1. IB Chemistry Topic 8: Acids & Bases Theories, Properties, pH, Strong vs. Weak and Acid Deposition 1

2. 22 8.1 Theories of Acids & Bases A Brønsted–Lowry acid is a proton/H+ donor and a Brønsted–Lowry base is a proton/H+ acceptor." Amphiprotic species can act as both Brønsted–Lowry acids and bases. A pair of species differing by a single proton is called a conjugate acid-base pair.

3. 33 8.1.1Define acids and bases according to the Brǿnsted-Lowry and Lewis theories. Historical Aspects Acid: Comes from Latin word acidus, meaning sour or tart. Alkali (another term for basic): derived from the Arabic word for calcined ashes. Base: Comes from an old English meaning of the word, “to bring low.” When bases are added to acids, they lower the amount of acids.

4. 44 8.1.1Define acids and bases according to the Brǿnsted-Lowry and Lewis theories. Arrhenius Acids & Bases In the 1880’s Svante Arrhenius defined acids & bases as follows: Acids are substances that, when dissolved in water, increase the [H + ] / [H 3 O + ]. Bases are substances that, when dissolved in water, increase the [OH - ].

5. Arrhenius acid is a substance that produces H + (H 3 O + ) in water Arrhenius base is a substance that produces OH - in water 4.3 Arrhenius definition

6. 66 8.1.1Define acids and bases according to the Brǿnsted-Lowry and Lewis theories. Brǿnsted-Lowry Acids & Bases In 1923, Johannes Brǿnsted & Thomas Lowry defined acids & bases: An acid is a proton (H + ) donor A base is a proton (H + ) acceptor Lewis Acids & Bases Gilbert Lewis defined acids and bases: Lewis acid is an electron-pair acceptor Lewis base is an electron-pair donor

7. A Brønsted-Lowry acid is a proton donor A Brønsted-Lowry base is a proton acceptor acidbase

8. Lewis Definition Lewis acid: electron-pair acceptor Lewis base: electron-pair donor

9. 99 8.1.2Deduce whether or not a species could act as a Brǿnsted-Lowry and/or a Lewis acid or base. HCl is an acid HCl(g)  H + (aq) + Cl - (aq) (Arrhenius) –Forms H + ions in aqueous solution HCl(g) + H 2 O(l)  H 3 O + (aq) + Cl - (aq) (Br-Lowry) –H 3 O + (aq) is called the hydronium ion –HCl donates a proton to water H + accepts an electron pair from H 2 O (Lewis acid) –H O H + H

10. 10 8.1.2Deduce whether or not a species could act as a Brǿnsted-Lowry and/or a Lewis acid or base. NaOH & NH 3 are bases NaOH(s)  Na + (aq) + OH - (aq) (Arrhenius) –Forms OH - ions in aqueous solution NH 3 (aq) + H 2 O(l)  NH 4 + (aq) + OH - (aq) (Br-Lowry) –NH 3 accepts a proton from water NH 3 donates an electron pair to H + (Lewis base) H –H N H + H

11. 11 8.1.3Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) In any acid-base equilibrium, the forward reaction and the reverse reaction involve proton transfer: HX(g) + H 2 O(l)  H 3 O + (aq) + X - (aq) In the forward reaction, HX acts as an acid (donates a proton) and H 2 O acts as a base (accepts a proton). In the reverse reaction, H 3 O + acts as an acid (donates a proton) and X - acts as a base (accepts a proton). HX & X - are an acid-conjugate base pair. H 2 O & H 3 O + are a base-conjugate acid pair. They differ only in the presence or absence of a proton.

12. Conjugate Pairs Acids react to form bases and vice versa. Therefore, a pair of species differing by a single proton is called a conjugate acid- base pair.

