Bond enthalpy terms and electronegativities

It was through bond enthalpy terms that Linus Pauling arrived at the famous Pauling electronegativity scale. In the previous Section, you saw that S(H—Cl) = 432 kJ moFk The homonuclear molecules, H2 and CI2, formed by these two atoms have bond enthalpy terms of 436 and 243 kJ moF, respectively. The average is 340kJ moF, yet fi(H—Cl) exceeds this by 92 kJ moF. This shows that the bond in the HCl molecule is unusually strong. 

Pauling suggested that if the electron pair in the H—Cl bond were shared equally between the hydrogen and chlorine atoms, then B(H—Cl) would indeed be the average of B(H—H) and B(C1—Cl). The additional 92 kJ moF is caused by an 

Pauling therefore argued that, for single bonds between two different atoms, 

A and B, the quantity B(A—B) - 2[(B(A—A) + B(B—B)] is a measure of the electronegativity difference. He called this quantity the ionic resonance energy, and found that he got the best fit if he assumed that it was proportional to the square of the electronegativity difference (Xa Xb) where Xa nd Xb are the electronegativities of A and B  

Here C is a constant. The scale of electronegativities is tied to a fixed point by assigning a value of 96.5 kJ mol to C, and Xh is fixed at 2.1. Thus, rearranging Equation 4.11  

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Understand the relationship between bond enthalpy terms and electronegativity differences embodied in Pauling s electronegativity equation, and calculate one from the other. (Question 4.4)... 

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