If covalent bonding is a handshake, ionic bonding is a gift exchange — one atom gives its electrons to another, and they are held together by the resulting electrical attraction. This happens most often between metals (which easily lose electrons) and nonmetals (which eagerly accept them). The classic example is table salt, sodium chloride (NaCl). Sodium, an alkali metal, has one electron in its outer shell that it is practically itching to lose. Chlorine, a halogen, needs just one electron to complete its outer shell. Sodium transfers its electron to chlorine, creating Na⁺ and Cl⁻ ions that snap together through electrostatic attraction.
Ionic compounds do not form individual molecules. Instead, the ions arrange themselves in a repeating three-dimensional lattice — a crystal. In a single grain of table salt, billions of Na⁺ and Cl⁻ ions are stacked in a perfect cubic pattern, each sodium surrounded by six chlorines and vice versa. This regular arrangement is why salt crystals are naturally cubic. The strong electrostatic forces throughout the lattice give ionic compounds their characteristic properties: high melting points (NaCl melts at 801°C), brittleness (a sharp blow shifts ion layers, putting like charges next to each other and causing the crystal to shatter), and the ability to conduct electricity when dissolved in water or melted (because the ions become free to move).
Ionic compounds are everywhere in daily life beyond the salt shaker. Calcium carbonate (CaCO₃) forms limestone, marble, chalk, and the shells of marine organisms. Potassium chloride (KCl) is used as a salt substitute and a potassium supplement. Sodium bicarbonate (NaHCO₃) is baking soda. Calcium fluoride strengthens tooth enamel. Even the batteries in your smoke detector rely on ionic compounds. Every time you see a compound made from a metal and a nonmetal with a high melting point that dissolves clearly in water, you are almost certainly looking at ionic bonding in action.