22.E: Exercises

22.1: Periodic Trends in Bonding

Problems

1. Can an oxide be neither acidic nor basic?
2. \(Rb + O_2\: (excess) \rightarrow \:?\)
3. \(Na + O_2 \rightarrow \:?\)
4. BaO₂ is which of the following: hydroxide, peroxide, or superoxide?
5. What is an amphoteric solution?
6. Why is it difficult to obtain oxygen directly from water?

Solutions

1. Yes, an example is carbon monoxide (CO). CO doesn’t produce a salt when reacted with an acid or a base.
2. \( Rb + O_2 \; (excess) \rightarrow RbO_2 \)
   With the presence of excess oxygen, Rubidium forms a superoxide. Please review section regarding basic oxides above for more detail.
3. \( 2 Na + O_2 \rightarrow Na_2O \)
   Note: The problem does not specify that the oxygen was in excess, so it cannot be a peroxide. Please review section regarding basic oxides for more detail.
4. BaO₂ is a peroxide. Barium has an oxidation state of +2 so the oxygen atoms have oxidation state of -1. As a result, the compound is a peroxide, but more specifically referred to as barium peroxide.
5. An amphoteric solution is a substance that can chemically react as either acid or base. See section above on Properties of Amphoteric Oxides for more detail.
6. Water as such is a neutral stable molecule. It is difficult to break the covalent O-H bonds easily. Hence, electrical energy through the electrolysis process is applied to separate dioxygen from water. When a small amount of acid is added to water ionization is initiated which helps in electrochemical reactions as follows.

   \[[H_2O\:(acidulated)\rightleftharpoons H^+\,(aq)+OH]^-\times4\]

   At cathode:

   \[[H^+\,(aq)+e^-\rightarrow\dfrac{1}{2}H_2(g)]\times4\]

   At anode:

   \[4OH^-\,(aq)\rightarrow O_2+2H_2O + 4e^-\]

   Net reaction:

   \[2H_2O \xrightarrow{\large{electrolysis}} 2H_2\,(g) + O_2\,(g)\]

   Oxygen can thus be obtained from acidified water by its electrolysis.

22.2: Group 18: The Noble Gases

Conceptual Problems

1. The chemistry of the noble gases is largely dictated by a balance between two competing properties. What are these properties? How do they affect the reactivity of these elements?
2. Of the group 18 elements, only krypton, xenon, and radon form stable compounds with other atoms and then only with very electronegative elements. Why?
3. Give the type of hybrid orbitals used by xenon in each species.

1. XeF₂
2. XeF₄
3. XeO₃
4. XeOF₄
5. XeO₄
6. XeO₆⁴⁻

4. Which element is the least metallic—B, Ga, Tl, Pb, Ne, or Ge?
5. Of Br, N, Ar, Bi, Se, He, and S, which would you expect to form positive ions most easily? negative ions most easily?
6. Of BCl₃, BCl₄⁻, CH₄, H₃N·BF₃, PCl₃, PCl₅, XeO₃, H₂O, and F⁻, which species do you expect to be

1. electron donors?
2. electron acceptors?
3. neither electron donors nor acceptors?
4. both electron donors and acceptors?

7. Of HCl, HClO₄, HBr, H₂S, HF, KrF₂, and PH₃, which is the strongest acid?
8. Of CF₄, NH₃, NF₃, H₂O, OF₂, SiF₄, H₂S, XeF₄, and SiH₄, which is the strongest base?

Structure and Reactivity

1. Write a balanced chemical equation showing how you would prepare each compound from its elements and other common compounds.
   1. XeF₂
   2. XeF₄
   3. XeF₆
   4. XeOF₄
   5. XeO₃

2. Write a balanced chemical equation showing how you would make each compound.
   1. XeF₂ from Xe gas
   2. NaXeF₇ from its elements
   3. RnO₃ from Rn

3. In an effort to synthesize XeF₆, a chemist passed fluorine gas through a glass tube containing xenon gas. However, the product was not the one expected. What was the actual product?
4. Write a balanced chemical equation to describe the reaction of each species with water.

