IB DP Chemistry Structure 2.1: The ionic model – Practice Questions

These are original questions written for Marlbridge, for revision and practice on this content. They are not reproduced past-paper questions, and they do not replicate the exam’s exact structure, question count or mark tariffs – the IB holds copyright in its own papers. Use these alongside the official past papers available through your school or the IB store.

This set practises Structure 2.1, The ionic model, following the IB’s Diploma Programme Chemistry guide (issued February 2023; first assessment 2025). Questions 1 to 6 cover understandings 2.1.1 and 2.1.2; questions 7 to 11 cover 2.1.3, and the longer ones draw on all three. The whole sub-topic is shared by SL and HL, since it has no AHL extension.

If the content is new to you, read the study guide first; for a quick refresh, the revision notes are enough. The IB DP chemistry hub and the course checklist to print list the rest of the course.

Questions

1. Magnesium has atomic number 12.

(a) State what is meant by the term cation. [1] (b) Give the formula of the ion formed by magnesium and the number of electrons it contains. [2]

2. The electron configurations of three atoms, J, L and M, are shown. J, L and M are not their chemical symbols.

J: [Ne] 3s² 3p⁵ L: [Ar] 4s² M: [Ar] 3d¹⁰ 4s² 4p³

(a) Deduce the charge on the ion formed by each atom. [3] (b) Deduce the formula of the compound formed between L and M, using the letters given. [1]

3. Iron and copper are transition elements.

(a) Explain why the name of an iron compound includes a Roman numeral but the name of a sodium compound does not. [1] (b) Name Fe(OH)₂ and FePO₄. [2] (c) Write the formula of copper(I) sulfide. [1]

4. Write the formula of each compound.

(a) ammonium sulfate [1] (b) scandium nitrate [1] (c) potassium hydrogencarbonate [1] (d) iron(II) phosphate [1]

5. Name each compound.

(a) Cs₂S [1] (b) Ca(OH)₂ [1] (c) Cr₂(CO₃)₃ [1]

6. A technician makes up a solution containing 0.150 mol of gallium sulfate, Ga₂(SO₄)₃. Take the Avogadro constant as 6.02 × 10²³ mol⁻¹.

(a) State the number of ions in one formula unit of gallium sulfate. [1] (b) Calculate the total number of ions present in the solution. [2] (c) Calculate the number of sulfate ions present. [1]

7. The formula of solid potassium iodide is written as KI.

(a) Describe the structure of solid potassium iodide. [2] (b) Explain why KI is described as an empirical formula. [1]

8. A student tests lithium chloride for electrical conductivity in three states: as a solid, as a liquid at high temperature, and as a solution in water.

(a) Predict, for each state, whether lithium chloride conducts. [1] (b) Explain your predictions in terms of the particles present. [3]

9. A sample of rubidium bromide dissolves in water but not in hexane. Explain this difference. [3]

10. Use these values for the lattice enthalpies: NaF 926 kJ mol⁻¹, MgO 3791 kJ mol⁻¹. Use these ionic radii (pm): Na⁺ 102, Mg²⁺ 72, F⁻ 133, O²⁻ 140.

(a) Write an equation, including state symbols, for the process whose enthalpy change is the lattice enthalpy of magnesium oxide. [2] (b) Explain why the lattice enthalpy of MgO is roughly four times that of NaF. [3] (c) Predict which of the two compounds has the higher melting point, and explain your answer. [2]

11. Use these ionic radii (pm): Li⁺ 76, Cs⁺ 167, Ca²⁺ 100, Ba²⁺ 135, Cl⁻ 181, O²⁻ 140.

(a) Rank caesium chloride, lithium chloride, barium oxide and calcium oxide in order of decreasing lattice enthalpy. Explain your order. [4] (b) Lithium chloride is described as having low volatility. Explain this property. [2] (c) Calcium oxide forms when calcium burns in oxygen. Use electron configurations to explain the charge on each ion formed. [2]

Answers

1. (a) A positive ion, formed when an atom loses one or more electrons [1] (b) Mg²⁺ [1]; 12 − 2 = 10 electrons [1] Examiner insight: The charge must be written with both number and sign, as 2+; “Mg2” or “Mg⁺²” is not the correct way to write the ion.

2. (a) J has 7 outer electrons and gains 1: J⁻ [1]. L has 2 outer electrons and loses them: L²⁺ [1]. M has 5 outer electrons (4s² 4p³) and gains 3: M³⁻ [1] (b) LCM of 2 and 3 is 6, so 3 L²⁺ and 2 M³⁻: L₃M₂ [1] Examiner insight: For M, count only the n = 4 electrons; the full 3d¹⁰ sublevel is not part of the outer level, and counting it leads to a wrong charge.

3. (a) Iron forms ions of more than one charge (Fe²⁺ and Fe³⁺), so the numeral shows which ion is present; sodium only forms Na⁺ [1] (b) iron(II) hydroxide [1]; iron(III) phosphate [1] (c) Cu⁺ and S²⁻ in a 2 : 1 ratio: Cu₂S [1] Examiner insight: In (b), the numeral comes from the anion: PO₄³⁻ carries 3−, so the single iron ion must be 3+. Show that link if a question asks you to explain the name.

