Chapter 22

Identify each of the following elements as a metal, nonmetal, or metalloid: (a) boron, (b) barium, (c) argon, \((\mathbf{d})\) caesium, \((\mathbf{e})\) yttrium, \((\mathbf{f})\) astatine.

Identify each of the following elements as a metal, nonmetal, or metalloid: (a) germanium, \((\mathbf{b})\) bismuth, (c) sulphur, \((\mathbf{d})\) calcium, \((\mathbf{e})\) rhenium, \((\mathbf{f})\) tin.

Consider the elements \(\mathrm{N}, \mathrm{F}\), Si, Rb, Te, and Ir. From this list, select the element that (a) is most electronegative, \((\mathbf{b})\) exhibits a maximum oxidation state of \(+6,(\mathbf{c})\) loses an electron most readily, (d) forms \(\pi\) bonds most readily, \((\mathbf{e})\) is a transition metal, (f) forms four covalent bonds to achieve octet.

Consider the elements \(\mathrm{Ba}, \mathrm{Na}, \mathrm{O}, \mathrm{B}, \mathrm{P},\) and \(\mathrm{Kr}\). From this list, select the element that (a) is most electronegative, (b) has the greatest metalliccharacter, \((\mathbf{c})\) most readily forms a positive ion, \((\mathbf{d})\) exhibits a maximum oxidation sate of +5 , (e) exists as monoatomic gas at room temperature, (f) has multiple allotropes.

Which of the following statements are true? (a) Si can form an ion with six fluorine atoms, \(\mathrm{SiF}_{6}^{2-}\), whereas carbon cannot. (b) Si can form three stable compounds containing two Si atoms each, \(\mathrm{Si}_{2} \mathrm{H}_{2}, \mathrm{Si}_{2} \mathrm{H}_{4},\) and \(\mathrm{Si}_{2} \mathrm{H}_{6}\) (c) In \(\mathrm{HNO}_{3}\) and \(\mathrm{H}_{3} \mathrm{PO}_{4}\) the central atoms, \(\mathrm{N}\) and \(\mathrm{P}\), have different oxidation states. (d) \(\mathrm{S}\) is more electronegative than Se.

Complete and balance the following equations: (a) \(\mathrm{Mg}_{3} \mathrm{~N}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (b) \(\mathrm{C}_{3} \mathrm{H}_{7} \mathrm{OH}(I)+\mathrm{O}_{2}(g) \longrightarrow\) (c) \(\mathrm{MnO}_{2}(s)+\mathrm{C}(s) \stackrel{\Delta}{\longrightarrow}\) (d) \(\mathrm{AlP}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (e) \(\mathrm{Na}_{2} \mathrm{~S}(s)+\mathrm{HCl}(a q) \longrightarrow\)

(a) Give the names and chemical symbols for the three isotopes of hydrogen. (b) List the isotopes in order of decreasing natural abundance. (c) Which hydrogen isotope is radioactive? (d) Write the nuclear equation for the radioactive decay of this isotope.

The physical properties of \(\mathrm{D}_{2} \mathrm{O}\) differ from those of \(\mathrm{H}_{2} \mathrm{O}\) because (a) D has a different electron configuration than \(\mathrm{O}\). (b) D is radioactive. (c) D forms stronger bonds with \(\mathrm{O}\) than \(\mathrm{H}\) does. (d) \(\mathrm{D}\) is much more massive than \(\mathrm{H}\).

Give a reason why hydrogen might be placed along with the group 1 elements of the periodic table.

What does hydrogen have in common with the halogens? Explain.

Complete and balance the following equations: (a) \(\mathrm{NaH}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (b) \(\mathrm{Fe}(s)+\mathrm{H}_{2} \mathrm{SO}_{4}(a q) \longrightarrow\) (c) \(\mathrm{H}_{2}(g)+\mathrm{Br}_{2}(g) \longrightarrow\) (d) \(\mathrm{Na}(l)+\mathrm{H}_{2}(g) \longrightarrow\) (e) \(\mathrm{PbO}(s)+\mathrm{H}_{2}(g) \longrightarrow\)

Write balanced equations for each of the following reactions (some of these are analogous to reactions shown in the chapter). (a) Aluminum metal reacts with acids to form hydrogen gas. (b) Steam reacts with magnesium metal to give magnesium oxide and hydrogen. (c) Manganese(IV) oxide is reduced to manganese(II) oxide by hydrogen gas. (d) Calcium hydride reacts with water to generate hydrogen gas.

