Three types of strong bond

Whenever atoms join, they're trying to reach a full outer shell (a stable noble-gas arrangement). They do this in one of three ways — and which one depends on whether the atoms are metals or non-metals:

• Ionic — a metal + a non-metal. Electrons are transferred from metal to non-metal, making charged ions.
• Covalent — non-metal + non-metal. Atoms share pairs of electrons.
• Metallic — in a metal (or alloy). Positive ions sit in a sea of delocalised electrons.

Quick rule: spot the elements first. Metal + non-metal → ionic. Two non-metals → covalent. Metal only → metallic. Get this and you can predict bonding before you draw a thing.

Which bond will form?

Transfer: making ions

In ionic bonding, a metal atom loses its outer electron(s) and a non-metal atom gains them. Both end up with full outer shells, but now they're charged ions:

• A metal atom loses electrons → a positive ion (cation), e.g. Na → Na⁺, Mg → Mg²⁺.
• A non-metal atom gains electrons → a negative ion (anion), e.g. Cl → Cl⁻, O → O²⁻.

The charge usually matches the group: Group 1 → 1+, Group 2 → 2+, Group 6 → 2−, Group 7 → 1−.

Misconception: ionic bonding is transfer, not sharing. The opposite charges then attract by a strong electrostatic force — and that force is the ionic bond.

Ionic dot-and-cross: NaCl

A dot-and-cross diagram shows where the electrons came from. Sodium (2,8,1) gives its one outer electron to chlorine (2,8,7), which then has 8. We show the ions in square brackets with their charge:

Read it right: dots vs crosses only label which atom an electron started on — all electrons are identical. The outer shell of Na⁺ is shown empty; Cl⁻ has the full eight.

Ionic structure → properties

Ions don't pair off — they stack into a huge, repeating giant ionic lattice, with each ion surrounded by oppositely-charged neighbours. The electrostatic forces act in all directions throughout the lattice.

That structure explains the properties:

• High melting & boiling points — many strong electrostatic forces need lots of energy to overcome.
• Conduct electricity only when molten or dissolved — then the ions are free to move and carry charge. When solid, the ions are locked in place, so it does not conduct.
• Often hard but brittle — shifting a layer puts like charges together, which repel.

Misconception: solid ionic compounds do not conduct. It's not "ionic = conductor" — it's "ionic conducts only when molten or dissolved", because only then can the ions move.

Explain the property

Sharing: covalent bonds

Two non-metals can't both grab electrons — so instead they share. A covalent bond is a shared pair of electrons, with one electron from each atom. The shared pair sits between the two nuclei and holds them together.

Misconception: covalent = sharing, not transfer — no ions form. Common small molecules you should be able to draw: H₂, Cl₂, O₂, HCl, H₂O, NH₃ and CH₄.

Simple molecular substances

Substances like water, methane, oxygen and CO₂ are made of small, separate molecules. The atoms inside a molecule are joined by strong covalent bonds — but the molecules are only weakly attracted to each other by intermolecular forces.

• Low melting & boiling points — melting or boiling only has to overcome the weak intermolecular forces, not the strong covalent bonds. Many are gases or liquids at room temperature.
• Bigger molecules → higher m.p./b.p. — larger molecules have stronger intermolecular forces.
• Do not conduct electricity — molecules are neutral, with no overall charge and no free electrons or ions.

Key wording: when a simple molecular substance melts, it is the intermolecular forces that break, not the covalent bonds. Examiners love this distinction.

Name that structure

Polymers

A polymer is a very large molecule made of many small units (monomers) joined in a long chain by strong covalent bonds. We draw a polymer as a repeat unit inside brackets, using a single line for each single bond, with an n for "many repeats".

Because the molecules are so long, the intermolecular forces between polymer chains add up to be relatively strong. That's why polymers (like poly(ethene)) are solids at room temperature — but their melting points are still much lower than giant ionic or giant covalent structures, since the chains themselves aren't covalently joined to each other.

Watch out: the strong covalent bonds run along each chain. It's the forces between chains that you overcome on melting — stronger than in a small molecule, but weaker than the bonds within the chain.

Giant covalent: diamond vs graphite

In a giant covalent (macromolecular) structure, a huge network of atoms is joined entirely by strong covalent bonds — e.g. diamond and silicon dioxide. Diamond and graphite are both pure carbon, but their structures differ, so their properties do too:

All giant covalent structures have very high melting points (you must break many strong covalent bonds). But:

• Diamond — every carbon bonded to 4 others, no spare electrons → very hard, and it does not conduct electricity.
• Graphite — every carbon bonded to only 3 others, in layers. Each carbon has one delocalised electron, so graphite conducts electricity; the layers slide, so it is soft and slippery (used as a lubricant and in pencils).

