Last updated: 16 August 2026
Knowing how two atoms bond is only the beginning. In a solid, enormous numbers of particles are arranged together, and the way these particles are organised determines how the substance behaves on a macroscopic scale.
This is why sodium chloride is brittle, copper can be hammered into shape, diamond is extremely hard, graphite conducts electricity and ice floats on liquid water. Each property can be traced back to the particles present, their arrangement and the attractions acting between them.
Rather than memorising a separate list of properties for every substance, it is more useful to develop one reasoning method: identify the particles, identify what holds them together, and then ask what must move or what attractions must be overcome.
Try these five questions before reading the chapter. They test the distinctions between particles, bonding, lattices and electrical conductivity that you will use throughout this topic.
- Check your understanding before studying the chapter.
- Spot common misconceptions about lattices and bonding.
- Get an explanation immediately after each answer.
Use your result as a starting point. The chapter below develops the reasoning behind each idea rather than asking you simply to memorise properties.
What is a lattice structure?
A lattice is a regular, repeating arrangement of particles in a crystalline solid. The term describes how particles are arranged in the solid state; it does not, by itself, identify the type of bonding present.
The particles occupying the lattice may be ions, atoms, positive metal ions with delocalised electrons, or separate molecules. For example, solid sodium chloride has an ionic lattice, copper has a metallic lattice, diamond has a giant covalent network, and solid iodine contains I₂ molecules arranged in a molecular crystal lattice.
Once a substance melts, its particles are no longer fixed in the regular, repeating positions characteristic of the crystalline solid. We therefore describe the lattice of solid iodine, for example, rather than saying that liquid iodine has the same lattice.
A lattice is an ordered arrangement of particles in a crystalline solid. The word giant means that the bonded or electrostatically connected structure extends continuously through a very large number of particles. A molecular solid can form a regular lattice while still containing separate, discrete molecules.
The two questions that unlock most solid-structure problems
When describing or comparing solid structures, begin with two questions.
- What particles or components are present?
- What bonds or forces hold them together?
Once these questions have been answered, properties such as melting point, conductivity, hardness and brittleness can usually be explained logically.
| Type of solid | Particles or components | Main bonding or attraction |
|---|---|---|
| Giant ionic | Positive and negative ions | Strong electrostatic attractions between oppositely charged ions |
| Metallic | Positive metal ions and delocalised electrons | Electrostatic attraction between positive ions and delocalised electrons |
| Giant covalent | Atoms | Covalent bonds extending throughout the structure |
| Simple molecular | Separate molecules | Intermolecular forces between molecules |
| Hydrogen-bonded molecular | Separate molecules | Hydrogen bonds between molecules |
Giant ionic lattices
An ionic solid contains a giant three-dimensional lattice of oppositely charged ions. The ions are arranged so that positive ions have negative ions around them and negative ions have positive ions around them.
The lattice is held together by strong electrostatic attractions between oppositely charged ions. These attractions act in all directions throughout the structure.
Sodium chloride as an ionic lattice
Solid sodium chloride contains Na⁺ ions and Cl⁻ ions arranged in a regular repeating pattern. It does not contain separate NaCl molecules.
Incorrect description: Sodium chloride consists of NaCl molecules held together in a lattice.
Correct description: Sodium chloride consists of a giant lattice of Na⁺ and Cl⁻ ions held together by strong electrostatic attractions.
Why do ionic solids usually have high melting points?
Melting an ionic solid requires the ions to become mobile enough to move relative to one another. A large amount of energy is needed to overcome enough of the strong electrostatic attractions holding the ions in their lattice positions.
Ionic solids therefore generally have high melting and boiling points. The exact strength of the lattice depends particularly on the charges and sizes of the ions.
Enough attractions must be overcome for the ions to leave their fixed lattice positions and move relative to one another. The ions themselves remain charged in the molten state.
Why does magnesium oxide have a higher melting point than sodium chloride?
Sodium chloride contains Na⁺ and Cl⁻ ions, whereas magnesium oxide contains Mg²⁺ and O²⁻ ions.
