Chapter 18 Chemistry of the Environment

1. Chapter 18Chemistry of the Environment Chapter 18

2. Introduction Human activities impact our environment. Economic growth depends of chemical processes • e.g. clean water, energy usage, chemical synthesis, etc. Earth Summit – Brazil 1992 1997 – Kyoto meeting 2001 – Bonn – Signing of “Kyoto Protocols” Protocols designed to internationally address environmental concerns and establish regulations. 2007 – UN meeting to establish further regulations and discuss monetary help for developing countries. Chapter 18

3. Topics Earth’s Atmosphere • Ozone • Acid rain • Greenhouse effect The Ocean Freshwater Soils and Weathering: A little basic geochemistry Green Chemistry Find the topic interesting? Some easy reading… 1. “An Introduction to Environmental Chemistry” by JE Andrews, P Brimblecombe, TD Jickells and PS Liss. 2. “A Short History of Nearly Everything” by Bill Bryson. 3. Movie: An Inconvenient Truth Chapter 18

4. Earth’s Atmosphere Four regions. Remember region and boundary names Notice relationship between altitude and T and P. (Colder in JNB than KNP AND why jets are pressurized.) Jets fly at this level. Chapter 18

5. General Interest Thin clouds made of mixtures of ice, HNO3 and H2SO4 form in the upper atmosphere (stratosphere) over the poles when temperatures drop below -78°C (-109°F). Ozone depletion occurs in such polar stratospheric clouds. Aurora Borealis: collisions of charged particles above 80 km with O, O2, N, and N2. Clouds from Earth’s surface to the stratosphere. Chapter 18

6. Earth’s Atmosphere Earth’s atmosphere is affected by temperature and pressure and gravity. Lighter molecules and atoms are found at higher altitudes. Less O2, air is “thinner”. There is slow mixing of gases between regions in the atmosphere (important for ozone and pollution). Troposphere and stratosphere account for 99.9% of the atmosphere’s mass, with about 75% of that in the troposphere. Two major components of the atmosphere are nitrogen, N2, and oxygen, O2. Chapter 18

7. A Note on N2 and O2 Reactivity N2 has a triple bond (bond energy of 941kJ/mol) O2 double bond (bond energy of 495kJ/mol) Oxygen is more reactive than nitrogen. Generally, oxides of non-metals form acidic solutions with water and oxides of metals form basic solutions in water. Chapter 18

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9. Parts Per Million 1 part per million (ppm) refers to 1 part in 1 million units of the whole. (One red pencil in a million yellow pencils) If considering PV = nRT, volume fraction which is usually used for ppm can be interchanged with mole fraction. So 1ppm = 1 mole in 1 million moles of total gas. concentration in ppm = mole fraction x 106 Other common ppm designations: solids in mg kg-1, liquids in mg L-1, and most engineers use mg dm-3 (remember that 1L = 1dm3). Chapter 18

10. Example Calculate the concentration of methane in ppm if mole fraction = 0.000002. ppm = mole fraction x 106 = 0.000002 x 106 = 2 ppm Calculate the concentration of xenon in ppb (parts per billion) if the mole fraction = 0.000000087. ppm = mole fraction x 106 ppb = mole fraction x 109 = 0.000000087 x 109 = 87 ppb Chapter 18

11. Outer Regions of the Atmosphere Contains only a small percentage of the atmospheric mass. Low pressure! Forms the outer defense against radiation and high-energy particles from space. (If Earth was the size of a desk globe, our atmosphere’s thickness would be ~ two layers of varnish. Very thin, but important, layer!) Outer atmospheric reactions: photodissociation and photoionization Chapter 18

12. Solar Spectrum Chapter 18

13. Photodissociation • Wave equation: The higher the frequency, the shorter the wavelength and the higher the energy of radiation. For photodissociation reactions to occur, photons must have sufficient energy to break the required bonds and the molecules must also absorb these photons. Defn: the rupture of a chemical bond induced by radiation. Chapter 18

14. Photodissociation In the upper atmosphere (above 120km), photodissociation causes the formation of oxygen atoms: O2(g) + h 2O(g) Minimum energy required determined by the bond dissociation energy of O2 (495kJ/mol). Dissociation of O2 is very extensive at high elevations: at 400km only 1% of oxygen is O2 while at 130km, about 50% is O2. Chapter 18

15. Photoionization Defn: ionization of molecules (and atoms) caused by radiation. 1924: electrons discovered in the upper atmosphere. Therefore, cations must be present in the upper atmosphere (for charge balance). Photoionization occurs when a molecule absorbs a photon of sufficient energy to remove an electron. Wavelengths of light that cause photoionization and photodissociation are filtered by the atmosphere (occur more readily further from Earth). Chapter 18

