1. Write equations that show NH 3 as both a conjugate acid and a conjugate base. 2. Write equations that show acting both as an acid and as a base. 3. Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry acid: (f) 4. Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry acid: (f) HS − 5. Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry base: (f) 6. Show by suitable net ionic equations that each of the following species can act as a Brønsted-Lowry base: (f) 7. What is the conjugate acid of each of the following? What is the conjugate base of each? OH − H 2 O NH 3 H 2 O 2 HS − 8. What is the conjugate acid of each of the following? What is the conjugate base of each? H 2 S PH 3 HS − H 4 N 2 CH 3 OH 9. Identify and label the Brønsted-Lowry acid, its conjugate base, the Brønsted-Lowry base, and its conjugate acid in each of the following equations: (f) 10. Identify and label the Brønsted-Lowry acid, its conjugate base, the Brønsted-Lowry base, and its conjugate acid in each of the following equations: (f) 11. What are amphiprotic species? Illustrate with suitable equations. 12. State which of the following species are amphiprotic and write chemical equations illustrating the amphiprotic character of these species: H 2 O S 2− 13. State which of the following species are amphiprotic and write chemical equations illustrating the amphiprotic character of these species. NH 3 Br − 14. Is the self-ionization of water endothermic or exothermic? The ionization constant for water ( K w ) is 2.9 \times 10^{−14} at 40 °C and 9.3 \times 10^{−14} at 60 °C. 15. Explain why a sample of pure water at 40 °C is neutral even though [H 3 O + ] = 1.7 \times 10^{−7} M . K w is 2.9 \times 10^{−14} at 40 °C. 16. The ionization constant for water ( K w ) is 2.9 \times 10^{−14} at 40 °C. Calculate [H 3 O + ], [OH − ], pH, and pOH for pure water at 40 °C. 17. The ionization constant for water ( K w ) is 9.311 \times 10^{−14} at 60 °C. Calculate [H 3 O + ], [OH − ], pH, and pOH for pure water at 60 °C. 18. Calculate the pH and the pOH of each of the following solutions at 25 °C for which the substances ionize completely: 0.200 M HCl 0.0143 M NaOH 3.0 M HNO 3 0.0031 M Ca(OH) 2 19. Calculate the pH and the pOH of each of the following solutions at 25 °C for which the substances ionize completely: 0.000259 M HClO 4 0.21 M NaOH 0.000071 M Ba(OH) 2 2.5 M KOH 20. What are the pH and pOH of a solution of 2.0 M HCl, which ionizes completely? 21. What are the hydronium and hydroxide ion concentrations in a solution whose pH is 6.52? 22. Calculate the hydrogen ion concentration and the hydroxide ion concentration in wine from its pH. See Figure 14.2 for useful information. 23. Calculate the hydronium ion concentration and the hydroxide ion concentration in lime juice from its pH. See Figure 14.2 for useful information. 24. The hydronium ion concentration in a sample of rainwater is found to be 1.7 \times 10^{−6} M at 25 °C. What is the concentration of hydroxide ions in the rainwater? 25. The hydroxide ion concentration in household ammonia is 3.2 \times 10^{−3} M at 25 °C. What is the concentration of hydronium ions in the solution? 26. Explain why the neutralization reaction of a strong acid and a weak base gives a weakly acidic solution. 27. Explain why the neutralization reaction of a weak acid and a strong base gives a weakly basic solution. 28. Use this list of important industrial compounds (and Figure 14.8) to answer the following questions regarding: Ca(OH) 2 , CH 3 CO 2 H, HCl, H 2 CO 3 , HF, HNO 2 , HNO 3 , H 3 PO 4 , H 2 SO 4 , NH 3 , NaOH, Na 2 CO 3 . Identify the strong Brønsted-Lowry acids and strong Brønsted-Lowry bases. Identify the compounds that can behave as Brønsted-Lowry acids with strengths lying between those of H 3 O + and H 2 O. Identify the compounds that can behave as Brønsted-Lowry bases with strengths lying between those of H 2 O and OH − . 29. The odor of vinegar is due to the presence of acetic acid, CH 3 CO 2 H, a weak acid. List, in order of descending concentration, all of the ionic and molecular species present in a 1- M aqueous solution of this acid. 30. Household ammonia is a solution of the weak base NH 3 in water. List, in order of descending concentration, all of the ionic and molecular species present in a 1- M aqueous solution of this base. 31. Explain why the ionization constant, K a , for H 2 SO 4 is larger than the ionization constant for H 2 SO 3 . 