Understand Models to Materials — Understand Structure...

SECTION 1

Historical Context & Motivation

For centuries, people used materials—bronze, glass, steel—without truly understanding why they behaved the way they did. A bronze sword is rigid and strong, while wax is soft and melts easily. Salt dissolves in water, but diamond does not. The question of why materials possess specific properties required scientists to peer beneath the surface and build bonding models that connect the invisible world of atoms to the tangible world of materials.

1916

Lewis Dot Structures

Gilbert N. Lewis proposed that atoms share or transfer electrons to achieve stable octets. His electron-dot diagrams became the first widely used model for covalent and ionic bonding.

1927

Quantum Mechanical Model

Heitler and London applied quantum mechanics to the hydrogen molecule, providing a mathematical foundation for understanding covalent bond formation through orbital overlap.

1950s

Band Theory & Materials Science

Scientists developed band theory to explain why metals conduct electricity, semiconductors sometimes conduct, and insulators never do—all from bonding models.

1985

Discovery of Fullerenes

The discovery of C₆₀ (buckminsterfullerene) showed that bonding models could predict entirely new allotropes of carbon with unique properties.

2004

Graphene Isolated

Geim and Novoselov isolated graphene—a single layer of graphite. Its extraordinary strength and conductivity were predicted by bonding models before experimental confirmation.

The central question this lesson addresses is: How can we use models of bonding and structure to predict and explain the physical properties of real materials? By mastering the connection between four types of bonding—ionic, covalent, metallic, and intermolecular—and the structures they produce, you gain the power to explain everything from why ice floats to why copper wires carry electricity.

SECTION 2

Core Principles — From Bonding to Properties

Structure 2.4 in the IB Chemistry syllabus focuses on four major categories of bonding and the structures they create. Each bonding type produces a characteristic structural arrangement, and each structural arrangement generates a predictable set of macroscopic properties. Understanding this chain—bonding → structure → properties—is the key to the entire topic.

Ionic Compounds

Electrostatic attractions between oppositely charged ions form a giant ionic lattice. These compounds have high melting points, are brittle, and conduct electricity only when molten or dissolved.

Covalent Network Structures

Atoms share electrons in a continuous three-dimensional covalent network. Examples include diamond (C) and silicon dioxide (SiO₂). They are extremely hard with very high melting points.

Metallic Structures

A sea of delocalized electrons surrounds positive metal cations in a lattice. This explains metals' electrical conductivity, malleability, and high thermal conductivity.

Molecular Covalent Substances

Discrete molecules held together by weak intermolecular forces (London dispersion, dipole–dipole, or hydrogen bonding). These substances have low melting and boiling points and generally do not conduct electricity.

The Bridge: Models to Properties

The type and strength of bonding in a substance directly determines whether it is hard or soft, conducts or insulates, melts at 100 °C or 3000 °C. Recognizing the bonding model lets you predict the property.

✦ KEY TAKEAWAY

Think of bonding models like architectural blueprints. Just as a blueprint tells you whether a building will be a skyscraper or a cottage—and whether it will withstand an earthquake or collapse in a breeze—a bonding model tells you what properties a material will have. Strong, extended bonding = high melting point and hardness. Weak intermolecular forces = low melting point and softness.

SECTION 3

Visual Explanation — The Four Structural Types

The four structural types: ionic lattices (top-left), covalent networks (top-right), metallic lattices (bottom-left), and molecular structures (bottom-right). Notice how bonding extends in all directions for ionic, network, and metallic types, but molecular structures consist of discrete units held together by weak dashed-line intermolecular forces.

In the diagram above, pay close attention to the difference between intra-molecular bonds (the solid lines connecting atoms within a molecule) and inter-molecular forces (the dashed lines between molecules). When a molecular substance melts, it is the intermolecular forces that break, not the covalent bonds. This is why molecular substances have relatively low melting points compared to ionic or network covalent substances, where breaking the structure means breaking strong bonds throughout the entire lattice.

SECTION 4

How Bonding Type Determines Properties

Melting and Boiling Points

The melting point of a substance tells you how much energy is needed to break the forces holding the solid together. In an ionic compound like sodium chloride (NaCl), you must overcome strong electrostatic attractions between every ion and all of its neighbors in three dimensions. This requires a great deal of energy, so NaCl melts at 801 °C. In contrast, a molecular substance like iodine (I₂) is held together by weak London dispersion forces, so it sublimes at only 184 °C.

LATTICE ENERGY (IONIC)

E ∝ (q⁺ × q⁻) / r

Where q⁺ and q⁻ are the charges on the ions and r is the distance between ion centers. Higher charges and smaller ions produce stronger lattice energies and higher melting points.

