The Chemistry of Air Pollution – A Molecular Analysis

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CHEM 131
Experiment
5
The Chemistry of Air Pollution – A Molecular Analysis
Your goals in this experiment are to use the concepts of chemical bonding, including Lewis Structures and bond
energies, in order to gain a better understanding of various air pollution problems.
For this lab, you will draw Lewis structures and make calculations directly on the report sheets (in pencil). Note
that the questions are numbered and indicated in boldface type throughout the pollution discussion.
Corresponding numbers are indicated on the report sheet, next to the location where you should provide your
answers or show your calculations.
In order to carry out this lab, you will need to review a number of conceptual and problem-solving skills, that
span chapters 5 through 8 in your textbook. The most important of these skills are addressed in the following
pre-lab questions.
For more background information regarding air pollution and other environmental chemistry issues, see The
Chemistry of the Environment interchapter in your textbook on pages 949-961.
Pre-Lab Questions
1. Many reactions involving chlorine gas are photoinitiated; that is, they start when light breaks a chlorine
molecule into atoms.
(a) Using the standard enthalpies of formation given in Appendix L of your textbook, calculate ΔH° for the
reaction:
Cl2 (g) ⎯→ 2 Cl (g)
(b) Look up the bond enthalpy of a Cl−Cl bond, D(Cl−Cl), in Table 8.9 in your textbook. How does
D(Cl−Cl) compare to the ΔH° value calculated in part (a)?
(c) Show that the maximum wavelength of light needed to break a Cl−Cl bond is 491 nm. (Hint: note that 1
mol of photons is needed to break 1 mol of Cl−Cl bonds. Work backwards through the calculation
scheme on p. 274 to solve for the wavelength that corresponds to this amount of energy.)
2. Average bond enthalpies can be used to estimate ΔH for reactions of more complex molecules, such as the
combustion of methane.
CH4 (g) + 2 O2 (g) ⎯→ CO2 (g) + 2 H2O (g)
(a) Using the standard enthalpies of formation given in Appendix L of your textbook, calculate ΔH° for the
combustion of methane.
(b) Draw Lewis structures for CH4, O2, CO2, and H2O. (See section 8.2 in your textbook for an explanation
of how to draw Lewis structures.)
(c) Using the average bond enthalpies given in Table 8.9 in your textbook, obtain a new estimate for ΔH°
for the combustion of methane. (See Example 8.14 in your textbook.)
(d) Calculate the percent error in the ΔH° value obtained from the average bond enthalpies. Why does this
ΔH° value contain an error?
CHEM 131
Experiment 5
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The Chemistry of Air Pollution – A Molecular Analysis
[This] most excellent canopy, the air, look you, this brave o’erhanging firmament, this majestical roof fretted
with golden fire—why, it appears no other thing to me than a foul and pestilent congregation of vapours.1
While a human can survive without food for weeks and without water for days, we can only last a matter of
minutes without air. A constant supply of oxygen, provided by the earth’s atmosphere, is absolutely essential to
our basic cellular metabolic functions. But the atmosphere is important not only for its supply of oxygen; it also
serves as a protective blanket that shields us from damaging radiation, maintains relatively constant, comfortable
temperatures, and serves as an important component of the earth’s hydrologic cycle. Unfortunately, some
human activities (especially the operation of motor vehicles) can disrupt the abilities of the atmosphere to
protect us and can thereby expose us to a variety of health hazards.
The atmosphere can be divided into a number of different layers, based upon temperature profiles (see p 534 in
your textbook). Generation of air pollutants primarily takes place at or near the surface of the earth in the
troposphere. There are a variety of important pollution problems that consequently result in the troposphere,
including photochemical smog, acid rain, and the enhanced greenhouse effect. Some of the pollutants gradually
make their way up into the stratosphere and there disrupt a sequence of naturally occurring chemical reactions
that involve the life cycle of the ozone molecule. These naturally occurring reactions, together with some
additional reactions in the thermosphere (or ionosphere), are important in protecting us from damaging forms of
solar radiation.
GENERATION OF POLLUTANTS
Most of the serious air pollution problems result directly or indirectly from the combustion of fossil fuels to
provide for a variety of energy needs in society. Fossil fuels are primarily hydrocarbons (such as methane),
which react exothermically with oxygen to produce carbon dioxide and water vapor. The combustion of
methane was considered in a pre-lab question.
