22.E: Chemistry of the Nonmetals (Exercises)

22.1: General Concepts: Periodic Trends and Reactions

22.2: Hydrogen

22.3: Group 18: Nobel Gases

Questions

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)

22.4: Group 7A: The Halogens

Questions

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)

22.5: Oxygen

22.6: The Other Group 6A Elements: S, Se, Te, and Po

22.7: Nitrogen

Problems

• Complete and balance the following equations

N₂ + ___H₂→ ___NH_{_}

H₂N₂O₂ → ?

2NH₃ + CO₂ → ?

__Mg + N₂ → Mg_{_}N_{_}

N₂H₅ + H₂O → ?

• What are the different isotopes of Nitrogen?

• List the oxiadation states of various nitrogen oxides: N₂O, NO, N₂O_3, N₂O_4, N₂O₅

• List the different elements that Nitrogen will react with to make it basic or acidic....

• Uses of nitrogen

Answers

• Complete and balance the following equations

N₂ + 3H₂→ 2NH₃(Haber process)

H₂N₂O₂ → HNO

2NH₃ + CO₂ → (NH₂)₂CO + H₂O

2Mg + 3N₂ → Mg₃N₂

N₂H₅ + H₂O → N₂+ H⁺ + H₂O

• What are the different isotopes of Nitrogen?

Stable forms include nitrogen-14 and nitrogen-15

• List the oxidation states of various nitrogen oxides: +1, +2, +3, +4, +5 respectively

• List the different elements that Nitrogen will react with to make it basic or acidic :Nitride ion is a strong base when reacted with water, Ammonia is generally a weak acid
• Uses of nitrogen include anesthetic, Refrigerant, metal protector

22.8: The Other Group 5A Elements: P, AS, Sb, and Bi

22.9: Carbon

22.10: The Other Group 4A Elements: Si, Ge, Sn, and Pb

22.11: Boron