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3 Bohr's Theory of the Hydrogen Atom

3 Bohr's Theory of the Hydrogen Atom

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Questions and Problems



16.68 Which of the following will be more soluble in acid

solution than in pure water? (a) CuI, (b) Ag2SO4,

(c) Zn(OH)2, (d) BaC2O4, (e) Ca3(PO4)2

16.69 Compare the molar solubility of Mg(OH)2 in water

and in a solution buffered at a pH of 9.0.

16.70 Calculate the molar solubility of Fe(OH)2 in a solution buffered at (a) pH 8.00, (b) pH 10.00.

16.71 The solubility product of Mg(OH)2 is 1.2 3 10211.

What minimum OH2 concentration must be attained

(for example, by adding NaOH) to decrease the

Mg21 concentration in a solution of Mg(NO3)2 to

less than 1.0 3 10210 M?

16.72 Calculate whether or not a precipitate will form if

2.00 mL of 0.60 M NH3 are added to 1.0 L of

1.0 3 1023 M FeSO4.



Complex Ion Equilibria and Solubility

Review Questions

16.73 Explain the formation of complexes in Table 16.4 in

terms of Lewis acid-base theory.

16.74 Give an example to illustrate the general effect of

complex ion formation on solubility.



Problems

16.75 If 2.50 g of CuSO4 are dissolved in 9.0 3 102 mL of

0.30 M NH3, what are the concentrations of Cu21,

Cu(NH3)21

4 , and NH3 at equilibrium?

16.76 Calculate the concentrations of Cd21, Cd(CN)422,

and CN2 at equilibrium when 0.50 g of Cd(NO3)2

dissolves in 5.0 3 102 mL of 0.50 M NaCN.

16.77 If NaOH is added to 0.010 M Al31, which will be the

predominant species at equilibrium: Al(OH)3 or

Al(OH)24 ? The pH of the solution is 14.00. [Kf for

33

Al(OH)2

4 5 2.0 3 10 .]

16.78 Calculate the molar solubility of AgI in a 1.0 M NH3

solution.

16.79 Both Ag1 and Zn21 form complex ions with NH3.

Write balanced equations for the reactions. However, Zn(OH)2 is soluble in 6 M NaOH, and AgOH is

not. Explain.

16.80 Explain, with balanced ionic equations, why (a) CuI2

dissolves in ammonia solution, (b) AgBr dissolves in

NaCN solution, (c) HgCl2 dissolves in KCl solution.



Qualitative Analysis

Review Questions

16.81 Outline the general procedure of qualitative analysis.

16.82 Give two examples of metal ions in each group

(1 through 5) in the qualitative analysis scheme.



Problems

16.83 In a group 1 analysis, a student obtained a precipitate containing both AgCl and PbCl2. Suggest one



761



reagent that would enable her to separate AgCl(s)

from PbCl2(s).

16.84 In a group 1 analysis, a student adds HCl acid to the

unknown solution to make [Cl2] 5 0.15 M. Some

PbCl2 precipitates. Calculate the concentration of

Pb21 remaining in solution.

16.85 Both KCl and NH4Cl are white solids. Suggest one

reagent that would enable you to distinguish between

these two compounds.

16.86 Describe a simple test that would enable you to distinguish between AgNO3(s) and Cu(NO3)2(s).



Additional Problems

16.87 The buffer range is defined by the equation

pH 5 pKa 6 1. Calculate the range of the ratio

[conjugate base]/[acid] that corresponds to this

equation.

16.88 The pKa of the indicator methyl orange is 3.46. Over

what pH range does this indicator change from

90 percent HIn to 90 percent In2?

16.89 Sketch the titration curve of a weak acid versus a

strong base like the one shown in Figure 16.5. On

your graph indicate the volume of base used at the

equivalence point and also at the half-equivalence

point, that is, the point at which half of the acid has

been neutralized. Show how you can measure the pH

of the solution at the half-equivalence point. Using

Equation (16.4), explain how you can determine the

pKa of the acid by this procedure.

