What you'll learn
Bonding, structure and the properties of matter is the unit of AQA GCSE Combined Science: Trilogy that explains why substances behave as they do. The central principle is that properties follow from structure and bonding, and almost every question in this unit is an application of that one idea. There are three types of chemical bond — ionic, covalent and metallic — and which one forms depends entirely on whether the elements involved are metals or non-metals. By the end you should be able to predict the bonding in a substance from the elements present, describe and explain the three states of matter and the changes between them, explain the properties of ionic compounds, small molecules, giant covalent structures, polymers and metals in terms of their structure, and describe the structure and uses of diamond, graphite, graphene and fullerenes. This unit is assessed on Chemistry Paper 1.
Key terms and definitions
Ionic bond — the strong electrostatic force of attraction between oppositely charged ions
Covalent bond — a shared pair of electrons between two non-metal atoms
Metallic bond — the strong electrostatic attraction between positive metal ions and delocalised electrons
Ion — an atom or group of atoms with an overall charge, formed by losing or gaining electrons
Giant ionic lattice — a regular three-dimensional arrangement of oppositely charged ions held by strong electrostatic forces in all directions
Intermolecular forces — the weak forces of attraction between separate molecules, much weaker than covalent bonds
Giant covalent structure — a structure in which all the atoms are linked by strong covalent bonds throughout
Delocalised electron — an electron free to move through a structure, found in metals and in graphite
Polymer — a very large molecule made of many small repeating units joined by covalent bonds
Alloy — a mixture of a metal with other elements, harder than the pure metal because the layers are distorted
Core concepts
The three states of matter
The three states are solid, liquid and gas. In the simple particle model, particles are represented as small solid spheres with no forces between them, which is useful but has limitations: real particles are not solid spheres, they are not all the same size, and there are forces between them that the model ignores.
In a solid, particles are held closely together in fixed positions and can only vibrate, so solids keep their shape and volume. In a liquid, particles are still close together but can move past one another, so liquids flow and take the shape of the container. In a gas, particles are far apart and move quickly in all directions, so gases fill their container.
Melting and freezing occur at the melting point; boiling and condensing occur at the boiling point. In each case energy is supplied to overcome the forces between particles, or released as those forces re-form. The amount of energy needed depends on the strength of those forces and on the type of bonding, which is why melting point is such a good clue to structure.
Ionic bonding
Ionic bonding occurs between a metal and a non-metal. The metal atom loses its outer electrons to become a positive ion, and the non-metal atom gains them to become a negative ion. Both end up with the electronic structure of a noble gas.
The charge on the ion follows directly from the group. Group 1 metals form ions with a charge of plus 1, Group 2 plus 2, Group 6 non-metals minus 2 and Group 7 minus 1. In sodium chloride, one sodium atom transfers one electron to one chlorine atom, giving sodium ions with a charge of plus 1 and chloride ions with a charge of minus 1.
The oppositely charged ions attract one another strongly in all directions, forming a giant ionic lattice. The bond is the electrostatic attraction itself, and describing it in those words is what earns the mark.
Properties of ionic compounds
Ionic compounds have high melting and boiling points because the many strong electrostatic forces between oppositely charged ions throughout the lattice need a great deal of energy to overcome.
They do not conduct electricity when solid, because the ions are in fixed positions and cannot move. When melted or dissolved in water the ions become free to move, so a charge can flow and the compound conducts. This distinction between the solid and molten or aqueous states is one of the most frequently examined points in the unit.
Covalent bonding and small molecules
Covalent bonding occurs between non-metals. Atoms share pairs of electrons so that each achieves a full outer shell. The covalent bond itself is very strong.
Substances made of small molecules, such as hydrogen, chlorine, oxygen, water, ammonia, methane and hydrogen chloride, have low melting and boiling points, and many are gases or liquids at room temperature. The reason is critical and often stated wrongly: melting or boiling breaks the weak intermolecular forces between molecules, not the strong covalent bonds within them. As the size of the molecule increases, the intermolecular forces increase and so do the melting and boiling points.
Small molecules do not conduct electricity, because the molecules have no overall charge and there are no free electrons or ions.
Giant covalent structures
In a giant covalent structure every atom is joined to its neighbours by strong covalent bonds throughout the whole structure. Diamond, silicon dioxide and graphite are the examples. These substances have very high melting and boiling points because a large amount of energy is needed to break the many strong covalent bonds.
In diamond, each carbon atom forms four covalent bonds in a rigid three-dimensional structure, which makes it very hard and gives it a very high melting point. It has no delocalised electrons, so it does not conduct electricity.
In graphite, each carbon atom forms only three covalent bonds, producing layers of hexagonal rings with no covalent bonds between the layers. The layers can slide over each other, so graphite is soft and slippery and is used as a lubricant. Because each carbon uses only three of its four outer electrons for bonding, one electron per atom is delocalised, so graphite conducts electricity and thermal energy — the only non-metal you meet that does.
Graphene and fullerenes
Graphene is a single layer of graphite, one atom thick. It is very strong, conducts electricity because of its delocalised electrons, and has potential uses in electronics and composites.
