10.9: End of Chapter Problems

1. Describe the differences in behavior between NaOH and CH₃OH in aqueous solution. Which solution would be a better conductor of electricity? Explain your reasoning.
2. What is the relationship between the strength of the electrostatic attraction between oppositely charged ions and the distance between the ions? How does the strength of the electrostatic interactions change as the size of the ions increases?
3. Which will result in the release of more energy: the interaction of a gaseous sodium ion with a gaseous oxide ion or the interaction of a gaseous sodium ion with a gaseous bromide ion? Why?

Answer

The interaction of a sodium ion and an oxide ion. The electrostatic attraction energy between ions of opposite charge is directly proportional to the charge on each ion (Q₁ and Q₂ in Equation 9.1). Thus, more energy is released as the charge on the ions increases (assuming the internuclear distance does not increase substantially). A sodium ion has a +1 charge; an oxide ion, a −2 charge; and a bromide ion, a −1 charge. For the interaction of a sodium ion with an oxide ion, Q₁ = +1 and Q₂ = −2, whereas for the interaction of a sodium ion with a bromide ion, Q₁ = +1 and Q₂ = −1. The larger value of Q₁ × Q₂ for the sodium ion–oxide ion interaction means it will release more energy.

4. Which will result in the release of more energy: the interaction of a gaseous chloride ion with a gaseous sodium ion or a gaseous potassium ion? Explain your answer.
5. What are the predominant interactions when oppositely charged ions are
   1. far apart?
   2. at internuclear distances close to r₀?
   3. very close together (at a distance that is less than the sum of the ionic radii)?
6. Several factors contribute to the stability of ionic compounds. Describe one type of interaction that destabilizes ionic compounds. Describe the interactions that stabilize ionic compounds.
7. What is the relationship between the electrostatic attractive energy between charged particles and the distance between the particles?
8. How does the energy of the electrostatic interaction between ions with charges +1 and −1 compare to the interaction between ions with charges +3 and −1 if the distance between the ions is the same in both cases? How does this compare with the magnitude of the interaction between ions with +3 and −3 charges?

Answer

According to Equation 9.1, in the first case Q₁Q₂ = (+1)(−1) = −1; in the second case, Q₁Q₂ = (+3)(−1) = −3. Thus, E will be three times larger for the +3/−1 ions. For +3/−3 ions, Q₁Q₂ = (+3)(−3) = −9, so E will be nine times larger than for the +1/−1 ions.

9. How many grams of gaseous MgCl₂ are needed to give the same electrostatic attractive energy as 0.5 mol of gaseous LiCl? The ionic radii are Li⁺ = 76 pm, Mg⁺² = 72 pm, and Cl⁻ = 181 pm.
10. Sketch a diagram showing the relationship between potential energy and internuclear distance (from r = ∞ to r = 0) for the interaction of a bromide ion and a potassium ion to form gaseous KBr. Explain why the energy of the system increases as the distance between the ions decreases from r = r₀ to r = 0.

Answer

At r < r₀, the energy of the system increases due to electron–electron repulsions between the overlapping electron distributions on adjacent ions. At very short internuclear distances, electrostatic repulsions between adjacent nuclei also become important.

11. Calculate the magnitude of the electrostatic attractive energy (E, in kilojoules) for 85.0 g of gaseous SrS ion pairs. The observed internuclear distance in the gas phase is 244.05 pm.
12. What is the electrostatic attractive energy (E, in kilojoules) for 130 g of gaseous HgI₂? The internuclear distance is 255.3 pm.
13. Why is it incorrect to speak of a molecule of solid NaCl?

Answer

NaCl consists of discrete ions arranged in a crystal lattice, not covalently bonded molecules.

