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Ionic bonds form when electrons are transferred from a metal to a nonmetal, creating positively charged cations and negatively charged anions. These ions attract each other electrostatically. Ionic compounds form when there is a large difference in electronegativity (typically > 1.7), making electron transfer favorable.
From periodic trends, you know that metals on the left side of the periodic table have low ionization energies (they give up electrons easily) and nonmetals on the right have high electron affinities (they readily accept electrons). Ionic bonding is what happens when these two tendencies meet: a metal atom transfers one or more valence electrons to a nonmetal atom, and the resulting oppositely charged ions are held together by electrostatic attraction — the same force described by Coulomb's law.
Consider sodium and chlorine. Sodium (Group 1) has one valence electron and a low ionization energy — removing that electron leaves it with a stable noble-gas electron configuration (like neon). Chlorine (Group 17) needs one electron to complete its octet. When sodium transfers its valence electron to chlorine, two ions form: Na⁺ (a cation, positively charged because it lost an electron) and Cl⁻ (an anion, negatively charged because it gained one). The electrostatic attraction between these opposite charges is the ionic bond. This is not a bond between two specific atoms in the way a covalent bond is — each Na⁺ attracts every surrounding Cl⁻, and vice versa, forming an extended three-dimensional crystal lattice rather than discrete molecules.
The electronegativity difference between the two atoms predicts whether bonding will be ionic or covalent. When the difference is large (conventionally greater than about 1.7), the more electronegative atom pulls the shared electrons so completely toward itself that the transfer is effectively complete — ionic bonding results. Sodium (electronegativity 0.9) and chlorine (3.2) differ by 2.3, well into ionic territory. When the difference is small, electrons are shared rather than transferred, and covalent bonding results. This is a continuum, not a sharp boundary: bonds with intermediate electronegativity differences have partial ionic character.
The properties of ionic compounds follow directly from the lattice structure and the strength of electrostatic forces. High melting points result because pulling ions apart from a stable lattice requires overcoming many strong Coulomb attractions simultaneously. Ionic compounds are brittle because displacing one layer of ions shifts like charges next to each other, causing powerful repulsion that shatters the crystal. They conduct electricity when dissolved or melted because the ions become free to move and carry charge, but they are insulators as solids because the ions are locked in place. These macroscopic properties are not separate facts to memorize — they are direct, logical consequences of the electrostatic nature of the bond.