Last updated: 29 September 2026
When sodium reacts with chlorine, electrons are transferred and ions are formed. But many non-metal atoms do not bond in this way. Hydrogen, oxygen, nitrogen, chlorine and carbon can instead join together by sharing electrons. This gives us covalent bonding.
Understanding covalent bonding is more than learning how to draw dots and crosses. The same bonding idea explains why methane is a gas, why water is a liquid, why diamond is extremely hard and why graphite can conduct electricity even though most covalent substances cannot.
The key is to keep two questions separate: how are the atoms bonded? and how are those bonded atoms arranged in the substance? Once that distinction is clear, many structure-and-property questions become much easier.
Try these five questions before reading the chapter. They are designed to reveal whether the basic ideas about electrons, bonding and structure are already secure.
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What is a covalent bond?
A covalent bond is formed when two atoms share a pair of electrons. Covalent bonding occurs mainly between non-metal atoms.
More precisely, the shared pair of negatively charged electrons is attracted to the positively charged nuclei of both bonded atoms. This strong electrostatic attraction holds the atoms together.
The two electrons in the shared pair are counted as part of the outer electron shell of both bonded atoms.
Why do atoms share electrons?
Non-metal atoms often have outer electron shells that are nearly full. Instead of transferring several electrons between one another, they can share electrons.
Consider chlorine. A chlorine atom has the electronic configuration 2,8,7. It has seven outer-shell electrons and needs one more for a stable outer-shell arrangement.
When two chlorine atoms bond, each contributes one electron to a shared pair. Both chlorine atoms can then count the shared pair as part of their outer shells, so each effectively has eight outer-shell electrons.
Useful model: We often say that atoms form bonds to obtain a “full outer shell”.
This is a useful O-Level model for predicting simple structures. It should not be interpreted as atoms consciously “wanting” or “choosing” electrons.
Hydrogen is different
Hydrogen has only one occupied electron shell. Its first shell is complete with two electrons, not eight.
A hydrogen atom therefore needs only one additional electron. This is why hydrogen normally forms one covalent bond.
Predicting the number of covalent bonds
For many common non-metal atoms, you can predict the usual number of covalent bonds by considering how many more electrons the atom needs to complete its outer shell.
| Atom | Outer-shell electrons | Electrons usually needed | Usual number of covalent bonds |
|---|---|---|---|
| Hydrogen | 1 | 1 | 1 |
| Carbon | 4 | 4 | 4 |
| Nitrogen | 5 | 3 | 3 |
| Oxygen | 6 | 2 | 2 |
| Chlorine | 7 | 1 | 1 |
This pattern is especially useful when drawing the common molecules encountered at O-Level.
Drawing dot-and-cross diagrams
A dot-and-cross diagram shows the outer-shell electrons of the atoms involved in bonding. Electrons from one atom may be represented by dots, while electrons from another atom are represented by crosses.
The dot and cross symbols do not represent different kinds of electron. All electrons are identical. The different symbols simply allow us to see which atom each electron originally came from.
A reliable method
- Write down the atoms present. Use the molecular formula to determine how many atoms of each element are present.
- Determine the outer-shell electrons. Use the element's electronic structure or position in the Periodic Table.
- Work out how many more electrons each atom needs. This helps you predict the likely number of covalent bonds.
- Place the shared pairs between bonded atoms. Each shared pair forms one covalent bond.
- Add the remaining outer-shell electrons. Electrons not involved in bonding must still be shown.
- Check every atom. Hydrogen should normally have two electrons around it, while atoms such as carbon, nitrogen, oxygen and chlorine should normally have eight.
A dot-and-cross diagram must show the relevant outer-shell electrons, not only those in the shared pairs. Unshared pairs of electrons are called lone pairs.
Examples of simple covalent molecules
Hydrogen, H₂
Each hydrogen atom has one electron and needs one more electron to complete its first shell.
The two hydrogen atoms each contribute one electron to form one shared pair. Therefore, H₂ contains one single covalent bond.
Chlorine, Cl₂
Each chlorine atom has seven outer-shell electrons and needs one more.
Two chlorine atoms share one pair of electrons. Each chlorine atom can therefore count eight electrons in its outer shell.
Each chlorine atom also has three lone pairs that are not involved in the bond.
Water, H₂O
Oxygen has six outer-shell electrons and needs two more. Each hydrogen atom needs one more.
Oxygen therefore shares one pair of electrons with each of two hydrogen atoms, producing two O–H covalent bonds.
The oxygen atom also has two lone pairs of electrons.
