2.1 Bond Types

Bond Type Spectrum

Pick two elements from a live electronegativity map and watch a real electron density map redraw itself. Covalent, ionic, and metallic bonding, named by the metal and nonmetal rule rather than a ΔEN cutoff.

Electron DensityΔEN ContinuumIonic vs CovalentMetallic Bonding
Topic 2.1

Types of Chemical Bonds

Explain the relationship between the type of bonding and the properties of the elements participating in the bond.

Electronegativity is an atom's pull on the electrons in a bond. For the representative elements it increases left to right across a period and decreases down a group — for exactly the reasons from Unit 1: rising effective nuclear charge across, added shells and shielding down.

The difference in electronegativity between two bonded atoms tells you how the electrons are distributed:

  • Nonpolar covalent — atoms of similar electronegativity share equally. Cl–Cl is the obvious case, but the CED singles out C–H as effectively nonpolar even though carbon is slightly more electronegative.
  • Polar covalent — unequal sharing. The more electronegative atom develops a partial negative charge (δ−) and the other a δ+. In single bonds, a greater electronegativity difference means a greater bond dipole.
  • Ionic — the transfer end of the same continuum.

The single most important idea here is EK 2.1.A.3(iii): all polar bonds have some ionic character, and the difference between ionic and covalent bonding is not a sharp boundary but a continuum. There is no magic ΔEN value that flips a bond from covalent to ionic.

That said, EK 2.1.A.4 gives practical guidance: electronegativity difference is not the only factor. Generally, metal + nonmetal → ionic, and nonmetal + nonmetal → covalent. And the definitive test is not a number at all — it is examining the properties of the compound (melting point, conductivity, brittleness).

Metallic bonding is the third type: in a metallic solid the valence electrons are delocalized, belonging to the whole lattice rather than to any individual atom.

Key points

  • Ionic and covalent are the two ends of one continuum; every polar bond has partial ionic character.
  • C–H bonds are treated as nonpolar on the AP Exam despite a small electronegativity difference.
  • The best way to classify bonding in an unknown compound is to look at its physical properties, not to compute ΔEN.
  • In metals, valence electrons are delocalized across the entire lattice — not shared between two specific atoms.

Common mistakes

  • Do not quote a ΔEN cutoff (like 1.7) as if it were a rule. The CED explicitly frames bonding as a continuum, and rubrics reward that language.
  • Polar bond ≠ polar molecule. CO₂ has two very polar bonds and zero net dipole because the geometry cancels them.
  • Electronegativity is not electron affinity. Electronegativity describes an atom in a bond; electron affinity is a measured energy change for an isolated gaseous atom.
  • Metallic bonding is not "sharing between two atoms". It is delocalization over the whole solid.

Worked example

Rank the bonds C–F, C–O, C–C, and C–H from most to least polar, and explain your ranking.

Electronegativity values: F (4.0) > O (3.5) > C (2.5) ≈ H (2.1).

C–F (ΔEN 1.5) > C–O (ΔEN 1.0) > C–H (ΔEN 0.4) > C–C (ΔEN 0)

Bond polarity scales with the electronegativity difference between the bonded atoms, because a larger difference means a more uneven electron distribution and a larger partial charge separation. C–C has identical atoms, so the sharing is perfectly equal and the bond is nonpolar. C–H is small enough that the AP course treats it as effectively nonpolar.

Full notes for topic 2.1 →