Structure Types
A concise revision guide to the four types of solid lattice, giant ionic, giant metallic, giant covalent and simple molecular: which particles they contain, what holds them together, and how that decides melting point, conductivity, solubility and the energy changes on melting and boiling.
- 3.1.3.4i-a
- 3.1.3.4i-b
- 3.1.3.4i-c
- 3.1.3.4i-d
- 3.1.3.4iii-a
- 3.1.3.4iii-b
- 3.1.3.4iv
What these spec points say
- 3.1.3.4i-a know ionic crystal structures
- 3.1.3.4i-b know metallic crystal structures
- 3.1.3.4i-c know macromolecular (giant covalent) crystal structures
- 3.1.3.4i-d know molecular crystal structures
- 3.1.3.4iii-a relate the melting point of materials to the type of structure and bonding present
- 3.1.3.4iii-b relate the conductivity of materials to the type of structure and bonding present
- 3.1.3.4iv explain the energy changes associated with changes of state
The Four Types of Solid Lattice
Choose a step to watch each type of solid lattice form, see what holds its particles together, then press Heat it to see what melting that solid has to overcome.
© Dr. Mohammed Al-Fatah – onlinelearningsystem.net
GCSE Recap: Bonding Types and Changes of State
Answer three quick GCSE questions before you start.
The Four Types of Solid Lattice
Every crystalline solid is a regular, repeating arrangement of particles: a lattice. What decides a solid’s properties is not the lattice pattern itself but which particles sit at the lattice points and which forces hold them there. There are four cases to know.
| Lattice type | Particles | Forces holding the lattice together | Examples |
|---|---|---|---|
| Giant ionic | Positive and negative ions | Strong electrostatic attraction between oppositely charged ions, acting in all directions | NaCl, MgO, CaF2 |
| Giant metallic | Positive metal ions and delocalised electrons | Strong electrostatic attraction between the ions and the sea of delocalised electrons | Mg, Cu, Na |
| Giant covalent (macromolecular) | Atoms | Strong covalent bonds throughout the whole structure | Diamond, graphite, SiO2 |
| Simple molecular | Molecules | Weak intermolecular forces between molecules (the covalent bonds inside each molecule are strong but are not broken on melting) | I2, ice, C60, CO2 |
The word giant means the bonding continues throughout the crystal, so there is no separate molecule to point to.
In a simple molecular solid there are two different forces: strong covalent bonds within each molecule and weak forces between them, and only the weak forces are overcome when the solid melts.
Key idea: Three giant lattices, one molecular lattice. Melting a giant lattice breaks strong bonds; melting a molecular lattice only overcomes weak intermolecular forces.
Quick Check: Spot the Molecular Solids
Click every substance that would form a simple molecular lattice.
Melting Point Follows the Force That Must Be Overcome
Melting a solid means giving the particles enough energy to break free from their lattice positions. The energy needed depends on the force holding them, so melting points fall into two clear bands.
| Substance | Lattice type | What must be overcome on melting | Melting point / °C |
|---|---|---|---|
| Sodium chloride | Giant ionic | Strong electrostatic attraction between ions | 801 |
| Magnesium | Giant metallic | Strong attraction between ions and delocalised electrons | 650 |
| Diamond | Giant covalent | Strong covalent bonds | above 3500 (sublimes) |
| Iodine | Simple molecular | Weak van der Waals forces between I2 molecules | 114 |
| Ice | Simple molecular | Hydrogen bonds between H2O molecules | 0 |
Within a band the details matter.
| Type | What gives a higher melting point |
|---|---|
| Ionic compounds | A higher charge and a smaller ionic radius give stronger attraction and a higher melting point (MgO melts at 2852 °C, NaCl at 801 °C) |
| Metals | More delocalised electrons per ion and a smaller ion give a higher melting point (Mg above Na) |
Among simple molecular substances two things decide the melting point. For molecules of similar size, that is with similar numbers of electrons, the order is hydrogen bonds above permanent dipoles above van der Waals forces.
Size matters as well, because a larger molecule has more electrons and therefore stronger van der Waals forces.
