Presentation Slides for Air Pollution and Global Warming: History, Science, and Solutions Chapter 10: Acid Deposition

By Mark Z. Jacobson Cambridge University Press (2012)

Last update: February 23, 2012 The photographs shown here either appear in the textbook or were obtained from the internet and are provided to facilitate their display during course instruction. Permissions for publication of photographs must be requested from individual copyright holders. The source of each photograph is given below the figure and/or in the textbook.

History of Acid Deposition

1200s.

Coal used in lime kilns and forges produces sulfur and nitrogen oxides that result in sulfuric and nitric acid deposition.

1750s.

Expanded use of the steam engine at the beginning of the Industrial Revolution increases sulfuric and nitric acid deposition.

1783.

the most efficient and economical method to produce soda ash Na French Academy of Sciences offers a prize to a person who finds 2 CO 3 (aq), which along with animal fat, is used to make soap.

1789.

Nicolas LeBlanc develops two-step soda ash process.

The Rise of Alkali Factories

Soda ash + Animal fat = Soap Img.alibaba.com

Keetsa.com

www.nutritionfx.com

LeBlanc Process

High temp erature 2NaCl(s) + H 2 SO 4 (aq) Sodium chloride (salt) Sulfuric acid Na 2 SO 4 (aq) + 2HCl(g) Sodium sulfate Hy drochloric acid High temp erature Na 2 SO 4 (aq) + 2C + CaCO 3 (s) Na 2 CO 3 (aq) + CaS(s) + 2CO 2 (g) Sodium sulfate Carbon from Calcium carbonate Sodium carbonate Calcium sulfide Carbon dioxide charcoal from chalk (10.1) - (10.2) Sulfuric acid produced by burning sulfur and saltpeter in water 7S(s)+6KNO 3 (s)+4H 2 O(aq) --> 3K 2 S(s)+6NO(g)+4H 2 SO 4 (aq)

Byproducts of Leblanc Process

HCl(g) evaporated H 2 SO 4 (g) NO(g), which produces HNO 3 (g) Soot (C) CO 2 (g) --> widespread acid deposition due mostly to HCl(g) CaS(s), yellow-gray powder, forms hydrogen sulfide and gypsum CaS(s) + H 2 O(aq) --> CaO(s) + H 2 S(g) CaS(s) + O 2 (g) + 3H 2 O(aq) --> CaSO 4 -2H 2 O(s) + H 2 (g) Piles of gypsum (galligu) still exist near old soda-ash factories.

Gypsum/Galligu Piles

Halton Borough Council, UK

www2.halton.gov.uk

Gypsum/Galligu Piles

The Rise of Alkali Factories

Widnes in Cheshire, 1800s, under cloud of LeBlanc Process Halton Borough Council, UK

History of Acid Deposition

1791.

LeBlanc patents technique 1793.

Loses patent to the state during French Revolution 1806.

LeBlanc commits suicide 1838.

Liverpool landowner files complaint against an alkali factory: destroyed crops and interfered with hunting.

Early 1860s.

William Gossage, who built alkali factory in 1830, develops scrubber for HCl(g) by converting a windmill into a tower, filling tower with brushwood, and letting water drip down the brush as smoke rose from the bottom.

www.engpro.com

HCl Scrubber Today

History of Acid Deposition

1861.

Ernst Solvay (1838-1922) develops another way to form soda ash, but it replaces Leblanc process only by 1880s.

1863.

Due to devastation from Alkali factories and Gossage’s scrubber, the 1863 Alkali Act passes in the U.K.

France.

Regulation takes the form of planning laws controlling location of alkali factories.

1866, 1868.

Inventions allow chlorine to be recycled for bleaching powder.

1881.

U.K. Alkali Act is modified since significant pollution of other chemicals from alkali factories still occurs.

Solvay Process

1861.

Solvay rains a salt made of sodium chloride and ammonium down a tower over an upcurrent of CO 2 (g). Only inputs: limestone, salt. Only byproduct: calcium chloride.

Production of carbon dioxide CaCO 3 (s) + heat --> CaO(s) + CO 2 (g) Production of sodium bicarbonate NaCl(s)+NH 3 (g)+CO 2 (g)+H 2 O(aq)-->NaHCO 3 (aq)+NH 4 Cl(s) Production of sodium carbonate, recycling of carbon dioxide 2NaHCO 3 (aq) + heat --> Na 2 CO 3 (aq) + H 2 O(aq) + CO 2 (g) Recycling of ammonia 2NH 4 Cl(s) + CaO(s) --> 2NH 3 (g) + CaCl 2 (s) + H 2 O(aq)

Robert Angus Smith (1817-1884)

