Crystal Structures and Physical Properties — AQA A-level Chemistry

SPEC 3.1.3.4 • BONDING

Bonding & Physical Properties

How structure and bonding determine melting point, conductivity, hardness and solubility

Why Structure Matters

The physical properties of a substance depend on two key factors:

1. Type of Bonding

The forces holding particles together: ionic bonds, covalent bonds, metallic bonds, or intermolecular forces.

2. Crystal Structure

How particles are arranged: giant ionic, giant metallic, giant covalent (macromolecular), or simple molecular.

Key Properties to Explain

Melting/Boiling Point: Energy needed to overcome forces between particles

Electrical Conductivity: Presence of mobile charged particles (ions or delocalised electrons)

Hardness: Resistance to deformation; depends on bond strength and arrangement

Solubility: Whether particles can interact with solvent molecules

Four Crystal Structure Types

Structure TypeParticlesForces to OvercomeExample

Giant IonicCations and anionsStrong electrostatic attractionsNaCl, MgO

Giant MetallicMetal cations + delocalised e⁻Metallic bondsMg, Fe, Cu

Giant CovalentAtoms in 3D networkStrong covalent bondsDiamond, graphite, SiO₂

Simple MolecularSmall moleculesWeak intermolecular forcesI₂, ice, CO₂

Common Exam Mistake

When explaining melting points, always state what forces are being broken. For simple molecular substances, you break intermolecular forces (NOT covalent bonds) when melting.

Giant Ionic Structures

Ions form a lattice structure — an evenly distributed crystalline arrangement. Ions are arranged in a regular repeating pattern so that positive charges cancel out negative charges. The attraction between cations and anions occurs in all directions — each ion is attracted to all of the oppositely charged ions around it.

Key Points

The electrostatic forces of attraction bring cations and anions together to form the lattice structure. Ions are arranged in a regular repeating pattern so that positive charges cancel out negative charges - the lattice has a neutral charge overall. The attraction between cations and anions occurs in all directions - each ion is attracted to all of the oppositely charged ions around it.

Structure Details

Anions are larger than cations — when metals lose electrons, the ion shrinks; when non-metals gain electrons, the ion expands. The type of lattice formed depends on the relative sizes of the ions, which are arranged in an alternating fashion. The ionic lattices of NaCl and MgO are both cubic.

Properties of Ionic Compounds

PropertyObservationExplanation

Melting PointHigh (NaCl: 801°C)Strong electrostatic attractions between oppositely charged ions require large amounts of energy to overcome

Conductivity (solid)Does not conductIons are in fixed positions in the lattice and cannot move

Conductivity (molten/dissolved)Conducts electricityIons are free to move and carry charge

HardnessHard but brittleDisplacing layers brings like charges together, causing repulsion and fracture

SolubilityOften soluble in waterPolar water molecules can surround and stabilise ions (hydration)

Exam Tip

When asked why ionic compounds conduct when molten, always say: "The ions are free to move and can carry charge."

Charge Density Effect on Melting Point

Melting and boiling points increase with charge density of the ions due to greater electrostatic attraction. For example, Mg²⁺O²⁻ has a higher melting point than Na⁺Cl⁻ because of the higher ionic charges. Ionic compounds are brittle — when a force is applied, ions of like charge are brought together causing repulsion and the crystal splits apart.

Worked Example01

Explain why magnesium oxide (MgO) has a higher melting point than sodium chloride (NaCl). [3 marks]

1

MgO contains Mg²⁺ and O²⁻ ions, while NaCl contains Na⁺ and Cl⁻ ions.

2

The ions in MgO have higher charges (2+ and 2−) compared to NaCl (1+ and 1−).

3

The Mg²⁺ ion is also smaller than Na⁺, so the charge density is higher.

4

Therefore, the electrostatic attractions between ions are stronger in MgO, requiring more energy to overcome.

Key points: Higher ionic charges → stronger electrostatic attraction → higher melting point

Giant Metallic Structures

Metal atoms with electrons in the outer shell

When metals are closely packed, their outer electrons dissociate forming positive metal ions.

Positive metal ions in a "sea" of delocalised electrons

Once the outer electrons are delocalised, the positive metal ions are suspended in the "sea" of electrons. The delocalised electrons are free to move through the metal. Metallic bonding involves very strong electrostatic forces of attraction between the positive metal ions and the delocalised electrons.

