IB Chemistry Quiz: Understand The Ionic Model
20 questions · exam conditions
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Understand The Ionic ModelQuestion 1 of 20

The chemical formula for scandium molybdate is Sc₂(MoO₄)₃. Given that scandium is in Group 3 and typically forms a stable 3+ ion (Sc³⁺), what is the charge on the molybdate polyatomic ion?

2-
3-
1-
6-
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IB Chemistry Quiz

IB Chemistry Quiz: Understand The Ionic Model

Practice Understand The Ionic Model in IB Chemistry with focused quiz questions that help you check what you know, review explanations, and build confidence with test-style prompts.

What this quiz covers

This quiz focuses on Understand The Ionic Model, giving you a quick way to practice the rules, question types, and explanations that matter most for IB Chemistry.

How to use this quiz

Try each quiz question before looking at the correct answer. Use the explanations to review missed ideas, then come back to similar questions until the pattern feels familiar.

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Question 1

The chemical formula for scandium molybdate is Sc₂(MoO₄)₃. Given that scandium is in Group 3 and typically forms a stable 3+ ion (Sc³⁺), what is the charge on the molybdate polyatomic ion?

  1. 2- (correct answer)
  2. 3-
  3. 1-
  4. 6-
Explanation: An ionic compound must be electrically neutral. The formula contains two scandium ions. Since each scandium ion has a charge of 3+, the total positive charge from the scandium ions is 2 × (+3) = +6. To balance this, the total negative charge from the three molybdate ions must be -6. Therefore, the charge on a single molybdate ion (MoO₄) is (-6) / 3 = -2. The molybdate ion is MoO₄²⁻.

Question 2

Ammonium nitrate, NH₄NO₃, is an ionic compound. Which statement best explains its electrical conductivity?

  1. It conducts electricity in the solid state because electrons are delocalized throughout the crystal lattice.
  2. It does not conduct electricity in any state because it is composed entirely of non-metal atoms.
  3. It conducts electricity when molten or in aqueous solution because its ions become mobile charge carriers. (correct answer)
  4. It conducts electricity only when dissolved in a non-polar solvent that allows the ions to be fully separated.
Explanation: Ionic compounds are characterized by a crystal lattice structure in the solid state where ions are held in fixed positions and are not free to move, making them electrical insulators. When molten or dissolved in a polar solvent like water, the ions (NH₄⁺ and NO₃⁻) dissociate and become mobile, allowing them to move and carry an electrical current. Distractor A describes metallic bonding. Distractor B incorrectly assumes that the absence of a metal atom precludes ionic bonding (NH₄⁺ is a polyatomic cation). Distractor D is incorrect as ionic compounds are generally insoluble in non-polar solvents.

Question 3

A solid substance is a good electrical insulator at 25 °C but becomes a good conductor when heated above its melting point of 801 °C. The solid is also hard and brittle. Which substance fits this description?

  1. Graphite
  2. Sodium chloride (correct answer)
  3. Silicon dioxide
  4. Iodine
Explanation: The properties described—hard, brittle, high melting point, electrical insulator as a solid, and conductor as a liquid—are characteristic of an ionic compound. In solid sodium chloride (NaCl), ions are in fixed positions in the lattice. When molten (melting point = 801 °C), the Na⁺ and Cl⁻ ions are free to move and act as charge carriers. Graphite (A) is a conductor in the solid state. Silicon dioxide (C) is a covalent network solid with a very high melting point, but it does not conduct electricity when molten. Iodine (D) is a molecular solid with a low melting point and does not conduct electricity in either solid or liquid state.

Question 4

The ionic model assumes a complete transfer of electrons. In reality, all ionic bonds have some degree of covalent character. Which compound is expected to have the bonding with the most ionic character?

  1. Caesium fluoride, CsF (correct answer)
  2. Beryllium chloride, BeCl₂
  3. Aluminium oxide, Al₂O₃
  4. Lithium iodide, LiI
Explanation: The degree of ionic character is related to the difference in electronegativity (ΔEN) between the two bonding atoms. The largest ΔEN will result in the most ionic character. Caesium is the least electronegative element (or among the least), and fluorine is the most electronegative element. Therefore, the bond between Cs and F will have the greatest possible electronegativity difference and thus the most ionic character. BeCl₂, Al₂O₃, and LiI all have smaller electronegativity differences and exhibit more significant covalent character.

