Apply The Ionic Model — Apply Structure 2.1—The ionic model in problem-solving and explanations - IB Chemistry

SECTION 1

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.

1807 • Humphry Davy's Electrolysis

Davy used electrolysis to isolate potassium and sodium from their molten compounds, demonstrating that these substances contained electrically attracted components.

1884 • Arrhenius & Ionic Dissociation

Svante Arrhenius proposed that salts dissociate into ions when dissolved in water, explaining why salt solutions conduct electricity. His theory was initially controversial but later earned a Nobel Prize.

1916 • Kossel's Ionic Bond Theory

Walther Kossel described ionic bonding as the complete transfer of electrons from a metal to a non-metal, with the resulting ions held together by electrostatic forces.

1918 • Born–Landé Equation

Max Born and Alfred Landé developed a mathematical equation to calculate the lattice energy of ionic crystals, providing a quantitative foundation for the ionic model.

1920s–Today • Modern Applications

The ionic model remains central to chemistry education and industry. It is used to design ceramics, batteries, water treatment systems, and countless other technologies.

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.

SECTION 2

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

Metals lose electrons to achieve a stable noble gas electron configuration, forming cations. Non-metals gain electrons to complete their valence shell, forming anions.

Electrostatic Attraction

Oppositely charged ions attract each other strongly. The force depends on the magnitude of the charges and the distance between the ion centres — smaller, more highly charged ions attract more strongly.

Ionic Lattice Structure

Ions are not paired in isolated molecules. Instead, each ion is surrounded by ions of opposite charge in a repeating 3D lattice, maximising attractive forces and minimising repulsive ones.

Lattice Energy

The energy released when gaseous ions come together to form one mole of ionic lattice. Higher lattice energy means stronger ionic bonds and generally higher melting points.

Charge Density

Charge density is the ratio of an ion's charge to its size. Ions with high charge density (small and/or highly charged) create stronger attractions, leading to higher melting points and greater lattice energies.

✦

KEY TAKEAWAY

Think of an ionic lattice like a tightly packed stadium crowd where every person wearing a red shirt is surrounded by people in blue shirts, and vice versa. Each person is "held in place" by the people of the opposite colour around them — not just one partner. That is why ionic compounds form extended structures rather than discrete molecules, and why it takes a lot of energy (a high melting point) to pull the whole structure apart.

SECTION 3

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.

A 2D cross-section of the NaCl lattice. The smaller cyan circles represent Na⁺ cations and the larger pink circles represent Cl⁻ anions. Notice how each ion is surrounded by ions of the opposite charge, creating a continuous, non-directional bonding network.

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.

SECTION 4

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.

COULOMB'S LAW (ELECTROSTATIC FORCE)

F = k × (q⁺ × q⁻) / r²

F = electrostatic force between two ions; k = Coulomb's constant (8.99 × 10⁹ N·m²·C⁻²); q⁺ and q⁻ = charges on the cation and anion; r = distance between the centres of the two ions (sum of ionic radii).

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.

LATTICE ENERGY TREND

Lattice energy ∝ (q⁺ × q⁻) / (r⁺ + r⁻)

Lattice energy is directly proportional to the product of the ionic charges and inversely proportional to the sum of the ionic radii. Higher lattice energy means a more stable lattice and typically a higher melting point.

💡

IB Exam Tip

You are not expected to memorise the Born–Landé equation for IB Chemistry. Instead, you must be able to use the proportionality relationship above to compare lattice energies and predict which ionic compound will have a higher melting point. Always compare both charge and size — sometimes one factor dominates, and you need to identify which.

PREDICTING RELATIVE MELTING POINTS

Higher lattice energy → Higher melting point

Example: MgO (Mg²⁺ and O²⁻, both small ions with 2+ and 2− charges) has a much higher melting point (2852 °C) than NaCl (Na⁺ and Cl⁻, 1+ and 1− charges, melting point 801 °C). The charges in MgO are doubled and the ions are smaller, so the lattice energy is significantly greater.

SECTION 5

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.

Five key properties of ionic compounds, each explained using the ionic model. The top row covers melting point, electrical conductivity, and solubility. The bottom row covers brittleness and hardness.

