Last updated: 4 October 2026
Diamond, graphite, sodium chloride and copper are all solids, but they behave in completely different ways. Diamond is extremely hard, graphite is soft enough to leave a mark on paper, sodium chloride crystals shatter when struck and copper can be drawn into long electrical wires. The explanation lies in a world that we cannot see directly: the arrangement and bonding of their particles.
This topic brings together ionic, covalent and metallic bonding and asks the question that matters most in materials Chemistry: how does microscopic structure produce macroscopic properties? Once you learn to move from particles → bonding → movement → property, many facts that once seemed unrelated become part of one logical story.
Try these five questions before reading the chapter. They test the bonding ideas that you will use to explain the properties of materials.
- 5 questions
- Instant feedback
- Explanation for every option
Structure determines the properties of a material
When an examination question asks why a substance has a particular property, simply naming its structure is not enough. You need to connect the structure to the property.
Follow the structure–property reasoning chain
A reliable explanation usually follows this sequence:
- Identify the particles — atoms, molecules, ions or delocalised electrons.
- Describe their arrangement — separate molecules, a giant lattice, layers or a continuous network.
- State the bonding or attractive forces acting between the particles.
- Ask what can move — ions, electrons or whole layers.
- Connect this directly to the observed property.
Suppose you are asked why sodium chloride has a high melting point. Writing only “because it has a giant ionic lattice” identifies the structure, but does not explain the property. The important next step is to state that there are strong electrostatic forces of attraction between oppositely charged ions throughout the lattice. A large amount of energy is needed to overcome these attractions.
Elements, compounds and mixtures
Before comparing different material structures, it helps to distinguish three basic ways in which matter can be classified.
| Type | What it contains | Composition | How it can be separated | Example |
|---|---|---|---|---|
| Element | Only one type of atom | One chemical element | Cannot be broken down into simpler substances by chemical methods | Copper, oxygen, carbon |
| Compound | Atoms of two or more elements chemically combined | Fixed ratio | Can be broken down by chemical methods | H₂O, CO₂, NaCl |
| Mixture | Two or more substances not chemically combined | Can vary | Components can be separated by physical methods | Air, seawater, brass |
Compound or mixture?
In a compound, the constituent elements have reacted and formed chemical bonds. The compound therefore has its own properties and a fixed composition.
In a mixture, the substances remain chemically distinct. Their proportions may vary, and suitable physical methods can be used to separate them.
Giant ionic structures
Ionic compounds are formed from positive and negative ions. In many familiar examples, electrons are transferred from metal atoms to non-metal atoms. The metal atoms lose electrons and become positive ions, while the non-metal atoms gain electrons and become negative ions.
For sodium chloride:
Formation of Na⁺ and Cl⁻ ions
Sodium: Na → Na⁺ + e⁻
Chlorine: Cl + e⁻ → Cl⁻
Once the ions have formed, Na⁺ and Cl⁻ attract one another because they have opposite charges. This attraction is the ionic bond.
What is ionic bonding?
Ionic bonding is the strong electrostatic force of attraction between oppositely charged ions.
Notice the distinction: electron transfer forms the ions, while the electrostatic attraction between the ions is the ionic bond.
An ionic crystal is a giant lattice
A sodium chloride crystal is not made from separate NaCl molecules or isolated Na⁺–Cl⁻ pairs. Instead, enormous numbers of positive and negative ions are arranged in a regular, three-dimensional repeating structure called a giant ionic lattice.
Each ion is surrounded by ions of opposite charge. As a result, strong electrostatic attractions act in many directions throughout the crystal.
Why do ionic compounds have high melting and boiling points?
Melting an ionic solid requires the ions to become able to move past one another. To make this happen, enough of the strong electrostatic attractions holding the giant lattice together must be overcome.
Because these attractions are strong and extend throughout the lattice, a large amount of thermal energy is required. Ionic compounds therefore generally have high melting and boiling points.
Build the full examination explanation
A weak answer is: “Sodium chloride has a high melting point because its ionic bonds are strong.”
A stronger answer explains the particles and forces: there are strong electrostatic forces of attraction between oppositely charged ions throughout the giant ionic lattice, so a large amount of energy is required to overcome these attractions.
Why does solid sodium chloride not conduct electricity?