16.90 A 200-mL volume of NaOH solution was added to

400 mL of a 2.00 M HNO2 solution. The pH of the

mixed solution was 1.50 units greater than that of

the original acid solution. Calculate the molarity of

the NaOH solution.

16.91 The pKa of butyric acid (HBut) is 4.7. Calculate Kb

for the butyrate ion (But2).

16.92 A solution is made by mixing 5.00 3 102 mL of

0.167 M NaOH with 5.00 3 102 mL 0.100 M

HCOOH. Calculate the equilibrium concentrations

of H1, HCOOH, HCOO2, OH2, and Na1.

16.93 Cd(OH)2 is an insoluble compound. It dissolves in excess NaOH in solution. Write a balanced ionic equation for this reaction. What type of reaction is this?

16.94 A student mixes 50.0 mL of 1.00 M Ba(OH)2 with

86.4 mL of 0.494 M H2SO4. Calculate the mass of

BaSO4 formed and the pH of the mixed solution.

16.95 For which of the following reactions is the equilibrium constant called a solubility product?

(a) Zn(OH) 2 (s) 1 2OH 2(aq) Δ

Zn(OH) 22

4 (aq)

21

32

(b) 3Ca (aq) 1 2PO 4 (aq) Δ Ca3(PO 4 ) 2(s)

(c) CaCO 3 (s) 1 2H 1 (aq) Δ

Ca21(aq) 1 H 2O(l) 1 CO 2 (g)

(d) PbI2 (s) Δ Pb21 (aq) 1 2I2 (aq)



762



Acid-Base Equilibria and Solubility Equilibria



Mass of HgI2 formed



16.96 A 2.0-L kettle contains 116 g of boiler scale (CaCO3).

How many times would the kettle have to be completely filled with distilled water to remove all of the

deposit at 258C?

16.97 Equal volumes of 0.12 M AgNO3 and 0.14 M ZnCl2

solution are mixed. Calculate the equilibrium concentrations of Ag1, Cl2, Zn21, and NO2

3.

16.98 Calculate the solubility (in g/L) of Ag2CO3.

16.99 Find the approximate pH range suitable for separating

Mg21 and Zn21 by the precipitation of Zn(OH)2 from

a solution that is initially 0.010 M in Mg21 and Zn21.

16.100 A volume of 25.0 mL of 0.100 M HCl is titrated

against a 0.100 M CH3NH2 solution added to it from

a buret. Calculate the pH values of the solution

(a) after 10.0 mL of CH3NH2 solution have been

added, (b) after 25.0 mL of CH3NH2 solution have

been added, (c) after 35.0 mL of CH3NH2 solution

have been added.

16.101 The molar solubility of Pb(IO 3) 2 in a 0.10 M

NaIO3 solution is 2.4 3 10211 mol/L. What is Ksp

for Pb(IO3)2?

16.102 When a KI solution was added to a solution of mercury(II) chloride, a precipitate [mercury(II) iodide]

formed. A student plotted the mass of the precipitate

versus the volume of the KI solution added and obtained the following graph. Explain the appearance

of the graph.



Volume of KI added



16.103 Barium is a toxic substance that can seriously impair

heart function. For an X ray of the gastrointestinal

tract, a patient drinks an aqueous suspension of 20 g

BaSO4. If this substance were to equilibrate with the

5.0 L of the blood in the patient’s body, what would

be [Ba21]? For a good estimate, we may assume that

the temperature is at 258C. Why is Ba(NO3)2 not

chosen for this procedure?

16.104 The pKa of phenolphthalein is 9.10. Over what pH

range does this indicator change from 95 percent HIn

to 95 percent In2?

16.105 Solid NaBr is slowly added to a solution that is 0.010 M

in Cu1 and 0.010 M in Ag1. (a) Which compound

will begin to precipitate first? (b) Calculate [Ag1]

when CuBr just begins to precipitate. (c) What percent of Ag1 remains in solution at this point?