Misconception: graphite conducts because of its delocalised electrons (one per carbon) — not because "all carbon conducts". Diamond, with no spare electrons, does not conduct.

Predict the bonding

From the elements involved, tap the bonding type each substance has.

Graphene & fullerenes

Carbon forms other remarkable structures, all from the same covalent network idea:

• Graphene — a single layer of graphite, just one atom thick. It conducts electricity (delocalised electrons), is extremely strong and light — useful in electronics and composite materials.
• Fullerenes — hollow carbon molecules built from rings of carbon (mostly hexagons, sometimes pentagons). The first discovered was Buckminsterfullerene, C₆₀, a hollow sphere shaped like a football.
• Carbon nanotubes — fullerenes rolled into cylinders with a very high length-to-diameter ratio. They're strong, conduct, and are used to reinforce materials, in electronics, and in drug delivery.

Link it up: graphene = one graphite layer; C₆₀ and nanotubes are fullerenes. All rely on carbon's covalent bonding plus, in many cases, delocalised electrons that let them conduct.

Which carbon structure?

Metals: a sea of electrons

In a metal, atoms pack into a regular lattice and their outer electrons become delocalised — they leave the atoms and are free to move through the whole structure. So a metal is a lattice of positive ions sitting in a sea of delocalised electrons. The strong attraction between the positive ions and this sea is the metallic bond.

• Good conductors of heat and electricity — the delocalised electrons are free to move and carry charge/energy.
• Malleable (bendy) — the layers of ions can slide over each other, so metals can be bent and shaped.
• High melting points — strong metallic bonds need lots of energy to break.

Alloys: harder metals

An alloy is a mixture of a metal with one or more other elements. The added atoms are usually a different size, so they distort the regular layers of ions.

In a pure metal, identical layers can slide easily — that's why pure metals are soft. In an alloy the distorted layers can't slide over each other so easily, so the alloy is harder and stronger than the pure metal. (Steel, brass and bronze are everyday examples.)

Why harder: different-sized atoms disrupt the layers, blocking them from sliding. The metallic bonding itself is unchanged — it's the broken-up arrangement that does it.

Structure ↔ property

Tap a structure on the left, then the property it explains on the right.

Limitations of bonding models

Dot-and-cross diagrams, ball-and-stick models and 2-D drawings are all useful, but none is perfect — each leaves something out:

• Dot-and-cross shows where electrons came from, but not the 3-D shape, the sizes of atoms, or that electrons are really identical and don't sit still.
• Ball-and-stick shows shape and bond angles, but exaggerates the space between atoms and makes bonds look like solid rods.
• 2-D drawings of giant lattices only show a small part — the real structure repeats in all directions and is three-dimensional.

Exam point: if asked to "evaluate a model", say what it shows well and what it can't show. No single representation captures everything.

Nanoparticles

Nanoparticles are tiny particles with diameters from about 1 nm to 100 nm (a nanometre is 10⁻⁹ m) — only a few hundred atoms across. They are smaller than fine particles, and far smaller than anything you can see.

Their key feature is a very high surface-area-to-volume ratio: as a particle gets smaller, a much bigger fraction of its atoms is on the surface. (Useful scale fact: if the side of a cube decreases by a factor of 10, the surface-area-to-volume ratio increases by a factor of 10.) Because of this, nanoparticles can have different properties from the same material in bulk — and only a tiny quantity is needed to be effective.

Don't muddle the units: milli- = 10⁻³, micro- = 10⁻⁶, nano- = 10⁻⁹. Nanoparticles sit at the nanometre scale — 1 to 100 nm.

Nanoparticles: uses & risks

That huge surface area makes nanoparticles useful in many ways:

• Catalysts — a large surface area means even a small amount speeds up reactions efficiently.
• Sun creams — titanium dioxide nanoparticles block harmful UV without leaving a white film; silver nanoparticles are antibacterial (used in deodorants and wound dressings).
• Stronger, lighter materials (e.g. carbon-nanotube composites) and tiny electrical circuits / sensors.

But there are risks: because they are so small, nanoparticles could enter the body — even the bloodstream — and their long-term effects on health and the environment are not yet fully known. Many people argue more testing is needed before they're used widely.

Balanced view: a good answer weighs the benefits (high surface area, small amounts needed) against the risks (unknown effects on health/environment).

Surface-area-to-volume ratio

Smaller = bigger ratio

Which ion forms?

Why does it conduct?