14. What information can you use to predict whether a bond between two atoms is covalent or ionic?

15. Predict which of the following compounds are ionic and which are covalent, based on the location of their constituent atoms in the periodic table:
    1. Cl₂CO
    2. MnO
    3. NCl₃
    4. CoBr₂
    5. K₂S
    6. CO
    7. CaF₂
    8. HI
    9. CaO
    10. IBr
    11. CO₂

Answer

ionic: (b), (d), (e), (g), and (i); covalent: (a), (c), (f), (h), (j), and (k)

16. Explain the difference between a nonpolar covalent bond, a polar covalent bond, and an ionic bond.

17. From its position in the periodic table, determine which atom in each pair is more electronegative:
    1. Br or Cl
    2. N or O
    3. S or O
    4. P or S
    5. Si or N
    6. Ba or P
    7. N or K

Answer

Cl; O; O; S; N; P; N

18. From its position in the periodic table, determine which atom in each pair is more electronegative:
    1. N or P
    2. N or Ge
    3. S or F
    4. Cl or S
    5. H or C
    6. Se or P
    7. C or Si

19. From their positions in the periodic table, arrange the atoms in each of the following series in order of increasing electronegativity:
    1. C, F, H, N, O
    2. Br, Cl, F, H, I
    3. F, H, O, P, S
    4. Al, H, Na, O, P
    5. Ba, H, N, O, As

Answer

H, C, N, O, F; H, I, Br, Cl, F; H, P, S, O, F; Na, Al, H, P, O; Ba, H, As, N, O

20. From their positions in the periodic table, arrange the atoms in each of the following series in order of increasing electronegativity:
    1. As, H, N, P, Sb
    2. Cl, H, P, S, Si
    3. Br, Cl, Ge, H, Sr
    4. Ca, H, K, N, Si
    5. Cl, Cs, Ge, H, Sr
21. Which atoms can bond to sulfur so as to produce a positive partial charge on the sulfur atom?

N, O, F, and Cl

22. Which is the most polar bond?
    1. C–C
    2. C–H
    3. N–H
    4. O–H
    5. Se–H

23. Identify the more polar bond in each of the following pairs of bonds:
    1. HF or HCl
    2. NO or CO
    3. SH or OH
    4. PCl or SCl
    5. CH or NH
    6. SO or PO
    7. CN or NN

Answer

HF; CO; OH; PCl; NH; PO; CN

24. Which of the following molecules or ions contain polar bonds?
    1. O₃
    2. S₈
    3. \(\ce{O2^2-}\)
    4. \(\ce{NO3-}\)
    5. CO₂
    6. H₂S
    7. \(\ce{BH4-}\)

25. Write the Lewis symbols for each of the following ions:
    1. As^3–
    2. I^–
    3. Be²⁺
    4. O^2–
    5. Ga³⁺
    6. Li⁺
    7. N^3–

Answer

eight electrons:

eight electrons:

no electrons

Be²⁺;

eight electrons:

no electrons

Ga³⁺;

no electrons

Li⁺;

eight electrons:

26. Many monatomic ions are found in seawater, including the ions formed from the following list of elements. Write the Lewis symbols for the monatomic ions formed from the following elements:
    1. Cl
    2. Na
    3. Mg
    4. Ca
    5. K
    6. Br
    7. Sr
    8. F

27. Write the Lewis symbols of the ions in each of the following ionic compounds and the Lewis symbols of the atom from which they are formed:
    1. MgS
    2. Al₂O₃
    3. GaCl₃
    4. K₂O
    5. Li₃N
    6. KF

Answer

(a)

;

(b)

;

(c)

;

(d)

;

(e)

;

(f)

28. In the Lewis structures listed here, M and X represent various elements in the third period of the periodic table. Write the formula of each compound using the chemical symbols of each element:
    1. 
    2. 
    3. 
    4. 

29. Write the Lewis structure for the diatomic molecule P₂, an unstable form of phosphorus found in high-temperature phosphorus vapor.

30. Write Lewis structures for the following:
    1. H₂
    2. HBr
    3. PCl₃
    4. SF₂
    5. H₂CCH₂
    6. HNNH
    7. H₂CNH
    8. NO^–
    9. N₂
    10. CO
    11. CN^–

31. Write Lewis structures for the following:
    1. O₂
    2. H₂CO
    3. AsF₃
    4. ClNO
    5. SiCl₄
    6. H₃O⁺
    7. \(\ce{NH4+}\)
    8. (h) \(\ce{BF4-}\)
    9. (i) HCCH
    10. (j) ClCN
    11. (k) \(\ce{C2^2+}\)

Answer

(a)

In this case, the Lewis structure is inadequate to depict the fact that experimental studies have shown two unpaired electrons in each oxygen molecule.