1. B₂H₆
2. F₂
3. C⁴⁺

5. Using heavy water (D₂O) as the source of deuterium, how could you prepare each compound?
   1. LiAlD₄
   2. D₂SO₄
   3. SiD₄
   4. DF
6. Predict the product(s) of each reaction and write a balanced chemical equation for each reaction.
   1. Al₂O₃(s) in OH⁻(aq)
   2. Ar(g) + F₂(g)
   3. PI₃(s) + H₂O(l)
   4. H₃PO₃(l) + OH⁻(aq)
   5. Bi(s) + excess Br₂(l)

Answers

1. Xe(g) + F₂(g) → XeF₂(s)
2. Xe(g) + 2F₂(g) → XeF₄(s)
3. Xe(g) + 3F₂(g) → XeF₆(s)
4. 2XeF₆(s) + SiO₂(s) → 2XeOF₄(l) + SiF₄(l)
5. XeF₆(s) + 3H₂O(l) → XeO₃(s) + 6HF(aq)

3. SiF₄; SiO₂(s) + 2F₂(g) → SiF₄(l) + 2O₂(g)

5. 1. 2Na(s) + 2D₂O(l) → D₂(g) + 2NaOD(aq)

2Li(s) + D₂(g) → 2LiD(s)

4LiD(s) + AlCl₃(soln) → LiAlD₄(s) + 3LiCl(soln)

2. D₂O(l) + SO₃(g) → D₂SO₄(l)
3. SiCl₄(l) + LiAlD₄(s) [from part (a)] → SiD₄(g) + LiCl(s) + AlCl₃(s)
4. CaF₂(s) + D₂SO₄(l) [from part (b)] → 2DF(g) + CaSO₄(s)

22.3: Group 17: The Halogens

Conceptual Problems

1. The lightest elements of groups 15, 16, and 17 form unusually weak single bonds. Why are their bonds so weak?
2. Fluorine has an anomalously low F–F bond energy. Why? Why does fluorine form compounds only in the −1 oxidation state, whereas the other halogens exist in multiple oxidation states?
3. Compare AlI₃, InCl₃, GaF₃, and LaBr₃ with respect to the type of M–X bond formed, melting point, and solubility in nonpolar solvents.
4. What are the formulas of the interhalogen compounds that will most likely contain the following species in the indicated oxidation states: I (+3), Cl (+3), I (−1), Br (+5)?
5. Consider this series of bromides: AlBr₃, SiBr₄, and PBr₅. Does the ionic character of the bond between the Br atoms and the central atom decrease or increase in this series?
6. Chromium forms compounds in the +6, +3, and +2 oxidation states. Which halogen would you use to produce each oxidation state? Justify your selections.
7. Of ClF₇, BrF₅, IF₇, BrF₃, ICl₃, IF₃, and IF₅, which one is least likely to exist? Justify your selection.

Answers

1. Electrostatic repulsions between lone pairs on adjacent atoms decrease bond strength.

5. Ionic character decreases as Δχ decreases from Al to P.

7. ClF₇

Structure and Reactivity

1. SiF₄ reacts easily with NaF to form SiF₆²⁻. In contrast, CF₄ is totally inert and shows no tendency to form CF₆²⁻ under even extreme conditions. Explain this difference.
2. Predict the products of each reaction and then balance each chemical equation.

1. Xe(g) + excess F₂(g) →
2. Se(s) + Cl₂(g) →
3. SO₂(g) + Br₂(g) →
4. NaBH₄(s) + BF₃(soln) →

3. Write a balanced chemical equation for the reaction of aqueous HF with
   1. SiO₂.
   2. Na₂CO₃.
   3. CaO.

4. Oxyhalides of sulfur, such as the thionyl halides (SOX₂, where X is F, Cl, or Br), are well known. Because the thionyl halides react vigorously with trace amounts of water, they are used for dehydrating hydrated metal salts. Write a balanced chemical equation to show the products of reaction of SOCl₂ with water.

5. Write a balanced chemical equation describing each reaction.
   1. the burning of sulfur in a chlorine atmosphere
   2. the dissolution of iodine in a potassium iodide solution
   3. the hydrolysis of PCl₃
   4. the preparation of HF from calcium fluoride and sulfuric acid
   5. the thermal decomposition of KClO₃
   6. the oxidation of sulfide ion by elemental iodine

6. Write the complete Lewis electron structure, the type of hybrid used by the central atom, and the number of lone pair electrons present on the central atom for each compound.
   1. CF₄
   2. PCl₃
   3. XeF₄

Answers

1. Carbon has no low energy d orbitals that can be used to form a set of d²sp³ hybrid orbitals. It is also so small that it is impossible for six fluorine atoms to fit around it at a distance that would allow for formation of strong C–F bonds.

