The Pauling Electronegativity Scale: Part 2, Inspired by Biology

Linus Pauling’s electronegativity scale was inspired by Biology. In the early 1930s, chromosomal genes were being mapped out by measuring how frequently two independent traits were inherited together; the idea being that the closer the genes were, the more likely that they would stay linked during genetic crossover.

Pauling tested this idea with chemical compounds, finding that bonds between similar elements were not as strong as bonds between dissimilar elements. He attributed this discovery to ionic contributions in the stronger bonds, and correlated the ionic nature of certain elements with further spreads on his electronegativity scale. For example, the bond between Lithium and Fluorine was almost one-hundred percent ionic – therefore, he placed Lithium on one end of his scale and Fluorine on the other end.

From there, Pauling assigned arbitrary values for each known element based upon their position on the accepted ‘map’ of ionic bond proclivities. Later, he explained that his calculation of each electronegativity value was an estimate of the covalent contribution to an element’s bond subtracted from the actual bond energy, as per the following formula:


In this formula,  Δ is a measure of excess ionic energy – the value that Pauling used to arbitrarily assign electronegativity values to elements. Again, the higher the ionic bond energy measured within an element, the more electronegative the element was to be considered.

In terms of chemistry, Pauling’s electronegativity scale was one of his least theoretically well-founded theories.  On the very same token, it was also one of his most influential ideas in that it allowed chemists to make assumptions about bonds and molecules that could give rise to new interesting and useful correlations.

Indeed, Pauling’s electronegativity scale was very practical.  He used electronegativity to explain chemical bonding characteristics, including the changes in the energy of atoms that occur as electrons rearrange their placement in the atoms’ orbitals. By comparing these values, researchers could predict the properties of a given bond without ever needing to know the bond’s complicated wave equation from quantum mechanics.

Pauling’s faith in his scale was such that he used it to theorize that Fluorine was so electronegative, it would form compounds with an inert gas – something that, at the time, was thought to be impossible. Inert gasses simply did not bond.  However, he couldn’t prove the relationship, and it frustrated Pauling. Eventually though, some thirty years later, he was proven correct by another team of scientists. In their discussions of Pauling’s “stochastic method,” biographers have shown that much of Pauling’s research followed the Fluorine example:  more often that not, his intuition about chemical systems was correct, despite his inability to empirically prove his ideas with hard data.

Electronegativity data, 1930s.

Electronegativity data, 1930s.

In 1932, Linus Pauling published his original paper proposing a thermochemical method of assigning relative electronegativity values. He applied his system to ten nonmetallic elements.  As with Berzelius’ earlier attempts at developing an electronegativity scale, Pauling failed to clearly define how he established his proposed values.  (For the contemporary student, his later calculations regarding electronegativity are contained in his 37th research journal, though to most readers the journal is difficult to follow due to it being more of a stream-of-conscious study as opposed to a series of well-explained experimental argument.)

All of this noted, and despite many additional attempts at determining a rigorous electronegativity scale in the years following his work, Pauling’s 1932 scale is still the one most-commonly in use today.

Learn more about electronegativity on the website Linus Pauling and the Nature of the Chemical Bond, available at the Linus Pauling Online portal.

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