(b)

;

(c)

;

(d)

;

(e)

;

(f)

;

(g)

;

(h)

;

(i)

;

(j)

;

(k)

32. Write Lewis structures for the following:
    1. ClF₃
    2. PCl₅
    3. BF₃
    4. \(\ce{PF6-}\)

33. Write Lewis structures for the following:
    1. SeF₆
    2. XeF₄
    3. \(\ce{SeCl3+}\)
    4. Cl₂BBCl₂ (contains a B–B bond)

Answer

SeF₆:

;

XeF₄:

;

\(\ce{SeCl3+}\):

;

Cl₂BBCl₂:

34. Write Lewis structures for:
    1. \(\ce{PO4^3-}\)
    2. \(\ce{ICl4-}\)
    3. \(\ce{SO3^2-}\)
    4. HONO

35. Correct the following statement: “The bonds in solid PbCl₂ are ionic; the bond in a HCl molecule is covalent. Thus, all of the valence electrons in PbCl₂ are located on the Cl^– ions, and all of the valence electrons in a HCl molecule are shared between the H and Cl atoms.”

Answer

Two valence electrons per Pb atom are transferred to Cl atoms; the resulting Pb²⁺ ion has a 6s² valence shell configuration. Two of the valence electrons in the HCl molecule are shared, and the other six are located on the Cl atom as lone pairs of electrons.

36. Write Lewis structures for the following molecules or ions:
    1. SbH₃
    2. XeF₂
    3. Se₈ (a cyclic molecule with a ring of eight Se atoms)

37. Methanol, H₃COH, is used as the fuel in some race cars. Ethanol, C₂H₅OH, is used extensively as motor fuel in Brazil. Both methanol and ethanol produce CO₂ and H₂O when they burn. Write the chemical equations for these combustion reactions using Lewis structures instead of chemical formulas.

Answer

38. Many planets in our solar system contain organic chemicals including methane (CH₄) and traces of ethylene (C₂H₄), ethane (C₂H₆), propyne (H₃CCCH), and diacetylene (HCCCCH). Write the Lewis structures for each of these molecules.
39. Carbon tetrachloride was formerly used in fire extinguishers for electrical fires. It is no longer used for this purpose because of the formation of the toxic gas phosgene, Cl₂CO. Write the Lewis structures for carbon tetrachloride and phosgene.

Answer

40. Identify the atoms that correspond to each of the following electron configurations. Then, write the Lewis symbol for the common ion formed from each atom:
    1. 1s²2s²2p⁵
    2. 1s²2s²2p⁶3s²
    3. 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰
    4. 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p⁴
    5. 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰4p¹
41. The arrangement of atoms in several biologically important molecules is given here. Complete the Lewis structures of these molecules by adding multiple bonds and lone pairs. Do not add any more atoms.

the amino acid serine:

urea:

pyruvic acid:

uracil:

carbonic acid:

Answer

(a)

;

(b)

;

(c)

;

(d)

;

(e)

42. A compound with a molar mass of about 28 g/mol contains 85.7% carbon and 14.3% hydrogen by mass. Write the Lewis structure for a molecule of the compound.
43. A compound with a molar mass of about 42 g/mol contains 85.7% carbon and 14.3% hydrogen by mass. Write the Lewis structure for a molecule of the compound.

44. Two arrangements of atoms are possible for a compound with a molar mass of about 45 g/mol that contains 52.2% C, 13.1% H, and 34.7% O by mass. Write the Lewis structures for the two molecules.
45. How are single, double, and triple bonds similar? How do they differ?

Answer

Each bond includes a sharing of electrons between atoms. Two electrons are shared in a single bond; four electrons are shared in a double bond; and six electrons are shared in a triple bond.

46. Write resonance forms that describe the distribution of electrons in each of these molecules or ions.
    1. selenium dioxide, OSeO
    2. nitrate ion, \(\ce{NO3-}\)
    3. nitric acid, HNO₃ (N is bonded to an OH group and two O atoms)
    4. benzene, C₆H₆:

47. Write resonance forms that describe the distribution of electrons in each of these molecules or ions.
    1. sulfur dioxide, SO₂
    2. carbonate ion, \(\ce{CO3^2-}\)
    3. hydrogen carbonate ion, \(\ce{HCO3-}\) (C is bonded to an OH group and two O atoms)
    4. pyridine:

5. the allyl ion:

Answer

(a)

;

(b)

;

(c)

;

(d)

;

(e)

48. Write the resonance forms of ozone, O₃, the component of the upper atmosphere that protects the Earth from ultraviolet radiation.
49. Sodium nitrite, which has been used to preserve bacon and other meats, is an ionic compound. Write the resonance forms of the nitrite ion, \(\ce{NO2-}\).

