1. Unit 13 Acids and Bases Slides adapted from Nivaldo Tro

2. Tro - Chapter 142 Types of Electrolytes salts = water soluble ionic compounds all strong electrolytes acids = form H +1 ions in water solution bases = combine with H +1 ions in water solution increases the OH -1 concentration  may either directly release OH -1 or pull H +1 off H 2 O

3. Tro - Chapter 143 Properties of Acids Sour taste react with “active” metals i.e. Al, Zn, Fe, but not Cu, Ag or Au 2 Al + 6 HCl  AlCl 3 + 3 H 2 corrosive react with carbonates, producing CO 2 marble, baking soda, chalk, limestone CaCO 3 + 2 HCl  CaCl 2 + CO 2 + H 2 O change color of vegetable dyes blue litmus turns red react with bases to form ionic salts

4. Tro - Chapter 144 Acid Reactions Acids React with Metals acids react with many metals but not all!! when acids react with metals, they produce a salt and hydrogen gas 3 H 2 SO 4 (aq) + 2 Al(s) → Al 2 (SO 4 ) 3 (aq) + 3 H 2 (g)

5. Tro - Chapter 145 Common Acids

6. Tro - Chapter 146 Structures of Acids binary acids have acid hydrogens attached to a nonmetal atom HCl, HF Hydrofluoric acid

7. Tro - Chapter 147 Structure of Acids oxy acids have acid hydrogens attached to an oxygen atom H 2 SO 4, HNO 3

8. Tro - Chapter 148 Structure of Acids carboxylic acids have COOH group HC 2 H 3 O 2, H 3 C 6 H 5 O 3 only the first H in the formula is acidic the H is on the COOH

9. Tro - Chapter 149 Properties of Bases also known as alkalis taste bitter alkaloids = plant product that is alkaline  often poisonous solutions feel slippery change color of vegetable dyes different color than acid red litmus turns blue react with acids to form ionic salts neutralization

10. Tro - Chapter 1410 Base Reactions the reaction all bases have is common is neutralization of acids Ex: acid base volcano

11. Tro - Chapter 1411 Common Bases

12. Tro - Chapter 1412 Structure of Bases most ionic bases contain OH ions NaOH, Ca(OH) 2 some contain CO 3 2- ions CaCO 3 NaHCO 3 molecular bases contain structures that react with H + mostly amine groups

13. Tro - Chapter 1413 Arrhenius Theory bases dissociate in water to produce OH - ions and cations ionic substances dissociate in water NaOH(aq) → Na + (aq) + OH – (aq) acids ionize in water to produce H + ions and anions because molecular acids are not made of ions, they cannot dissociate they must be pulled apart, or ionized, by the water HCl(aq) → H + (aq) + Cl – (aq) in formula, ionizable H written in front HC 2 H 3 O 2 (aq) → H + (aq) + C 2 H 3 O 2 – (aq)

14. Tro - Chapter 1414 Arrhenius Theory HCl ionizes in water, producing H + and Cl – ions NaOH dissociates in water, producing Na + and OH – ions

15. Tro - Chapter 1415 Arrhenius Acid-Base Reactions the H + from the acid combines with the OH - from the base to make a molecule of H 2 O it is often helpful to think of H 2 O as H-OH the cation from the base combines with the anion from the acid to make a salt acid + base → salt + water HCl(aq) + NaOH(aq) → NaCl(aq) + H 2 O(l)

16. Tro - Chapter 1416 Problems with Arrhenius Theory does not explain why molecular substances, like NH 3, dissolve in water to form basic solutions – even though they do not contain OH – ions does not explain acid-base reactions that do not take place in aqueous solution H + ions do not exist in water. Acid solutions contain H 3 O + ions H + = a proton! H 3 O + = hydronium ions

17. Tro - Chapter 1417 Arrow Conventions chemists commonly use two kinds of arrows in reactions to indicate the degree of completion of the reactions a single arrow indicates all the reactant molecules are converted to product molecules at the end a double arrow indicates the reaction stops when only some of the reactant molecules have been converted into products  in these notes

