Why does diamond cut glass while graphite rubs off on your finger? Why does salt melt at 800°C when oxygen is a gas? In every case the answer comes down to the type of bonding between the particles and the structure that bonding creates. Once you can link a structure to its properties, you can reason your way through melting points, conductivity, strength and solubility for almost any substance an exam puts in front of you.
Most of C2 is for everyone — only a few parts are tier-specific:
- Everyone — states of matter; ionic bonding and giant ionic lattices; covalent bonding (simple molecules and giant covalent structures); metallic bonding and alloys; polymers; carbon allotropes; and how structure and bonding explain conductivity, melting point and solubility.
- Higher Tier H — limitations of the simple particle model (Section 1) and half equations for ion formation.
- Triple (Chemistry only) T — nanoparticles (Section 12).
1States of Matter & Particle Theory
All substances exist in one of three states of matter: solid, liquid, or gas. The state of a substance depends on how strongly its particles are held together and how much energy they have.
| State | Shape | Volume | Can be compressed? | Particle arrangement |
|---|---|---|---|---|
| Solid | Fixed | Fixed | No | Regular, close-packed — particles vibrate about fixed positions |
| Liquid | Not fixed — takes the shape of container | Fixed | No (virtually) | Close but irregular — particles can move past each other |
| Gas | Not fixed — fills the container | Not fixed | Yes — easily | Widely spaced, random movement — particles move rapidly in all directions |
Particle arrangements in the three states of matter. In a solid, particles are held in fixed positions and can only vibrate; in a liquid, particles are close together but can slide past each other; in a gas, particles are far apart and move rapidly in random directions.
Changes of State
When a substance is heated or cooled, it can change state. Changes of state are physical changes — they are reversible and no new substance is formed.
| Change of state | Direction | Energy change |
|---|---|---|
| Melting | Solid → Liquid | Energy absorbed |
| Freezing | Liquid → Solid | Energy released |
| Evaporation / Boiling | Liquid → Gas | Energy absorbed |
| Condensation | Gas → Liquid | Energy released |
| Sublimation | Solid → Gas (directly) | Energy absorbed — e.g. dry ice (solid CO₂) |
Melting: particles gain enough energy for their vibrations to overcome the forces holding them in fixed positions. They break free and can move past each other.
Boiling: particles gain enough energy to completely overcome the forces holding them close together and escape into the gas phase.
During melting and boiling, the temperature stays constant — all the energy supplied is used to break the forces between particles, not to raise the temperature.
Heating Curve
Heating curve for a pure substance. The temperature stays constant during melting and boiling because the energy supplied is used to break forces between particles rather than raise the temperature.
Students often think energy is not being supplied during the flat sections of a heating curve. Energy is still being supplied — it is just being used to break the forces between particles rather than to raise the temperature. The flat sections occur at the melting point and the boiling point.
State Symbols
In chemical equations, the physical state of each substance is shown using a state symbol in brackets after the formula:
| State symbol | Meaning | Example |
|---|---|---|
| (s) | Solid | NaCl(s) |
| (l) | Liquid | H₂O(l) |
| (g) | Gas | CO₂(g) |
| (aq) | Aqueous solution (dissolved in water) | NaCl(aq) |
When state symbols are asked for, include them carefully. They are especially important in ionic equations, precipitation reactions and electrolysis. They also often explain observations of a reaction e.g. bubbles may be caused by a gas, or a precipitate may appear if a product is solid.
The Particle Model and its Limitations
The simple particle model represents particles as small solid spheres. It is a useful model, but it is important to recognise that atoms themselves do not have the bulk properties of the materials they make up. For example, a single water molecule is not wet, does not flow, and does not have a surface tension — these are properties that only emerge when vast numbers of molecules interact together.
Limitations of the Simple Particle ModelH
The simple particle model treats particles as identical, solid spheres with no forces between them. In reality:
- Particles are not all the same size — atoms, molecules and ions vary in size.
- Particles are not solid spheres — atoms are mostly empty space.
- Forces do exist between particles (even in gases, intermolecular forces affect behaviour at high pressure and low temperature).
- The model does not explain energy changes during changes of state at the atomic level.
A student heats a pure, solid substance and measures its temperature every minute. The temperature rises steadily, then stays constant for several minutes, then starts rising again.
(a) What change of state is taking place during the flat section?
(b) Explain, in terms of particles, why the temperature does not change during this time.
Show answer
- (a) Melting — the substance is changing from solid to liquid at its melting point.
- (b) During melting, the energy supplied is used to break the forces holding particles in their fixed positions, rather than increasing the temperature. In this topic we will be learning the different possibilities for what these forces could be.
2Types of Chemical Bond
Chemical bonds form because atoms tend to become more stable by achieving a full outer electron shell — like the noble gases. There are three main types of chemical bond, and which type forms depends on the elements involved.
To work out which type of bond forms, you first need to identify whether each element is a metal or a non-metal.
Periodic table
Need a clue? Use the staircase on the periodic table: metals are generally on the left and centre; non-metals are on the right. Hydrogen is the exception.
Use the periodic table above. For each element in the compounds below, click Metal or Non-metal to get immediate feedback.
| Elements involved | Type of bond |
|---|---|
| Metal + non-metal | Ionic |
| Non-metal + non-metal | Covalent |
| Metals only | Metallic |
Using the rules above, classify each substance as ionic, covalent, or metallic. Click a button for immediate feedback and an explanation.
The noble gases (Group 0) already have full outer shells — this is why they do not normally form chemical bonds and are so unreactive.