16.106 Cacodylic acid is (CH3)2AsO2H. Its ionization constant is 6.4 3 1027. (a) Calculate the pH of 50.0 mL

of a 0.10 M solution of the acid. (b) Calculate the pH

of 25.0 mL of 0.15 M (CH3)2AsO2Na. (c) Mix the

solutions in part (a) and part (b). Calculate the pH of

the resulting solution.

16.107 Radiochemical techniques are useful in estimating

the solubility product of many compounds. In one

experiment, 50.0 mL of a 0.010 M AgNO3 solution

containing a silver isotope with a radioactivity of

74,025 counts per min per mL were mixed with

100 mL of a 0.030 M NaIO3 solution. The mixed solution was diluted to 500 mL and filtered to remove

all of the AgIO3 precipitate. The remaining solution

was found to have a radioactivity of 44.4 counts per

min per mL. What is the Ksp of AgIO3?

16.108 The molar mass of a certain metal carbonate, MCO3,

can be determined by adding an excess of HCl

acid to react with all the carbonate and then “backtitrating” the remaining acid with NaOH. (a) Write

an equation for these reactions. (b) In a certain

experiment, 20.00 mL of 0.0800 M HCl were added

to a 0.1022-g sample of MCO3. The excess HCl

required 5.64 mL of 0.1000 M NaOH for neutralization. Calculate the molar mass of the carbonate

and identify M.

16.109 Acid-base reactions usually go to completion. Confirm this statement by calculating the equilibrium

constant for each of the following cases: (a) A strong

acid reacting with a strong base. (b) A strong acid

reacting with a weak base (NH3). (c) A weak acid

(CH3COOH) reacting with a strong base. (d) A weak

acid (CH3COOH) reacting with a weak base (NH3).

(Hint: Strong acids exist as H1 ions and strong bases

exist as OH2 ions in solution. You need to look up

Ka, Kb, and Kw.)

16.110 Calculate x, the number of molecules of water in

oxalic acid hydrate, H2C2O4 ? xH2O, from the following data: 5.00 g of the compound is made up to

exactly 250 mL solution, and 25.0 mL of this solution requires 15.9 mL of 0.500 M NaOH solution for

neutralization.

16.111 Describe how you would prepare a 1-L 0.20 M

CH3COONa/0.20 M CH3COOH buffer system by

(a) mixing a solution of CH3COOH with a solution

of CH3COONa, (b) reacting a solution of CH3COOH

with a solution of NaOH, and (c) reacting a solution

of CH3COONa with a solution of HCl.

16.112 Phenolphthalein is the common indicator for the titration of a strong acid with a strong base. (a) If the

pKa of phenolphthalein is 9.10, what is the ratio of

the nonionized form of the indicator (colorless) to

the ionized form (reddish pink) at pH 8.00? (b) If



Questions and Problems



2 drops of 0.060 M phenolphthalein are used in a titration involving a 50.0-mL volume, what is the concentration of the ionized form at pH 8.00? (Assume

that 1 drop 5 0.050 mL.)

16.113 Oil paintings containing lead(II) compounds as constituents of their pigments darken over the years.

Suggest a chemical reason for the color change.

16.114 What reagents would you employ to separate the following pairs of ions in solution? (a) Na1 and Ba21,

(b) K1 and Pb21, (c) Zn21 and Hg21.

16.115 Look up the Ksp values for BaSO4 and SrSO4 in Table 16.2. Calculate the concentrations of Ba21, Sr21,

and SO22

4 in a solution that is saturated with both

compounds.

16.116 In principle, amphoteric oxides, such as Al2O3 and

BeO can be used to prepare buffer solutions because

they possess both acidic and basic properties (see

Section 15.11). Explain why these compounds are of

little practical use as buffer components.

