Electronegativity measures how hard an atom pulls on shared electrons. When the two atoms in a covalent bond pull unequally, the shared pair spends more time near one of them.
Key idea
The electronegativity difference $\Delta EN$ tells how unevenly a bond is shared. A common guide: below 0.4, nonpolar covalent; 0.4 to 1.8, polar covalent; above 1.8, mostly ionic. In a polar bond the more electronegative atom carries a partial negative charge ($\delta^-$) and the other a partial positive charge ($\delta^+$).
Why it works
Ionic and covalent bonds are two ends of one range. In $\mathrm{Cl_2}$ the two atoms are identical, so the pair is shared evenly. In H–Cl, chlorine (3.16) pulls harder than hydrogen (2.20), so electron density shifts toward Cl, leaving it slightly negative and H slightly positive. In NaCl the difference (2.23) is so large that the electron has practically moved over: ions. The cutoffs 0.4 and 1.8 are only signposts on that smooth range.
Worked examples
Worked example: Classifying
Classify the O–H bond (O 3.44, H 2.20).
- $\Delta EN = 3.44 - 2.20 = 1.24$
- Between 0.4 and 1.8: polar covalent.
- Oxygen is more electronegative: O is $\delta^-$, H is $\delta^+$.
Answer: Polar covalent, with $\delta^-$ on O.
Worked example: Ranking
Rank C–H, C–O and C–F from least to most polar (C 2.55, H 2.20, O 3.44, F 3.98).
- C–H: $0.35$. C–O: $0.89$. C–F: $1.43$.
- Bigger difference, more polar.
Answer: C–H < C–O < C–F
Common mistake
Treating ionic and covalent bonds as two separate boxes with a sharp cutoff. A bond with $\Delta EN = 1.7$ is not completely covalent while one with 1.9 is completely ionic; both are strongly polar. Use the guide to name a bond, but think of polarity as a smooth scale.