The ionic charges in MgO are greater, giving stronger electrostatic attractions between oppositely charged ions. More energy is therefore required to overcome the attractions in the MgO lattice.
Exam reasoning: Mg²⁺ and O²⁻ have greater ionic charges than Na⁺ and Cl⁻. The electrostatic attractions between oppositely charged ions are therefore stronger in MgO, so more energy is required to overcome them.
Ionic radius also matters. For ions of comparable charge, smaller ions allow the opposite charges to approach more closely, strengthening their electrostatic attraction.
Electrical conductivity of ionic substances
An electric current requires charged particles that are free to move.
- In a solid ionic lattice, the ions are charged but fixed in position. The solid therefore does not conduct electricity.
- When the substance is molten, the regular solid lattice has broken down and the ions can move through the liquid. The melt conducts electricity.
- When a soluble ionic substance is dissolved in water, mobile aqueous ions can carry charge through the solution.
The charged particles must also be able to move. Solid sodium chloride contains ions, but the ions cannot move freely through the rigid lattice.
Why are ionic solids brittle?
If a force causes one layer of ions to shift relative to another, ions with the same charge may become aligned beside one another.
Strong electrostatic repulsion then develops between these like-charged ions. Instead of bending gradually, the crystal tends to split. Ionic solids are therefore brittle rather than malleable.
Metallic lattices
A metal consists of a giant structure of positive metal ions surrounded by delocalised electrons. These electrons are not confined to one particular atom and can move throughout the metallic structure.
A metallic bond is the electrostatic attraction between positive metal ions and delocalised electrons.
Describing a metal merely as “closely packed atoms” leaves out the electrostatic attraction responsible for holding the metallic structure together.
Why do metals conduct electricity?
The delocalised electrons are mobile. When a potential difference is applied, they undergo a net movement through the metal and carry charge.
Metals can therefore conduct electricity in both the solid and molten states. When the metal melts, the ordered solid lattice is disrupted, but delocalised electrons are still available to move.
Why are many metals malleable and ductile?
Metallic bonding is non-directional. If layers of positive metal ions slide past one another, the ions remain attracted to the surrounding delocalised electrons.
The structure can therefore change shape without the bonding being broken completely. This allows metals to be hammered into sheets and drawn into wires.
Ionic lattice: shifting layers can bring like-charged ions beside one another, producing repulsion and fracture.
Metallic structure: shifted metal ions remain attracted to the surrounding delocalised electrons.
Why do metals often have high melting points?
A large amount of energy may be required to overcome enough of the strong electrostatic attractions between the positive metal ions and delocalised electrons.
Metallic melting points vary considerably. Their values depend on factors including the charge and size of the metal ions and the number of electrons contributing to the delocalised electron system.
Giant covalent lattices
In a giant covalent structure, atoms are connected by covalent bonds throughout an extended network. There are no separate molecules within the structure.
Many strong covalent bonds must be broken or substantially disrupted when the structure melts. Giant covalent substances therefore generally have very high melting points.
Diamond: a three-dimensional covalent network
In diamond, each carbon atom forms four covalent bonds with four neighbouring carbon atoms. These bonds are arranged tetrahedrally, producing a rigid three-dimensional network.
- Each carbon atom is bonded to four neighbouring carbon atoms.
- Strong covalent bonds extend throughout the structure.
- The rigid network makes diamond extremely hard.
- A large amount of energy is required to disrupt the covalent network, giving diamond a very high melting point.
Why does diamond not conduct electricity?
Each carbon atom uses all four of its outer-shell electrons in covalent bonds. There are no mobile ions or delocalised electrons available to carry charge through the structure.
Diamond is therefore an electrical insulator.
The mere presence of electrons does not make a substance conductive. Electrons localised in covalent bonds cannot move freely throughout the structure.
Graphite: covalent layers with delocalised electrons
Graphite is also made entirely of carbon atoms, but its structure is different from diamond. Each carbon atom forms three covalent bonds with three neighbouring carbon atoms, producing flat hexagonal layers.
Strong covalent bonds hold the carbon atoms together within each layer. The fourth outer-shell electron from each carbon atom becomes delocalised across the layer.
