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.

The Pauling Electronegativity Scale: Part 1, Historical Background

Linus Pauling lecturing on Amedeo Avogadro, Rome, Italy, June 6, 1956

Linus Pauling lecturing on Amedeo Avogadro, Rome, Italy, June 6, 1956

The development of an accurate electronegativity scale was one of Linus Pauling’s many major contributions to the study of chemistry.  In this two part series, we’ll first look at the electronegativity research that preceded Pauling’s breakthrough, before analyzing the details of the scale that Pauling ultimately derived.

The concept of electronegativity is measured along a relative scale that compares the degree to which atoms of different elements tend to attract electrons from their surrounding environment. Because the electronegativity scale is a qualitative measurement – meaning that there is no measurable constant value for electronegativity – the scale itself has been both difficult and interesting to develop. The electronegativity scale we use today was formalized by Linus Pauling, and was first published in 1932. However, the idea of electronegativity existing between atoms was established well before Pauling, dating back to the early 1800s.

In 1809, Amedeo Avogadro published a paper connecting the correlations between the neutralization that occurs with acids and bases, and the neutralization that occurs between positive and negative electrical charges. Avogadro claimed that these cancellation relationships could be applied to all chemical interactions; between both simple substances and more complex compounds. From this, he proposed the creation of what he termed an “oxygenicity scale” on which every element could be placed – its location dependent upon the element’s tendency to react with other elements – in order to compare the properties of elements that had not yet been tested together.  This was, of course, the forerunner of the modern electronegativity scale.

To determine the relative “oxygenicity” values of elements, Avogadro relied upon contact electrification experiments published by two fellow scientific giants, Humphrey Davy and Alessandro Volta, as well as the work of a German-Danish researcher named Christian Heinrich Pfaff (pdf link).  These experiments found that when two bodies are electrified on contact, the potential between them becomes a value that can be measured.  These sets of values were, in turn, the units that Avogadro used to develop his oxygeniticity scale.

As it turned out, a significant problem with Avogadro’s method is that measures of contact electricity are very easily affected by outside factors, such as moisture or impurities.  As a result, Avogadro’s oxygenicity values turned out to be inconsistent and inaccurate.  Into this void stepped the important Swedish chemist Jöns Jakob Berzelius.

Portrait of Jöns Jacob Berzelius.  Image courtesy of the Michigan State University department of Chemistry.

Portrait of Jöns Jacob Berzelius. Image courtesy of the Michigan State University department of Chemistry.

In 1811, Berzelius published an article detailing his own ideas on electrochemistry. He utilized much of the same groundwork as Avogadro, but, crucially, used the term “electronegativity” instead of “oxygenicity.”

Besides their names, a major difference between the two scales lies in their focus on heat evolution in chemical reactions – while Avogadro never mentions the concept, it is central to Berzelius’ theory, which, indeed, he presented as a new theory of chemical combustion. Berzelius assumed that both heat (or “caloric“, as it was conceived of at the time) and electricity were fluids.  As such, Berzelius attempted to connect heat to his electronegativity scale because he believed that caloric was created by the combination of negative and positive electricity.

Unfortunately for the theory, this assumed connection failed to account for half of all possible chemical reactions (endothermic association and exothermic dissociation), and was eventually discarded in favor of more modern views of the electronegativity scale. However, Berzelius did provide an almost-complete listing of his measured electronegativities, which coordinate remarkably well with both Pauling’s modern thermochemical definition as well as the current Allred-Rochow force definition. Berzelius’ electrochemical theory eventually failed despite its similarities with current systems because it could not account for increasingly complex organic molecules, and was incompatible with Michael Faraday‘s laws of electrolysis – laws that were already generally-accepted during his time.

Learn more about electronegativity on the website Linus Pauling and the Nature of the Chemical Bond, available at the Linus Pauling Online portal.  For more on the early history of electrochemistry, see Dr. Roderick MacKinnon’s lecture “Ion Channel Chemistry: The Electrical System of Life.”