Electrical conduction requires mobile charged particles.
Solid sodium chloride certainly contains charged particles: Na⁺ and Cl⁻ ions. However, the ions are held in fixed positions in the lattice. They can vibrate about their positions, but they cannot travel through the solid carrying charge.
Solid sodium chloride therefore does not conduct electricity.
Why do molten ionic compounds conduct electricity?
When an ionic compound melts, the ions are no longer held in fixed lattice positions. The positive and negative ions can move through the liquid.
Because these mobile ions carry electrical charge, the molten ionic compound conducts electricity.
Why do aqueous ionic compounds conduct electricity?
When a soluble ionic compound dissolves in water, its ions separate and become free to move through the solution.
An aqueous solution containing mobile ions can therefore conduct electricity.
Ions must be mobile
- Solid ionic compound: ions present but fixed → does not conduct.
- Molten ionic compound: ions mobile → conducts.
- Aqueous ionic compound: dissolved ions mobile → conducts.
Why are ionic crystals brittle?
Ionic solids are hard, but they are also brittle. This means that they tend to crack rather than bend when a sufficiently large force is applied.
Imagine one layer of ions being pushed sideways relative to another. The shift may bring positive ions next to other positive ions and negative ions next to other negative ions.
Like charges repel. The resulting repulsion can split the lattice, causing the crystal to fracture.
Simple molecular substances
Covalent substances do not all have the same type of structure. Some consist of separate molecules, while others form giant covalent networks. This distinction is essential because it determines which forces must be overcome when the substance melts or boils.
Simple molecular substances include H₂, O₂, H₂O, NH₃, CH₄, CO₂ and I₂.
There are strong bonds inside a molecule but weaker forces between molecules
The atoms within each molecule are held together by strong covalent bonds. However, the attractions between one molecule and another are much weaker. These attractions are called intermolecular forces.
Do not confuse covalent bonds with intermolecular forces
Inside a molecule: strong covalent bonds hold atoms together.
Between molecules: weaker intermolecular forces attract neighbouring molecules.
During ordinary melting and boiling, it is the intermolecular forces that are overcome. The molecules themselves remain intact.
Why do simple molecular substances have low melting and boiling points?
When a simple molecular substance melts or boils, the molecules move further apart from one another. The atoms inside each molecule do not need to separate.
Therefore, the strong covalent bonds within the molecules are not broken. Only the relatively weak intermolecular forces between molecules need to be overcome.
Because comparatively little energy is required, simple molecular substances generally have low melting and boiling points.
Why do simple molecular substances usually not conduct electricity?
Most simple molecular substances consist of neutral molecules and contain neither mobile ions nor delocalised electrons.
With no mobile charged particles available to carry current, they generally do not conduct electricity.
Giant covalent structures
A giant covalent substance is very different from a simple molecular substance. Instead of containing separate small molecules, it contains an enormous continuous network of atoms joined by covalent bonds.
Important examples include diamond, graphite and silicon dioxide, SiO₂.
Why do giant covalent substances have very high melting points?
Strong covalent bonds extend throughout the giant structure. To melt the substance, many of these strong bonds must be broken as the network is disrupted.
A very large amount of energy is therefore required, giving giant covalent substances very high melting points.
Diamond and graphite: structure changes everything
Diamond and graphite are both made entirely of carbon. They are different structural forms of the same element, known as allotropes.
Because the carbon atoms are arranged differently in the two structures, diamond and graphite have dramatically different properties.
Structure and properties of diamond
In diamond, each carbon atom forms four strong covalent bonds with four other carbon atoms. The bonds extend in different directions to form a rigid three-dimensional giant covalent network.
Diamond is extremely hard because:
- each carbon atom is strongly bonded to four others;
- strong covalent bonds extend throughout the structure; and
- there are no layers that can easily slide over one another.
This explains why diamond is useful for cutting and drilling hard materials.
Why does diamond not conduct electricity?
Carbon has four outer-shell electrons. In diamond, all four are used in forming covalent bonds.
There are therefore no mobile ions or delocalised electrons available to carry electrical charge. Diamond does not conduct electricity.
Structure and properties of graphite
In graphite, each carbon atom forms three covalent bonds with three other carbon atoms. The atoms form flat layers containing hexagonal arrangements of carbon atoms.
