Molecular Geometry and Dipole Moments

VSEPR theory helps to determine electron-pair geometries and molecular geometries.

A series of steps is used to predict the geometry and bond angles of molecules, such as phosphorus trichloride.

The first step is to draw the Lewis structure of the molecule.

Next, count the total number of electron groups on the central atom. Around phosphorus, there are four electron groups: three bonding pairs and one lone pair.

Now, determine the electron-pair geometry. The electron-pair geometry is tetrahedral. However, because of the lone pair, the molecular geometry is trigonal pyramidal. The lone pair reduces the bond angle to less than 109.5°.

Bonding electron pairs are not always shared equally between the two bonding atoms.

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.

Dipole moments are commonly expressed in debyes. One debye is equal to 3.336 × 10⁻³⁰ coulomb-meters.

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.

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.