3Ionic Bonding
Ionic bonding — the strong electrostatic attraction between oppositely charged ions. It occurs when a metal reacts with a non-metal. Electrons are transferred from the metal atom to the non-metal atom.
| Atom type | What happens | Ion formed | Example H |
|---|---|---|---|
| Metal | Loses electrons from its outer shell | Positive ion (cation) — fewer electrons than protons | Na → Na⁺ + e⁻ Mg → Mg²⁺ + 2e⁻ Al → Al³⁺ + 3e⁻ |
| Non-metal | Gains electrons into its outer shell | Negative ion (anion) — more electrons than protons | Cl₂ + 2e⁻ → 2Cl⁻ O₂ + 4e⁻ → 2O²⁻ |
The equations in the Example column are half equations — they show the electrons (e⁻) lost or gained as the ions form. Writing half equations is a Higher Tier skill; you'll meet them again in more detail in electrolysis (C4).
Metals are losers — they lose electrons and form positive ions.
Non-metals are gainers — they gain electrons and form negative ions.
The number of electrons lost or gained equals the charge on the ion. Atoms lose/gain enough electrons to achieve a full outer shell. For metals in Groups 1 and 2 and non-metals in Groups 6 and 7, the ions formed have the electronic structure of a noble gas (Group 0).
Anion → think A Negative ion. The start of the word spells out exactly what it is.
Ca+ion → the letter t looks like a + sign — a positive ion. Or if that doesn't stick: cats are pawsitive — and a cation is a positive ion.
The outer electron transfers from sodium to chlorine, forming Na⁺ and Cl⁻ — oppositely charged ions that attract each other electrostatically.
Example — Sodium Chloride (NaCl)
Sodium (Group 1) has 1 electron in its outer shell. Chlorine (Group 7) has 7 electrons in its outer shell and needs 1 more to be full.
The sodium atom transfers its outer electron to the chlorine atom. Sodium becomes Na⁺ (2, 8) and chlorine becomes Cl⁻ (2, 8, 8). The oppositely charged ions attract each other — this electrostatic attraction is the ionic bond.
Ionic compounds are also known as salts — in chemistry, a "salt" simply means a compound made of positive and negative ions held together by ionic bonds. There are thousands of them, but the example above is the most familiar of all: sodium chloride (NaCl) is ordinary table salt, the stuff you sprinkle on your food.
Remember that compounds have different physical and chemical properties to their respective elements. Sodium is a soft metal that reacts violently with water, and chlorine is a poisonous green gas, yet the compound they form is something completely safe that you eat every day.
AQA may ask you to draw a dot-and-cross diagram showing the outer electron shells of both atoms before and after ion formation. Show electrons transferred as a dot moving from the metal to the non-metal. After transfer, place the ions in square brackets with their charge shown outside: [Na]⁺ and [Cl]⁻.
Ionic Formulae
The overall charge on an ionic compound is always zero. The charges on the ions must balance. Use this to work out the formula.
But where do those charges come from? The charge on most ions can be worked out straight from the periodic table. Only the molecular ions follow no pattern — those you have to learn by heart.
Charge = the group number
Charge = the group number − 8
No pattern and no shortcut — you have to learn these off by heart.
For Groups 1, 2, 3, 6 and 7 you can always work the charge out from the periodic table, so there is nothing to memorise there — this is the skill AQA asks for directly. Transition metals break the group rule, but the Roman numeral in the name tells you the charge (e.g. iron(III) is Fe³⁺). The molecular ions in the amber box follow no pattern at all, so these are the ones to learn off by heart — you will need them to write formulae throughout the course.
Mg²⁺ has a 2+ charge — it needs exactly two Cl⁻ ions to balance.
| Compound | Ions present | Formula | Why |
|---|---|---|---|
| Sodium chloride | Na⁺ and Cl⁻ | NaCl | +1 and −1 balance exactly |
| Magnesium oxide | Mg²⁺ and O²⁻ | MgO | +2 and −2 balance exactly |
| Lithium oxide | Li⁺ and O²⁻ | Li₂O | Need 2× Li⁺ to balance one O²⁻ |
| Magnesium chloride | Mg²⁺ and Cl⁻ | MgCl₂ | Need 2× Cl⁻ to balance one Mg²⁺ |
| Aluminium oxide | Al³⁺ and O²⁻ | Al₂O₃ | 2× Al³⁺ (+6) balanced by 3× O²⁻ (−6) |
Students often describe ionic bonding as "sharing electrons". This is wrong. In ionic bonding, electrons are transferred (given away), not shared. Sharing electrons is covalent bonding. The ionic bond itself is the electrostatic attraction between oppositely charged ions — not the transfer of electrons.
🧪 Exam-style questions
Calcium and chlorine react to produce calcium chloride.
Describe what happens to calcium atoms and chlorine atoms when the ionic compound calcium chloride is formed.
Show answer
The official AQA mark scheme awards 1 mark for each of the four points below:
- Each calcium atom loses two electrons 1 mark (calcium is in Group 2)
- (and) each chlorine atom gains one electron 1 mark (chlorine is in Group 7)
- (so) one calcium atom reacts with two chlorine atoms 1 mark (each chlorine atom accepts only one of calcium’s two electrons)
- (to form) Ca²⁺ ions and Cl⁻ ions — or calcium ion(s) and chloride ion(s) 1 mark
Allow: just 1 mark for “calcium atoms lose electrons and chlorine atoms gain electrons” if the numbers are not given; “ions with full outer shells” for the final mark; and “energy level” in place of “shell”.