32. Explain why the ionization constant, K a , for HI is larger than the ionization constant for HF. 33. Gastric juice, the digestive fluid produced in the stomach, contains hydrochloric acid, HCl. Milk of Magnesia, a suspension of solid Mg(OH) 2 in an aqueous medium, is sometimes used to neutralize excess stomach acid. Write a complete balanced equation for the neutralization reaction, and identify the conjugate acid-base pairs. 34. Nitric acid reacts with insoluble copper(II) oxide to form soluble copper(II) nitrate, Cu(NO 3 ) 2 , a compound that has been used to prevent the growth of algae in swimming pools. Write the balanced chemical equation for the reaction of an aqueous solution of HNO 3 with CuO. 35. What is the ionization constant at 25 °C for the weak acid the conjugate acid of the weak base CH 3 NH 2 , K b = 4.4 \times 10^{−4}. 36. What is the ionization constant at 25 °C for the weak acid the conjugate acid of the weak base (CH 3 ) 2 NH, K b = 5.9 \times 10^{−4}? 37. Which base, CH 3 NH 2 or (CH 3 ) 2 NH, is the stronger base? Which conjugate acid, or , is the stronger acid? 38. Which is the stronger acid, or HBrO? 39. Which is the stronger base, (CH 3 ) 3 N or 40. Predict which acid in each of the following pairs is the stronger and explain your reasoning for each. H 2 O or HF B(OH) 3 or Al(OH) 3 or NH 3 or H 2 S H 2 O or H 2 Te 41. Predict which compound in each of the following pairs of compounds is more acidic and explain your reasoning for each. or NH 3 or H 2 O PH 3 or HI NH 3 or PH 3 H 2 S or HBr 42. Rank the compounds in each of the following groups in order of increasing acidity or basicity, as indicated, and explain the order you assign. acidity: HCl, HBr, HI basicity: H 2 O, OH − , H − , Cl − basicity: Mg(OH) 2 , Si(OH) 4 , ClO 3 (OH) (Hint: Formula could also be written as HClO 4 .) acidity: HF, H 2 O, NH 3 , CH 4 43. Rank the compounds in each of the following groups in order of increasing acidity or basicity, as indicated, and explain the order you assign. (f) basicity: BrO − , 44. Both HF and HCN ionize in water to a limited extent. Which of the conjugate bases, F − or CN − , is the stronger base? 45. The active ingredient formed by aspirin in the body is salicylic acid, C 6 H 4 OH(CO 2 H). The carboxyl group (−CO 2 H) acts as a weak acid. The phenol group (an OH group bonded to an aromatic ring) also acts as an acid but a much weaker acid. List, in order of descending concentration, all of the ionic and molecular species present in a 0.001- M aqueous solution of C 6 H 4 OH(CO 2 H). 46. Are the concentrations of hydronium ion and hydroxide ion in a solution of an acid or a base in water directly proportional or inversely proportional? Explain your answer. 47. What two common assumptions can simplify calculation of equilibrium concentrations in a solution of a weak acid or base? 48. Which of the following will increase the percent of NH 3 that is converted to the ammonium ion in water? addition of NaOH addition of HCl addition of NH 4 Cl 49. Which of the following will increase the percentage of HF that is converted to the fluoride ion in water? addition of NaOH addition of HCl addition of NaF 50. What is the effect on the concentrations of HNO 2 , and OH − when the following are added to a solution of KNO 2 in water: HCl HNO 2 NaOH NaCl KNO 51. What is the effect on the concentration of hydrofluoric acid, hydronium ion, and fluoride ion when the following are added to separate solutions of hydrofluoric acid? HCl KF NaCl KOH HF 52. Why is the hydronium ion concentration in a solution that is 0.10 M in HCl and 0.10 M in HCOOH determined by the concentration of HCl? 53. From the equilibrium concentrations given, calculate K a for each of the weak acids and K b for each of the weak bases. (a) CH 3 CO 2 H: = 1.34 \times 10^{−3} M ; = 1.34 \times 10^{−3} M ; [CH 3 CO 2 H] = 9.866 \times 10^{−2} M ; (b) ClO − : [OH − ] = 4.0 \times 10^{−4} M ; [HClO] = 2.38 \times 10^{−4} M ; [ClO − ] = 0.273 M ; (c) HCO 2 H: [HCO 2 H] = 0.524 M ; = 9.8 \times 10^{−3} M ; = 9.8 \times 