When two atoms join together, they form a chemical bond by sharing their outermost electrons, known as valence electrons. This process, called covalent bonding, allows atoms to achieve a more stable electron configuration. While “sharing” suggests an equal arrangement, most covalent bonds involve a distribution of electrons that is far from balanced. The electrons are not always positioned exactly halfway between the two atomic nuclei. Instead, the sharing exists on a continuous spectrum, where the electrons may be pulled closer to one atom than the other, leading to unequal sharing.
The Atomic Property That Governs Sharing
The primary factor determining whether the shared electrons are distributed equally or unequally is an inherent atomic property called electronegativity. Electronegativity is an atom’s measure of its ability to attract a shared electron pair toward itself in a chemical bond. It essentially describes the atom’s “pulling power” over the bonding electrons, like an ongoing tug-of-war where the shared electrons are the prize.
Atoms with a higher nuclear charge and a smaller atomic radius tend to have greater electronegativity. A smaller atom means the bonding electrons are closer to the attractive nucleus, increasing the pull. Conversely, larger atoms, where the valence electrons are farther from the nucleus, exert a weaker pull on the shared electrons. Linus Pauling developed a widely used scale to quantify this atomic property, assigning a numerical value to each element’s electron-attracting ability.
If the two atoms involved in a bond have identical electronegativity, the electrons are shared equally. When two atoms with different electronegativity values bond, the atom with the higher value attracts the electron cloud closer to its nucleus. This unequal attraction shifts the shared electrons toward the more electronegative atom. The magnitude of this difference dictates the degree of unequal sharing and the character of the resulting bond.
Defining the Difference Between Bond Types
The difference in electronegativity (\(\Delta\text{EN}\)) between the two bonded atoms is used to classify the type of bond formed. This difference establishes a continuum, moving from equal sharing to a complete transfer of electrons. When the \(\Delta\text{EN}\) is very small (typically less than 0.4), the electrons are shared nearly equally, resulting in a nonpolar covalent bond. For example, in a hydrogen gas molecule (\(\text{H}_2\)), the identical atoms have the same electronegativity, leading to perfect electron sharing.
A moderate difference in electronegativity, generally falling between approximately 0.4 and 1.7, signifies unequal sharing and results in a polar covalent bond. In a molecule like hydrogen chloride (\(\text{HCl}\)), the chlorine atom has a much stronger attraction for the shared electrons than the hydrogen atom. This stronger pull means the electron pair spends more time closer to the chlorine nucleus, creating a bond with distinct electrical ends.
When the difference in electronegativity is very large (often greater than 1.7 or 2.0), the sharing becomes so unequal that the more electronegative atom strips the electron away. This complete transfer creates oppositely charged ions held together by electrostatic forces, forming an ionic bond. Sodium chloride (\(\text{NaCl}\)) is a common example, where chlorine gains an electron from sodium. Although these numerical cutoffs are useful guidelines, bond character is a continuum, not separated by a sharp line.
How Unequal Sharing Affects the Molecule
The unequal sharing of electrons in a polar covalent bond creates localized partial charges within the molecule. The atom that exerts the stronger pull acquires a partial negative charge (\(\delta-\)). Conversely, the atom that loses electron density acquires a partial positive charge (\(\delta+\)). These partial charges establish a bond dipole moment, a vector quantity that points from the positive pole toward the negative pole.
To determine the overall polarity of the entire molecule, one must consider both the polarity of its individual bonds and the molecule’s three-dimensional shape, known as its molecular geometry. The vector nature of the bond dipoles means they can either add up or cancel each other out, depending on the spatial arrangement of the atoms. If the molecule is highly symmetrical, the individual bond dipoles may perfectly balance, resulting in a net molecular dipole moment of zero.
For instance, carbon dioxide (\(\text{CO}_2\)) has two polar carbon-oxygen bonds. However, the molecule is linear, meaning the two bond dipoles point in opposite directions and cancel each other out, making \(\text{CO}_2\) nonpolar. In contrast, a water molecule (\(\text{H}_2\text{O}\)) also contains polar bonds, but its bent shape prevents the bond dipoles from canceling. The molecular dipole moments combine to create a significant net molecular dipole, making the water molecule polar overall.
The presence of this net dipole moment is a defining characteristic of a polar molecule, greatly influencing its physical properties like solubility, melting point, and interaction with other substances.