Electronegativity is the ability of an atom to attract a shared pair of electrons toward itself within a chemical bond. It is quantified through comparative scales, as this tendency is not directly measurable. Determining these values allows chemists to predict the distribution of electron density in a molecule, which governs its physical and chemical properties.
The Pauling Scale
The Pauling scale, introduced by Linus Pauling, is the original and most commonly referenced method for quantifying electronegativity. This approach is comparative, relying on the measurement of energy differences observed in chemical bonds.
The scale is based on the principle that if two atoms, A and B, form a bond, the energy required to break that bond (\(E_{A-B}\)) is often greater than the energy expected for a purely covalent bond. The difference between the measured bond dissociation energy and the energy of a purely covalent bond (calculated as the average of the energies for the A-A and B-B bonds) is attributed to the partial ionic character of the A-B bond. This extra energy, known as the ionic-covalent resonance energy, is directly related to the difference in the electronegativity values (\(\chi_A – \chi_B\)). Pauling arbitrarily set the electronegativity of Fluorine, the most electron-attracting element, at 4.0, which anchored the entire scale.
The Mulliken Scale
The Mulliken scale offers an alternative, more absolute approach to determining electronegativity by focusing on the intrinsic properties of a single, isolated atom. Developed by Robert S. Mulliken, this scale defines electronegativity as the average of an atom’s ionization energy (IE) and its electron affinity (EA).
Ionization energy is the energy required to remove an electron from a neutral atom, while electron affinity is the energy released when an electron is added to a neutral atom. This method is derived from fundamental, measurable atomic properties rather than the bond energies of molecules. The Mulliken values offer a direct physical interpretation of an atom’s electron-holding and electron-gaining tendencies.
The Allred-Rochow Scale
The Allred-Rochow scale bases its calculation on classical physics and the concept of electrostatic force. This method conceptualizes electronegativity as the force exerted by the nucleus on a valence electron at the atom’s boundary.
The scale’s calculation incorporates two physical parameters: the effective nuclear charge (\(Z_{eff}\)) and the covalent radius of the atom. The effective nuclear charge is the net positive charge experienced by the outermost electrons, accounting for the shielding effect of inner electrons. Electronegativity increases as the effective nuclear charge rises and decreases as the covalent radius increases, linking electronegativity directly to atomic size and nuclear charge.
Using Determined Values to Predict Bonding
The primary application of determined electronegativity values is predicting the nature of the chemical bond formed between two atoms. This prediction is made by calculating the difference (\(\Delta EN\)) between the electronegativity values of the two bonded atoms. A zero or very small difference indicates that the electrons are shared almost equally, resulting in a nonpolar covalent bond.
As the difference increases, the sharing of electrons becomes unequal, creating a polar covalent bond where the electrons spend more time near the more electronegative atom. For instance, the difference between Carbon (\(\sim\)2.5) and Hydrogen (\(\sim\)2.1) is small, leading to a nonpolar bond in hydrocarbons. A moderate difference, often cited between 0.4 and 1.7, signifies a bond with significant polarity, such as the bond in water between Oxygen and Hydrogen. A large electronegativity difference, typically greater than 1.7, indicates that the more electronegative atom has essentially stripped the electron away from the other, forming an ionic bond. This happens when a metal with low electronegativity, like Sodium (\(\sim\)0.9), bonds with a nonmetal with high electronegativity, like Chlorine (\(\sim\)3.0).