Question 5

The overall formation of an ionic compound from its elements can be exothermic, even if the formation of gaseous ions from atoms is highly endothermic. For which compound's formation is the lattice enthalpy most critical in compensating for a highly endothermic ion-formation process?

  1. Potassium chloride, KCl
  2. Calcium fluoride, CaF₂
  3. Magnesium oxide, MgO (correct answer)
  4. Sodium bromide, NaBr
Explanation: The formation of gaseous ions can be very endothermic, especially when forming ions with charges greater than 1. To form Mg²⁺(g) requires both the first and second ionization energies of magnesium, a large energy input. To form O²⁻(g), the first electron affinity is exothermic, but the second is endothermic (adding an electron to an already negative ion). The combined energy to form Mg²⁺(g) and O²⁻(g) is highly endothermic. The formation of MgO(s) is only feasible because of the extremely large, exothermic lattice enthalpy associated with the 2+ and 2- ions, which compensates for the initial energy investment. The other options involve forming +1 or +2/+1 ions, which are less endothermic processes.

Question 6

The dissolution of an ionic salt in water involves overcoming its lattice enthalpy. Which statement best explains why silver chloride (AgCl) is sparingly soluble in water, while sodium chloride (NaCl) is readily soluble?

  1. The lattice enthalpy of AgCl is significantly larger than that of NaCl, and the hydration enthalpy released is insufficient to overcome it. (correct answer)
  2. NaCl is more soluble because sodium is a more reactive metal than silver, leading to stronger interactions with polar water molecules.
  3. The Ag⁺ ion is much larger than the Na⁺ ion, which physically prevents water molecules from effectively surrounding it during solvation.
  4. The covalent character in the Ag-Cl bond is negligible, making it a purely ionic compound that strongly resists interaction with polar water.
Explanation: Solubility depends on the balance between the energy required to break the ionic lattice (lattice enthalpy) and the energy released when ions are hydrated (hydration enthalpy). AgCl has a higher lattice enthalpy than NaCl, partly due to significant covalent character in its bonding which strengthens the lattice. This high lattice enthalpy is not sufficiently compensated for by the hydration enthalpies of Ag⁺ and Cl⁻ ions, making it sparingly soluble. Distractor B confuses chemical reactivity with hydration enthalpy. Distractor C is factually incorrect as the ionic radii are similar (Ag⁺ ≈ 115 pm, Na⁺ ≈ 102 pm) and not the primary reason. Distractor D incorrectly states AgCl has negligible covalent character; it is the presence of covalent character that increases its lattice enthalpy and decreases its solubility.

Question 7

The theoretical lattice enthalpy of silver iodide (AgI), calculated using a purely ionic model, is significantly smaller in magnitude than the experimental value determined from a Born-Haber cycle. What is the most likely reason for this discrepancy?

  1. The silver iodide lattice adopts a non-standard crystal structure, which is not accounted for in the ionic model.
  2. There is a significant degree of covalent character in the bonding, which adds to the overall lattice strength. (correct answer)
  3. The experimental values for ionization energy and electron affinity used in the Born-Haber cycle have high uncertainty.
  4. The Ag⁺ and I⁻ ions are not perfect spheres, leading to packing inefficiencies that weaken the lattice.
Explanation: A purely ionic model assumes ions are perfect spheres with point charges. When the experimental lattice enthalpy (from a Born-Haber cycle) is greater than the theoretical value, it indicates that the bonding is stronger than predicted by the ionic model alone. This extra stability comes from a degree of covalent character in the bond. The Ag⁺ ion is polarizing and the large I⁻ ion is highly polarizable, leading to significant sharing of electron density (covalent character). This strengthens the overall bonding. Distractor A is a minor effect. Distractor C is incorrect as these values are known precisely. Distractor D is related to the model's assumptions, but covalent character is the specific chemical reason for the large energy difference.

Question 8

A mineral's formula can be represented as Caₓ(OH)ₐ(PO₄)ₙ. If analysis shows the molar ratio of calcium ions to phosphate ions is 5:3, what is the simplest whole number value for 'a', the subscript for the hydroxide ion, that results in a neutral formula unit?