13. Conjugate Pairs

14. A Brønsted-Lowry acid is a proton donor A Brønsted-Lowry base is a proton acceptor acid conjugate base base conjugate acid

15. 15`15 8.1.3Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) HNO 2 (aq) + H 2 O(l)  H 3 O + (aq) + NO 2 - (aq) acidbase conj acid conj base A B-L acid loses a proton (H + ) to form the conjugate base. A B-L base gains a proton (H + ) to form the conjugate acid. remove H + add H +

16. 16`16 8.1.3Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) H 2 O(l) + NH 3 (aq)  NH 4 + (aq) + OH - (aq) acidbase conj acid conj base A B-L acid loses a proton (H + ) to form the conjugate base. A B-L base gains a proton (H + ) to form the conjugate acid. remove H + add H +

17. Amphiprotic Substances Amphiprotic species can both donate hydrogen ions (protons) and accept them. Just to confuse you, there is also a term, amphoteric: can act as both an acid and a base –All amphiprotic substances are also amphoteric, but the converse is not true. –Aluminum oxide is amphoteric, but not amphiprotic. 17

18. 18`18 8.1.3Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) Consider the reaction: H 2 O(l) + HSO 3 - (aq)  H 3 O + (aq) + SO 3 2- (aq) Identify the acid, the base, the conjugate acid and the conjugate base.

19. 19`19 8.1.3Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) Consider the reaction: H 2 O(l) + HSO 3 - (aq)  H 3 O + (aq) + SO 3 2- (aq) Identify the acid: HSO 3 - (aq) the base: H 2 O(l) the conjugate acid: H 3 O + (aq) the conjugate base: SO 3 2- (aq)

20. 20`20 8.1.3Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) What is the conjugate base for each of the following acids: HClO 4 ; H 2 S; PH 4 + ; HCO 3 - ?

21. 21`21 8.1.3Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) What is the conjugate base for each of the following acids: HClO 4 ; H 2 S; PH 4 + ; HCO 3 - ? The conjugate base for HClO 4 is ClO 4 - since a B-L acid loses a proton (H + ). What is the conjugate acid of each of the following bases: CN - ; SO 4 2- ; H 2 O; HCO 3 - ? The conjugate acid for CN- is HCN since a B-L base gains a proton (H + ).

22. 22`22 8.1.3Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) What is the conjugate acid for each of the following bases: HClO 4 ; H 2 S; PH 4 + ; HCO 3 - ?

23. 23`23 8.1.3Deduce the formula of the conjugate acid (or base) of any Brǿnsted-Lowry base (or acid) What is the conjugate base for each of the following acids: HClO 4 ; H 2 S; PH 4 + ; HCO 3 - ? The conjugate base for HClO 4 is ClO 4 - since a B-L acid loses a proton (H + ). What is the conjugate acid of each of the following bases: CN - ; SO 4 2- ; H 2 O; HCO 3 - ? The conjugate acid for CN- is HCN since a B-L base gains a proton (H + ).

24. 24 IB Topic 8: Acids & Bases 8.2 Properties of Acids & Bases Most acids have observable characteristic chemical reactions with reactive metals, metal oxides, metal hydroxides, hydrogen carbonates and carbonates. Salt and water are produced in exothermic neutralization reactions.

25. Review Minute Dissolving –A physical process –Separates ionic compounds into ions Ex. CaCl 2 separates into 1 Ca +2 ion and 2 Cl -1 ions We also call this dissociation Salt –Any ionic compound (formed when the hydrogen of an acid is replaced by a metal or another positive ion) 25

26. 26 8.2.1Outline the characteristic properties of acids and bases in aqueous solutions Acids Sour taste Electrolytes… Huh?? –  Conductive due to dissociation 3 main types of reactions to form salts: 1.Acid + metal (single displacement) 2.Acid + base (double displacement / neutralization) 3.Acid + carbonate (double displacement with effervescence)

27. 27 8.2.1Outline the characteristic properties of acids and bases in aqueous solutions Acids 1.Acid + metal (single displacement) HCl (aq) + Zn (s)  H 2 SO 4 (aq) + Fe (s)  assume Fe (II) CH 3 COOH (aq) + Mg (s) 

28. 28 8.2.1Outline the characteristic properties of acids and bases in aqueous solutions Acids 1.Acid + metal (single displacement) HCl(aq) + Zn(s)  ZnCl 2 (aq) + H 2 (g) H 2 SO 4 (aq) + Fe(s)  FeSO 4 (aq) + H 2 (g) CH 3 COOH(aq) + Mg(s)  Mg(CH 3 COO) 2 (aq) + H 2 (g) All form salt and hydrogen (gas).

