Electronegativity

Electronegativity is a dimensionless measure of an atom's tendency to attract electrons toward itself in a bond. Unlike electron affinity, which is a measurable physical quantity expressed in kJ/mol representing energy released when an isolated gas-phase atom gains an electron, electronegativity is calculated rather than measured. Linus Pauling derived the first electronegativity scale using thermochemical data from bond energies.

Electronegativity differences between bonded atoms determine how shared electrons distribute. In polar covalent bonds, electrons shift toward the more electronegative atom, creating partial charges. The greater the electronegativity difference, the more polarized the electron distribution and the larger the partial charges. This relationship between bond polarity dipole moment and percent ionic character helps predict bond characteristics.

In HCl, electrons are unequally shared between hydrogen and chlorine rather than transferred completely. Chlorine's higher electronegativity attracts the shared electrons more strongly, creating higher electron density on chlorine and a partial negative charge, while hydrogen carries a partial positive charge. This unequal sharing defines a polar covalent bond, distinct from ionic bonding where electrons transfer entirely from metals to nonmetals.

Electronegativity increases from left to right across a period and decreases down a group in the periodic table. Nonmetals in the upper right have the highest electronegativities, with fluorine being the most electronegative element at 3.98. Metals are less electronegative, and group 1 metals have the lowest values. Electronegativity decreases with increasing atomic size because larger atoms attract electrons less effectively.

Pauling studied bond energies in heteronuclear molecules like HF and compared them to predicted values based on homonuclear bond energies. He discovered that HF's experimental bond energy (565 kJ/mol) exceeded predictions, suggesting ionic character in the bond. To explain this discrepancy, Pauling developed an electronegativity scale ranging from 0 to 4, based on thermochemical data from bond breaking energies.

Noble gases are excluded because they typically do not share electrons with other atoms due to having a full valence shell. Since electronegativity measures an atom's tendency to attract electrons in bonds, and noble gases rarely form bonds under normal conditions, assigning them electronegativity values is not practical. Although rare noble gas compounds exist under extreme conditions, they do not fit the general electronegativity model.

When gaseous HCl is placed in an electric field, hydrogen orients toward the cathode (negative pole) and chlorine toward the anode (positive pole), indicating partial charges on each atom. This behavior demonstrates that HCl is a polar molecule with unequal electron distribution. In contrast, nonpolar molecules like N2 orient equally between poles, showing no preferential charge distribution.