  1. 1 (correct answer)
  2. 2
  3. 3
  4. 5
Explanation: First, determine the charge from the calcium and phosphate ions based on the 5:3 ratio. The formula contains Ca₅(PO₄)₃. Calcium ions have a 2+ charge (Ca²⁺) and phosphate ions have a 3- charge (PO₄³⁻). The total positive charge is 5 × (+2) = +10. The total negative charge from phosphate is 3 × (−3) = −9. The net charge from this portion of the formula is (+10) + (−9) = +1. To make the overall formula unit neutral, the hydroxide portion must contribute a charge of −1. Since each hydroxide ion (OH⁻) has a charge of 1-, only one hydroxide ion is needed. Therefore, a = 1. The full formula is Ca₅(PO₄)₃(OH), known as hydroxylapatite.

Question 9

A student observes that compound X has a melting point of 801°C, conducts electricity when molten but not when solid, and is soluble in polar solvents. However, when X-ray crystallography is performed, the electron density maps show significant overlap between adjacent atoms rather than distinct spherical ions. Which statement best explains this observation?

  1. The compound exhibits partial covalent character due to polarization effects, indicating limitations of the pure ionic model (correct answer)
  2. The compound is purely ionic, and the electron density overlap is an artifact of the X-ray crystallography technique used
  3. The compound is actually covalent network, and the electrical conductivity when molten indicates metallic character
  4. The compound demonstrates perfect ionic character, and electron density overlap is expected for all ionic compounds
Explanation: The observed properties (high melting point, electrical conductivity when molten) suggest ionic character, but the electron density overlap indicates partial covalent character due to polarization. This demonstrates the limitations of the pure ionic model, which assumes complete electron transfer and spherical ions. Real ionic compounds often show some degree of covalent character, especially when formed between small, highly charged ions.

Question 10

When comparing the hydration enthalpies of Li⁺ and Cs⁺ ions, experimental values show that Li⁺ has a much more negative (favorable) hydration enthalpy despite having the same +1 charge. If the ionic model's predictions about ion-dipole interactions are applied, which statement best explains this trend?

  1. Li⁺ can coordinate with more water molecules in its hydration sphere, leading to a greater number of favorable interactions
  2. Cs⁺ experiences greater ion-dipole repulsion due to its larger size, making hydration energetically unfavorable compared to Li⁺
  3. Li⁺ has higher charge density, creating stronger ion-dipole interactions with water molecules and more favorable hydration (correct answer)
  4. The smaller Li⁺ ion creates less disruption of water's hydrogen bonding network, resulting in more efficient hydration
Explanation: When you encounter questions about hydration enthalpies, focus on how ionic charge density affects ion-dipole interactions with water molecules. The key principle is that smaller ions with the same charge create stronger electrostatic attractions. Li⁺ has a much more negative hydration enthalpy than Cs⁺ because of charge density differences. Both ions carry a +1 charge, but Li⁺ is significantly smaller (ionic radius ~0.76 Å vs ~1.67 Å for Cs⁺). This means Li⁺ has much higher charge density - the same charge concentrated in a smaller volume creates a stronger electric field that attracts water dipoles more powerfully. The ion-dipole interaction energy is proportional to chargedistance2\frac{charge}{distance^2}, so the shorter Li⁺-water distances result in much stronger, more favorable interactions. Option A is incorrect because coordination number (number of water molecules) doesn't necessarily correlate with hydration enthalpy strength - it's the individual interaction strength that matters most. Option B misunderstands the physics - larger ions don't experience "repulsion" from their size; they simply have weaker attractive forces due to lower charge density. Option D incorrectly suggests that disruption of water's hydrogen bonding network is the primary factor, when it's actually the direct ion-dipole attraction strength that dominates. Remember this pattern: for ions with the same charge, smaller size means higher charge density, which creates stronger electrostatic interactions with polar molecules like water. This principle applies broadly to ion-dipole interactions in solution chemistry.

Question 11

An ionic compound AB crystallizes in a structure where each A²⁺ ion is surrounded by 8 B²⁻ ions, and each B²⁻ ion is surrounded by 8 A²⁺ ions. If the ionic model perfectly described this compound, what would be the most likely reason for this high coordination number?

