Electronegativity is a fundamental concept in chemistry, representing an atom’s inherent tendency to attract shared electrons towards itself when forming a chemical bond. It provides insight into the distribution of electron density within a bond, which influences a molecule’s overall behavior and reactivity.
Understanding Electronegativity
Electronegativity quantifies an atom’s “pulling power” on the electrons it shares in a bond. This ability is primarily influenced by two atomic characteristics: nuclear charge and atomic radius. A greater nuclear charge, meaning more protons in the nucleus, results in a stronger positive attraction for electrons.
Conversely, a larger atomic radius, which indicates that an atom’s valence electrons are further from the nucleus, diminishes this attractive force. The increased distance and the shielding effect from inner electrons reduce the nucleus’s pull on the shared electrons. Therefore, atoms with smaller radii and higher nuclear charges tend to exhibit greater electronegativity.
Locating Electronegativity Values
Electronegativity values for elements are commonly found on periodic tables or in chemistry reference charts. These values are not directly measured but are calculated from other atomic or molecular properties.
Electronegativity generally increases as you move from left to right across a period. This is because elements across a period have an increasing nuclear charge while their valence electrons reside in the same primary energy level, leading to a stronger pull on bonding electrons. Conversely, electronegativity typically decreases as you move down a group. Elements further down a group have more electron shells, increasing their atomic radius and the shielding of outer electrons from the nucleus, thus weakening the nuclear attraction for shared electrons. Fluorine, located in the upper right, is the most electronegative element, while francium and cesium, in the lower left, are among the least electronegative.
Different Electronegativity Scales
Scientists have developed several scales to quantify electronegativity, each based on different theoretical or experimental approaches. The most widely used is the Pauling scale, which assigns a dimensionless value to elements based on bond energies. Linus Pauling arbitrarily set fluorine’s electronegativity at 4.0, and other elements are scaled relative to this.
Another approach is the Mulliken scale, which calculates electronegativity as the average of an atom’s ionization energy and electron affinity. The Allred-Rochow scale, in contrast, bases its values on the effective nuclear charge and the covalent radius of an atom. While the absolute numerical values may differ between these scales, they consistently show the same general periodic trends in electronegativity.
Applying Electronegativity Values
The difference in electronegativity between two bonded atoms determines the type of bond that forms: nonpolar covalent, polar covalent, or ionic. When this difference is very small (typically less than 0.5), electrons are shared relatively equally, forming a nonpolar covalent bond. A moderate difference (usually between 0.5 and 1.7) indicates unequal sharing, leading to a polar covalent bond with partial charges. If the difference is large (generally 1.7 or greater), electrons are essentially transferred, forming an ionic bond. These predictions help understand molecular polarity and reactivity.