Question

What is meant by the term average bond enthalpy? Why is there difference in bond enthalpy of \[\ce{O - H}\] bond in ethanol \[\ce{(C2H5OH)}\] and water?

Answer 1

The average bond enthalpy is defined as the ratio of total bond dissociation enthalpy to the number of bonds broken in the structure.

The identical \[\ce{O - H}\] bonds in water molecule does not have the same bond enthalpies. According to the structure of a water molecule, there are two \[\ce{O - H}\] bonds, but there is a change in the breaking of the first \[\ce{O - H}\] bond than the second because of the different charge. 

Hence in water molecule the average bond enthalpy will be:

The average \[\ce{O - H}\] bond enthalpy = (502 + 427)/2 = 464 kJ mol^–1.

In ethanol \[\ce{C2H5OH}\] the bond enthalpy of \[\ce{O - H}\] is different because the electronic environment around oxygen atom is different. In ethanol the \[\ce{O - H}\] is attached to the carbon atom and in water molecule the \[\ce{O - H}\] is attached to the hydrogen atom.