IB DP Chemistry Structure 2.1: The ionic model – Revision Notes

These notes condense Structure 2.1, The ionic model, for the final weeks before your exams. They follow the IB’s Diploma Programme Chemistry guide (February 2023 edition, first assessment 2025) and cover understandings Structure 2.1.1, 2.1.2 and 2.1.3. All of it is common to SL and HL, because this sub-topic has no AHL extension.

Need the full explanations and longer worked examples? Go to the Structure 2.1 study guide. Afterwards, try the practice set. Course context sits on the hub page and the course checklist. The free diagnostics show which topics need more time.

Key definitions

Term	Meaning
Cation	Positive ion, formed when an atom (usually a metal) loses electrons
Anion	Negative ion, formed when an atom (usually a non-metal) gains electrons
Ionic bond	Electrostatic attraction between oppositely charged ions
Polyatomic ion	A charged group of covalently bonded atoms that acts as one unit
Binary compound	A compound made of only two elements
Lattice	A regular three-dimensional array of ions, repeated in all directions
Empirical formula	The simplest whole-number ratio of ions in the lattice
Lattice enthalpy	Enthalpy change to separate one mole of solid ionic compound into gaseous ions; positive, and larger for stronger ionic bonding

2.1.1 Ion charges from electron configurations

• Count the electrons in the outer main level. Lose them (if there are 1 to 3) or gain enough to fill the p sublevel (if there are 5 to 7).
• Charge = protons − electrons. To count electrons in an ion, subtract a positive charge from Z or add the size of a negative charge to it.
• Main-group ions usually end up with a noble gas configuration.

Atom	Configuration	Ion	Electrons in ion
Li (Z = 3)	1s² 2s¹	Li⁺	2
N (Z = 7)	1s² 2s² 2p³	N³⁻	10
Al (Z = 13)	[Ne] 3s² 3p¹	Al³⁺	10
S (Z = 16)	[Ne] 3s² 3p⁴	S²⁻	18
K (Z = 19)	[Ar] 4s¹	K⁺	18

Transition elements lose 4s electrons first, then may lose 3d electrons. So one element gives several ions: Fe²⁺ and Fe³⁺; Cu⁺ and Cu²⁺. Show the charge with a Roman numeral in the name.

2.1.2 Formulas and names

Method in steps: formula from a name

1. Write both ions with charges.
2. Lowest common multiple of the charge sizes = total charge each side.
3. Divide to get the number of each ion.
4. Cation first; brackets round a polyatomic ion that appears more than once.

Method in steps: name from a formula

1. Name the cation. For a transition metal, find its charge from the anion total and add a Roman numeral.
2. Name the anion: element stem + -ide for a single atom; the ion’s own name if polyatomic.

Polyatomic ions (learn all seven)

Ion	Formula	Ion	Formula
ammonium	NH₄⁺	carbonate	CO₃²⁻
hydroxide	OH⁻	sulfate	SO₄²⁻
nitrate	NO₃⁻	phosphate	PO₄³⁻
hydrogencarbonate	HCO₃⁻

Small worked reminders

rubidium nitride:        Rb⁺, N³⁻        -> 3 Rb⁺ : 1 N³⁻        Rb₃N
indium oxide:            In³⁺, O²⁻       -> 2 : 3                In₂O₃
iron(III) carbonate:     Fe³⁺, CO₃²⁻     -> 2 : 3                Fe₂(CO₃)₃
ammonium hydrogencarbonate: NH₄⁺, HCO₃⁻  -> 1 : 1                NH₄HCO₃
Mn(NO₃)₂: two NO₃⁻ give 2−, so Mn²⁺     -> manganese(II) nitrate
Li₂Se: binary, anion Se²⁻               -> lithium selenide

2.1.3 Lattices and properties

• Structure: giant 3D lattice; each ion surrounded by oppositely charged ions; attractions act in all directions.
• Formula: empirical, because a crystal has no fixed number of ions and no discrete molecules.

Property	What you see	Reason
Volatility	Low; melting and boiling points are high	Strong attractions between a huge number of ions have to be overcome
Conductivity, solid	Does not conduct	Ions fixed in the lattice
Conductivity, molten or aqueous	Conducts	Ions can drift, carrying charge
Solubility in water	Often soluble	Polar water molecules attract and surround the ions
Solubility in non-polar solvents	Generally insoluble	No attraction strong enough to separate the ions

Lattice enthalpy: what changes it

Change	Effect on lattice enthalpy	Why
Higher ion charges	Increases (charge usually dominates)	Stronger attraction between ions
Smaller ions	Increases	Ion centres closer together
Larger ions	Decreases	Ion centres further apart

Method for a ranking question: sort by charge product first, then by the sum of ionic radii within each group.

Model explanations to adapt

Aim for short, complete answers. Each frame below names the particles, the force and the consequence.

Why is the melting point high?