Electrical Conductivity

Electrical conductivity requires mobile charged particles. Metals conduct in the solid state because their delocalized electrons can flow freely. Ionic compounds only conduct when molten or dissolved in water, because the ions must be free to move. Molecular and covalent network substances (with exceptions like graphite) do not conduct electricity because they lack mobile charges.

Solubility

The general rule is "like dissolves like". Polar substances and ionic compounds tend to dissolve in polar solvents like water. Non-polar molecular substances dissolve in non-polar solvents like hexane. When NaCl dissolves, water molecules surround each ion in a process called hydration, with the energy released by ion–dipole interactions compensating for the lattice energy that must be overcome.

Hardness and Brittleness

Ionic crystals are hard but brittle: when struck, layers of ions shift, placing like charges next to each other, and the resulting repulsion shatters the crystal. Metals are malleable because their layers of cations can slide past one another while the sea of electrons continuously adjusts to maintain bonding.

SECTION 5

Allotropes of Carbon — Same Element, Different Properties

Carbon provides a fascinating case study for connecting bonding models to material properties. The same element can form dramatically different allotropes—forms of the same element with different structural arrangements. Diamond, graphite, graphene, and fullerenes are all pure carbon, yet their properties could not be more different. The difference lies entirely in how the carbon atoms are bonded.

Three allotropes of carbon compared: diamond has a tetrahedral covalent network (sp³), graphite has planar layers with delocalized π electrons (sp²), and fullerene C₆₀ forms discrete spherical molecules held together by weak intermolecular forces. The dashed lines between graphite layers represent weak van der Waals forces.

Diamond is the hardest naturally occurring substance because every carbon atom is bonded to four neighbors in a rigid three-dimensional covalent network. To break diamond, you have to break covalent bonds—an enormous energy cost. Graphite, however, consists of flat sheets of carbon atoms arranged in hexagonal rings. Within each layer, strong covalent bonds and delocalized π electrons make the sheets incredibly strong, but the layers are held together by only weak London dispersion forces. This allows layers to slide over one another, making graphite soft and slippery—perfect for pencil lead and lubricants.

Graphite is unusual among covalent network structures because it conducts electricity. Each carbon uses only three of its four valence electrons for σ bonds, leaving one electron per atom delocalized across the π system of each layer. These mobile electrons can carry current along the plane of the layers.

💡 IB Exam Tip

You should be able to explain why graphite conducts electricity but diamond does not, even though both are covalent network structures of carbon. The key difference is the delocalized π electrons present in graphite's sp² layers that are absent in diamond's sp³ structure.

SECTION 6

Worked Example — Predicting Properties from Structure

Let's walk through a typical IB-style question that asks you to predict and explain the properties of a substance based on its bonding model.

Comparing MgO and CO₂

Step 1 — Identify the bonding type of each substanceMagnesium oxide (MgO) is formed between a metal (Mg) and a non-metal (O). Magnesium transfers two electrons to oxygen, creating Mg²⁺ and O²⁻ ions. This is ionic bonding, and MgO forms a giant ionic lattice. Carbon dioxide (CO₂) is composed entirely of non-metals. Carbon shares electrons with oxygen through double covalent bonds. CO₂ exists as discrete molecules—a molecular covalent structure.

MgO = ionic lattice; CO₂ = molecular covalent

Step 2 — Predict melting pointsMgO has very high charges (2+ and 2−) and small ionic radii, so the electrostatic attraction between ions is extremely strong. Its melting point is 2852 °C. In CO₂, the molecules are linear and non-polar, so they are held together only by weak London dispersion forces. CO₂ sublimes at −78.5 °C.

MgO m.p. = 2852 °C >> CO₂ sublimation point = −78.5 °C

Step 3 — Predict electrical conductivityMgO does not conduct electricity as a solid because the ions are fixed in the lattice and cannot move. However, when molten or dissolved in water, the ions become mobile and can carry charge—so it conducts. CO₂ has no charged particles that can move, so it does not conduct electricity in any state.

MgO: conducts when molten/aqueous. CO₂: non-conductor in all states.

Step 4 — Predict solubilityMgO is an ionic compound and dissolves (slightly) in water because the polar water molecules can hydrate the ions, though the very high lattice energy of MgO limits its solubility. CO₂ is non-polar, so it has limited solubility in water (though it does react with water to form carbonic acid, H₂CO₃). CO₂ is more soluble in non-polar organic solvents.

MgO: slightly soluble in water (ionic). CO₂: slightly soluble in water (reacts), more soluble in non-polar solvents.

SECTION 7

Comparing the Four Structural Types

The table below summarizes the key properties associated with each structural type. Being able to reproduce and explain this table is essential for success on the IB Chemistry exam.