29. 29 8.2.1Outline the characteristic properties of acids and bases in aqueous solutions Acids 2. Acid + base (double displacement / neutralization) HCl(aq) + NaOH(aq)  HNO 3 (aq) + NH 4 OH(aq)  CH 3 COOH(aq) + CuO(s) 

30. 30 8.2.1Outline the characteristic properties of acids and bases in aqueous solutions Acids 2. Acid + base (double displacement / neutralization) HCl(aq) + NaOH(aq)  NaCl(aq) + H 2 O(l) HNO 3 (aq) + NH 4 OH(aq)  NH 4 NO 3 (aq) + H 2 O(l) CH 3 COOH(aq) + CuO(s)  Cu(CH 3 COO) 2 (aq) + H 2 O(l) All are exothermic (release heat) and form salt and water.

31. 31 8.2.1Outline the characteristic properties of acids and bases in aqueous solutions Acids 3. Acid + carbonate (double displacement  decomp) HCl(aq) + CaCO 3 (s)  H 2 SO 4 (aq) + Na 2 CO 3 (aq)  HCl(aq) + NaHCO 3 (aq)  Just a heads up… carbonic acid is VERY unstable and breaks down into two very common substances.

32. 32 8.2.1Outline the characteristic properties of acids and bases in aqueous solutions Acids 3. Acid + carbonate (double displacement  decomp) HCl(aq) + CaCO 3 (s)  CaCl 2 (aq) + H 2 O(l) + CO 2 (g) H 2 SO 4 (aq) + Na 2 CO 3 (aq)  Na 2 SO 4 (aq) + H 2 O(l) + CO 2 (g) HCl(aq) + NaHCO 3 (aq)  NaCl(aq) + H 2 O(l) + CO 2 (g) All form salt, water and carbon dioxide.

33. 33 8.2.1Outline the characteristic properties of acids and bases in aqueous solutions Acids Reactions with indicators. Indicators are substances that change color when the concentration of hydrogen ions changes. Ex. phenolpthalein, bromothymol blue, litmus –H + + In -  HIn –In - is one color, HIn is a different color –Universal indicator is a mixture of different indicators and produces a range of colors.

34. pH Indicators 34

35. 35 8.2.1Outline the characteristic properties of acids and bases in aqueous solutions Bases Bitter taste Feel slippery due to reaction with the oils on your skin. This forms a soap. Conductive due to dissociation Neutralize acids (see acid properties). Reactions with indicators. Bases DO NOT react with indicators but the addition of a base changes the [H + ] which affects the indicator. –H + (aq) + OH-(aq)  H 2 O(l)

36. 36 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3. Metallic Oxides in Period 3 Sodium oxide: Na 2 Oionic Magnesium oxide: MgOionic Aluminum oxide: Al 2 O 3 ionic Metalloid oxide in Period 3 Silicon dioxide: SiO 2 covalent Nonmetallic oxides in Period 3 Tetraphosphorus decoxide: P 4 O 10 covalent Sulfur trioxide: SO 3 covalent Dichlorine heptoxide: Cl 2 O 7 covalent