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.

Summary of ionic compound properties explained by the ionic model

SECTION 6

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.

Compare the melting points of NaCl (801 °C) and MgO (2852 °C). Explain the difference using the ionic model.

1
Step 1 — Identify the Ions and Their Charges

NaCl contains Na⁺ (1+) and Cl⁻ (1−). MgO contains Mg²⁺ (2+) and O²⁻ (2−). The charges in MgO are double those in NaCl.

NaCl: 1+/1−; MgO: 2+/2−
2
Step 2 — Compare the Ionic Radii

Na⁺ has an ionic radius of 102 pm, while Mg²⁺ has an ionic radius of 72 pm — Mg²⁺ is smaller because it has lost more electrons relative to its nuclear charge. Cl⁻ has an ionic radius of 181 pm, while O²⁻ has an ionic radius of 140 pm. Both ions in MgO are smaller than their NaCl counterparts.

Mg²⁺ (72 pm) < Na⁺ (102 pm); O²⁻ (140 pm) < Cl⁻ (181 pm)
3
Step 3 — Apply the Lattice Energy Proportionality

Lattice energy is proportional to (q⁺ × q⁻) / (r⁺ + r⁻). For NaCl: (1 × 1) / (102 + 181) = 1/283. For MgO: (2 × 2) / (72 + 140) = 4/212. The ratio for MgO is approximately 5.3 times larger than for NaCl, indicating a much higher lattice energy.

MgO lattice energy ≈ 5.3 × NaCl lattice energy
4
Step 4 — State the Conclusion

MgO has a much higher melting point than NaCl because its ions have higher charges (2+ and 2− vs. 1+ and 1−) and smaller radii. Both factors increase the electrostatic attraction between the ions, resulting in a stronger ionic lattice that requires more energy to break apart.

MgO (2852 °C) >> NaCl (801 °C) due to higher charges and smaller ions → stronger electrostatic attractions → greater lattice energy

💡

Exam Writing Strategy

In IB exam answers, always mention both ionic charge and ionic radius when comparing melting points. Then explicitly link these to 'electrostatic attraction' and 'lattice energy.' Use comparative language: 'greater charge,' 'smaller radius,' 'stronger attraction,' 'higher melting point.'

SECTION 7

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.

Strengths and limitations of the ionic model

✦

KEY TAKEAWAY

Think of the ionic model like a weather forecast. It is incredibly useful for day-to-day predictions — it will tell you whether to bring an umbrella and what temperature to expect. But it cannot predict every detail perfectly, such as the exact minute rain will start or the shape of each cloud. Similarly, the ionic model gives reliable predictions about melting points, conductivity, and solubility, but it cannot explain every nuance of every ionic compound. When the model falls short, chemists use more advanced models such as Fajans' rules and crystal field theory.

SECTION 8

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).

The ionic model vs. more advanced bonding theories

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.

SECTION 9

Practice Problems

Question 1: CONCEPTUAL

Sodium chloride (NaCl) does not conduct electricity as a solid but does conduct when dissolved in water. Using the ionic model, explain why.

Show answer

In solid NaCl, the Na⁺ and Cl⁻ ions are held in fixed positions within the ionic lattice by strong electrostatic attractions. Because the ions cannot move, there are no mobile charge carriers and the solid does not conduct electricity. When NaCl dissolves in water, polar water molecules surround the individual ions (hydration), breaking the lattice apart. The Na⁺ and Cl⁻ ions are now free to move throughout the solution, acting as mobile charge carriers that can carry electrical current. This behaviour is a direct prediction of the ionic model, which states that conduction requires the movement of charged particles.

Question 2: BASIC CALCULATION

Using the proportionality relationship for lattice energy (∝ q⁺ × q⁻ / (r⁺ + r⁻)), calculate the approximate ratio of lattice energies for LiF and NaCl. Use ionic radii: Li⁺ = 76 pm, F⁻ = 133 pm, Na⁺ = 102 pm, Cl⁻ = 181 pm. Both have 1+ and 1− charges.