16.117 CaSO4 (Ksp 5 2.4 3 1025) has a larger Ksp value than

that of Ag2SO4 (Ksp 5 1.4 3 1025). Does it follow

that CaSO4 also has greater solubility (g/L)?

16.118 When lemon juice is squirted into tea, the color becomes lighter. In part, the color change is due to dilution, but the main reason for the change is an

acid-base reaction. What is the reaction? (Hint: Tea

contains “polyphenols” which are weak acids and

lemon juice contains citric acid.)

16.119 How many milliliters of 1.0 M NaOH must be added

to a 200 mL of 0.10 M NaH2PO4 to make a buffer

solution with a pH of 7.50?

16.120 The maximum allowable concentration of Pb21 ions in

drinking water is 0.05 ppm (that is, 0.05 g of Pb21 in

one million grams of water). Is this guideline exceeded

if an underground water supply is at equilibrium with

the mineral anglesite, PbSO4 (Ksp 5 1.6 3 1028)?

16.121 One of the most common antibiotics is penicillin G

(benzylpenicillinic acid), which has the structure

shown next:

O

B

H

COOH

O

G D

J

C

H3C

G D NOC H

A A A

C

D G COC ONOCOCH2O

H3C

S A A

B

H H

O



It is a weak monoprotic acid:

HP Δ H1 1 P2



Ka 5 1.64 3 1023



where HP denotes the parent acid and P2 the conjugate base. Penicillin G is produced by growing molds



16.122



16.123



16.124



16.125



763



in fermentation tanks at 258C and a pH range of

4.5 to 5.0. The crude form of this antibiotic is obtained

by extracting the fermentation broth with an organic

solvent in which the acid is soluble. (a) Identify the

acidic hydrogen atom. (b) In one stage of purification, the organic extract of the crude penicillin G is

treated with a buffer solution at pH 5 6.50. What is

the ratio of the conjugate base of penicillin G to the

acid at this pH? Would you expect the conjugate

base to be more soluble in water than the acid?

(c) Penicillin G is not suitable for oral administration,

but the sodium salt (NaP) is because it is soluble.

Calculate the pH of a 0.12 M NaP solution formed

when a tablet containing the salt is dissolved in a

glass of water.

Which of the following solutions has the highest

[H1]? (a) 0.10 M HF, (b) 0.10 M HF in 0.10 M NaF,

(c) 0.10 M HF in 0.10 M SbF5. (Hint: SbF5 reacts

with F2 to form the complex ion SbF2

6 .)

Distribution curves show how the fractions of nonionized acid and its conjugate base vary as a function

of pH of the medium. Plot distribution curves for

CH3COOH and its conjugate base CH3COO2 in solution. Your graph should show fraction as the y axis

and pH as the x axis. What are the fractions and pH

at the point where these two curves intersect?

Water containing Ca21 and Mg21 ions is called hard

water and is unsuitable for some household and industrial use because these ions react with soap to

form insoluble salts, or curds. One way to remove

the Ca21 ions from hard water is by adding washing

soda (Na2CO3 ? 10H2O). (a) The molar solubility of

CaCO3 is 9.3 3 1025 M. What is its molar solubility

in a 0.050 M Na2CO3 solution? (b) Why are Mg21

ions not removed by this procedure? (c) The Mg21

ions are removed as Mg(OH)2 by adding slaked

lime [Ca(OH)2] to the water to produce a saturated

solution. Calculate the pH of a saturated Ca(OH)2

solution. (d) What is the concentration of Mg21 ions

at this pH? (e) In general, which ion (Ca21 or Mg21)

would you remove first? Why?

Consider the ionization of the following acid-base

indicator:

HIn(aq) Δ H1(aq) 1 In2 (aq)



The indicator changes color according to the ratios

of the concentrations of the acid to its conjugate base

as described on p. 733. Show that the pH range over

which the indicator changes from the acid color to

the base color is pH 5 pKa 6 1, where Ka is the ionization constant of the acid.

