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261
Chapter 18. Chemistry of the Environment
Media Resources
Figures and Tables in Transparency Pack: Section:
Figure 18.1 Temperature and Pressure in the 18.1 Earth’s Atmosphere
Atmosphere Vary as a Function of Altitude
Above Sea Level
Figure 18.6 Ozone Present in the Southern 18.2 Human Activities and Earth’s Atmosphere
Animations: Section:
Stratospheric Ozone 18.1 Earth’s Atmosphere
CFCs and Stratospheric Ozone 18.2 Human Activities and Earth’s Atmosphere
Catalytic Destruction of Stratospheric Ozone 18.2 Human Activities and Earth’s Atmosphere
Movies: Section:
Carbon Dioxide Behaves as an Acid in Water 18.2 Human Activities and Earth’s Atmosphere
3-D Models: Section:
Nitrogen 18.1 Earth’s Atmosphere
Other Resources
Further Readings: Section:
Introducing Atmospheric Reactions: A Systematic 18.1 Earth’s Atmosphere
Approach for Students
Chapter 18
262
Chemists Celebrate Earth Day 2009: Air—The 18.1 Earth’s Atmosphere
Sky’s the Limit. JCE Resources for Chemistry
and the Atmosphere: An Update
Consumption and CO2 Emissions to Illustrate
Chemical Principles
The Expiration of Respiration: Oxygen–The 18.4 Human Activities and Earth’s Water
Missing Ingredient in Many Bodies of Water
A Discovery-Based Experiment Illustrating How 18.4 Human Activities and Earth’s Water
Iron Metal Is Used to Remediate Contaminated
Groundwater
Live Demonstrations: Section:
Acid-Neutralizing Capacity of Lake Beds 18.4 Human Activities and Earth’s Water
Going Green: Lecture Assignments and Lab 18.5 Green Chemistry
Experiences for the College Curriculum
Chemistry of the Environment
263
Chapter 18. Chemistry of the Environment
Common Student Misconceptions
• Students often confuse the toxicity of ozone in the troposphere and the beneficial effects of ozone in
Teaching Tips
• This chapter gives an excellent opportunity to review Dalton’s law of partial pressure, Planck’s law
calculations, ionization energy, equilibrium, reaction mechanisms, and some acid-base chemistry.
Lecture Outline
18.1 Earth’s Atmosphere
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• The temperature of the atmosphere varies with altitude.
• The atmosphere is divided into four regions based on the temperature profile.
• The troposphere (below an altitude of 12 km)
• The temperature decreases from 290 K to 215 K as altitude increases.
• This region is where we spend most of our time.
Composition of the Atmosphere
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• The composition of the atmosphere is not uniform.
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“Introducing Atmospheric Reactions: A Systematic Approach for Students” from Further Readings
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Figure 18.1 from Transparency Pack
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“Thermal Physics (and Some Chemistry) of the Atmosphere” from Further Readings
4
Chapter 18
264
• Temperature and pressure vary over a wide range with altitude.
• Gases in the atmosphere are bombarded by radiation and energetic particles from the sun.
• Gravity also plays a role.
• Lighter molecules and atoms are found at higher altitudes.
Photochemical Reactions in the Atmosphere
• Recall: E = h = hc/
• Thus, the higher the frequency, the shorter the wavelength and the higher the energy of radiation.
• For a chemical reaction induced by radiation to occur, the photons must have sufficient energy to
break the required bonds, and the molecules must absorb the photons.
• Photoionization is the ionization of molecules (and atoms) caused by radiation.
• The molecule absorbs energy, causing the loss of an electron.
• Thus, the photon must have sufficient energy to remove an electron when it is absorbed by a
molecule.
• Wavelengths of light that cause photoionization and photodissociation are absorbed by the upper
atmosphere.
• This filters them out and prevents them from reaching the Earth.
Ozone in the Upper Atmosphere
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• Ozone absorbs photons with wavelengths between 240 and 310 nm.
