Chemical polarity describes how electrons are distributed within a molecule. This uneven sharing creates regions of slight positive and negative charge, which influences molecular behavior. Understanding polarity is fundamental as it dictates how molecules interact, affecting properties like solubility, boiling points, and chemical reactivity.
The Foundation: Electronegativity
Electronegativity is an atom’s inherent ability to attract shared electrons towards itself in a chemical bond. It varies predictably across the periodic table, generally increasing from left to right across a period and decreasing down a group. The Pauling scale assigns numerical values to electronegativity, standardizing electron-attracting power comparison. The differences in these values drive the unequal sharing of electrons in chemical bonds.
Calculating Bond Polarity
Bond polarity is determined by the difference in electronegativity between two bonded atoms. A larger difference indicates a more unequal sharing of electrons, creating a more polar bond. To classify a bond, subtract the smaller electronegativity value from the larger one. For instance, hydrogen has an electronegativity of approximately 2.20, while chlorine has a value of about 3.16.
Bonds are generally classified into three main types based on this difference. If the electronegativity difference is very small, typically between 0 and 0.4, the bond is considered nonpolar covalent, with electrons shared almost equally. An example is the bond in a hydrogen molecule (H-H), where the difference is 2.20 – 2.20 = 0, indicating equal sharing of electrons.
When the difference in electronegativity falls within a moderate range, usually between 0.5 and 1.7, the bond is classified as polar covalent. In a hydrogen chloride molecule (H-Cl), the difference is 3.16 – 2.20 = 0.96, meaning electrons are unequally shared, with chlorine attracting them more strongly. This unequal sharing creates a partial negative charge on the more electronegative atom and a partial positive charge on the less electronegative atom. Finally, if the electronegativity difference is greater than 1.7, the bond is considered ionic, where electrons are transferred rather than shared. An example is the bond in sodium chloride (Na-Cl), where sodium has an electronegativity of 0.93 and chlorine has 3.16, resulting in a difference of 2.23.
Determining Molecular Polarity
Determining overall molecular polarity requires considering both individual bond polarities and the molecule’s three-dimensional shape. Even if a molecule contains polar bonds, its overall polarity can be nonpolar if the molecule’s geometry causes these individual bond dipoles to cancel each other out. This cancellation occurs in highly symmetrical molecules.
Consider carbon dioxide (CO2), which has two polar carbon-oxygen bonds. Oxygen is more electronegative than carbon, making each C=O bond polar. However, the carbon dioxide molecule has a linear shape. Because the two bond dipoles point in opposite directions and are of equal magnitude, they cancel, resulting in a nonpolar molecule.
Another example is carbon tetrachloride (CCl4), which has four polar C-Cl bonds arranged symmetrically in a tetrahedral shape. Despite individual bond polarity, the symmetrical arrangement leads to a net cancellation of bond dipoles, making the molecule nonpolar.
Conversely, molecules with polar bonds that are arranged asymmetrically result in an overall molecular dipole moment, making the molecule polar. Water (H2O) has two polar oxygen-hydrogen bonds. The water molecule has a bent shape due to lone pairs on the oxygen atom. This bent geometry prevents bond dipoles from canceling, leading to a net molecular dipole moment and an overall polar molecule.
Similarly, ammonia (NH3) has three polar nitrogen-hydrogen bonds. The nitrogen atom is more electronegative than hydrogen, and its lone pair contributes to an asymmetrical charge distribution, making ammonia polar.