Covalent 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.
🧪 Exam-style questions
A chlorine molecule, Cl2, contains a covalent bond. Describe how this covalent bond is formed.
Show answer
- The two chlorine atoms share a pair of electrons 1 mark
- so that each atom has a full outer shell (the shared pair counts for both atoms) 1 mark
Allow: “one electron from each atom is shared”; reference to a single bond / one shared pair.
A nitrogen molecule, N2, is held together by a triple bond. How many pairs of electrons are shared in this bond? Tick (✓) one box.
The displayed formula of ethane is shown below, where each line is a covalent bond.
Deduce the molecular formula of ethane from this structure.
Show answer
- Count the atoms: 2 carbon and 6 hydrogen, so the molecular formula is C2H6 1 mark
Simple 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.
🧪 Exam-style questions
Oxygen, O2, is a simple molecular substance. Explain why oxygen has a very low boiling point.
Show answer
- Oxygen consists of simple molecules with weak intermolecular forces between them 1 mark
- Only a small amount of energy is needed to overcome / break these forces 1 mark
- It is the intermolecular forces that are overcome on boiling, not the (strong) covalent bonds inside the molecules 1 mark
Do not accept: “the covalent bonds are weak” — the covalent bonds are strong; it is the forces between molecules that are weak.
As the molecules in a series get larger, their boiling points increase. Explain why.
Show answer
- Larger molecules have stronger intermolecular forces 1 mark
- so more energy is needed to overcome them, giving a higher boiling point 1 mark
Why do simple molecular substances not conduct electricity? Tick (✓) one box.
Polymers
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.
🧪 Exam-style questions
Which feature in a diagram tells you that a substance is a polymer? Tick (✓) one box.
Poly(ethene) is a polymer. Explain why poly(ethene) is solid at room temperature.
Show answer
- The long polymer chains are held together by relatively strong intermolecular forces 1 mark
- so a lot of energy is needed to overcome them — more than for small molecules — so it is solid at room temperature 1 mark
Do not accept: answers that say the covalent bonds inside the chains are overcome on melting — only the intermolecular forces between chains are overcome.
Giant 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.
🧪 Exam-style questions
Silicon dioxide is a giant covalent structure. Explain why it has a very high melting point.
Show answer
- It has a giant covalent structure with many strong covalent bonds 1 mark
- A large amount of energy is needed 1 mark
- to break these (many strong covalent) bonds 1 mark
Do not accept: overcoming intermolecular forces — a giant covalent structure has no molecules.
Carbon dioxide (CO2) is a gas at room temperature, but silicon dioxide (SiO2) is a solid with a very high melting point. Explain this difference in terms of their structures.
Show answer
- CO2 is a simple molecule with weak intermolecular forces, so little energy is needed to separate the molecules (gas at room temperature) 1 mark
- SiO2 is a giant covalent structure, so many strong covalent bonds must be broken to melt it (very high melting point) 1 mark
Carbon 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 and the fullerenes — a family of hollow carbon structures that includes buckminsterfullerene (C60) and carbon nanotubes (which are cylindrical fullerenes).
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.
Buckminsterfullerene (C60)
Buckminsterfullerene was the first fullerene to be discovered. It 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.
Buckminsterfullerene (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.
🧪 Exam-style questions
Graphite conducts electricity and is soft, but diamond does not conduct electricity and is hard. Both are forms of carbon. Explain these differences in terms of their structure and bonding.
Show answer
- In graphite, each carbon forms 3 covalent bonds, leaving one delocalised electron per atom that is free to move and carry charge — so graphite conducts 1 mark
- In diamond, each carbon forms 4 covalent bonds, so there are no delocalised electrons — so diamond does not conduct 1 mark
- Graphite is arranged in layers held together by weak intermolecular forces, so the layers can slide over each other — making it soft / slippery 1 mark
- Diamond is a rigid 3D giant covalent network with strong bonds in all directions, so it is hard 1 mark
Allow: “like a metal, graphite has delocalised electrons” for the conduction mark.
Describe the structure of graphene, and give one property that makes it useful.
Show answer
- Graphene is a single layer of graphite — one atom thick — made of carbon atoms in a sheet of hexagonal rings 1 mark
- Any one useful property: it is very strong / very light / conducts electricity (delocalised electrons) 1 mark
Allow: any one named property linked to a use (e.g. electronics, composites).
Buckminsterfullerene was the first fullerene to be discovered. What is its chemical formula? Tick (✓) one box.
Which statement about carbon nanotubes is correct? Tick (✓) one box.
Carbon nanotubes are cylindrical fullerenes. Explain the properties of carbon nanotubes. Answer in terms of structure and bonding. This is a levels-of-response question — you are marked on how well your ideas are organised and linked, not just the number of points.
Show a model answer
How it is marked (levels of response):
- Level 3 (5–6): a detailed and coherent explanation applying knowledge of the properties of nanotubes, with clear and logical links to the reasons why carbon nanotubes have these properties.
- Level 2 (3–4): relevant statements showing clear knowledge of the properties of nanotubes, with an attempt to link properties to why they occur — but the logic may be unclear.
- Level 1 (1–2): simple relevant statements of the properties of nanotubes, with no linking to an explanation of why they occur.
Indicative content — properties:
- high tensile strength
- high electrical / thermal conductivity
- high melting point
Explanations:
- nanotubes are fullerenes based on hexagonal rings of carbon atoms
- which means each carbon forms three covalent bonds with three other carbon atoms
- covalent bonds are strong (need a lot of energy to break them)
- so nanotubes are strong / have a high tensile strength, and have a high melting point
- the structure means one electron from each carbon atom is delocalised
- as in metals and graphite, the delocalised electrons can move throughout the structure
- allowing the carbon nanotube / fullerene to conduct thermal energy and electricity
Source: AQA GCSE Chemistry.