The covalent bonds within each layer are strong, but the forces between neighbouring layers are much weaker.
Why is graphite soft and slippery?
Because only weak forces act between neighbouring layers, the layers can slide over one another relatively easily.
This makes graphite soft and slippery, explaining its use in pencil cores and lubricants.
Why does graphite conduct electricity?
Each carbon atom uses three of its four outer-shell electrons in covalent bonding. The remaining electron becomes delocalised.
These delocalised electrons can move through the graphite layers and carry electrical charge. Graphite therefore conducts electricity.
| Feature | Diamond | Graphite |
|---|---|---|
| Bonds formed by each carbon atom | Four covalent bonds | Three covalent bonds |
| Arrangement | Rigid three-dimensional network | Layers of carbon atoms |
| Hardness | Extremely hard | Soft and slippery |
| Electrical conductivity | Does not conduct | Conducts |
| Reason for conductivity | No mobile charged particles | Contains mobile delocalised electrons |
Do not say that graphite has weak covalent bonds
The covalent bonds within each graphite layer are strong. Graphite is soft because the forces between neighbouring layers are weak, allowing the layers to slide.
Structure and properties of silicon dioxide
Silicon dioxide, SiO₂, is another giant covalent substance. Although its formula is written as SiO₂, a solid sample does not consist of separate SiO₂ molecules.
Instead, silicon and oxygen atoms are joined by covalent bonds throughout a giant network. Each silicon atom is bonded to four oxygen atoms, while each oxygen atom connects two silicon atoms.
Because many strong covalent bonds extend throughout the structure, silicon dioxide has a high melting point. It is also hard and does not conduct electricity because it lacks mobile charged particles.
Metallic structure and metallic bonding
Metals have a giant structure containing positive metal ions surrounded by delocalised electrons. These electrons are not associated with one particular atom or ion; they can move throughout the metallic structure.
What is metallic bonding?
Metallic bonding is the strong electrostatic force of attraction between positive metal ions and delocalised electrons.
The attraction between the positive ions and the delocalised electrons acts throughout the metal and holds its giant structure together.
Why do metals conduct electricity?
The delocalised electrons in a metal are mobile. When a potential difference is applied, they can move through the metal and carry electrical charge.
This is why a metal can conduct electricity even while it is solid. The positive metal ions do not need to move through the structure; the electrons are the charge carriers.
Why do metals conduct heat well?
Mobile delocalised electrons can transfer energy rapidly from hotter regions of a metal to cooler regions. Vibrations of the closely packed particles also transfer energy through the structure.
This helps explain why metals such as copper and aluminium are widely used where efficient heat transfer is required.
Why do metals generally have high melting points?
There are strong electrostatic attractions between the positive metal ions and the delocalised electrons.
A considerable amount of energy is required to overcome enough of these attractions for the metal to melt. Metals therefore generally have relatively high melting points, although the strength of metallic bonding differs from one metal to another.
Why are metals malleable and ductile?
A material is malleable if it can be hammered or pressed into different shapes. It is ductile if it can be drawn into wires.
When a force is applied to a metal, layers of positive metal ions can move past one another. The delocalised electrons continue to attract the ions even after their positions have changed.
Metallic bonding is therefore maintained as the layers slide, so the metal can change shape instead of immediately cracking.
Why does an ionic crystal crack while a metal bends?
Ionic crystal: shifting layers may bring ions of the same charge next to one another. Their repulsion causes the lattice to fracture.
Metal: layers of positive ions can slide while remaining attracted to the surrounding delocalised electrons, so the material can change shape.
Why are alloys harder than pure metals?
In a pure metal, atoms of the same element have similar sizes and form relatively regular layers. When a force is applied, these layers can slide over one another.
An alloy is a mixture containing a metal and one or more other elements. Examples include:
- Brass: mainly copper and zinc.
- Bronze: mainly copper and tin.
- Steel: mainly iron with carbon and, in some steels, other elements.
The added atoms are often different in size from the atoms of the main metal. This disturbs the regular arrangement of layers and makes it more difficult for one layer to slide over another.
Alloys are therefore generally harder than the corresponding pure metals.
How to explain why an alloy is harder
Alloys contain atoms of different sizes. These atoms disrupt the regular arrangement of atoms in the pure metal and make it more difficult for the layers to slide over one another. Therefore, the alloy is harder.
