Ionic bonding = transfer of electrons

An ionic bond forms between a metal and a non-metal. The metal atom transfers its outer electron(s) to the non-metal. Both atoms end up with a full outer shell, but they are now charged ions:

• The metal loses electrons → a positive ion (cation), e.g. Na → Na⁺.
• The non-metal gains electrons → a negative ion (anion), e.g. Cl → Cl⁻.

Watch out: in ionic bonding electrons are transferred, never shared. The "dot-and-cross" just tracks where each electron came from — there is still only one kind of electron.

Which ion forms?

Ions build a giant lattice

Millions of oppositely-charged ions pack together in a regular 3-D giant ionic lattice, held by strong electrostatic attraction in every direction. In sodium chloride each Na⁺ is surrounded by 6 Cl⁻, and each Cl⁻ by 6 Na⁺.

Properties explained by the lattice:

• High melting & boiling points — lots of energy is needed to break the many strong ionic bonds.
• Conducts when molten or dissolved, but NOT when solid — the ions only become free to move (and carry charge) once the lattice breaks down.
• Most ionic compounds are soluble in water.

Watch out: a solid ionic compound does not conduct — the ions are locked in place. It only conducts when molten or dissolved, when the ions are free to move.

Name the structure from its clues

Read the properties, then tap the structure they describe.

Covalent bonding = sharing electrons

A covalent bond forms between two non-metals. Instead of transferring electrons, the atoms share a pair of electrons — one from each atom. A shared pair = one covalent bond.

Watch out: covalent = sharing, ionic = transfer. No ions form in a covalent bond, so a single covalent molecule has no overall charge.

Predict the bonding

Simple molecular substances

Many covalent substances exist as small, separate molecules — e.g. carbon dioxide (CO₂) and iodine (I₂). The bonds inside each molecule are strong, but the forces between molecules are only weak van der Waals' forces.

Properties:

• Low melting & boiling points — only the weak van der Waals' forces between molecules need to be broken (so many are gases/liquids at room temperature).
• Do not conduct electricity — there are no free charged particles (the molecules are neutral).
• Many are insoluble in water.

Giant covalent: diamond & graphite

Some covalent substances are not small molecules — they are giant lattices of thousands of atoms joined by strong covalent bonds. Carbon forms allotropes (different structures of the same element): diamond and graphite (and graphene). Carbon can form four covalent bonds.

Both have very high melting/boiling points (many strong covalent bonds), but they differ:

• Diamond — very hard; does not conduct (no free electrons). Used in cutting tools.
• Graphite — soft & slippery (layers slide); conducts electricity. Used as a lubricant and in pencils.

Watch out: graphite conducts because each carbon has one delocalised electron (it only uses 3 of its 4 bonds), free to move and carry charge. Diamond uses all four bonds, so it has no free electrons and cannot conduct.

Why does graphite conduct?

Graphene — a single layer

Graphene is a single layer of graphite — one atom thick — with each carbon covalently bonded to three others in hexagons. It is:

• Very strong and light (strong covalent bonds, only one atom thick).
• A good electrical conductor (one delocalised electron per carbon, as in graphite).

Because it is strong, light, cheap and conducts, graphene has many future uses — for example in batteries and solar cells. The term allotrope means a different structural form of the same element: diamond, graphite and graphene are all allotropes of carbon.

Link: graphene shares graphite's conducting trick (delocalised electrons) but, being a single sheet, is far stronger per atom.

Metallic bonding = sea of electrons

In a metal, the atoms lose their outer electrons to become a giant lattice of positive ions sitting in a sea of delocalised electrons. The strong electrostatic attraction between the positive ions and the negative delocalised electrons is the metallic bond.

Properties explained by this model:

• Good electrical conductors — the delocalised electrons are free to move and carry charge.
• Malleable & ductile — layers of ions can slide over each other while the sea of electrons holds them together, so the metal can be hammered into shape or drawn into wire.
• High melting & boiling points — much energy is needed to break the strong metallic bonds.

Predict the bonding type

Drag each substance into the bonding type its elements give. Tap a chip, then tap a box.

Alloys — harder than pure metals

An alloy is a mixture of two or more elements, at least one a metal, and the mixture has metallic properties. The different-sized atoms distort the layers, so the layers can no longer slide over each other easily — making alloys harder than the pure metal.

CCEA example: jewellery gold is an alloy with silver, copper and zinc. Purity is in carats — 24 carat = pure gold; 18 carat = 75% gold (18 ÷ 24 × 100).

Your turn — carat gold

Structure decides the properties

The whole topic comes down to matching a structure to its properties — especially melting point and electrical conductivity:

• Giant ionic lattice — high mp; conducts only when molten/dissolved; usually soluble.
• Covalent molecular — low mp; never conducts; often insoluble.
• Giant covalent — very high mp; usually doesn't conduct (graphite/graphene are the exceptions).
• Metallic — high mp; conducts (solid or molten); malleable & ductile.

Match structure to property

Tap a structure on the left, then its key property on the right.

What structure is it?

The points to know

You've covered all of CCEA Unit 1.3 — the three bonds, four structures, and how structure explains properties. Press Finish to see your score.