Chemical Bonds and Electronegativity

One Spectrum, Three Names

A chemical bond forms when an arrangement of nuclei and electrons is lower in energy than the separated atoms. How the electrons are shared out gives the three classic bond types.

An ionic bond transfers electrons outright, producing cations and anions held by electrostatic attraction in a giant lattice. Its strength is captured by the Coulomb expression for lattice energy,

$$E \propto \frac{q_1 q_2}{d}$$

so lattice energy rises sharply with ionic charge and falls as the ions get bigger. MgO ($2+/2-$) is bound roughly four times as strongly as NaF ($1+/1-$) at similar spacing, which is why MgO melts at 2852 °C and NaF at 993 °C.

A covalent bond shares a pair of electrons between two nuclei. A metallic bond pools valence electrons into a delocalised sea across a lattice of cations, which explains conduction, malleability and lustre in one stroke.

Electronegativity is the tendency of an atom in a bond to attract the shared electrons. On the Pauling scale it runs from about 0.7 (Cs) to 4.0 (F), rising across a period and falling down a group, the same $Z_{\text{eff}}$ story as before. The electronegativity difference $\Delta$EN predicts the character of the bond:

These boundaries are conventions on a continuum, not physical walls. A polar bond has a dipole moment $\mu = q \times d$, drawn as an arrow pointing toward the more electronegative atom.

Common pitfall: treating ionic and covalent as a strict either/or. Real bonds sit on a sliding scale of shared-to-transferred; HF at $\Delta$EN $= 1.9$ is conventionally "ionic" yet exists as discrete molecules, and even NaCl retains a little covalent character.