Organic Chemistry/Introduction to reactions/Bond dissociation energy

> Introduction to reactions

The bond dissociation energy, or bond enthalpy, for a diatomic molecule X-Y is defined as the energy required to break one mole of X-Y bonds, as illustrated in the following process...


 * X-Y(g) &rarr; X(g) + Y(g)

Bond enthalpies always refer to breaking bonds under gaseous conditions.

The mean molar bond enthalpy is an average value that is quoted for a bond that can occur in different molecular environments. An example is methane, CH4


 * CH4(g) &rarr; C(g) + 4H(g)

Bond enthalpy values are used in Hess's Law Calculations.

The standard enthalpy of a reaction can be found by considering the bond enthalpies of the products and reactants of the reaction -


 * Standard Enthalpy = &Sigma; Enthalpy of formation of the products - &Sigma; Enthalpy of formation of reactants

For stronger bonds, bond dissociation energy is higher as more energy is needed to break the bond.

A carbon-carbon double bond is stronger than a single bond and requires more energy to be broken. However, a carbon-carbon double bond is not twice as strong as a single one, it is only 1.5 times stronger.

All chemical bonds need an input of energy to be broken, as bonds allow a lower energy state for the component atoms. If a bond did not offer a lower energy state for the atoms that form it, a bond would not form.