Chapter 6: Neutralizing the Threat of Acid Rain

1. Chapter 6: Neutralizing the Threat of Acid Rain Is normal rain acidic? Is acid rain worse in some parts of the country? Is there a way to “neutralize” acid rain?

2. What does the word acid mean to you?

3. There are several ways to define an acid. The most common definition for a acid is: Any substance that causes hydrogen ions to be formed in an aqueous solution. Consider hydrogen chloride gas, dissolved in water: H2O HCl(g) H+(aq) + Cl–(aq) (a proton) 6.1

4. A H+ ion (a proton) is very reactive and does not actually exist in water. Instead, when the ion is formed it reacts with a water molecule to form a hydronium ion, H3O+. H+ (aq) + H2O → H3O+ (aq) The chloride ion (Cl-) remains unchanged, but is surrounded by water molecules to keep it dissolved. The hydronium ion is often abbreviated: H3O+ (aq) = H+ (aq) and in many cases, aqueous is assumed, so the (aq) is often not added. + 6.1

5. Therefore, the same reaction between gaseous HCl and water could be written in two different ways: H2O HCl(g) H+(aq) + Cl–(aq) (a proton) or: HCl(g) + H2O(l) H3O+(aq) + Cl–(aq) hydronium ion

6. The most common definition for a base is: Any substance that causes hydroxide ions (OH-) to be produced in aqueous solution. NaOH (s) + H2O → Na+ (aq) + OH- (aq) Notice that when written in this manner, the equation cannot be balanced. In this case water is the solvent, not a reactant, and more properly written above the arrow to indicate reaction conditions: H2O NaOH(s) Na+(aq) + OH–(aq) Sodium hydroxide Hydroxide ion Sodium ion

7. Calcium hydroxide produces 2 equivalents of OH– : H2O Ca(OH)2(s) Ca2+(aq) + 2OH–(aq) What about ammonia (NH3)? It is a base, but has no OH group. NH3(aq) + H2O(l) NH4OH(aq) ammonium hydroxide H2O NH4OH(aq) NH4+(aq) + OH–(aq) 6.2

8. Acid/Base Strength • The chemistry definition of strength refers only to how easily the acid or base forms ions in water: • Strong acids/bases completely form ions • Weak acids/bases only partially form ions, so stay as the initial compound • Strength has nothing to do with concentration or reactivity

9. Strong acids: fully ionize to form H+ ions: HCl (aq) → H+1 (aq) + Cl-1 (aq) 99.9% of HCl molecules split apart • Weak acids: only a small percentage of the molecules form hydronium ions: HF (aq) ↔ H+1 (aq) + F-1 (aq) less than 1% of HF molecules split apart

10. Soluble ionic compounds completely split apart into individual ions in water. • Therefore, soluble metal hydroxides fully split apart and are strong bases: NaOH (aq) → Na+1 (aq) + OH-1 (aq) • For 'insoluble' metal hydroxides, only a very small amount of ions are formed: FeOH (s) ↔ Fe+1 (aq) + OH-1 (aq) most of the compound remains as a solid

11. Strong/weak – refers to % ionization Concentrated/dilute – refers to the amount of acid/base per liter of solution Reactivity – depends on the acid or base, and with what it is reacting

12. HCl • Strong acid • Concentrated in your stomach • Does not react readily with stomach lining • Does react to digest food particles

13. HF • Weak acid • Extremely reactive with most substances • Even dilute solutions eat through glass

14. HC2H3O2 – acetic acid • Weak acid • Dilute in vinegar • used to make pickles, salad dressing... • Concentrated solutions would be hazardous to ingest.

15. When acids and bases react with each other, we call this aneutralizationreaction. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) acid + base salt + water In neutralization reactions, hydrogen ions from an acid combine with the hydroxide ions from a base to form molecules of water (hydrogen hydroxide). The other product is a salt (an ionic compound). 6.3

16. There are three ways chemical equations are typically written: • Chemical equation (standard form): shows reactants and products properly balanced. • Ionic equation: shows ions that are present in solution. • Net ionic equation: shows only ions and compounds that participate in the reaction. • Spectator ions – appear on both sides of the ionic equation, and not shown in net ionic equations.