4Giant Ionic Structures
Giant ionic lattice — a regular, repeating three-dimensional arrangement of alternating positive and negative ions, held together by strong electrostatic attractions in all directions. Ionic compounds always form giant ionic lattices — there are no small, discrete molecules.
In sodium chloride, each Na⁺ ion is surrounded by 6 Cl⁻ ions, and each Cl⁻ ion is surrounded by 6 Na⁺ ions. This extends in all three dimensions throughout the crystal.
Ball-and-stick model of the NaCl giant ionic lattice — each ion is surrounded by six ions of the opposite charge.
AQA often shows a diagram of the NaCl lattice and asks you to describe it. Key phrases: regular repeating structure, alternating positive and negative ions, electrostatic attraction between oppositely charged ions.
Different representations of ionic structures all have limitations:
- Dot-and-cross diagrams show electron transfer clearly, but only represent two ions — they give no sense of the giant, repeating three-dimensional lattice.
- Ball-and-stick models show the arrangement of ions in 3D, but the sticks imply that bonds point in fixed directions, and the gaps between spheres are misleading — ions in the real lattice are touching.
- Two-dimensional diagrams of the lattice cannot fully represent the three-dimensional nature of the structure.
- Three-dimensional diagrams can show the repeating pattern but may be hard to interpret and still only represent a small portion of the lattice.
AQA may ask you to describe the limitations of a given type of diagram in the context of ionic structures.
Properties of Ionic Compounds
| Property | Observation | Explanation |
|---|---|---|
| Melting and boiling point | Very high — ionic compounds are solids at room temperature | A lot of energy is required to overcome the strong forces of electrostatic attraction between the oppositely charged ions. |
| Electrical conductivity (solid) | Does not conduct electricity | The ions are held in fixed positions in the lattice and cannot move. Charge cannot flow. |
| Electrical conductivity (molten/liquid) | Does conduct electricity | When melted, the lattice breaks apart. The ions are free to move and can carry charge through the liquid. |
| Electrical conductivity (dissolved in water) | Does conduct electricity | Water breaks up the lattice (it dissolves). The ions are free to move and carry charge. |
| Solubility | Many dissolve in water | Water molecules pull ions away from the lattice one by one, surrounding them and allowing them to disperse. |
Explain why sodium chloride does not conduct electricity as a solid, but does conduct when it is dissolved in water. Use ideas about ions in your answer.
Show answer
Solid sodium chloride: the Na⁺ and Cl⁻ ions are held in fixed positions in the giant ionic lattice by strong electrostatic forces of attraction. They cannot move, so charge cannot flow — the solid does not conduct electricity.
Dissolved in water: when sodium chloride dissolves, the lattice breaks apart and the ions separate. The Na⁺ and Cl⁻ ions are now free to move through the solution and so a charge can flow— the solution does conduct electricity.
5Covalent Bonding
Covalent bond — a shared pair of electrons between two non-metal atoms. Both atoms contribute one electron to the shared pair, and both atoms count the shared electrons as their own. This allows both atoms to achieve a full outer electron shell.
Covalent bonding occurs when non-metal atoms bond to non-metal atoms. Unlike ionic bonding, no electrons are transferred — instead, electrons are shared.
Single, Double and Triple Bonds
| Bond type | Shared pairs | Examples |
|---|---|---|
| Single bond | 1 shared pair (2 electrons) | H₂, Cl₂, HCl, H₂O, NH₃ |
| Double bond | 2 shared pairs (4 electrons) | O₂, CO₂ |
| Triple bond | 3 shared pairs (6 electrons) | N₂ |
Dot-and-Cross Diagrams
Dot-and-cross diagrams show the outer-shell electrons of each atom. One atom's electrons are shown as dots (·) and the other's as crosses (×). Only the outer shells are drawn. Shared pairs appear in the overlapping region between the two atoms' outer shells.
Non-bonding electrons (also called lone pairs) must be shown in dot-and-cross diagrams — don't forget them. In water, oxygen has two pairs of non-bonding electrons as well as its two bonding pairs. Leaving them out loses marks.
Build each molecule yourself by dragging electrons into place — nothing tells you how many to use.
- Aim: give every atom a full outer shell of 8 electrons. Hydrogen is the exception — it has only one shell, and the first shell is full with just 2 electrons.
- Bonding pair — drag an electron into the overlap between two atoms. Shared electrons count for both atoms.
- Non-bonding electrons (lone pairs) — drag onto an atom's shell, away from the overlap.
- Use dots or crosses — either is fine. Only the number and position are marked.
- Tap a placed electron to remove it.
- Marked like the exam: 1 mark for the bonding electrons, 1 mark for the non-bonding electrons.
Hydrogen gives one electron (a cross) and chlorine one electron (a dot) to the shared pair in the overlap — that pair scores the bonding mark. Chlorine's three lone pairs (six electrons) score the non-bonding mark.
Make sure you can draw a dot-and-cross diagram for each of these from memory — they are the molecules AQA requires: H₂, Cl₂, O₂, N₂, HCl, H₂O, NH₃ and CH₄.
Describe the covalent bonding in a water molecule (H₂O). Explain how both the oxygen and hydrogen atoms achieve a full outer shell.
Show answer
Oxygen has 6 electrons in its outer shell. Each hydrogen has 1 electron. The oxygen atom forms two covalent bonds — one with each hydrogen atom — by sharing one pair of electrons with each.
Each shared pair counts as 2 electrons for both the O and the H: hydrogen gains 1 shared electron to achieve a full shell of 2; oxygen gains 2 shared electrons to achieve a full shell of 8. The oxygen also has two pairs of non-bonding electrons (lone pairs).