Adjacent layers are held together by weak instantaneous dipole–induced dipole attractions, which are a type of van der Waals’ attraction.
Why is graphite soft and slippery?
The attractions between neighbouring graphite layers are much weaker than the covalent bonds within each layer. The layers can therefore slide over one another relatively easily.
This explains why graphite is soft and can be used as a lubricant and in pencil leads.
Why does graphite conduct electricity?
The delocalised electrons can move across each carbon layer. They can therefore carry charge along the layers.
The electrons are delocalised within the layers, not in the spaces between the layers.
| Feature | Diamond | Graphite |
|---|---|---|
| Bonds formed by each carbon atom | Four | Three |
| Overall arrangement | Rigid three-dimensional tetrahedral network | Flat hexagonal layers |
| Delocalised electrons | None | One electron per carbon atom contributes to the delocalised system |
| Electrical conductivity | Does not conduct | Conducts along the layers |
| Hardness | Very hard | Soft and slippery |
| Melting point | Very high | Very high |
Diamond and graphite both have very high melting points because strong covalent bonds extend throughout their giant structures.
Their different hardness and conductivity arise from differences in how the carbon atoms and electrons are arranged.
Simple molecular solids
A simple molecular substance consists of separate, discrete molecules. When the substance is solid, these molecules occupy positions in a regular molecular crystal lattice.
Strong covalent bonds hold the atoms together within each molecule, while intermolecular forces act between neighbouring molecules. When the solid melts, the ordered lattice breaks down because intermolecular forces are overcome, but the covalent bonds within the molecules normally remain intact.
The molecules are arranged regularly while the substance is solid, but each molecule remains a separate particle. There is no continuous network of covalent bonds extending from one molecule to another.
| Location | What is present? | Attraction |
|---|---|---|
| Within each molecule | Atoms joined together | Strong covalent bonds |
| Between neighbouring molecules | Separate molecules | Intermolecular forces |
Solid iodine as a molecular crystal
Solid iodine contains separate I₂ molecules arranged in a regular molecular crystal lattice.
- Each I₂ molecule contains two iodine atoms joined by a covalent bond.
- Instantaneous dipole–induced dipole attractions act between neighbouring I₂ molecules.
- These intermolecular attractions hold the molecular crystal together.
What happens when iodine melts or sublimes?
When iodine melts or sublimes, the ordered solid lattice breaks down as instantaneous dipole–induced dipole attractions between neighbouring I₂ molecules are overcome.
The covalent I—I bond inside each I₂ molecule is not broken. The particles in liquid iodine and iodine vapour remain I₂ molecules.
Incorrect: The covalent bonds in iodine are broken when iodine melts.
Correct: Intermolecular attractions between I₂ molecules are overcome. The I—I covalent bond within each molecule remains intact.
Why do simple molecular substances usually have lower melting points?
Melting a molecular solid requires intermolecular forces between its molecules to be overcome. These forces are generally much weaker than ionic attractions, metallic bonding or covalent bonds extending throughout a giant covalent structure.
Less energy is therefore usually required to melt a simple molecular solid.
This does not mean every molecular substance has an extremely low melting point. Larger, more polarisable molecules can experience stronger instantaneous dipole–induced dipole attractions than smaller molecules.
Why do simple molecular substances usually not conduct electricity?
The molecules are electrically neutral, and there are normally no mobile ions or delocalised electrons. Simple molecular substances therefore generally do not conduct electricity in either the solid or liquid state.
The hydrogen-bonded lattice of ice
Ice is a molecular solid. It contains separate H₂O molecules arranged in an ordered solid lattice rather than a continuous network of covalent O—H bonds.
Within each H₂O molecule, oxygen and hydrogen atoms are joined by strong covalent bonds. Between neighbouring molecules, hydrogen bonds form between the δ⁺ hydrogen atom of one molecule and a lone pair on the δ⁻ oxygen atom of another.
How are water molecules arranged in ice?
In the ice lattice, each water molecule is hydrogen-bonded to four neighbouring water molecules in an approximately tetrahedral arrangement.
