Chemical Bonding

Why Do Atoms Bond?

Isolated atoms are generally unstable because their outer electron shells are not full. Atoms achieve stability by reaching a full outer shell — usually 8 electrons (the octet rule), or 2 electrons for hydrogen and helium. They do this in three main ways: by transferring electrons (ionic bonding), sharing electrons (covalent bonding), or pooling electrons in a sea (metallic bonding).

The type of bonding depends on which elements are involved. Metal + non-metal → ionic. Non-metal + non-metal → covalent. Metal + metal → metallic. These three rules cover the vast majority of O Level bonding questions.

Ionic Bonding

Ionic bonding occurs when a metal transfers one or more electrons to a non-metal. The metal atom loses electrons and becomes a positive ion (cation). The non-metal atom gains those electrons and becomes a negative ion (anion). The two oppositely charged ions attract each other strongly — this electrostatic attraction is the ionic bond.

Worked Example — Sodium Chloride (NaCl)

• Sodium (Na) has electron arrangement 2, 8, 1 — one outer electron.
• Chlorine (Cl) has electron arrangement 2, 8, 7 — seven outer electrons, one short of a full shell.
• Sodium transfers its outer electron to chlorine.
• Na becomes Na⁺ (2, 8 — now has a full outer shell).
• Cl becomes Cl⁻ (2, 8, 8 — now has a full outer shell).
• Na⁺ and Cl⁻ attract each other and form an ionic bond.

Properties of Ionic Compounds

• High melting and boiling points: strong electrostatic forces between ions require a lot of energy to break.
• Conduct electricity when molten or dissolved in water: ions are free to move and carry charge. Solid ionic compounds cannot conduct because ions are held in fixed positions.
• Soluble in water (most): water molecules surround and separate the ions.
• Hard but brittle: striking a crystal shifts layers so that like charges align and repel, shattering the structure.

Giant Ionic Lattice

Ionic compounds do not exist as individual pairs of ions. Instead, they form a giant ionic lattice — a three-dimensional arrangement where each positive ion is surrounded by negative ions and vice versa. This is why ionic compounds are solid at room temperature and have high melting points.

Covalent Bonding

Covalent bonding occurs when two non-metal atoms share one or more pairs of electrons. Each atom provides one electron to the shared pair. By sharing, both atoms achieve a full outer shell without gaining or losing electrons entirely.

Types of Covalent Bonds

• Single bond: one shared pair of electrons (e.g., H–H in H₂, H–Cl in HCl).
• Double bond: two shared pairs of electrons (e.g., O=O in O₂, C=O in CO₂).
• Triple bond: three shared pairs of electrons (e.g., N≡N in N₂).

Worked Example — Water (H₂O)

• Oxygen has 6 outer electrons — needs 2 more for a full shell.
• Each hydrogen has 1 outer electron — needs 1 more.
• Oxygen shares one electron with each hydrogen atom.
• Both hydrogens achieve 2 electrons (full for hydrogen); oxygen achieves 8 electrons.
• Result: two O–H single covalent bonds.

Simple Molecular Structures

Small covalently bonded molecules (H₂O, CO₂, CH₄, HCl) form simple molecular structures. The covalent bonds within each molecule are strong, but the forces between separate molecules (intermolecular forces) are weak. This gives simple molecular compounds low melting and boiling points, and they are often gases or liquids at room temperature.

Giant Covalent Structures

Some covalent substances form giant covalent lattices where covalent bonds extend throughout the entire structure. Diamond and silicon dioxide (SiO₂) are examples. These substances have very high melting points because thousands of strong covalent bonds must all be broken simultaneously. Diamond does not conduct electricity because all electrons are held in bonds; graphite does conduct because each carbon atom forms only three bonds, leaving one delocalised electron free to move.

Properties of Simple Covalent Molecules

• Low melting and boiling points: weak intermolecular forces.
• Do not conduct electricity: no charged particles free to move.
• Insoluble in water (most): but polar covalent molecules like HCl dissolve readily.

Metallic Bonding

Metallic bonding occurs in metals. Metal atoms release their outer electrons into a shared "sea" of delocalised electrons. The positive metal ions (cations) are held in a regular lattice, surrounded by and attracted to this sea of free electrons. The electrostatic attraction between the positive ions and the electron sea is the metallic bond.