1. SiO₂(s) + 6HF(aq) → SiF₆²⁻(aq) + 2H⁺(aq) + 2H₂O(l)
2. Na₂CO₃(s) + 2HF(aq) → CO₂(g) + 2NaF(aq) + H₂O(l)
3. CaO(s) + 2HF(aq) → CaF₂(s) + H₂O(l)

1. S₈(s) + 4Cl₂(g) → 4S₂Cl₂(l)
2. I₂(s) + KI(aq) → I₃⁻(aq) + K⁺(aq)
3. PCl₃(l) + 3H₂O(l) → H₃PO₃(aq) + 3HCl(aq)
4. CaF₂(s) + H₂SO₄(aq) → 2HF(aq) + CaSO₄(s)
5. 2KClO₃(s) \(\xrightarrow{\Delta}\) 2KCl(s) + 3O₂(g)
6. 8S²⁻(aq) + 8I₂(aq) → S₈(s) + 16I⁻(aq)

22.4: Group 16: The Oxygen Family

Conceptual Problems

1. Unlike the other chalcogens, oxygen does not form compounds in the +4 or +6 oxidation state. Why?
2. Classify each oxide as basic, acidic, amphoteric, or neutral.

1. CaO
2. SO₂
3. NO
4. Rb₂O
5. PbO₂

3. Classify each oxide as basic, acidic, amphoteric, or neutral.
   1. BaO
   2. Br₂O
   3. SnO
   4. B₂O₃
   5. Sb₂O₃

4. Polarization of an oxide affects its solubility in acids or bases. Based on this, do you expect RuO₂ to be an acidic, a basic, or a neutral oxide? Is the compound covalent? Justify your answers.
5. Arrange CrO₃, Al₂O₃, Sc₂O₃, and BaO in order of increasing basicity.
6. As the atomic number of the group 16 elements increases, the complexity of their allotropes decreases. What factors account for this trend? Which chalcogen do you expect to polymerize the most readily? Why?
7. Arrange H₃BO₃, HIO₄, and HNO₂ in order of increasing acid strength.
8. Of OF₂, SO₂, P₄O₆, SiO₂, and Al₂O₃, which is most ionic?
9. Of CO₂, NO₂, O₂, SO₂, Cl₂O, H₂O, NH₃, and CH₄, which do you expect to have the

1. most polar covalent bond(s)?
2. least polar covalent bond(s)?

10. Of Na₂O₂, MgO, Al₂O₃, and SiO₂, which is most acidic?

11. Give an example of
    1. a covalent hydride that engages in strong hydrogen bonding.
    2. an amphoteric oxide.

12. The Si–O bond is shorter and stronger than expected. What orbitals are used in this bond? Do you expect Si to interact with Br in the same way? Why or why not?

Answers

1. Oxygen has the second highest electronegativity of any element; consequently, it prefers to share or accept electrons from other elements. Only with fluorine does oxygen form compounds in positive oxidation states.

1. basic
2. acidic
3. amphoteric
4. acidic
5. amphoteric

5. CrO₃ < Al₂O₃ < Sc₂O₃ < BaO

7. H₃BO₃ < HNO₂ < HIO₄

9. Most polar: H₂O; least polar: O₂

1. H₂O, HF, or NH₃
2. SnO or Al₂O₃

Structure and Reactivity

1. Considering its position in the periodic table, predict the following properties of selenium:
   1. chemical formulas of its most common oxide, most common chloride, and most common hydride
   2. solubility of its hydride in water, and the acidity or basicity of the resulting solution
   3. the principal ion formed in aqueous solution

2. Using arguments based on electronegativity, explain why ZnO is amphoteric. What product would you expect when ZnO reacts with an aqueous
   1. acid?
   2. base?

3. Write a balanced chemical equation for the reaction of sulfur with
   1. O₂(g).
   2. S²⁻(aq).
   3. F₂(g).
   4. HNO₃(aq).