That is why iodine, with only van der Waals forces, melts at 114 °C, well above ice at 0 °C, although ice has hydrogen bonds: an I2 molecule is far larger than an H2O molecule.
Exam sentence: Sodium chloride has a high melting point because a lot of energy is needed to overcome the strong electrostatic attraction between oppositely charged ions in the giant lattice.
Iodine has a low melting point because only weak van der Waals forces between I2 molecules need to be overcome.
Quick Check: Explain Two Melting Points
Write a short explanation, then compare it with the mark points and the model answer.
Electrical Conductivity Needs Mobile Charged Particles
A substance conducts electricity only if it contains charged particles that are free to move. Apply that one test to each lattice type in each state.
| Lattice type | Solid | Molten or dissolved | Reason |
|---|---|---|---|
| Giant ionic | Does not conduct | Conducts | Ions are fixed in the solid lattice; when molten or in solution they are free to move and carry charge |
| Giant metallic | Conducts | Conducts | Delocalised electrons are free to move through the lattice in both states |
| Giant covalent | Does not conduct (except graphite) | Does not conduct | All outer electrons are held in localised covalent bonds; graphite alone has one delocalised electron per carbon |
| Simple molecular | Does not conduct | Does not conduct | Molecules are neutral and there are no delocalised electrons or ions |
The ionic case is the one examiners test most: the solid is an insulator, the melt is a conductor, and the reason is the mobility of the ions, not whether ions exist.
Exam focus: Never write “ionic compounds conduct electricity”. Write “ionic compounds conduct when molten or aqueous because the ions are free to move; in the solid the ions are held in fixed positions”.
Quick Check: When Does Potassium Bromide Conduct?
Use the data to choose the structure and the reason for the conductivity.
Solubility and Hardness
- Giant ionic: many dissolve in water because the polar water molecules attract the ions and hydrate them; insoluble in non-polar solvents. Hard but brittle, because shifting a layer brings like charges together and the layers repel.
- Giant metallic: insoluble in all common solvents; malleable and ductile because layers of ions slide past each other while the electron sea keeps the attraction.
- Giant covalent: insoluble in everything; very hard (diamond) or soft and slippery (graphite, where layers slide).
- Simple molecular: soft, easily crushed; solubility follows polarity (“like dissolves like”): iodine dissolves in hexane, not in water; sugar and ethanol dissolve in water because they can hydrogen bond with it.
Key idea: Solubility is about the forces between solute and solvent; hardness is about how easily the lattice can be deformed without the forces breaking.
Quick Check: Shatter, Bend or Dissolve
Drag the words into place to explain three test results.
Energy Changes on Changes of State
When a solid melts or a liquid boils, energy is absorbed to overcome the forces between particles while the temperature stays constant: the enthalpy of fusion and enthalpy of vaporisation. The size of each depends on the lattice type.
| Change | Giant lattices | Simple molecular |
|---|---|---|
| Melting | Large energy input: strong ionic, metallic or covalent bonds are partly broken | Small energy input: some weak intermolecular forces are overcome |
| Boiling | Very large: the remaining strong bonds are broken so the particles separate completely | Small: all the remaining intermolecular forces are overcome; the covalent bonds inside the molecules survive |
| Sign of the enthalpy change | Endothermic (positive) | Endothermic (positive) |
| Reverse change (freezing, condensing) | Exothermic: the same amount of energy is released as the forces re-form | Exothermic |
On a heating curve the two flat plateaus are melting and boiling: the energy supplied is going into overcoming forces, not into raising the temperature. The plateaus are longer and sit at higher temperatures for giant lattices than for simple molecular substances.
Energy supplied during a plateau overcomes forces between particles; the temperature does not rise until they are overcome.
Exam sentence: During melting the temperature stays constant because the energy supplied is used to overcome the forces between the particles rather than to increase their kinetic energy.
Quick Check: The Plateau for Sulfur
Decide where the energy goes while the temperature stays constant.
Watch: deducing a structure from data
Four substances, their melting points, conductivity and solubility. What is the structure of each?
