Historical Context & Motivation
For thousands of years, people knew that certain substances like table salt dissolved easily in water, formed hard crystals, and could conduct electricity when melted. However, no one could explain why these substances behaved so differently from metals or waxes. The breakthrough came when scientists realized that atoms could gain or lose electrons to form charged particles called ions, and that the electrostatic attraction between oppositely charged ions could explain a whole family of compounds. This idea became the foundation of the ionic model — a powerful framework that lets us predict melting points, solubility, electrical conductivity, and more.
These discoveries raised a central question in chemistry: how can a simple model of positive and negative ions arranged in a lattice explain so many observable properties? The answer lies in understanding the electrostatic forces between ions and the way those forces shape the macroscopic behavior of ionic compounds. In this lesson, you will learn to apply the ionic model to predict properties, explain experimental observations, and solve quantitative problems — exactly the skills tested in IB Chemistry.
Core Principles of the Ionic Model
The ionic model is built on a few foundational ideas that connect atomic structure to the observable properties of ionic compounds. When a metal atom loses one or more electrons, it becomes a positively charged cation. When a non-metal atom gains one or more electrons, it becomes a negatively charged anion. These oppositely charged ions attract each other through electrostatic attraction, forming a three-dimensional arrangement called an ionic lattice. The strength of this attraction determines many of the compound's physical properties.
Electron Transfer
Electrostatic Attraction
Ionic Lattice Structure
Lattice Energy
Charge Density
Visualising the Ionic Lattice
Seeing the arrangement of ions in a crystal lattice is one of the best ways to understand why ionic compounds behave the way they do. The diagram below shows a simplified 2D cross-section of a sodium chloride (NaCl) lattice, where each Na⁺ cation is surrounded by Cl⁻ anions and vice versa. Notice how no single pair of ions "belongs" to each other — the bonding is non-directional and extends in all directions throughout the crystal.
In the diagram above, there are several key features to notice. First, the Na⁺ ions (cyan) are drawn smaller than the Cl⁻ ions (pink), reflecting the real difference in ionic radii — sodium loses an electron and its electron cloud shrinks, while chlorine gains an electron and its electron cloud expands. Second, the alternating pattern means every cation is surrounded by anions and every anion is surrounded by cations. This arrangement maximises the attractive forces while keeping like-charged ions as far apart as possible. Third, there are no distinct "NaCl molecules" — just an extended three-dimensional lattice, which is why we write ionic compounds as empirical formulas (NaCl) rather than molecular formulas.
Mathematical Framework — Coulomb's Law and Lattice Energy
The ionic model's predictive power comes from the physics of electrostatic attraction. Coulomb's law tells us that the force between two charged particles depends on the magnitude of their charges and the distance between them. When applied to ionic compounds, it allows us to predict which compounds will have higher melting points, greater lattice energies, and stronger bonds.
From this equation, two key relationships emerge. First, the force increases when ionic charges are larger — a compound containing Mg²⁺ and O²⁻ (charges of 2+ and 2−) will have much stronger attractions than NaCl (charges of 1+ and 1−). Second, the force increases when ionic radii are smaller, because the ions can get closer together. These two factors together determine the lattice energy of the compound.
Properties of Ionic Compounds Explained by the Model
One of the most useful aspects of the ionic model is that it lets us explain — and predict — the physical properties of ionic compounds from first principles. The diagram below summarises how the ionic lattice determines behaviour under different conditions.
| Property | Observation | Explanation Using the Ionic Model |
|---|---|---|
| Melting point | Generally high (typically > 300 °C) | A large amount of energy is needed to overcome the strong electrostatic attractions between many ions throughout the lattice. |
| Conductivity (solid) | Does not conduct | Ions are fixed in position in the lattice and cannot move to carry charge. |
| Conductivity (molten/aq) | Conducts electricity | When melted or dissolved, ions are free to move and carry charge through the liquid. |
| Solubility | Many dissolve in water | Polar water molecules surround individual ions (hydration), stabilising them and overcoming lattice forces. |
| Brittleness | Shatters when struck | A mechanical force can shift ion layers so that like-charged ions align, producing repulsion and fracture. |
Worked Example — Comparing Melting Points
A common IB Chemistry question asks you to compare the melting points of two ionic compounds and explain the difference using the ionic model. Let's work through a full example step by step.
