Understand The Ionic Model - IB Chemistry

1. 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?
   Answer: Magnesium oxide, MgO 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.
2. Which of the following ionic compounds is predicted to have the lattice enthalpy with the largest magnitude?
   Answer: Magnesium oxide, MgO 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.
3. 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?
   Answer: There is a significant degree of covalent character in the bonding, which adds to the overall lattice strength. 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.
4. In the crystal lattice of an ionic compound, what is meant by the term 'coordination number' of an ion?
   Answer: The number of nearest-neighboring ions of opposite charge that surround a specific ion. The coordination number in an ionic crystal lattice is a key feature of its geometry. It specifically refers to the number of ions of opposite charge that are in direct contact with (i.e., are the nearest neighbors of) a given central ion. For example, in the sodium chloride lattice, each Na⁺ ion is surrounded by 6 Cl⁻ ions, and each Cl⁻ ion is surrounded by 6 Na⁺ ions, so the coordination number for both ions is 6. The other options describe different concepts related to ionic compounds but do not define coordination number.
5. A neutral atom of an element has the electron configuration 1s²2s²2p⁶3s²3p⁴. What is the charge on the most stable ion it typically forms, and what is the reason?
   Answer: 2−, because gaining two electrons fills the 3p sublevel and achieves a noble gas configuration. The element with this electron configuration is sulfur, which is in Group 16. It has six valence electrons (3s²3p⁴). The most stable ion is formed by achieving a noble gas electron configuration, which is an octet of valence electrons. For sulfur, the most energetically favorable way to do this is to gain two electrons to fill its 3p sublevel, forming the configuration 1s²2s²2p⁶3s²3p⁶ (same as argon). This results in a sulfide ion with a 2− charge (S²⁻). Losing six electrons (B) would require a prohibitively large amount of energy (sum of six ionization energies).
6. A student dissolves solid magnesium chloride, MgCl₂, in distilled water. Which statement best describes the particles present in the resulting aqueous solution?
   Answer: The solution contains hydrated Mg²⁺ ions and hydrated Cl⁻ ions, which are mobile. When a soluble ionic compound like MgCl₂ dissolves in water, it dissociates into its constituent ions. The solid lattice breaks down, and the individual ions are surrounded by water molecules in a process called hydration. The solution will contain mobile, hydrated magnesium ions (Mg²⁺(aq)) and hydrated chloride ions (Cl⁻(aq)). The mobility of these ions allows the solution to conduct electricity. Distractor A is incorrect because dissociation occurs. Distractor C shows incorrect ion charges. Distractor D is incorrect because the solution is conducting and the ions are solvated, not neutralized.
7. 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?
   Answer: Caesium fluoride, CsF 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.
8. 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?
   Answer: 2- 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⁻.
9. Ammonium nitrate, NH₄NO₃, is an ionic compound. Which statement best explains its electrical conductivity?
   Answer: It conducts electricity when molten or in aqueous solution because its ions become mobile charge carriers. 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.
10. 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?
    Answer: XY, ionic 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.
11. Consider the compounds NaF and MgO. The ions Na⁺, F⁻, Mg²⁺, and O²⁻ are all isoelectronic. Which statement best explains the significant difference in their lattice enthalpies?
    Answer: The lattice enthalpy of MgO has a much larger magnitude because the product of the ionic charges (+2 × −2) is four times greater than that of NaF (+1 × −1). Lattice enthalpy is principally determined by two factors: ionic charge and ionic radius, as described by Coulomb's Law. The force of attraction is proportional to the product of the charges (q₁q₂) and inversely proportional to the square of the distance between them. The product of charges in MgO is (−2) × (+2) = −4, while in NaF it is (−1) × (+1) = −1. This four-fold increase in charge product makes the electrostatic attraction in the MgO lattice much stronger, leading to a significantly larger magnitude of lattice enthalpy. While radii also play a role, the effect of charge is dominant.
12. Iron can form both Fe²⁺ and Fe³⁺ ions. What are the correct chemical formulas for iron(II) phosphate and iron(III) sulfate, respectively?
    Answer: Fe₃(PO₄)₂ and Fe₂(SO₄)₃ 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.
13. What is the correct IUPAC name for the compound with the formula (NH₄)₂CrO₄?
    Answer: Ammonium chromate This is an ionic compound composed of a polyatomic cation and a polyatomic anion. The cation is NH₄⁺, which is named ammonium. The anion is CrO₄²⁻, which is named chromate. When naming ionic compounds, the cation is named first, followed by the anion. Prefixes indicating the number of ions (like 'di-') are not used. Roman numerals are only used for cations (usually transition metals) that can have multiple oxidation states, which is not the case for ammonium. Therefore, the correct name is simply ammonium chromate.
14. 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?
    Answer: The ionic model overestimates lattice energies because it doesn't account for partial covalent character in real ionic compounds 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.
15. 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?
    Answer: The ionic model fails because it cannot account for mixed oxidation states that arise from partial electron transfer between ions 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.
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?
    Answer: The model fails to predict that larger ions will exhibit weaker electrostatic attractions and different solvation behavior 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.
17. 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?
    Answer: The ions have similar sizes, allowing efficient packing with maximum electrostatic attractions and minimal repulsions 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.
18. 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?
    Answer: The compound exhibits partial covalent character due to polarization effects, indicating limitations of the pure ionic model 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.
19. 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?
    Answer: Li⁺ has higher charge density, creating stronger ion-dipole interactions with water molecules and more favorable hydration 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 $\frac{charge}{distance^2}$, 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.
20. Which of the following substances consists of positive and negative ions held together by electrostatic forces in its solid state?
    Answer: Potassium sulfate, K₂SO₄ An ionic solid consists of ions held in a crystal lattice by electrostatic forces. Potassium sulfate, K₂SO₄, is an ionic compound formed from potassium cations (K⁺) and sulfate polyatomic anions (SO₄²⁻). Silicon dioxide (A) is a covalent network solid. Sucrose (B) is a molecular covalent solid. Brass (D) is a metallic alloy characterized by metallic bonding (a lattice of cations in a sea of delocalized electrons).

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