18. Tro - Chapter 1418 Brønsted-Lowery Theory in a Brønsted-Lowery Acid-Base reaction, an H + is transferred does not have to take place in aqueous solution broader definition than Arrhenius acid is H donor, base is H acceptor base structure must contain an atom with an unshared pair of electrons in the reaction, the acid molecule gives an H + to the base molecule H–A + :B  :A – + H–B +

19. Tro - Chapter 1419 Amphoteric Substances amphoteric substances can act as either an acid or a base have both transferable H and atom with lone pair HCl(aq) is acidic because HCl transfers an H + to H 2 O, forming H 3 O + ions water acts as base, accepting H + HCl(aq) + H 2 O(l) → Cl – (aq) + H 3 O + (aq) NH 3 (aq) is basic because NH 3 accepts an H + from H 2 O, forming OH – (aq) water acts as acid, donating H + NH 3 (aq) + H 2 O(l)  NH 4 + (aq) + OH – (aq)

20. Tro - Chapter 1420 Brønsted-Lowery Acid-Base Reactions one of the advantages of Brønsted-Lowery theory is that it allows reactions to be reversible H–A + :B → :A – + H–B + the original base has an extra H + after the reaction – so it could act as an acid in the reverse process and the original acid has a lone pair of electrons after the reaction – so it could act as a base in the reverse process :A – + H–B + → H–A + :B a double arrow, , is usually used to indicate a process that is reversible

21. Tro - Chapter 1421 Conjugate Pairs In a Brønsted-Lowery Acid-Base reaction, the original base becomes an acid in the reverse reaction, and the original acid becomes a base in the reverse process each reactant and the product it becomes is called a conjugate pair the original base becomes the conjugate acid; and the original acid becomes the conjugate base

22. Tro - Chapter 1422 Brønsted-Lowery Acid-Base Reactions H–A + :B  :A – +H–B + acidbaseconjugate conjugate base acid HCHO 2 + H 2 O  CHO 2 – +H 3 O + acid baseconjugate conjugate base acid H 2 O + NH 3  HO – +NH 4 + acid baseconjugate conjugate base acid

23. Tro - Chapter 1423 Conjugate Pairs In the reaction H 2 O + NH 3  HO – + NH 4 + H 2 O and HO – constitute an Acid/Conjugate Base pair NH 3 and NH 4 + constitute a Base/Conjugate Acid pair

24. Tro - Chapter 1424 Practice – Identify the Brønsted-Lowery Acids and Bases and their Conjugates in each Reaction H 2 SO 4 + H 2 O  HSO 4 – +H 3 O + HCO 3 – + H 2 O  H 2 CO 3 +HO –

25. Tro - Chapter 1425 Practice – Identify the Brønsted-Lowery Acids and Bases and their Conjugates in each Reaction H 2 SO 4 + H 2 O  HSO 4 – +H 3 O + acid baseconjugate conjugate base acid HCO 3 – + H 2 O  H 2 CO 3 +HO – base acidconjugate conjugate acid base

26. Tro - Chapter 1426 Neutralization Reactions H + + OH -  H 2 O acid + base  salt + water double displacement reactions salt = cation from base + anion from acid cation and anion charges stay constant H 2 SO 4 + Ca(OH) 2 → CaSO 4 + 2 H 2 O some neutralization reactions are gas evolving where H 2 CO 3 decomposes into CO 2 and H 2 O H 2 SO 4 + 2 NaHCO 3 → Na 2 SO 4 + 2 H 2 O + 2 CO 2

27. Acid Strength and Equilibrium Tro - Chapter 1427

28. Tro - Chapter 1428 Strong or Weak a strong acid is a strong electrolyte practically all the acid molecules ionize, → a strong base is a strong electrolyte practically all the base molecules form OH– ions, either through dissociation or reaction with water, → a weak acid is a weak electrolyte only a small percentage of the molecules ionize,  a weak base is a weak electrolyte only a small percentage of the base molecules form OH– ions, either through dissociation or reaction with water, 

29. Tro - Chapter 1429 Strong Acids The stronger the acid, the more willing it is to donate H use water as the standard base strong acids donate practically all their H’s 100% ionized in water strong electrolyte [H 3 O + ] = [strong acid] [ ] = molarity HCl  H + + Cl - HCl + H 2 O  H 3 O + + Cl -

