Chapter 22

Which of the following species (there may be more than one) is/ are likely to have the structure shown below: (a) \(\mathrm{XeF}_{4}\), (b) \(\mathrm{BrF}_{4}{\underline{\phantom{xx}}}^{+},(\mathrm{c}) \mathrm{SiF}_{4}\) (d) \(\mathrm{TeCl}_{4}\), \((\mathrm{e}) \mathrm{HClO}_{4} ?\) (The colors shown do not reflect the identity of any element.) [Sections 22.3, 22.4, 22.6, and 22.10]

The atomic and ionic radii of the first three group \(6 \mathrm{~A}\) elements are (a) Explain why the atomic radius increases in moving downward in the group. (b) Explain why the ionic radii are larger than the atomic radii. (c) Which of the three anions would you expect to be the strongest base in water? Explain. [Sections \(22.5\) and 22.6]

(a) Draw the Lewis structures for at least four species that have the general formula $$ [: \mathrm{X} \equiv \mathrm{Y}:]^{n} $$ where \(X\) and Y may be the same or different, and \(n\) may have a value from \(+1\) to \(-2\). (b) Which of the compounds is likely to be the strongest Brønsted base? Explain. [Sections 22.1, 22.7, and 22.9]

Identify each of the following elements as a metal, nonmetal, or metalloid: (a) phosphorus, (b) strontium, (c) manganese, (d) selenium, (e) rhodium, (f) krypton.

Identify each of the following elements as a metal, nonmetal, or metalloid: (a) gallium, (b) molybdenum, (c) tellurium, (d) arsenic, (e) xenon, (f) cadmium.

Consider the elements \(\mathrm{O}, \mathrm{Ba}, \mathrm{Co}, \mathrm{Be}, \mathrm{Br}\), and Se. From this list select the element that (a) is most electronegative, (b) exhibits a maximum oxidation state of \(+7\), (c) loses an electron most readily, (d) forms \(\pi\) bonds most readily, (e) is a transition metal.

Consider the elements \(\mathrm{Li}, \mathrm{K}, \mathrm{Cl}, \mathrm{C}, \mathrm{Ne}\), and Ar. From this list select the element that (a) is most electronegative, (b) has the greatest metallic character, (c) most readily forms a positive ion, (d) has the smallest atomic radius, (e) forms \(\pi\) bonds most readily.

Explain the following observations: (a) The highest fluoride compound formed by nitrogen is \(\mathrm{NF}_{3}\), whereas phosphorus readily forms \(\mathrm{PF}_{5} .\) (b) Although \(\mathrm{CO}\) is a well-known compound, SiO does not exist under ordinary conditions. (c) \(\mathrm{AsH}_{3}\) is a stronger reducing agent than \(\mathrm{NH}_{3}\)

Explain the following observations: (a) \(\mathrm{HNO}_{3}\) is a stronger oxidizing agent than \(\mathrm{H}_{3} \mathrm{PO}_{4}\). (b) Silicon can form an ion with six fluorine atoms, \(\mathrm{SiF}_{6}{\underline{\phantom{xx}}}^{2-}\), whereas carbon is able to bond to a maximum of four, \(\mathrm{CF}_{4}\). (c) There are three compounds formed by carbon and hydrogen that contain two carbon atoms each \(\left(\mathrm{C}_{2} \mathrm{H}_{2}, \mathrm{C}_{2} \mathrm{H}_{4}\right.\), and \(\left.\mathrm{C}_{2} \mathrm{H}_{6}\right)\), whereas silicon forms only one analogous compound \(\left(\mathrm{Si}_{2} \mathrm{H}_{6}\right)\).

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

Complete and balance the following equations: (a) \(\mathrm{NaOCH}_{3}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (b) \(\mathrm{CuO}(s)+\mathrm{HNO}_{3}(a q) \longrightarrow\) (c) \(\mathrm{WO}_{3}(s)+\mathrm{H}_{2}(g) \stackrel{\Delta}{\longrightarrow}\) (d) \(\mathrm{NH}_{2} \mathrm{OH}(l)+\mathrm{O}_{2}(g) \longrightarrow\) (e) \(\mathrm{Al}_{4} \mathrm{C}_{3}(\mathrm{~s})+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\)

(a) Give the names and chemical symbols for the three isotopes of hydrogen. (b) List the isotopes in order of decreasing natural abundance.