16.126 Amino acids are building blocks of proteins. These

compounds contain at least one amino group (—NH2)

and one carboxyl group (—COOH). Consider



764



Acid-Base Equilibria and Solubility Equilibria



glycine (NH2CH2COOH). Depending on the pH of

the solution, glycine can exist in one of three possible forms:

ϩ



Fully protonated: NH3—CH2—COOH

ϩ

Dipolar ion: NH —CH —COO–

3



2



Fully ionized: NH2—CH2—COO–



Predict the predominant form of glycine at pH 1.0,

7.0, and 12.0. The pKa of the carboxyl group is 2.3

and that of the ammonium group (—NH1

3 ) is 9.6.

16.127 (a) Referring to Figure 16.6, describe how you

would determine the pKb of the base. (b) Derive an



analogous Henderson-Hasselbalch equation relating

pOH to pKb of a weak base B and its conjugate acid

HB1. Sketch a titration curve showing the variation

of the pOH of the base solution versus the volume of

a strong acid added from a buret. Describe how you

would determine the pKb from this curve. (Hint: pKb 5

2log Kb.)

16.128 A 25.0-mL of 0.20 M HF solution is titrated with a

0.20 M NaOH solution. Calculate the volume of

NaOH solution added when the pH of the solution is

(a) 2.85, (b) 3.15, (c) 11.89. Ignore salt hydrolysis.



Special Problems



O

9.17

B 1.82

ϩ

H3NOCHOCOOH

A



CH2

6.00 A

ϩ

HN

NH



(a) Show stepwise ionization of histidine in solution

(Hint: The H1 ion will first come off from the strongest acid group followed by the next strongest acid

group and so on.) (b) A dipolar ion is one in which

the species has an equal number of positive and negative charges. Identify the dipolar ion in (a). (c) The

pH at which the dipolar ion predominates is called

the isoelectric point, denoted by pI. The isoelectric



point is the average of the pKa values leading to and

following the formation of the dipolar ion. Calculate

the pI of histidine. (d) The histidine group plays an

important role in buffering blood (see Chemistry in

Action on p. 724). Which conjugate acid-base pair

shown in (a) is responsible for this action?

16.132 A sample of 0.96 L of HCl at 372 mmHg and 228C is

bubbled into 0.034 L of 0.57 M NH3. What is the pH

of the resulting solution? Assume the volume of solution remains constant and that the HCl is totally

dissolved in the solution.

16.133 A 1.0-L saturated silver carbonate solution at 58C is

treated with enough hydrochloric acid to decompose the compound. The carbon dioxide generated

is collected in a 19-mL vial and exerts a pressure

of 114 mmHg at 258C. What is the Ksp of Ag2CO3

at 58C?

16.134 The titration curve shown here represents the titration of a weak diprotic acid (H2A) versus NaOH.

(a) Label the major species present at the marked

points. (b) Estimate the pKa1 and pKa2 values of the

acid.

14

12

10

pH



16.129 Draw distribution curves for an aqueous carbonic

acid solution. Your graph should show fraction

of species present as the y axis and pH as the x

axis. Note that at any pH, only two of the three

species (H2CO3, HCO23 , and CO322) are present in

appreciable concentrations. Use the pKa values in

Table 15.5.

16.130 One way to distinguish a buffer solution with an acid

solution is by dilution. (a) Consider a buffer solution made of 0.500 M CH3COOH and 0.500 M

CH3COONa. Calculate its pH and the pH after it has

been diluted 10-fold. (b) Compare the result in

(a) with the pHs of a 0.500 M CH3COOH solution

before and after it has been diluted 10-fold.

16.131 Histidine is one of the 20 amino acids found in proteins. Shown here is a fully protonated histidine molecule where the numbers denote the pKa values of

the acidic groups.



8

6

4

2

Volume of NaOH added



Answers to Practice Exercises



765



Answers to Practice Exercises

16.1 4.01; 2.15. 16.2 (a) and (c). 16.3 9.17; 9.20.