50. In terms of the bonds present, explain why acetic acid, CH₃CO₂H, contains two distinct types of carbon-oxygen bonds, whereas the acetate ion, formed by loss of a hydrogen ion from acetic acid, only contains one type of carbon-oxygen bond. The skeleton structures of these species are shown:

51. Write the Lewis structures for the following, and include resonance structures where appropriate. Indicate which has the strongest carbon-oxygen bond.
    1. CO₂
    2. CO

Answer

(a)

(b)

CO has the strongest carbon-oxygen bond because there is a triple bond joining C and O. CO₂ has double bonds.

52. Toothpastes containing sodium hydrogen carbonate (sodium bicarbonate) and hydrogen peroxide are widely used. Write Lewis structures for the hydrogen carbonate ion and hydrogen peroxide molecule, with resonance forms where appropriate.
53. Determine the formal charge of each element in the following:
    1. HCl
    2. CF₄
    3. PCl₃
    4. PF₅

Answer

H: 0, Cl: 0; C: 0, F: 0; P: 0, Cl 0; P: 0, F: 0

54. Determine the formal charge of each element in the following:
    1. H₃O⁺
    2. \(\ce{SO4^2-}\)
    3. NH₃
    4. \(\ce{O2^2-}\)
    5. H₂O₂

55. Calculate the formal charge of chlorine in the molecules Cl₂, BeCl₂, and ClF₅.

Answer

Cl in Cl₂: 0; Cl in BeCl₂: 0; Cl in ClF₅: 0

56. Calculate the formal charge of each element in the following compounds and ions:
    1. F₂CO
    2. NO^–
    3. \(\ce{BF4-}\)
    4. \(\ce{SnCl3-}\)
    5. H₂CCH₂
    6. ClF₃
    7. SeF₆
    8. \(\ce{PO4^3-}\)

57. Draw all possible resonance structures for each of these compounds. Determine the formal charge on each atom in each of the resonance structures:
    1. O₃
    2. SO₂
    3. \(\ce{NO2-}\)
    4. \(\ce{NO3-}\)

Answer

(a)

(b)

(c)

(d)

58. Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in nitrosyl chloride: ClNO or ClON?
59. Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in hypochlorous acid: HOCl or OClH?

Answer

HOCl

60. Based on formal charge considerations, which of the following would likely be the correct arrangement of atoms in sulfur dioxide: OSO or SOO?
61. Draw the structure of hydroxylamine, H₃NO, and assign formal charges; look up the structure. Is the actual structure consistent with the formal charges?

The structure that gives zero formal charges is consistent with the actual structure:

62. Iodine forms a series of fluorides (listed here). Write Lewis structures for each of the four compounds and determine the formal charge of the iodine atom in each molecule:
    1. IF
    2. IF₃
    3. IF₅
    4. IF₇
63. Write the Lewis structure and chemical formula of the compound with a molar mass of about 70 g/mol that contains 19.7% nitrogen and 80.3% fluorine by mass, and determine the formal charge of the atoms in this compound.

Answer

NF₃;

64. Which of the following structures would we expect for nitrous acid? Determine the formal charges:

65. Sulfuric acid is the industrial chemical produced in greatest quantity worldwide. About 90 billion pounds are produced each year in the United States alone. Write the Lewis structure for sulfuric acid, H₂SO₄, which has two oxygen atoms and two OH groups bonded to the sulfur.

Answer

66. Which bond in each of the following pairs of bonds is the strongest?
    1. C–C or \(\mathrm{C=C}\)
    2. C–N or \(\mathrm{C≡N}\)
    3. \(\mathrm{C≡O}\) or \(\mathrm{C=O}\)
    4. H–F or H–Cl
    5. C–H or O–H
    6. C–N or C–O

67. Using the bond energies in Table, determine the approximate enthalpy change for each of the following reactions:
    1. \(\ce{H2}(g)+\ce{Br2}(g)⟶\ce{2HBr}(g)\)
    2. \(\ce{CH4}(g)+\ce{I2}(g)⟶\ce{CH3I}(g)+\ce{HI}(g)\)
    3. (c) \(\ce{C2H4}(g)+\ce{3O2}(g)⟶\ce{2CO2}(g)+\ce{2H2O}(g)\)

