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3 Bohr's Theory of the Hydrogen Atom
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
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.
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,
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
4 5 2.0 3 10 .]
16.78 Calculate the molar solubility of AgI in a 1.0 M NH3
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
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.
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.
16.83 In a group 1 analysis, a student obtained a precipitate containing both AgCl and PbCl2. Suggest one
reagent that would enable her to separate AgCl(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).
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
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
(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)
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
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
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
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,
4 in a solution that is saturated with both
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
G D NOC H
A A A
D G COC ONOCOCH2O
S A A
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
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
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
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
Acid-Base Equilibria and Solubility Equilibria
glycine (NH2CH2COOH). Depending on the pH of
the solution, glycine can exist in one of three possible forms:
Fully protonated: NH3—CH2—COOH
Dipolar ion: NH —CH —COO–
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
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.
(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
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
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
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.
Volume of NaOH added
Answers to Practice Exercises
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
0.017 M, [NH3] 5 0.23 M. 16.16 3.5 3 1023 mol/L.
A Hard-Boiled Snack
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
What is the chemical basis for the changes brought about by boiling an egg?
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
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
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
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?
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).
Chemistry in the
Lightning causes atmospheric
nitrogen and oxygen to form nitric
oxide, which is eventually
converted to nitrates. The models
show nitrogen, oxygen, and nitric
A Look Ahead
We begin by examining the regions and composition of Earth’s atmosphere.
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
Next, we study the depletion of ozone in the stratosphere and its detrimental
effects and ways to slow the progress. (17.3)
Focusing on events in the troposphere, we first examine volcanic eruptions.
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)
Phenomena in the Outer
Layers of the Atmosphere
The Greenhouse Effect
Example Practice Problems
End of Chapter Problems
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.
Chemistry in the Atmosphere
17.1 Earth’s Atmosphere
Composition of Dry Air
at Sea Level
(% by Volume)
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.
Plant and animal wastes,
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.
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
N2 (g) 1 O2 (g) OO¡
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
4FeO + O2
H2O + CO2
O2 + 2CO
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.