Summary of properties for the four main structural types in IB Chemistry

Property	Ionic	Metallic	Covalent Network	Molecular
Melting Point	High	Variable (generally high)	Very high	Low
Hardness	Hard but brittle	Variable; malleable	Very hard	Soft
Conductivity (solid)	No	Yes	No (except graphite)	No
Conductivity (molten)	Yes	Yes	No	No
Solubility in water	Often soluble	Insoluble	Insoluble	Polar: yes; Non-polar: no
Volatility	Low	Low	Very low	High
Particles present	Cations & anions	Cations & delocalized e⁻	Atoms (covalently bonded)	Molecules

✦ KEY TAKEAWAY

Think of melting a substance as "breaking apart the structure." For ionic, metallic, and covalent network substances, this means breaking strong bonds throughout an extended lattice—so melting points are high. For molecular substances, only the weak intermolecular forces between discrete molecules need to be overcome, so the covalent bonds within molecules remain intact and the melting points are low.

SECTION 8

Connecting to Advanced Theory — Beyond the Four Types

The four structural types you have learned provide an excellent framework for predicting properties, but reality is sometimes more nuanced. Some substances blur the boundaries between categories, and advanced chemistry introduces more sophisticated models to explain these edge cases.

Connections between Structure 2.4 concepts and higher-level theory

Concept at This Level	Advanced Extension
Ionic vs. covalent bonding as distinct categories	Bonding is a spectrum: Fajans' rules predict when ionic bonds have significant covalent character based on ion charge density and polarizability
Graphite conducts electricity due to delocalized electrons	Band theory explains conductivity through valence bands, conduction bands, and band gaps — graphite has overlapping bands
Lattice energy proportional to q⁺ × q⁻ / r	The Born–Landé equation quantifies lattice energy including the Madelung constant and the Born exponent
Molecular substances have low melting points	Polymers (very large molecules) can have properties similar to covalent networks due to chain entanglement and cross-linking

An interesting example is silicon carbide (SiC), which has a covalent network structure similar to diamond. It is one of the hardest known materials, with a melting point of approximately 2730 °C. In advanced materials science, engineers exploit the bonding model → property prediction chain to design new materials: need something extremely hard? Build a covalent network structure. Need something that conducts? Ensure mobile charged particles are present. This design principle is the essence of modern materials science.

SECTION 9

Practice Problems

PROBLEM 1 — CONCEPTUAL

Explain why sodium chloride (NaCl) is brittle, while copper metal is malleable. Your answer should reference the type of bonding and the arrangement of particles in each substance.

PROBLEM 2 — BASIC CALCULATION

Two ionic compounds are compared: NaF (Na⁺ and F⁻) and MgO (Mg²⁺ and O²⁻). Both have similar ionic radii. Using the relationship E ∝ (q⁺ × q⁻) / r, predict which compound has the higher melting point and explain your reasoning.

PROBLEM 3 — INTERMEDIATE

Silicon dioxide (SiO₂) and carbon dioxide (CO₂) both contain the same elements combined with oxygen, yet SiO₂ melts at 1713 °C while CO₂ sublimes at −78.5 °C. Explain this enormous difference in terms of structure and bonding.

PROBLEM 4 — APPLIED

A materials engineer needs to select a substance for use as an electrode. The material must conduct electricity in the solid state and withstand temperatures above 2000 °C. Which of the four structural types (ionic, metallic, covalent network, or molecular) would be most suitable? Suggest a specific substance and justify your choice.

PROBLEM 5 — CRITICAL THINKING

Ice (solid H₂O) has a melting point of 0 °C, while solid hydrogen sulfide (H₂S) melts at −85 °C. Both are molecular substances with similar molar masses. However, ice is less dense than liquid water, which is highly unusual. Explain both observations — the difference in melting points and the anomalous density of ice — in terms of intermolecular forces and structure.

SUMMARY

Lesson Summary

Structure 2.4 teaches us to bridge the gap between bonding models and material properties. The four main structural types are ionic lattices (high melting points, brittle, conduct when molten), metallic structures (conductive, malleable, lustrous), covalent network structures (extremely hard, very high melting points, usually non-conductive), and molecular covalent substances (low melting points, soft, non-conductive). The key to predicting properties is identifying the type and strength of the forces that must be broken when a substance melts, dissolves, or conducts.

Remember that allotropes of carbon demonstrate how the same element can have wildly different properties depending on its bonding arrangement: diamond (sp³, network, hardest material) versus graphite (sp², layered, conducts electricity) versus fullerenes (molecular, low melting point). For the IB exam, always use the chain: identify bonding type → determine structure → predict properties. Practice applying this reasoning to unfamiliar substances, and you will be well prepared.