Properties of Metals (Magnesium)

PropertyObservationExplanation

Melting PointVariable but often high (Mg: 650°C)Strong metallic bonds between cations and delocalised electrons require energy to break

ConductivityGood conductors (solid and liquid)Delocalised electrons are free to move and carry charge throughout the structure

Malleability & DuctilityCan be hammered/drawn into wiresLayers of ions can slide over each other without breaking bonds — the sea of electrons adjusts

Thermal ConductivityGood conductors of heatWhen heated, cations vibrate more vigorously and transfer kinetic energy to delocalised electrons, which carry it rapidly through the structure

SolubilityInsoluble in waterMetallic bonds are very strong — water cannot break the electrostatic attractions

Factors Affecting Metallic Bond Strength

1. Charge on the cation: Higher charge → stronger attraction (Mg²⁺ > Na⁺)

2. Size of the cation: Smaller ion → higher charge density → stronger attraction

3. Number of delocalised electrons: More electrons → stronger "sea" → stronger bonding

Alloys

When other atoms (such as carbon) are added to a metal structure, this creates an alloy. Alloys are much stronger than pure metals because the different-sized atoms disrupt the regular layers, preventing them from sliding over each other as easily.

Giant Covalent: Diamond

In giant covalent structures (also called macromolecular), atoms are bonded by strong covalent bonds extending throughout the entire structure. There are no discrete molecules.

Each carbon atom forms 4 very strong covalent bonds with others in a tetrahedral arrangement. This creates a rigid 3D network.

Diamond Structure

Each carbon atom is covalently bonded to 4 other carbon atoms in a tetrahedral arrangement. This creates a rigid 3D network with no weak points.

Allotropes

Diamond, graphite, and buckminsterfullerene (C₆₀) are all allotropes of carbon — different structural forms of the same element. They have different physical properties because their atoms are arranged differently.

Properties of Diamond

PropertyObservationExplanation

Melting PointExtremely high (~3550°C)Many strong covalent bonds must be broken; bonds extend in all directions

HardnessHardest natural substanceRigid tetrahedral structure with strong bonds in all directions — no planes of weakness

Electrical ConductivityDoes not conductAll 4 outer electrons on each carbon are used in bonding — no delocalised electrons

Thermal ConductivityGood conductor of heatStrong interatomic bonds transmit thermal vibrations rapidly through the rigid structure

SolubilityInsolubleCovalent bonds are too strong to be broken by solvation — no interaction with water

Exam Language

For diamond's high melting point, say: "Many strong covalent bonds must be broken, which requires a large amount of energy."

Giant Covalent: Graphite

Graphite is also pure carbon but has a completely different structure — layers of hexagonal rings with weak forces between layers.

Each carbon atom forms 3 strong covalent bonds in a planar hexagonal pattern, with each carbon contributing 1 delocalised electron. Multiple stacked layers are held by weak van der Waals forces.

Properties of Graphite

Each carbon in graphite bonds to only 3 others, leaving one electron per carbon delocalised across the layers.

PropertyObservationExplanation

Melting PointVery high (~3700°C)Strong covalent bonds within layers must be broken

HardnessSoft and slipperyWeak van der Waals forces between layers allow them to slide easily

Electrical ConductivityConducts along the layersOne electron per carbon is delocalised and free to move within layers

DensityLower than diamondWeak intermolecular forces between layers mean greater separation — less tightly packed

Diamond vs Graphite

Same element, different structures:

Diamond: 4 bonds per C, tetrahedral, hard, insulator

Graphite: 3 bonds per C, layered hexagons, soft, conducts (parallel to layers)

Worked Example02

Graphite is used as a lubricant in machinery, but diamond is one of the hardest known materials. Explain why these two allotropes of carbon have such different mechanical properties. [4 marks]

1

In diamond, each carbon atom forms 4 strong covalent bonds in a tetrahedral arrangement, creating a rigid 3D network.

2

To deform diamond, you must break these strong covalent bonds, which requires a very large amount of energy — hence diamond is extremely hard.

3

In graphite, each carbon forms only 3 covalent bonds within flat hexagonal layers.

4

The layers are held together by weak van der Waals forces, so layers can easily slide over each other when force is applied.

5

This sliding action makes graphite soft and slippery — ideal as a lubricant.

Key: Diamond = 3D network of strong bonds (hard); Graphite = layers with weak interlayer forces (soft, slippery)

Worked Example03

Substance X has a melting point of 801°C. When solid, it is an insulator, but when molten it conducts electricity well. Predict the structure of substance X and explain your reasoning. [4 marks]

1

The high melting point (801°C) indicates strong forces must be overcome — this rules out simple molecular (low MP).

2

The substance does not conduct when solid but conducts when molten.

3

This pattern is characteristic of ionic compounds — ions are fixed in the solid lattice but free to move when molten.

4

It is NOT giant covalent because giant covalent substances (like diamond/graphite) do not conduct when molten.

Answer: Substance X has an ionic crystal lattice structure.

Simple Molecular Structures

Simple molecular substances consist of small discrete molecules. Within each molecule, atoms are held by strong covalent bonds. Between molecules, there are only weak intermolecular forces.

Iodine (I₂)

Structure: Diatomic molecules

Forces: Weak van der Waals

MP: 114°C (low)

I₂ molecules in face-centred cubic lattice

Ice (H₂O)

Structure: Open hexagonal lattice

Forces: Hydrogen bonds

MP: 0°C (low)

ICE (H₂O)

Critical Point

When simple molecular substances melt or boil, you overcome the weak intermolecular forces — NOT the strong covalent bonds within molecules. The molecules remain intact.

