Bond Polarity and a Continuum of Bonding

When covalent bonds form between atoms with different electronegativities, the bonding electrons are shared unequally, forming a polar bond.

When two atoms of the same element bond, they form a nonpolar covalent bond, in which the nuclei of the bonded atoms exert equal or nearly equal force on the shared electrons. In this configuration the shared electrons are likely to be found near either of the atoms at a given time; the electrons are shared equally.

When two atoms with different electronegativities bond, they form a polar covalent bond, which is a covalent bond in which the electron density is more localized on one end of the bond. One end is slightly positive, and one end is slightly negative. The electrons are more likely to be found closer to the nucleus that exerts a greater force on the electrons. Unequal charge distribution within a polar covalent bond gives a permanent bond dipole moment, a mathematical measure of the amount of charge difference, so that electrons are shared unequally.

Consider hydrogen chloride (HCl), a molecular gas that dissociates into ions in water, forming hydrochloric acid. Hydrogen has an electronegativity of 2.20 on the Pauling scale, a dimensionless ranking of electronegativity. Chlorine has an electronegativity of 3.16. In the hydrogen chloride molecule, chlorine exerts a greater force on the shared electrons, causing them to have a greater probability of being found near the chlorine atom. The chlorine side of hydrogen chloride molecules, therefore, has a partial negative charge, and the hydrogen side of the molecule has a partial positive charge. These partial charges are represented by the lowercase Greek letter delta ($\delta{+}$ and $\delta{-}$). A molecule with partial positive and partial negative charges is called a dipole.

Hydrogen Chloride Dipole

Ionic and covalent bonds are often described as two different bonding schemes. However, one way of considering an ionic bond is to think of an extremely polar covalent bond. In reality, bonds do not carry pure ionic or pure covalent characteristics. Each bond is part covalent and part ionic. It is appropriate to think of bonding as a continuum of states, from pure nonpolar electron sharing to a complete electron transfer.

Scientists classify bonds as ionic or covalent to make things simpler. Electronegativity differences play a key role in this classification. When two atoms of the same element bond, the electronegativity difference is zero, and the bond is nonpolar. A bond between two atoms with an electronegativity difference of up to 0.4 according to the Pauling scale is still considered to be a nonpolar bond. A bond between two atoms with an electronegativity difference of 0.4 to 1.8 is considered a polar covalent bond. A bond between two atoms with an electronegativity difference of over 1.8 is an ionic bond. These numbers vary according to source; some sources use 0.5 instead of 0.4 or 1.7 instead of 1.8. Keep in mind these numbers indicate positions on a spectrum. A bond between atoms with a 1.7 electronegativity difference will be similar to a bond between atoms with a 1.8 electronegativity difference, even though one is considered covalent and the other ionic. The periodic table is the best guide to identifying whether elements form ionic or covalent bonds. Combinations of atoms from the lower left and far right generally have ionic bonds, while bonds between elements adjacent to each other are covalent. In general, bonds between two nonmetals tend to be covalent, and those between a metal and nonmetal are often ionic.

Electronegativity Ranges for Different Types of Bonds

Bond Electronegativity Difference
Covalent < 0.4
Polar covalent 0.4 to 1.8
Ionic > 1.8

Bonds become more ionic as the electronegativity difference between two atoms increases.