Library and Information Centre, Royal Society of Chemistry

Acidity

Measure of the concentration of hydrogen ions (H + ) in solution pH = -log 10 [H + ] [H + ] = molarity (M, moles of H + per liter of solution) Higher [H + ] --> lower pH --> more acidic solution (10.3)

Acidity

In dilute water, the only source of H + is H 2 O(aq) Liquid water H + + OH Hydrogen Hydroxide ion ion [H + ][OH ] = 10 -14 M 2 --> [H + ]=[OH ]=10 -7 M --> pH = -log 10 [10 -7 M] = 7 (10.4)

Acid/Base

Acid Substance that, when added to a solution, dissociates, increasing [H + ], decreasing pH Strong acid: Substances that dissociate readily (e.g., H 2 SO 4 , HCl, HNO 3 ) Weak acids: Substances that dissociate less readily (e.g., H 2 CO 3 ) Base (alkalis) Substances that, when added to a solution, reduce [H + ], increasing pH. (e.g., NH 3 (aq), Ca(OH) 2 (aq))

pH 0 1 2 3 Acid rain, fog (2-5.6)

pH Scale

Natural rainwater (5-5.6) Sea water (7.8-8.3) 4 5 6 7 8 9 10 11 12 13 14 Battery acid (1.0) Lemon juice (2.2) Apples (3.1) Vinegar CH 3 COOH(aq) (2.8) More acidic Distilled water (7.0) Milk (6.6) Baking soda NaHCO 3 (aq) (8.2) Ammonium Lye hydroxide NaOH(aq) NH 4 OH(aq) (11.1) (13.0) Slaked lime Ca(OH) 2 (aq) (12.4) More basic or alkaline Figure 10.3

pH of Food

Lime Lemon Apple Apple juice Tomato juice Buttermilk Banana Cheddar cheese Beef Cabbage Asparagus Milk Crab Egg white 1.9

2.1-2.2

2.9-3.3

4.0

4.4

4.5

4.6

5.3

5.5

5.7

5.9

6.4-6.6

7.0

8.0

Acid Dissociation

Addition of acid to solution increases [H + ], decreasing pH Carbonic acid CO 2 (aq) + H 2 O(aq) Dissolved carbon dioxide Liquid water H 2 CO 3 (aq) Dissolved carbonic acid H + + HCO 3 Hydrogen ion Bicarbonate ion 2H + + CO Hydrogen ion 3 Carbonate ion 2 Sulfuric acid H 2 SO 4 (g) Sulfuric acid gas H 2 SO 4 (aq) Dissolved sulfuric acid H + + HSO 4 Hydrogen Bisulfate ion ion 2H + + SO 4 2 Hydrogen Sulfate ion ion (10.5) - (10.6)

Acid Dissociation

Hydrochloric acid HCl(g) HCl(aq) Hydrochloric acid gas Dissolved hydrochloric acid H + + Cl Hydrogen Chloride ion ion Nitric acid HNO 3 (g) Nitric acid gas HNO 3 (aq) Dissolved nitric acid H + + NO Hydrogen 3 Nitrate ion ion (10.7) - (10.8)

S(IV) and S(VI) Families

S(IV) Family S(VI) Family Table 10.1

Sulfur dioxide SO 2 (g,aq) Sulfurous acid H 2 SO 3 (aq) Bisulfite ion HSO 3 Sulfite ion SO 3 2 O S O O S HO OH O S HO O O S O O Sulfuric acid H 2 SO 4 (g,aq) Bisulfate ion HSO 4 Sulfate ion SO 4 2 O O O O S OH O OH S OH O O S O O

Mechanisms of Converting S(IV) to S(VI)

Why is converting to S(VI) important?

It allows sulfuric acid to enter or form within cloud drops and aerosol particles, increasing their acidity Mechanisms 1. Gas-phase oxidation of SO 2 (g) to H 2 SO 4 (g) followed by condensation of H 2 SO 4 (g) 2. Dissolution of SO 2 (g) into liquid water to form H 2 SO 3 (aq) and its dissociation products, which convert chemically to H 2 SO 4 (aq) and its dissociation products.

Gas-Phase Oxidation of S(IV) to S(VI)

SO 2 (g) + OH(g), M HSO 3 (g) Sulfur dioxide Bisulfite + O 2 (g) SO 3 (g) + H 2 O(g) HO 2 (g) Sulfur trioxide H 2 SO 4 (g) Sulfuric acid (10.9)

Aqueous-Phase Oxidation of S(IV)

Step 1. S(IV) dissolution and dissociation SO 2 (g) Sulfur dioxide gas SO 2 (aq) Dissolved sulfur dioxide SO 2 (aq) + H 2 O(aq) Dissolved sulfur dioxide Liquid water H 2 SO 3 (aq) Sulfurous acid H + + HSO 3 Hydrogen ion Bisulfite ion 2H + + SO 3 2 Hydrogen ion Sulfite ion (10.10) - (10.11)

