Electronegativity is a fundamental concept in chemistry: an atom’s ability to attract electrons when forming a chemical bond. It is often described as an element’s “electron greed.” Contrary to the idea of a low value, nonmetals generally exhibit remarkably high electronegativity. This strong electron-attracting power is a defining characteristic of nonmetallic elements. Understanding this property requires a closer look at how this tendency is measured and the atomic structures that drive it.
What Electronegativity Measures
Electronegativity quantifies an atom’s tendency to attract the electrons involved in a chemical bond. It is not a measure of energy like ionization energy or electron affinity, but rather a relative measure of attraction. The most widely used system is the Pauling scale, developed by chemist Linus Pauling. This scale assigns a dimensionless number to each element, providing a standardized way to compare their electron-attracting abilities. Values typically range from about 0.7 for Francium to 4.0 for Fluorine. The difference in the electronegativity values between two bonded atoms predicts the nature of the bond they will form.
Nonmetals and High Electron Affinity
Nonmetals are situated on the right side of the periodic table, excluding the noble gases, and their atomic structure is the direct cause of their high electronegativity. These elements typically possess valence shells that are nearly full, often having five, six, or seven valence electrons. This configuration means the atoms are just a few electrons short of achieving a stable, full octet.
This inherent need for a small number of electrons creates a powerful pull on any shared electrons from an adjacent atom. The nonmetal’s nucleus has a strong effective positive charge acting on the valence shell, which is only minimally shielded by the inner electrons. Consequently, nonmetals have a strong tendency to attract electrons to complete their outer shell. This strong attraction is why nonmetals tend to gain electrons and form negative ions, or anions, when they react with metals.
How Periodic Trends Determine Electronegativity
The placement of an element on the periodic table systematically dictates its electronegativity through predictable trends. Electronegativity values generally increase as one moves from left to right across any given period. This trend occurs because atoms across a period experience an increasing nuclear charge, while their valence electrons remain in the same principal energy level. The greater positive charge in the nucleus pulls the electron cloud more tightly, increasing the atom’s ability to attract bonding electrons.
Conversely, moving down a group causes electronegativity to decrease. This reduction is a consequence of increasing atomic size and the shielding effect. As one moves down a group, each successive element adds a new electron shell, placing the valence electrons further from the attractive force of the nucleus. The inner electron shells effectively “shield” the valence electrons from the full nuclear charge, weakening the atom’s pull on bonding electrons. The combination of these two trends results in the most electronegative elements, the nonmetals, being clustered in the upper-right corner of the periodic table, closest to Fluorine.
The Role of Electronegativity in Chemical Bonding
The disparity in electronegativity between two atoms determines the fundamental nature and polarity of the chemical bond that forms between them. When the difference in electronegativity (\(\Delta EN\)) is relatively small (less than 0.4), the electrons are shared almost equally, resulting in a nonpolar covalent bond. This type of bond often occurs between two identical nonmetal atoms, like in a molecule of oxygen gas.
If the electronegativity difference is moderate (ranging from 0.4 to approximately 1.7), the bond is classified as polar covalent. In this case, the shared electron pair is pulled more strongly toward the nonmetal atom with the higher electronegativity, creating partial negative and partial positive charges on the respective atoms. A large difference in electronegativity (exceeding 1.7) indicates that the electron is essentially transferred from the less electronegative atom (typically a metal) to the highly electronegative nonmetal. This transfer results in the formation of distinct positive and negative ions held together by electrostatic attraction, defining an ionic bond.