1	Chapter 18 Chemistry of the Environment
2	The Environment There are a number of general problems in the environment thought to be caused by human activities Focuses basically on the earth’s atmosphere and the aquatic environment called the hydrosphere.
3	Atmosphere Pressure decreases regularly with increasing altitude. Temperature fluctuates with altitude: 0-~12 km: Temperature decreases with altitude (298 K – 215K) ~12 km-50 km: Temperature increases with altitude (215 K-275 K) 50 km – 85 km: Temperature decreases with altitude (275 K –180 K) Above 85 km: Temperature increases with altitude
4
5	Atmosphere Atmosphere consists of layers like an onion. Troposphere (0-12 km) Temperature decreases with altitude. Most weather occurs in this layer. Contains 75% of the mass of the atmosphere Airplanes fly near the top of the troposphere called the tropopause. Gases mix across very slowly across the boundaries between layers.
6	Atmosphere Stratosphere (12 km – 50 km) Temperature increases with altitude. Together, the troposphere and the stratosphere account for 99.9% of the mass of the atmosphere. Upper boundary called the stratopause. Mesosphere (50 km – 85 km) Temperature decreases with altitude. Upper boundary called the mesopause. Thermosphere (Above 85 km) Temperature increases with altitude.
7	Composition of the atmosphere Mostly N₂ and O₂ (99%) with CO₂(g) and noble gases comprising most of the rest of the atmosphere.
8	Chemistry of the atmosphere N₂(g) Contains a triple bond between the two N atoms. (bond energy 941 kJ/mol) Makes N₂(g) an extremely unreactive molecule. O₂(g)’ Contains a double bond between the two O atoms. (bond energy 495 kJ/mol) Oxygen is much more reactive than N₂.
9	Outer regions of the atmosphere The outer atmosphere is bombarded with large amounts of cosmic radiation. Very energetic subatomic particles and high energy electromagnetic radiation. These particles are absorbed by molecules which cause photodissociation and photoionization reactions.
10
11	Photodissociation reactions A gas molecule absorbs a photon of light which breaks the bond. O₂(g) + hn ¾® 2 O(g) NO(g) + hn ¾® N(g) + O(g) Absorbs much of the very high energy cosmic radiation. At 400 km: 99% O atoms, 1% O₂ molecule At 130 km: 50% O atoms, 50% O₂ molecules At sea level: 99.9% O₂ molecules. N₂(g) does not undergo photodissociation
12	Photoionization reactions Photons of light may be absorbed by gas molecules which then eject electrons producing ions.
13	Photoionization reactions Absorb more of the higher energy cosmic radiation. (wavelengths smaller than 240 nm) These processes occur mostly above 90 km (mesosphere and thermosphere). Radiowaves bounce off the layer of ions produced by these reactions.
14
15	Ozone layer Ozone is created by the reaction of O₂ with O: O₂(g) + O(g) ¾® O₃(g) Ozone absorbs solar radiation which decomposes it back to O₂(g) and O(g). Absorbs light of wavelengths 200-310 nm. Ozone concentration reaches a maximum around an altitude of 50 km.
16
17	Ozone layer Protects us from harmful UV radiation. Ozone depletion A class of compounds called chlorofluorocarbons (CFC’s) decomposed when exposed to light in the stratosphere to produce Cl atoms. Cl atoms catalyze the destruction of ozone. Nitrogen oxides also catalyze the destruction of ozone. Both reactions were hypothesized in 1974 and 1970.
18	Ozone Layer Satellites were placed in orbit that could measure the amount of ozone in the ozone layer in the 1980’s and 1990’s Severe depletion of ozone in the stratosphere has been measured over the South pole and a less severe depletion over the North pole. Montreal Protocol: use of CFC’s phased out. Replacements: hydrofluorocarbons HFC’s
19	Pollution in the Troposphere CO₂(g) 375 ppm CO(g) 0.05 ppm in clean air, 1-50 ppm in urban environments. CH₄(g) 1.77 ppm NO(g) 0.01 ppm in clean air, 0.2 ppm in smog O₃(g) 0-0.01 ppm in clean air, 0.5 ppm in photochemical smog. SO₂(g) 0-0.01 ppm in clean air, 0.1-2 ppm in urban environments.
20	Acid Rain Sulfur-containing compounds present naturally in the atmosphere. Organic decay Volcanic activity Burning fossil fuels, especially coal, releases SO₂(g) into the atmosphere. SO₂(g) very toxic.
21	Acid Rain SO₂(g) undergoes a series of reactions in the atmosphere 2 SO₂(g) + O₂(g) ¾® 2 SO₃(g) SO₃(g) + H₂O(l) ¾® H₂SO₄(aq) Contributes to a phenomenon known as “acid rain”
22	Acid Rain Rainwater is naturally slightly acidic due to the presence of CO₂(g): pH around 5.6 Acid rain: pH around 4 or lower. Kills most aquatic life. Dissolves statues and stone monuments. Besides SO₂(g), nitrogen oxides (NO(g)) also contributes to acid rain.
23
24
25
26	Carbon Monoxide CO(g) Formed by the incomplete combustion of carbon-containing compounds (fossil fuels) Most abundant of pollutant gases. It is very toxic to humans. It binds to the active site in hemoglobin making it incapable of carrying oxygen.
