4.2.2.1The three states of matter
The three states of matter are solid, liquid and gas. Melting and freezing take place at the melting point; boiling and condensing take place at the boiling point.
The three states of matter can be represented by a simple model. In this model, particles are represented by small solid spheres. Particle theory can help to explain melting, boiling, freezing and condensing.
The amount of energy needed to change state from solid to liquid, and from liquid to gas, depends on the strength of the forces between the particles of the substance. The nature of the particles involved depends on the type of bonding and the structure of the substance. The stronger the forces between the particles, the higher the melting point and boiling point of the substance.
That sentence is the hinge of the whole topic. Every property in the rest of this sub-topic — why sodium chloride melts at 801 °C and oxygen at −218 °C — is the same argument applied to different particles and different forces.
Worked example
A substance has a melting point of −7 °C and a boiling point of 59 °C. Predict its state at 20 °C and at 100 °C.
- Below the melting point the substance is solid; between melting and boiling point it is liquid; above the boiling point it is gas.
- 20 °C is above −7 °C, so it has melted.
- 20 °C is below 59 °C, so it has not yet boiled.
- 100 °C is above 59 °C, so it has boiled.
Answer: Liquid at 20 °C and gas at 100 °C. (The substance is bromine.)
You should be able to predict the states of substances at different temperatures given appropriate data, explain the different temperatures at which changes of state occur in terms of energy transfers and types of bonding, and recognise that atoms themselves do not have the bulk properties of materials.
That last point matters more than it looks. Melting point, hardness and conductivity belong to the structure, not to the atom — a single copper atom is not shiny, malleable or a conductor. Those are properties of many atoms bonded together.
Common mistake
When a molecular substance boils, the covalent bonds inside the molecules are not broken — only the forces between the molecules. Water vapour is still H₂O. Saying "the covalent bonds break when water boils" loses the mark.
- States of matter
- solid, liquid and gas; melting and freezing take place at the melting point, boiling and condensing at the boiling point
4.2.2.1Limitations of the simple particle modelHigher tier
Limitations of the simple model above include that in the model there are no forces, that all particles are represented as spheres, and that the spheres are solid.
You should be able to explain the limitations of particle theory in relation to changes of state when particles are represented by solid inelastic spheres which have no forces between them.
Common mistake
The limitation is not that the drawing is small or two-dimensional. Answer in terms of the three assumptions the specification names — no forces, spheres, solid spheres — and say what each one prevents the model from explaining.
4.2.2.2State symbols
In chemical equations, the three states of matter are shown as (s), (l) and (g), with (aq) for aqueous solutions.
You should be able to include appropriate state symbols in chemical equations for the reactions in this specification.
- (s)
- Solid.
- (l)
- Liquid — a pure liquid, such as water itself.
- (g)
- Gas.
- (aq)
- Aqueous — dissolved in water.
Worked example
Add state symbols to: Mg + 2HCl → MgCl₂ + H₂, where the acid and the salt solution are in water.
- Magnesium is a metal added as a solid ribbon or powder: (s).
- Hydrochloric acid is dissolved in water: (aq).
- Magnesium chloride stays dissolved in the solution: (aq).
- Hydrogen bubbles off as a gas: (g).
Answer: Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g)
Common mistake
(l) and (aq) are not interchangeable. (l) means the pure substance is a liquid; (aq) means it is dissolved in water. Hydrochloric acid is (aq), not (l), because it is hydrogen chloride dissolved in water.
4.2.2.3Properties of ionic compounds
Ionic compounds have regular structures — giant ionic lattices — in which there are strong electrostatic forces of attraction in all directions between oppositely charged ions.
These compounds have high melting points and high boiling points because of the large amounts of energy needed to break the many strong bonds.
When melted or dissolved in water, ionic compounds conduct electricity because the ions are free to move and so charge can flow.
Knowledge of the structures of specific ionic compounds other than sodium chloride is not required.
Common mistake
Ionic compounds do not conduct as solids because the ions cannot move — not because they have no charge. The charged particles are there all along; it is the movement that is missing.
4.2.2.4Properties of small molecules
Substances that consist of small molecules are usually gases or liquids that have relatively low melting points and boiling points.
These substances have only weak forces between the molecules — intermolecular forces. It is these intermolecular forces that are overcome, not the covalent bonds, when the substance melts or boils.
The intermolecular forces increase with the size of the molecules, so larger molecules have higher melting and boiling points.
These substances do not conduct electricity because the molecules do not have an overall electric charge.