4. (a) (NH₄)₂SO₄ [1] (b) Sc(NO₃)₃ [1] (c) KHCO₃ [1] (d) Fe²⁺ and PO₄³⁻, LCM 6: Fe₃(PO₄)₂ [1] Examiner insight: Brackets are part of the formula. Without them, a formula such as Fe3PO42 is ambiguous, so write the brackets every time a polyatomic ion appears more than once.

5. (a) caesium sulfide [1] (b) calcium hydroxide [1] (c) Three CO₃²⁻ give 6−, shared by two chromium ions: chromium(III) carbonate [1] Examiner insight: Check the ending of the anion name: “sulfate” for Cs₂S or “carbide” for CO₃²⁻ names a different species altogether.

6. (a) 2 Ga³⁺ + 3 SO₄²⁻ = 5 ions [1] (b) moles of ions = 5 × 0.150 = 0.750 mol [1]; number = 0.750 × 6.02 × 10²³ = 4.52 × 10²³ [1] (c) 3 × 0.150 × 6.02 × 10²³ = 2.71 × 10²³ sulfate ions [1] Examiner insight: Write the ion count per formula unit in your working, so your method stays clear to the reader even if the arithmetic slips.

7. (a) A giant three-dimensional lattice of K⁺ and I⁻ ions in a regular repeating arrangement [1]; each ion is surrounded by ions of the opposite charge, held by electrostatic attraction in all directions [1] (b) There are no discrete molecules; the formula shows only the simplest ratio of ions, 1 : 1 [1] Examiner insight: Words such as “molecule” or “intermolecular forces” in a description of an ionic solid contradict the model, so avoid them even in passing.

8. (a) Solid: does not conduct; liquid (molten): conducts; aqueous: conducts [1] (b) Lithium chloride contains Li⁺ and Cl⁻ ions [1]. In the solid, every ion is locked into its place in the lattice, so no charged particle can travel through it [1]. Once the solid melts or dissolves, the Li⁺ and Cl⁻ ions can drift through the liquid and carry charge to the electrodes [1] Examiner insight: Name ions as the charge carriers; an answer that mentions moving electrons for any of the three states describes metallic conduction instead.

9. Water is polar; its δ− oxygen is attracted to Rb⁺ and its δ+ hydrogen atoms are attracted to Br⁻ [1]. These attractions are strong enough to separate the ions from the lattice and surround them [1]. Hexane is non-polar, so it cannot attract the ions strongly enough to overcome the attractions in the lattice [1] Examiner insight: Get the orientation right: say which end of the water molecule faces which ion, because reversing them is a factual error.

10. (a) MgO(s) → Mg²⁺(g) + O²⁻(g): correct species and direction [1]; correct state symbols [1] (b) Mg²⁺ and O²⁻ carry charges of 2+ and 2−, against 1+ and 1− in NaF, so the charge product is four times larger [1]. The ions in MgO are also slightly closer (72 + 140 = 212 pm against 102 + 133 = 235 pm) [1]. A higher charge product and a shorter distance give a stronger electrostatic attraction, so far more energy is needed to separate the ions; the charge factor accounts for most of the difference (3791/926 ≈ 4.1) [1] (c) MgO has the higher melting point [1]; its ions attract each other more strongly, so more energy is needed to break down the lattice [1] Examiner insight: In (a), the lattice enthalpy refers to separating the solid into gaseous ions; writing atoms, or a reaction forming the solid from elements, describes a different enthalpy change.

11. (a) CaO > BaO > LiCl > CsCl [1]. CaO and BaO contain 2+ and 2− ions, so they have stronger attractions than the singly charged chlorides [1]. Ca²⁺ is smaller than Ba²⁺ (distances 240 pm and 275 pm), so CaO is larger than BaO [1]. Li⁺ is smaller than Cs⁺ (distances 257 pm and 348 pm), so LiCl is larger than CsCl [1] (b) Each ion is strongly attracted to the many oppositely charged ions around it, right across the lattice [1]; overcoming all of these attractions takes a lot of energy, which gives lithium chloride high melting and boiling points [1] (c) Calcium ([Ar] 4s²) loses its two 4s electrons to form Ca²⁺, with the configuration of argon [1]. Oxygen (1s² 2s² 2p⁴) gains two electrons to fill 2p, forming O²⁻ with the configuration of neon [1] Examiner insight: In (a), back up each placing with a reason that names charge or radius, and when the charges differ, make the charge comparison first.

Where marks are usually lost

• Writing charges with the sign first (Mg⁺²) or leaving out the number (Ca⁺ for calcium).
• Counting the full 3d¹⁰ sublevel as outer electrons for elements such as gallium, arsenic or selenium.
• Dropping brackets round polyatomic ions, or changing the formula of the polyatomic ion itself.
• Giving the Roman numeral without checking the charge carried by the anion.
• Multiplying by the number of formula units instead of the number of ions per formula unit in mole questions.
• Describing an ionic lattice using “molecules” or “intermolecular forces”.
• Stating that ionic compounds conduct because electrons move, or that the solid conducts at all.
• Describing water’s attraction to ions with the dipole the wrong way round.
• Ranking lattice enthalpies by radius alone when the ions carry different charges.
• Writing a lattice enthalpy equation with atoms, or with the solid as the product.

Next steps

• Structure 2.1 revision notes for a final read-through
• Structure 2.1 study guide for the full explanations
• The IB DP chemistry course hub and the printable checklist
• Try all free 10-minute diagnostics
• Book a free trial class

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.