Each H₂O molecule can donate two hydrogen bonds through its two hydrogen atoms and accept two hydrogen bonds through two lone pairs on its oxygen atom.
This produces an open three-dimensional molecular lattice containing considerable empty space. When viewed in suitable directions, the arrangement gives rise to hexagonal patterns.
The water molecules adopt positions that allow an extended hydrogen-bonding arrangement rather than simply packing as closely together as possible.
Why is ice less dense than liquid water?
Density depends on the mass contained within a given volume. The open hydrogen-bonded structure of ice keeps the water molecules relatively far apart.
When ice melts, part of this ordered structure is disrupted. The rigid lattice partially collapses and the water molecules can, on average, pack more closely together.
The same mass of liquid water therefore occupies a smaller volume than ice. Liquid water is denser, which is why ice floats.
Are hydrogen bonds broken when ice melts?
Some hydrogen bonds maintaining the rigid ice lattice are disrupted when ice melts. However, hydrogen bonds continue to form and break continually between molecules in liquid water.
The covalent O—H bonds within the H₂O molecules remain intact during ordinary melting.
Within each H₂O molecule: strong covalent O—H bonds.
Between neighbouring H₂O molecules: hydrogen bonds.
Melting ice disrupts the organised intermolecular hydrogen-bonded lattice; it does not decompose the H₂O molecules.
Comparing the main solid structures
The following comparison brings the main patterns together. Read each row from the particles to the bonding and then to the resulting properties.
| Structure | Particles or components | What holds the solid together? | Typical melting point | Electrical conductivity | Example |
|---|---|---|---|---|---|
| Giant ionic | Positive and negative ions | Strong electrostatic attractions between oppositely charged ions | High | Not when solid; conducts when molten or when mobile ions are present in solution | NaCl, MgO |
| Metallic | Positive metal ions and delocalised electrons | Electrostatic attraction between positive ions and delocalised electrons | Often high, but variable | Conducts when solid and molten | Cu |
| Giant covalent: diamond | Carbon atoms | Covalent bonds throughout a three-dimensional network | Very high | Does not conduct | Diamond |
| Giant covalent: graphite | Carbon atoms in layers | Covalent bonds within layers and weak attractions between layers | Very high | Conducts along the layers | Graphite |
| Simple molecular solid | Separate molecules | Intermolecular forces between molecules | Usually relatively low | Usually does not conduct | Solid I₂ |
| Hydrogen-bonded molecular solid | Separate molecules | Hydrogen bonds between molecules | Influenced strongly by hydrogen bonding | Pure ice does not conduct significantly | Ice, H₂O |
Which structures should be described as giant?
The word giant is appropriate when bonding or electrostatic attraction extends continuously through a very large number of particles.
- Giant ionic lattices contain continuous arrays of oppositely charged ions.
- Metallic bonding extends throughout a giant metallic structure.
- Giant covalent structures contain extended networks of covalently bonded atoms.
Solid iodine and ice contain separate molecules. Their molecules occupy regular positions in crystalline lattices, but the substances are still molecular solids.
Use lattice when discussing the regular arrangement of particles in a crystalline solid. Use giant when bonding or electrostatic attraction extends continuously through the structure rather than ending within separate molecules.
Reasoning from structure to property
Good explanations do not stop after naming a structure. They connect the microscopic structure to the observed property through a clear chain of reasoning.
Example 1: explaining melting point
Explain why sodium chloride has a much higher melting point than iodine.
Sodium chloride has a giant ionic lattice containing Na⁺ and Cl⁻ ions held together by strong electrostatic attractions. A large amount of energy is required to overcome enough of these attractions for the ions to move.
Solid iodine contains separate I₂ molecules arranged in a molecular crystal lattice. Only weaker instantaneous dipole–induced dipole attractions between the molecules are overcome during melting. Less energy is therefore required to melt iodine.
Example 2: explaining electrical conductivity
Explain why graphite conducts electricity but diamond does not.
In graphite, each carbon atom forms three covalent bonds and contributes one electron to a delocalised electron system. These electrons can move along the layers and carry charge.
In diamond, each carbon atom forms four covalent bonds. All four outer-shell electrons are localised in covalent bonds, so there are no mobile charged particles available to carry current.
