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Electronegativity is an atom's ability to attract bonding electrons. The difference in electronegativity between bonded atoms determines bond polarity, ranging from nonpolar covalent to ionic.
Compare electronegativities from a periodic table to classify bonds and predict molecular behavior.
Confusing electronegativity with electron affinity; thinking all C-H bonds are identical regardless of context.
From periodic trends, you know that atomic properties change systematically across periods and down groups. Electronegativity is one of the most consequential of these trends: it measures how strongly an atom attracts the electrons in a covalent bond toward itself. On the Pauling scale, fluorine is the most electronegative element (4.0), and electronegativity generally increases going up and to the right on the periodic table — the same direction as ionization energy, and for the same fundamental reason: smaller atoms with more protons hold their electrons more tightly.
When two atoms with different electronegativities form a covalent bond, the shared electrons are not shared equally. The more electronegative atom pulls the electron density toward itself, creating an uneven distribution of charge. This produces a polar covalent bond — a bond with a partial negative charge (δ−) on the more electronegative atom and a partial positive charge (δ+) on the less electronegative one. Think of it as a tug-of-war for electrons: if both sides pull equally, the rope stays centered (nonpolar); if one side is stronger, the rope shifts toward that side (polar).
The electronegativity difference (ΔEN) between bonded atoms provides a rough guide to bond character. When ΔEN is close to zero (typically < 0.4), the bond is essentially nonpolar covalent — as in H₂ or Cl₂, where identical atoms share electrons equally. When ΔEN is moderate (roughly 0.4 to 1.7), the bond is polar covalent — as in H–Cl (ΔEN = 0.9), where chlorine's greater electronegativity pulls electron density away from hydrogen. When ΔEN is large (typically > 1.7), the electron transfer is so lopsided that the bond is effectively ionic — as in NaCl (ΔEN = 2.1), where sodium essentially surrenders its electron to chlorine entirely.
These boundaries are guidelines, not sharp cutoffs — bond polarity exists on a continuous spectrum from purely covalent to purely ionic. What matters practically is that bond polarity determines much of a molecule's chemical behavior. Polar bonds create sites of partial charge that attract other polar molecules, influence reaction mechanisms by making certain atoms more susceptible to attack, and ultimately determine whether a molecule as a whole is polar — which you will explore when you study molecular polarity. The simple act of looking up two electronegativity values and taking their difference gives you predictive power over a molecule's bonding character, solubility behavior, and reactivity patterns.