Not all bonds are purely 'ionic' or purely 'covalent' — real bonds fall on a spectrum. Picture every bond as a tug-of-war over a pair of electrons. When the two atoms are identical, as in , they pull equally and the pair sits exactly in the middle. When one atom pulls a little harder, as in HCl, the pair is still shared but drifts toward the stronger puller. When one atom pulls far harder, as in NaCl, the weaker atom loses the contest completely and the electron is transferred, making ions. The tool for locating a bond on that spectrum is electronegativity difference, : the absolute difference between the electronegativities of the two bonded atoms, using the Pauling scale values you met in the periodic trends topic. One subtraction tells you whether a substance is likely to be a brittle, high-melting ionic solid or a molecule, and whether that molecule's bonds carry partial charges that will later control its shape, polarity, boiling point and solubility.
What electronegativity measures, and why it varies
Electronegativity is the ability of an atom in a bond to attract the shared pair of electrons toward itself. It has no units: Linus Pauling built a relative scale on which fluorine, the strongest puller, is about 4.0 and the alkali metals are below 1. The trend follows directly from the forces inside the atom. Moving right across a period, the nuclear charge grows while the electrons being added stay in the same energy level, so the valence shell is pulled closer and harder: electronegativity increases. Moving down a group, each new energy level puts the valence electrons farther from the nucleus and behind more inner electrons that shield them, so the nucleus's grip on a bonding pair weakens: electronegativity decreases. That is why the strongest pullers (F, O, N, Cl) sit at the top right, and the weakest (Cs, K, Na) sit at the bottom left. Noble gases are usually left off the scale because they rarely form bonds at all.
Calculating and classifying ΔEN
| ΔEN range | Bond type | Electron distribution | Example (ΔEN) |
|---|---|---|---|
| 0.0 – 0.4 | Nonpolar covalent | Electrons shared essentially equally | Cl–Cl (0.00), C–H (0.35) |
| 0.4 – 1.7 | Polar covalent | Electrons shared unequally; a partial charge () develops on each atom | H–Cl (0.96), O–H (1.24) |
| ≥ 1.7 | Ionic | Electrons effectively transferred, not shared | Na–Cl (2.23), K–F (3.16) |
- Look up the electronegativity of each bonded atom.
- Subtract the smaller value from the larger one, so is always positive.
- Compare the result with the cutoffs 0.4 and 1.7 to name the bond type.
- For a polar covalent bond, label the more electronegative atom and the other .
- Sanity-check against what you know: a metal with a nonmetal is usually ionic, two nonmetals are usually covalent. If ΔEN disagrees with that, say so and explain.
In a polar covalent bond, the more electronegative atom pulls the shared electron pair closer to itself, gaining a small negative charge () while the other atom becomes slightly positive (). The δ (delta) means "partial": the charges are fractions of an electron's charge, not the full and of real ions. This partial charge separation is called a bond dipole, and it's the building block of molecular polarity (covered in the next topic). A bond dipole is often drawn as an arrow along the bond pointing toward the end, with a small cross on the tail at the end. The bigger the ΔEN, the bigger the partial charges and the more polar the bond, so you can rank bonds by polarity just by ranking their ΔEN values.
- Electron pair is still shared, but lopsided
- Partial charges only: δ+ on H, δ− on Cl
- Atoms stay part of one molecule
- Pure HCl is a gas at room temperature
- Electron transferred from Na to Cl
- Full charges: Na⁺ and Cl⁻
- Ions pack into a crystal lattice
- Solid NaCl melts at 801 °C
Coordinate covalent (dative) bonds
In an ordinary covalent bond, each atom contributes one electron to the shared pair. In a coordinate covalent bond, both electrons of the shared pair come from the same atom — usually one with a lone pair to spare, donated to an atom (or ion) that needs one more bond to complete its octet.
How do you recognise when a coordinate covalent bond must be present? Look for a species with a lone pair (NH₃, H₂O) meeting a species with no electrons of its own to share (a bare ion has none). The only way a bond can form is if the donor supplies both electrons. Coordinate bonds are also why carbon monoxide can have a triple bond: oxygen donates one of its lone pairs to carbon to give both atoms an octet.