In chemistry, polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole moment, with a negatively charged end and a positively charged end. Polar molecules must contain one or more polar bonds due to a difference in electronegativity between the bonded atoms. Molecules containing polar bonds have no molecular polarity if the bond dipoles cancel each other out by symmetry. Polar molecules interact through dipole-dipole intermolecular forces and hydrogen bonds. Polarity underlies a number of physical properties including surface tension, solubility, and melting and boiling points.
Polarity of bonds
Not all atoms attract electrons with the same force. The amount of "pull" an atom exerts on its electrons is called its electronegativity. Atoms with high electronegativities – such as fluorine, oxygen, and nitrogen – exert a greater pull on electrons than atoms with lower electronegativities such as alkali metals and alkaline earth metals. In a bond, this leads to unequal sharing of electrons between the atoms, as electrons will be drawn closer to the atom with the higher electronegativity. Because electrons have a negative charge, the unequal sharing of electrons within a bond leads to the formation of an electric dipole: a separation of positive and negative electric charge. Because the amount of charge separated in such dipoles is usually smaller than a fundamental charge, they are called partial charges, denoted as δ+ (delta plus) and δ− (delta minus). These symbols were introduced by Sir Christopher Ingold and Edith Hilda (Usherwood) Ingold in 1926. The bond dipole moment is calculated by multiplying the amount of charge separated and the distance between the charges. These dipoles within molecules can interact with dipoles in other molecules, creating dipole-dipole intermolecular forces.
Classification Bonds can fall between one of two extremes – completely nonpolar or completely polar. A completely nonpolar bond occurs when the electronegativities are identical and therefore possess a difference of zero. A completely polar bond is more correctly called an ionic bond, and occurs when the difference between electronegativities is large enough that one atom actually takes an electron from the other. The terms "polar" and "nonpolar" are usually applied to covalent bonds, that is, bonds where the polarity is not complete. To determine the polarity of a covalent bond using numerical means, the difference between the electronegativity of the atoms is used. Bond polarity is typically divided into three groups that are loosely based on the difference in electronegativity between the two bonded atoms. According to the Pauling scale:
Nonpolar bonds generally occur when the difference in electronegativity between the two atoms is less than 0.5 Polar bonds generally occur when the difference in electronegativity between the two atoms is roughly between 0.5 and 2.0 Ionic bonds generally occur when the difference in electronegativity between the two atoms is greater than 2.0 Pauling based this classification scheme on the partial ionic character of a bond, which is an approximate function of the difference in electronegativity between the two bonded atoms. He estimated that a difference of 1.7 corresponds to 50% ionic character, so that a greater difference corresponds to a bond which is predominantly ionic. As a quantum-mechanical description, Pauling proposed that the wave function for a polar molecule AB is a linear combination of wave functions for covalent and ionic molecules: ψ = aψ(A:B) + bψ(A+B−). The amount of covalent and ionic character depends on the values of the squared coefficients a2 and b2.
Bond dipole moments
The bond dipole moment uses the idea of electric dipole moment to measure the polarity of a chemical bond within a molecule. It occurs whenever there is a separation of positive and negative charges. The bond dipole μ is given by:
μ = δ d {\displaystyle \mu =\delta \,d} . The bond dipole is modeled as δ+ — δ– with a distance d between the partial charges δ+ and δ–. It is a vector, parallel to the bond axis, pointing from minus to plus, as is conventional for electric dipole moment vectors. Chemists often draw the vector pointing from plus to minus. This vector can be physically interpreted as the movement undergone by electrons when the two atoms are placed a distance d apart and allowed to interact, the electrons will move from their free state positions to be localised more around the more electronegative atom. The SI unit for electric dipole moment is the coulomb–meter. This is too large to be practical on the molecular scale. Bond dipole moments are commonly measured in debyes, represented by the symbol D, which is obtained by measuring the charge δ {\displaystyle \delta } in units of 10−10 statcoulomb and the distance d in Angstroms. Based on the conversion factor of 10−10 statcoulomb being 0.208 units of elementary charge, so 1.0 debye results from an electron and a proton separated by 0.208 Å. A useful conversion factor is 1 D = 3.335 64×10−30 C m. For diatomic molecules there is only one (single or multiple) bond so the bond dipole moment is the molecular dipole moment, with typical values in the range of 0 to 11 D. At one extreme, a symmetrical molecule such as bromine, Br2, has zero dipole moment, while near the other extreme, gas phase potassium bromide, KBr, which is highly ionic, has a dipole moment of 10.41 D. For polyatomic molecules, there is more than one bond. The total molecular dipole moment may be approximated as the vector sum of the individual bond dipole moments. Often bond dipoles are obtained by the reverse process: a known total dipole of a molecule can be decomposed into bond dipoles. This is done to transfer bond dipole moments to molecules that have the same bonds, but for which the total dipole moment is not yet known. The vector sum of the transferred bond dipoles gives an estimate for the total (unknown) dipole of the molecule.
Polarity of molecules
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