37. 37 3.3.2 Discuss the changes in nature, from ionic to covalent and from basic to acidic, of the oxides across period 3. Acidic/Basic Metallic oxides in Period 3 are basic Sodium oxide: Na 2 O + H 2 O  2 NaOH basic Magnesium oxide: MgO + H 2 O  Mg(OH) 2 basic Aluminum oxide: Al 2 O 3 + H 2 O  2 Al(OH) 3 amphoteric Metalloid oxide in Period 3 is acidic Silicon dioxide:SiO 2 + H 2 O  H 2 SiO 3 acidic Nonmetallic oxides in Period 3 are acidic Tetraphosphorus decoxide: P 4 O 10 + 6H 2 O  4H 3 PO 4 acidic Sulfur trioxide: SO 3 + H 2 O  H 2 SO 4 acidic Dichlorine heptoxide: Cl 2 O 7 + H 2 O  2HClO 4 acidic Argon does not form an oxide

38. 38 IB Topic 8: Acids & Bases 8.3 The pH scale pH = − log[H+(aq)] and [H + ] = 10 −pH A change of one pH unit represents a 10-fold change in the hydrogen ion concentration [ + ] pH values distinguish between acidic, neutral and alkaline solutions The ionic product constant, = [H + ][OH − ] = 10 −14 at 298 K

39. The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H + (or OH - ) ion. Under 7 = acid 7 = neutral Over 7 = base

40. pH of Common Substances

41. 41 8.4.1Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale. pH is a measure of the [H + (aq)] in a solution. pH = -log [H + (aq)] [H + ] = 10 −pH In a neutral solution the [H + (aq)] = 1 x 10 -7 so the pH = -log(1 x 10 -7 ) = 7 Addition of an acid to water increases the [H + (aq)] so the pH < 7. Low pH is acidic. Addition of a base to water decreases the [H + (aq)] so the pH > 7. High pH is basic.

42. Calculating the pH pH = - log [H+] (Remember that the [ ] means Molarity aka mol dm -3 ) Example: If [H + ] = 1 X 10 -10 pH = - log 1 X 10 -10 pH = - (- 10) pH = 10 Example: If [H + ] = 1.8 X 10 -5 pH = - log 1.8 X 10 -5 pH = - (- 4.74) pH = 4.74

43. 43 8.4.1Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale. pH is a measure of the [H + (aq)] in a solution. pH = -log [H + (aq)] pH 7 is basic Is a solution with a [H + (aq)] = 1 x 10 -5 acidic, basic or neutral? Find the pH: pH = -log 1 x 10 -5 = 5 so it is acidic Is a solution with a [H + (aq)] = 1 x 10 -10 acidic, basic or neutral? Find the pH: pH = -log 1 x 10 -10 = 10 so it is basic

44. Try These! Find the pH of these: 1)A 0.15 M solution of Hydrochloric acid 2) A 3.00 X 10 -7 M solution of Nitric acid pH = - log [H+] pH = - log 0.15 pH = - (- 0.82) pH = 0.82 pH = - log 3 X 10-7 pH = - (- 6.52) pH = 6.52

45. 45 8.4.1Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale. You can also calculate the [H + (aq)] if you know the pH [H + (aq)] = inverse log(-pH) In a solution with a pH = 6, what is the [H + (aq)]? [H + (aq)] = inverse log(-6) = 1 x 10 -6 Is a solution with a pH = 11, what is the [H + (aq)]? [H + (aq)] = inverse log(-11) = 1 x 10 -11

46. 46 8.4.2Identify which of two or more aqueous solutions is more acidic or alkaline using pH values. Acid/AlkalinepH More basic14 12 10 Less basic8 Neutral7 Less acidic6 4 2 More acidic0

47. 47 8.4.3State that each change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H + (aq)]. An increase in one pH unit decreases the [H + (aq)] by a factor of 10. A decrease in one pH unit increases the [H + (aq)] by a factor of 10. Acid/AlkalinepH [H + (aq)] More basic111 x 10 -11 101 x 10 -10 91 x 10 -9 Less basic81 x 10 -8 Neutral71 x 10 -7 Less acidic61 x 10 -6 51 x 10 -5 41 x 10 -4 More acidic31 x 10 -3