Representing Simple Covalent Molecules
The covalent bonds in molecules and giant structures can be represented in the following forms:
Ammonia (NH₃) shown four ways — a dot-and-cross diagram, a 3-D space-filling model, a 2-D displayed formula and a ball-and-stick model. Each makes some features clear while hiding others, which is the basis of exam questions on the limitations of diagrams.
- Dot-and-cross diagrams show which electrons are shared, but give no information about the 3D shape of the molecule, bond angles, or the relative sizes of atoms.
- Ball-and-stick models show 3D shape and bond angles, but the sticks exaggerate the space between atoms, and the spheres are not drawn to a consistent scale.
- Two-dimensional structural formulae (lines for bonds) are clear and compact but cannot represent 3D geometry or non-bonding electrons.
- Three-dimensional diagrams of giant structures (e.g. diamond, SiO₂) show the repeating pattern, but represent only a tiny fragment of what is a continuous lattice.
You may be asked to describe the limitations of a specific type of diagram when representing a molecule or giant covalent structure.
6Simple Molecules
Many covalently bonded substances form simple molecular structures — small, discrete molecules with a fixed number of atoms. Examples include hydrogen (H₂), water (H₂O), ammonia (NH₃), methane (CH₄), and carbon dioxide (CO₂).
Properties of Simple Molecular Substances
| Property | Observation | Explanation |
|---|---|---|
| Melting and boiling point | Usually low — many are gases or liquids at room temperature | The covalent bonds inside the molecules are strong, but the intermolecular forces (forces between molecules) are weak. Only a little energy is needed to overcome the weak intermolecular forces and separate the molecules. |
| Electrical conductivity | Does not conduct electricity | Simple molecules have no overall electric charge and no free electrons or ions. There are no charged particles to carry charge. |
| Solubility | Many dissolve in water; some do not | Varies from one molecule to another. (Beyond GCSE — useful context: this depends on molecular polarity; polar molecules like HCl tend to dissolve, non-polar ones like CH₄ tend not to.) |
Animation — Intermolecular Forces vs Covalent Bonds
This is the key idea behind the low melting and boiling points of simple molecules. Watch what happens to a small group of water molecules (H₂O) when they are heated — notice which forces are overcome, and which ones stay intact.
Five water molecules (H₂O) sit close together. The solid lines are covalent bonds holding each H to its O inside the molecule.
The animation above is liquid water boiling — turning into steam (a gas). For pure water this happens at just 100 °C, because only the weak intermolecular forces between the molecules need to be overcome.
The strong covalent O–H bonds inside each molecule are not broken by boiling. Splitting water into hydrogen and oxygen by heat alone (breaking those covalent bonds) would need a temperature of over 2,000 °C — showing just how much stronger the covalent bonds are than the forces between the molecules.
Students often write: "Simple molecules have low melting points because the covalent bonds are weak." This is wrong. The covalent bonds inside the molecules are actually strong. The low melting point is because the intermolecular forces between the molecules are weak — these are what need to be overcome to melt or boil the substance. The covalent bonds themselves are not broken during melting or boiling.
Larger molecules have more electrons, which means stronger intermolecular forces between them. This is why larger simple molecules (like octane, C₈H₁₈) have higher boiling points than smaller ones (like methane, CH₄). Both are still simple molecular substances, but the intermolecular forces are stronger in the larger molecule.
Explain why hydrogen chloride (HCl) is a gas at room temperature. Use ideas about bonding and forces in your answer.
Show answer
Hydrogen chloride consists of simple molecules. There is a strong covalent bond between the H and Cl atoms within each molecule. However, the intermolecular forces between HCl molecules are weak.
Only a small amount of energy is needed to overcome these weak intermolecular forces and separate the molecules from each other. This means the boiling point of HCl is well below room temperature, so it exists as a gas at room temperature.
Water itself is a simple molecule (H₂O) with no overall charge and no free electrons or ions — so on its own it has nothing to carry a current. Pure distilled water does not conduct electricity.
Tap water, though, does conduct. This is because it is not pure: it contains soluble salts (ionic compounds) dissolved in it, picked up from rocks, soil and the pipes it travels through. Just as we saw for ionic compounds dissolved in water, the lattice breaks apart as these salts dissolve and the ions become free to move. These mobile ions carry charge through the water, so it conducts.
The key point: it is the dissolved ions that do the conducting, not the water molecules themselves. Remove the salts — as in pure distilled water — and there is nothing left to carry the charge.
7Polymers
Polymer — a substance made of very large molecules built from a small group of atoms (the repeating unit) joined end-to-end many times by covalent bonds. The letter n stands for the large number of repeating units. Polymers can be represented as:
[repeat unit]n
where n is a large number.
How a polymer is represented: the repeating unit is drawn inside brackets with subscript n, and a bond extends through each bracket to show the chain continues. Single covalent bonds are shown as lines. The example shown is poly(ethene).
Structure and Bonding
In polymers, the atoms within each polymer chain are joined by strong covalent bonds. The polymer chains themselves are held together by intermolecular forces between the chains. These intermolecular forces are relatively strong (stronger than those between small molecules) because the chains are long and have many points of contact.
Properties
| Property | Explanation |
|---|---|
| Solid at room temperature | The relatively strong intermolecular forces between the long polymer chains require more energy to overcome than those between small molecules — so polymers are solid at room temperature. |
| Does not conduct electricity | No free electrons or ions — the molecules have no overall charge. |
| Variable melting point | Depends on the strength of the intermolecular forces between chains. Longer or more entangled chains have stronger forces and higher melting points. |
Polymers are made of very large molecules, so an answer must mention molecules and the intermolecular forces between the chains. Treating a polymer like diamond or silicon dioxide (a giant covalent structure with bonds broken on melting) is wrong. Also avoid "covering all bases" by listing both covalent bonds and intermolecular forces as the thing overcome on melting — only the intermolecular forces between chains are overcome.