30. Tro - Chapter 1430 Strong Acids Pure WaterHCl solution

31. Tro - Chapter 1431 Weak Acids weak acids donate a small fraction of their H’s most of the weak acid molecules do not donate H to water much less than 1% ionized in water [H 3 O + ] << [weak acid] HF  H + + F - HF + H 2 O  H 3 O + + F -

32. Tro - Chapter 1432 Weak Acids Pure WaterHF solution

33. Tro - Chapter 1433 Strong Bases The stronger the base, the more willing it is to accept H use water as the standard acid strong bases, practically all molecules are dissociated into OH – or accept H’s strong electrolyte multi-OH bases completely dissociated [HO – ] = [strong base] x (# OH) NaOH  Na + + OH -

34. Tro - Chapter 1434 Weak Bases in weak bases, only a small fraction of molecules accept H’s weak electrolyte most of the weak base molecules do not take H from water much less than 1% ionization in water [HO – ] << [strong base] NH 3 + H 2 O  NH 4 + + OH -

35. Tro - Chapter 1435 Relationship between Strengths of Acids and their Conjugate Bases the stronger an acid is, the weaker the attraction of the ionizable H for the rest of the molecule is the better the acid is at donating H, the worse its conjugate base will be at accepting a H strong acid HCl + H 2 O → Cl – + H 3 O + weak conj. base weak acid HF + H 2 O  F – + H 3 O + strong conj. base

36. Tro - Chapter 1436 Autoionization of Water Water is actually an extremely weak electrolyte therefore there must be a few ions present about 1 out of every 10 million water molecules form ions through a process called autoionization H 2 O  H + + OH – H 2 O + H 2 O  H 3 O + + OH – all aqueous solutions contain both H + and OH – the concentration of H + and OH – are equal in water [H + ] = [OH – ] = 10 -7 M @ 25°C

37. Tro - Chapter 1437 Ion Product of Water the product of the H + and OH – concentrations is always the same number the number is called the ion product of water and has the symbol K w [H + ] x [OH – ] = 1 x 10 -14 = K w as [H + ] increases the [OH – ] must decrease so the product stays constant inversely proportional

38. Tro - Chapter 1438 Acidic and Basic Solutions neutral solutions have equal [H + ] and [OH – ] [H + ] = [OH – ] = 1 x 10 -7 acidic solutions have a larger [H + ] than [OH – ] [H + ] > 1 x 10 -7 ; [OH – ] < 1 x 10 -7 basic solutions have a larger [OH – ] than [H + ] [H + ] 1 x 10 -7

39. Tro - Chapter 1439 Example - Determine the [H +1 ] for a 0.00020 M Ba(OH) 2 and determine whether the solution is acidic, basic or neutral Ba(OH) 2 = Ba 2+ + 2 OH – therefore [OH – ] = 2 x 0.00020 = 0.00040 = 4.0 x 10 -4 M [H + ] = 2.5 x 10 -11 M

40. Tro - Chapter 1440 Practice - Determine the [H +1 ] concentration and whether the solution is acidic, basic or neutral for the following [OH – ] = 0.000250 M [OH – ] = 3.50 x 10 -8 M Ca(OH) 2 = 0.20 M

41. Tro - Chapter 1441 Complete the Table [H + ] vs. [OH - ] OH - H+H+ H+H+ H+H+ H+H+ H+H+ [OH - ] [H + ] 10 0 10 -1 10 -3 10 -5 10 -7 10 -9 10 -11 10 -13 10 -14

42. Tro - Chapter 1442 Complete the Table [H + ] vs. [OH - ] OH - H+H+ H+H+ H+H+ H+H+ H+H+ [OH - ]10 -14 10 -13 10 -11 10 -9 10 -7 10 -5 10 -3 10 -1 10 0 [H + ] 10 0 10 -1 10 -3 10 -5 10 -7 10 -9 10 -11 10 -13 10 -14 even though it may look like it, neither H + of OH - will ever be 0 the sizes of the H + and OH - are not to scale because the divisions are powers of 10 rather than units Acid Base