16. Ozone and the Upper Atmosphere Ozone (O3) absorbs photons with a wavelength between 240 and 310 nm. (N2, O2 and O absorb wavelengths shorter than 240nm.) Most of the ozone is present in the stratosphere, 12-50 km altitude. Maximum ozone concentration at an altitude of ~20 km. Between 30-90km photodissociation of oxygen is possible: O2(g) + h 2O(g) Chapter 18

17. Ozone and the Upper Atmosphere Oxygen atoms can collide with oxygen molecules to form ozone with excess energy, O3*: O(g) + O2(g)  O3*(g) (releases 105kJ/mol) The excited ozone loses energy by decomposing to oxygen atoms and oxygen molecules (the reverse reaction) or by transferring the energy to M (usually N2 or O2): O3*(g) + M(g)  O3(g) + M*(g) Chapter 18

18. Ozone and the Stratosphere The formation of ozone in the atmosphere depends on the presence of O(g): At low altitudes, most radiation with sufficient energy to form O(g) has been absorbed by upper atmosphere. Release of energy from O3* depends on collisions, which generally occur at lower altitudes (more gas molecules). Combining altitude and O(g) concentration means maximum ozone formation in the stratosphere. Chapter 18

19. Ozone’s Natural Cycle A combination of photodissociation and photodecomposition reactions – makes and breaks ozone naturally. O2(g) + hv O(g) + O(g) O(g) + O2(g)  O3(g) + M*(g) (heat released) O3(g) + hv O2(g) + O(g) O(g) + O(g) O2(g) (heat released) O3 also formed by lightning strikes (can split O2) – can detect metallic smell of ozone during thunderstorms. Chapter 18

20. Ozone Depletion: CFCs Crutzen 1970: naturally occurring nitrogen oxides catalytically destroy ozone. In 1974, Rowland and Molina showed that chlorine from chlorofluorocarbons (CFCs) deplete the ozone layer by catalyzing the formation of ClO and O2. CFC Characteristics CFCl3 (Freon-11TM) and CF2Cl2 (Freon-12TM) used as propellants in spray cans, as refrigerant gases, and foaming agents for plastics. Unreactive in lower atmosphere, insoluble in water. Diffuse slowly into stratosphere. Several million tons now present in the atmosphere. Chapter 18

21. CFC Depletion of Ozone • In the stratosphere, CFCs undergo photodissociation of C-Cl: • CF2Cl2(g) + h CF2Cl(g) + Cl(g) (optimal at 30km) • Subsequently: Cl(g) + O3(g)  ClO(g) + O2(g) • rate = k[Cl][O3], k = 7.2  109M-1s-1 at 298K • In addition, the ClO generated produces Cl as well: • 2ClO(g)  O2(g) + 2Cl(g) • The overall reaction: 2O3(g)  3O2(g). Think about this… The rate at which ozone is destroyed increases with the amount of Cl, thus the greater the amount of CFCs that diffuse into the stratosphere, the faster the destruction of the ozone layer. Chapter 18

22. The Ozone Hole and CFC Replacements 1992: ~100 nations agreed to ban production and use of CFCs by 1996. October 1994: Ozone map in the Southern hemisphere showed a hole over Antarctica. In 1995 the Nobel Prize for chemistry was awarded to F. Sherwood Rowland, Mario Molina, and Paul Crutzen for their studies of ozone depletion. Replacement for CFCs are HFCs (e.g. CH2FCF3), where a C-H bond is broken instead of a C-Cl. Chapter 18

23. Antarctic O3 Hole: How it forms Circulation of Earth’s atmosphere = “polar vortex” during Antarctic winter. Extreme cold (-80˚C) dissolves more gas. Formation of polar stratospheric clouds (PSCs) (ice and acid). Dark for 3 months – allows reactants to build up. Sept-Nov: sun returns and photochemical reactions go crazy! Estimated >50% of O3 in polar vortex destroyed. PSCs break apart, O3-depleted stratosphere disperses and O3 can start rebuilding naturally. Only an ozone “dimple” above Arctic pole (WHY?). Chapter 18

24. Down to Earth: Chemistry of the Troposphere The troposphere consists mostly of O2 and N2 (~99%), but also CO, CO2, H2O, CH4, NOx, and SOx. Even low to trace concentrations of other gases can have profound effects on the environment. Most of these environmentally unfriendly gases are produced by the combustion of fossil fuels, industry and everyday life. Chapter 18