  1. The crystal structure maximizes hydrogen bonding between the ions and any residual water molecules in the lattice
  2. The high charges on both ions create strong enough attractions to overcome repulsions between like charges in the high coordination environment
  3. The compound exhibits significant covalent character, requiring directional bonding that favors the 8-coordinate geometry
  4. The ions have similar sizes, allowing efficient packing with maximum electrostatic attractions and minimal repulsions (correct answer)
Explanation: When you encounter questions about crystal structures and coordination numbers, focus on the fundamental factors that determine how ions arrange themselves: size relationships, charge effects, and electrostatic principles. The key to understanding this 8:8 coordination lies in recognizing what enables such high coordination numbers. For ions to surround themselves with 8 nearest neighbors of opposite charge, the size ratio must be favorable. When cation and anion radii are similar (ratio close to 1), they can pack efficiently in cubic arrangements like the cesium chloride structure, maximizing attractive forces while minimizing repulsions between like charges. This geometric compatibility allows the maximum number of stabilizing electrostatic interactions. Option A is incorrect because the question specifies we're using the ionic model, which assumes purely electrostatic interactions—hydrogen bonding with residual water isn't part of this idealized framework. Option B misunderstands the relationship between charge and coordination. While higher charges do create stronger attractions, they actually tend to favor lower coordination numbers because the increased repulsions between like charges in crowded environments become prohibitive. Option C contradicts the premise entirely—the question asks us to assume the ionic model applies perfectly, meaning no covalent character exists. Remember this pattern: high coordination numbers in ionic compounds primarily result from favorable size ratios, not high charges. When ions are similar in size, they can achieve efficient close-packing arrangements that maximize the number of stabilizing electrostatic contacts while maintaining reasonable distances between like charges.

Question 12

A research team measures the Born-Landé lattice energy of several alkali halides and compares them to experimental values obtained from Born-Haber cycles. They find that the theoretical values are consistently 5-15% higher than experimental values. Which aspect of the ionic model most likely accounts for this systematic discrepancy?

  1. The ionic model assumes point charges, but real ions have finite size and charge distribution, leading to weaker actual interactions
  2. The experimental Born-Haber cycle method consistently underestimates lattice energies due to measurement errors in sublimation enthalpies
  3. The ionic model overestimates lattice energies because it doesn't account for partial covalent character in real ionic compounds (correct answer)
  4. The theoretical calculations fail to include van der Waals attractions, which would increase the predicted lattice energies further
Explanation: The systematic overestimation of lattice energies by the ionic model occurs because real 'ionic' compounds have some degree of covalent character, which makes the bonds somewhat weaker than pure electrostatic interactions would predict. The ionic model assumes complete electron transfer and purely electrostatic interactions, but real compounds show electron sharing that reduces the electrostatic contribution. Option A is incorrect because finite ion size is accounted for in Born-Landé calculations. Option B wrongly blames experimental error rather than model limitations. Option D is incorrect because van der Waals forces would be much smaller than lattice energies and wouldn't explain the systematic discrepancy.

Question 13

A student calculates the theoretical lattice energy of CuCl using the ionic model and assuming Cu⁺ and Cl⁻ ions. However, the experimental lattice energy is significantly lower than predicted. X-ray photoelectron spectroscopy reveals that the copper has some Cu²⁺ character. How does this observation relate to the limitations of the ionic model?

  1. The lower experimental lattice energy confirms that ionic interactions are weaker than the model predicts when transition metals are involved
  2. The presence of Cu²⁺ indicates that the compound is actually CuCl₂, and the student used the wrong chemical formula for calculations
  3. The ionic model correctly predicts behavior, but experimental techniques like XPS introduce artifacts that make Cu⁺ appear to have Cu²⁺ character
  4. The ionic model fails because it cannot account for mixed oxidation states that arise from partial electron transfer between ions (correct answer)
Explanation: When you encounter questions about lattice energy discrepancies involving transition metals, you're dealing with the fundamental limitations of the ionic bonding model. The ionic model assumes complete electron transfer and purely electrostatic interactions between discrete ions. The correct answer is D because the ionic model cannot account for the complex reality of transition metal bonding. CuCl is supposed to contain Cu⁺ and Cl⁻ ions, but XPS shows some Cu²⁺ character. This suggests that electron transfer isn't complete or uniform - some copper atoms have given up more electron density than others, creating a mixed oxidation state environment. The ionic model fails here because it treats bonding as black-and-white (complete transfer) rather than recognizing the continuum of electron sharing that actually occurs. Option A is incorrect because the issue isn't that ionic interactions are inherently weaker with transition metals - it's that the bonding isn't purely ionic at all. Option B misses the point entirely; the compound is still CuCl, not CuCl₂. The problem isn't the formula but the bonding model. Option C incorrectly dismisses experimental evidence as artifacts when XPS is actually revealing the true electronic structure that the ionic model cannot explain. Remember that transition metals often show significant covalent character in their "ionic" compounds due to their variable oxidation states and d-orbital involvement. When you see lattice energy discrepancies with transition metals, consider whether the ionic model's assumption of complete electron transfer actually holds true.

