S2.1 The ionic model
Practice exam-style IB Chemistry questions for The ionic model, aligned with the syllabus and grouped by topic.
Question 1
A magnesium atom forms an ion. What change occurs?
A.
It gains two protons to form Mg²⁺.
B.
It loses two electrons to form Mg²⁺.
C.
It gains two electrons to form Mg²⁻.
D.
It loses two neutrons to form Mg²⁺.
Question 2
The electron configuration of an oxygen atom is 1s²2s²2p⁴. What is the electron configuration of the oxide ion?
A.
1s²2s²2p⁴
B.
1s²2s²2p⁶
C.
1s²2s²2p⁶3s²
D.
1s²2s²2p²
Question 3
What is the formula of aluminium sulfate?
A.
Al₂SO₄
B.
Al₃(SO₄)₂
C.
AlSO₄
D.
Al₂(SO₄)₃
Question 4
What is the name of Ca₃N₂?
A.
tricalcium dinitride
B.
calcium nitrate
C.
calcium nitrogen
D.
calcium nitride
Question 5
What describes an ionic bond?
A.
Attraction between a shared pair of electrons and two nuclei
B.
Attraction between a metal nucleus and delocalized electrons
C.
Repulsion between ions with the same charge
D.
Attraction between oppositely charged ions
Question 6
State the meaning of the term cation.
[1]
State the charge on the ion commonly formed by sulfur.
[1]
Write the half-equation for the formation of Na⁺ from Na.
[1]
Question 7
Name MgO.
[1]
Name FeCl₃.
[1]
Explain the purpose of the Roman numeral in the name of FeCl₃.
[1]
Question 8
Solid sodium chloride does not conduct electricity but molten sodium chloride does. What accounts for this difference?
A.
Electrons are delocalized only in the molten state.
B.
Ions are mobile only in the molten state.
C.
Sodium atoms are formed only in the molten state.
D.
Covalent bonds break only in the molten state.
Question 9
NaF has a larger lattice enthalpy than KF. What is the best explanation?
A.
F⁻ has a smaller charge in NaF than in KF.
B.
KF contains covalent bonds while NaF is ionic.
C.
Na⁺ has a larger charge than K⁺.
D.
Na⁺ has a smaller ionic radius than K⁺.
Question 10
What is the usual orientation of water molecules around a chloride ion in aqueous solution?
A.
Water molecules transfer protons to Cl⁻.
B.
Oxygen atoms point towards Cl⁻.
C.
Water molecules remain randomly oriented.
D.
Hydrogen atoms point towards Cl⁻.
Question 11
An iron atom is represented as [Ar]4s²3d⁶. What is the electron configuration of Fe²⁺?
A.
[Ar]3d⁶
B.
[Ar]4s²3d⁴
C.
[Ar]3d⁴
D.
[Ar]4s¹3d⁵
Question 12
A conjugate base is more stable. What is the expected effect on the acid dissociation constant, Kₐ, of its conjugate acid?
A.
Kₐ decreases because fewer ions can form.
B.
Kₐ increases because dissociation is more favourable.
C.
Kₐ becomes zero because the acid cannot dissociate.
D.
Kₐ is unrelated to conjugate-base stability.
Question 13
Transition-element ion charges are less predictable from group number than main-group ion charges. What is the reason?
A.
The ns and (n−1)d electrons are close in energy.
B.
The nuclear charge changes during ion formation.
C.
Transition elements gain electrons to reach noble gases.
D.
Transition elements do not contain valence electrons.
Question 14
The electron configuration of chlorine is 1s²2s²2p⁶3s²3p⁵.
Deduce the charge on the ion formed when chlorine reaches a noble gas configuration.
[1]
State the electron configuration of this ion.
[1]
Identify the noble gas with which this ion is isoelectronic.
[1]
Question 15
State the formula of the nitrate ion.
[1]
Deduce the formula of calcium nitrate.