17. Consider the reaction of hydrobromic acid with barium hydroxide. This reaction may be represented with a molecular, ionic, or net ionic equation: Molecular (standard chemical equation): 2 HBr(aq) + Ba(OH)2(aq) BaBr2(aq) + 2 H2O(l) Ionic: 2 H+(aq) + 2 Br–(aq) + Ba2+(aq) + 2 OH–(aq) Ba2+(aq) + 2 Br–(aq) + 2 H2O(l) Net Ionic: 2 H+(aq) + 2 OH–(aq) 2 H2O(l) or by dividing both sides of the equation by 2 to simplify it: H+(aq) + OH–(aq) H2O(l) 6.3

18. 2 HBr(aq) + Ba(OH)2(aq) BaBr2(aq) + 2 H2O(l) How did we go from Ionic to Net Ionic? Ionic: 2 H+(aq) + 2 Br–(aq) +Ba2+(aq)+ 2 OH–(aq)Ba2+(aq)+ 2 Br–(aq) +2H2O(l) Remove the species that appear unchanged on both sides of the reaction – these are calledspectator ions. 6.3

19. 2 HBr(aq) + Ba(OH)2(aq) BaBr2(aq) + 2 H2O(l) How did we go from Ionic to Net Ionic? Ionic: 2 H+(aq) + 2 Br–(aq)+ Ba2+(aq) + 2 OH–(aq)Ba2+(aq) + 2 Br–(aq) + 2 H2O(l) spectatorions Net Ionic: H+(aq) + OH–(aq) H2O(l) 6.3

20. The hydroxide and hydronium concentrations in water are related by the expression: Kw = [H+][OH–] = 1 x 10–14 (at 25 oC) where Kw is theion-product constantfor water, and remember that [x] means molarity (moles/liters) of x. Since Kw is a constant, as the hydronium concentration is increased, the hydroxide concentration decreases, and vice versa. Knowing the hydroxide ion concentration, we can calculate the [H+] and vice versa. 6.3

21. Kw = [H+][OH–] = 1 x 10–14 (at 25 oC) This expression shows that both ions, H+ and OH-, are always present in water. For acidic solutions, the hydronium concentration is greater. For basic solutions, the hydroxide concentration is greater. For neutral solutions, the two concentrations are equal (10-7 M), not zero. 6.3

22. pHis used to compare the concentration of the H+ ions present in aqueous solutions (easier to compare than using exponential values). The mathematical expression for pH is a log-based scale and is represented as: pH = –log[H+] So for a solution with a [H+] = 1.0 x 10–3 M, the pH = –log(1.0 x 10–3), or –(–3.0) = 3.0 Since pH is a log scale based on 10, a pH change of 1 unit represents a power of 10 change in [H+]. That is, a solution with a pH of 2 has a [H+] ten times that of a solution with a pH of 3. 6.4

23. Measuring pH with a pH meter Or with litmus paper Base applied to red litmus Acid applied to blue litmus 6.6

24. To measure the pH in basic solutions, we make use of the expression below to calculate [H+] from [OH–]. Kw = [H+][OH–] = 1 x 10–14 (at 25 oC) The three possible aqueous solution situations are: [H+] = [OH–] a neutral solution (pH = 7) [H+] > [OH–] an acidic solution (pH < 7) [H+] < [OH–] a basic solution (pH > 7) 6.4

25. Common substances and their pH values Note that “normal” rain is slightly acidic. 6.4

26. Rain water is naturally slightly acidic due to dissolved carbon dioxide: CO2 (g) ↔ CO2 (aq) CO2 (aq) + H2O (l) ↔ H2CO3 (aq) H2CO3 (aq) ↔ H+1 (aq) + HCO3-1 (aq) In fact any pure water left exposed to air will become slightly acidic due to small amounts of dissolved carbon dioxide

27. Rainwater is naturally acidic, but ocean water is basic. Three chemical species responsible for maintaining ocean pH: Base Base Acid 6.5

28. Many ocean creatures, coral, mollusks, urchins, have shells made out of CaCO3 They get the CO3-2 to make their shells from dissolved ions in the ocean The shells from these dead animals are buried in the ocean floor, but help keep the CO3-2 (aq) concentration constant and neutralize acid from dissolved CO2

29. Over the past 200 years, the amount of carbon dioxide in the atmosphere has increased, so more carbon dioxide is dissolving in the ocean and forming carbonic acid. Ocean acidification –the lowering of ocean pH due to increased atmospheric carbon dioxide The H + produced from the dissociation of carbonic acid reacts with carbonate ion in seawater to form the bicarbonate ion: H+(aq) + CO32–(aq) → HCO3–(aq) This reduces the concentration of carbonate ion in seawater, so the calcium carbonate in the shells of sea creatures begins to dissolve to maintain the concentration of carbonate ions in seawater: CaCO3(s) → Ca2+(aq) + CO32–(aq) The buffering from the buried shells is too slow of a process to compensate for the carbon dioxide changes over the last 200 years

30. Chemistry of Carbon Dioxide in the Ocean 6.5

31. Why is rain naturally acidic? Carbon dioxide in the atmosphere dissolves to a slight extent in water and reacts with it to produce a slightly acidic solution of carbonic acid: CO2(g) + H2O(l) H2CO3(aq) carbonic acid H2CO3(aq) H+(aq) + HCO3–(aq) The carbonic acid dissociates slightly leading to rain with a pH around 5.3. 6.6