Deducing the Structure from Data
Exam questions often give a table of properties for unnamed substances and ask for the structure and bonding. Work through the tests in order.
| Question to ask | If yes | If no |
|---|---|---|
| Does it conduct as a solid? | Metallic (or graphite) | Go on |
| Does it conduct when molten but not as a solid? | Giant ionic | Go on |
| Is the melting point very high (above about 1000 °C) with no conduction in any state? | Giant covalent | Go on |
| Is the melting point low and there is no conduction? | Simple molecular | Re-check the data |
Then confirm with solubility: ionic solids often dissolve in water, molecular solids dissolve in solvents of similar polarity, giant covalent and metallic solids dissolve in neither.
Keep to that order. Plenty of ionic solids also melt far above 1000 °C, MgO at 2852 °C, so a high melting point on its own does not prove giant covalent. It is the conduction test, done first, that separates the two.
One warning about data tables. A few simple molecular substances react with water to form ions, so their solutions conduct although the pure substance does not.
Hydrogen chloride is the example: it is a molecular gas, but hydrochloric acid conducts. A low melting point with conduction only in aqueous solution points to this, not to a giant ionic lattice.
Exam focus: State the structure, then justify it with two properties from the data, naming the particles and the forces each time.
Test conductivity first, then melting behaviour, to decide between metallic, ionic, giant covalent and simple molecular.
Quick Check: Six Unknown Solids
Read each set of data and pick the structure as quickly as you can.
Quick Check: The Whole Page in Six Rounds
Pick the one accurate statement in each round to pull the four lattice types together.
Common Exam Mistakes
- Saying a simple molecular solid melts when “the covalent bonds break”. Only the intermolecular forces are overcome; the molecules stay intact.
- Saying ionic compounds conduct electricity without stating the condition (molten or aqueous) and the reason (ions free to move).
- Calling the attraction in a metal “between atoms” or “between ions”. It is between positive ions and delocalised electrons.
- Describing diamond as a molecule. It is a giant covalent lattice with no separate molecules.
- Explaining a high melting point by “strong bonds” alone. Name the type of bond or force and the particles it acts between.
Exam sentence: Particles, force between them, consequence: every structure-and-bonding answer follows that pattern.
Where This Sits in the AQA Specification
The specification asks for the four crystal structures (ionic, metallic, macromolecular and molecular) with sodium chloride, magnesium, diamond, graphite, ice and iodine as the named examples.
You need to relate melting point and conductivity to the structure and bonding present, and to explain the energy changes on changes of state. You may also be asked to draw a lattice fragment with a specified number of particles.
Diamond and graphite are on the Giant Covalent Structures page, and ice and iodine on Simple Molecular Structures.
AQA focus: Name the particles, name the force between them, then state what that force means for melting point and conductivity. Three steps, every time.
FAQs
These questions cover the structure-type mistakes that cost marks most often.
Why does iodine melt above ice if hydrogen bonds are stronger than van der Waals forces?
Hydrogen bonds are stronger only for molecules of similar size. An I2 molecule has 106 electrons against 10 in H2O, so its van der Waals forces add up to more. Iodine melts at 114 °C, ice at 0 °C.
Do covalent bonds break when a simple molecular solid melts?
No. Only the intermolecular forces between the molecules are overcome. The covalent bonds inside each molecule stay intact, which is why iodine vapour still contains I2 molecules.
Why does solid sodium chloride not conduct when it contains ions?
The ions are fixed in the lattice and cannot move. Sodium chloride conducts only when molten or dissolved in water, because the ions are then free to move and carry charge.
How can I tell a giant covalent solid from an ionic one using data?
Both have very high melting points, so check conductivity. An ionic solid conducts when molten or in aqueous solution. A giant covalent solid such as diamond or silicon dioxide conducts in no state; graphite is the exception.
Why does a solution of hydrogen chloride conduct if HCl is a simple molecule?
Hydrogen chloride reacts with water to form H+(aq) and Cl–(aq) ions. The pure substance is molecular and does not conduct; the ions only appear in solution.
Copyright notice: This OLS revision content, including the explanations, layout, diagrams, tables and embedded learning structure, is authored for Online Learning System by Dr. Mohammed Al-Fatah. It may not be copied, reproduced, redistributed or adapted without written permission.
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