Question 14

Two compounds, KF and KI, both adopt the rock salt crystal structure. However, when their ionic conductivities are measured in the solid state at 500°C, KF shows negligible conductivity while KI shows measurable ionic conduction. According to the ionic model, both should behave identically. What factor best explains this difference?

  1. KI has weaker electrostatic attractions due to larger ionic radii, allowing greater ionic mobility at elevated temperatures (correct answer)
  2. KF exhibits more covalent character than KI, creating directional bonding that restricts ionic movement in the crystal lattice
  3. The larger size difference between K⁺ and I⁻ creates more lattice defects in KI, facilitating ionic conduction pathways
  4. KI undergoes partial decomposition at 500°C, creating vacancies that enhance ionic conductivity compared to the more stable KF
Explanation: The larger I⁻ ion results in longer K-I distances compared to K-F distances, leading to weaker electrostatic attractions in KI. According to Coulomb's law, force decreases with distance squared, so the weaker interactions in KI allow ions to move more easily at elevated temperatures, enabling ionic conduction. Option B is incorrect because KF actually has more ionic character than KI due to the larger electronegativity difference. Option C incorrectly suggests that size differences create defects. Option D assumes decomposition without evidence and doesn't relate to the ionic model.

Question 15

The lattice energy of MgO is significantly higher than that of NaCl, despite both compounds having the same crystal structure. If the ionic model perfectly described both compounds, which factor would be the primary reason for this difference?

  1. The difference in ionic radii, with smaller ions in MgO leading to shorter interionic distances and stronger attractions
  2. The higher charge density of the ions in MgO, resulting in electrostatic forces that are four times stronger than in NaCl (correct answer)
  3. The greater electronegativity difference between Mg and O compared to Na and Cl, leading to more complete electron transfer
  4. The difference in coordination numbers, with MgO having higher coordination leading to more favorable packing arrangements
Explanation: According to Coulomb's law (F∝q1q2r2F ∝ \frac{q_1q_2}{r^2}), the electrostatic force depends on the product of charges. In MgO, the charges are +2 and -2, giving a charge product of 4. In NaCl, the charges are +1 and -1, giving a charge product of 1. This makes the electrostatic attractions in MgO four times stronger than in NaCl, which is the primary factor affecting lattice energy. While ionic size differences (A) do matter, the charge effect is dominant. Both compounds have the same structure, eliminating (D), and electronegativity differences (C) don't directly determine lattice energy.

Question 16

Three compounds with similar lattice energies (MgO, CaO, and BaO) show dramatically different properties. MgO has moderate solubility in water and high melting point, CaO is more soluble with slightly lower melting point, and BaO is highly soluble with the lowest melting point. Which factor most clearly demonstrates a limitation of the ionic model in explaining these differences?

  1. The ionic model cannot account for how different ionic sizes affect both lattice energies and hydration energies
  2. The model fails to predict that larger ions will exhibit weaker electrostatic attractions and different solvation behavior (correct answer)
  3. The ionic model incorrectly assumes that compounds with similar lattice energies should have identical properties
  4. The model cannot explain why compounds with similar charges show different solubility behaviors in polar solvents
Explanation: The ionic model can predict that larger ions (Mg²⁺ < Ca²⁺ < Ba²⁺) will have weaker electrostatic attractions due to increased distance, leading to lower melting points and easier solvation. However, the model's limitation is in quantitatively predicting how these size effects will balance lattice energy and hydration energy to determine solubility. The trend from MgO to BaO shows increasing solubility as the larger cations become more easily solvated, but the ionic model struggles to predict the relative magnitudes of these competing effects.

