The VSEPR theory can be used to determine the electron pair geometries and molecular structures as follows:
When atoms with different electronegativities form a bond, the electrons are pulled toward the more electronegative atom, leaving one atom with a partial positive charge (δ+) and the other atom with a partial negative charge (δ–). Such bonds are called polar covalent bonds, and the separation of charge gives rise to a bond dipole moment. The magnitude of a bond dipole moment is represented by the Greek letter µ and is given by:
μ = Qr
where Q is the magnitude of the partial charges (determined by the electronegativity difference), and r is the distance between them. Dipole moments are commonly expressed in debyes, where one debye is equal to 3.336 × 10−30 C·m.
The bond dipole moment is a vector represented by an arrow pointing along the bond from the less electronegative toward the more electronegative atom, with a small plus sign on the less electronegative end.
A whole molecule may also have a separation of charge, depending on its molecular structure and the polarity of each of its bonds. Such molecules are said to be polar. The dipole moment measures the extent of net charge separation in the molecule as a whole. In diatomic molecules, the bond dipole moment determines the molecular polarity.
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 the vector sum of their bond moments equals zero, then the molecule is nonpolar (e.g., CO2). The water molecule has a bent molecular structure, and the two bond moments do not cancel. Therefore, water is a polar molecule with a net dipole moment.
This text has been adapted fromOpenstax, Chemistry 2e, Section 7.6 Molecular Structure and Polarity.
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