10^{−3} M ; (d) = 0.233 M ; [C 6 H 5 NH 2 ] = 2.3 \times 10^{−3} M ; = 2.3 \times 10^{−3} M 54. From the equilibrium concentrations given, calculate K a for each of the weak acids and K b for each of the weak bases. (a) NH 3 : [OH − ] = 3.1 \times 10^{−3} M ; = 3.1 \times 10^{−3} M ; [NH 3 ] = 0.533 M ; (b) HNO 2 : = 0.011 M ; = 0.0438 M ; [HNO 2 ] = 1.07 M ; (c) (CH 3 ) 3 N: [(CH 3 ) 3 N] = 0.25 M ; [(CH 3 ) 3 NH + ] = 4.3 \times 10^{−3} M ; [OH − ] = 3.7 \times 10^{−3} M ; (d) = 0.100 M ; [NH 3 ] = 7.5 \times 10^{−6} M ; [H 3 O + ] = 7.5 \times 10^{−6} M 55. Determine K b for the nitrite ion, In a 0.10- M solution this base is 0.0015% ionized. 56. Determine K a for hydrogen sulfate ion, In a 0.10- M solution the acid is 29% ionized. 57. Calculate the ionization constant for each of the following acids or bases from the ionization constant of its conjugate base or conjugate acid: (f) (as a base) 58. Calculate the ionization constant for each of the following acids or bases from the ionization constant of its conjugate base or conjugate acid: (f) (as a base) 59. Using the K a value of 1.4 60. Calculate the concentration of all solute species in each of the following solutions of acids or bases. Assume that the ionization of water can be neglected, and show that the change in the initial concentrations can be neglected. 0.0092 M HClO, a weak acid 0.0784 M C 6 H 5 NH 2 , a weak base 0.0810 M HCN, a weak acid 0.11 M (CH 3 ) 3 N, a weak base 0.120 M a weak acid, K a = 1.6 \times 10^{−7} 61. Propionic acid, C 2 H 5 CO 2 H ( K a = 1.34 \times 10^{−5}), is used in the manufacture of calcium propionate, a food preservative. What is the pH of a 0.698- M solution of C 2 H 5 CO 2 H? 62. White vinegar is a 5.0% by mass solution of acetic acid in water. If the density of white vinegar is 1.007 g/cm 3 , what is the pH? 63. The ionization constant of lactic acid, CH 3 CH(OH)CO 2 H, an acid found in the blood after strenuous exercise, is 1.36 \times 10^{−4}. If 20.0 g of lactic acid is used to make a solution with a volume of 1.00 L, what is the concentration of hydronium ion in the solution? 64. Nicotine, C 10 H 14 N 2 , is a base that will accept two protons ( K b1 = 7 \times 10^{−7}, K b2 = 1.4 \times 10^{−11}). What is the concentration of each species present in a 0.050- M solution of nicotine? 65. The pH of a 0.23- M solution of HF is 1.92. Determine K a for HF from these data. 66. The pH of a 0.15- M solution of is 1.43. Determine K a for from these data. 67. The pH of a 0.10- M solution of caffeine is 11.70. Determine K b for caffeine from these data: 68. The pH of a solution of household ammonia, a 0.950 M solution of NH 3, is 11.612. Determine K b for NH 3 from these data. 69. Determine whether aqueous solutions of the following salts are acidic, basic, or neutral: Al(NO 3 ) 3 RbI KHCO 2 CH 3 NH 3 Br 70. Determine whether aqueous solutions of the following salts are acidic, basic, or neutral: FeCl 3 K 2 CO 3 NH 4 Br KClO 4 71. Novocaine, C 13 H 21 O 2 N 2 Cl, is the salt of the base procaine and hydrochloric acid. The ionization constant for procaine is 7 \times 10^{−6}. Is a solution of novocaine acidic or basic? What are [H 3 O + ], [OH − ], and pH of a 2.0% solution by mass of novocaine, assuming that the density of the solution is 1.0 g/mL. 72. Which of the following concentrations would be practically equal in a calculation of the equilibrium concentrations in a 0.134- M solution of H 2 CO 3 , a diprotic acid: [OH − ], [H 2 CO 3 ], No calculations are needed to answer this question. 73. Calculate the concentration of each species present in a 0.050- M solution of H 2 S. 74. Calculate the concentration of each species present in a 0.010- M solution of phthalic acid, C 6 H 4 (CO 2 H) 2 . 