficulty’ – this is similar to how mean bond enthalpy provides an averaged strength across identical bonds.

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