48. 48 8.4.4Deduce changes in [H + (aq)] when the pH of a solution changes by more than one pH unit. An increase in three pH units decreases the [H + (aq)] by a factor of 1,000. A decrease in five pH units increases the [H + (aq)] by a factor of 10 5. Acid/AlkalinepH [H + (aq)] More basic141 x 10 -14 121 x 10 -12 101 x 10 -10 Less basic81 x 10 -8 Neutral71 x 10 -7 Less acidic61 x 10 -6 41 x 10 -4 21 x 10 -2 More acidic01 x 10 0

49. 49 IB Topic 8: Acids & Bases 8.4 Strong and weak acids & bases Strong and weak acids and bases differ in the extent of ionization. Strong acids and bases of equal concentrations have higher conductivities than weak acids and bases. A strong acid is a good proton donor and has a weak conjugate base. A strong base is a good proton acceptor and has a weak conjugate acid.

50. Review Minute Dissociation –Separates ionic compounds into ions Ex. CaCl 2 separates into 1 Ca +2 ion and 2 Cl -1 ions We also call this ionization –Separates acids into positive hydrogen ions and anions Ex. H 2 SO 4 separates into 2 H +1 ions and 1 SO 4 -2 ion 50

51. Strong vs. Weak everythingClassifying an acid or base has everything to do with how many of the acid or base compounds dissociate in solution –Strong = MOST –Weak = FEW https://www.youtube.com/watch?v=rKqYE5sZi1s Sim: https://phet.colorado.edu/sims/html/acid-base- solutions/latest/acid-base-solutions_en.htmlhttps://phet.colorado.edu/sims/html/acid-base- solutions/latest/acid-base-solutions_en.html 51

52. 52 Strong acids & bases Strong acids and bases dissociate completely in water Acid:HX(aq) + H 2 O(l) H 3 O + (aq) + X - (aq) Base: YOH(aq) Y + (aq) + OH - (aq) When strong acids and bases dissolve in water, the solution consists of almost entirely of ions with a negligible amount of molecules. Write the dissociation reactions for H 2 SO 4 & Ba(OH) 2

53. 53 8.4.1Distinguish between aqueous solutions that are acidic, neutral or alkaline using the pH scale. Weak acids and bases only partly dissociate in water. Acid:HX(aq) + H 2 O(l) H 3 O + (aq) + X - (aq) Base: NH 3 (aq) + H 2 O(l) NH 4 + (aq) + OH - (aq) When weak acids and bases dissolve in water, the solution consists of almost entirely of molecules with a negligible amount of ions. Kc (Ka or Kb) << 1 Write the dissociation reactions for H 2 CO 3 & C 2 H 5 NH 2

54. 54 Distinguishing between strong and weak acids and bases. 3 properties that can be used to differentiate a strong acid or base from a weak one: 1.Electrical conductivity 2.Rate of reaction 3.pH Strong acids dissociate completely. A 0.010M HCl solution has a [H + (aq)] = 0.010M so the pH = 2

55. 55 Testing based on electrical conductivity Strong acids and bases are good conductors of electricity Since strong acids and bases dissociate completely into ions, these solutions are excellent conductors of electricity. However, there can be weak solutions of strong acids and bases which lowers the ion concentration and lowers the conductivity. Weak acids and bases are poor conductors of electricity Since weak acids and bases dissociate very little, few ions are present in solution so these are weak conductors of electricity.

