Substance types, physical properties and models

Ionic compounds contain positive and negative ions. An ion is an atom or group of atoms with an overall electrical charge. Typical ionic compounds form between metals and non-metals.

In a solid ionic compound, the ions form a giant ionic lattice: a regular, repeating arrangement extending throughout the crystal. Strong electrostatic attractions act between oppositely charged ions in all directions. There are no separate small molecules.

Ionic compounds have high melting and boiling points because a large amount of energy is needed to overcome these strong attractions. This is very different from separating molecules held together by weak forces.

Their electrical conductivity depends on their state:

• Solid: the ions have charge, but they are fixed in position and cannot move through the lattice. The solid does not conduct electricity.
• Molten: melting allows the ions to move and carry charge, so the liquid conducts electricity.
• Aqueous solution: if the compound dissolves in water, its ions can move through the solution and carry charge.

Many ionic compounds dissolve in water, although not all do. Dissolving separates ions from the lattice and allows them to mix throughout the water. A molten compound is the compound itself in its liquid state; an aqueous solution contains the compound dissolved in water.

Simple molecular substances: strong bonds within, weak forces between

A simple molecular substance contains separate small molecules. Within each molecule, atoms are joined by strong covalent bonds, formed by shared pairs of electrons. Between molecules there are much weaker intermolecular forces.

When a simple molecular substance melts or boils, molecules move apart from one another. The intermolecular forces are overcome, but the covalent bonds within the molecules remain intact. Relatively little energy is needed, so these substances typically have low melting and boiling points. Many are gases or liquids at room temperature.

For example, when water boils, water molecules separate from one another: they do not split into hydrogen and oxygen atoms. This shows why strong covalent bonds can exist in a substance with a relatively low boiling point.

Typical simple molecular substances are poor electrical conductors. Their molecules are uncharged, and they have neither mobile ions nor delocalised electrons to carry charge. Being able to move is not enough: the moving particles must also carry electrical charge.

Many simple molecular substances are insoluble in water, but some dissolve because their molecules can form sufficient attractions with water molecules. Carbon dioxide and ammonia are examples that dissolve. Dissolving alone does not guarantee electrical conduction: a solution needs mobile charged particles. Some molecular substances form ions when they interact with water, allowing their solutions to conduct.

Giant covalent substances: an extended bonded network

In a giant covalent substance, atoms are joined by strong covalent bonds throughout an extended network. This is not a collection of separate small molecules.

These substances have very high melting points because many strong covalent bonds must be broken to disrupt the structure. Large amounts of energy are needed. Diamond and graphite are examples; their detailed structures explain their different properties.

Giant covalent substances are generally insoluble in water. Attractions to water cannot overcome the strong covalent bonding that holds the network together.

Electrical conductivity depends on whether the structure contains mobile charged particles. Diamond does not conduct because it has no delocalised electrons or mobile ions. Graphite does conduct because it has delocalised electrons. ‘Covalent’ therefore does not automatically mean ‘non-conducting’: the particular structure matters.

Metals: mobile electrons and sliding layers

A metal has a giant structure of closely packed positive metal ions surrounded by delocalised electrons. These electrons come from the outer shells of metal atoms and are not attached to one particular atom. Metallic bonding is the strong electrostatic attraction between the positive ions and the delocalised electrons.

Most metals have high melting and boiling points because considerable energy is needed to overcome strong metallic bonding. Metals conduct electricity as solids and when molten because their delocalised electrons can move through the structure and carry charge. Unlike a solid ionic compound, a solid metal already has mobile charge carriers.

Metals are malleable: they can be hammered or bent into different shapes without breaking. Layers of metal ions can slide past one another while attraction to the delocalised electrons continues to hold the structure together.

Metals are insoluble in water, although some react with it. Reacting is a chemical change that produces new substances; it is not simply dissolving an unchanged metal.

Most metals are shiny solids at room temperature, have high density and conduct electricity well. High density means a large mass in a given volume. Most non-metals, by contrast, have low boiling points and are poor electrical conductors. These are general patterns, not definitions: solid giant covalent non-metals and conducting graphite show why exceptions matter.

What models show—and what they leave out

Atoms and electrons are too small to see directly in ordinary classroom observations. Models make selected features understandable, but no single model gives a complete picture.

A dot and cross diagram shows electrons involved in bonding. Different symbols indicate which atom the electrons originally came from; they do not mean that there are two kinds of electron. In ammonia, three shared pairs represent three covalent bonds, while nitrogen also has a lone pair. However, this diagram does not show the molecule’s three-dimensional shape or the relative sizes of its atoms accurately.

A two-dimensional displayed formula uses element symbols and lines for covalent bonds. It clearly shows which atoms are joined. For ammonia, it shows one nitrogen atom connected to three hydrogen atoms. It does not accurately show relative atom sizes, bond lengths or the molecule’s three-dimensional shape. A flat drawing should not be taken as evidence that the real molecule is flat.

A three-dimensional ball and stick model helps show the spatial arrangement of atoms and the shape of a molecule. The balls represent atoms and the sticks represent bonds. However, real bonds are not physical rods, and the model exaggerates the gaps between atoms to make connections visible. It also does not show electron movement.

Three-dimensional models can make spatial arrangements clearer than flat diagrams, but they are still simplified and usually static. A small model of a giant structure shows only a fragment, not the whole lattice. Choose a model according to the information needed: electron sharing, connections between atoms, or three-dimensional arrangement.

Metals and non-metals

• Metals are malleable: layers slide while metallic bonding remains.
• Most metals: shiny solids, high melting points, high density, good electrical conductivity.
• Most non-metals: low boiling points and poor electrical conductivity. Exceptions exist.

Model limitations

• Dot and cross: shows bonding electrons and their origins, not accurate 3D shape or relative atom sizes.
• 2D displayed formula: shows atom connections, not accurate 3D shape, sizes or bond lengths.
• 3D ball and stick: shows spatial arrangement, but exaggerates gaps and does not show electron movement.
• A model of a giant structure usually shows only a small fragment.