In addition to the production of CO2, which itself acts as a greenhouse gas, incomplete combustion of fossil
fuels can result in the release of CO (which is a potent respiratory toxin) and various small reactive
hydrocarbons. Less obvious, however, is the generation of other nonmetal oxides, specifically the oxides of
nitrogen and sulfur.
NOx: In the extremely high temperatures present in automobile engines, molecular nitrogen and oxygen gases,
present in the air, can react with each other to produce nitrogen monoxide. The nitrogen monoxide then reacts
quickly with additional oxygen to form nitrogen dioxide. Both of these nitrogen oxides (known symbolically as
NOx) are very toxic to humans because of their high chemical reactivities.
1. Draw Lewis structures for each of these 2 compounds (Hint: in nitrogen dioxide, the N has the odd
number of electrons, and experimental evidence has shown that there are 2 equal bond lengths in the
molecule). What about their Lewis structures indicates that they would be very chemically reactive?
(As an interesting side note, although NO is extremely toxic to humans at even moderate concentrations, it is
vital to normal physiological function at very low concentrations. The body itself synthesizes very small
amounts of NO, which acts as a vasodilator to increase blood flow when needed. The drug Viagara blocks the
body’s normal breakdown of NO, thereby allowing one to maintain, for an increased period of time, the
increased blood flow that is necessary to maintain an erection.)
1
Hamlet, Act II, Scene 2, Hamlet, William Shakespeare
CHEM 131
Experiment 5
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The Chemistry of Air Pollution – A Molecular Analysis
While NOx form readily in the high temperatures of internal combustion engines (and during lightning storms),
they do not form normally in the air, even though large quantities of N2 and O2 are constantly present. Notice
that, in order for the reaction to occur, the bond in the N2 molecule must first be broken before the products can
be formed.
2. Draw the Lewis structure for N2. What is the relative strength of the bond in this molecule compared
to others in the table?
Another reason that NOx do not form readily in the air at normal temperatures has to do with the energetics of
the reaction.
3. Write a balanced reaction equation for the formation of NO. Use bond dissociation enthalpies to
estimate the reaction enthalpy (the bond enthalpy in NO is 631 kJ/mol). Which are more stable (i.e.,
at a lower energy state), the reactants or the products?
One common oxide of nitrogen is not toxic. N2O, known as nitrous oxide (or “laughing gas”), is commonly
used as an anesthetic. N2O is not a major product of internal combustion engines; however, it is released from
the oceans and from soils, and its presence in the atmosphere appears to be increasing due to the use of nitrogenbased fertilizers in agriculture (and from the production of nylon). While it does not pose a direct health hazard
to humans due to toxicity, it is an important greenhouse gas.
4. Draw the Lewis structures for N2O (note that there are 2 ways the atoms can be connected). Use
formal charges to help decide on the most important Lewis structure. What do you notice about its
Lewis structure, compared to those of the other NOx, that might suggest its lesser chemical reactivity?
SOx: Globally, the primary sources of sulfur oxides are volcanoes and the decomposition of plants. However,
since these natural generations take place either high up in the troposphere or far from urban areas, they are not
major contributors to the pollution problems that plague cities. These problems are due primarily to
anthropogenic (human caused) generation of sulfur oxides through the combustion of fossil fuels (especially
coal and oil) and the smelting of sulfide metal ores. Sulfur oxides themselves pose health hazards to humans by
causing and aggravating various respiratory ailments. Additionally, they are a major contributor to acid rain,
which is discussed in more detail below.
5. Draw Lewis structures for all the compounds that can be formed from a single S atom, and any
number of O atoms directly attached to the S. Which of these compounds do you predict to be stable
based upon their Lewis Structures? Look in the Handbook of Chemistry and Physics, or Merck
Index, to verify that these compounds commonly exist. You will notice that one of the compounds you
predicted to be stable does not appear. This compound can react exothermically with itself to re-form
elemental sulfur and oxygen, or one of the other oxide compounds.
PHOTOCHEMICAL SMOG
Photochemical smog is caused by a complex sequence of reactions, many initiated by sunlight, that convert a
mixture of NO, O2, and volatile organic compounds (VOC’s) into ozone (O3), nitric acid, and partially oxidized
organics (such as formaldehyde and acetaldehyde). The most important VOC’s, which result from auto
emissions and evaporation of solvents, are hydrocarbons that are particularly reactive. These hydrocarbons owe
their reactivity to a particularly reactive type of bond that is present in many organic molecules, including the
molecule ethene (or ethylene). (It is interesting to note that, while ethylene is a VOC that is emitted by
combustion processes, it is also released by plants in fairly significant quantities. However, the ethylene emitted
by plants occurs primarily in rural areas and therefore contributes very little to the problems of photochemical
smog in urban environments.)