25. Sulfur Compounds and Acid Rain • Sulfur dioxide, SO2, is largely produced by the combustion of oil and coal and isa serious health hazard especially to people with respiratory difficulties. • SO2 is oxidized to SO3 by reacting with O2 or O3 which then reacts with water to produce sulfuric acid (acid rain): • SO3(g) + H2O(l)  H2SO4(aq) • Note: nitrogen oxides also contribute to acid rain (nitric acid). • Normal rainwater has a pH of about 5.6 (due to dissolved H2CO3 from CO2 in the air). Chapter 18

26. Sulfur Compounds and Acid Rain Acid rain has a pH around 4 (pH of natural waters containing living organisms is 6.5 to 8.5). Natural waters with pH below 4 cannot sustain life. Acid rain: corrosive to metals and stone building materials (will slowly dissolve limestone). Too expensive to remove sulfur from oil and coal prior to its use. Therefore, the SO2 is removed from fuel upon combustion. More than 30MT per year of SO2 are released into the atmosphere in the USA. Even more in developing countries. Chapter 18

27. Sulfur Removal from Coal Combustion SO2 is commonly removed from fuel (oil and coal) as follows: 2. CaSO3 and unreacted SO2 are passed into a scrubber (purification chamber) where the excess SO2 is converted to CaSO3 by jets of CaO. • 1. Powdered limestone decomposes into CaO which reacts with SO2 to form CaSO3 in a furnace. 3. CaSO3 is precipitated into a watery slurry. Chapter 18

28. Carbon Monoxide Produced by incomplete combustion of carbon-containing materials, e.g. fossil fuels. About 1014 g of CO is produced in the United States per year (mostly from automobile exhaust). CO binds irreversibly to the Fe in hemoglobin (about 210 times more strongly than oxygen). Extremely dangerous and potentially fatal if inhaled in large quantities. Chapter 18

29. Nitrogen Oxides Photochemical smog (“the brown cloud”) is the result of photochemical reactions on pollutants. In car engines, NO forms as follows: In air, rapid oxidation of NO takes place: At 393nm (sunlight), NO2 decomposes NO2(g) + h NO(g) + O(g) Chapter 18

30. Nitrogen Oxides The O produced by photodissociation of NO2 can react with O2 to form O3, which is the key component of smog. O(g) + O2(g)  O3(g) + M*(g) Ozone is undesirable in the troposphere because it’s toxic and reactive. Ozone problem…too much of it in smog, not enough in the stratosphere. Chapter 18

31. Water Vapor, CO2, and Climate Thermal balance between Earth and its surroundings. Ideally, radiation is emitted from Earth at the same rate as it is absorbed. The troposphere is transparent to visible light (comes through), but not to IR radiation (heat). CO2 and H2O absorb IR radiation escaping from Earth’s surface at night keeping us warm. Called the Greenhouse Effect. Chapter 18

32. Water Vapor, CO2, and IR Radiation Chapter 18

33. Historical CO2 Profile The carbon dioxide level on Earth has been increasing over the years. Current CO2 levels = 380 ppm (a 35% increase from pre-industrial levels). www.physorg.com Majority of the increase is from burning of fossil fuels. Chapter 18

34. No Quick Solution for CO2 Between 2050 and 2100, the CO2 concentration is expected to double, possibly resulting in a global temperature increase of 1 to 3˚C. International Energy Agency reports China will surpass US in CO2 emissions by 2009 (heavy reliance on coal). NY Times, 7 Nov 2006 CO2 levels higher now than in the past 650,000 years based on comparision to tiny bubbles trapped in Antarctic ice cores. 25 Nov 2005, Science Could result in melting of glaciers and a subsequent sea level rise. Chapter 18

35. Increased CO2 Implications Glacial recession in Glacier National Park, Montana (USA). Wikipedia Adapted from Gormitz and Lebedeff, 1967 by UNFCCC. Submerged Florida peninsula (USA) shown with 1 meter sea level rise (in red). NASA Antarctic meltoff, 2005. NASA Chapter 18

36. Methane: CH4 Of the three main greenhouse gases…receives least attention. About 25x stronger than CO2 in greenhouse effect contribution! Industrial age: increased from 0.3ppm to 1.8ppm. Reactions: Stratosphere: oxidized – water byproduct. Troposphere: attacked – ozone byproduct. Where it comes from: anaerobic bacterial processes, ruminant animals, volcanic activity. Australia – 14% country’s total greenhouse gas emissions from sheep and cows!! Chapter 18