Answer

3. 1. S₈ + 8O₂ \(\xrightarrow{\Delta}\) 8SO₂(g)
   2. S₈(s) + 8S²⁻(aq) → 8S₂²⁻(aq)
   3. S₈(s) + 24F₂(g) → 8SF₆(g)
   4. S₈(s) + 48HNO₃(aq) → 8H₂SO₄(aq) + 48NO₂(g) + 16H₂O(l)

22.5: Group 15: The Nitrogen Family

Conceptual Problems

1. Nitrogen is the first diatomic molecule in the second period of elements. Why is N₂ the most stable form of nitrogen? Draw its Lewis electron structure. What hybrid orbitals are used to describe the bonding in this molecule? Is the molecule polar?
2. The polymer (SN)_n has metallic luster and conductivity. Are the constituent elements in this polymer metals or nonmetals? Why does the polymer have metallic properties?
3. Except for NF₃, all the halides of nitrogen are unstable. Explain why NF₃ is stable.
4. Which of the group 15 elements forms the most stable compounds in the +3 oxidation state? Explain why.
5. Phosphorus and arsenic react with the alkali metals to produce salts with the composition M₃Z₁₁. Compare these products with those produced by reaction of P and As with the alkaline earth metals. What conclusions can you draw about the types of structures favored by the heavier elements in this part of the periodic table?

Structure and Reactivity

1. PF₃ reacts with F₂ to produce PF₅, which in turn reacts with F⁻ to give salts that contain the PF₆⁻ ion. In contrast, NF₃ does not react with F₂, even under extreme conditions; NF₅ and the NF₆⁻ ion do not exist. Why?
2. Red phosphorus is safer to handle than white phosphorus, reflecting their dissimilar properties. Given their structures, how do you expect them to compare with regard to reactivity, solubility, density, and melting point?
3. Bismuth oxalate [Bi₂(C₂O₄)₃] is a poison. Draw its structure and then predict its solubility in H₂O, dilute HCl, and dilute HNO₃. Predict its combustion products. Suggest a method to prepare bismuth oxalate from bismuth.
4. Small quantities of NO can be obtained in the laboratory by reaction of the iodide ion with acidic solutions of nitrite. Write a balanced chemical equation that represents this reaction.
5. Although pure nitrous acid is unstable, dilute solutions in water are prepared by adding nitrite salts to aqueous acid. Write a balanced chemical equation that represents this type of reaction.
6. Metallic versus nonmetallic behavior becomes apparent in reactions of the elements with an oxidizing acid, such as HNO₃. Write balanced chemical equations for the reaction of each element of group 15 with nitric acid. Based on the products, predict which of these elements, if any, are metals and which, if any, are nonmetals.
7. Predict the product(s) of each reaction and then balance each chemical equation.

1. P₄O₁₀(s) + H₂O(l) →
2. AsCl₃(l) + H₂O(l) →
3. Bi₂O₃(s) + H₂O(l) →
4. Sb₄O₆(s) + OH⁻(aq) →
5. (C₂H₅)₃Sb(l) + O₂(g) \(\xrightarrow{\Delta}\)
6. SbCl₃(s) + LiAlH₄(soln) →
7. Ca(s) + N₂O(g) \(\xrightarrow{\Delta}\)

8. Write a balanced chemical equation to show how you would prepare each compound.
   1. H₃PO₄ from P
   2. Sb₂O₅ from Sb
   3. SbH₃ from Sb

Answers

5. NaNO₂(s) + HCl(aq) → HNO₂(aq) + NaCl(aq)

1. P₄O₁₀(s) + 6H₂O(l) → 4H₃PO₄(aq)
2. AsCl₃(l) + 3H₂O(l) → H₃AsO₃(aq) + 3HCl(aq)
3. Bi₂O₃(s) + 3H₂O(l) → 2Bi(OH)₃(s)
4. Sb₄O₆(s) + 4OH⁻(aq) + 2H₂O(l) → 4H₂SbO₃⁻(aq)
5. 2(C₂H₅)3Sb(l) + 21O₂(g) \(\xrightarrow{\Delta}\) 12CO₂(g) + 15H₂O(g) + Sb₂O₃(s)
6. 4SbCl₃(s) + 3LiAlH₄(soln) → 4SbH₃(g) + 3LiCl(soln) + 3AlCl₃(soln)
7. Ca(s) + N₂O(g) \(\xrightarrow{\Delta}\) CaO(s) + N₂(g)

