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22.E: Exercises

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  • 22.1: Periodic Trends in Bonding


    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?


    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:


      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−
    1. Which element is the least metallic—B, Ga, Tl, Pb, Ne, or Ge?
    2. Of Br, N, Ar, Bi, Se, He, and S, which would you expect to form positive ions most easily? negative ions most easily?
    3. 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?
    1. Of HCl, HClO4, HBr, H2S, HF, KrF2, and PH3, which is the strongest acid?
    2. 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
    1. 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
    1. 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?
    2. Write a balanced chemical equation to describe the reaction of each species with water.
    1. B2H6
    2. F2
    3. C4+
    1. Using heavy water (D2O) as the source of deuterium, how could you prepare each compound?
      1. LiAlD4
      2. D2SO4
      3. SiD4
      4. DF
    2. 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)


    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)
    1. SiF4; SiO2(s) + 2F2(g) → SiF4(l) + 2O2(g)
      1. 2Na(s) + 2D2O(l) → D2(g) + 2NaOD(aq)

    2Li(s) + D2(g) → 2LiD(s)

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

    1. D2O(l) + SO3(g) → D2SO4(l)
    2. SiCl4(l) + LiAlD4(s) [from part (a)] → SiD4(g) + LiCl(s) + AlCl3(s)
    3. 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.


    1. Electrostatic repulsions between lone pairs on adjacent atoms decrease bond strength.
    1. Ionic character decreases as Δχ decreases from Al to P.
    1. 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) →
    1. Write a balanced chemical equation for the reaction of aqueous HF with
      1. SiO2.
      2. Na2CO3.
      3. CaO.
    1. 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.
    1. 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
    1. 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


    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
    1. Classify each oxide as basic, acidic, amphoteric, or neutral.
      1. BaO
      2. Br2O
      3. SnO
      4. B2O3
      5. Sb2O3
    1. 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.
    2. Arrange CrO3, Al2O3, Sc2O3, and BaO in order of increasing basicity.
    3. 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?
    4. Arrange H3BO3, HIO4, and HNO2 in order of increasing acid strength.
    5. Of OF2, SO2, P4O6, SiO2, and Al2O3, which is most ionic?
    6. 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)?
    1. Of Na2O2, MgO, Al2O3, and SiO2, which is most acidic?
    1. Give an example of
      1. a covalent hydride that engages in strong hydrogen bonding.
      2. an amphoteric oxide.
    1. 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?


    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
    1. CrO3 < Al2O3 < Sc2O3 < BaO
    1. H3BO3 < HNO2 < HIO4
    1. 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
    1. 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?
    1. Write a balanced chemical equation for the reaction of sulfur with
      1. O2(g).
      2. S2−(aq).
      3. F2(g).
      4. HNO3(aq).


      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) \(\xrightarrow{\Delta}\)
    1. 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


    1. 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


    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
    1. Which has the higher ionization energy—H or H? Why?
    1. 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?
    1. If H2O were a linear molecule, would the density of ice be less than or greater than that of liquid water? Explain your answer.
    1. 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.


    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.
      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.
    1. 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.
    1. 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).

    1. 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) →
    1. Using heavy water (D2O) as the source of deuterium, how could you conveniently prepare
      1. D2SO4?
      2. LiD?
    1. What are the products of reacting NaH with D2O? Do you expect the same products from reacting NaD and H2O? Explain your answer.
    1. 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?
    1. 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?
    1. 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.
    1. 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?
    1. 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.