16.4 Weigh out Na2CO3 and NaHCO3 in mole ratio of

0.60 to 1.0. Dissolve in enough water to make up a 1-L

solution. 16.5 (a) 2.19, (b) 3.95, (c) 8.02, (d) 11.39.

16.6 5.92. 16.7 (a) Bromophenol blue, methyl orange,

methyl red, and chlorophenol blue; (b) all except thymol

blue, bromophenol blue, and methyl orange; (c) cresol red



and phenolphthalein. 16.8 2.0 3 10214. 16.9 1.9 3 1023 g/L.

16.10 No. 16.11 (a) . 1.6 3 1029 M, (b) . 2.6 3 1026 M.

16.12 (a) 1.7 3 1024 g/L, (b) 1.4 3 1027 g/L. 16.13 (a) More

soluble in acid solution, (b) more soluble in acid solution,

(c) about the same. 16.14 Zn(OH)2 precipitate will

form. 16.15 [Cu21] 5 1.2 3 10213 M, [Cu(NH3)21

4 ]5

0.017 M, [NH3] 5 0.23 M. 16.16 3.5 3 1023 mol/L.



CHEMICAL



Mystery

A Hard-Boiled Snack



M



ost of us have eaten hard-boiled eggs. They are easy to cook and nutritious. But when was

the last time you thought about the process of boiling an egg or looked carefully at a hardboiled egg? A lot of interesting chemical and physical changes occur while an egg cooks.

A hen’s egg is a complicated biochemical system, but here we will focus on the three

major parts that we see when we crack open an egg: the shell, the egg white or albumen, and

the yolk. The shell protects the inner components from the outside environment, but it has

many microscopic pores through which air can pass. The albumen is about 88 percent water

and 12 percent protein. The yolk contains 50 percent water, 34 percent fat, 16 percent protein,

and a small amount of iron in the form of Fe21 ions.

Proteins are polymers made up of amino acids. In solution, each long chain of a protein

molecule folds in such a way that the hydrophobic parts of the molecule are buried inside and

the hydrophilic parts are on the exterior, in contact with the solution. This is the stable or native

state of a protein which allows it to perform normal physiological functions. Heat causes protein molecules to unfold, or denature. Chemicals such as acids and salt (NaCl) can also denature proteins. To avoid contact with water, the hydrophobic parts of denatured proteins will

clump together, or coagulate to form a semirigid opaque white solid. Heating also decomposes

some proteins so that the sulfur in them combines with hydrogen to form hydrogen sulfide

(H2S), an unpleasant smelling gas that can sometimes be detected when the shell of a boiled

egg is cracked.

The accompanying photo of hard-boiled eggs shows an egg that has been boiled for about

12 minutes and one that has been overcooked. Note that the outside of the overcooked yolk is

green.

What is the chemical basis for the changes brought about by boiling an egg?



Chemical Clues

1.



Schematic diagram of an egg. The chalazae are the cords that anchor the yolk

to the shell and keep it centered.



One frequently encountered problem with hard-boiled eggs is that their shells crack in

water. The recommended procedure for hard boiling eggs is to place the eggs in cold water



Shell

Membrane

Albumen

Yolk membrane

Yolk

Chalaza

Air space



766



and then bring the water to a boil. What causes the shells to crack in this case? How does

pin holing, that is, piercing the shell with a needle, prevent the shells from cracking? A

less satisfactory way of hard boiling eggs is to place room-temperature eggs or cold eggs

from the refrigerator in boiling water. What additional mechanism might cause the shells

to crack?

2.



When an eggshell cracks during cooking, some of the egg white leaks into the hot water

to form unsightly “streamers.” An experienced cook adds salt or vinegar to the water prior

to heating eggs to minimize the formation of streamers. Explain the chemical basis for

this action.



3.



