Electronegativity is an atom’s ability to attract shared electrons in a chemical bond. It helps explain why some bonds share electrons evenly while others pull electron density toward one atom. This idea matters because bond polarity affects molecular shape, solubility, boiling point, and many chemical reactions.
A small difference in electron attraction can change how a molecule behaves in water, in cells, or in a lab reaction.
When two atoms form a bond, the more electronegative atom pulls the shared electron cloud closer to itself. This gives that atom a partial negative charge, written δ-, while the other atom gets a partial positive charge, written δ+. The electronegativity difference, often written ΔEN, is used to estimate whether a bond is nonpolar covalent, polar covalent, or ionic.
Across the periodic table, electronegativity generally increases from left to right and decreases from top to bottom.
Key Facts
- Electronegativity is an atom’s ability to attract shared electrons in a chemical bond.
- ΔEN = |EN atom 1 - EN atom 2| measures the electronegativity difference between two bonded atoms.
- A nonpolar covalent bond usually has ΔEN from 0 to about 0.4, so electrons are shared nearly equally.
- A polar covalent bond usually has ΔEN from about 0.5 to 1.7, so electrons are shared unequally.
- An ionic bond usually has ΔEN greater than about 1.7, so electron transfer becomes a useful model.
- Electronegativity generally increases left to right across a period and decreases down a group.
Vocabulary
- Electronegativity
- Electronegativity is the ability of an atom to attract shared electrons in a chemical bond.
- Bond polarity
- Bond polarity is the uneven distribution of electron density between two bonded atoms.
- Partial charge
- A partial charge is a small charge imbalance in a polar bond, shown as δ+ or δ-.
- Dipole
- A dipole is a separation of positive and negative charge within a bond or molecule.
- Electron density
- Electron density describes where electrons are most likely to be found around atoms or within a bond.
Common Mistakes to Avoid
- Treating all covalent bonds as nonpolar is wrong because many covalent bonds have unequal electron sharing when the atoms have different electronegativities.
- Pointing the dipole arrow toward the less electronegative atom is wrong because electron density is pulled toward the more electronegative atom.
- Confusing partial charges with full ionic charges is wrong because δ+ and δ- show unequal sharing, not complete electron transfer.
- Using electronegativity trends backward is wrong because electronegativity generally increases across a period from left to right and decreases down a group.
Practice Questions
- 1 The electronegativity of H is 2.20 and the electronegativity of Cl is 3.16. Calculate ΔEN for the H-Cl bond and classify the bond as nonpolar covalent, polar covalent, or ionic.
- 2 The electronegativity of C is 2.55 and the electronegativity of O is 3.44. Calculate ΔEN for a C-O bond and identify which atom has δ-.
- 3 A molecule has two identical polar bonds arranged in opposite directions in a straight line. Explain why the molecule may be nonpolar even though each bond is polar.