Chemistry of the Environment
265
• The oxygen atoms can collide with oxygen molecules to form ozone with excess energy, O3*:
O(g) + O2(g) → O3*(g)
• The excited ozone (O3*) can lose energy by decomposing to oxygen atoms and oxygen molecules
(the reverse reaction) or by transferring the energy to M (usually N2 or O2):
O(g) + O2(g) → O3*(g)
O3*(g) + M(g) → O3(g) + M*(g)
• Why does maximum ozone formation occur in the stratosphere?
18.2 Human Activities and Earth’s Atmosphere
The Ozone Layer and Its Depletion
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• 1970: Cratzen demonstrated that naturally occurring nitrogen oxides can catalytically degrade ozone.
• 1974: Rowland and Molina demonstrated that chlorine atoms from chlorofluorocarbons (CFCs)
deplete the ozone layer.
• Rowland and Molina were awarded the Nobel prize in 1995.
• Hydrofluorocarbons are the main alternatives to hydrofluorocarbons.
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“Understanding Ozone” from Further Readings
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Figure 18.6 from Transparency Pack
Chapter 18
266
Sulfur Compounds and Acid Rain
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• Sulfur dioxide, SO2, is produced by natural events (volcanic gases, bacterial action, forest fires).
• The major source is linked to human activities such as the combustion of sulfur-containing fuels.
• Combustion of coal accounts for approximately 65% of the SO2 released in the United States.
• The amount of SO2 produced depends on the sulfur content of the coal or oil, which varies with
the source of the coal or oil.
• The SO2 can be oxidized to SO3, which dissolves in water to produce sulfuric acid (a component
of acid rain): SO3(g) + H2O(l) → H2SO4(aq)
• The acids in acid rain are problematic.
• They react with metals and promote corrosion.
• They react with carbonates (such as the calcium carbonate in marble and limestone).
• How can we reduce the amount of SO2 produced from fuel combustion?
• It is too expensive to remove sulfur from oil and coal prior to its use.
Nitrogen Oxides and Photochemical Smog
• Photochemical smog is the result of photochemical reactions on pollutants.
• Oxides of nitrogen are the primary components of smog.
• In car engines, NO forms as follows:
2NO(g) + O2(g) → 2NO2(g) ∆H = –113.1 kJ
• In air NO is rapidly oxidized: NO2(g) + h → NO(g) + O(g)
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Figure 18.9 from Transparency Pack
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“Acid Rain Effects on Stone Monuments” from Further Readings
Chemistry of the Environment
267
Greenhouse Gases: Water Vapor, Carbon Dioxide, and Climate
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• There is a thermal balance between the Earth and its surroundings.
• Therefore, radiation is emitted from the Earth at the same rate as it is absorbed by the Earth.
• The troposphere is transparent to visible light.
• However, the troposphere is not transparent to IR radiation (heat).
• Therefore, the troposphere insulates the Earth, making it appear colder from the outside than it is on
the surface.
fossil fuels in the present manner.)
FORWARD REFERENCES
• Other methods of energy production, such as direct methanol fuel cells, will be discussed in
Chapter 20 (section 20.7).
18.3 The Earth’s Water
• 72% of the Earth’s surface is covered with water.
• Water plays an important role in our environment.
• The properties of water are important.
The Global Water Cycle
• All water on Earth is connected in a global water cycle.
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Figure 18.11 from Transparency Pack
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Figure 18.12 from Transparency Pack
Chapter 18
268
• Water is warmed by the sun.
• Liquid water in the oceans evaporates into the atmosphere as water vapor.
• The water vapor condenses into water droplets.
• Water falls to the ground as water, snow, or rain.
Saltwater: Earth’s Oceans and Seas
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• The volume of oceans in the world is 1.35 109 km3.
Freshwater and Groundwater
• Freshwater : natural waters that have low concentrations of dissolved salts and solids.
• Includes waters of lakes, rivers, ponds, and streams.
• An adult needs about 2 L of water a day for drinking.
• In the United States, the average person uses about 300 L of freshwater per day.
• Industry uses even more freshwater than this (e.g., about 105 L of water is used to make 1000 kg of
steel).
18.4 Human Activities and Earth’s Water
Dissolved Oxygen and Water Quality
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• Water fully saturated with air at 1 atm and 20°C has 9 ppm of O2 dissolved in it.
• Cold-water fish require about 5 ppm of dissolved oxygen for life.