Identify the green substance on the outer layer of the yolk of an overcooked egg and write

an equation representing its formation. The unsightly “green yolk” can be eliminated or

minimized if the overcooked egg is rinsed with cold water immediately after it has been

removed from the boiling water. How does this action remove the green substance?



4.



The way to distinguish a raw egg from a hard-boiled egg, without cracking the shells, is

to spin the eggs. How does this method work?



A 12-minute egg (left) and an overcooked hardboiled egg (right).



Iron(II) sulfide.



767



Chemistry in the

Atmosphere



Lightning causes atmospheric

nitrogen and oxygen to form nitric

oxide, which is eventually

converted to nitrates. The models

show nitrogen, oxygen, and nitric

oxide molecules.



Chapter Outline



A Look Ahead





We begin by examining the regions and composition of Earth’s atmosphere.

(17.1)







We then study a natural phenomenon—aurora borealis—and a human-made

phenomenon—the glow of space shuttles—in the outer layers of the atmosphere. (17.2)



Depletion of Ozone in the

Stratosphere







Next, we study the depletion of ozone in the stratosphere and its detrimental

effects and ways to slow the progress. (17.3)



Volcanoes







Focusing on events in the troposphere, we first examine volcanic eruptions.

(17.4)







We study the cause and effect of greenhouse gases and ways to curtail the

emission of carbon dioxide and other harmful gases. (17.5)







We see that acid rain is largely caused by human activities such as the burning of fossil fuels and roasting of metal sulfides. We discuss ways to minimize sulfur dioxide and nitrogen oxides productions. (17.6)







Another human-made pollution is smog formation, which is the result of the

heavy use of automobiles. We examine mechanisms of smog formation and

ways to reduce the pollution. (17.7)







Finally, we consider some examples of indoor pollutants such as radon,

carbon dioxide and carbon monoxide, and formaldehyde. (17.8)



17.1

17.2



Earth’s Atmosphere



17.3

17.4

17.5

17.6

17.7

17.8



Phenomena in the Outer

Layers of the Atmosphere



The Greenhouse Effect

Acid Rain

Photochemical Smog

Indoor Pollution



Student Interactive

Activities

Media Player

Chapter Summary

ARIS

Example Practice Problems

End of Chapter Problems



W



e have studied basic definitions in chemistry, and we have examined the

properties of gases, liquids, solids, and solutions. We have discussed

chemical bonding and intermolecular forces and seen how chemical kinetics and

chemical equilibrium concepts help us understand the nature of chemical reactions. It is appropriate at this stage to apply our knowledge to the study of one

extremely important system: the atmosphere. Although Earth’s atmosphere is

fairly simple in composition, its chemistry is very complex and not fully understood. The chemical processes that take place in our atmosphere are induced by

solar radiation, but they are intimately connected to natural events and human

activities on Earth’s surface.

In this chapter, we will discuss the structure and composition of the atmosphere, together with some of the chemical processes that occur there. In addition,

we will take a look at the major sources of air pollution and prospects for controlling them.



769



770



Chemistry in the Atmosphere



17.1 Earth’s Atmosphere



TABLE 17.1

Composition of Dry Air

at Sea Level

Gas



Composition

(% by Volume)



N2

O2

Ar

CO2

Ne

He

Kr

Xe



78.03

20.99

0.94

0.033

0.0015

0.000524

0.00014

0.000006



Earth is unique among the planets of our solar system in having an atmosphere that is

chemically active and rich in oxygen. Mars, for example, has a much thinner atmosphere

that is about 90 percent carbon dioxide. Jupiter, on the other hand, has no solid surface;

it is made up of 90 percent hydrogen, 9 percent helium, and 1 percent other substances.