Example 3: explaining mechanical behaviour
Explain why copper is malleable but sodium chloride is brittle.
In copper, layers of positive metal ions can slide past one another while remaining attracted to surrounding delocalised electrons. The metal can therefore change shape without immediately breaking.
In sodium chloride, shifting the ion layers can bring ions with the same charge beside one another. Their strong electrostatic repulsion causes the crystal to split.
A reliable structure–property sequence
- Name the type of structure.
- Identify the particles or components.
- State the bonding or forces present.
- Explain what must move or what attractions must be overcome.
- Link this directly to the observed property.
The important reasoning lies between them: identify the particles, identify the attractions and decide whether suitable charged particles can move.
Common mistakes with solid structures
| Common mistake | Why it is wrong | Correct approach |
|---|---|---|
| Treating “lattice” and “giant” as interchangeable | A lattice is an ordered solid-state arrangement; a giant structure extends continuously through many particles | Recognise that molecular solids can form lattices while still containing separate molecules |
| Calling every crystalline solid a giant structure | Simple molecular solids contain discrete molecules | Identify whether the particles remain separate molecules |
| Describing solid sodium chloride as NaCl molecules | An ionic lattice does not contain discrete NaCl molecules | Describe a giant lattice of Na⁺ and Cl⁻ ions |
| Saying that ionic solids conduct because they contain ions | The ions in the solid are fixed in position | State that mobile ions are required for conduction |
| Saying that metallic bonding is attraction between metal atoms | This omits the positive ions and delocalised electrons | Define metallic bonding as electrostatic attraction between positive metal ions and delocalised electrons |
| Saying that diamond is hard simply because each carbon forms four bonds | The important feature is the rigid extended covalent network | State that strong covalent bonds extend throughout a three-dimensional structure |
| Saying that graphite has weak covalent bonds | The covalent bonds within each layer are strong | State that the attractions between neighbouring layers are weak |
| Saying that graphite’s electrons are delocalised between its layers | The delocalised electrons extend across each carbon layer | Describe delocalisation within the layers |
| Saying that iodine melts because its covalent bonds break | The I—I bonds remain intact during melting | State that intermolecular attractions between I₂ molecules are overcome |
| Describing ice as a giant covalent structure | Ice contains separate H₂O molecules | Describe a hydrogen-bonded molecular lattice |
| Saying that all hydrogen bonds disappear when ice melts | Hydrogen bonding continues in liquid water | State that the rigid, ordered hydrogen-bonded lattice is disrupted |
| Explaining a high melting point only by saying “the bonds are strong” | The particles and type of attraction have not been identified | Name the particles, attraction and why substantial energy is required |
Solid-structure checklist
Use “lattice” correctly
- A lattice is a regular arrangement in a crystalline solid.
- The term does not identify one particular type of bonding.
- A molecular solid can have a molecular crystal lattice.
- The ordered lattice is disrupted when the solid melts.
Identify the particles
- Ionic structures contain positive and negative ions.
- Metallic structures contain positive ions and delocalised electrons.
- Giant covalent structures contain atoms.
- Simple molecular solids contain separate molecules.
Explain the attractions
- State electrostatic attraction for ionic and metallic structures.
- State covalent bonds for giant covalent structures.
- Identify the correct intermolecular force for molecular solids.
- Distinguish forces between molecules from bonds within molecules.
Connect structure to property
- For melting point, explain what attractions must be overcome.
- For conductivity, identify mobile charged particles.
- For hardness, consider the rigidity of the structure.
- For brittleness or malleability, consider what happens when layers move.
These ten questions go beyond simple definitions. Use particle-level reasoning to identify structures and explain melting point, conductivity and mechanical behaviour.
- Apply structure–property reasoning to unfamiliar examples.
- Distinguish bonding from intermolecular forces.
- Test common examination misconceptions.
Strong answers in this topic identify the particles, state the correct attraction and then connect that microscopic structure directly to the observed property.
Frequently asked questions
Questions students often ask about solid lattice structures
These questions clarify the distinctions that matter most when linking microscopic structures to observable physical properties.