[1]
Deduce the formula of ammonium phosphate.
[1]
Explain why brackets are used in the formula of calcium nitrate.
[1]
Question 16
Sodium chloride forms when sodium reacts with chlorine.
State which element is oxidized.
[1]
State which element is reduced.
[1]
Explain why formation of sodium chloride from its elements is a redox reaction.
[1]
Question 17
The graph shows electrical conductivity for potassium iodide under different conditions.
Identify the condition with the highest conductivity.
[1]
Describe the change in conductivity when solid potassium iodide is melted.
[1]
Explain the difference in conductivity between the solid and molten samples.
[1]
Question 18
A table lists ions present in four unknown ionic compounds.
| Compound | Cation | Anion |
|---|---|---|
| A | Ba²⁺ | OH⁻ |
| B | Al³⁺ | CO₃²⁻ |
| C | NH₄⁺ | NO₃⁻ |
| D | Ca²⁺ | PO₄³⁻ |
Deduce the formula of the compound containing Ba²⁺ and OH⁻.
[1]
Deduce the formula of the compound containing Al³⁺ and CO₃²⁻.
[1]
Name the compound containing NH₄⁺ and NO₃⁻.
[1]
Explain why ionic formulas must be electrically neutral overall.
[1]
Question 19
Successive ionization energies of a first-row transition element increase gradually for several electrons before a much larger increase. What does this support?
A.
Variable oxidation states can occur.
B.
The nucleus loses protons during ionization.
C.
Transition metals form anions more readily than cations.
D.
Only one fixed cation charge is possible.
Question 20
For sulfate, which Lewis representation is preferred using formal charge considerations?
A.
The one in which sulfur has no bonds to oxygen.
B.
The one with the most reasonable distribution of formal charges.
C.
The one with the largest number of atoms carrying formal charge.
D.
The one in which the overall charge is zero.
Question 21
Aluminium chloride has a much lower melting point than predicted by a purely ionic model. What is the best explanation?
A.
Aluminium chloride has significant covalent character.
B.
Aluminium ions have no charge.
C.
Chloride ions are fixed in the liquid.
D.
Aluminium chloride contains only metallic bonds.
Question 22
Across NaCl, MgCl₂ and AlCl₃, what factor becomes increasingly important in explaining the melting point trend?
A.
Increasing number of neutrons in chloride ions
B.
Increasing covalent character of the chloride compounds
C.
Formation of separate NaCl molecules only
D.
Decreasing mobility of gaseous ions
Question 23
The acid H₂CO₃ has carbonate as its conjugate base after loss of two protons. What is the formula of the carbonate ion?
A.
HCO₃²⁻
B.
CO₃⁻
C.
CO₃²⁻
D.
C₂O₃²⁻
Question 24
State what is meant by an ionic lattice.
[1]
Explain why ionic compounds are represented by empirical formulas.
[1]
Explain why solid potassium bromide does not conduct electricity.
[1]
Explain why molten potassium bromide conducts electricity.
[1]
Question 25
The lattice enthalpy of MgO is greater than that of NaF.
State the meaning of lattice enthalpy as used in the IB data booklet.
[1]
Suggest why MgO has the greater lattice enthalpy.
[1]
Question 26
A small amount of sodium chloride is added to water.
State the type of attraction formed between Na⁺ ions and water molecules.
[1]
Describe the orientation of water molecules around Na⁺.
[1]
Explain why many ionic compounds dissolve in water but not in hexane.
[1]
Question 27
Vanadium can form several cations.
State why transition-element ion charges are not predicted simply from group number.
[1]
Explain how successive ionization energy data can support variable oxidation states in a transition element.
[1]
Question 28
Chromium is represented as [Ar]4s¹3d⁵.
Deduce the electron configuration of Cr²⁺.
[1]
Deduce the electron configuration of Cr³⁺.
[1]
State the order in which 4s and 3d electrons are removed during ionization of first-row transition elements.