The solid is a giant lattice of oppositely charged ions. Every ion is pulled strongly towards all of the oppositely charged neighbours around it. Breaking up that network takes a great deal of energy, which is why the melting point is high.

Why does it conduct only when molten or dissolved?

In the solid, each ion is locked into its lattice site, so nothing charged can travel. Melting or dissolving releases the ions, which then drift towards the electrodes and carry the current.

Why does it dissolve in water?

Water is polar. The δ− oxygen atoms are attracted to cations and the δ+ hydrogen atoms are attracted to anions. These attractions separate ions from the lattice and surround them in solution.

Why does compound X have the larger lattice enthalpy?

State the charges of both ions in each compound. State which ions are smaller. Then link: higher charge and smaller radius mean a shorter distance between ion centres and a stronger electrostatic attraction, so more energy is needed to separate the ions.

Worked reminder: an unusual pair

Compare potassium sulfide, K₂S, with calcium sulfide, CaS. Both contain S²⁻. Ca²⁺ is doubly charged and smaller than K⁺ (radii of 100 pm against 138 pm). Both factors raise the attraction, so CaS has the larger lattice enthalpy and the higher melting point. Note that K₂S needs two K⁺ for each S²⁻, while CaS is 1 : 1.

Must-know distinctions

• Ion vs atom: an ion has a different number of electrons; protons never change.
• Binary vs polyatomic naming: sodium sulfide is Na₂S; sodium sulfate is Na₂SO₄.
• Empirical formula vs molecular formula: ionic compounds only have the first. Calling CaO a molecule is wrong.
• Molten vs aqueous vs solid: the first two conduct, the last does not. Ions move; electrons do not.
• Lattice enthalpy vs melting point: related, but lattice enthalpy refers to gaseous ions, while melting gives a liquid.

Where marks are usually lost

• Giving Ga or Al a 5− charge “to fill the shell” instead of losing three electrons.
• Forgetting that the electron count of an anion is Z plus the charge size, and so miscounting Se²⁻ or N³⁻.
• Writing (NH₄)PO₄ or NH₄₃PO₄ instead of (NH₄)₃PO₄.
• Naming SO₄²⁻ as sulfide or sulfite. The guide lists seven polyatomic ions to know by name and formula; use those names exactly.
• Writing iron oxide for Fe₂O₃ without the (III), or adding a numeral to calcium or potassium.
• Saying molten ionic compounds conduct “because electrons are delocalized”.
• Describing ionic bonding as “sharing” or as an attraction between just two ions.
• Using radius to rank compounds that have different charges, and missing that the doubly charged compound wins.
• Stating that all ionic compounds dissolve in water.

Quick self-test

1. Give the formula of the ion formed by Z = 20.
2. How many electrons are in a phosphide ion?
3. Write the formula of strontium nitrate.
4. Write the formula of aluminium selenide.
5. Name Ni(OH)₂.
6. Name Co₃(PO₄)₂.
7. Why is the formula of potassium iodide an empirical formula?
8. Explain why molten calcium chloride conducts but solid calcium chloride does not.
9. Suggest why an ionic compound does not dissolve in hexane.
10. Use these radii (pm) to rank RbF, RbI and SrO by decreasing lattice enthalpy: Rb⁺ 152, Sr²⁺ 118, F⁻ 133, I⁻ 220, O²⁻ 140.
11. Copper forms two oxides. Give the name of each, with formulas Cu₂O and CuO.
12. State the total number of ions in one formula unit of Ga₂Se₃.

Answers

1. Ca²⁺ (calcium, [Ar] 4s², loses two electrons).
2. 18 (15 protons, gains 3 electrons).
3. Sr(NO₃)₂.
4. Al₂Se₃.
5. Nickel(II) hydroxide.
6. Cobalt(II) phosphate (two PO₄³⁻ give 6−, shared by three cobalt ions, so each is 2+).
7. A crystal contains a vast, variable number of ions with no separate molecules, so the formula gives only the 1 : 1 ratio of K⁺ to I⁻.
8. Melting releases the Ca²⁺ and Cl⁻ ions from their lattice sites so they can drift and carry charge; in the solid they stay locked in place.
9. Hexane is non-polar, so it cannot attract the ions strongly enough to pull them out of the lattice.
10. SrO > RbF > RbI. SrO has 2+ and 2− ions. Between the singly charged pair, RbF has a smaller distance between ion centres (285 pm against 372 pm).
11. Cu₂O is copper(I) oxide; CuO is copper(II) oxide.
12. Five: two Ga³⁺ and three Se²⁻.

Next steps

• Study guide for Structure 2.1 and Structure 2.1 practice questions
• Wider course pages: the syllabus guide, the IB chemistry hub and the printable checklist

Official syllabus

Source: the International Baccalaureate’s Diploma Programme Chemistry guide, issued in February 2023 for courses examined from 2025. This page follows its sub-topic Structure 2.1, The ionic model: understandings 2.1.1, 2.1.2 and 2.1.3, which are all common to SL and HL.