68. Using the bond energies in Table, determine the approximate enthalpy change for each of the following reactions:
    1. \(\ce{H2}(g)+\ce{Br2}(g)⟶\ce{2HBr}(g)\)
    2. \(\ce{CH4}(g)+\ce{I2}(g)⟶\ce{CH3I}(g)+\ce{HI}(g)\)
    3. \(\ce{C2H4}(g)+\ce{3O2}(g)⟶\ce{2CO2}(g)+\ce{2H2O}(g)\)

Answer

(a) −114 kJ;

(b) 30 kJ;

(c) −1055 kJ

69. Using the bond energies in Table, determine the approximate enthalpy change for each of the following reactions:
    1. \(\ce{Cl2}(g)+\ce{3F2}(g)⟶\ce{2ClF3}(g)\)
    2. \(\mathrm{H_2C=CH_2}(g)+\ce{H2}(g)⟶\ce{H3CCH3}(g)\)
    3. (c) \(\ce{2C2H6}(g)+\ce{7O2}(g)⟶\ce{4CO2}(g)+\ce{6H2O}(g)\)
70. When a molecule can form two different structures, the structure with the stronger bonds is usually the more stable form. Use bond energies to predict the correct structure of the hydroxylamine molecule:

Answer

The greater bond energy is in the figure on the left. It is the more stable form.

71. How does the bond energy of HCl differ from the standard enthalpy of formation of HCl(g)?
72. Using the standard enthalpy of formation data in Appendix G, show how the standard enthalpy of formation of HCl(g) can be used to determine the bond energy.

Answer

\(\ce{HCl}(g)⟶\dfrac{1}{2}\ce{H2}(g)+\dfrac{1}{2}\ce{Cl2}(g)\hspace{20px}ΔH^\circ_1=−ΔH^\circ_{\ce f[\ce{HCl}(g)]}\\
\dfrac{1}{2}\ce{H2}(g)⟶\ce{H}(g)\hspace{105px}ΔH^\circ_2=ΔH^\circ_{\ce f[\ce H(g)]}\\
\underline{\dfrac{1}{2}\ce{Cl2}(g)⟶\ce{Cl}(g)\hspace{99px}ΔH^\circ_3=ΔH^\circ_{\ce f[\ce{Cl}(g)]}}\\
\ce{HCl}(g)⟶\ce{H}(g)+\ce{Cl}(g)\hspace{58px}ΔH^\circ_{298}=ΔH^\circ_1+ΔH^\circ_2+ΔH^\circ_3\)

\(\begin{align}
D_\ce{HCl}=ΔH^\circ_{298}&=ΔH^\circ_{\ce f[\ce{HCl}(g)]}+ΔH^\circ_{\ce f[\ce H(g)]}+ΔH^\circ_{\ce f[\ce{Cl}(g)]}\\
&=\mathrm{−(−92.307\:kJ)+217.97\:kJ+121.3\:kJ}\\
&=\mathrm{431.6\:kJ}
\end{align}\)

73. Using the standard enthalpy of formation data in Appendix G, calculate the bond energy of the carbon-sulfur double bond in CS₂.
74. Using the standard enthalpy of formation data in Appendix G, determine which bond is stronger: the S–F bond in SF₄(g) or in SF₆(g)?

Answer

The S–F bond in SF₄ is stronger.

75. Using the standard enthalpy of formation data in Appendix G, determine which bond is stronger: the P–Cl bond in PCl₃(g) or in PCl₅(g)?
76. Complete the following Lewis structure by adding bonds (not atoms), and then indicate the longest bond:

Answer

The C–C single bonds are longest.

77. Use the bond energy to calculate an approximate value of ΔH for the following reaction. Which is the more stable form of FNO₂?

78. Use principles of atomic structure to answer each of the following:¹

1. The radius of the Ca atom is 197 pm; the radius of the Ca²⁺ ion is 99 pm. Account for the difference.
2. The lattice energy of CaO(s) is –3460 kJ/mol; the lattice energy of K₂O is –2240 kJ/mol. Account for the difference.
3. Given these ionization values, explain the difference between Ca and K with regard to their first and second ionization energies.