CHEM 131
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The Chemistry of Air Pollution – A Molecular Analysis
6. Draw the Lewis structure of ethylene (C2H4). What bond is present between the C’s in this molecule?
It is at first surprising that this bond should be more reactive than a C-C single bond, since it has greater bond
energy. However, when molecules with this type of bond react, it is usually only a portion of the bond (the
“second” bond of the double bond, known in chemistry as a “pi” bond) that is broken.
7. Use the bond energies for C-C single bonds and C=C double bonds to estimate the bond energy of the
pi bond (the second bond in the double bond)? How does the energy compare to that of a C-C single
bond or a C-H bond?
Although organic compounds such as acetaldehyde and formaldehyde do result from photochemical smog, their
concentrations are sufficiently low that they are not considered a major health hazard in typical outdoor air
pollution. Formaldehyde (CH2O) is a concern, however, in indoor air pollution, as it is released from a large
number of household materials. These include various foams (including foam insulation), adhesives in plywood
and particleboard, carpet glues and dyes, and various other fabrics. Although the actual health consequences of
low-level formaldehyde exposure (typical in homes) are still being debated and investigated, some believe that it
may lead to eye irritation, increases in childhood asthma, and higher risks of developing cancer.
8. Draw the Lewis structure of formaldehyde, CH2O (also known as methanal).
A related compound, C2H4O, acetaldehyde (or ethanal), can be produced as a result of the incomplete
combustion of ethanol, and this fact is one of the arguments against the use of ethanol-based fuels.
Interestingly, our own bodies produce this compound in the process of breaking down the ethanol in alcoholic
beverages. The health effects of acetaldehyde exposure (headache and nausea) are familiar to many people who
have, at one time or another, had too much to drink; these effects are more commonly known as a “hangover.”
9. Draw the Lewis structures of ethanol, C2H5OH (Hint: one of the H’s is connected to the O) and
acetaldehyde, C2H4O (Hint: note that the C to O bonding is the same as in formaldehyde).
Another product of photochemical smog that was mentioned is ozone (O3). Ozone both impairs respiratory
function in humans and causes damage to plants. While ozone is an undesirable constituent of air in the
troposphere where we live, it is better known for its beneficial attributes in the upper atmosphere (described
below).
10. Draw the Lewis Structure of the ozone molecule.
A NATURAL SUNSCREEN
One important feature of our atmosphere is its ability to protect us from damaging electromagnetic (EM) energy
coming from the sun. Atoms and molecules can absorb photons of EM energy in order to promote electrons to
higher energy levels. If the photon is of sufficiently short wavelength, an electron can be completely removed
from an atom or molecule, thereby creating an ion. This process, known as photoionization, occurs in the
highest level of the atmosphere (the thermosphere or “ionosphere”), and most of the radiation in this wavelength
range has been screened out below 90 km. The first ionizations of N2 and O2 molecules is indicated by the
reaction equations below:
CHEM 131
N2 (g) + hν Æ N2+ (g) + e-
ΔE = 1495 kJ/mol
O2 (g) + hν Æ O2+ (g) + e-
ΔE = 1205 kJ/mol
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The Chemistry of Air Pollution – A Molecular Analysis
11. Calculate the longest wavelength of light that is capable of ionizing each of these molecules.
From about 30-90 km (the upper stratosphere and mesosphere), the primary mechanism for screening out EM
radiation is photodissociation of O2. In this process, absorption of photons of sufficient energy can break the
bond in the molecule, splitting the molecule into its atoms:
O2 (g) + hν Æ 2O (g)
12. Use the bond energy of an oxygen molecule to calculate the longest wavelength photon of light that
could cause photodissociation of O2.
In addition to photoionization and photodissociation, one other important factor in screening high energy
radiation is the “ozone layer,” a region of high O3 concentration in the stratosphere. The free oxygen atoms that
are produced from photodissociation of O2 quickly react with other oxygen molecules to create ozone:
O (g) + O2 (g) Æ O3 (g)
The ozone itself undergoes photodissociation that effectively reverses the above reaction:
O3 (g) + hν Æ O (g) + O2 (g)
However, much less energy (and therefore longer wavelengths) is required for this dissociation than that for O2.
As a result, O3 effectively screens out the remaining radiation in the “UV-C” range (defined as 200-280 nm).