Methane, CH₄
Carbon has four outer-shell electrons and needs four more. Each hydrogen atom needs one more electron.
Carbon shares one pair of electrons with each of four hydrogen atoms. Carbon therefore forms four single covalent bonds.
Ammonia, NH₃
Nitrogen has five outer-shell electrons and needs three additional electrons to complete its outer shell.
Nitrogen shares one pair with each of three hydrogen atoms, forming three N–H covalent bonds. One pair of electrons remains unshared on the nitrogen atom.
Single, double and triple covalent bonds
Sometimes one shared pair is not enough to give both atoms a complete outer-shell arrangement. Atoms can therefore share more than one pair of electrons.
| Bond | Number of shared pairs | Number of shared electrons | Example |
|---|---|---|---|
| Single bond | 1 | 2 | H–H |
| Double bond | 2 | 4 | O=O |
| Triple bond | 3 | 6 | N≡N |
Oxygen, O₂
Each oxygen atom has six outer-shell electrons and needs two more. The two oxygen atoms therefore share two pairs of electrons.
The O₂ molecule therefore contains a double covalent bond. Each oxygen atom also has two lone pairs.
Nitrogen, N₂
Each nitrogen atom has five outer-shell electrons and needs three more. The two nitrogen atoms share three pairs of electrons.
N₂ therefore contains a triple covalent bond. Each nitrogen atom also has one lone pair.
Carbon dioxide, CO₂
Carbon has four outer-shell electrons and needs four more. Each oxygen atom needs two more electrons.
Carbon shares two pairs of electrons with one oxygen atom and another two pairs with the other oxygen atom.
Carbon dioxide therefore contains two C=O double covalent bonds.
In O₂, the two oxygen atoms are connected by two shared pairs. In N₂, the two nitrogen atoms are connected by three shared pairs.
Molecules and covalent substances
A molecule is a group of atoms held together by covalent bonds. H₂, O₂, N₂, H₂O, NH₃, CH₄ and CO₂ are examples of molecules.
However, being covalently bonded does not always mean that a substance consists of separate molecules. Some substances form enormous networks of covalently bonded atoms instead.
This leads to one of the most important distinctions in the topic: simple molecular structures versus giant covalent structures.
Simple molecular and giant covalent structures
The type of structure determines which attractions must be overcome when a substance melts or boils.
| Feature | Simple molecular structure | Giant covalent structure |
|---|---|---|
| Basic arrangement | Separate molecules | Continuous network of atoms |
| Within the structure | Strong covalent bonds within each molecule | Strong covalent bonds throughout the structure |
| During melting | Intermolecular forces between molecules are overcome | Many strong covalent bonds must be broken |
| Melting and boiling points | Usually relatively low | Very high |
| Electrical conductivity | Usually does not conduct | Usually does not conduct; graphite is an important exception |
| Examples | H₂, Cl₂, H₂O, CH₄, CO₂ | Diamond, graphite, silicon dioxide |
Do not confuse strong covalent bonds with intermolecular forces.
A simple molecular substance can have strong covalent bonds inside each molecule and yet still have a low boiling point because the forces between its molecules are much weaker.
Simple molecular substances
Simple molecular substances consist of separate molecules. Strong covalent bonds hold the atoms together inside each molecule.
However, there are weaker forces of attraction between neighbouring molecules. These are called intermolecular forces.
Why do simple molecular substances have low melting and boiling points?
When a simple molecular substance melts or boils, the individual molecules do not normally break apart into atoms.
Instead, the molecules move further apart. This means that the weaker intermolecular forces between the molecules are overcome.
Because relatively little energy is required to overcome these forces, simple molecular substances generally have relatively low melting and boiling points.
Liquid H₂O becomes gaseous H₂O. The molecules separate from one another, but each molecule remains H₂O.
Why do simple molecular substances usually not conduct electricity?
Electrical conduction requires charged particles that are free to move.
Simple molecular substances normally consist of neutral molecules and do not contain mobile ions or freely moving delocalised electrons.
They therefore usually do not conduct electricity.
Giant covalent structures
In a giant covalent structure, a very large number of atoms are connected together by strong covalent bonds to form a continuous network.
There are no separate small molecules that can simply move away from one another. Melting such a substance requires many strong covalent bonds within the network to be broken.
This requires a large amount of energy, so giant covalent substances have very high melting points.
Diamond
Diamond is a giant covalent form of carbon. Each carbon atom forms four covalent bonds with four other carbon atoms.
The atoms are arranged tetrahedrally around each carbon atom and the bonding extends throughout a giant three-dimensional structure.
