Polarity of Bonds

Partial ionic character in covalent bonds. When a covalent bond forms between two IDENTICAL atoms (as in H₂, O₂, Cl₂), both atoms have an equal tendency to attract the shared electron pair, so the shared pair sits exactly midway between the two nuclei -- such a bond is purely (100%) covalent, non-polar. But when the two combining atoms have DIFFERENT electronegativities, the more electronegative atom attracts the shared pair more strongly toward itself, distorting the electron cloud. Consider H-F: the Pauling electronegativities of hydrogen and fluorine are 2.1 and 4.0 respectively, so fluorine attracts the shared pair roughly TWICE as strongly as hydrogen does -- fluorine develops a partial NEGATIVE charge (δ−) and hydrogen a partial POSITIVE charge (δ+). Such a bond is called a polar covalent bond. In H-F specifically, this small, equal-and-opposite charge separation, spread over the H-F bond distance (91 pm), constitutes what is called a dipole.

Dipole moment. The polarity of a covalent bond is measured by its dipole moment, defined as:

where $\mu$ is the dipole moment, $q$ is the magnitude of the (partial) charge, and $2d$ is the distance separating the two charges (Fig 10.15). Dipole moment is a VECTOR quantity; the figures in this chapter (e.g. Fig 10.16, water's dipole) draw the vector pointing from the positive end of the bond toward the negative end -- i.e. from the less-electronegative to the more-electronegative atom, in the direction the electron density has shifted.

Units. The SI unit of dipole moment is the coulomb-metre (C m), but it is usually expressed in the practical unit Debye (D), where 1 Debye = 3.336 × 10⁻³⁰ C m.

Nonpolar vs polar molecules. Homonuclear diatomic molecules such as H₂, O₂ and F₂ have ZERO dipole moment (nonpolar molecules, since there is no electronegativity difference to create a charge separation). Heteronuclear molecules such as HF, HCl, CO and NO have a NON-ZERO dipole moment (polar molecules). (Table 10.5 gives measured dipole moments for HF, HCl, H₂O, NH₃ and CHCl₃.)

A polar BOND does not guarantee a polar MOLECULE. This is the crucial subtlety: a molecule's overall dipole moment is the VECTOR SUM of all its individual bond dipoles, and geometric symmetry can make that sum zero even when every individual bond is polar. Linear carbon dioxide is the textbook example: it has two polar C=O bonds, but because the molecule is linear (180°), the two bond-dipole vectors are equal in magnitude and point in exactly opposite directions, so

and CO₂ has zero net dipole moment despite its polar bonds. Water, by contrast, is bent (104.5°, not linear) because oxygen carries two lone pairs, so its two O-H bond dipoles do NOT cancel -- their vector sum gives water a substantial net dipole moment of 1.85 D (Fig 10.16).

Percentage ionic character from electronegativity difference. For a typical polar bond written A(δ−)-B(δ+), the electronegativity difference ($\chi_A - \chi_B$) predicts roughly how much ionic character the bond has: if the difference equals 1.7, the bond has 50% ionic character; if GREATER than 1.7, MORE than 50% ionic character; if LESS than 1.7, LESS than 50% ionic character.

(Try Q5 of the 'Evaluate Yourself' box -- comparing the dipole moments of OCS and CS₂ -- using this vector-sum reasoning.)

Partial covalent character in ionic bonds — the flip side. Just as covalent bonds can carry partial ionic character, IONIC compounds can show partial COVALENT character. Lithium chloride is the classic example: though formally ionic, LiCl is soluble in organic solvents such as ethanol, a property normally associated with covalent compounds. The explanation is a phenomenon called polarisation: in an ionic compound, the positively-charged cation attracts the valence electrons of the neighbouring anion (while repelling the anion's nucleus), distorting the anion's electron cloud so that some electron density drifts toward the cation -- producing partial SHARING of that electron density, i.e. partial covalent character. The cation's ability to cause this distortion is its polarising power (or polarising ability); the anion's tendency to BE distorted is its polarisability.

Fajans' rules describe how much polarisation (and hence covalent character) results:

1. Higher charge, more covalent character. The higher the positive charge on the cation, the greater its pull on the anion's electron cloud; likewise, the higher the negative charge on the anion, the greater its own polarisability. So increasing the charge on EITHER ion increases covalent character. Example: since cationic charge increases Na⁺ < Mg²⁺ < Al³⁺, covalent character increases in the same order, NaCl < MgCl₂ < AlCl₃. …