Ionic, Covalent and Metallic Bonding
Ionic Bonding
When a metal reacts with a non-metal, electrons are transferred from the metal atom to the non-metal atom. The metal atom loses one or more electrons to form a positive ion (cation), and the non-metal atom gains those electrons to become a negative ion (anion). Both ions then have full outer electron shells and are stable.
The strong electrostatic attraction between the oppositely charged ions is called an ionic bond. In sodium chloride (NaCl), sodium loses one electron to become \(Na^+\) and chlorine gains one electron to become \(Cl^-\). The resulting ions pack together in a regular three-dimensional arrangement called a giant ionic lattice.
Ionic bonding: metal atoms lose electrons (form positive ions) and non-metal atoms gain electrons (form negative ions). Strong electrostatic attraction between the opposite charges holds ions together in a giant ionic lattice.
Naming Ionic Compounds
The name of an ionic compound gives useful information about which ions are present. When a metal combines with a non-metal element only, the non-metal ion takes an -ide ending — for example, sodium chloride, magnesium oxide, potassium iodide. When the compound contains oxygen as part of a polyatomic ion, it usually takes an -ate ending — for example, sodium sulfate, copper(II) nitrate, calcium carbonate. Recognising these endings tells you instantly whether a simple or complex ion is involved.
Dot-and-cross diagrams show only the outer shell electrons. Use dots for one atom and crosses for the other. For sodium chloride:
- Draw Na with one outer electron (dot). Draw Cl with seven outer electrons (crosses).
- Show the dot crossing into Cl's outer shell, so Cl now has eight electrons. Draw square brackets around each ion and write the charge outside: \([Na]^+\) and \([Cl]^-\).
For magnesium oxide, Mg loses two electrons (2,8,2 → 2,8, becoming \(Mg^{2+}\)) and oxygen gains two (2,6 → 2,8, becomes \(O^{2-}\)).
Covalent Bonding
When non-metals bond with other non-metals, atoms share pairs of electrons rather than transferring them. Each shared pair of electrons is a covalent bond. Both atoms count the shared pair as part of their own outer shell, so both achieve a full, stable electron configuration.
Common examples:
- \(H_2\) — a single covalent bond (one shared pair).
- \(O_2\) — a double covalent bond (\(O=O\), two shared pairs).
- \(H_2O\) — two single covalent bonds (oxygen bonds to two hydrogen atoms).
- \(CH_4\) — four single covalent bonds (carbon bonds to four hydrogen atoms).
- \(NH_3\) — three single covalent bonds plus one lone pair on nitrogen.
For water (\(H_2O\)): draw oxygen in the centre with six outer-shell electrons. Draw a hydrogen on each side with one electron. Show one dot from O and one cross from each H overlapping in a shared pair. Each H now has a full shell of two; oxygen has a full shell of eight.
For methane (\(CH_4\)): carbon has four outer electrons. Each of the four hydrogen atoms contributes one electron. Show four shared pairs around carbon, giving it eight; each hydrogen has two.
Only draw outer-shell electrons. Never add electrons that aren't there or remove electrons to make the diagram look neater — the total electron count must balance.
Metallic Bonding
In metals, the outer electrons leave their parent atoms and become delocalised electrons — free to move throughout the entire metal structure. What remains are positive metal ions. The metallic bond is the strong electrostatic attraction between these positive ions and the surrounding sea of electrons.
Metallic bonding: positive metal ions in a lattice surrounded by a sea of delocalised electrons. The electrons carry charge → metals conduct electricity. Layers of ions can slide over each other → metals are malleable and ductile.
Ionic: electrons transferred, metal + non-metal, giant ionic lattice, high melting point, conducts when molten or dissolved.
Covalent: electrons shared, non-metal + non-metal, simple molecules (low m.p.) or giant covalent (very high m.p.), mostly poor conductors.
Metallic: delocalised electrons, metal atoms only, giant metallic lattice, good conductors, malleable, generally high melting points (though mercury is liquid at room temperature).
The key exam skill is justifying properties from structure: ionic compounds have high melting points because many strong electrostatic attractions must be overcome simultaneously. Simple covalent molecules like water have low boiling points because intermolecular forces, not covalent bonds, are broken during boiling. Giant covalent structures like diamond and silicon dioxide have very high melting points because many strong covalent bonds must be broken throughout the entire lattice.