Identify the following hydrides as ionic, metallic, or molecular: (a) \(\mathrm{H}_{2} \mathrm{~S},(\mathbf{b}) \mathrm{LiH},(\mathbf{c}) \mathrm{VH}_{0.56}\)

Identify the following hydrides as ionic, metallic, or molecular: \((\mathbf{a}) \mathrm{MgH}_{2},(\mathbf{b}) \mathrm{HI},(\mathbf{c}) \mathrm{LaNi}_{4.7} \mathrm{Al}_{0.3} \mathrm{H}_{6}\)

Describe two characteristics of hydrogen that are favorable for its use as a general energy source in vehicles.

The \(\mathrm{H}_{2} / \mathrm{O}_{2}\) fuel cell converts elemental hydrogen and oxygen into water, producing, theoretically, \(1.23 \mathrm{~V}\). What is the most sustainable way to obtain hydrogen to run a large number of fuel cells? Explain.

Why does xenon form stable compounds with fluorine, whereas argon does \(\underline{\text { not } ?}\)

A friend tells you that the "neon" in neon signs is a compound of neon and aluminum. Can your friend be correct? Explain.

Write the chemical formula for each of the following, and indicate the oxidation state of the halogen or noble-gas atom in each: (a) xenon trioxide difluoride, (b) chlorine dioxide, (c) molybdenum hexafluoride, (d) iodic acid, (e) sodium hypobromite, (f) magnesium iodite.

Write the chemical formula for each of the following, and indicate the oxidation state of the halogen or noble-gas atom in each: (a) krypton tetrafluoride, (b) hexafluoroantimonate ion, (c) sodium hypoiodite, (d) perbromic acid, (e) aluminum perchlorate, (f) iron(II) iodite.

Name the following compounds and assign oxidation states to the halogens in them: (a) \(\mathrm{Co}\left(\mathrm{IO}_{3}\right)_{3},\) (b) \(\mathrm{Ca}\left(\mathrm{lO}_{4}\right)_{2}\), (c) \(\mathrm{PF}_{6}^{-},(\mathbf{d}) \mathrm{ICl}_{4}^{-}\) (e) HBrO, (f) \(\mathrm{AtH}\)

Name the following compounds and assign oxidation states to the halogens in them: (a) \(\mathrm{BCl}_{3},\) (b) \(\operatorname{Sr}\left(\mathrm{IO}_{4}\right)_{2}\), (c) LiOCl, (d) \(\mathrm{HClO}\) (e) \(\mathrm{CuClO}\) (f) \(\mathrm{Mg}\left(\mathrm{IO}_{2}\right)\)

Explain the following observations: (a) For a given oxidation state, the acid strength of the oxyacid in aqueous solution decreases in the order chlorine \(>\) bromine \(>\) iodine. (b) Hydrofluoric acid cannot be stored in glass bottles. (c) HI cannot be prepared by treating Nal with sulfuric acid. (d) The interhalogen \(\mathrm{ICl}_{3}\) is known, but \(\mathrm{BrCl}_{3}\) is not.

Write balanced equations for each of the following reactions. (a) When mercury(II) oxide is heated, it decomposes to form \(\mathrm{O}_{2}\) and mercury metal. (b) When copper(II) nitrate is heated strongly, it decomposes to form copper(II) oxide, nitrogen dioxide, and oxygen. (c) Lead(II) sulfide, PbS(s), reacts with ozone to form \(\mathrm{PbSO}_{4}(s)\) and \(\mathrm{O}_{2}(g) .\) (d) When heated in air, \(\mathrm{ZnS}(s)\) is converted to \(\mathrm{ZnO}\). (e) Potassium peroxide reacts with \(\mathrm{CO}_{2}(g)\) to give potassium carbonate and \(\mathrm{O}_{2}\). (f) Oxygen is converted to ozone in the upper atmosphere.

Complete and balance the following equations: (a) \(\mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (b) \(\mathrm{Al}_{2} \mathrm{O}_{3}(s)+\mathrm{H}^{+}(a q) \longrightarrow\) (c) \(\mathrm{Na}_{2} \mathrm{O}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (d) \(\mathrm{N}_{2} \mathrm{O}_{3}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (e) \(\mathrm{KO}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (f) \(\mathrm{NO}(g)+\mathrm{O}_{3}(g) \longrightarrow\)

Predict whether each of the following oxides is acidic, basic, amphoteric, or neutral: (a) \(\mathrm{SO}_{3},(\mathbf{b}) \mathrm{L}_{2} \mathrm{O},(\mathbf{c}) \mathrm{SnO},(\mathbf{d}) \mathrm{ZnO}\).