Element	First Ionization Energy (kJ/mol)	Second Ionization Energy (kJ/mol)
K	419	3050
Ca	590	1140

4. The first ionization energy of Mg is 738 kJ/mol and that of Al is 578 kJ/mol. Account for this difference.

Answer

When two electrons are removed from the valence shell, the Ca radius loses the outermost energy level and reverts to the lower n = 3 level, which is much smaller in radius. The +2 charge on calcium pulls the oxygen much closer compared with K, thereby increasing the lattice energy relative to a less charged ion. (c) Removal of the 4s electron in Ca requires more energy than removal of the 4s electron in K because of the stronger attraction of the nucleus and the extra energy required to break the pairing of the electrons. The second ionization energy for K requires that an electron be removed from a lower energy level, where the attraction is much stronger from the nucleus for the electron. In addition, energy is required to unpair two electrons in a full orbital. For Ca, the second ionization potential requires removing only a lone electron in the exposed outer energy level. In Al, the removed electron is relatively unprotected and unpaired in a p orbital. The higher energy for Mg mainly reflects the unpairing of the 2s electron.

79. The lattice energy of LiF is 1023 kJ/mol, and the Li–F distance is 200.8 pm. NaF crystallizes in the same structure as LiF but with a Na–F distance of 231 pm. Which of the following values most closely approximates the lattice energy of NaF: 510, 890, 1023, 1175, or 4090 kJ/mol? Explain your choice.
80. For which of the following substances is the least energy required to convert one mole of the solid into separate ions?
    1. MgO
    2. SrO
    3. (c) KF
    4. CsF
    5. MgF₂
81. The reaction of a metal, M, with a halogen, X₂, proceeds by an exothermic reaction as indicated by this equation: \(\ce{M}(s)+\ce{X2}(g)⟶\ce{MX2}(s)\). For each of the following, indicate which option will make the reaction more exothermic. Explain your answers.
    1. a large radius vs. a small radius for M⁺²
    2. a high ionization energy vs. a low ionization energy for M
    3. an increasing bond energy for the halogen
    4. a decreasing electron affinity for the halogen
    5. an increasing size of the anion formed by the halogen
82. The lattice energy of LiF is 1023 kJ/mol, and the Li–F distance is 201 pm. MgO crystallizes in the same structure as LiF but with a Mg–O distance of 205 pm. Which of the following values most closely approximates the lattice energy of MgO: 256 kJ/mol, 512 kJ/mol, 1023 kJ/mol, 2046 kJ/mol, or 4008 kJ/mol? Explain your choice.

Answer

4008 kJ/mol; both ions in MgO have twice the charge of the ions in LiF; the bond length is very similar and both have the same structure; a quadrupling of the energy is expected based on the equation for lattice energy

83. Which compound in each of the following pairs has the larger lattice energy? Note: Mg²⁺ and Li⁺ have similar radii; O^2– and F^– have similar radii. Explain your choices.
    1. MgO or MgSe
    2. LiF or MgO
    3. (c) Li₂O or LiCl
    4. Li₂Se or MgO
84. Which compound in each of the following pairs has the larger lattice energy? Note: Ba²⁺ and K⁺ have similar radii; S^2– and Cl^– have similar radii. Explain your choices.
    1. K₂O or Na₂O
    2. K₂S or BaS
    3. (c) KCl or BaS
    4. BaS or BaCl₂

Answer

Na₂O; Na⁺ has a smaller radius than K⁺; BaS; Ba has a larger charge than K; (c) BaS; Ba and S have larger charges; BaS; S has a larger charge

85. Which of the following compounds requires the most energy to convert one mole of the solid into separate ions?
    1. MgO
    2. SrO
    3. KF
    4. CsF
    5. MgF₂
86. Which of the following compounds requires the most energy to convert one mole of the solid into separate ions?
    1. K₂S
    2. K₂O
    3. CaS
    4. Cs₂S
    5. CaO
87. The lattice energy of KF is 794 kJ/mol, and the interionic distance is 269 pm. The Na–F distance in NaF, which has the same structure as KF, is 231 pm. Which of the following values is the closest approximation of the lattice energy of NaF: 682 kJ/mol, 794 kJ/mol, 924 kJ/mol, 1588 kJ/mol, or 3175 kJ/mol? Explain your answer.