Why is diamond extremely hard?
Strong covalent bonds extend throughout the three-dimensional structure. The carbon atoms are held in fixed positions and there are no layers that can easily slide over one another.
Diamond is therefore extremely hard and can be used in cutting and drilling equipment.
Why does diamond have a very high melting point?
Many strong covalent bonds must be broken before the giant structure can melt. A large amount of energy is therefore required.
Why does diamond not conduct electricity?
Each carbon atom has four outer-shell electrons and uses all four of them to form four covalent bonds.
There are no mobile ions and no delocalised electrons available to carry electrical charge through the structure.
Graphite
Graphite is also made entirely of carbon, but its atoms are arranged differently. Each carbon atom forms three covalent bonds with three neighbouring carbon atoms in the same layer.
The atoms form flat layers containing hexagonal rings. Strong covalent bonds hold the carbon atoms together within each layer, while much weaker forces act between neighbouring layers.
Why is graphite soft and slippery?
The forces between neighbouring layers are weak compared with the strong covalent bonds within each layer.
The layers can therefore slide over one another relatively easily. This makes graphite soft and slippery and allows it to be used as a lubricant.
Why does graphite conduct electricity?
Each carbon atom has four outer-shell electrons but uses only three of them to form covalent bonds.
The remaining electron becomes delocalised. These delocalised electrons can move along the layers and carry electrical charge.
Graphite can therefore conduct electricity even though most covalent substances cannot.
Diamond: each carbon forms four covalent bonds in a rigid three-dimensional structure. There are no mobile electrons.
Graphite: each carbon forms three covalent bonds in layers. One electron per carbon atom is delocalised and can move along the layers.
Diamond and graphite: same element, different properties
| Property | Diamond | Graphite |
|---|---|---|
| Atoms present | Carbon only | Carbon only |
| Bonds per carbon atom | 4 | 3 |
| Structure | Rigid three-dimensional network | Flat hexagonal layers |
| Hardness | Very hard | Soft and slippery |
| Electrical conductivity | Does not conduct | Conducts |
| Reason for electrical behaviour | No mobile charged particles | Contains mobile delocalised electrons |
Diamond and graphite show an important principle in Chemistry: knowing which element is present is not enough to predict properties. The arrangement and bonding of the atoms matter as well.
Silicon dioxide, SiO₂
Silicon dioxide is another important giant covalent substance. It is found in substances such as quartz and is a major component of many types of sand.
In silicon dioxide, each silicon atom is covalently bonded to four oxygen atoms, while each oxygen atom is bonded to two silicon atoms. These bonds continue throughout a giant three-dimensional network.
In giant covalent silicon dioxide, the formula SiO₂ gives the simplest ratio of silicon atoms to oxygen atoms in the network: one silicon atom for every two oxygen atoms.
Why does silicon dioxide have a high melting point?
Many strong Si–O covalent bonds extend throughout its giant structure.
A large amount of energy is needed to break enough of these bonds for the substance to melt.
Why does silicon dioxide not conduct electricity?
Silicon dioxide has no mobile ions and no freely moving delocalised electrons.
It therefore does not conduct electricity.
How to explain structure and properties
Examination questions often ask you to explain a property rather than simply state it. A useful approach is to build your answer in a logical chain.
First identify the type of structure. Then identify the bonds, forces or mobile particles that matter. Finally, connect these ideas to the observed property.
Example: why does methane have a low boiling point?
Methane has a simple molecular structure. There are weak intermolecular forces between CH₄ molecules.
Only a small amount of energy is needed to overcome these forces, so methane has a low boiling point.
Example: why does diamond have a high melting point?
Diamond has a giant covalent structure. Many strong covalent bonds extend throughout the structure.
A large amount of energy is required to break these bonds, so diamond has a very high melting point.
Example: why does graphite conduct electricity?
Each carbon atom in graphite forms three covalent bonds, leaving one delocalised electron per carbon atom.
These electrons can move along the layers and carry electrical charge, so graphite conducts electricity.
