Types of Chemical Bonds

Every bond is a tug-of-war over electrons. Electronegativity measures how hard an atom pulls on shared electrons, and the gap between two atoms' electronegativities, delta-EN, tells you who wins. When the gap is huge, like sodium and chlorine, chlorine wins outright: it rips the electron off sodium and both atoms become ions, held together by the electrical attraction between a positive and negative charge. That's an ionic bond. When the gap is small or zero, neither atom can fully win, so they share the electrons instead, forming a covalent bond. If the sharing is uneven, the bond is polar covalent, with a slightly negative end and a slightly positive end. If the atoms are identical or nearly so, like two sodium atoms or two chlorine atoms, the sharing is even and the bond is nonpolar covalent.

Look at the electrons themselves. In NaCl, the electron has fully transferred, leaving Na+ and Cl- ions in a rigid lattice. In HCl, the electron pair sits closer to chlorine but still belongs to both atoms, giving a polar covalent bond with a dipole. In a block of copper, electrons aren't owned by any single atom at all: they move freely through a 'sea' of electrons shared by every copper nucleus, which is metallic bonding and explains why metals conduct electricity.

Textbooks hand you cutoffs, delta-EN below 0.4 nonpolar, 0.4 to 1.7 polar, above 1.7 ionic, but electronegativity difference is a continuous number, not a switch. Bond character slides smoothly from fully shared to fully transferred, and these labels are just convenient zones on that slide. This spectrum sets up the next idea: how bond type and the resulting structure, ionic lattice versus discrete molecule, control properties like melting point.