75. Salicylic acid, HOC 6 H 4 CO 2 H, and its derivatives have been used as pain relievers for a long time. Salicylic acid occurs in small amounts in the leaves, bark, and roots of some vegetation (most notably historically in the bark of the willow tree). Extracts of these plants have been used as medications for centuries. The acid was first isolated in the laboratory in 1838. Both functional groups of salicylic acid ionize in water, with K a = 1.0 \times 10^{−3} for the—CO 2 H group and 4.2 \times 10^{−13} for the −OH group. What is the pH of a saturated solution of the acid (solubility = 1.8 g/L). Aspirin was discovered as a result of efforts to produce a derivative of salicylic acid that would not be irritating to the stomach lining. Aspirin is acetylsalicylic acid, CH 3 CO 2 C 6 H 4 CO 2 H. The −CO 2 H functional group is still present, but its acidity is reduced, K a = 3.0 \times 10^{−4}. What is the pH of a solution of aspirin with the same concentration as a saturated solution of salicylic acid (See Part a). 76. The ion HTe − is an amphiprotic species; it can act as either an acid or a base. What is K a for the acid reaction of HTe − with H 2 O? What is K b for the reaction in which HTe − functions as a base in water? Demonstrate whether or not the second ionization of H 2 Te can be neglected in the calculation of [HTe − ] in a 0.10 M solution of H 2 Te. 77. Explain why a buffer can be prepared from a mixture of NH 4 Cl and NaOH but not from NH 3 and NaOH. 78. Explain why the pH does not change significantly when a small amount of an acid or a base is added to a solution that contains equal amounts of the acid H 3 PO 4 and a salt of its conjugate base NaH 2 PO 4 . 79. Explain why the pH does not change significantly when a small amount of an acid or a base is added to a solution that contains equal amounts of the base NH 3 and a salt of its conjugate acid NH 4 Cl. 80. What is [H 3 O + ] in a solution of 0.25 M CH 3 CO 2 H and 0.030 M NaCH 3 CO 2 ? 81. What is [H 3 O + ] in a solution of 0.075 M HNO 2 and 0.030 M NaNO 2 ? 82. What is [OH − ] in a solution of 0.125 M CH 3 NH 2 and 0.130 M CH 3 NH 3 Cl? 83. What is [OH − ] in a solution of 1.25 M NH 3 and 0.78 M NH 4 NO 3 ? 84. What is the effect on the concentration of acetic acid, hydronium ion, and acetate ion when the following are added to an acidic buffer solution of equal concentrations of acetic acid and sodium acetate: HCl KCH 3 CO 2 NaCl KOH CH 3 CO 2 H 85. What is the effect on the concentration of ammonia, hydroxide ion, and ammonium ion when the following are added to a basic buffer solution of equal concentrations of ammonia and ammonium nitrate: KI NH 3 HI NaOH NH 4 Cl 86. What will be the pH of a buffer solution prepared from 0.20 mol NH 3 , 0.40 mol NH 4 NO 3 , and just enough water to give 1.00 L of solution? 87. Calculate the pH of a buffer solution prepared from 0.155 mol of phosphoric acid, 0.250 mole of KH 2 PO 4 , and enough water to make 0.500 L of solution. 88. How much solid NaCH 3 CO 2 •3H 2 O must be added to 0.300 L of a 0.50- M acetic acid solution to give a buffer with a pH of 5.00? (Hint: Assume a negligible change in volume as the solid is added.) 89. What mass of NH 4 Cl must be added to 0.750 L of a 0.100- M solution of NH 3 to give a buffer solution with a pH of 9.26? (Hint: Assume a negligible change in volume as the solid is added.) 90. A buffer solution is prepared from equal volumes of 0.200 M acetic acid and 0.600 M sodium acetate. Use 1.80 \times 10^{−5} as K a for acetic acid. What is the pH of the solution? Is the solution acidic or basic? What is the pH of a solution that results when 3.00 mL of 0.034 M HCl is added to 0.200 L of the original buffer? 91. A 5.36–g sample of NH 4 Cl was added to 25.0 mL of 1.00 M NaOH and the resulting solution diluted to 0.100 L. What is the pH of this buffer solution? Is the solution acidic or basic? What is the pH of a solution that results when 3.00 mL of 0.034 M HCl is added to the solution? 92. Explain how to choose the appropriate acid-base indicator for the titration of a weak base with a strong acid. 93. Explain why an acid-base indicator changes color over a range of pH values rather than at a specific pH. 94. Calculate the pH at the following points in a titration of 40 mL (0.040 L) of 0.100 M barbituric acid ( K a = 9.8 \times 10^{−5}) with 0.100 M KOH. no KOH added 20 mL of KOH solution added 39 mL of KOH solution added 40 mL of KOH solution added 41 mL of KOH solution added 95. The indicator dinitrophenol is an acid with a K a of 1.1 \times 10^{−4}. In a 1.0 \times 10^{−4}- M solution, it is colorless in acid and yellow in base. Calculate the pH range over which it goes from 10% ionized (colorless) to 90% ionized (yellow).