Ionic, Covalent and Metallic Bonding
Why Atoms Bond
Atoms bond together because doing so lowers their energy and gives each atom a full outer electron shell — the same stable arrangement as a noble gas. There are three main types of chemical bond, each arising from a different way of achieving this stability: ionic bonding, covalent bonding, and metallic bonding. The type of bonding determines a substance's properties — its melting point, conductivity, and physical strength.
Ionic: metal + non-metal, electrons transferred. Covalent: non-metal + non-metal, electrons shared. Metallic: metal atoms, electrons delocalised. Bond type determines properties — knowing which type applies is the first step in any bonding question.
Ionic Bonding
When a metal reacts with a non-metal, electrons are transferred from the metal atom to the non-metal atom. The metal loses one or more electrons to form a cation (positive ion), while the non-metal gains electrons to form an anion (negative ion). Both ions now have full outer shells and are stable.
The strong electrostatic attraction between the oppositely charged ions is the ionic bond. In sodium chloride (NaCl), for example, sodium loses one electron to become \(Na^+\) and chlorine gains one electron to become \(Cl^-\). These ions pack together in a regular three-dimensional arrangement called a giant ionic lattice.
In a dot-and-cross diagram for ionic bonding, draw the outer-shell electrons of each atom (dots for one, crosses for the other). Show an arrow from the metal to the non-metal indicating the electron being transferred, then draw the resulting ions with square brackets and the charge written outside. Only outer-shell electrons are shown — inner shells are omitted for clarity.
For magnesium oxide (MgO): magnesium (2,8,2) loses 2 electrons to become \(Mg^{2+}\), and oxygen (2,6) gains 2 electrons to become \(O^{2-}\). Both now have an outer shell of 8.
Properties from Ionic Bonding
Giant ionic lattices have characteristic properties. Their melting and boiling points are high because large amounts of energy are needed to overcome the strong electrostatic forces acting in all directions. When an ionic compound is solid, the ions are fixed in position — it does not conduct electricity. When melted or dissolved in water, the ions are free to move and carry charge, so the compound does conduct electricity.
Ionic compounds: high melting points (strong forces in all directions), brittle (layers of ions shift so like charges repel), conduct when molten or dissolved (ions move freely), but not when solid (ions fixed in lattice).
Covalent Bonding
Non-metals bond with other non-metals by sharing pairs of electrons. Each shared pair counts as one covalent bond. Both atoms contribute one electron to each shared pair, and by sharing, both achieve a full outer shell. Examples of simple covalent molecules include:
- \(H_2\) — a single covalent bond (one shared pair)
- \(O_2\) — a double covalent bond (two shared pairs: \(O=O\))
- \(N_2\) — a triple covalent bond (three shared pairs: \(N \equiv N\))
- \(H_2O\) — two single bonds; \(CO_2\) — two double bonds (\(O=C=O\))
Simple covalent molecules have strong bonds within each molecule but only weak intermolecular forces between molecules. This is why many simple covalent substances are gases or liquids at room temperature — it takes very little energy to separate the molecules from each other.
A student is given two white solids — sodium chloride (NaCl) and glucose (\(C_6H_{12}O_6\)) — and asked to identify which is which using only a conductivity test on dissolved samples and a melting-point check.
Step 1: Predict the bonding. NaCl is metal + non-metal → ionic. Glucose is non-metal only → covalent.
Step 2: Predict the properties. NaCl solution conducts electricity (free ions). Glucose solution does not conduct (no ions). NaCl has a very high melting point; glucose melts at a much lower temperature.
Step 3: Carry out the tests and match observations to predictions. The solid that conducts in solution and resists melting is NaCl.
Metallic Bonding
In metals, the outer electrons leave their atoms and form a sea of delocalised electrons. The metal is then a lattice of positive ions surrounded by this mobile electron sea. The metallic bond is the strong electrostatic attraction between these positive ions and the delocalised electrons.
This structure directly explains the key properties of metals. The delocalised electrons carry charge easily, so metals are excellent electrical conductors. They also transfer thermal energy readily, making them good thermal conductors. When a force is applied, layers of ions can slide over each other without breaking the bond — the electron sea simply flows around them — so metals are malleable (can be hammered into shape) and ductile (can be drawn into wires).
Metallic bonding: positive metal ions in a sea of delocalised electrons. This explains conductivity (mobile electrons carry charge), malleability (layers slide without breaking), and high melting point (strong attraction throughout the lattice).
A very common error is to say that aluminium rusts. It does not. Rust is specifically iron(III) oxide formed on iron or steel. Aluminium does react with oxygen in the air, but it instantly forms a thin, hard layer of aluminium oxide (\(Al_2O_3\)) on its surface. This oxide layer is tightly bonded to the metal beneath and acts as a protective barrier, preventing further oxidation. This is why aluminium products — bike frames, aircraft fuselage panels, drinks cans — appear unchanged after years of exposure to air and moisture.
Iron's oxide (rust) is flaky and porous, allowing oxygen and water to continue reaching the metal underneath. Aluminium's oxide layer is self-sealing — once formed, it stops further attack. This is a key difference for any compare-or-contrast question on metal properties.