ELECTRONEGATIVITY

How Is Electronegativity Related To Bond Enthalpies?

Pauling wanted to know how to estimate the bond enthalpy of a heteronuclear bond for a diatomic molecule, e.g. X-Y. First he thought that perhaps it was the average of the bond enthalpies for X-X and Y-Y. He tried this for a few diatomics and compared his answer to the experimental value. Look at the table below (all energies in KJ mol^-1:

As you can see this technique does not work. The estimate for H-I is quite close to the experimental value, but as you go down the group, the estimate is increasingly further from the experimental value. This led Pauling to the concept of electronegativity, he proposed that not all atoms have equal power to attract electrons to themselves. He also proposed that DD is a measure of the ionicity of the bond X-Y.
Pauling then proposed the following relationship: c(Y) - c(X) = 0.102DD^½ (0.102 is just a constant to convert to KJ mol^-1 because pauling worked in eV (electron volts))

He then came up with a consistent set of electronegativities which predict bond enthalpies quite well. To prove this we will estimate the bond dissociation enthalpy for ICl:
D(Cl₂) = 242 KJ mol^-1,        c(Cl)=3.2
D(I₂ = 151 KJ mol^-1,            c(I)=2.7
0.102DD^½ = c(Cl) - c(I) = 3.2 - 2.7 = 0.5
0.102DD = (0.5)² = 0.25
DD = 0.25/0.102 = 24.1 KJ mol^-1

Problems With Pauling Electronegativities

• Pauling electronegativities are purely empirical, i.e. it works, but we don't know why.

• What is the significance of having to square root DD. Pauling found that he had to use a square root by plotting graphs and getting straight lines

• How do the electronegativities relate to electron distributions in molecules?

Move on to the next page: A brief biography of Linus Pauling