Chapter 22

SOLUTION Analyze We are asked to determine the nucleus that results when radium-226 loses an alpha particle. Plan We can best do this by writing a balanced nuclear reaction for the process. Solve The periodic table shows that radium has an atomic number of 88 . The complete chemical symbol for radium- 226 is therefore \({ }_{85}^{226} \mathrm{Ra}\). An alpha particle is a helium-4 nucleus, and so its symbol is \({ }_{2}^{4} \mathrm{He}\). The alpha particle is a product of the nuclear reaction, and so the equation is of the form $$ { }_{8}^{226} \mathrm{Ra} \longrightarrow{ }_{2}^{A} \mathrm{X}+{ }_{2}^{4} \mathrm{He} $$ where \(A\) is the mass number of the product nucleus and \(Z\) is its atomic number. Mass numbers and atomic numbers must balance, so $$ 226=A+4 $$ and $$ 88=Z+2 $$ Hence, $$ A=222 \text { and } Z=86 $$ Again, from the periodic table, the element with \(Z=86\) is radon (Rn). The product, therefore, is \({ }_{86}^{222} \mathrm{Rn}\), and the nuclear equation is $$ { }_{88}^{226} \mathrm{Ra} \longrightarrow{ }_{86}^{222} \mathrm{Rn}+{ }_{2}^{4} \mathrm{He} $$

SOLUTION Analyze We must write balanced nuclear equations in which the masses and charges of reactants and products are equal. Plan We can begin by writing the complete chemical symbols for the nuclei and decay particles that are given in the problem. Solve (a) The information given in the question can be summarized as $$ { }_{80}^{201} \mathrm{Hg}+{ }_{-1}^{0} \mathrm{e} \longrightarrow{ }_{Z}^{A} \mathrm{X} $$ The mass numbers must have the same sum on both sides of the equation: $$ 201+0=A $$ Thus, the product nucleus must have a mass number of 201. Similarly, balancing the atomic numbers gives $$ 80-1=Z $$ Thus, the atomic number of the product nucleus must be 79 , which identifies it as gold (Au): $$ { }_{80}^{201} \mathrm{Hg}+{ }_{-1}^{0} \mathrm{e} \longrightarrow{ }_{79}^{201} \mathrm{Au} $$ (b) In this case we must determine what type of particle is emitted in the course of the radioactive decay: $$ { }_{90}^{231} \mathrm{Th} \longrightarrow{ }_{91}^{231} \mathrm{~Pa}+{ }_{\mathrm{Z}}^{\mathrm{X}} \mathrm{X} $$ From \(231=231+A\) and \(90=91+Z\), we deduce \(A=0\) and \(Z=-1\). According to Table \(21.2\), the particle with these characteristics is the beta particle (electron). We therefore write $$ { }_{90}^{231} \mathrm{Th} \longrightarrow{ }_{91}^{231} \mathrm{~Pa}+{ }_{-1}^{0} \mathrm{e} \text { or } \quad{ }_{90}^{231} \mathrm{Th} \longrightarrow{ }_{91}^{231} \mathrm{~Pa}+\beta^{-} $$

You have two glass bottles, one containing oxygen and one filled with nitrogen. How could you determine which one is which? [Sections \(22.5\) and \(22.7\) ]

\text { Write the balanced nuclear equation for the process summarized as }{\underline{\phantom{xx}}}_{13}^{27} \mathrm{Al}(\mathrm{n}, \alpha)_{11}^{24} \mathrm{Na} \text {. }SOLUTION Analyze We must go from the condensed descriptive form of the reaction to the balanced nuclear equation. Plan We arrive at the balanced equation by writing \(\mathrm{n}\) and \(\alpha\), each with its associated subscripts and superscripts. Solve The \(\mathrm{n}\) is the abbreviation for a neutron \(\left({ }_{0}^{1} \mathrm{n}\right)\) and \(\alpha\) represents an alpha particle ( \(\left.{ }_{2}^{4} \mathrm{He}\right)\). The neutron is the bombarding particle, and the alpha particle is a product. Therefore, the nuclear equation is $$ { }_{13}^{27} \mathrm{Al}+{ }_{0}^{1} \mathrm{n} \longrightarrow{ }_{11}^{24} \mathrm{Na}+{ }_{2}^{4} \mathrm{He} \text { or } \quad{ }_{13}^{27} \mathrm{Al}+\mathrm{n} \longrightarrow{ }_{11}^{24} \mathrm{Na}+\alpha $$

The half-life of cobalt-60 is \(5.27\) yr. How much of a \(1.000-\mathrm{mg}\) sample of cobalt-60 is left after \(15.81 \mathrm{yr}\) ? SOLUTION Analyze We are given the half-life for cobalt-60 and asked to calculate the amount of cobalt-60 remaining from an initial \(1.000\)-mg sample after \(15.81\) yr. Plan We will use the fact that the amount of a radioactive substance decreases by \(50 \%\) for every half-life that passes. Solve Because \(5.27 \times 3=15.81,15.81\) yr is three half-lives for cobalt-60. At the end of one half-life, \(0.500 \mathrm{mg}\) of cobalt-60 remains, \(0.250 \mathrm{mg}\) at the end of two half-lives, and \(0.125 \mathrm{mg}\) at the end of three half-lives.

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 $$ [: X \equiv Y:]^{\pi} $$ 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 Bronsted base? Explain. [Sections 22.1, \(22.7\), and \(22.91\) eriodic Trends and Chemical Reactions Section 22.1)

Identify cach of the following elements as a metal, nonmetal, or metalloid: (a) phosphorus, (b) strontium, (c) manganese, (d) sclenium, (c) sodium, (f) krypton.