43. pH

44. Tro - Chapter 1444 pH the acidity/basicity of a solution is often expressed as pH pH = -log[H + ], [H + ] = 10 -pH exponent on 10 with a positive sign pH water = -log[10 -7 ] = 7 need to know the [H + ] concentration to find pH pH 7 is basic, pH = 7 is neutral

45. Tro - Chapter 1445 pH the lower the pH, the more acidic the solution; the higher the pH, the more basic the solution 1 pH unit corresponds to a factor of 10 difference in acidity normal range 0 to 14 pH 0 is [H + ] = 1 M, pH 14 is [OH – ] = 1 M pH can be negative (very acidic) or larger than 14 (very alkaline)

46. Tro - Chapter 1446 pH of Common Substances SubstancepH 1.0 M HCl0.0 0.1 M HCl1.0 stomach acid1.0 to 3.0 lemons2.2 to 2.4 soft drinks2.0 to 4.0 plums2.8 to 3.0 apples2.9 to 3.3 cherries3.2 to 4.0 unpolluted rainwater5.6 human blood7.3 to 7.4 egg whites7.6 to 8.0 milk of magnesia (sat’d Mg(OH) 2 )10.5 household ammonia10.5 to 11.5 1.0 M NaOH14

47. Tro - Chapter 1447 Example - Calculate the pH of a 0.0010 M Ba(OH) 2 solution & determine if is acidic, basic or neutral [H + ] = 1 x 10 -14 2.0 x 10 -3 = 5.0 x 10 -12 M pH > 7 therefore basic Ba(OH) 2 = Ba 2+ + 2 OH - therefore [OH - ] = 2 x 0.0010 = 0.0020 = 2.0 x 10 -3 M pH = -log [H + ] = -log (5.0 x 10 -12 ) pH = 11.3

48. Tro - Chapter 1448 Practice - Calculate the pH of the following strong acid or base solutions 0.0020 M HCl 0.0050 M Ca(OH) 2 0.25 M HNO 3

49. Tro - Chapter 1449 OH - H+H+ H+H+ H+H+ H+H+ H+H+ [OH - ]10 -14 10 -13 10 -11 10 -9 10 -7 10 -5 10 -3 10 -1 10 0 [H + ] 10 0 10 -1 10 -3 10 -5 10 -7 10 -9 10 -11 10 -13 10 -14 pH Complete the Table pH

50. Tro - Chapter 1450 Complete the Table pH OH - H+H+ H+H+ H+H+ H+H+ H+H+ [OH - ]10 -14 10 -13 10 -11 10 -9 10 -7 10 -5 10 -3 10 -1 10 0 [H + ] 10 0 10 -1 10 -3 10 -5 10 -7 10 -9 10 -11 10 -13 10 -14 pH 0 1 3 5 7 9 11 13 14 Acid Base

51. Tro - Chapter 1451 Sample - Calculate the concentration of [H + ] for a solution with pH 3.7 [H + ] = 10 -pH [H + ] = 10 -3.7 means 0.0001 < [H +1 ] < 0.001 [H + ] = 2 x 10 -4 M = 0.0002 M

52. Tro - Chapter 1452 Practice - Determine the [H + ] for each of the following pH = 2.7 pH = 12 pH = 0.60

53. Titration Tro - Chapter 1453

54. Tro - Chapter 1454 Titration using reaction stoichiometry to determine the concentration of an unknown solution Titrant (known solution) added from a buret indicators are chemicals added to help determine when a reaction is complete the endpoint of the titration occurs when the reaction is complete

55. Tro - Chapter 1455 Titration

56. Tro - Chapter 1456 Titration The base solution is the titrant in the buret. As the base is added to the acid, the H + reacts with the OH – to form water. But there is still excess acid present so the color does not change. At the titration’s endpoint, just enough base has been added to neutralize all the acid. At this point the indicator changes color.