Aqueous-Phase Oxidation of S(IV)

Step 2. S(IV) oxidation by hydrogen peroxide in solution HSO 3 + H 2 O 2 (aq) + H + Bisulfite ion Dissolved Hydrogen peroxide SO 4 2- + H 2 O(aq) + 2H + Sulfate ion Alternate Step 2. S(IV) oxidation by ozone in solution SO 3 2 + O 3 (aq) Sulfite ion Dissolved ozone SO 4 2- + O 2 (aq) Sulfate ion Dissolved oxygen (10.12) - (10.13)

Production of Nitric Acid Gas

OH(g) + NO 2 (g) Hydroxyl radical Nitrogen dioxide M HNO 3 (g) Nitric acid (10.14)

Effects of Acid Deposition

Lakes and streams Lake pH decreased 1 pH unit between the 1950s and 1960s in Scandinavia. 25,000 out of 90,000 lakes in Sweden were acidified in the 1970s. 17,000 of these were due to anthropogenic pollution.

Fish and microorganisms are sensitive to pH.

Plants Acids damage trees at their roots and erodes cuticle wax.

Acidified water carries away minerals.

Acids release metals Al 3+ , Fe 3+ from the minerals Al(OH) 3 (s), Fe(OH) 3 (s). Metals are toxic to roots.

Prunerov Power Plant, Sulfur Triangle, 2005

www.chmi.cz

Dead Norway Spruce, Bohemian Forest, Czech Republic, 2005

www.chmi.cz

Dead Norway Spruce, Nacetin Forest, Czech Republic, 1995

www.chmi.cz

Dead Spruce, Mt. Mumlava, Czech Republic

www.chmi.cz

Acidified forest near Most, Czechoslovakia (1987)

Owen Bricker, United States Geological Survey

Acidified Forest, Oberwiesenthal, Germany (1991)

Stefan Rosengren/Naturbild

Acidified Creek, New South Wales, 2007

www.ramsar.org

Neutralizing Acids

Add ammonium hydroxide to a lake NH 4 OH(aq) + H + Ammonium Hydrogen hydroxide ion Add slaked lime to a lake Ca(OH) 2 (aq) + 2H + Calcium Hydrogen hydroxide ion NH 4 + + H 2 O(aq) Ammonium Liquid ion water Ca 2+ + 2H 2 O(aq) Calcium Liquid ion water Calcium carbonate is a natural neutralizing agent in soil CaCO 3 (s) + 2H + Calcium carbonate Hydrogen ion Ca 2+ + CO 2 (g) + H 2 O(aq) Calcium ion Carbon dioxide gas Liquid water (10.15) - (10.17)

Liming of a Lake in Sweden

Tero Niemi / Naturbild

Neutralizing Acids

Sea salt is a natural neutralizing agent near the coast NaCl(s) + H + Na + + HCl(g) Sodium chloride Hydrogen ion Sodium ion Hydrochloric acid Ammonia is a neutralizing agent NH 3 (aq) + H + Dissolved Hydrogen ammonia ion NH 4 + Ammonium ion (10.19) - (10.20)

Visibility Trend Sulfur Triangle

Stjern et al. (2011)

1994

U.S. Rainwater pH

2009 U.S. National Atmospheric Deposition Program/National Trends Network

Effects of Acid Deposition

Materials The addition of sulfuric acid to marble or limestone produces gypsum, which forms a crust that creates pits when rain loosens the crust.

Crusting of marble and limestone CaCO 3 (s) + 2H + Calcium carbonate Hydrogen ion Ca 2+ + CO 2 (g) + H 2 O(aq) Calcium ion Carbon dioxide gas Liquid water Ca 2+ + SO 4 2 + 2H 2 O(aq) Calcium Sulfate Liquid ion ion water CaSO 4 -2H 2 O(s) Calcium sulfate dihydrate (gypsum) (10.17) - (10.18)

Cathedral of Learning, Pittsburgh, 1930 and 1934

University Archives, University of Pittsburgh

Sandstone Figure in 1908 and 1968, Westphalia, Germany

Herr Schmidt-Thomsen

Recent Regulation of Acid Deposition

1970. U.S. Clean Air Act Amendments 1977. U.S. National Atmospheric Deposition Program 1979.

Geneva Convention on Long-Range Transboundary Air Pollution 1985. Sulfur Protocol 1988. Nitrogen Oxide Protocol 1990. U.S. Clean Air Act Amendments 1994. Second Sulfur Protocol

Methods of Controlling Emission

Use of low-sulfur coal Remove sulfur from high-sulfur coal Advanced scrubbers (e.g., flue-gas desulfurization)