While ozone does absorb some of the radiation in the UV-B range (280-320 nm), most of this radiation, as well
as that in the UV-A range (320-400 nm), reaches the earth’s surface (hence the need for UVA/UVB sunscreen).
Decreases in the amount of O3 in the stratosphere can lead to large increases in the amount of UVB and UVC
radiation that reaches us. Because radiation of these wavelengths is effective at disrupting the structure of DNA
(kinks in the DNA chain are created when adjacent bases become covalently bonded), damage to the ozone layer
leads to increases in the incidence of skin cancer.
CFC’s and Ozone Depletion: In the 1970’s and 1980’s it was discovered that chlorofluorocarbons (CFC’s),
used as refrigerant fluids and spray can propellants, were catalyzing the decomposition of ozone in the
stratosphere. One example of a CFC is CF2Cl2 (Freon-12).
13. Draw the Lewis structure for CF2Cl2.
The reaction mechanism for the accelerated decomposition was elucidated by F. Sherwood Rowland and Mario
Molina in 1974. The first step in the process is the photodissociation of a CFC molecule, which generates a free
Cl atom:
CF2Cl2 (g) + hν Æ CF2Cl (g) + Cl (g)
The free Cl atoms act as a catalyst by reacting with O3 and converting it to O2. Subsequent reactions regenerate
free Cl atoms, which then go on to destroy more ozone molecules. A small increase in the number of Cl atoms
can therefore lead to a major reduction in ozone.
14. Use bond energies to determine the longest wavelength of light capable of causing the
photodissociation above. Replacement of CFC’s with substitutes that do not contain Cl will help
prevent this dissociation from occurring. Why is the presence of F in the molecule less of an issue
than Cl?
CHEM 131
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The Chemistry of Air Pollution – A Molecular Analysis
PRODUCTION OF ACID RAIN
Nonmetal oxides react with water to produce oxyacids. CO2, NOx, and SOx are no exception. The CO2, NO2,
SO2, and SO3 in the atmosphere react with water in raindrops, producing acid rain.
When a nonmetal oxide reacts with water to produce an acid, the oxidation number of the nonmetal does not
change. This allows you to predict which acid will be produced. (Review pages 200-201 in your textbook,
which show how to assign oxidation numbers to the atoms in a compound.)
As an example, consider two other oxides of nitrogen, N2O3 and N2O5.
N2O3 (oxidation numbers: N = +3, O = −2)
N2O5 (oxidation numbers: N = +5, O = −2)
We know that nitrogen has two common oxyacids, HNO3 and HNO2.
HNO3 (oxidation numbers: H = +1, N = +5, O = −2)
HNO2 (oxidation numbers: H = +1, N = +3, O = −2)
By matching the oxidation numbers, we see that N2O3 will react with water to form HNO2, and N2O5
will react to form HNO3. The balanced equations are:
N2O3 + H2O ⎯→ 2 HNO2
N2O5 + H2O ⎯→ 2 HNO3
Knowing that a substance is an acid tells you something about its Lewis structure. Acids release H+ ions, so the
bond to H must be highly polar. This means the H must be bonded to the most electronegative atom in the
molecule, which is usually O.
Here are two possible Lewis structures for HNO2:
H
H
H
O
N
O
O
N
O
O
(1)
N
O
(2)
Since we know HNO2 is an acid, the structure (1), with H attached to O, must be correct.
15. Think of two more reasons why structure (1) is favored over structure (2) for HNO2. The first reason
has to do with formal charge. The second reason has to do with bond strengths.
Even without human activities, rainwater is slightly acidic due to the CO2 naturally present in the atmosphere.
The pH of pure water is 7.0, and atmospheric CO2 can lower it to about 5.6. Here in the Bay Area, our rainwater
has a pH of about 6.0, very close to neutral.
16. Write a balanced equation showing CO2 reacting with water to produce an acid. Draw the Lewis
structure for the acid.
Due to SO2 and SO3 emissions from coal-burning power plants, rainwater in some areas of the northeastern US
is quite a bit more acidic. In New England, downwind from coal-burning power plants in Pennsylvania, Ohio,
and West Virginia, the pH of rainwater is about 4.5, quite a bit more acidic.
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The Chemistry of Air Pollution – A Molecular Analysis
17. Write two balanced equations showing SO2 and SO3 reacting with water to produce acids. Draw
Lewis structures for each acid.