Identify cach of the following clements as a metal, nonmetal, or metalloid: (a) gallium, (b) molybdenum, (c) tellurium, (d) arsenic, (c) xenon, (f) ruthenium.

Consider the clements \(\mathrm{O}, \mathrm{Ba}, \mathrm{Co}, \mathrm{Bc}, \mathrm{Br}\), and Se. From this list, select the element that (a) is most electronegative, (b) exhibits a maximum oxidation state of \(+7\), (c) loses an clcctron most rcadily, (d) forms \(\pi\) bonds most readily, (c) is a transition metal, ( \(f\) ) is a liquid at room temperature and pressure.

Consider the elements L.i, K, Cl, C, Ne, and Ar. From this list, select the element that (a) is most clectronegative, (b) has the greatest metallic character, (c) most readily forms a positive ion, (d) has the smallest atomic radius, (c) forms \(\pi\) bonds most readily, (f) has multiple allotropes.

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) \(\mathrm{NaOCH}_{3}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (b) \(\mathrm{CuO}(s)+\mathrm{HNO}_{3}(a q) \longrightarrow\) (c) \(\mathrm{WO}_{3}(s)+\mathrm{H}_{2}(g) \stackrel{\Delta}{\longrightarrow}\) (d) \(\mathrm{NH}_{2} \mathrm{OH}(l)+\mathrm{O}_{2}(g) \longrightarrow\) (c) \(\mathrm{Al}_{1} \mathrm{C}_{1}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\)

Complete and balance the following equations: (a) \(\mathrm{Mg}_{3} \mathrm{~N}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l)\) (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) \(\mathrm{AlP}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (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.

Are the physical properties of \(\mathrm{H}_{2} \mathrm{O}\) different from \(\mathrm{D}_{2} \mathrm{O}\) ? Explain.

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

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

Write a balanced equation 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 \(\mathrm{H}_{2}\).

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) \(\mathrm{Na}(l)+\mathrm{H}_{2}(g)\) (c) \(\mathrm{PbO}(s)+\mathrm{H}_{2}(g)\)

Write balanced equations for each of the following reactions (some of these are analogous to reactions shown in the chapter). (a) Aluminum metal rcacts 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 manganesc(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.

Why does xenon form stable compounds with fluorine, whereas argon does 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) calcium hypobromite, (b) bromic acid, (c) xenon trioxide, (d) perchlorate ion, (e) iodous acid, (f) iodine pentafluoride.

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.

Name the following compounds and assign oxidation states to the halogens in them: (a) \(\mathrm{Fe}\left(\mathrm{ClO}_{3}\right)_{3}\), (b) \(\mathrm{HClO}_{2}\), (c) \(\mathrm{XeF}_{60}\) (d) \(\mathrm{BrF}_{5}\), (e) \(\mathrm{XeOF}_{4}\), (f) \(\mathrm{HIO}_{3}\),

Name the following compounds and assign oxidation states to the halogens in them: (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 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) \(\mathrm{Hy}\) drofluoric 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. Oxygen and the Other Group \(6 \mathrm{~A}\) Elements (Sections \(22.5\) and 22.6)

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, \(\mathrm{PbS}(s)\), 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) \(\mathrm{CaO}(s)+\mathrm{H}_{2} \mathrm{O}(l)\) (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) \(\mathrm{NO}_{2}\), (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 \(\mathrm{SO}_{3}\) v (d) \(\mathrm{SiO}_{2}\) and \(\mathrm{SO}_{2}\), (e) \(\mathrm{Ga}_{2} \mathrm{O}_{3}\) and \(\mathrm{In}_{2} \mathrm{O}_{3}\), (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 \(6 \mathrm{~A}\) element in each: (a) selenous acid, (b) potassium hydrogen sulfite, (c) hydrogen telluride, (d) carbon disulfide, (e) calcium sulfate, (f) cadmium sulfide, (g) zinc telluride.

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, (f) sulfur dioxide, (g) mercury telluride.

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) \(\mathrm{SeO}_{3}{\underline{\phantom{xx}}}^{2-} ;\) (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 \(\mathrm{SO}_{4}{\underline{\phantom{xx}}}^{2-}\) and elemental sulfur, while \(\mathrm{Cl}_{2}\) is reduced to \(\mathrm{Cl}^{-}\). Nitrogen and the Other Group 5A Elements (Sections 22.7 and 22.8)

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, (g) nitrogen, (h) boron nitride.

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, describe its geometry, and indicate the oxidation state of the nitrogen: (a) \(\mathrm{HNO}_{2}\), (b) \(\mathrm{N}_{3}^{-}\), (c) \(\mathrm{N}_{2} \mathrm{H}_{5}^{+}\), (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: (a) \(\mathrm{NH}_{4}^{+}\), (b) \(\mathrm{NO}_{2}^{-}\), (c) \(\mathrm{N}_{2} \mathrm{O}\), (d) \(\mathrm{NO}_{2}\).

Complete and balance the following equations: (a) \(\mathrm{Mg}_{3} \mathrm{~N}_{2}(s)+\mathrm{H}_{2} \mathrm{O}(l) \longrightarrow\) (b) \(\mathrm{NO}(g)+\mathrm{O}_{2}(g) \longrightarrow\) (c) \(\mathrm{N}_{2} \mathrm{O}_{5}(g)+\mathrm{H}_{2} \mathrm{O}(l)\) (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?