Question 17

An element X from Group 2 of the periodic table reacts completely with an element Y from Group 16. What is the most likely empirical formula and bonding type of the resulting compound?

  1. XY, ionic (correct answer)
  2. XY₂, covalent
  3. X₂Y, ionic
  4. XY, covalent
Explanation: Elements from Group 2 (alkaline earth metals) typically lose two electrons to form cations with a 2+ charge (X²⁺). Elements from Group 16 (chalcogens) typically gain two electrons to form anions with a 2- charge (Y²⁻). The reaction is between a metal and a non-metal, so the bonding will be ionic. To form a neutral compound, one X²⁺ ion will balance one Y²⁻ ion. Therefore, the empirical formula of the compound is XY.

Question 18

Which of the following ionic compounds is predicted to have the lattice enthalpy with the largest magnitude?

  1. Lithium fluoride, LiF
  2. Sodium chloride, NaCl
  3. Magnesium oxide, MgO (correct answer)
  4. Calcium sulfide, CaS
Explanation: Lattice enthalpy is directly proportional to the product of the ionic charges and inversely proportional to the sum of the ionic radii. The product of charges for MgO (+2 and -2) is 4, and for CaS (+2 and -2) is also 4. For LiF and NaCl, the product of charges (+1 and -1) is 1. Since ionic charge is the dominant factor, MgO and CaS will have significantly larger lattice enthalpies than LiF and NaCl. Between MgO and CaS, Mg²⁺ is smaller than Ca²⁺ (period 3 vs period 4) and O²⁻ is smaller than S²⁻ (period 2 vs period 3). Therefore, the internuclear distance is smaller in MgO, resulting in the lattice enthalpy with the largest magnitude.

Question 19

Iron can form both Fe²⁺ and Fe³⁺ ions. What are the correct chemical formulas for iron(II) phosphate and iron(III) sulfate, respectively?

  1. FePO₄ and FeSO₄
  2. Fe₂(PO₄)₃ and Fe₃(SO₄)₂
  3. Fe₃(PO₄)₂ and Fe₂(SO₄)₃ (correct answer)
  4. FePO₄ and Fe₂(SO₄)₃
Explanation: To form a neutral compound, the total positive charge must balance the total negative charge. Phosphate is PO₄³⁻ and sulfate is SO₄²⁻. For iron(II) phosphate, we need to balance Fe²⁺ and PO₄³⁻. The lowest common multiple of 2 and 3 is 6. This requires three Fe²⁺ ions (3 x +2 = +6) and two PO₄³⁻ ions (2 x -3 = -6), giving the formula Fe₃(PO₄)₂. For iron(III) sulfate, we need to balance Fe³⁺ and SO₄²⁻. The lowest common multiple of 3 and 2 is 6. This requires two Fe³⁺ ions (2 x +3 = +6) and three SO₄²⁻ ions (3 x -2 = -6), giving the formula Fe₂(SO₄)₃.

Question 20

Why do ionic compounds like potassium bromide (KBr) typically have very low volatility compared to molecular compounds like methane (CH₄)?

  1. The intermolecular forces holding KBr formula units together are much stronger than the forces between CH₄ molecules.
  2. Strong electrostatic forces of attraction between oppositely charged ions in the lattice require a large amount of energy to overcome. (correct answer)
  3. KBr has a much higher molar mass than CH₄, which makes it more difficult for its particles to enter the gas phase.
  4. The triple covalent bonds within the KBr structure are significantly stronger than the single covalent bonds present in CH₄.
Explanation: Volatility is the tendency of a substance to vaporize. For an ionic compound to vaporize, the strong, omnidirectional electrostatic forces of attraction between the cations (K⁺) and anions (Br⁻) in the entire crystal lattice must be overcome. This requires a very large amount of thermal energy, resulting in low volatility and a high boiling point. In contrast, methane is a molecular substance with weak intermolecular forces (London dispersion forces) between its molecules, which are easily overcome. Distractor A uses imprecise terminology ('intermolecular forces' for an ionic lattice). Distractor C correctly states a fact about molar mass but misattributes it as the primary cause for the vast difference in volatility. Distractor D incorrectly describes KBr as having covalent bonds.