Electronegativity is a fundamental property in chemistry that describes an atom’s ability to attract a shared pair of electrons toward itself when participating in a chemical bond. This property is essential for understanding the nature of chemical substances, particularly those involving nonmetals. It provides a framework for predicting the distribution of electron density within any given molecule.
Understanding Electronegativity
Electronegativity is not a directly measured physical property; instead, it is a calculated value on a relative scale. The most common system is the Pauling scale, introduced in 1932 and based on the energy required to break chemical bonds. This scale assigns a dimensionless number to each element, typically ranging from 0.7 (Francium) to 3.98 (Fluorine). A higher number indicates a stronger attraction for shared electrons.
The Pauling scale allows chemists to quantify the electron-attracting tendencies of different elements relative to one another. The values determined using this scale are helpful in estimating the polarity, or charge separation, that exists within a chemical bond. This measure predicts how the electrons will be distributed between two bonded atoms.
Nonmetals and Their Position on the Periodic Table
Nonmetals are generally located on the upper right side of the periodic table, with the exception of hydrogen. This position is directly correlated with their high electronegativity values. Electronegativity follows predictable trends across the periodic table, increasing as you move from left to right across a period. This increase is due to a rising number of protons in the nucleus, which creates a stronger positive charge that more effectively attracts the electrons in the outermost shell.
Electronegativity decreases as you move down a group. This decrease is a result of increasing atomic radius, where the outermost valence electrons are farther from the attractive positive nucleus and are shielded by inner electron shells. Consequently, elements in the top-right corner, such as Oxygen, Chlorine, and especially Fluorine, possess the highest electronegativity values. These nonmetals have valence shells that are nearly full, giving them a strong tendency to attract or gain the few extra electrons needed to achieve a stable, noble gas configuration.
Electronegativity and Chemical Bonding
The difference in electronegativity (\(\Delta EN\)) between two bonded atoms determines the nature of the chemical bond formed between them. When a nonmetal with a high electronegativity bonds with a metal that has a very low electronegativity, the difference is substantial. For instance, a \(\Delta EN\) greater than approximately 1.7 to 2.0 typically results in an ionic bond, where the nonmetal completely strips the electron from the metal.
When two nonmetals bond, the difference in their electronegativity is usually small to moderate. If the \(\Delta EN\) is near zero, the electrons are shared almost equally, forming a nonpolar covalent bond. However, a moderate difference, such as between 0.5 and 1.6, means the electrons are shared unequally, creating a polar covalent bond. In a polar covalent bond, the higher electronegativity nonmetal pulls the electron pair closer to itself, resulting in a partial negative charge on that atom and a partial positive charge on the other.