Select the more acidic member of each of the following pairs: (a) \(\mathrm{V}_{2} \mathrm{O}_{5}\) and VO, (b) \(\mathrm{Pb} \mathrm{O}\) and \(\mathrm{PbO}_{2}\), (c) \(\mathrm{N}_{2} \mathrm{O}_{3}\) and \(\mathrm{N}_{2} \mathrm{O}_{2},\) (d) \(\mathrm{SO}_{2}\) and \(\mathrm{SeO}_{2}\), (e) \(\mathrm{Al}_{2} \mathrm{O}_{3}\) and \(\mathrm{SeO}_{2},\) (f) \(\mathrm{CO}_{2}\) and \(\mathrm{B}_{2} \mathrm{O}_{3} .\)

Write the chemical formula for each of the following compounds, and indicate the oxidation state of the group 16 elements in each: (a) potassium peroxide, (b) potassium thiosulfate, \((\mathbf{c})\) selenium oxychloride, \((\mathbf{d})\) sodium telluride, (e) magnesium sulfite, \((\mathbf{f})\) selenium hexafluoride.

Write the chemical formula for each of the following compounds, and indicate the oxidation state of the group 16 elements in each: (a) selenious acid, (b) sulphur trioxide, \((\mathbf{c})\) selenium dichloride, \((\mathbf{d})\) aluminium selenide, (e) iron(II) sulfate, (f) tellurium trioxide.

In aqueous solution, hydrogen sulfide reduces (a) \(\mathrm{Fe}^{3+}\) to \(\mathrm{Fe}^{2+},(\mathbf{b}) \mathrm{Br}_{2}\) to \(\mathrm{Br}^{-},(\mathbf{c}) \mathrm{MnO}_{4}^{-}\) to \(\mathrm{Mn}^{2+},(\mathbf{d}) \mathrm{HNO}_{3}\) to \(\mathrm{NO}_{2}\) In all cases, under appropriate conditions, the product is elemental sulfur. Write a balanced net ionic equation for each reaction.

An aqueous solution of \(\mathrm{SO}_{2}\) reduces (a) aqueous \(\mathrm{KMnO}_{4}\), to \(\mathrm{MnSO}_{4}(a q),(\mathbf{b})\) acidic aqueous \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) to aqueous \(\mathrm{Cr}^{3+}\) (c) aqueous \(\mathrm{Hg}_{2}\left(\mathrm{NO}_{3}\right)_{2}\) to mercury metal. Write balanced equations for these reactions.

The \(\mathrm{SF}_{5}^{-}\) ion is formed when \(\mathrm{SF}_{4}(g)\) reacts with fluoride salts containing large cations, such as CsF(s). Draw the Lewis structures for \(\mathrm{SF}_{4}\) and \(\mathrm{SF}_{5}^{-}\), and predict the molecular structure of each.

Write a balanced equation for each of the following reactions: (a) Sulfur dioxide reacts with water. (b) Solid zinc sulfide reacts with hydrochloric acid. (c) Elemental sulfur reacts with sulfite ion to form thiosulfate. (d) Sulfur trioxide is dissolved in sulfuric acid.

Write a balanced equation for each of the following reactions. (You may have to guess at one or more of the reaction products, but you should be able to make a reasonable guess, based on your study of this chapter.) (a) Hydrogen selenide can be prepared by reaction of an aqueous acid solution on aluminum selenide. \((\mathbf{b})\) Sodium thiosulfate is used to remove excess \(\mathrm{Cl}_{2}\) from chlorine-bleached fabrics. The thiosulfate ion forms \(\mathrm{SO}_{4}^{2-}\) and elemental sulfur, while \(\mathrm{Cl}_{2}\) is reduced to \(\mathrm{Cl}^{-}\).

Write the chemical formula for each of the following compounds, and indicate the oxidation state of nitrogen in each: (a) sodium azide, (b) ammonium ion, (c) nitrous acid, (d) magnesium nitride, \((\mathbf{e})\) diazene, \((\mathbf{f})\) sodium nitrate, (g) nitrogen trifluoride, \((\mathbf{h})\) nitric acid.

Write the chemical formula for each of the following compounds, and indicate the oxidation state of nitrogen in each: (a) dinitrogen trioxide, (b) dinitrogen pentoxide, (c) hydrogen cyanide, \((\mathbf{d})\) ammonium nitrate, (e) iron(III) nitrite, \((\) f) nitrogen trichloride.

Write the Lewis structure for each of the following species, describe its geometry, and indicate the oxidation state of the nitrogen: \((\mathbf{a}) \mathrm{H} \mathrm{NO}_{2},(\mathbf{b}) \mathrm{N}_{3}^{-},(\mathbf{c}) \mathrm{N}_{2} \mathrm{H}_{5}^{+},(\mathbf{d}) \mathrm{NO}_{3}^{-}\).

Write the Lewis structure for each of the following species, describe its geometry, and indicate the oxidation state of the nitrogen: \((\mathbf{a}) \mathrm{N}_{2} \mathrm{O}_{3},(\mathbf{b}) \mathrm{NOCl},(\mathbf{c}) \mathrm{NO}_{2} \mathrm{Cl},(\mathbf{d}) \mathrm{N}_{2} \mathrm{O}_{4}\).

