Types of Chemical Bonding: KS3 Chemistry Guide

Picture an atom as a tiny positive nucleus surrounded by a cloud of negative electrons — that mental model is the starting point for all chemical bonding. Bonds form because atoms are more stable when their outermost electron shell is full, and they achieve this either by transferring electrons or sharing them. At KS3, three types of bonding matter: ionic, covalent, and metallic.

Why do atoms form bonds at all?

Most atoms have outer electron shells that are not full. A full outer shell is more stable (lower energy) than a partially filled one. Atoms achieve a full outer shell in two main ways:

Transfer electrons — one atom gives electrons to another (ionic bonding).

Share electrons — two atoms each contribute electrons to a shared pair (covalent bonding).

Metals have a third option: they pool their electrons into a shared "sea" (metallic bonding).

What is ionic bonding?

Ionic bonding occurs between a metal and a non-metal. The metal atom loses one or more electrons, becoming a positively charged ion (cation). The non-metal atom gains those electrons, becoming a negatively charged ion (anion). Opposite charges attract, forming a strong electrostatic bond.

Think of it like a donation: the metal is generous with electrons; the non-metal is eager to receive them. Neither atom shares — there is a complete handover.

Step | What happens (sodium chloride example)
1 | Sodium (Na) has 1 electron in its outer shell
2 | Chlorine (Cl) has 7 electrons in its outer shell (needs 1 more for a full shell of 8)
3 | Na transfers its 1 outer electron to Cl
4 | Na becomes Na⁺ (positive); Cl becomes Cl⁻ (negative)
5 | Na⁺ and Cl⁻ attract strongly — an ionic bond forms

The result is a giant ionic lattice — millions of alternating positive and negative ions packed in a regular 3D grid, which is why ionic compounds form crystals and have high melting points.

What is covalent bonding?

Covalent bonding occurs between two non-metals. Instead of transferring electrons, the atoms share pairs of electrons. Each shared pair forms one covalent bond, and both atoms count the shared electrons toward filling their outer shell.

Picture two atoms reaching toward each other and clasping hands — the shared electrons are the handshake neither atom lets go of.

Simple molecule examples:

Molecule | Bond type | Shared pairs
H₂ (hydrogen gas) | Single covalent | 1 pair
O₂ (oxygen gas) | Double covalent | 2 pairs
N₂ (nitrogen gas) | Triple covalent | 3 pairs
H₂O (water) | Single covalent | 1 pair per H–O bond
CO₂ (carbon dioxide) | Double covalent | 2 pairs per C=O bond

Simple covalent molecules have low melting points because the forces between molecules are weak (even though the bonds within the molecule are strong).

What is metallic bonding?

In a metallic solid, metal atoms release their outer electrons into a shared pool — a "sea of delocalised electrons" that flows freely throughout the whole structure. The positive metal ions (now stripped of their outer electrons) are held together in a lattice, and the sea of electrons holds them in place like a glue.

This model explains several key properties of metals:

Electrical conductivity: The delocalised electrons can move through the lattice, carrying charge.

Thermal conductivity: Moving electrons transfer energy (heat) quickly through the metal.

Malleability and ductility: The positive ions can slide past each other without the bonds breaking, because the electron sea moves with them. This is why metals can be hammered into sheets or drawn into wires.

Shiny appearance: Free electrons interact with light and reflect it, giving metals their lustre.

How do the three types of bonding compare?

Feature | Ionic | Covalent | Metallic
Between | Metal + non-metal | Two non-metals | Metal atoms only
Electron behaviour | Transferred | Shared in pairs | Delocalised (sea)
Typical state at room temp | Solid (crystal lattice) | Often gas or liquid | Solid
Melting point | High | Low (simple molecules) | Variable (mostly high)
Conducts electricity? | Only when dissolved/melted | No (usually) | Yes, always
Example | NaCl (table salt) | H₂O (water) | Fe (iron)

What clues tell you which type of bond is present?

Three questions to ask about any substance:

What elements are involved? Metal + non-metal → likely ionic. Two non-metals → likely covalent. Pure metal or metal alloy → metallic.

Does it conduct electricity? Ionic compounds do not conduct when solid (ions are fixed) but do when dissolved or melted. Metals always conduct. Simple covalent molecules usually do not.

What is the melting point? Very high → likely ionic or giant metallic. Low (below 300 °C) → likely simple covalent molecules.

Frequently asked questions

What is the difference between an atom and an ion?

An atom is electrically neutral — it has equal numbers of protons (positive) and electrons (negative). An ion is a charged particle formed when an atom gains or loses electrons. A metal atom loses electrons to become a positive ion (cation); a non-metal atom gains electrons to become a negative ion (anion). The number of protons does not change — only the electron count.

Why do ionic compounds have such high melting points?

Ionic compounds form giant lattices where millions of ions are bound together by strong electrostatic attraction (opposite charges attracting). To melt the compound, you need enough energy to overcome ALL of those attractions simultaneously. That takes a very high temperature — for example, sodium chloride melts at 801 °C.

Can a molecule have both ionic and covalent bonds?

Yes. Many polyatomic ions (ions made of more than one atom) contain covalent bonds within the ion AND ionic bonds between the ion and its partner. For example, in sodium nitrate (NaNO₃), the nitrate ion (NO₃⁻) is held together by covalent bonds, but the ionic bond is between Na⁺ and NO₃⁻.

Why are metals good conductors of electricity but ionic compounds are not (in solid form)?

In a metal, the delocalised electrons are already free to move through the lattice, so a small voltage causes a current immediately. In a solid ionic compound, the ions are locked in fixed positions in the lattice and cannot move — there is no charge flow. Only when the ionic compound is dissolved in water or melted do the ions become free to move and conduct electricity.

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