• Aerobic bacteria consume oxygen to oxidize biodegradable organic material.
Chemistry of the Environment
269
Water Purification: Desalination
• Seawater has a salt concentration too high for drinking.
• Water used for drinking should contain less than 500 ppm dissolved salts (United States
municipal water).
• Desalination is the removal of salts from seawater or brackish water.
• Common methods for isolation of drinking water include distillation (small scale) and reverse
Water Purification: Municipal Treatment
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• There are five steps:
• Coarse filtration
• This occurs as water is taken from a lake, river, or reservoir and passed through a screen.
• Sedimentation
• Water is allowed to stand so that solid particles (e.g., sand) can settle out.
• To remove small components (like bacteria), CaO and Al2(SO4)3 are added.
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Water Softening
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• Water containing a high concentration of Ca2+ and Mg2+ and other divalent cations is called hard
water.
• The presence of these ions may cause the water to be unsuitable for some uses.
• For example, soaps form an insoluble soap "scum" and water in water heaters forms a deposit
FORWARD REFERENCES
• The role of water in redox reactions will be discussed throughout Chapter 20.
• Teflon will be mentioned in Chapter 22 (section 22.4).
• Lead chemistry, lead poisoning, and chelating agents will be mentioned in Chapter 23.
18.5 Green Chemistry
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• 12 basic principles of green chemistry: an initiative that promotes the design and application of
chemical products and processes that are compatible with human health and that preserve the
environment:
• 1. Prevention
• 2. Atom Economy
• 3. Less Hazardous Chemical Syntheses
• 4. Designing Safer Chemicals
• 5. Safer Solvents and Auxiliaries
Demonstrations
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271
Supercritical Solvents
• The release of toxic volatile solvents to the atmosphere must be avoided.
• Alternative environmentally friendly methods need to be developed.
Greener Reagents and Processes
• Earth-friendly alternatives are being developed for many processes important to modern society.
• Where possible, synthetic routes are adjusted to result in a high level of ‘atom economy”:
• A high percentage of the atoms from the starting materials end up in the product.
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Further Readings:
1. N. Colin Baird, “Introducing Atmospheric Reactions: A Systematic Approach for Students,” J. Chem.
Educ., Vol. 72, 1995, 153–157.
2. Stephen K. Lower, “Thermal Physics (and Some Chemistry) of the Atmosphere,” J. Chem. Educ., Vol.
75, 1998, 837–840.
7. F. Sherwood Rowland and Mario J. Molina, “Ozone Depletion: 20 Years after the Alarm,” Chemical
and Engineering News, August 15, 1994, 8–15.
8. Lynn Diener, “News from Online: Stratospheric Chemistry,” J. Chem. Educ., Vol. 86, 2009, 153–155.
9. Muhammad Hanif, “Understanding Ozone,” The Science Teacher, December 1995, 20–23.
14. Tina Adler, “The Expiration of Respiration: Oxygen—The Missing Ingredient in Many Bodies of
Water,” Science News, February 10, 1996, 88–89.
15. Barbara A. Balko and Paul G Tratnyek, “A Discovery-Based Experiment Illustrating How Iron Metal
Is Used to Remediate Contaminated Groundwater,” J. Chem. Educ., Vol. 78, 2001, 1661–1663.
16. Richard S. Treptow, “Carbon Footprint Calculations: An Application of Chemical Principles,” J.
Chem. Educ., Vol. 87, 2010, 168–171.
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17. Maria T. Oliver-Hoyo and Gabriel Pinto, “Using the Relationship between Vehicle Fuel
Consumption and CO2 Emissions to Illustrate Chemical Principles,” J. Chem. Educ., Vol. 85, 2008, 218–
220.
21. Stephan K. Ritter, “Green Chemistry Gets Greener,” Chemical and Engineering News, May 20,
2002, 38–42.
22. Bette Hileman, “Climate Change,” Chemical and Engineering News, December 15, 2003, 27–37.
Live Demonstrations:
1. Bassam Z. Shakhashiri, “Acid-Neutralizing Capacity of Lake Beds,” Chemical Demonstrations: A
Handbook for Teachers of Chemistry, Volume 3 (Madison: The University of Wisconsin Press, 1989), pp.
125–127.