Worked example
Explain why the boiling point rises from methane (−162 °C) to butane (−1 °C).
- Both are small molecules held together by covalent bonds inside and intermolecular forces between.
- Boiling overcomes the intermolecular forces, not the covalent bonds.
- Butane molecules are larger than methane molecules.
- Intermolecular forces increase with the size of the molecules, so more energy is needed to separate butane molecules.
Answer: Butane has larger molecules, so stronger intermolecular forces, so more energy is needed to overcome them and the boiling point is higher.
You should be able to use the idea that intermolecular forces are weak compared with covalent bonds to explain the bulk properties of molecular substances. Say both halves: the covalent bonds inside are strong, and the forces between the molecules are weak. Only the second half is overcome on melting or boiling, and that is why the melting points are low.
- Intermolecular forces
- the weak forces between molecules, which are what is overcome when a substance made of small molecules melts or boils — not the covalent bonds inside the molecules
4.2.2.5Polymers
Polymers have very large molecules. The atoms in the polymer molecules are linked to other atoms by strong covalent bonds. The intermolecular forces between polymer molecules are relatively strong, and so these substances are solids at room temperature.
A polymer sits between the small molecules and the giant covalent structures. It is still made of separate molecules, so the forces between them can still be overcome — but the molecules are so long that those forces add up along the whole chain, which is why a polymer is a solid where a small molecule of the same elements would be a gas.
You should be able to recognise polymers from diagrams showing their bonding and structure — look for a repeating unit, drawn in brackets, with the bonds continuing out of each side.
4.2.2.6Giant covalent structures
Substances that consist of giant covalent structures are solids with very high melting points. All of the atoms in these structures are linked to other atoms by strong covalent bonds. These bonds must be overcome to melt or boil these substances.
Diamond and graphite (forms of carbon) and silicon dioxide (silica) are examples of giant covalent structures.
That is the difference from a small molecule in one line: here there is nothing weak to overcome. Melting means breaking covalent bonds themselves, which is why the melting points are measured in thousands of degrees.
You should be able to recognise giant covalent structures from diagrams showing their bonding and structure — every atom is joined by lines to its neighbours, and the bonds run off the edge of the picture rather than stopping.
- Giant covalent structure
- a solid in which all the atoms are linked by strong covalent bonds, giving a very high melting point; diamond, graphite and silicon dioxide are examples
4.2.2.7Properties of metals and alloys
Metals have giant structures of atoms with strong metallic bonding. This means that most metals have high melting and boiling points.
In pure metals, atoms are arranged in layers, which allows metals to be bent and shaped. Pure metals are too soft for many uses and so are mixed with other metals to make alloys, which are harder.
You should be able to explain why alloys are harder than pure metals in terms of distortion of the layers of atoms in the structure of a pure metal.
The explanation runs like this. In a pure metal the atoms are all the same size and sit in regular layers, so the layers can slide over one another and the metal bends easily. An alloy contains atoms of a different size, which disrupt the regular arrangement and distort the layers. The layers can no longer slide over each other, so the alloy is harder.
Common mistake
Alloys are harder because the layers cannot slide, not because the bonds are stronger. The metallic bonding is the same kind of bonding — it is the geometry of the layers that changes.
4.2.2.8Metals as conductors
Metals are good conductors of electricity because the delocalised electrons in the metal carry electrical charge through the metal.
Metals are good conductors of thermal energy because energy is transferred by the delocalised electrons.
Both answers are the same sentence with one word changed, and both depend on the delocalised electrons being free to move through the whole structure — which is exactly what metallic bonding provides.
Before you practise
- Melting point depends on the strength of the forces between the particles: the stronger the forces, the higher the melting and boiling points.
- Higher tier: the simple model has no forces, treats all particles as spheres, and makes those spheres solid.
- State symbols: (s), (l), (g), and (aq) for aqueous solutions.
- Ionic compounds are giant lattices with strong electrostatic attraction in all directions; they conduct only when molten or dissolved.
- Small molecules have low melting points because only weak intermolecular forces are overcome — the covalent bonds stay intact.
- Intermolecular forces increase with the size of the molecules.
- Polymers are very large molecules with relatively strong intermolecular forces, so they are solids at room temperature.
- Giant covalent structures — diamond, graphite, silicon dioxide — need covalent bonds broken to melt, so melting points are very high.
- Alloys are harder than pure metals because different-sized atoms distort the layers so they cannot slide.
- Metals conduct electricity and heat because of their delocalised electrons.