10.4: Molecular Shape and Polarity

Dipole Moment of a Molecule

Polar covalent bonds connect two atoms with differing electronegativities, leaving one atom with a partial positive charge (δ⁺) and the other atom with a partial negative charge (δ^–), as the electrons are pulled toward the more electronegative atom. This separation of charge gives rise to a bond dipole moment. The magnitude of a bond dipole moment is represented by the Greek letter mu (µ) and is given by the formula shown here, where Q is the magnitude of the partial charges (determined by the electronegativity difference) and r is the distance between the charges: μ = Qr.

This bond moment can be represented as a vector, a quantity having both direction and magnitude. Dipole vectors are shown as arrows pointing along the bond from the less electronegative atom toward the more electronegative atom. A small plus sign is drawn on the less electronegative end to indicate the partially positive end of the bond. The length of the arrow is proportional to the magnitude of the electronegativity difference between the two atoms.

Factors Determining Polarity of a Molecule

A whole molecule may also have a separation of charge, depending on its molecular structure and the polarity of each of its bonds. If such a charge separation exists, the molecule is said to be a polar molecule; otherwise, the molecule is said to be nonpolar. The dipole moment measures the extent of net charge separation in the molecule as a whole. We determine the dipole moment by adding the bond moments in three-dimensional space, taking into account the molecular structure.

For diatomic molecules, there is only one bond, so its bond dipole moment determines the molecular polarity. Homonuclear diatomic molecules such as Br₂ and N₂ have no difference in electronegativity, so their dipole moment is zero. For heteronuclear molecules such as CO, there is a small dipole moment. For HF, there is a larger dipole moment because there is a larger difference in electronegativity.

When a molecule contains more than one bond, the geometry must be taken into account. If the bonds in a molecule are arranged such that their bond moments cancel (vector sum equals zero), then the molecule is nonpolar. This is the situation in CO₂. Each of the bonds is polar, but the molecule as a whole is nonpolar. From the Lewis structure, and using VSEPR theory, we determine that the CO₂ molecule is linear with polar C=O bonds on opposite sides of the carbon atom. The bond moments cancel because they are pointed in opposite directions. In the case of the water molecule, the Lewis structure again shows that there are two bonds to a central atom, and the electronegativity difference again shows that each of these bonds has a nonzero bond moment. In this case, however, the molecular structure is bent because of the lone pairs on O, and the two bond moments do not cancel. Therefore, water does have a net dipole moment and is a polar molecule (dipole).

In an OCS molecule, the structure is similar to CO₂, but a sulfur atom replaces one of the oxygen atoms.

The C-O bond is considerably polar. Although C and S have very similar electronegativity values, S is slightly more electronegative than C, and so the C-S bond is just slightly polar. Because oxygen is more electronegative than sulfur, the oxygen end of the molecule is the negative end.

Chloromethane, CH₃Cl, is a tetrahedral molecule with three slightly polar C-H bonds and a more polar C-Cl bond. The relative electronegativities of the bonded atoms are H < C < Cl, and so the bond moments all point toward the Cl end of the molecule and sum to yield a considerable dipole moment (the molecules are relatively polar).

For molecules of high symmetry such as BF₃ (trigonal planar), CH₄ (tetrahedral), PF₅ (trigonal bipyramidal), and SF₆ (octahedral), all the bonds are of identical polarity (same bond moment) .and they are oriented in geometries that yield nonpolar molecules (dipole moment is zero). Molecules of less geometric symmetry, however, may be polar even when all bond moments are identical. For these molecules, the directions of the equal bond moments are such that they sum to give a nonzero dipole moment and a polar molecule. Examples of such molecules include hydrogen sulfide, H₂S (nonlinear), and ammonia, NH₃ (trigonal pyramidal)

To summarize, to be polar, a molecule must:

1. Contain at least one polar covalent bond.
2. Have a molecular structure such that the sum of the vectors of each bond dipole moment does not cancel.

This text has been adapted from Openstax, Chemistry 2e, Section 7.6 Molecular Structure and Polarity.

In a covalent bond like that of hydrofluoric acid, the electrons are pulled toward the more electronegative atom, indicated by a partial charge. Such bonds are called polar bonds.

The charge separation creates a vector called the bond dipole moment, which is indicated by the Greek letter µ. Its value is the product of the magnitude of the partial charges and the distance between them.

The vector points from the less to the more electronegative atom and indicates the bond dipole moment. Its length is proportional to the magnitude of the electronegativity difference between the two atoms. Most diatomic molecules containing atoms of different elements have dipole moments and therefore are polar molecules. 

Electrostatic potential maps indicate the high and low electron density regions in the compound with red and blue colors, respectively. Colors in between depict moderate electron density.

In polyatomic compounds, the net dipole moment is determined by the individual bond dipole moments and geometry of the compound.

Consider a water molecule with two polar bonds. It has a bent shape and is a polar molecule.

In contrast, a carbon dioxide molecule is linear. The two carbon-oxygen bonds are polar but are oriented in opposite directions, canceling out each other's dipole moment  and making the overall molecule nonpolar. 

Carbonyl sulfide molecules are also linear. However, the dipole moments of the carbon-oxygen and carbon-sulfur bonds do not cancel each other, and the molecule has a net dipole moment. 

Boron trifluoride is a trigonal planar compound. The dipole moments of the boron-fluorine bonds cancel each other owing to the molecular symmetry, and the compound is nonpolar. 

However, the dipole moments of the three polar bonds in trigonal pyramidal phosphorus trichloride molecules do not cancel each other, making it a polar compound.

Tetrafluoromethane is a tetrahedral molecule, which is nonpolar, as the dipole moments of the four identical polar bonds cancel each other. 

Fluoromethane is also a tetrahedral molecule. However, it possesses a net dipole moment as the C-F bond has a large dipole moment compared to the C-H bonds, and the bond dipole moments do not cancel each other.

In an electric field, polar molecules align the positive end toward the negative plate and the negative end toward the positive plate. In contrast, nonpolar molecules stay unaffected by an electric field. 

Generally, polar solvents dissolve polar solutes, and nonpolar solvents dissolve nonpolar solutes because similar types of molecules tend to have more favorable interactions. Water is polar and easily dissolves polar compounds such as sucrose, commonly known as table sugar. Oil is nonpolar and remains immiscible in water.