Bond polarity describes the electrical charge distribution within a chemical bond. It refers to how equally or unequally shared electrons are distributed between the bonded atoms. This distribution determines if one end of the bond is more negative and the other more positive. Understanding this charge separation is foundational because it dictates how molecules interact, influencing a substance’s physical properties like melting point, boiling point, and solubility.
The Origin of Polarity: Unequal Electron Sharing
Chemical bonds involving the sharing of electrons are called covalent bonds. When two identical atoms, such as two oxygen atoms in O2, form a bond, the electron pair is shared symmetrically. However, when two different atoms join, the sharing is almost never equal, leading to an asymmetrical electron distribution.
This unequal sharing creates a polar covalent bond. The atom that pulls the shared electrons closer acquires a slight excess of negative charge, while the other atom acquires a slight positive charge. Chemists represent these slight, unbalanced charges using the lowercase Greek letter delta, resulting in a partial negative charge on the electron-rich atom and a partial positive charge on the electron-poor atom. This charge separation creates a bond dipole moment, which measures the direction and magnitude of the separation. It is often visualized as an arrow pointing from the partial positive atom toward the partial negative atom.
Quantifying Polarity: The Electronegativity Scale
The reason some atoms pull on electrons more strongly than others is a property known as electronegativity. Electronegativity is defined as the power of an atom within a molecule to attract electrons to itself. This inherent atomic property allows chemists to predict which atom in a bond will bear the partial negative charge. The most widely used system for quantifying this property is the Pauling scale, which assigns a numerical value to nearly every element. The higher the electronegativity value, the greater the atom’s pull on the shared electrons. By calculating the difference in electronegativity (EN difference) between two bonded atoms, the type of chemical bond can be classified.
Classifying Bond Types
A bond between two identical atoms, such as hydrogen (H2), has an EN difference of zero and is classified as nonpolar covalent because the electrons are shared equally. If the EN difference is small, typically less than 0.5, the bond is still considered nonpolar covalent, like the bond between carbon and hydrogen. When the EN difference is moderate, usually falling between 0.5 and approximately 1.7, the bond is polar covalent. For example, the bond in hydrogen chloride (HCl) has an EN difference of 0.9, meaning the electron density is significantly shifted toward the chlorine atom. When the difference is very large, generally greater than 1.7 or 2.0, the electron transfer is so complete that the bond is considered an ionic bond, as seen in table salt (NaCl). The EN difference provides a continuum of bond types that range from purely nonpolar to fully ionic.
From Bond Polarity to Molecular Polarity
While the polarity of an individual bond is determined by the electronegativity difference, the overall polarity of an entire molecule is a distinct concept. A molecule’s total electrical character, known as its molecular polarity, depends not only on the existence of polar bonds but also on their spatial arrangement. The individual bond dipole moments are vector quantities, meaning they have both magnitude and direction. The geometry of the molecule dictates whether these individual bond dipoles add up or cancel one another out.
Symmetrical Molecules
Symmetrical molecules, such as carbon dioxide (CO2), often have highly polar bonds. However, because the molecule is linear, the two bond dipoles point in opposite directions and are of equal magnitude. This arrangement causes the dipoles to cancel each other completely, resulting in a net molecular dipole moment of zero and a nonpolar molecule. Similarly, molecules with a tetrahedral geometry, like carbon tetrachloride (CCl4), are nonpolar if all four outer atoms are identical. The highly symmetrical arrangement of the four polar C-Cl bonds ensures that their electrical effects perfectly balance.
Asymmetrical Molecules
In contrast, molecules with asymmetrical shapes, such as water (H2O) or ammonia (NH3), are polar. The water molecule has a bent shape, which means the two O-H bond dipoles do not point directly away from each other. Instead, they are angled, and their vector sum results in a net molecular dipole moment. This net polarity means the water molecule has a distinct negative side and a positive side, which significantly impacts its interactions. Molecular polarity is what allows water to dissolve many substances, a principle summarized by the saying “like dissolves like,” where polar solvents dissolve polar solutes.