Comparing the main structures
The easiest way to see the whole topic is to compare the particles, bonding and properties side by side.
| Structure | Particles and arrangement | Melting point | Electrical conductivity | Other important properties |
|---|---|---|---|---|
| Giant ionic | Positive and negative ions in a regular giant lattice | High because strong electrostatic attractions must be overcome | Does not conduct when solid; conducts when molten or aqueous because ions become mobile | Hard and brittle |
| Simple molecular | Separate molecules with intermolecular forces between them | Usually low because relatively weak intermolecular forces are overcome | Usually does not conduct because there are no mobile charged particles | Often gases, liquids or low-melting solids |
| Giant covalent | Atoms joined by covalent bonds throughout a continuous network | Very high because many strong covalent bonds must be broken | Usually does not conduct; graphite is an important exception | Usually hard; graphite is soft because of its layered structure |
| Metallic | Positive metal ions surrounded by delocalised electrons | Generally high because metallic bonding is strong | Conducts in the solid and molten states because delocalised electrons are mobile | Malleable and ductile |
How to identify a structure from its properties
Examination questions sometimes give you a set of properties for an unknown substance and ask you to deduce its structure. The key is to use several pieces of evidence together.
High melting point
Suggests strong bonding or attractions in a giant structure. It could be ionic, metallic or giant covalent, so more evidence is needed.
Low melting or boiling point
Usually suggests a simple molecular substance with relatively weak intermolecular forces.
Conducts only when molten
Strong evidence for an ionic substance because the ions become mobile after melting.
Conducts as a solid
Suggests mobile electrons, as found in metals and graphite.
Hard but brittle
Often characteristic of an ionic crystal because displacement can bring like-charged ions together and cause the lattice to fracture.
Malleable and ductile
Strongly suggests metallic bonding because layers can slide while metallic attraction is maintained.
Worked reasoning: identify the likely structure
| Observation | Likely structure | Reasoning |
|---|---|---|
| High melting point; does not conduct when solid; conducts when molten | Giant ionic | Strong attractions give a high melting point. Ions are fixed when solid but mobile when molten. |
| Low boiling point; does not conduct electricity | Simple molecular | Weak intermolecular forces require little energy to overcome, and there are no mobile charged particles. |
| Very high melting point; extremely hard; does not conduct | Giant covalent | Strong covalent bonds extend throughout the structure, and there are no mobile charged particles. |
| Conducts as a solid and can be hammered into shape | Metallic | Mobile delocalised electrons carry charge, while layers of metal ions can slide without destroying the metallic bonding. |
Do not identify a structure from one property alone
A high melting point does not automatically mean that a substance is ionic. Metals and giant covalent substances can also have high melting points.
Combine evidence from melting point, conductivity in different physical states, hardness, brittleness and malleability.
Common mistakes and why they are wrong
“Ionic bonding is the transfer of electrons.”
Electron transfer forms the ions. Ionic bonding is the strong electrostatic attraction between oppositely charged ions.
“Solid ionic compounds have no charged particles.”
They contain charged ions, but the ions are fixed in the lattice and cannot move through the solid to carry charge.
“Covalent bonds break when water boils.”
Boiling water overcomes intermolecular forces between H₂O molecules. The covalent O–H bonds within each molecule remain intact.
“All covalent substances have low melting points.”
Simple molecular substances generally have low melting points, but giant covalent substances have very high melting points.
“Graphite is soft because its covalent bonds are weak.”
The covalent bonds within each layer are strong. Graphite is soft because weak forces between the layers allow them to slide.
“Graphite conducts because carbon ions move.”
The carbon atoms remain in the structure. Mobile delocalised electrons are responsible for electrical conduction.
“Metals are malleable because metallic bonds are weak.”
Metallic bonding is strong. Layers can move while the positive ions remain attracted to delocalised electrons.
“An alloy is a compound.”
An alloy is a mixture. Its composition can vary and it does not have one fixed chemical formula.
Structure and properties checklist
Before finishing this chapter, make sure you can move from a structure to its properties and explain every connection using particles and bonding.
Giant ionic structures
I can explain high melting points, solid-state non-conductivity, molten or aqueous conductivity and brittleness using ions and electrostatic attraction.
