Structure and Properties of Materials

students, in this lesson you will explore how the arrangement of particles inside a material controls what that material can do ✨. The same idea helps explain why salt dissolves in water, why metals bend instead of shatter, and why diamond is so hard. By the end of this lesson, you should be able to explain structure-property relationships, compare the main types of bonding, and connect these ideas to real-world materials used in everyday life.

Objectives

Explain how structure affects the properties of materials.
Compare ionic, covalent, metallic, and intermolecular forces.
Use particle-level evidence to predict properties such as melting point, conductivity, hardness, and solubility.
Relate materials science to the IB Chemistry SL topic Structure 2 — Models of Bonding and Structure.

A key idea in chemistry is that structure determines properties. A material is not just “made of stuff”; it has an arrangement of atoms, ions, molecules, or electrons, and that arrangement explains its behavior. For example, a copper wire conducts electricity because of mobile electrons, while plastic does not because its electrons are held tightly in bonds.

1. What do we mean by structure and properties?

In chemistry, structure means how particles are arranged and what forces hold them together. Properties are the observable characteristics of a material, such as melting point, boiling point, hardness, conductivity, flexibility, and solubility.

A useful IB Chemistry idea is that every material can be understood by asking two questions:

What particles are present?
What forces or bonds hold those particles together?

The answer to these questions lets us predict behavior. For example, if a substance has strong attractions between particles, a lot of energy is needed to separate them, so it usually has a high melting point. If particles can move freely, the material may conduct electricity.

Real-world example 🌍: table salt, $\text{NaCl}$, and sugar, $\text{C}_{12}\text{H}_{22}\text{O}_{11}$, both dissolve in water, but they behave differently in solution. Salt solution conducts electricity because it contains ions, while sugar solution does not because sugar molecules remain neutral.

2. Ionic structures: strong attractions in a crystal lattice

Ionic compounds are made of positive and negative ions arranged in a giant repeating structure called a lattice. In $\text{NaCl}$, each $\text{Na}^+$ ion is surrounded by $\text{Cl}^-$ ions and vice versa. The electrostatic attraction between oppositely charged ions is strong and acts in all directions.

Because of these strong attractions, ionic compounds often have:

high melting and boiling points,
hard but brittle crystals,
electrical conductivity when molten or dissolved in water,
poor conductivity as solids.

Why are they brittle? If layers in the lattice shift, ions with the same charge can end up next to each other. Like charges repel, so the crystal can crack.

Example: magnesium oxide, $\text{MgO}$, has a very high melting point because the ions have $2+$ and $2-$ charges, creating stronger attraction than in $\text{NaCl}$. This is a good example of how charge affects structure-property relationships.

In water, many ionic compounds dissolve because polar water molecules can surround and separate ions. This is called hydration. The ability to dissolve depends on the balance between lattice attraction and attraction to water molecules.

3. Covalent structures: molecules and giant covalent networks

Covalent bonding happens when atoms share pairs of electrons. Covalent substances can be divided into two major types: simple molecular substances and giant covalent structures.

Simple molecular substances

These are made of small molecules such as $\text{H}_2\text{O}$, $\text{CO}_2$, $\text{CH}_4$, and $\text{I}_2$. The atoms within each molecule are held by strong covalent bonds, but the forces between molecules are much weaker and are called intermolecular forces.

Because intermolecular forces are relatively weak, simple molecular substances usually have:

low melting and boiling points,
low hardness,
no electrical conductivity,
high volatility for some substances.

For example, carbon dioxide is a gas at room temperature because only weak intermolecular forces hold its molecules together. In contrast, water has a higher boiling point than expected because hydrogen bonding is a strong intermolecular force.

Giant covalent structures

These materials have atoms joined in a large network of covalent bonds. Examples include diamond, graphite, and silicon dioxide, $\text{SiO}_2$.

Diamond: each carbon atom forms four covalent bonds in a rigid 3D structure. This makes diamond extremely hard and gives it a very high melting point.
Graphite: each carbon atom forms three covalent bonds in layers. One electron per carbon is delocalized, so graphite conducts electricity along the layers. The layers are held together by weak forces, so they can slide over each other, making graphite soft and slippery.
Silicon dioxide: a giant covalent network with strong bonds throughout the structure, giving it a high melting point and hardness.