Properties of Metals Explained by Bonding

• Good electrical conductors: the delocalised electrons carry charge through the structure.
• Good thermal conductors: delocalised electrons transfer heat energy quickly.
• High melting points: strong attraction between positive ions and the electron sea (though this varies — mercury melts at −39 °C because it is unusual).
• Malleable and ductile: layers of ions can slide over each other without breaking the bond, because the electron sea continues to hold the structure together. Ionic compounds cannot do this — shifting layers brings like charges together and the structure shatters.

Comparing the Three Bond Types

• Ionic: metal + non-metal; electron transfer; high melting point; conducts when molten or dissolved.
• Covalent (simple): non-metal + non-metal; electron sharing; low melting point; does not conduct.
• Covalent (giant): non-metal + non-metal; extensive sharing; very high melting point; does not conduct (except graphite).
• Metallic: metal + metal; electron sea; high melting point; conducts in all states.

Fully Worked Exam Examples

Example 1 — Identify Bond Type

Question: Magnesium oxide (MgO) has a melting point of 2852 °C and conducts electricity when molten. What type of bonding does it have? Explain.

Answer: Ionic bonding. Magnesium is a metal and oxygen is a non-metal — metal + non-metal indicates ionic. The high melting point shows strong electrostatic forces between Mg²⁺ and O²⁻ ions. Conduction when molten confirms ions become free to move.

Example 2 — Dot and Cross Diagram Description

Question: Describe the bonding in a molecule of chlorine (Cl₂).

Answer: Each chlorine atom has 7 outer electrons. Each atom shares one electron with the other, forming one single covalent bond. Both atoms now have 8 outer electrons (a full outer shell). The shared pair of electrons is the covalent bond between the two chlorine atoms.

Example 3 — Properties from Structure

Question: Substance A melts at 3550 °C and does not conduct electricity in any state. Suggest its bonding and structure.

Answer: Giant covalent structure. The extremely high melting point indicates very strong bonds extending throughout the structure (not just weak intermolecular forces). The inability to conduct in any state rules out ionic and metallic bonding. Substance A is likely diamond (carbon).

Example 4 — Why Metals Are Malleable

Question: Explain why metals are malleable but ionic compounds are brittle.

Answer: In metals, layers of positive ions can slide over each other while still being surrounded by the electron sea, so the bonding is maintained. In ionic compounds, shifting the layers brings ions of the same charge next to each other — like charges repel strongly, so the structure shatters.

Common Mistakes and How to Avoid Them

• Saying ionic compounds conduct electricity as solids. They do not — ions are fixed in the lattice. They only conduct when molten or dissolved.
• Confusing intermolecular forces with covalent bonds. Covalent bonds are strong. The weak forces between molecules (not within them) are what give simple covalent substances low melting points.
• Forgetting graphite is an exception. Graphite is a giant covalent structure that does conduct — always mention this if the question involves carbon allotropes.
• Drawing dot-and-cross diagrams with wrong electron counts. Always count outer electrons first, then pair up the sharing.
• Stating that metals are brittle. Metals are malleable; ionic compounds and giant covalent structures are hard and brittle.

Exam Tips

• Identify the elements first — metal + non-metal = ionic; non-metal + non-metal = covalent; metal + metal = metallic. This rule alone answers most "type of bonding" questions.
• Use melting point and conductivity together to narrow down structure: high melting point + conducts when liquid = ionic; low melting point + never conducts = simple covalent; very high melting point + never conducts = giant covalent; conducts in all states = metallic.
• For dot-and-cross questions, write down each atom's outer electron count before drawing anything.
• Always describe the ion charges in ionic bonding — Na⁺ and Cl⁻, not just "sodium and chloride."
• Mention the electron sea explicitly when explaining metallic bonding — it earns the mark that vague answers miss.

Quick Summary

• Ionic: metal transfers electrons to non-metal; forms ions; giant lattice; high melting point; conducts when molten/dissolved.
• Covalent: non-metals share electrons; simple molecules have low melting points; giant covalent has very high melting points.
• Metallic: metals share electron sea; high melting point; conducts in all states; malleable.

Bonding determines every physical property of a substance. If you can identify the bonding type from the elements involved, you can predict the melting point, conductivity, solubility, and hardness before being told — which is exactly what examiners test.