Strengths and Limitations of the Ionic Model
Like all models in science, the ionic model is a simplification of reality. It works extremely well for many purposes, but it has boundaries where it breaks down. Understanding these boundaries is essential for the IB — examiners frequently ask you to evaluate a model's limitations.
| Strengths | Limitations |
|---|---|
| Accurately predicts high melting points and boiling points for compounds formed between metals and non-metals. | Assumes complete electron transfer — in reality, many ionic bonds have some degree of covalent character (electron sharing), especially when cations are small and highly charged. |
| Correctly explains why ionic compounds conduct electricity when molten or dissolved but not as solids. | Cannot explain why some ionic compounds are coloured (e.g., CuSO₄ is blue) — this requires a more advanced model (crystal field theory). |
| Explains brittleness and hardness of ionic crystals using the lattice structure. | Treats ions as hard spheres with fixed charges — does not account for polarisation or distortion of electron clouds. |
| Allows quantitative comparison of lattice energies using Coulomb's law. | The electronegativity difference between bonding atoms is a continuum — there is no sharp boundary between ionic and covalent bonding. |
Connection to Advanced Theory — From Ionic to Covalent Character
In the IB programme, you will encounter the idea that bonding is a spectrum, not a set of discrete categories. At one end, you have pure ionic bonding (complete electron transfer), and at the other, pure covalent bonding (equal electron sharing). Most real compounds fall somewhere in between. A compound like NaCl is well-described by the ionic model because sodium and chlorine have a large electronegativity difference (≈ 2.1). However, a compound like AlCl₃, which also contains a metal and a non-metal, shows significant covalent character because the small, highly charged Al³⁺ ion strongly polarises the electron cloud of Cl⁻.
| Feature | Ionic Model (Structure 2.1) | Advanced: Polarisation / Fajans' Rules |
|---|---|---|
| Electron transfer | Assumed to be complete — cation and anion have full charges. | Recognised as partial — small, highly charged cations pull anion electron density back toward themselves. |
| Ion shape | Treated as perfect, non-deformable spheres. | Electron clouds can be distorted (polarised), especially for large anions paired with small cations. |
| Bonding type | Binary classification: ionic or covalent. | Spectrum of bonding: degree of ionic vs. covalent character varies continuously. |
| When it works best | Compounds with large electronegativity differences (> 1.7) between Group 1/2 metals and Group 16/17 non-metals. | Required when ionic model predictions deviate from experimental values (e.g., AlCl₃ has lower melting point than expected). |
For now, the key point is this: the ionic model is the starting point for understanding compounds formed between metals and non-metals. It provides a robust framework for predicting and explaining properties. As you advance in chemistry, you will learn to refine this model by considering polarisation and covalent character — but those refinements build on, rather than replace, the ionic model you are learning here.
Practice Problems
Lesson Summary
The ionic model describes compounds formed by electron transfer from metals to non-metals, producing positively charged cations and negatively charged anions held together by electrostatic attraction in a three-dimensional ionic lattice. The strength of these attractions, quantified by lattice energy, depends on ionic charge (higher charge → stronger attraction) and ionic radius (smaller radius → stronger attraction), as described by Coulomb's law.
This model successfully explains the high melting points, electrical conductivity when molten or dissolved, solubility in water, brittleness, and hardness of ionic compounds. However, it has limitations — it assumes complete electron transfer and treats ions as hard spheres, which breaks down when small, highly charged cations polarise anion electron clouds, introducing covalent character. In IB Chemistry, you must be able to apply this model to compare compounds, explain experimental data, and recognise when its assumptions fail.