27
28	Carbon monoxide Produced whenever a fossil fuel is burned. Catalytic converter in your car catalyzes the reaction of CO(g) with O₂(g) to produce CO₂(g) If you burn wood in your fireplace or use a kerosene heater, make sure a window is cracked open to prevent carbon monoxide poisoning.
29
30	Photochemical smog Starts with the production of NO(g) N₂(g) + O₂(g) ¾® 2 NO(g) Reaction occurs in internal combustion engines and industrial combustion processes. NO is a toxic gas.
31	Photochemical smog NO reacts further in the atmosphere 2 NO(g) + O₂(g) ¾® 2 NO₂(g) NO₂(g) + hn ¾® NO(g) + O(g) O₂(g) + O(g) ¾® O₃(g) Called photochemical smog. Air turns a yellow-brown color (NO₂) Ozone is extremely toxic.
32
33	Global warming During the day, the earth’s surface is heated by the sun. During the night, the earth’s surface loses this heat mostly by radiational cooling. It emits infrared radiation. Both H₂O and CO₂ absorb a significant part of this radiation, thereby trapping the heat near the surface.
34
35
36
37	Global warming Since the industrial revolution, humans have drastically increased the amount of CO₂(g) in the atmosphere. Over the past century, global temperatures have increased from 0.3-0.6°C It has been noticed that the polar ice caps are melting at a much greater rate. This raises sea level.
38
39	Global warming Also referred to as the “greenhouse effect.” Methane, CH₄(g), is also thought to contribute to global warming. Kyoto protocol: reduce the amount of CO₂ released to the atmosphere.
40	Hydrosphere Water 97.2% contained in the world’s oceans 2.1% ice caps and glaciers 0.6% fresh water 0.1% brackish (slightly salty) water
41	Ocean water Saltwater Salinity: 35 grams of dry salt per kg of seawater 3.5% dissolved salts
42
43
44	Ocean water It is generally too expensive to use seawater as a source of raw materials. Compounds extracted from seawater: Sodium chloride Bromine (from bromide ions) Magnesium Oceans play an important role in absorbing CO₂(g) from the atmosphere.
45	Saltwater Too salty for human consumption Osmotic pressure too high. Bodily fluids of saltwater fish are isotonic with seawater so they can drink it. Desalination: removing salt from seawater Distillation Reverse osmosis
46
47	Fresh water 0.6% of total water on Earth. 9 ´ 10¹¹ L per day used in U.S. 41% agriculture 39% hydroelectric power 6% industry 6% household needs 1% drinking
48	Fresh water Average person uses about 300 L per day of fresh water 8 L cooking and drinking 120 L for cleaning 80 L for toilets 80 L for watering lawns
49	Fresh water As water runs off the land, it dissolves a variety of cations: Na⁺, K⁺, Ca²⁺, Fe²⁺, and anions: Cl¯, SO₄^2-, HCO₃¯. Quality of water supplies measured by the amount of dissolved O₂. 9 ppm saturated O₂ at 20°C Fish need at least 5 ppm
50	Water quality Add biodegradable material (animal wastes) These materials are food for bacteria in the water which decompose it. When bacteria are well fed, they multiply. As they multiply, they consume more of the dissolved O₂ in the water. Fish and other aquatic life die off. This is more food for the bacteria. Bacteria multiply more and the amount of dissolved O₂ decreases more. These bacteria produce CO₂, HCO₃¯, H₂O, NO₃¯, SO₄^2-, and PO₄^3-.
51	Water quality Eventually, the amount of dissolved O₂ gets too low for bacteria to live. They die off. This is food for bacteria which live in the absence of O₂. Called anaerobic bacteria. The anaerobic bacteria produce hydrogen containing compounds: H₂S, NH₃, CH₄, PH₃ These gases produce generally offensive odors.
52	Water quality Fertilizer runoff into water Causes excess plant growth. The lack of sunlight on the bottom kills plants on the bottom which feeds bacteria. Amount of dead material leads to the same result as animal waste. Eutrophication.
53	Treatment of municipal water Filter the water from the source. Sedimentation: let the water stand in tanks to allow solid particles to settle to bottom. Pass it through a finer filter. Aerate the water Sterilize the water with either Cl₂(g) or O₃(g)
54
55	Equilibrium Problem The value of K_eq is 0.15 at 25°C for the reaction:
56	Strong Acid Calculate the pH of 0.040 M HClO₄.
57	Weak Acid Calculate the pH of 0.025 M HC₇H₅O₂ (benzoic acid). The K_a for benzoic acid is equal to 6.3 ´ 10^-5.
58	Strong Base Calculate the pH of 0.073 M Ba(OH)₂.
59	Weak Base Calculate the pH of 0.15 M CH₃NH₂. The K_b for CH₃NH₂ is equal to 5.0 ´ 10^-4.
60	Salt Solution Calculate the pH of 0.40 M KClO. The K_a for HClO is equal to 3.5 ´ 10^-8.
61	Common Ion Effect What is the pH of a solution that is 1.00 M in HC₂H₃O₂ and 0.75 M in NaC₂H₃O₂? The K_a for HC₂H₃O₂ is equal to 1.8 ´ 10^-5.
62	Solubility What is the solubility, in grams per liter, of calcium fluoride? The K_sp for CaF₂ is equal to 5.3 ´ 10^-11.
63	Common Ion-Solubility What is the solubility, in grams per liter, of calcium carbonate in 0.10 M CaCl₂ solution. The K_sp of CaCO₃ is equal to 3.7 ´ 10^-9.