56. 56 Testing based on reaction rate Reaction rate of acids depends on [H + ]. Reaction rate will, therefore, increase with an increased [H + ] Strong acids  faster rates of reaction / more vigorous reaction –Gas is produced more quickly Weak acids  slower rates of reaction / less vigorous reaction Do not say that strong acids produce more gas, just that they produce it at a faster rate

57. 57 Testing based on pH Strong acids dissociate completely. A 0.010M HCl solution has a [H + (aq)] = 0.010M so the pH = 2 Weak acids do not dissociate completely so their [H + (aq)] concentration will be less than that of an equal concentration of a strong acid  higher pH A 0.010M ethanoic acid (CH 3 COOH) solution has a [H + (aq)] = 0.00042M so the pH = 3.4 A 0.01M carbonic acid (H 2 CO 3 ) solution has a [H + (aq)] = 6.6 x 10 -5 so the pH = 4.2

58. 58 Strong Acids & Bases Strong acids and bases completely dissociate into their ions in aqueous solutions. Strong acidsStrong bases Hydrochloric acid, HCl(aq)Lithium hydroxide, LiOH(aq) Nitric acid, HNO 3 (aq)Sodium hydroxide, NaOH(aq) Sulfuric acid, H 2 SO 4 (aq)Potassium hydroxide, KOH(aq) Barium hydroxide, Ba(OH) 2 (aq)

59. 59 Weak Acids & Bases Weak acids and bases only slightly dissociate into their ions in aqueous solution Weak acids Weak bases Ethanoic (acetic) acid CH 3 COOH(aq) Ammonia, NH 3 (aq) Other organic acids (look like train sounds… choo choo!!) Ethylamine, C 2 H 5 NH 2 (aq) Carbonic acid H 2 CO 3 (aq) (CO 2 in water) Other amines (organic nitrogen- containing compounds) Phosphoric acid, H 3 PO 4 (aq)

60. 60 8.3.3Distinguish between strong and weak acids and bases, and determine the relative strengths of acids and bases, using experimental data. A strong acid is a good proton donor and has a weak conjugate base. A strong base is a good proton acceptor and has a weak conjugate acid. Which acid in each of the following pairs has the stronger conjugate base? A) H 2 CO 3 or H 2 SO 4 B) HCl or HCOOH

61. 61 8.3.3Distinguish between strong and weak acids and bases, and determine the relative strengths of acids and bases, using experimental data. Hank #8: http://www.youtube.com/watch?v=ANi709MYnWg&inde x=8&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr http://www.youtube.com/watch?v=ANi709MYnWg&inde x=8&list=PL8dPuuaLjXtPHzzYuWy6fYEaX9mQQ8oGr Lab & Discussion: Properties of Acids & Bases

62. Solve problems involving concentration, amount of solute and volume of solution. Dilutions are typically used to make lower concentrations of concentrated acids. Solutions used in lab are prepped by diluting the concentrated versions with water. C 1 x V 1 = C 2 x V 2 AAA = Always Add Acid… REMEMBER THIS!!! The concentrated acid needs to always be added to the water. Water should NEVER be added to concentrated acids!!!

63. Solve problems involving concentration, amount of solute and volume of solution. Dilutions What volume of 3.00 mol dm -3 HCl(aq) would you need to use to make 500 cm 3 of 1.00 mol dm -3 HCl (aq)?

64. Titrations A titration is a method for determining the concentration of a solution by reacting a known volume of that solution with a solution of known concentration If you wish to find the concentration of an acid solution, you would titrate the acid solution with a solution of a base of known concentration. You could also titrate a base of unknown concentration with an acid of known concentration. C a x V a = C b x V b Multiply appropriate side by subscript 4.7

65. In the titration of an acid by a base, the pH meter measures the pH of the acid solution in the beaker as a solution of a base with a known concentration is added from the buret. http://wps.prenhall.com/wps/media/objects/3 312/3392202/blb1703.html 65

66. Titrations In a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Equivalence point – the point at which the reaction is complete Indicator – substance that changes color at the endpoint (hopefully close to the equivalence point) Slowly add base to unknown acid UNTIL The indicator changes color (pink) 4.7

67. Strong Acid-Strong Base Titrations NaOH (aq) + HCl (aq) H 2 O (l) + NaCl (aq) OH - (aq) + H + (aq) H 2 O (l) 16.4 100% ionization! No equilibrium