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38. Signs of global warming • A rise in average global surface air temperature • Winters have become shorter by about 11 days • The earth’s ice cover is shrinking fast. Glaciers, polar icecaps, and polar sea ice are melting and disappearing at unprecedented rates due to global warming. • Warm water is killing much of the coral in ocean reefs and threatening sea life. • The mosquito-borne diseases have reached higher altitudes. Because of warmer temperatures, mosquitoes are now able to survive in regions where they formerly were not viable. • Rising sea levels are threatening to engulf Pacific islands. • Extreme weather is becoming more common. Chapter 18

39. The World Ocean 72% of Earth’s surface is covered with water. 97.2% of Earth’s water is seawater, with a volume of 1.35  109 km3. Only 0.6% of Earth’s total water is in rivers, lakes, and groundwater! Salinity: mass in grams of dry salts in 1 kg of seawater. Seawater salinity averages about 35. Most elements in seawater are only present in small quantities (trace). Commercially, NaCl, Br- and Mg2+ are obtained from seawater. Chapter 18

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41. Desalination • Water used for drinking should contain less than 500 ppm dissolved salts (United States water regulation). Joburg water is at ~350 ppm. • Desalination: removal of salts from seawater. • Common method: reverse osmosis (energy intensive). • Osmosis: transport across a semipermeable membrane, where solvent moves from dilute to concentrated. • No good for desalination. • Reverse osmosis: under applied pressure, solvent moves from more concentrated solution to more dilute solution. Chapter 18

42. Desalination Seawater is introduced under pressure and water passes through the fiber walls and is separated from the ions. • Largest desalination plant in Saudi Arabia responsible for 50% of the country’s drinking water by reverse osmosis from Persian Gulf. Small scale reverse osmosis desalinators used for camping, traveling, and at sea. Chapter 18

43. Freshwater An adult needs about 2 L a day for drinking. The average person uses about 300 L of freshwater per day. Industry uses about 105 L of water are used to make enough steel for one car! As water flows over the terrain it dissolves many substances. Freshwater usually contains some ions (Na+, K+, Mg2+, Ca2+, Fe2+, Cl-, SO42-, and HCO3-) and dissolved gases (O2, N2, and CO2). Chapter 18

44. Dissolved O2 and Water Quality 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. Chapter 18

45. Bacterial Activity in Groundwater Aerobic bacteria require oxygen to biodegrade organic material, e.g., sewage, industrial waste from food-processing plants and paper mills, and effluent from meat packing plants. Oxidize organic material into CO2, HCO3-, H2, NO3-, SO42-, and phosphates. Once the oxygen level has been depleted, aerobic bacteria cannot survive. Anaerobic bacteria complete the decomposition process forming CH4, NH3, H2S, PH3, and other smelly products. Chapter 18

46. Municipal Water Treatment There are five steps 5.Sterilization. Chlorine used, forms HClO(aq) in solution which kills bacteria. 1.Coarse filtration. Occurs as water is taken up from lake, river or reservoir. 4.Aeration. Air oxidizes any organic material. • 2.Sedimentation. Water is allowed to stand so that solid particles (e.g. sand) can settle out. Gelatinous precipitate of Al(OH)3 settles out slowly. 3.Sand Filtration. Filteration through a sand bed to remove Al(OH)3 and anything it trapped. Chapter 18

47. Green Water Purification When Cl2(g) is used to treat water, some trihalomethanes (THM) are often produced, which can go undetected. THMs (CHCl3, CHClBr2) are suspected carcinogens. Ozone and ClO2 could be used as alternatives, but they are not completely safe. Green water purification is an open problem. Chapter 18

48. Hard Water Hard water contains relatively high amounts of Ca2+, Mg2+ and other divalent cations. Unsuitable for most household and industrial uses. -Forms insoluble soap scum and water “stains” Mineral deposits form when hard water is heated Ca2+(aq) + 2HCO3-(aq) CaCO3(s) + CO2(g) + H2O(l) “scale” Chapter 18

49. Water Softening Water from underground sources with considerable contact with CaCO3 and other minerals containing Ca2+, Mg2+ and Fe3+ requires softening. Lime-soda process used for large scale municipal water-softening operations. Water treated with lime, CaO (or slaked lime, Ca(OH)2), and soda ash, Na2CO3. Ca2+(aq) + CO32-(aq) CaCO3(s) Mg2+(aq) + 2OH-(aq) Mg(OH)2(s) precipitates Chapter 18

50. Water Softening Ion exchange used for household water softening. Hard water is passed through a bed of ion exchange resin: plastic beads with covalently bound anion groups (–COO- or –SO3-). These anion groups have Na+ attached to counter their charges. The Ca2+ and other cations in the hard water are exchanged with the Na+. 2Na+(R-COO-)(s) + Ca2+(aq) Ca2+(R-COO-)2(s) + 2Na+(aq) Chapter 18