Which isotope of hydrogen is radioactive? Write the nuclear equation for the radioactive decay of this isotope.

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

Why are the properties of hydrogen different from those of either the group \(1 \mathrm{~A}\) or \(7 \mathrm{~A}\) elements?

Write a balancedequation for the preparation of \(\mathrm{H}_{2}\) using (a) \(\mathrm{Mg}\) and an acid, (b) carbon and steam, (c) methane and steam.

List (a) three commercial means of producing \(\mathrm{H}_{2}\), (b) three industrial uses of \(\mathrm{H}_{2}\).

Completeand 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) \(\mathrm{PbO}(\mathrm{s})+\mathrm{H}_{2}(g) \longrightarrow\)

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) \(\mathrm{BaH}_{2}\), (b) \(\mathrm{H}_{2} \mathrm{Te}\), (c) \(\mathrm{TiH}_{1.7}\).

Identify the following hydrides as ionic, metallic, or molecular: (a) \(\mathrm{B}_{2} \mathrm{H}_{6}\), (b) \(\mathrm{RbH}\), (c) \(\mathrm{Th}_{4} \mathrm{H}_{1.5}\)

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

Would hydrogen be a satisfactory basis for a fuel econo\(\mathrm{my}\) if the only available sources were cracking of natural gas and conversion of grain- based ethanol? Explain.

Why does xenon form stable compounds with fluorine, whereas argon does not?

Why are there so few compounds of the noble gases?

Write the chemical formula for each of the following compounds, and indicate the oxidation state of the halogen or noble-gas atom in each: (a) chlorate ion, (b) hydroiodic acid, (c) iodine trichloride, (d) sodium hypochlorite, (e) perchloric acid, (f) xenon tetrafluoride.

Write the chemical formula for each of the following, and indicate the oxidation state of the halogen or noble-gas atom in each, (a) calcium hypobromite, (b) bromic acid, (c) xenon trioxide, (d) perchlorate ion, (e) iodous acid, (f) iodine pentafluoride.

Name the following compounds: (a) \(\mathrm{Fe}\left(\mathrm{ClO}_{3}\right)_{3}\), (b) \(\mathrm{HClO}_{2},(\mathrm{c}) \mathrm{XeF}_{6},(\mathrm{~d}) \mathrm{BrF}_{5},(\mathrm{e}) \mathrm{XeOF}_{4}\), (f) \(\mathrm{HIO}_{3}\) (named as an acid).

Name the following compounds: (a) \(\mathrm{KClO}_{3}\), (b) \(\mathrm{Ca}\left(\mathrm{IO}_{3}\right)_{2}\), (c) \(\mathrm{AlCl}_{3}\), (d) \(\mathrm{HBrO}_{3}\), (e) \(\mathrm{H}_{5} \mathrm{IO}_{6}\), (f) \(\mathrm{XeF}_{4}\).

Explain each of the following observations: (a) At room temperature \(\mathrm{I}_{2}\) is a solid, \(\mathrm{Br}_{2}\) is a liquid, and \(\mathrm{Cl}_{2}\) and \(\mathrm{F}_{2}\) are both gases. (b) \(\mathrm{F}_{2}\) cannot be prepared by electrolytic oxidation of aqueous \(\mathrm{F}^{-}\) solutions. (c) The boiling point of \(\mathrm{HF}\) is much higher than those of the other hydrogen halides. (d) The halogens decrease in oxidizing power in the order \(\mathrm{F}_{2}>\mathrm{Cl}_{2}>\mathrm{Br}_{2}>\mathrm{I}_{2}\)

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 NaI with sulfuric acid. (d) The interhalogen \(\mathrm{ICl}_{3}\) is known, but \(\mathrm{BrCl}_{3}\) is not.

Ammonium perchlorate remains a significant pollutant in soils and water supplies even decades after it is placed into the environment. What can you conclude from this about the stability of the perchlorate anion? What feature of the anion's structure might promote this high stability, even though perchlorate is a strong oxidizing agent?