AQA may show you a diagram of a polymer structure and ask you to identify it. Look for: a long repeating chain structure with a bracket and subscript n; all bonds are covalent (no ions). Compare this to giant ionic lattices (charged ions, regular 3D arrangement) and simple molecules (small, discrete units).
In this topic you only need to recognise polymers from diagrams and explain their properties in terms of bonding — you do not need to know how they are made. Triple Chemistry students will explore polymers further in the Organic Chemistry topic later in the course, including how small molecules (monomers) join together by addition and condensation polymerisation.
8Giant Covalent Structures
Giant covalent structure — a substance in which a very large number of atoms are joined together by covalent bonds in a regular, repeating network. There are no individual molecules — the entire crystal is one giant molecule. Examples include diamond, graphite, silicon dioxide (SiO₂), and graphene.
General Properties
| Property | Observation | Explanation |
|---|---|---|
| Melting and boiling point | Very high | A huge number of strong covalent bonds must be broken to melt or boil the substance. This requires a very large amount of energy. |
| Hardness | Usually very hard | The rigid, 3D network of covalent bonds resists deformation. (Exception: graphite — see below.) |
| Electrical conductivity | Usually does not conduct | There are no free electrons or ions to carry charge. (Exception: graphite and graphene — see below.) |
| Solubility | Insoluble in water | The covalent bonds are too strong to be broken by water molecules. |
Silicon Dioxide (SiO₂)
Silicon dioxide (also known as silica or quartz) is a giant covalent structure in which each silicon atom is bonded to four oxygen atoms by covalent bonds, and each oxygen is bonded to two silicon atoms. This gives the formula SiO₂. The structure extends in all three dimensions, forming a very hard, high-melting-point solid used in glass manufacture and furnace linings.
Silicon dioxide (silica): each silicon atom (larger, purple) is bonded to four oxygen atoms (smaller, red), and each oxygen joins two silicons — a giant 3D covalent network with the formula SiO₂.
CO₂ is a simple covalent molecule — it is a gas at room temperature with a very low boiling point. SiO₂ is a giant covalent structure — it is a solid with a very high melting point. Both have similar-looking formulae, but their structures (and therefore properties) are completely different.
Explain why silicon dioxide (SiO₂) has a very high melting point.
Show answer
Silicon dioxide has a giant covalent structure. Silicon and oxygen atoms are joined by a very large number of strong covalent bonds throughout the entire crystal lattice.
To melt silicon dioxide, all of these strong covalent bonds must be broken. This requires a very large amount of energy, which means the melting point is very high.
9Carbon Allotropes
Allotropes — different structural forms of the same element, in which the atoms are bonded together in different ways. Carbon has several allotropes, including diamond, graphite, graphene, carbon nanotubes, and buckminster fullerene.
Diamond
In diamond, each carbon atom forms four covalent bonds to four other carbon atoms in a tetrahedral arrangement. This creates an extremely rigid, three-dimensional giant covalent network.
- Very hard — the 3D network of strong covalent bonds in all directions resists any attempt to break or scratch it. Diamond is the hardest natural substance.
- Very high melting point — all four bonds around each carbon must be broken to melt diamond; this requires an enormous amount of energy.
- Does not conduct electricity — all four outer electrons of each carbon are used in covalent bonds. There are no delocalised electrons free to move and carry charge.
- Insoluble in water — water cannot overcome the strong giant covalent network structure.
Uses of diamond: drill bits and cutting tools (hardness), jewellery (appearance and durability).
Diamond is a giant covalent structure — there are no molecules and therefore no intermolecular forces. Explaining diamond's high melting point by saying you "overcome intermolecular forces" contradicts the structure and loses the mark. You melt diamond by breaking many strong covalent bonds, not weak forces between molecules.
Diamond structure: every carbon bonds to four others, creating a rigid 3D network. All bonds are strong covalent bonds.
Graphite
In graphite, each carbon atom forms three covalent bonds to three other carbon atoms, creating flat hexagonal rings that form layers. The layers are held together by weak intermolecular forces. The fourth outer electron from each carbon atom is delocalised — free to move between the layers. Like metals, graphite has delocalised electrons, which is what lets it conduct electricity.
- Soft and slippery — the weak intermolecular forces between the layers are easily overcome. The layers can slide over each other, making graphite feel slippery and useful as a lubricant.
- High melting point — the strong covalent bonds within each layer still require a lot of energy to break.
- Conducts electricity — each carbon atom only uses 3 of its 4 outer electrons in covalent bonds. The fourth electron is delocalised and free to move along the layers, carrying charge.
- Conducts heat — the delocalised electrons also transfer thermal energy efficiently.
Uses of graphite: pencil 'lead' (soft and marks paper), lubricants (layers slide), electrodes in electrolysis (conducts electricity).
Graphite structure: layers of hexagonal rings. Strong covalent bonds within each layer; weak intermolecular forces between the layers, shown as dashed lines.