Show answer

For LiF: (1 × 1) / (76 + 133) = 1/209. For NaCl: (1 × 1) / (102 + 181) = 1/283. The ratio of lattice energies is (1/209) ÷ (1/283) = 283/209 ≈ 1.35. This means LiF has an approximately 35% higher lattice energy than NaCl. This is consistent with the experimental observation that LiF has a higher melting point (845 °C) than NaCl (801 °C). Since both compounds have ions with the same charges, the difference arises entirely from the smaller ionic radii of Li⁺ and F⁻, which allow the ions to pack more closely and increase the electrostatic attraction.

Question 3: INTERMEDIATE

Arrange the following compounds in order of increasing melting point and justify your ranking: NaCl, MgO, KBr. Consider the ionic charges and approximate ionic radii (K⁺ = 138 pm, Na⁺ = 102 pm, Mg²⁺ = 72 pm, Br⁻ = 196 pm, Cl⁻ = 181 pm, O²⁻ = 140 pm).

Show answer

The order of increasing melting point is: KBr < NaCl < MgO. KBr has the lowest melting point because K⁺ (138 pm) and Br⁻ (196 pm) are both large ions, and both carry only single charges (1+/1−). The large distance between ion centres weakens the electrostatic attraction, giving a relatively low lattice energy. NaCl has a higher melting point than KBr because Na⁺ (102 pm) and Cl⁻ (181 pm) are smaller ions, so they pack more closely. However, the charges are still 1+ and 1−. MgO has the highest melting point by a large margin because Mg²⁺ (2+) and O²⁻ (2−) both carry double charges, and both ions are small. The product of charges (2 × 2 = 4) is four times that of NaCl (1 × 1 = 1), and the sum of radii is smaller. Together, these factors produce a much higher lattice energy and thus a much higher melting point (2852 °C vs. 801 °C for NaCl vs. 734 °C for KBr).

Question 4: APPLIED

A student designs an experiment to test whether a white crystalline solid is ionic. They test its melting point, its electrical conductivity as a solid, and its conductivity when dissolved in water. The solid has a melting point of 614 °C, does not conduct as a solid, and conducts when dissolved. Another unknown compound melts at 45 °C and does not conduct in any state. Using the ionic model, explain which compound is likely ionic and which is not.

Show answer

The first compound is most likely ionic. Its high melting point (614 °C) indicates strong electrostatic attractions between ions in a lattice, which require significant energy to overcome. Its failure to conduct as a solid is consistent with ions locked in fixed positions, while its conductivity in aqueous solution shows that mobile ions are present after dissolution — exactly what the ionic model predicts. The second compound is likely molecular (covalent). Its low melting point (45 °C) suggests weak intermolecular forces rather than strong ionic bonds, and its inability to conduct in any state means it contains no mobile ions or free electrons. The ionic model helps us distinguish between these compound types by linking observable properties (melting point, conductivity) to the underlying structure (lattice of ions vs. discrete molecules).

Question 5: CRITICAL THINKING

Aluminium chloride (AlCl₃) contains a metal and a non-metal, yet it has a melting point of only 192.4 °C — far lower than expected for a typical ionic compound. Using your understanding of the ionic model and its limitations, propose an explanation for this anomaly. What does this suggest about the bonding in AlCl₃?

Show answer

The ionic model predicts that AlCl₃ should have a high melting point because it is formed from a metal (Al) and a non-metal (Cl). If it were purely ionic, we would expect Al³⁺ (a small, highly charged cation) to form very strong electrostatic attractions, similar to MgO. However, the very high charge density of Al³⁺ means it strongly polarises the electron cloud of Cl⁻, pulling electron density back toward itself. This creates significant covalent character — the electrons are partially shared rather than fully transferred. As a result, AlCl₃ forms covalent molecules (actually Al₂Cl₆ dimers) rather than an extended ionic lattice. With only weak intermolecular forces between these molecules, the melting point is low. This is a limitation of the simple ionic model: it does not account for polarisation. This example demonstrates that bonding is a spectrum, and compounds near the boundary between ionic and covalent cannot be accurately described by the ionic model alone. More advanced concepts, such as Fajans' rules, are needed to predict when covalent character will dominate.

SUMMARY

Lesson Summary

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.