The Electronic Theory of Bonding
Atoms combine in fixed ratios to form new substances, mainly through interactions between their valence (outermost) electrons. The electronic theory of bonding states that, when forming chemical bonds, atoms lose, gain or share valence electrons to reach a stable electronic structure like that of a noble gas — a complete outer shell of 2 electrons (a duplet) or 8 electrons (an octet).
A chemical bond is a strong electrostatic force of attraction created by the complete transfer or sharing of valence electrons. Only valence electrons take part in bonding. There are three main types.
1. Ionic (Electrovalent) Bonding
Ionic bonding is a strong electrostatic attraction between a metal atom and a non-metal atom, created when the metal atom transfers its valence electron(s) completely to the non-metal atom. The metal becomes a positive ion (cation); the non-metal becomes a negative ion (anion). Ionic bonds typically form between Group I/II metals and Group VI/VII non-metals — sodium and chlorine, potassium and oxygen, calcium and fluorine, and so on.
Note: although hydrogen is sometimes placed in Group I because it has one valence electron, it doesn’t form ionic bonds with metals — hydrogen–non-metal compounds are covalent. Hydrogen can form ionic compounds with metals, though; these are called hydrides.
Example — sodium chloride: sodium (configuration 2,8,1) loses its single outer electron to become Na⁺ (2,8); chlorine (configuration 2,8,7) gains that electron to become Cl⁻ (2,8,8). The resulting NaCl is a giant ionic crystal — sodium and chloride ions stacked in a regular lattice, held together by strong electrostatic forces.
Properties of ionic compounds: high melting points (strong electrostatic attraction throughout the lattice); poor conductors as solids (no free ions or electrons to carry charge) but good conductors molten or dissolved (ions become free to move); soluble in water and other polar solvents; insoluble in organic solvents.
2. Covalent Bonding
Covalent bonding is a strong electrostatic attraction between non-metal atoms, formed by sharing valence electrons rather than transferring them.
Simple (ordinary) covalent bonds form when each atom contributes electrons to a shared pair — seen in H₂, O₂, Cl₂, CH₄, NH₃, CO₂ and many more. In ammonia (NH₃), nitrogen (configuration 2,5) shares three of its five valence electrons with three hydrogen atoms, forming three covalent bonds; nitrogen ends up with a stable octet, each hydrogen with a stable duplet, and nitrogen keeps one unused lone pair.
Dative (coordinate) covalent bonds also share a pair of electrons, but both electrons come from just one atom — the donor — while the other atom, the acceptor, contributes none. This needs a donor atom with a lone pair, and an acceptor willing to receive it; it’s usually drawn as an arrow from donor to acceptor. The ammonium ion (NH₄⁺) is a classic example: ammonia’s lone pair on nitrogen is donated to an H⁺ ion (which has an empty shell), giving nitrogen a stable octet and hydrogen a stable duplet. Phosphorus oxychloride (POCl₃) and hydrogen peroxide (H₂O₂) also contain dative bonds.
Covalent compounds form either simple molecular structures (e.g. ice, iodine, camphor) or giant atomic/covalent structures (e.g. diamond, silicon dioxide).
Properties of covalent compounds: poor conductors of electricity (no free ions or electrons); low melting and boiling points, and often liquid or gaseous at room temperature (held together only by weak intermolecular forces, even though the bonds within each molecule are strong); insoluble in polar solvents like water but soluble in non-polar solvents like ethanol or methylbenzene.
3. Metallic Bonding
Metallic bonding is a strong electrostatic attraction between the positively charged nuclei of metal atoms and a surrounding “sea” of delocalised (mobile) valence electrons. In aluminium, for instance, each atom loses its outer electrons into this shared sea, becoming a positive ion held in place by strong attraction to the delocalised electrons — forming a giant metallic lattice. Sodium, magnesium, iron, copper and zinc are further examples.
Properties of metals: high melting and boiling points (strong metallic bonding); malleable (layers of atoms can slide over each other, so metals can be hammered into sheets); ductile (can be drawn into wires); good conductors of heat and electricity (thanks to the sea of free-moving electrons).