Comparison: Diamond vs Graphite
| Feature | Diamond | Graphite |
|---|---|---|
| Bonds per C atom | 4 covalent bonds | 3 covalent bonds |
| Structure | 3D tetrahedral network | 2D hexagonal layers |
| Hardness | Extremely hard | Soft and slippery |
| Conducts electricity? | No — no delocalised electrons | Yes — delocalised electrons |
| Melting point | Very high | Very high |
| Uses | Cutting tools, jewellery | Pencils, lubricants, electrodes |
Graphene
Graphene is a single layer of graphite — one atom thick — forming a flat sheet of hexagonal rings. It is a giant covalent structure with remarkable properties:
Graphene: a single layer of graphite — one atom thick. Each carbon atom is covalently bonded to three others, forming a flat sheet of hexagonal rings.
- Extremely strong — the covalent bonds form an uninterrupted 2D network; gram for gram, graphene is stronger than steel.
- Very lightweight — just one atom thick.
- Conducts electricity — delocalised electrons move freely across the sheet.
- Transparent — only one atom thick, so absorbs very little light.
- Potential uses: flexible electronics, transparent conductors, sensors, high-strength composite materials.
Carbon Nanotubes
Carbon nanotubes are cylindrical fullerenes — sheets of graphene rolled into a cylinder — with very high length to diameter ratios. They are giant covalent structures that are:
Carbon nanotube: a sheet of graphene rolled into a hollow cylinder.
- Very strong and lightweight — same reasons as graphene.
- Electrical conductors — delocalised electrons move along the nanotube.
- Potential uses: reinforcing materials (sports equipment, aircraft), nanotechnology, electronics.
Buckminster Fullerene (C₆₀)
Buckminster fullerene is a simple covalent molecule — unlike diamond and graphite, it is not a giant covalent structure. It contains 60 carbon atoms arranged in a hollow sphere, with formula C₆₀. Each carbon atom forms 3 covalent bonds, and the structure consists of 20 hexagons and 12 pentagons — like a football. More generally, fullerenes are based on hexagonal rings of carbon atoms, but may also contain rings with five or seven carbon atoms.
Buckminster fullerene (C₆₀): 60 carbon atoms in a hollow sphere — like a football.
- Formula: C₆₀. Simple (not giant) covalent molecule — discrete molecules, not an extended network.
- Hollow centre can trap other molecules — useful for drug delivery (medicine can be stored inside the cage and transported to a target in the body).
- Potential uses: catalysts, lubricants, drug delivery.
- Buckminster fullerene is a nanoparticle — its diameter is in the nanometre range.
Explain why graphite conducts electricity but diamond does not, even though both are allotropes of carbon.
Show answer
In diamond, each carbon atom forms four covalent bonds using all four of its outer electrons. There are no electrons left over — all electrons are held in bonds and cannot move freely. There are no delocalised electrons, so diamond cannot conduct electricity.
In graphite, each carbon atom forms only three covalent bonds. The fourth outer electron from each carbon atom is delocalised — free to move along the layers. These mobile electrons can carry charge, so graphite conducts electricity.
10Metallic Bonding & Alloys
Metallic bonding — the strong electrostatic attraction between a lattice of positive metal ions and the sea of delocalised electrons that surrounds them. Metal atoms lose their outer electrons, which become free to move throughout the entire metallic structure.
A regular lattice of positive metal ions sits in a ‘sea’ of delocalised electrons. Press a button to see how this structure explains a metal’s properties — each demonstration runs until you stop it.
Metallic bonding. Positive metal ions are arranged in a regular lattice, surrounded by a ‘sea’ of delocalised electrons that are free to move throughout the structure. The electrostatic attraction between the positive ions and the negative electron sea holds the metal together.
⇄ Why metals are malleable
The positive ions are arranged in layers. When a force is applied, the layers can slide over each other into new positions. This is why a metal can be bent and hammered into shape instead of shattering.
⚡ Why metals conduct electricity
Metals have delocalised electrons that are free to move through the whole structure. These electrons carry charge through the metal — this flow of electrons is the electric current.
🔥 Why metals conduct heat
The same delocalised electrons are free to move through the structure. They gain energy at the hot end and transfer this energy quickly to the cooler end.
Properties of Metals Explained by Metallic Bonding
| Property | Explanation |
|---|---|
| Conduct electricity | The delocalised electrons are free to move through the metal and carry charge, so they flow as an electric current. |
| Conduct heat | The delocalised electrons are free to move and transfer energy from the hot end to the cooler end. |
| Malleable and ductile (can be bent and shaped) | The layers of positive ions can slide over each other, so the metal can be hammered into shape (malleable) and drawn into wires (ductile). |
| High melting and boiling points | There are strong electrostatic attractions between the positive ions and the delocalised electrons. A large amount of energy is needed to overcome them, so most metals are solid at room temperature. |
| Shiny lustre | Delocalised electrons interact with light, reflecting it — giving metals their shiny appearance. |
Alloys
Alloy — a mixture of two or more elements where at least one is a metal. Common examples: steel (iron + carbon), brass (copper + zinc), bronze (copper + tin).
Pure metals have a regular arrangement of identical atoms. The layers of atoms can slide over each other easily, which makes pure metals soft and easy to deform.
In an alloy, atoms of a different element (with a different size) are introduced into the lattice. These different-sized atoms distort the regular layers, making it harder for them to slide over each other. This is why alloys are harder and stronger than the pure metals they are made from.
Two metal structures side by side — a pure metal and an alloy. Press the button to apply a force and watch whether the layers can slide over each other.
Why alloys are harder. Pure metal (left): identical ions in regular layers slide over each other easily. Alloy (right): atoms of a different size distort the layers, so they bump into each other and cannot slide — making the alloy harder and stronger.
