Chemical bonds are the forces that hold atoms together, forming every molecule and substance. Understanding how atoms join is foundational to chemistry and biology. The nature of the bond depends on how atoms share or transfer their valence electrons. This interaction determines the stability, shape, and properties of the resulting compound.
Defining Electron Sharing and Pulling Power
Covalent bonding is the mechanism by which two atoms, typically nonmetals, achieve a stable electron configuration by mutually sharing a pair of outer electrons. The shared electron pair orbits both nuclei, effectively linking the atoms to form a molecule.
Electronegativity (EN) quantifies an atom’s inherent power to attract the shared electrons toward itself when participating in a chemical bond. This “pulling power” is a relative measure, meaning it is only meaningful when comparing two atoms bonded together. Atoms with a greater number of protons and a smaller atomic size generally exhibit higher electronegativity. This is because their nucleus can exert a stronger attractive force on the shared electrons.
The Electronegativity Scale and Bond Polarity
The relationship between electronegativity and covalent bonding is quantified by calculating the difference (\(\Delta\text{EN}\)) between the values of the two bonded atoms. This difference predicts how the shared electrons will be distributed. The Pauling scale is the most common system, assigning numerical values to elements ranging from about 0.7 to 4.0, with fluorine being the highest.
The greater the \(\Delta\text{EN}\) between two atoms, the more unequally the electron pair is shared, which introduces the concept of bond polarity. Polarity arises when one atom has a significantly stronger pull, causing the shared electron density to shift toward it. This unequal distribution is often visualized as a tug-of-war where one team is much stronger than the other.
The atom with the higher electronegativity gains a partial negative charge (\(\delta^-\)) because the electrons spend more time near its nucleus. Conversely, the less electronegative atom acquires a partial positive charge (\(\delta^+\)). This charge separation creates an electric dipole within the bond. Bonding exists along a spectrum, where a small \(\Delta\text{EN}\) results in a mostly shared bond and a very large difference transitions the bond toward an ionic character.
Nonpolar vs. Polar Covalent Bonds
The spectrum of covalent bonding is primarily divided into two categories based on the degree of electron sharing: nonpolar and polar. Nonpolar covalent bonds occur when the \(\Delta\text{EN}\) between the two atoms is zero or very close to zero, typically less than 0.4. In these bonds, such as in the diatomic molecule \(\text{O}_2\), the electrons are shared perfectly equally because both atoms possess the same pulling power.
Polar covalent bonds form when the \(\Delta\text{EN}\) is significant, generally falling within a range of approximately 0.4 to 1.7. This difference causes the unequal sharing of the electron cloud, creating a distinct positive end and a negative end to the bond. An example is the bond between hydrogen and chlorine in \(\text{HCl}\), where chlorine’s higher electronegativity pulls the electron density closer to itself.
Chemists use a dipole moment vector to visually represent the direction of the electron pull. This vector is an arrow pointing from the \(\delta^+\) atom to the \(\delta^-\) atom, showing how the charge separation is oriented. The magnitude of this dipole moment directly correlates with the size of the \(\Delta\text{EN}\). When the difference in electronegativity increases beyond the polar covalent range, the bond crosses into ionic bonding, where one atom transfers the electron entirely.