[1]
Question 29
Polyatomic anions such as nitrate and carbonate are conjugate bases of acids.
Define a conjugate base in this context.
[1]
State the relationship between the stability of a conjugate base and the Kₐ of its conjugate acid.
[1]
Explain the relationship stated in (b).
[1]
Question 30
Compare main-group and transition-element cation formation.
State how main-group ion charge is often predicted.
[1]
State why transition elements often form more than one cation.
[1]
Give one example of a transition element forming two different cations.
[1]
Question 31
Hydrogencarbonate and carbonate are related conjugate bases.
State the formula of hydrogencarbonate.
[1]
Deduce the formula of magnesium hydrogencarbonate.
[1]
Explain why carbonate has a greater negative charge than hydrogencarbonate.
[1]
Question 32
The table gives lattice enthalpies and cation radii for several group 1 chlorides.
| Compound | Cation radius / pm | Lattice enthalpy / kJ mol⁻¹ |
|---|---|---|
| LiCl | 90 | 853 |
| NaCl | 116 | 786 |
| KCl | 152 | 715 |
| RbCl | 166 | 689 |
| CsCl | 181 | 657 |
State the trend in cation radius down group 1.
[1]
Describe the trend in lattice enthalpy shown by the data.
[1]
Explain the trend in lattice enthalpy using the ionic model.
[1]
Suggest why all compounds in the table have the same anion charge.
[1]
Question 33
A solubility investigation used water and hexane as solvents for three crystalline solids.
| Solid | Solubility in water | Solubility in hexane | Conductivity of aqueous mixture |
|---|---|---|---|
| A | Dissolves | Does not dissolve | Conducts |
| B | Dissolves | Does not dissolve | Does not conduct |
| C | Does not dissolve | Dissolves | Does not conduct |
Identify the solid most likely to be ionic from the data.
[1]
State one observation that supports your answer to (a).
[1]
Explain why water is a better solvent than hexane for many ionic compounds.
[1]
Suggest one safety precaution for this investigation.
[1]
Question 34
The diagram shows the numbers of protons and electrons in four particles.
| Particle | Number of protons | Number of electrons |
|---|---|---|
| A | 12 | 12 |
| B | 13 | 10 |
| C | 17 | 18 |
| D | 6 | 6 |
Identify the particle that is a cation.
[1]
Identify the particle that is an anion.
[1]
Explain your answer to (a).
[1]
Deduce the charge on a particle with 13 protons and 10 electrons.
[1]
Question 35
Sulfate is a polyatomic ion.
State the formula and charge of sulfate.
[1]
State what formal charge is used for in comparing Lewis representations of sulfate.
[1]
Explain two features of a preferred sulfate representation in terms of formal charge.
[1]
Question 36
The melting points of NaCl, MgCl₂ and AlCl₃ do not increase simply with cation charge.
State one factor, other than charge, that affects lattice enthalpy.
[1]
Explain why a purely ionic model is inadequate for AlCl₃.
[1]
State the term used to describe the gradual change from ionic to covalent bonding.
[1]
Question 37
A student claims: “A larger electronegativity difference always proves that a compound is a simple ionic lattice.”
State the general relationship between electronegativity difference and ionic character.
[1]
Give one limitation of using electronegativity difference alone.
[1]
Use aluminium chloride as an example to evaluate the claim.
[1]
Question 38
The graph shows successive ionization energies for a first-row transition element.
Identify the first ionization for which there is a much larger increase than the previous value.
[1]
State what this large increase indicates about the electron being removed.
[1]
Explain how the earlier gradual increases support variable oxidation states.
[1]
Suggest why group number alone is not sufficient to predict the common ions of this element.
[1]
Question 39
The graph shows melting points of selected period 3 chlorides.
Identify the chloride with the lowest melting point in the data.
[1]
Describe the change from NaCl to MgCl₂.
[1]
Explain why lattice enthalpy alone is insufficient to explain the value for AlCl₃.