Removal of perchlorate from water supplies is difficult. Naturally occurring microorganisms are, however, capable of destroying perchlorate in solution in minutes. What do you think might be the type of reaction occurring in the microorganisms, and what do you predict might be the fate of the perchlorate ion in the reaction?

(a) List three industrial uses of \(\mathrm{O}_{2}\). (b) List two industrial uses of \(\mathrm{O}_{3}\)

Draw the Lewis structure of ozone. Explain why the \(\mathrm{O}-\mathrm{O}\) bond \((1.28 \AA)\) is longer in ozone than in \(\mathrm{O}_{2}(1.21 \AA)\).

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, \(\mathrm{PbS}(\mathrm{s})\), reacts with ozone to form \(\mathrm{PbSO}_{4}(\mathrm{~s})\) and \(\mathrm{O}_{2}(\mathrm{~g}) .\) (d) When heated in air, \(\mathrm{ZnS}(s)\) is converted to ZnO. (e) Potassium peroxide reacts with \(\mathrm{CO}_{2}(g)\) to give potassium carbonate and \(\mathrm{O}_{2}\).

Complete and balance the following equations: (a) \(\mathrm{CaO}(\mathrm{s})+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (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) \(\mathrm{NO}(\mathrm{g})+\mathrm{O}_{3}(\mathrm{~g}) \longrightarrow\)

Predict whether each of the following oxides is acidic, basic, amphoteric, or neutral: (a) \(\mathrm{NO}_{2}\), (b) \(\mathrm{CO}_{2}\) (c) \(\mathrm{Al}_{2} \mathrm{O}_{3}\), (d) \(\mathrm{CaO}\).

Select the more acidic member of each of the following pairs: (a) \(\mathrm{Mn}_{2} \mathrm{O}_{7}\) and \(\mathrm{MnO}_{2}\), (b) \(\mathrm{SnO}\) and \(\mathrm{SnO}_{2}\), (c) \(\mathrm{SO}_{2}\) and \(\mathrm{SO}_{3},(\mathrm{~d}) \mathrm{SiO}_{2}\) and \(\mathrm{SO}_{2}\), (e) \(\mathrm{Ga}_{2} \mathrm{O}_{3}\) and \(\mathrm{In}_{2} \mathrm{O}_{3},(\mathbf{f}) \mathrm{SO}_{2}\) and \(\mathrm{SeO}_{2}\)

Write the chemical formula for each of the following compounds, and indicate the oxidation state of the group 6 A element in each: (a) selenous acid, (b) potassium hydrogen sulfite, (c) hydrogen telluride, (d) carbon disulfide, (e) calcium sulfate.

Write the chemical formula for each of the following compounds, and indicate the oxidation state of the group \(6 \mathrm{~A}\) element in each: (a) sulfur tetrachloride, (b) selenium trioxide, (c) sodium thiosulfate, (d) hydrogen sulfide, (e) sulfuric acid.

In aqueous solution, hydrogen sulfide reduces (a) \(\mathrm{Fe}^{3+}\) to \(\mathrm{Fe}^{2+}\), (b) \(\mathrm{Br}_{2}\) to \(\mathrm{Br}^{-}\), (c) \(\mathrm{MnO}_{4}^{-}\) to \(\mathrm{Mn}^{2+}\), (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 \(\mathrm{MnSO}_{4}(a q)\), (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.

Write the Lewis structure for each of the following species, and indicate the structure of each: (a) \(\mathrm{SeO}_{3}^{2-}\); (b) \(\mathrm{S}_{2} \mathrm{Cl}_{2}\); (c) chlorosulfonic acid, \(\mathrm{HSO}_{3} \mathrm{Cl}\) (chlorine is bonded to sulfur).

The \(\mathrm{SF}_{5}^{-}\) ion is formed when \(\mathrm{SF}_{4}(g)\) reacts with fluoride salts containing large cations, such as \(\mathrm{CsF}(s) .\) 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. (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 nitrite, (b) ammonia, (c) nitrous oxide, (d) sodium cyanide, (e) nitric acid, (f) nitrogen dioxide.

Write the chemical formula for each of the following compounds, and indicate the oxidation state of nitrogen in each: (a) nitric oxide, (b) hydrazine, (c) potassium cyanide, (d) sodium nitrite, (e) ammonium chloride, (f) lithium nitride.

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