Simple molecular substances
I can distinguish strong covalent bonds within molecules from weaker intermolecular forces between molecules.
Giant covalent structures
I can explain why giant covalent substances generally have very high melting points.
Diamond and graphite
I can link their different arrangements to hardness, softness and electrical conductivity.
Metallic structures
I can explain metallic bonding, conductivity, malleability and ductility using positive ions and delocalised electrons.
Alloys
I can explain why differently sized atoms make alloys harder than the corresponding pure metals.
Unknown substances
I can combine melting point, conductivity and mechanical properties to deduce a likely structure.
Examination explanations
I can identify the particles, state the forces or bonds, describe particle mobility and connect these ideas directly to the property being explained.
These ten questions go beyond simple recall. Use melting point, conductivity, bonding and particle movement to work out what is happening inside each material.
- 10 application questions
- Instant marking
- Explanation for every option
Questions students often ask about structures and properties
The main ideas become easier to use once a few common apparent exceptions are understood.
Are all ionic compounds soluble in water?
No. Some ionic compounds are soluble in water while others are insoluble or only slightly soluble. If an ionic compound does dissolve, its ions become mobile in the solution and can carry electrical charge.
Use the O-Level solubility rules when deciding whether a particular ionic compound is soluble.
Why does an ionic compound need to be molten before its ions can conduct?
In a solid ionic lattice, the ions are held in fixed positions. They may vibrate, but they cannot travel through the crystal.
Melting disrupts the fixed lattice arrangement. The ions then become mobile and can carry electrical charge through the liquid.
Are all covalent substances gases or liquids?
No. Many simple molecular substances are gases or liquids because the intermolecular forces between their molecules are relatively weak. Others, such as iodine, are solids at room temperature.
Giant covalent substances such as diamond and silicon dioxide are also solids, but their structures are completely different from simple molecular solids.
Why does melting a molecular substance not mean breaking its covalent bonds?
The covalent bonds hold atoms together inside each molecule. During melting, the molecules remain as molecules but gain enough energy to move past one another.
The attractions being overcome are the intermolecular forces between molecules, not the covalent bonds within them.
Why is graphite soft but still has a very high melting point?
These properties depend on different parts of the structure.
Graphite is soft because weak forces between neighbouring layers allow the layers to slide. However, each layer contains many strong covalent bonds. Melting graphite requires extensive disruption of these strong bonds, so a very large amount of energy is required.
Why does graphite conduct electricity but diamond does not?
In graphite, each carbon atom forms three covalent bonds, leaving one electron per carbon atom available to become delocalised. These electrons can move through the structure and carry charge.
In diamond, each carbon atom forms four covalent bonds. All four outer-shell electrons are involved in bonding, so there are no mobile delocalised electrons.
Does the formula SiO₂ mean that silicon dioxide contains separate SiO₂ molecules?
No. In solid silicon dioxide, the atoms form a continuous giant covalent network rather than separate SiO₂ molecules.
The formula SiO₂ tells us the overall ratio of silicon atoms to oxygen atoms in the structure: one silicon atom for every two oxygen atoms.
Why can metals conduct electricity while still solid?
The positive metal ions remain in their positions in the solid, but the delocalised electrons are already mobile. These electrons can move through the metal and carry electrical charge without the metal having to melt.
Is an alloy still metallic even though it is a mixture?
Yes. An alloy is a mixture rather than a compound, but it still has a metallic structure and metallic bonding.
Its properties differ from those of the pure metal because the added atoms disturb the regular arrangement of particles.
Can I identify a structure from its melting point alone?
Usually not. A high melting point may occur in ionic, metallic and giant covalent structures.
Use additional information such as electrical conductivity in the solid and molten states, brittleness, hardness, malleability and ductility before drawing a conclusion.
The most useful idea in this chapter is simple: properties are consequences of structure. Ionic compounds behave as they do because of oppositely charged ions in a giant lattice. Simple molecular substances behave differently because their molecules are separated by much weaker intermolecular forces. Giant covalent substances contain strong covalent bonding throughout a network, while metals combine positive ions with mobile delocalised electrons. Instead of memorising isolated statements, picture the particles and ask: what holds them together, what can move and what must be overcome? From there, the property usually follows.