Is the term “lattice” reserved for ionic substances?
No. A lattice describes a regular, repeating arrangement of particles in a crystalline solid. Ionic solids, metals, giant covalent substances and molecular crystals may all be described using lattice ideas.
Ionic lattices are particularly common in Chemistry explanations, which is probably why students sometimes associate the word only with ionic compounds.
Why do we usually describe iodine as simple molecular rather than simply calling it a molecular lattice?
Simple molecular identifies the type of structure: the substance contains separate molecules. Molecular lattice describes how those molecules are arranged specifically in the crystalline solid.
It is therefore useful to say that iodine is a simple molecular substance and that solid iodine contains I₂ molecules arranged in a molecular crystal lattice.
Does every solid have a lattice?
Crystalline solids possess long-range ordered arrangements that can be described using lattice structures. Amorphous solids do not possess the same long-range repeating crystalline order.
Is a lattice the same thing as a bond?
No. A lattice describes the ordered arrangement of particles. A bond or intermolecular force describes the attraction responsible for holding those particles together.
Does a liquid have a lattice?
A liquid does not possess the fixed, long-range repeating arrangement characteristic of a crystalline solid. Its particles remain close together but can move relative to one another.
Why is sodium chloride not described as containing NaCl molecules?
The formula NaCl gives the simplest ratio of Na⁺ ions to Cl⁻ ions in the lattice. It does not represent a discrete molecule containing one sodium ion bonded only to one chloride ion.
Do ionic bonds disappear completely when an ionic solid melts?
Melting provides enough energy for the ions to leave their fixed lattice positions and move relative to one another. The ions remain charged and continue to experience electrostatic interactions in the molten state.
Why can molten sodium chloride conduct but molten iodine cannot?
Molten sodium chloride contains mobile Na⁺ and Cl⁻ ions that can carry charge. Molten iodine contains neutral I₂ molecules and has no comparable mobile ions or delocalised electrons.
Are delocalised electrons in a metal completely free electrons?
They are free to move throughout the metallic structure rather than being confined to one particular atom, but they remain part of the metal and are electrostatically attracted to the positive metal ions.
Why does graphite have a high melting point if its layers are weakly attracted?
The weak attractions between graphite layers explain why the layers slide easily. However, each layer is itself an extended network of carbon atoms joined by strong covalent bonds.
Extensive strong covalent bonding must be disrupted to melt the giant structure, so graphite has a very high melting point.
Is graphite a simple molecular substance because it has separate layers?
No. Each layer is an extended network of covalently bonded carbon atoms. There are no discrete graphite molecules, so graphite is a giant covalent structure.
Are intermolecular forces always weak?
Intermolecular forces are weaker than typical covalent bonds within molecules, but their strength varies considerably. Hydrogen bonds are stronger than ordinary instantaneous dipole–induced dipole attractions between similar small molecules.
Instantaneous dipole–induced dipole attractions can also become substantial for large, highly polarisable molecules.
Is ice a giant hydrogen-bonded structure?
Hydrogen bonds extend throughout the ice crystal, but the fundamental particles remain separate H₂O molecules. It is therefore clearer to describe ice as a hydrogen-bonded molecular solid or molecular lattice, rather than as a giant covalent structure.
Does melting ice break the O—H bonds in water molecules?
No. Melting disrupts part of the ordered hydrogen-bonded arrangement between H₂O molecules. The strong covalent O—H bonds within each water molecule remain intact.
What should I mention first in an examination question about a solid’s properties?
Start by identifying the type of structure and the particles present. Then state the relevant bonding or intermolecular forces before explaining how these lead to the property being asked about.
A lattice is the ordered arrangement adopted by particles in a crystalline solid, but it should not automatically be confused with a giant structure. Ionic, metallic and giant covalent solids contain bonding or electrostatic attraction extending continuously through the structure. Simple molecular solids such as iodine and ice also form ordered lattices, but their particles remain separate molecules held together by intermolecular forces. Once this distinction is clear, properties such as melting point, conductivity, hardness, brittleness and density can be explained by identifying the particles, their attractions and the available mobile charge carriers.