Bond Properties

In periodic trends, we introduced electronegativity (χ) as the tendency of an atom to attract shared electrons toward itself. When two atoms form a bond, the electronegativity difference (Δχ) between them determines how the bonding electrons are distributed. This distribution is not an all-or-nothing choice between “shared” and “transferred.” It is a continuum from purely covalent to ionic.

Two identical atoms share electrons equally (pure covalent bond, Δχ = 0). A metal bonded to a nonmetal effectively transfers electrons (ionic bond, large Δχ). Most bonds fall somewhere in between, with electron density shifted toward the more electronegative atom.

Classifying Bonds by Δχ

Most textbooks classify bonds using Δχ thresholds:

  • Δχ < 0.4: Nonpolar covalent. Electrons are shared nearly equally.
  • 0.4 ≤ Δχ < 1.7: Polar covalent. Electrons are shifted toward the more electronegative atom.
  • Δχ ≥ 1.7: Ionic. Electrons are transferred.

These boundaries are approximate guidelines, not sharp dividers. The 0.4 boundary corresponds to ~4 % ionic character, and the 1.7 boundary corresponds to ~50 % ionic character. Some textbooks use 0.5/2.0 or 0.4/1.8 instead.

Visualizing Electron Cloud Deformation

The interactive visualization below shows how the electron cloud deforms as electronegativity difference increases, from symmetric (nonpolar) to skewed (polar) to fully separated (ionic).

Pauling’s Ionic Character Formula

Where do the 0.4 and 1.7 boundaries come from? Pauling’s empirical formula, derived from dipole moment measurements of gas-phase diatomic molecules, quantifies the percent ionic character of a bond:

\[\%\,\text{ionic character} = 100\% \times \left(1 - e^{-(\Delta\chi/2)^2}\right)\]

As Δχ increases, the percent ionic character rises smoothly toward 100%:

Table 1: Landmark values from Pauling’s ionic character formula

The value Δχ = 1.7 gives approximately 50 % ionic character, which is why many textbooks use it as the dividing line between “polar covalent” and “ionic.” The exact Δχ for 50 % is √(4 ln 2) ≈ 1.665. The commonly used 1.7 is a convenient round number.

Ionic Character Across Real Compounds

The plot below shows the Pauling curve with compounds placed according to their theoretical Pauling electronegativity differences. Hover over any point to see details.

HF and LiI sit close together on the x-axis (Δχ ≈ 1.7–1.8) but behave very differently.

WarningWhen Δχ Gets It Wrong: HF and LiI

The Pauling formula predicts similar ionic character for HF (Δχ = 1.78, ~55 %) and LiI (Δχ = 1.68, ~51 %). Their properties diverge:

HF is a molecular (covalent) compound:

  • Gas at room temperature (boiling point: −19.5 °C)
  • Exists as discrete HF molecules, not a crystal lattice
  • Dissolves in water as hydrofluoric acid (a weak acid, pKa = 3.17)
  • Both atoms are small nonmetals, so strong nuclear attraction prevents full electron transfer

LiI is an ionic compound:

  • Forms a NaCl-type face-centered cubic crystal lattice
  • Melting point: 469 °C, boiling point: 1171 °C
  • Fully dissociates in water into Li+ and I (strong electrolyte)
  • Conducts electricity when molten
  • Metal + nonmetal combination

HF exceeds the 1.7 threshold but behaves as covalent. LiI falls below it but behaves as ionic. The Δχ classification fails in both directions. Atom size, electron configuration, and whether the bond is between a metal and a nonmetal are often more reliable predictors than Δχ alone.

Bond Order

The bond order is the number of electron pairs shared between two atoms. A single bond has bond order 1, a double bond has bond order 2, and a triple bond has bond order 3.

For molecules with resonance structures, bond order can be fractional. It is calculated as:

\[\mathrm{bond~order} = \frac{\mathrm{total~bonding~pairs~between~two~atoms~across~all~resonance~structures}}{\mathrm{number~of~resonance~structures}}\]

For example, ozone has two resonance structures. In one, the left O–O bond is a double bond (2 pairs) and in the other it is a single bond (1 pair). The bond order is (2 + 1) / 2 = 1.5. Similarly, each C–O bond in carbonate has a bond order of (2 + 1 + 1) / 3 = 4/3.

Bond Length

Bond length is the distance between the nuclei of two bonded atoms. It is directly related to bond order:

Triple bonds are shorter than double bonds, which are shorter than single bonds.

More shared electron density between two nuclei pulls them closer together.

Bond Energy (Bond Dissociation Energy)

Bond energy (or bond dissociation energy) is the energy required to break one mole of a particular bond in the gas phase. The interactive plot below shows the potential energy of H2 as a function of internuclear distance. At the equilibrium bond length (0.74 Å), the energy is at a minimum. Pulling the atoms apart requires 432 kJ/mol of energy.

Bond energy follows the same trend as bond order:

Triple bonds are stronger than double bonds, which are stronger than single bonds.

More shared electrons means a stronger attraction between the nuclei and the bonding electrons, requiring more energy to break the bond.

Note

Bond energies are averages. The actual energy required to break a specific bond depends on the molecular context. For example, the C–H bond energy in methane (CH4) is slightly different from the C–H bond energy in ethane (C2H6).