These examples show that even substances made from the same element, like carbon, can have very different properties because their structures differ.

4. Metallic bonding: positive ions in a sea of electrons

Metals have a structure of positive ions surrounded by delocalized electrons. Metallic bonding is the attraction between the positive metal ions and these mobile electrons.

This model explains many metal properties:

Electrical conductivity: electrons can move and carry charge.
Thermal conductivity: electrons transfer energy quickly.
Malleability and ductility: layers of ions can slide past each other while the electron sea keeps the structure together.
Lustrous appearance: free electrons interact with light, giving metals their shine.

For example, copper is used in wiring because it conducts electricity well, and aluminum is used in aircraft because it is light but still strong enough for structural use.

Alloys are important real-world materials. An alloy is a mixture of metals, or a metal with another element. The different-sized atoms in an alloy make it harder for layers to slide, so alloys are often harder and stronger than pure metals. Steel is a common example.

5. Intermolecular forces and their effect on properties

Intermolecular forces are attractions between molecules, not within them. They are weaker than covalent or ionic bonds, but they strongly influence physical properties.

The main intermolecular forces you should know are:

London dispersion forces: present in all molecules; caused by temporary dipoles.
Permanent dipole-dipole forces: present between polar molecules.
Hydrogen bonding: a particularly strong dipole-dipole interaction when hydrogen is bonded to $\text{N}$, $\text{O}$, or $\text{F}$.

Stronger intermolecular forces usually mean:

higher melting and boiling points,
lower volatility,
higher viscosity,
greater surface tension.

Example: ethanol, $\text{C}_2\text{H}_5\text{OH}$, boils at a higher temperature than ethane, $\text{C}_2\text{H}_6$, because ethanol can form hydrogen bonds while ethane only has dispersion forces.

This matters in everyday life. Water’s unusually high boiling point helps life on Earth exist in liquid form. Hydrogen bonding also explains why water has high surface tension, allowing small insects to walk on it 💧.

6. Materials, models, and evidence

Chemistry uses models to represent tiny particles we cannot see directly. A model is useful if it explains observations and helps make predictions. However, models are simplified and do not show every detail.

In Structure 2, the bonding model helps explain evidence from experiments:

high melting point suggests strong attractions,
conductivity suggests mobile charged particles or electrons,
brittleness suggests a structure that breaks when layers shift,
solubility gives clues about particle interactions with solvents.

For instance, if a solid does not conduct electricity, that may mean it lacks mobile charges. But there are exceptions and more than one possible structure, so chemists use several pieces of evidence together.

This is one reason materials science is important. Engineers choose materials based on their structure and properties. For example:

ceramics are hard and heat-resistant but often brittle,
polymers are light and flexible but may soften at lower temperatures,
metals are strong and conductive,
composite materials combine useful properties from more than one material.

A smartphone contains many materials chosen for specific properties: glass for transparency and hardness, metal for conductivity and strength, and polymers for insulation and flexibility.

Conclusion

students, the central idea of this lesson is simple but powerful: the structure of a material explains its properties. Ionic solids have strong electrostatic attractions, simple molecular substances depend on intermolecular forces, giant covalent structures are rigid and often very hard, and metals contain mobile electrons that explain conductivity and malleability. By using models of bonding and structure, you can predict how a material will behave and connect chemistry to real-world applications. This is exactly the kind of reasoning expected in IB Chemistry SL 🔬.

Study Notes

Structure means how particles are arranged and what forces hold them together.
Properties are observable features such as melting point, hardness, conductivity, and solubility.
Ionic compounds form giant lattices of oppositely charged ions and usually have high melting points.
Ionic solids do not conduct electricity, but molten ionic compounds and aqueous ionic solutions do.
Simple molecular substances have strong covalent bonds within molecules but weak intermolecular forces between molecules.
Stronger intermolecular forces usually lead to higher boiling points and lower volatility.
Giant covalent structures like diamond and $\text{SiO}_2$ have very high melting points because of many covalent bonds.
Graphite conducts electricity because of delocalized electrons.
Metals conduct electricity and heat because of delocalized electrons.
Alloys are usually harder than pure metals because different-sized atoms hinder layer movement.
Models help explain data, but they are simplified representations.
In IB Chemistry SL, always connect a structure to an observed property using particle-level reasoning.