The redistribution of electrons in an atom, bond, or molecule creates two ends (or poles), one electron-rich negative and the other electron-deficient positive; such an atom, bond, or molecule is said to have a dipole (two poles).
The poles are always equal in magnitude but oppositely charged. The negative and the positive atom or molecular end is denoted with a delta minus (δ-) or a delta plus (δ+) sign. At times an arrow is used to denote the dipoles. The arrowhead rests above the negative end, and the tail with a plus sign on the positive end.
The formation of opposite poles can be a permanent, inherent feature of a molecule, or it can be induced.
The dipole is permanent when the atoms forming a covalent bond have an electronegativity difference between 0.5-1.7. In such situations, one more electronegative atom strongly attracts the electron density towards itself, creating natural poles.
The dipoles can also be induced when molecules with pre-existing dipoles (polar molecules) interact with molecules with no poles (nonpolar molecules). Or when the electrons of the nonpolar molecules redistribute due to their continuous motion to create fluctuating, temporary poles.
Therefore, dipoles are of three types- permanent, instantaneous, and induced.
The strength of the pole separation or the dipole is measured using dipole moment (μ) having a unit Debye (D) where the lower numbers indicate lower separation of charges.
For example, there is a dipole in the H-O bond (1.53 D) than in the C-H bond (0.30 D). Or on comparing the dipoles of two molecules, the CO2 molecule is nonpolar (0 D) with no opposite poles, but H2CO is polar (2.3D).
The formation of poles is a primary reason the atoms and molecules interact and affect various physical, chemical, and biological properties.