Properties of Simple Molecular Substances

PropertyObservationExplanation

Melting/Boiling PointLowOnly weak intermolecular forces need to be overcome (not covalent bonds)

Electrical ConductivityDo not conductNo free electrons or ions — molecules are neutral with no mobile charged particles

Solubility in WaterVariablePolar molecules (like water) dissolve; non-polar molecules (like I₂) do not

Why Ice Floats: The Open Hexagonal Structure

In ice, each water molecule forms 4 hydrogen bonds with neighbouring molecules in a fixed tetrahedral arrangement. This creates an open hexagonal lattice with significant empty space between molecules. When ice melts, the rigid hydrogen bond network collapses, allowing molecules to move closer together — so liquid water is denser than ice. This is why ice floats.

Energy Changes in State Changes

When a substance changes state, energy is absorbed or released to overcome intermolecular forces:

Melting/Boiling (endothermic): Energy absorbed → breaks intermolecular forces → particles gain freedom of movement

Freezing/Condensing (exothermic): Energy released → intermolecular forces form → particles become more ordered

Exception: Conductivity in Solution

Some simple covalent compounds do conduct electricity in solution, but only if they react with water to form ions. For example, HCl dissolves in water to form H⁺ and Cl⁻ ions, which can then carry charge.

Comparing All Structure Types

PropertyGiant IonicGiant MetallicGiant CovalentSimple Molecular

Melting PointHighVariable (often high)Very highLow

Conductivity (solid)NoYesNo (except graphite)No

Conductivity (liquid)YesYesNoNo

HardnessHard, brittleVariable, malleableVery hard (diamond) or soft (graphite)Soft

SolubilitySoluble (forms ion-dipole bonds)InsolubleInsolubleUsually insoluble (unless polar)

Forces to OvercomeElectrostatic attraction between ionsSea of electrons attracting positive ionsStrong covalent bondsWeak intermolecular forces

ExampleNaClCu, MgDiamond, graphite, SiO₂I₂, ice, C₆₀

Exam Strategy

When given an unknown substance, use its properties to deduce its structure type. High MP + conducts when molten = ionic. Low MP + doesn't conduct = simple molecular.

Buckminsterfullerene (C₆₀)

C₆₀ is a simple molecular allotrope of carbon — 60 carbon atoms arranged in a spherical cage. Despite being made of carbon (like diamond and graphite), it has a low melting point because only weak van der Waals forces hold the C₆₀ molecules together.

Practice

Question 1: Silicon dioxide (SiO₂) has a melting point of 1710°C and does not conduct electricity. Suggest and explain the structure of silicon dioxide.

1

Structure type: Giant covalent (macromolecular)

2

High melting point: Many strong covalent bonds must be broken, requiring a large amount of energy.

3

Does not conduct: All outer electrons are involved in covalent bonding — no delocalised electrons or mobile ions to carry charge.

Each Si is bonded to 4 oxygen atoms in a tetrahedral arrangement (similar to diamond).

Practice

Question 2: Explain why copper is a good electrical conductor but copper(II) chloride only conducts when molten or dissolved.

1

Copper (metallic): Has delocalised electrons that are free to move through the structure and carry charge — conducts in solid and liquid states.

2

Copper(II) chloride (ionic): In solid state, ions are held in fixed positions in the lattice and cannot move.

3

When molten or dissolved, the ionic lattice breaks down and ions become free to move and carry charge.

Practice

Question 3: Substance X has a melting point of 54°C and does not conduct electricity in any state. Identify the structure type and explain your reasoning.

1

Structure type: Simple molecular

2

Evidence for low MP: Only weak intermolecular forces between molecules need to be overcome — requires little energy.

3

Evidence for non-conductivity: No free electrons or ions. Molecules are neutral, so there are no mobile charged particles to carry current in any state.

4

Rules out: Ionic (would conduct when molten), metallic (would conduct always), giant covalent (very high MP).

Summary

Diamond vs Graphite Comparison

AspectDiamondGraphite

Bonding4 strong 3D covalent bonds per carbon in tetrahedral arrangement3 strong planar covalent bonds + 1 delocalised electron per carbon; weaker interlayer forces

Structure3D network of tetrahedrally bonded carbon atomsStacked 2D hexagonal carbon sheets

PropertiesExtremely hard / Very high MP / Good thermal conductor / Electrical insulator / InsolubleSoft, layers slide / Conducts electricity in layers / Lower density / High sublimation temp

Exam Essentials

Always state what forces/bonds are broken when explaining melting points

Conductivity requires mobile charged particles (ions or delocalised electrons)

Know the specific examples: NaCl, Mg, diamond, graphite, ice, iodine

Graphite conducts due to delocalised electrons (1 per carbon)