[1]
State the bonding concept illustrated by the data.
[1]
Question 40
The table compares three oxoacids and the formulas of their conjugate bases.
| Acid | Conjugate base | Base stability / a.u. | Kₐ at 298 K / mol dm⁻³ |
|---|---|---|---|
| HClO | ClO⁻ | 2.1 | 3.0 × 10⁻⁸ |
| HClO₂ | ClO₂⁻ | 6.4 | 1.1 × 10⁻² |
| HClO₃ | ClO₃⁻ | 8.9 | 1.0 × 10¹ |
Identify the conjugate base with the greatest stability according to the data.
[1]
State the acid with the largest Kₐ.
[1]
Explain the relationship between conjugate-base stability and Kₐ.
[1]
Question 41
The table gives electron configurations for a transition metal atom and two of its ions.
| Species | Electron configuration |
|---|---|
| V atom (neutral) | [Ar] 4s² 3d³ |
| V²⁺ ion | [Ar] 3d³ |
| V³⁺ ion | [Ar] 3d² |
Identify the ion formed by loss of the 4s electrons only.
[1]
Deduce how many 3d electrons are present in the 3+ ion.
[1]
Explain why the 4s electrons are removed before 3d electrons.
[1]
Suggest why both 2+ and 3+ ions can be observed for this element.
[1]
Question 42
Ionic compounds can have high melting points and different electrical conductivities in different states.
Describe the structure and bonding in a solid ionic compound.
[1]
Explain the high melting point and the electrical conductivity of solid, molten and aqueous ionic compounds.
[1]
Question 43
Magnesium nitride forms from magnesium and nitrogen.
Deduce the ions formed from magnesium and nitrogen and the formula of magnesium nitride.
[1]
Discuss the formation of magnesium nitride from its elements in terms of electron transfer, redox and ionic bonding.
[1]
Question 44
Sodium, aluminium, oxygen and chlorine form common monatomic ions.
Deduce the charges of the common ions formed by sodium and oxygen.
[1]
Compare and contrast the formation of cations and anions, referring to electron configurations and the octet rule.
[1]
Question 45
Two possible Lewis representations of sulfate are shown with formal charges marked.
State the overall charge on sulfate.
[1]
Identify the representation preferred by formal charge considerations.
[1]
Explain one reason for your choice in (b).
[1]
Explain why sulfate is treated as a single unit when writing ionic formulas.
[1]
Deduce the formula of iron(III) sulfate.
[1]
Question 46
A student uses melting point, solubility and conductivity data to decide whether an unknown crystalline solid is ionic.
State two physical properties expected for many ionic compounds.
[1]
Evaluate how melting point, solubility in water and hexane, and conductivity in solid and aqueous states can be used to identify an ionic compound.
[1]
Question 47
Transition elements often form ions with different charges.
Describe how Fe²⁺ and Fe³⁺ are formed from Fe, represented as [Ar]4s²3d⁶.
[1]
Explain how subshell energies and successive ionization energy data account for variable oxidation states in first-row transition elements.
[1]
Question 48
The melting points of period 3 chlorides provide evidence for limitations of the simple ionic model.
Explain how lattice enthalpy is affected by ionic charge and radius.
[1]
Evaluate the use of the ionic model to explain the melting points of NaCl, MgCl₂ and AlCl₃.
[1]
Question 49
Sulfate and nitrate are polyatomic anions used in many ionic compounds.
State the formulas of sulfate and nitrate, and deduce the formula of aluminium sulfate.
[1]
Discuss how formal charge and conjugate-base stability are used to understand polyatomic anions.
[1]
Question 50
Models are used to describe ionic bonding and polyatomic ions.
Outline how electronegativity difference is used to estimate ionic character.
[1]
Evaluate the strengths and limitations of using simple ionic, formal charge and conjugate-base stability models to explain compounds containing carbonate or sulfate ions.
[1]
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