22.6: Hydrogen: A Unique Element

Problems

1. Some periodic tables include hydrogen as a group 1 element, whereas other periodic tables include it as a group 17 element. Refer to the properties of hydrogen to propose an explanation for its placement in each group. In each case, give one property of hydrogen that would exclude it from groups 1 and 17.
2. If there were a planet where the abundances of D₂O and H₂O were reversed and life had evolved to adjust to this difference, what would be the effects of consuming large amounts of H₂O?
3. Describe the bonding in a hydrogen bond and the central B–H bond in B₂H₇⁻. Why are compounds containing isolated protons unknown?
4. With which elements does hydrogen form ionic hydrides? covalent hydrides? metallic hydrides? Which of these types of hydrides can behave like acids?
5. Indicate which elements are likely to form ionic, covalent, or metallic hydrides and explain your reasoning:

1. Sr
2. Si
3. O
4. Li
5. B
6. Be
7. Pd
8. Al

6. Which has the higher ionization energy—H or H⁻? Why?

7. The electronegativities of hydrogen, fluorine, and iodine are 2.20, 3.98, and 2.66, respectively. Why, then, is HI a stronger acid than HF?

8. If H₂O were a linear molecule, would the density of ice be less than or greater than that of liquid water? Explain your answer.

9. In addition to ion–dipole attractions, hydrogen bonding is important in solid crystalline hydrates, such as Na₄XeO₆·8H₂O. Based on this statement, explain why anhydrous Na₄XeO₆ does not exist.

Answers

1. H has one electron in an s orbital, like the group 1 metals, but it is also one electron short of a filled principal shell, like the group 17 elements. Unlike the alkali metals, hydrogen is not a metal. Unlike the halogens, elemental hydrogen is not a potent oxidant.

5. 1. ionic; it is an alkaline earth metal.
   2. covalent; it is a semimetal.
   3. covalent; it is a nonmetal.
   4. ionic; it is an alkali metal.
   5. covalent; it is a semimetal.
   6. covalent; it is a period 2 alkaline earth metal.
   7. metallic; it is a transition metal.
   8. covalent; it is a group 13 metal.

9. Hydrogen bonding with waters of hydration will partially neutralize the negative charge on the terminal oxygen atoms on the XeO₆⁴⁻ ion, which stabilizes the solid.

Structure and Reactivity

1. One of the largest uses of methane is to produce syngas, which is a source of hydrogen for converting nitrogen to ammonia. Write a complete equation for formation of syngas from methane and carbon dioxide. Calculate ΔG° for this reaction at 298 K and determine the temperature at which the reaction becomes spontaneous.

2. An alternative method of producing hydrogen is the water–gas shift reaction:

CO(g) + H₂O(g) → CO₂(g) + H₂(g)

Calculate ΔG° for this reaction at 298 K and determine the temperature at which the reaction changes from spontaneous to nonspontaneous (or vice versa).

3. Predict the products of each reaction at 25°C and then balance each chemical equation.
   1. CsH(s) + D₂O(l) →
   2. CH₃CO₂H(l) + D₂O(l) →
   3. H₃PO₄(aq) + D₂O(l) →
   4. NH₂CH₂CO₂H(s) + D₂O(l) →
   5. NH₄Cl(s) + D₂O(l) →

4. Using heavy water (D₂O) as the source of deuterium, how could you conveniently prepare
   1. D₂SO₄?
   2. LiD?

5. What are the products of reacting NaH with D₂O? Do you expect the same products from reacting NaD and H₂O? Explain your answer.

6. A 2.50 g sample of zinc metal reacts with 100.0 mL of 0.150 M HCl. What volume of H₂ (in liters) is produced at 23°C and 729 mmHg?

7. A chemical reaction requires 16.8 L of H₂gas at standard temperature and pressure. How many grams of magnesium metal are needed to produce this amount of hydrogen gas?

8. Seawater contains 3.5% dissolved salts by mass and has an average density of 1.026 g/mL. The volume of the ocean is estimated to be 1.35 × 10²¹ L. Using the data in Table \(\PageIndex{1}\), calculate the total mass of deuterium in the ocean.

9. From the data in Table \(\PageIndex{1}\), determine the molarity of DOH in water. Do you expect the molarity of D₂O in water to be similar? Why or why not?

10. From the data in Table \(\PageIndex{1}\), calculate how many liters of water you would have to evaporate to obtain 1.0 mL of TOD (tritium-oxygen-deuterium). The density of TOD is 1.159 g/mL.