Common mistakes with covalent bonding
| Common mistake | Why it is wrong | Correct approach |
|---|---|---|
| Saying electrons are transferred in a covalent bond | Electron transfer describes ionic bonding | State that a pair of electrons is shared |
| Drawing only the bonding electrons in a dot-and-cross diagram | Outer-shell electrons not involved in bonding must also be represented | Include the appropriate lone pairs |
| Saying one shared electron forms one covalent bond | A covalent bond contains a shared pair of electrons | Two shared electrons form one covalent bond |
| Saying O₂ contains only one shared pair | The oxygen atoms share two pairs of electrons | Describe O₂ as containing a double bond |
| Saying covalent bonds break when water boils | Water molecules remain intact during boiling | State that intermolecular forces between H₂O molecules are overcome |
| Saying all covalent substances have low melting points | Giant covalent structures contain many strong covalent bonds | Distinguish simple molecular from giant covalent structures |
| Saying graphite contains weak covalent bonds | The covalent bonds within each graphite layer are strong | State that the forces between the layers are weak |
| Saying graphite conducts because carbon is a conductor | Diamond also contains carbon but does not conduct | Explain conductivity using mobile delocalised electrons |
| Calling diamond or silicon dioxide molecules | They are continuous giant covalent networks | Describe them as giant covalent structures |
| Saying SiO₂ represents one silicon dioxide molecule | There are no separate SiO₂ molecules in the giant structure | Treat SiO₂ as the simplest silicon-to-oxygen ratio |
Covalent bonding checklist
Recognise covalent bonding
Covalent bonding usually occurs between non-metal atoms and involves sharing pairs of electrons.
Count outer-shell electrons
Use electronic structure to determine how many electrons each atom normally needs for a stable outer shell.
Count shared pairs
One shared pair gives a single bond, two shared pairs a double bond and three shared pairs a triple bond.
Show lone pairs
When drawing dot-and-cross diagrams, include the outer-shell electrons that are not used in bonding.
Identify the structure
Decide whether the substance consists of separate molecules or a continuous giant covalent network.
Explain melting and boiling
Simple molecular substances involve overcoming intermolecular forces, whereas giant covalent substances require strong covalent bonds to be broken.
Check for mobile charge
Electrical conduction requires ions or electrons that are free to move.
Remember graphite
Graphite has layers that can slide and delocalised electrons that allow it to conduct electricity.
These ten questions go beyond simple recall. Use electron arrangements, bonding and structure to reason through each situation.
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Questions students often ask about covalent bonding
A few ideas in covalent bonding can sound contradictory at first. These questions help connect the different parts of the chapter.
Do shared electrons belong to one atom or both atoms?
The shared electrons are attracted to both nuclei and are counted as part of the outer electron shell of both bonded atoms.
Are dots and crosses different types of electrons?
No. All electrons are identical. Dots and crosses are simply different symbols used to show which atom an electron came from before bonding.
Why can two atoms share more than one pair of electrons?
A single shared pair may not provide enough electrons for both atoms to achieve their stable outer-shell arrangements. Oxygen atoms share two pairs in O₂, while nitrogen atoms share three pairs in N₂.
If covalent bonds are strong, why do molecular substances often boil easily?
The covalent bonds inside each molecule are strong. During boiling these bonds remain intact. Instead, the weaker intermolecular forces between separate molecules are overcome.
Are all covalent substances insoluble in water?
No. Solubility depends on the particular substance and how strongly it interacts with water. It is therefore unsafe to memorise that every covalent substance is insoluble.
Why does graphite conduct electricity but diamond does not?
Each carbon atom in graphite forms three covalent bonds, leaving one delocalised electron that can move along the layers. In diamond, each carbon atom forms four covalent bonds, so there are no comparable mobile delocalised electrons.
Does graphite have weak covalent bonds because it is soft?
No. The covalent bonds within each graphite layer are strong. Graphite is soft because the forces between neighbouring layers are much weaker, so the layers can slide over one another.
Why is diamond hard while graphite is soft if both contain only carbon?
Their carbon atoms are arranged differently. Diamond contains a rigid three-dimensional network of covalent bonds. Graphite consists of flat layers with weaker forces between the layers. Different structures therefore produce different properties.
Is silicon dioxide made of SiO₂ molecules?
No. Solid silicon dioxide has a giant covalent network. The formula SiO₂ represents the simplest ratio of silicon atoms to oxygen atoms in that network rather than one discrete molecule.
How can I tell whether to discuss covalent bonds or intermolecular forces in an exam?
First identify the structure. For a simple molecular substance, melting and boiling usually involve overcoming intermolecular forces between molecules. For a giant covalent substance, melting requires many strong covalent bonds in the network to be broken.
Covalent bonding becomes much easier once you connect electron sharing to structure. Start by asking how many electrons each atom needs and how many pairs must be shared. Then ask whether those bonded atoms form separate molecules or one giant network. That second question explains the striking differences between substances such as methane, diamond, graphite and silicon dioxide — and turns property questions from memorisation into logical Chemistry.