Shape Memory Alloys (beyond the spec — useful context, not required knowledge)
Shape memory alloys can be deformed but then return to their original shape when heated. A well-known example is Nitinol (nickel–titanium alloy), used in dental braces — the alloy is bent to fit the teeth but slowly returns to its memorised shape as it warms to body temperature, gently straightening the teeth.
🧪 Exam-style questions
Iron is a metal. Describe how iron conducts thermal energy.
Show answer
The official AQA mark scheme awards 1 mark for each point:
- (Thermal) energy is transferred 1 mark (allow “heat is transferred”)
- by delocalised electrons 1 mark — the free electrons move through the structure, carrying energy from the hot end to the cooler end.
Pure iron is too soft for many uses.
Explain why mixing iron with other metals makes alloys which are harder than pure iron.
Show answer
The official AQA mark scheme awards 1 mark for each point:
- The alloy (mixture) contains different-sized atoms 1 mark
- (so the) layers are distorted 1 mark
- (so the) layers cannot easily slide over each other 1 mark (allow “atoms cannot slide over each other”)
Allow: “(positive / metal) ions” in place of “atoms” throughout.
11Does It Conduct Electricity?
Now that we have met every type of structure — ionic, simple molecular, giant covalent and metallic — we can pull the ideas together and answer one of the most common questions in this topic: will a given substance conduct electricity? You do not need to memorise a list. You can reason it out from first principles every time.
An electric current is a flow of charge. For a substance to conduct, it must contain charged particles that are free to move to carry that charge from one place to another.
There are only two kinds of mobile charge carrier you need to know:
- Delocalised electrons — negatively charged electrons that are not held in fixed positions and are free to move.
- Free-moving ions — charged ions that are able to move, rather than locked in place in a lattice.
So the question always reduces to one thing: does this substance contain delocalised electrons, or free-moving ions? If it has neither, it cannot conduct.
Substances That Conduct
| Substance | Charge carrier | Why it conducts |
|---|---|---|
| Metals (solid or molten) | Delocalised electrons | Metallic bonding leaves a ‘sea’ of delocalised electrons free to move throughout the lattice. |
| Graphite, graphene & carbon nanotubes | Delocalised electrons | Each carbon atom uses only 3 of its 4 outer electrons for bonds; the fourth is delocalised and free to move. |
| Ionic compounds — molten (liquid) | Free-moving ions | Melting breaks up the lattice, so the ions are free to move and carry charge. |
| Ionic compounds — dissolved in water | Free-moving ions | Dissolving breaks up the lattice, freeing the ions to move — this is why salt solution, and even tap water, conducts. |
Substances That Do Not Conduct
Everything else fails the test — it has no delocalised electrons and no free-moving ions, so there is nothing to carry the charge:
| Substance | Why it does not conduct |
|---|---|
| Ionic compounds — solid | The ions are present, but locked in fixed positions in the lattice — they cannot move. |
| Simple molecular substances (including pure water) | Molecules have no overall charge, no delocalised electrons and no free ions. |
| Giant covalent — diamond & silicon dioxide | All outer electrons are held in covalent bonds; none are delocalised. (Graphite is the exception.) |
| Polymers | Large covalent molecules — no delocalised electrons and no free ions. |
The classic trap: a solid ionic compound does not conduct, but the same compound conducts as soon as it is molten or dissolved. The bonding has not changed — what changes is whether the ions are free to move. Always check the state before you answer.
For each substance, state whether it conducts electricity and name the charge carrier responsible:
(a) copper wire (b) solid potassium chloride (c) potassium chloride dissolved in water (d) diamond
Show answer
- (a) Conducts — copper is a metal, so it has delocalised electrons that are free to move.
- (b) Does not conduct — the ions are locked in fixed positions in the solid lattice and cannot move.
- (c) Conducts — dissolving frees the ions, so the free-moving ions carry the charge.
- (d) Does not conduct — all of carbon’s outer electrons are held in covalent bonds; there are no delocalised electrons and no free ions.
12NanoparticlesT
Nanoparticle — a particle with a diameter in the range of 1–100 nm (nanometres). 1 nm = 1 × 10⁻⁹ m. Nanoparticles are only slightly larger than individual atoms and molecules.
Nanoscience — the study of nanoparticles and how they can be used.
| Scale | Typical size |
|---|---|
| Atom | ~0.1 nm = 1 × 10⁻¹⁰ m |
| Small molecule | ~0.5–2 nm |
| Nanoparticle | 1–100 nm = 1 × 10⁻⁹ – 1 × 10⁻⁷ m |
| Fine particles (PM2.5) | 100–2 500 nm (1 × 10⁻⁷ – 2.5 × 10⁻⁶ m) |
| Coarse particles (PM10) — dust | 2 500–10 000 nm (2.5 × 10⁻⁶ – 1 × 10⁻⁵ m) |
| Human hair | ~70 000 nm (70 µm) |
The SI prefixes on a logarithmic scale of size — each step down is ten times smaller.
Surface Area to Volume Ratio
As a particle gets smaller, its surface area to volume ratio increases dramatically. A nanoparticle has a very large surface area relative to its volume, compared to the same material as a large lump. For a cube, every time the side length decreases by a factor of 10, the surface area to volume ratio increases by a factor of 10.
Chemical reactions happen at surfaces. A nanoparticle exposes a much greater proportion of its atoms at the surface than a larger particle made of the same material. More surface atoms = more reaction sites = reactions happen much faster.
This is why nanoparticles are so effective as catalysts — a tiny mass of nanoparticles can catalyse a reaction that would need far more of the same material in bulk form.