Fullerenes are molecules of carbon atoms with hollow shapes, based mainly on hexagonal rings but also containing rings with five or seven carbon atoms. The first to be discovered was Buckminsterfullerene, with the formula C60, which has a spherical shape. Fullerenes are being investigated for delivering drugs into the body, as lubricants and as catalysts.
Carbon nanotubes are cylindrical fullerenes with a very high length to diameter ratio. They have useful electrical and thermal properties and very high tensile strength, so they are used in nanotechnology, electronics and materials.
Polymers
Polymers are very large molecules in which many small units are joined together by strong covalent bonds. The intermolecular forces between polymer molecules are relatively strong compared with those between small molecules, because the molecules are so long, so polymers are solids at room temperature.
Metallic bonding and the properties of metals
In a metal, the atoms are arranged in a giant structure of layers. The outer electrons of each atom are delocalised and free to move throughout the whole structure, and the strong electrostatic attraction between the positive metal ions and these delocalised electrons is the metallic bond.
This structure explains every property. Metals have high melting and boiling points because the metallic bonds are strong and much energy is needed to overcome them. They conduct electricity and thermal energy because the delocalised electrons can move through the structure carrying charge and energy. They are malleable and ductile because the layers of atoms can slide over one another.
Pure metals are often too soft for practical use, so most everyday metals are alloys. Mixing in atoms of a different size distorts the layers and makes it harder for them to slide, so the alloy is harder than the pure metal.
Worked examples
Example 1: Explaining conduction in ionic compounds (4 marks)
Explain why solid sodium chloride does not conduct electricity but molten sodium chloride does.
In the solid, the sodium ions and chloride ions are held in fixed positions in a giant lattice by strong electrostatic forces, so although the ions are charged they cannot move and no charge can flow. When the compound is melted, enough energy is supplied to overcome these forces so the ions are free to move. The charged ions can then move through the liquid, so a current flows.
Example 2: Comparing diamond and graphite (4 marks)
Both diamond and graphite are made only of carbon atoms. Explain why graphite conducts electricity but diamond does not, and why graphite is softer.
In diamond each carbon atom forms four covalent bonds, using all four of its outer electrons, so there are no delocalised electrons available to carry charge. In graphite each carbon atom forms only three covalent bonds, so one electron per atom is delocalised and free to move through the structure, allowing graphite to conduct. Graphite is softer because its layers are held together only by weak forces, with no covalent bonds between them, so the layers can slide over each other.
Example 3: Explaining a low boiling point (3 marks)
Methane boils at minus 161 degrees Celsius even though the bonds between carbon and hydrogen are strong. Explain why.
Methane consists of small separate molecules. Boiling requires only the weak intermolecular forces between the molecules to be overcome, not the strong covalent bonds within each molecule. Because these intermolecular forces are weak, very little energy is needed, so the boiling point is very low.
Common mistakes and how to avoid them
The single biggest error in this unit is saying that covalent bonds are broken when a simple molecular substance melts or boils. The covalent bonds stay intact; the intermolecular forces are overcome. Examiners look for this distinction specifically and it appears on nearly every paper.
Students often describe intermolecular forces as weak bonds. They are forces of attraction between molecules, and calling them bonds invites confusion with the covalent bonds inside the molecule.
Another regular slip is saying ionic compounds conduct when solid because they contain ions. Containing ions is not enough; the ions must be free to move.
In metallic bonding answers, many students write that electrons are free without saying delocalised, or describe metals as containing atoms rather than positive ions. Both cost marks.
Finally, students often say alloys are harder because they contain two metals. The reason is that the different-sized atoms distort the regular layers, preventing them from sliding.
Exam technique for "Chemistry: Bonding, Structure and the Properties of Matter"
Almost every question here can be answered with a three-step chain: identify the structure, describe the forces holding it together, then link the strength of those forces to the property asked about. Practise writing that chain and unfamiliar substances stop being a problem.
To identify bonding, look at the elements. Metal plus non-metal is ionic. Non-metal plus non-metal is covalent. Metal alone is metallic. This decides the answer before you think about anything else.
When a question gives melting point data, use it as evidence. A melting point of a few hundred degrees or more suggests a giant structure; below about 100 degrees Celsius suggests small molecules.
State symbols matter in this unit. If a question asks about conductivity, check whether the substance is solid, molten or aqueous before answering, because the answer changes completely.
Quick revision summary
Ionic bonding transfers electrons from metal to non-metal, giving a giant lattice held by strong electrostatic attraction in all directions, with high melting points and conduction only when molten or dissolved. Covalent bonding shares pairs of electrons between non-metals; small molecules have low melting and boiling points because only weak intermolecular forces are overcome, and they do not conduct. Giant covalent structures such as diamond, silicon dioxide and graphite have very high melting points because many strong covalent bonds must be broken. Diamond is hard with four bonds per carbon and no delocalised electrons; graphite has three bonds per carbon, sliding layers and one delocalised electron per atom so it conducts. Graphene is one layer of graphite, and fullerenes including C60 and nanotubes are hollow carbon structures. Polymers are long covalent molecules, solid at room temperature. Metals are positive ions in a sea of delocalised electrons, giving conduction, malleability and high melting points, and alloys are harder because distorted layers cannot slide.