It is generally believed that three billion or four billion years ago, Earth’s atmosphere consisted mainly of ammonia, methane, and water. There was little, if any, free

oxygen present. Ultraviolet (UV) radiation from the sun probably penetrated the atmosphere, rendering the surface of Earth sterile. However, the same UV radiation may

have triggered the chemical reactions (perhaps beneath the surface) that eventually led

to life on Earth. Primitive organisms used energy from the sun to break down carbon

dioxide (produced by volcanic activity) to obtain carbon, which they incorporated in

their own cells. The major by-product of this process, called photosynthesis, is oxygen.

Another important source of oxygen is the photodecomposition of water vapor by UV

light. Over time, the more reactive gases such as ammonia and methane have largely

disappeared, and today our atmosphere consists mainly of oxygen and nitrogen gases.

Biological processes determine to a great extent the atmospheric concentrations of these

gases, one of which is reactive (oxygen) and the other unreactive (nitrogen).

Table 17.1 shows the composition of dry air at sea level. The total mass of the

atmosphere is about 5.3 3 1018 kg. Water is excluded from this table because its

concentration in air can vary drastically from location to location.



Atmospheric nitrogen



Atmospheric fixation



Fixed juvenile

nitrogen



Igneous rocks



Industrial

fixation



Protein



Biological

fixation



Nitrate reduction

Plant and animal wastes,

dead organisms



Denitrification

Nitrous

oxide



Ammonium

Nitrite

Nitrate



Figure 17.1



To ground

water



The nitrogen cycle. Although the supply of nitrogen in the atmosphere is virtually inexhaustible, it must be combined

with hydrogen or oxygen before it can be assimilated by higher plants, which in turn are consumed by animals. Juvenile nitrogen is

nitrogen that has not previously participated in the nitrogen cycle.



771



17.1 Earth’s Atmosphere



Figure 17.1 shows the major processes involved in the cycle of nitrogen in nature.

Molecular nitrogen, with its triple bond, is a very stable molecule. However, through

biological and industrial nitrogen fixation, the conversion of molecular nitrogen into

nitrogen compounds, atmospheric nitrogen gas is converted into nitrates and other

compounds suitable for assimilation by algae and plants. Another important mechanism for producing nitrates from nitrogen gas is lightning. The steps are

electrical



N2 (g) 1 O2 (g) OO¡

2NO(g)

energy

2NO(g) 1 O2 (g) OO¡ 2NO2 (g)

2NO2 (g) 1 H2O(l) OO¡ HNO2 (aq) 1 HNO3 (aq)

About 30 million tons of HNO3 are produced this way annually. Nitric acid is converted to nitrate salts in the soil. These nutrients are taken up by plants, which in turn

are ingested by animals. Animals use the nutrients from plants to make proteins and

other essential biomolecules. Denitrification reverses nitrogen fixation to complete the

cycle. For example, certain anaerobic organisms decompose animal wastes as well as

dead plants and animals to produce free molecular nitrogen from nitrates.

The main processes of the global oxygen cycle are shown in Figure 17.2. This

cycle is complicated by the fact that oxygen takes so many different chemical forms.

Atmospheric oxygen is removed through respiration and various industrial processes

(mostly combustion), which produce carbon dioxide. Photosynthesis is the major mechanism by which molecular oxygen is regenerated from carbon dioxide and water.

Scientists divide the atmosphere into several different layers according to temperature variation and composition (Figure 17.3). As far as visible events are concerned, the most active region is the troposphere, the layer of the atmosphere that



High-energy ultraviolet radiation

O

O2



H

O2



H2O



H2O

O2



Ozone screen

O2



CO2



O2



Volcanism



Oxidative

weathering

4FeO + O2



CO2

Photic zone

H2O + CO2



H2CO3



+



H+



2CO2

CO



O



CO2



Phytoplankton



HCO3–



O2 + 2CO



OH



O3



2Fe2O3



Sediments

2HCO3–

H2O



CO23 –

Ca2+



CaCO3



Sediments



Figure 17.2



The oxygen cycle. The cycle is complicated because oxygen appears in so many chemical forms and combinations,

primarily as molecular oxygen, in water, and in organic and inorganic compounds.



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