Electronegativity is a fundamental concept in chemistry that describes an atom’s inherent power to attract electrons toward itself when it is part of a chemical bond. This property is not a simple physical measurement but rather a relative value that helps scientists understand how atoms share or transfer electrons. The most recognized and widely applied method for quantifying this tendency is the standardized, numerical scale developed by the American chemist Linus Pauling. This system provides a reliable way to predict the nature of chemical interactions between elements.
Defining the Pauling Scale
The Pauling scale is a relative and dimensionless measure, meaning it does not have standard units like mass or volume, but instead ranks elements against one another. Pauling derived the values from experimental energy measurements of chemical bonds, not by directly measuring the pull of an atom. He observed that the energy required to break a bond between two different atoms (A-B) was often greater than the average energy required to break the bonds between identical atoms (A-A and B-B).
This unexpected extra energy in the A-B bond was attributed to “ionic character,” resulting from the unequal sharing of electrons. Pauling quantified this difference in bond dissociation energies, which is the energy needed to break a chemical bond, and mathematically related it to the atoms’ differing abilities to attract electrons. By performing these calculations on numerous compounds, he assigned a numerical value to each element representing its relative electron-attracting power. The resulting values reflect an atom’s behavior when bonded to another atom.
The Numerical Range and Anchor Points
The Pauling scale spans a range of values from approximately 0.7 to 4.0, providing a clear numerical framework for comparing all elements. The maximum value of 4.0 was arbitrarily assigned to fluorine, which serves as the reference point for the entire scale. Fluorine has the strongest pull on bonding electrons because of its very small atomic size and a high effective nuclear charge, allowing its nucleus to exert a powerful force on shared electrons.
At the opposite end of the spectrum, the lowest values are found among the largest atoms, specifically the alkali metals. Francium (Fr) and Cesium (Cs) possess the lowest electronegativity values, hovering around 0.7. These elements are physically large, meaning their outermost valence electrons are shielded by many inner electron shells and located very far from the nucleus. This distance and shielding result in a weak attraction for shared electrons.
Interpreting Electronegativity Values
The practical utility of the Pauling scale lies in its ability to predict the type of chemical bond that will form between two atoms. This prediction is made by calculating the absolute difference between the electronegativity values of the two bonded atoms (\(\Delta\)EN). A small difference in electronegativity indicates a relatively equal sharing of electrons, resulting in a nonpolar covalent bond. Textbooks often use a difference of less than 0.4 as the threshold for this type of bond.
When the difference in electronegativity is moderate, typically between 0.4 and 1.7, the bond is classified as polar covalent. In this scenario, the electrons are shared unequally, with the atom having the higher Pauling value attracting the electron pair closer to itself and acquiring a partial negative charge.
A large difference in electronegativity, generally exceeding 1.7, signifies that the more electronegative atom has essentially stripped the electron away from the less electronegative atom. This complete transfer of an electron results in the formation of an ionic bond, where the atoms exist as fully charged ions held together by electrostatic attraction. The numerical range of the Pauling scale provides a continuous spectrum that links the electron-sharing nature of covalent bonds to the electron-transferring nature of ionic bonds.