68. Strong Acid-Strong Base Titrations NaOH (aq) + HCl (aq) H 2 O (l) + NaCl (aq) OH - (aq) + H + (aq) H 2 O (l) 16.4 100% ionization! No equilibrium

69. 69 IB Topic 8: Acids & Bases 8.5 Acid deposition Rain is naturally acidic because of dissolved CO 2 and has a pH of 5.6. Acid deposition has a pH below 5.6. Acid deposition is formed when nitrogen or sulfur oxides dissolve in water to form HNO 3, HNO 2, H 2 SO 4 and H 2 SO 3. Sources of the oxides of sulfur and nitrogen and the effects of acid deposition should be covered. Fuse: https://www.youtube.com/watch?v=Nf8cuvl62Vc

70. Acid deposition refers to how acidic particulars leave the atmosphere. A well known example is acid rain. The term “acid rain” was coined in 1872 by Robert Angus Smith, an English scientist who observed that acidic precipitation could damage plants and materials. However it wasn’t around the 1960’s or 70’s that acid deposition become a serious environmental issue, when scientist discover low pH level in lakes and streams.

71. Acid Rain Natural rain is acidic with a pH level around 5.6. H 2 O + CO 2  H 2 CO 3 Only a very small percentage is in rain. Typical acid rain has a pH level of 4.0 and pH level of 4.2 in lakes wouldn’t be able to support life.

72. DRY DEPOSITION Dry Deposition refers how the acidic particulars leave the atmosphere without the presence of precipitation. These particulars leave the atmosphere due to gravity, and these acidic gases such as sulfur dioxide have a direct harmful effect on the environment because the gases haven’t dissolved in the rain water. WET DEPOSITION Wet deposition refers how the acidic particulars leave the atmosphere through precipitation. Either by rain, snow, or fog.

73. The main source of acidity in the atmosphere is sulfur oxides produced from power plants. These sulfur acidity react with react in rain water. Two types of acids are formed from there sulfur oxide. SO 2 + H 2 O  H 2 SO 3 Sulfurous acid or SO 3 + H 2 O  H 2 SO 4 Sulfuric Acid Nitrogen oxides also contribute to acid rain. These nitrogen oxides are formed from vehicle engines. This gas combines with hydroxyl radical then forms with nitrous acid. HNO 2

74. Acid deposition effects the environment in 5 ways: 1.pH level of lakes and streams, and organism in them 2.The availability of metal ions in the soil, and therefore affects nearby plant life and water 3.Directly affects plant life 4.Affects buildings & other structures 5.Affects human health

76. A pH level below 5.5 would kill some species of fish like Salmon, also kill algae and zooplankton which depletes the food for larger organisms. At low pH levels eggs are unable to hatch. The pH of soil is a key factor whether or not if plants will grow. Not only does it damage the soil, it lowers the amount of nutrients that plants need. Acid deposition directly affects of plants by turning leaves brown and reduces photosynthetic ability of the plant

77. Majority of historical buildings are made of limestone and marble which are forms of calcium carbonate which acid rain erodes. CaCO 3 +H 2 SO 4  CaSO 4 +H 2 O+ CO 2 As for metallic building those made of iron or steel are readily attack by acid deposition by both dry and wet deposition. Dry deposition Fe+ SO 2+ O 2  FeSO 4 Wet deposition Fe + H 2 SO 4  FeSO 4 + H 2

78. To counteract acid deposition is by reducing the amount of sulfur and nitrogen oxides released in the atmosphere. Nitrogen oxides have been reduce from vehicle emissions using catalytic converters. EPA’s acid rain program focuses on power plants, the largest single source of SO 2 emissions, and a major source of NO x emissions by issuing permits to power plants of the amount of emissions being released. Limestone or calcium hydroxide is being use to increase the ph level in soil and lakes.

80. Terms to Know Brønsted-Lowry acid & base Lewis acid & base Strong acid & base Weak acid & base pH scale Buffer Conjugate acid & base Indicator