Complete and balance the following equations: (a) \(\mathrm{Mg}_{3} \mathrm{~N}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(I) \longrightarrow\) (b) \(\mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow\) (c) \(\mathrm{N}_{2} \mathrm{O}_{5}(g)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (d) \(\mathrm{NH}_{3}(a q)+\mathrm{H}^{+}(a q) \longrightarrow\) (e) \(\mathrm{N}_{2} \mathrm{H}_{4}(l)+\mathrm{O}_{2}(g) \longrightarrow\) Which ones of these are redox reactions?

Write a balanced net ionic equation for each of the following reactions: (a) Dilute nitric acid reacts with zinc metal with formation of nitrous oxide. (b) Concentrated nitric acid reacts with sulfur with formation of nitrogen dioxide. (c) Concentrated nitric acid oxidizes sulfur dioxide with formation of nitric oxide. (d) Hydrazine is burned in excess fluorine gas, forming \(\mathrm{NF}_{3}\). (e) Hydrazine reduces \(\mathrm{CrO}_{4}^{2-}\) to \(\mathrm{Cr}(\mathrm{OH})_{4}\) in base (hydrazine is oxidized to \(\mathrm{N}_{2}\) ).

Write complete balanced half-reactions for (a) oxidation of nitrous acid to nitrate ion in acidic solution, (b) oxidation of \(\mathrm{N}_{2}\) to \(\mathrm{N}_{2} \mathrm{O}\) in acidic solution.

Write complete balanced half-reactions for (a) reduction of nitrate ion to NO in acidic solution, \((\mathbf{b})\) oxidation of \(\mathrm{HNO}_{2}\) to \(\mathrm{NO}_{2}\) in acidic solution.

Write a molecular formula for each compound, and indicate the oxidation state of the group 15 element in each formula: (a) hexafluoroantimonate(V) ion, (b) calcium phosphate, (c) potassium dihydrogen phosphate, (d) arsenic trioxide, (e) tetraphosphorus hexaoxide, (f) arsenic trifluoride.

Account for the following observations: (a) Phosphorus forms a pentachloride, but nitrogen does not. (b) \(\mathrm{H}_{3} \mathrm{PO}_{2}\) is a monoprotic acid. (c) Phosphonium salts, such as \(\mathrm{PH}_{4} \mathrm{Cl}\), can be formed under anhydrous conditions, but they cannot be made in aqueous solution. (d) White phosphorus is more reactive than red phosphorus.

Account for the following observations: (a) \(\mathrm{H}_{3} \mathrm{PO}_{3}\) is a diprotic acid. (b) Nitric acid is a strong acid, whereas phosphoric acid is weak. (c) Phosphate rock is ineffective as a phosphate fertilizer. (d) Phosphorus does not exist at room temperature as diatomic molecules, but nitrogen does. (e) Solutions of \(\mathrm{Na}_{3} \mathrm{PO}_{4}\) are quite basic.

Write a balanced equation for each of the following reactions: (a) preparation of white phosphorus from calcium phosphate, \((\mathbf{b})\) hydrolysis of \(\mathrm{PBr}_{3},(\mathbf{c})\) reduction of \(\mathrm{PBr}_{3}\) to \(\mathrm{P}_{4}\) in the gas phase, using \(\mathrm{H}_{2}\).

Write a balanced equation for each of the following reactions: (a) hydrolysis of \(\mathrm{PCl}_{5},\) (b) dehydration of phosphoric acid (also called orthophosphoric acid) to form pyrophosphoric acid, \((\mathbf{c})\) reaction of \(\mathrm{P}_{4} \mathrm{O}_{10}\) with water.

Give the chemical formula for (a) copper(II) carbonate, (b) carbon monoxide, (c) magnesium hydrogen carbonate, (d) lithium acetylide, (e) carbon tetrafluoride.

Give the chemical formula for (a) fullerene, (b) potassium cyanide, \((\mathbf{c})\) zinc carbide, \((\mathbf{d})\) zinc acetylide, (e) carbon disulfide.

Complete and balance the following equations: (a) \(\mathrm{ZnCO}_{3}(s) \stackrel{\Delta}{\longrightarrow}\) (b) \(\mathrm{BaC}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (c) \(\mathrm{C}_{2} \mathrm{H}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow\) (d) \(\mathrm{CS}_{2}(g)+\mathrm{O}_{2}(g) \longrightarrow\) (e) \(\mathrm{Ca}(\mathrm{CN})_{2}(s)+\mathrm{HBr}(a q) \longrightarrow\)