57. Tro - Chapter 1457 Example: The titration of 10.00 mL of HCl solution of unknown concentration requires 12.54 mL of 0.100 M NaOH solution to reach the end point. What is the concentration of the unknown HCl solution? Answer =.125 M HCl this is a dilute solution of strong acid A concentrated solution would be greater than 1M Example: Acid-Base Titration Eq: M A x V A = M B x V B

58. Acids That Give More Than One H+ Polyprotic Acids Diprotic Acids – donate two protons H 2 SO 4 Twice as much base will be needed if the base is NaOH. An equal amount of base will be needed if the base has two OH-’s, such as Ba(OH) 2 Triprotic Acids – donate three protons H 3 PO 4 Three times as much base will be needed if the base is NaOH. An equal amount of base will be needed if the base has three OH-’s, such as Al(OH) 3 Tro - Chapter 1458

59. Buffers Tro - Chapter 1459

60. Tro - Chapter 1460 Buffers buffers are solutions that resist changing pH when small amounts of acid or base are added they resist changing pH by neutralizing added acid or base buffers are made by mixing together a weak acid and its conjugate base or weak base and it conjugate acid

61. Tro - Chapter 1461 How Buffers Work the weak acid present in the buffer mixture can neutralize added base the conjugate base present in the buffer mixture can neutralize added acid the net result is little to no change in the solution pH

62. Blood is a Buffer The kidneys and the lungs work together to help maintain a blood pH of 7.4 (range between 7.35 and 7.45)by affecting the components of the buffers in the blood. Therefore, to understand how these organs help control the pH of the blood, we must first discuss how buffers work in solution. Acid-base buffers provide resistance to a change in the pH of a solution when hydrogen ions (protons) or hydroxide ions are added or removed. An acid-base buffer typically consists of a weak acid, and its conjugate base Tro - Chapter 1462

63. Buffers work because the concentrations of the weak acid and its salt are large compared to the amount of protons or hydroxide ions added or removed When protons are added to the solution from an external source, some of the base component of the buffer is converted to the weak-acid component Tro - Chapter 1463

64. When hydroxide ions are added to the solution (or, equivalently, protons are removed from the solution; protons are dissociated from some of the weak-acid molecules of the buffer, converting them to the base of the buffer. 2 H 2 O + CO 2  H 2 CO 3 + H 2 O  H 3 O + + HCO 3 - The CO 2 can be removed by the lungs However, the change in acid and base concentrations is small relative to the amounts of these species present in solution. Hence, the ratio of acid to base changes only slightly. Thus, the effect on the pH of the solution is small, within certain limitations on the amount of H + or OH - added or removed. Tro - Chapter 1464

65. Tro - Chapter 1465 Acetic Acid/Acetate Buffer

66. Acid Rain Tro - Chapter 1466

67. Tro - Chapter 1467 What is Acid Rain? natural rain water has a pH of 5.6 naturally slightly acidic due mainly to CO 2 rain water with a pH lower than 5.6 is called acid rain acid rain is linked to damage in ecosystems and structures

68. Tro - Chapter 1468 What Causes Acid Rain? many natural and pollutant gases dissolved in the air are nonmetal oxides CO 2, SO 2, NO 2 nonmetal oxides are acidic CO 2 (g) + H 2 O(l) → H 2 CO 3 (aq) 2 SO 2 (g) + O 2 (g) + 2 H 2 O(l) → 2 H 2 SO 4 (aq) 4 NO 2 (g) + O 2 (g) + 2 H 2 O(l) → 4 HNO 3 (aq)

69. Tro - Chapter 1469 Nonmetal Oxides are Acidic processes that produce nonmetal oxide gases as waste increase the acidity of the rain natural – volcanoes and some bacterial action man-made – combustion of fuel weather patterns may cause rain to be acidic in regions other than where the nonmetal oxide is produced

70. Tro - Chapter 1470 pH of Rain in Different Regions

71. Tro - Chapter 1471 Sources of SO 2 from Utilities

72. Tro - Chapter 1472 Damage from Acid Rain acids react with metals, and materials that contain carbonates acid rain damages bridges, cars and other metallic structures acid rain damages buildings and other structures made of limestone or cement

73. Tro - Chapter 1473 Damage from Acid Rain circa 1935 circa 1995