Bonding, structure, properties.

Part 1Three bonds, three rules

There are only three types of strong chemical bond, and which one forms depends entirely on whether the atoms are metals or non-metals. Learn the three rules below and you can predict the bonding in almost anything.

Ionic bonding happens between a metal and a non-metal. The metal atom gives electrons to the non-metal atom. This makes charged particles called ions, held together by the attraction between opposite charges.

Covalent bonding happens between two non-metals. The atoms share pairs of electrons rather than transferring them.

Metallic bonding happens in metals and alloys — between metal atoms only.

Part 1Transferring electrons

When a metal reacts with a non-metal, the metal atoms lose their outer electrons and the non-metal atoms gain them. Both then have a full outer shell, like a noble gas. Losing or gaining electrons makes them charged — they become ions.

A metal atom that loses electrons becomes a positive ion (a cation). A non-metal atom that gains electrons becomes a negative ion (an anion). The number of electrons transferred matches the group: Group 1 forms 1+, Group 2 forms 2+, Group 6 forms 2−, Group 7 forms 1−.

A dot-and-cross diagram shows electrons from one atom as dots and from the other as crosses, so you can track exactly where each electron came from. For ionic bonding you draw the ions in square brackets with the charge written outside, top-right.

Part 2Giant ionic lattices & properties

Ionic compounds don't exist as little molecules. The ions pack into a giant ionic lattice — a regular, repeating 3-D arrangement where every positive ion is surrounded by negative ions and vice-versa. The electrostatic forces of attraction act in all directions.

Those forces are strong, so it takes a lot of energy to break them. That explains the properties:

High melting and boiling points — lots of energy is needed to overcome the many strong attractions. Conduct electricity when molten or dissolved, but not when solid — the ions are charged, so they can carry charge only when they are free to move. In a solid they're locked in place.

Part 1Sharing a pair of electrons

When two non-metal atoms bond, neither will give up its electrons — so instead they share. A covalent bond is a shared pair of electrons. Each shared pair lets both atoms count those electrons towards a full outer shell.

Each atom usually forms enough bonds to fill its outer shell. Hydrogen forms 1 bond, chlorine 1, oxygen 2, nitrogen 3 and carbon 4. A double bond is two shared pairs (four electrons), as in O₂ and CO₂.

Part 1Small molecules & polymers

Small molecules (like H₂O, CO₂, Cl₂) have strong covalent bonds inside each molecule, but only weak intermolecular forces between molecules. When these substances melt or boil, it's only the weak forces between molecules that break — not the covalent bonds. That's why they have low melting and boiling points and are often gases or liquids at room temperature.

As molecules get bigger, the intermolecular forces get stronger, so melting and boiling points rise. None of these substances conduct electricity, because the molecules are not charged and there are no free electrons.

Polymers are very large molecules — long chains of many repeating units joined by covalent bonds. The intermolecular forces between these long chains are large, so polymers are usually solid at room temperature.

Part 2Giant covalent structures

In a giant covalent structure, a huge number of atoms are joined by strong covalent bonds in a continuous network. Breaking it means breaking covalent bonds, so these substances have very high melting points. The three you must know are diamond, graphite and silicon dioxide.

Diamond: each carbon is covalently bonded to four others in a rigid 3-D lattice. This makes it very hard with a very high melting point. It does not conduct electricity — every outer electron is held in a bond.

Graphite: each carbon bonds to only three others, forming flat layers of hexagons. The layers are held together by weak forces, so they can slide over each other — that's why graphite is soft and slippery (a good lubricant). One electron per atom is delocalised, so graphite conducts electricity and thermal energy.

Silicon dioxide (silica, the structure of sand) is a giant covalent lattice of silicon and oxygen — hard, with a very high melting point.

Part 1Graphene — one layer

Graphene is a single layer of graphite — just one sheet of carbon atoms arranged in hexagons, one atom thick. Like graphite, each carbon bonds to three others, leaving a delocalised electron, so graphene conducts electricity. It is also very strong for its mass. These properties make it useful in electronics and in composite materials.

Part 2Fullerenes & nanotubes

Fullerenes are molecules of carbon shaped as hollow balls or tubes. Their structures are based mainly on hexagonal rings of carbon (with some rings of five or seven atoms to allow them to curve). The first one discovered was Buckminsterfullerene (C₆₀) — a hollow sphere of 60 carbon atoms, the most stable fullerene.

Fullerenes can be used to deliver drugs into the body (a molecule can be carried inside the cage), as lubricants, and as catalysts because of their huge surface area.

Carbon nanotubes are cylindrical fullerenes — long tubes with a very high length-to-diameter ratio. They have useful properties for nanotechnology, electronics and materials: they conduct both electricity and heat, and have very high tensile strength (they resist being pulled apart), so they can reinforce materials.

Part 1The sea of delocalised electrons

In a metal, the atoms are packed in a regular arrangement. The outer electrons leave their atoms and become delocalised — free to move throughout the whole structure. This leaves a lattice of positive metal ions sitting in a "sea" of delocalised electrons.

Metallic bonding is the strong electrostatic attraction between the positive ions and the delocalised electrons. It explains the key properties:

Good conductors of electricity and heat — the delocalised electrons are free to move and carry charge and energy through the metal. High melting and boiling points — the metallic bonds are strong. Malleable (can be bent and shaped) — the layers of ions can slide over each other without breaking the bonding.

Part 2Why alloys are harder

A pure metal has atoms of the same size arranged in neat layers, so the layers slide over each other easily — pure metals are soft. An alloy is a mixture of a metal with at least one other element (often another metal). The added atoms are a different size, so they distort the layers and stop them sliding so easily. This makes alloys harder than the pure metal, which is why most metals we use day to day are alloys.

Part 1Changes of state & the limits of models

In the solid, liquid and gas states, the particles are the same — only their arrangement, movement and the energy in them change. Heating gives particles more energy; at a high enough temperature the forces between them are overcome and the state changes. Melting and boiling happen when there is enough energy to overcome the forces of attraction between particles.

The more energy needed to break those forces, the higher the melting and boiling points. So the strength of the forces tells you the state at room temperature.

The diagram above is the particle model, and it has limits. It treats particles as solid spheres with no forces between them, but in reality particles are not solid, not all spheres, and the forces between them vary from substance to substance. State symbols (s), (l), (g) and (aq) tell you the state.

Part 2Nanoparticles & nanoscience

Nanoscience is the study of very small particles — nanoparticles, which are between 1 and 100 nanometres (nm) across (1 nm = 10⁻⁹ m). They contain only a few hundred atoms. For comparison: coarse particles (dust, PM10) are about 1×10⁻⁵ to 2.5×10⁻⁶ m; fine particles (PM2.5) are 100 nm to 2500 nm.

The special thing about nanoparticles is their surface area to volume ratio. As a particle gets smaller, this ratio gets much larger — in fact, if the side of a cube is divided by 10, the surface area to volume ratio increases by 10. This means nanoparticles can be very effective catalysts, and that smaller amounts may be needed than for the same material in bulk.

Uses include catalysts, sun creams, cosmetics, deodorants and medicine. But because nanoparticles are so new, their effects on health and the environment are not fully understood, so there are concerns about their long-term safety.