Electronegativity, Nonpolar, and Polar Covalent Bonds

In a nonpolar covalent bond the shared electrons are shared equally because the two atoms have the same or very similar electronegativity, while in a polar covalent bond the electrons are pulled closer to the more electronegative atom, which becomes partially negative (δ−) while the other becomes partially positive (δ+). In brief: the Cl–Cl bond is nonpolar, the H–Cl bond (ΔEN = 0.96) is polar with chlorine δ−, and the dipole moment of HCl is 1.08 debye, which corresponds to about 18% ionic character.

Bonds fall along a continuum from purely covalent to purely ionic, and the polarity of bonds explains the properties of molecules: solubility, boiling point and chemical reactivity. This article explains electronegativity differences, how to classify a bond and how to show polarity, introduces the dipole moment, distinguishes polar bonds from polar molecules, includes data tables, works through calculations and gives a short history.

What happens when electrons are shared unequally?

The electronegativity of an atom measures how strongly it attracts shared electrons. When two atoms of the same element form a bond, as in H₂, O₂ or Cl₂, the electronegativities are equal and the electrons are shared equally. When the atoms are different, the more electronegative atom pulls the electron density toward itself. The result is a polar bond, with a partial negative charge δ− on the more electronegative atom and a partial positive charge δ+ on the other. The partial charges are less than a full unit of charge. A polar bond can be shown with an arrow that points toward the more electronegative atom, with a small plus sign at the tail.

How do you classify a bond?

These limits are approximate. A common alternative is to use 0.5 and 2.0. A bond between a metal and a nonmetal is usually classed as ionic regardless of the exact number.

Electronegativity differences in common bonds

What is a dipole moment?

The dipole moment (μ) measures the separation of positive and negative charge in a molecule. For two charges of magnitude q separated by a distance d, μ = q × d. It is a vector pointing from the positive to the negative end. The SI unit is coulomb meter (C·m), but the debye (D) is commonly used: 1 D = 3.336 × 10⁻³⁰ C·m. Dipole moments are measured experimentally, for example by the effect of an electric field on a gas or liquid.

Are polar bonds the same as polar molecules?

No. A molecule is polar if it has polar bonds and its shape is such that the bond dipoles do not cancel. In carbon dioxide, each C=O bond is polar, but the molecule is linear and the two dipoles point in opposite directions and cancel, so the molecule is nonpolar. In carbon tetrachloride, four polar C–Cl bonds are arranged symmetrically in a tetrahedron, and the dipoles cancel. In water, the two O–H bonds are at an angle of 104.5°, so their dipoles add up and the molecule is polar. In ammonia, the pyramidal shape also gives a net dipole.

Worked examples

1. Classify the bonds. For N–H, ΔEN = 3.04 − 2.20 = 0.84, so the bond is polar covalent with N δ− and H δ+. For C–S, 2.58 − 2.55 = 0.03, so the bond is nonpolar.
2. Rank by polarity. C–H (0.35), N–H (0.84), O–H (1.24), F–H (1.78). The order of increasing polarity is C–H < N–H < O–H < F–H.
3. Percent ionic character of HCl. The bond length is 127.4 pm and the dipole moment is 1.08 D. If the bond were fully ionic, with charges of +e and −e, the dipole would be 1.602 × 10⁻¹⁹ C × 1.274 × 10⁻¹⁰ m = 2.04 × 10⁻²⁹ C·m, which is 6.12 D. The ionic character is 1.08 ÷ 6.12 × 100% = 17.6%.
4. Percent ionic character of HF. The bond length is 91.7 pm and μ = 1.82 D. The fully ionic value is 1.602 × 10⁻¹⁹ × 0.917 × 10⁻¹⁰ ÷ 3.336 × 10⁻³⁰ = 4.40 D. The ionic character is 1.82 ÷ 4.40 × 100% = 41%.
5. Polar or nonpolar molecule? BF₃ has three polar B–F bonds arranged in a flat triangle with 120° angles, so the dipoles cancel, and the molecule is nonpolar. NF₃ has a pyramidal shape, and the net dipole is small but not zero.
6. Direction of the dipole in CH₃Cl. The C–Cl bond is polar with Cl δ−, and the C–H bonds are almost nonpolar. The molecule has a net dipole of 1.9 D pointing toward the chlorine atom.

How does polarity affect properties?

• Solubility: polar molecules dissolve in polar solvents such as water, and nonpolar molecules in nonpolar solvents (“like dissolves like”).
• Boiling point: polar molecules attract each other by dipole-dipole forces and have higher boiling points than nonpolar molecules of similar mass. Hydrogen bonding in water, HF and ammonia gives even higher values.
• Reactivity: the partial positive atom is attacked by electron-rich reagents (nucleophiles), and the partial negative atom by electron-poor reagents (electrophiles).
• Behavior in electric fields: polar molecules orient in an electric field. This is used in microwave ovens, which make polar water molecules rotate.

How did the concept develop?

Linus Pauling introduced the electronegativity scale in 1932 and described the continuum between ionic and covalent bonds. The Dutch physicist Peter Debye measured the electric dipole moments of molecules in the 1910s and 1920s, and the unit of dipole moment is named after him. Debye received the Nobel Prize in Chemistry in 1936.

Which mistakes are common?

Mistake 1: Treating bond type as an either/or choice

Bonds form a continuum. Even a bond in NaCl has some covalent character.

Mistake 2: Assuming that every molecule with polar bonds is polar

Symmetrical molecules such as CO₂ and CCl₄ are nonpolar.

Mistake 3: Pointing the arrow the wrong way

The arrow of a bond dipole points toward the more electronegative atom, the δ− end.

Mistake 4: Using only the electronegativity difference for molecule polarity

The shape of the molecule is also needed.