ACID RAIN AND THE ENVIRONMENT
Plants obtain some essential nutrients, such as Ca2+ and Mg2+ ions, by releasing carbonic acid into the soil and
absorbing dissolved Ca2+ and Mg2+ ions through their roots. Sources of calcium and magnesium include CaCO3
(limestone), MgCO3 (magnesite), and CaMg(CO3)2 (dolomite). This reaction shows how plants extract Ca2+
ions from limestone (remember that if a product is indicated as aq, then it exists as ions in solution):
CaCO3 (s) + H2CO3 (aq) ⎯→ Ca(HCO3)2 (aq)
Highly acidic rainwater dissolves these minerals more quickly, and washes the needed nutrients out of the soil.
18. Draw Lewis structures for the ionic compounds CaCO3 and Ca(HCO3)2. Why is Ca(HCO3)2 more
soluble in water than CaCO3? (Think about the strengths of the ionic bonds holding the solids
together.)
Plants obtain K+ ions, another essential nutrient, in a similar way. In this case, the carbonic acid released by the
plants reacts with common silicate minerals (feldspar, KAlSi3O8) to produce clay (kaolinite, Al2Si2O5(OH)4),
releasing dissolved K+ ions.
2 KAlSi3O8 (s) + 2 H2CO3 (aq) + H2O (l) ⎯→ Al2Si2O5(OH)4 (s) + 4 SiO2 (s) + 2 KHCO3 (aq)
Highly acidic rainwater, containing H2SO4 or HNO3, not only washes the K+ ions out of the soil, it also reacts
with clay to release dissolved Al3+ ions, which are toxic to plants.
Al2Si2O5(OH)4 (s) + 3 H2SO4 (aq) ⎯→ Al2(SO4)3 (aq) + 2 SiO2 (s) + 5 H2O (l)
Al3+ ions are toxic to plants because they replace Ca2+ ions as cofactors in plant enzymes. Ca2+ ions normally
attach to enzymes by forming ionic bonds with negatively charged groups on the enzyme surface.
19. Give two reasons why Al3+ can form stronger ionic bonds than Ca2+.
Strongly acidic rainwater not only poisons plants by releasing Al3+ ions, it can also contaminate groundwater by
dissolving minerals that contain heavy metals. For example, H2SO4 in polluted rainwater can dissolve cinnabar,
HgS, an ore of mercury, releasing dissolved Hg2+ ions into the environment.
HgS (s) + H2SO4 (aq) ⎯→ H2S (g) + HgSO4 (aq)
20. Draw Lewis structures for the ionic compounds HgS and HgSO4. Why is HgSO4 more soluble in
water than HgS? (Think about the strengths of the ionic bonds holding the solids together.)
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The Chemistry of Air Pollution – A Molecular Analysis
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CHEM 131
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The Chemistry of Air Pollution – A Molecular Analysis
REPORT SHEET
NAME:
1.
NO
NO2
Explanation of Reactivity:
2.
N2
Relative bond strength of N2:
3. Balanced Equation:
Lewis Structures:
Enthalpy Calculation:
Lower in Energy?
4. Circle most likely structure(s):
Explanation of Reactivity:
CHEM 131
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The Chemistry of Air Pollution – A Molecular Analysis
5.
SO
SO2
SO3
SO4
Predicted stable compounds (Circle compound that did not appear):
6.
C2H4
7. Calculation of pi bond energy:
Relative strength?
8. CH2O
CHEM 131
9. C2H5OH
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The Chemistry of Air Pollution – A Molecular Analysis
9. C2H4O
10. O3
11. λmax (N2 ionization):
λmax (O2 ionization):
12. λmax (O2 dissociation):
13. CF2Cl2
14. λmax (C-Cl dissociation):
Reason why F is less of a problem?
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The Chemistry of Air Pollution – A Molecular Analysis
15. Two reasons why (1) is better than (2):
1.
2.
16. Balanced Equation:
Oxyacid with C
17. Balanced Equations:
Acid from SO2
Acid from SO3
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The Chemistry of Air Pollution – A Molecular Analysis
18.
CaCO3
Ca(HCO3)2
Why is Ca(HCO3)2 more soluble?
19. Reason 1:
Reason 2:
20.
Why is HgSO4 more soluble?
HgS
HgSO4
CHEM 131
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The Chemistry of Air Pollution – A Molecular Analysis
Additional Question (to answer at home after the lab period)
Draw Lewis structures for the following compounds and ions: NO+, NO2−, N2O, NO3−, NOCl.
Rank them in order from shortest to longest N−O bond. Explain the reasons for your ranking
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