32. 6.6

33. But acid rain can have pH levels lower than 4.3 (10 times the normal acidity) – where is the extra acidity coming from? The most acidic rain falls in the eastern third of the United States, with the region of lowest pH being roughly the states along the Ohio River valley. The extra acidity must be originating somewhere in this heavily industrialized part of the country. 6.6

34. But acid rain can have pH levels lower than 4.3 (10 times the normal acidity) Analysis of rain for specific compounds confirms that the chief culprits are the oxides of sulfur and nitrogen: sulfur dioxide (SO2), sulfur trioxide (SO3), nitrogen monoxide (NO), and nitrogen dioxide (NO2). These compounds are collectively designated SOx and NOx and are often referred to as “Sox and Nox.” 6.6

35. Oxides of sulfur and nitrogen are acid anhydrides, literally “acids without water.” SOx react with water to form acids: SO2(g) + H2O(l) H2SO3(aq) sulfurous acid SO3(g) + H2O(l) H2SO4(aq) sulfuric acid And then: H2SO4(aq) 2 H+(aq) + SO42–(aq) 6.6

36. What about the NOx? 4 NO2(g) + 2 H2O(l) + O2(g) 4 HNO3(aq)nitric acid Like sulfuric acid, nitric acid also dissociates to release the H+ ion: HNO3(aq)H+(aq) + NO3–(aq) 6.6

37. Sulfur dioxideemissions are highest in regions with many coal-fired electric power plants, steel mills, and other heavy industries that rely on coal. Allegheny County, in western Pennsylvania, is just such an area, and in 1990 it led the United States in atmospheric SO2 concentration. 6.7

38. How does the sulfur get into the atmosphere? The burning of coal. Coal contains 1–3% sulfur and coal burning power plants usually burn about 1 million metric tons of coal a year! Burning of sulfur with oxygen produces sulfur dioxide gas, which is poisonous. S(s) + O2(g) SO2(g) Once in the air, the SO2 can react with oxygen molecules to form sulfur trioxide, which acts in the formation of aerosols. 2 SO2(g) + O2(g) 2 SO3(g) 6.7

39. The highestNOxemissions are generally found in states with large urban areas, high population density, and heavy automobile traffic. Therefore, it is not surprising that the highest levels of atmospheric NO2 are measured over Los Angeles County, the car capital of the country. Nitrogen dioxide gas in the atmosphere reacts with the hydroxyl radical to form nitric acid. NO2(g) + ·OH(g) → HNO3(l) 6.8

40. Emissions of NOx and SO2 in the U.S. 6.10

41. These SOx and NOx compounds react with airborne moisture to acidify rain. But if no water vapor is present, they may be trapped in particulates on dry land. When rain does come, they then form acids. The term 'acid rain' is used to identify both processes of acidifying rain water.

42. Some effects of acid rain are: • Erosion of buildings • Limestone and marble are mainly CaCO3(a base), and react with acids just as coral shells do. • Rusting of metal in buildings and cars

43. Effects of acid rain: damage to marble 1944 At present These statues are made of marble, a form of limestone composed mainly of calcium carbonate, CaCO3. Limestone and marble slowly dissolve in the presence of H+ ions: CaCO3(s) + 2 H+(aq) Ca2+(aq) + CO2(g) + H2O(l) 6.11

44. Effects of acid rain: rusting metal Iron metal dissolves when exposed to hydrogen ions: 4 Fe(s) + 2 O2(g) + 8 H+(aq) → 4 Fe2+(aq) + 4 H2O(l) Now the aqueous Fe2+ ions react with oxygen to form rust (Fe2O3): 4 Fe2+(aq) + O2(g) + 4 H2O(l) → 2 Fe2O3(s) + 8 H+(aq) Billions of dollars are spent annually to protect bridges, cars, buildings, and ships from the reactions above. 6.11

45. Human health • Asthma • Bronchitis • Even death in some cases • Visibility • Causes haze or smog

46. “They stop at Overlooks, idling their engines as they read signs describing Shenandoah’s ruggedly scenic terrain: Hogback, Little Devil Stairs, Old Rag, Hawksbill. Jammed along the narrow road, the cars and motor homes add to the miasmic summer haze that cloaks the hills.” Ned Burks, Environmental writer describing Skyline Drive in Virginia 6.12

47. Harms surface fresh water • More acidic water can kill some or all plants and animals in lakes and rivers • Some regions have natural limestone bedrock which helps neutralize the acid • Regions with more granite bedrock do not have this buffering ability

48. Effects of acid rain: damage to lakes and streams 6.13

49. Harms forests • Stunts growth • Weakens/kills plants • Acid rain removes nutrients from the soil • Harms entire ecosystem

50. Simplified Nitrogen Cycle Plants need nitrogen in a chemical form that reacts more easily, such as the ammonium ion, ammonia, or the nitrate ion, in order to grow. 6.9