Changed Properties
Nanoparticles can have very different properties from the same material in bulk (large-scale) form. This is mainly because of their very high surface area to volume ratio. Examples:
- Gold nanoparticles appear red or purple rather than gold, depending on their size.
- Titanium dioxide nanoparticles are transparent (unlike bulk TiO₂, which is white) — used in sun creams.
- Some materials become much more reactive, stronger, or conductive at the nanoscale.
Beyond the spec (useful context): you are not required to recall these specific materials or their individual properties — only the general principle that a high surface-area-to-volume ratio can give nanoparticles different properties, together with the application areas named below.
Uses of Nanoparticles
| Application | Material | Reason nanoparticles are used |
|---|---|---|
| Sun creams | Titanium dioxide (TiO₂) | Absorbs UV radiation. Nanoparticles are transparent on skin (unlike bulk TiO₂ which is white and opaque). |
| Catalysts | Various metals (e.g. platinum) | Very high surface area dramatically increases the rate of reaction. |
| Medical drug delivery | Various | Nanoparticles can cross cell membranes and deliver drugs directly to target cells (e.g. cancer cells). Buckminster fullerene can store medicines in its hollow centre. |
| Electronics / sensors | Carbon nanotubes, graphene | Tiny size enables microscale circuitry; high conductivity; sensitive detection of small amounts of substances. |
| Stronger materials | Carbon nanotubes | Incredibly strong and lightweight — reinforcing composites for sports equipment and aerospace. |
| Deodorants | Silver nanoparticles | Silver nanoparticles are antibacterial; they kill bacteria on the skin that cause body odour. The high surface area makes them effective at low concentrations. |
| Other antibacterial products | Silver nanoparticles | Silver is antibacterial; nanoparticles increase the surface area in contact with bacteria — used in wound dressings, socks, and food packaging. |
Risks of Nanoparticles
The use of nanoparticles raises concerns because little is yet known about their long-term effects on human health and the environment. Scientists treat nanoparticles cautiously because:
- Their very small size means they can penetrate cell membranes and potentially damage cells from the inside.
- Some nanoparticles may be toxic to living organisms.
- Very small particles might cross the blood-brain barrier, potentially affecting the nervous system.
- If released into the environment, nanoparticles could act as nanopollutants, with unpredictable effects on ecosystems.
- They can speed up biological reactions in unpredictable ways.
AQA questions on nanoparticles often ask you to evaluate their use. Always acknowledge both the benefits and the risks. A complete answer mentions the high surface area to volume ratio (benefit: faster reactions / more effective catalysts), and acknowledges that their long-term health and environmental effects are not yet fully understood (risk). Avoid writing as though nanoparticles are either entirely safe or entirely dangerous.
🧪 Exam-style questions
What is the approximate number of atoms in a nanoparticle?Tick (✓) one box.
Nanoparticles of some elements can be used as catalysts. Which element is most likely to be used as a catalyst?Use the periodic table. Tick (✓) one box.
Figure 9 shows a cubic nanoparticle.
Figure 9 — a cubic nanoparticle, 4 nm along every edge.
Calculate (a) the surface area of the cubic nanoparticle, (b) its volume, and (c) the simplest whole number ratio of surface area : volume.
Show answer
The official AQA mark scheme awards all 6 marks as a method mark plus a result mark for each part:
- Surface area = 6 × 4² 1 mark = 96 nm² 1 mark (one face = 4² = 16 nm², then × 6)
- Volume = 4³ 1 mark = 64 nm³ 1 mark
- Ratio = 96 : 64 1 mark = 3 : 2 1 mark (divide both sides by 32)
Follow-through marks: the last two marks are still available even if an earlier value was wrong.
- Solids: fixed shape, fixed volume, close-packed regular particles. Liquids: fixed volume, flow, close irregular particles. Gases: no fixed shape/volume, compressible, widely spaced particles.
- Changes of state are physical and reversible. Temperature stays constant during melting and boiling — energy breaks forces, not raises temperature.
- Ionic bonding: metal transfers electrons to non-metal → oppositely charged ions → electrostatic attraction. Ionic compounds form giant ionic lattices with high MP. Conduct only when molten or dissolved (ions free to move).
- Covalent bond = shared pair of electrons. Occurs between non-metals. Single (1 pair), double (2 pairs), triple (3 pairs).
- Simple molecules: low MP/BP (weak intermolecular forces — not weak bonds). Do not conduct electricity. N.B. covalent bonds themselves are strong.
- Giant covalent structures: many strong covalent bonds → very high MP, very hard, insoluble. Examples: diamond, graphite, SiO₂.
- Diamond: 4 bonds per C, 3D network, hardest natural substance, does not conduct. Graphite: 3 bonds per C, layered, soft/slippery, conducts electricity (delocalised e⁻).
- Graphene: one layer of graphite — strong, lightweight, conducts. Nanotubes: cylindrical fullerenes with very high length:diameter ratio — strong, conducts. C₆₀: simple molecule, hollow sphere, drug delivery.
- Metallic bonding: lattice of positive ions + sea of delocalised electrons. Explains conductivity (e⁻ free to move), malleability (layers slide), high MP (strong attraction).
- Alloys: different-sized atoms distort layers → harder than pure metal. (Beyond the spec: shape memory alloys such as Nitinol return to their original shape on heating.)
- Nanoparticles T: diameter 1–100 nm. SA:V ratio increases by factor of 10 for every factor of 10 decrease in particle size. Properties differ from bulk material. Uses: medicine, electronics, cosmetics/sun creams, deodorants, catalysts. Risks: unknown long-term health/environmental effects.