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Chemistry | 12th Edition | ISBN: 9780078021510 | Authors: Raymond Chang; Kenneth Goldsby

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Chapter 16 Problems

Chapter: 16 Problem: 16.1

Chemistry 12
Use Le Chteliers principle to explain how the common ion effect affects the pH of a solution.

Chapter: 16 Problem: 16.2

Chemistry 12
Describe the effect on pH (increase, decrease, or no change) that results from each of the following additions: (a) potassium acetate to an acetic acid solution; (b) ammonium nitrate to an ammonia solution; (c) sodium formate (HCOONa) to a formic acid (HCOOH) solution; (d) potassium chloride to a hyd

Chapter: 16 Problem: 16.3

Chemistry 12
Define pKa for a weak acid. What is the relationship between the value of the pKa and the strength of the acid? Do the same for a weak base.

Chapter: 16 Problem: 16.4

Chemistry 12
The pKas of two monoprotic acids HA and HB are 5.9 and 8.1, respectively. Which of the two is the stronger acid?

Chapter: 16 Problem: 16.5

Chemistry 12
Determine the pH of (a) a 0.40 M CH3COOH solution, (b) a solution that is 0.40 M CH3COOH and 0.20 M CH3COONa.

Chapter: 16 Problem: 16.6

Chemistry 12
Determine the pH of (a) a 0.20 M NH3 solution, (b) a solution that is 0.20 M in NH3 and 0.30 M NH4Cl.

Chapter: 16 Problem: 16.7

Chemistry 12
What is a buffer solution? What constitutes a buffer solution?

Chapter: 16 Problem: 16.8

Chemistry 12
Which of the following has the greatest buffer capacity? (a) 0.40 M CH3COONa/0.20 M CH3COOH, (b) 0.40 M CH3COONa/0.60 M CH3COOH, and (c) 0.30 M CH3COONa/0.60 M CH3COOH.

Chapter: 16 Problem: 16.9

Chemistry 12
Which of the following solutions can act as a buffer? (a) KCl/HCl, (b) KHSO4/H2SO4, (c) Na2HPO4/ NaH2PO4, (d) KNO2/HNO2.

Chapter: 16 Problem: 16.1b

Chemistry 12
Which of the following solutions can act as a buffer? (a) KCN/HCN, (b) Na2SO4/NaHSO4, (c) NH3/ NH4NO3, (d) NaI/HI.

Chapter: 16 Problem: 16.11

Chemistry 12
Calculate the pH of the buffer system made up of 0.15 M NH3/0.35 M NH4Cl.

Chapter: 16 Problem: 16.12

Chemistry 12
Calculate the pH of the following two buffer solutions: (a) 2.0 M CH3COONa/2.0 M CH3COOH, (b) 0.20 M CH3COONa/0.20 M CH3COOH. Which is the more effective buffer? Why?

Chapter: 16 Problem: 16.13

Chemistry 12
The pH of a bicarbonate-carbonic acid buffer is 8.00. Calculate the ratio of the concentration of carbonic acid (H2CO3) to that of the bicarbonate ion (HCO3 2).

Chapter: 16 Problem: 16.14

Chemistry 12
Very concentrated NaOH solutions should not be stored in Pyrex glassware. Why? (Hint: See Section 11.7.

Chapter: 16 Problem: 16.15

Chemistry 12
The pH of a sodium acetateacetic acid buffer is 4.50. Calculate the ratio [CH3COO2]/[CH3COOH].

Chapter: 16 Problem: 16.16

Chemistry 12
The pH of blood plasma is 7.40. Assuming the principal buffer system is HCO3 2/H2CO3, calculate the ratio [HCO3 2]/[H2CO3]. Is this buffer more effective against an added acid or an added base?

Chapter: 16 Problem: 16.17

Chemistry 12
Calculate the pH of the 0.20 M NH3/0.20 M NH4Cl buffer. What is the pH of the buffer after the addition of 10.0 mL of 0.10 M HCl to 65.0 mL of the buffer?

Chapter: 16 Problem: 16.18

Chemistry 12
Calculate the pH of 1.00 L of the buffer 1.00 M CH3COONa/1.00 M CH3COOH before and after the addition of (a) 0.080 mol NaOH, (b) 0.12 mol HCl. (Assume that there is no change in volume.)

Chapter: 16 Problem: 16.19

Chemistry 12
A diprotic acid, H2A, has the following ionization constants: Ka1 5 1.1 3 1023 and Ka25 2.5 3 1026 . In order to make up a buffer solution of pH 5.80, which combination would you choose? NaHA/H2A or Na2A/NaHA

Chapter: 16 Problem: 16.2b

Chemistry 12
A student is asked to prepare a buffer solution at pH 5 8.60, using one of the following weak acids: HA (Ka 5 2.7 3 1023 ), HB (Ka 5 4.4 3 1026 ), HC (Ka 5 2.6 3 1029 ). Which acid should she choose? Why?

Chapter: 16 Problem: 16.21

Chemistry 12
Very concentrated NaOH solutions should not be stored in Pyrex glassware. Why? (Hint: See Section 11.7.

Chapter: 16 Problem: 16.22

Chemistry 12
Very concentrated NaOH solutions should not be stored in Pyrex glassware. Why? (Hint: See Section 11.7.

Chapter: 16 Problem: 16.23

Chemistry 12
Very concentrated NaOH solutions should not be stored in Pyrex glassware. Why? (Hint: See Section 11.7.

Chapter: 16 Problem: 16.24

Chemistry 12
How much HCl (in moles) must be added to 1 L of a buffer solution that is 0.84 M in ammonia and 0.96 M in ammonium chloride to result in a buffer solution of pH 8.56? Assume volume to remain constant.

Chapter: 16 Problem: 16.25

Chemistry 12
Briefly describe what happens in an acid-base titration.

Chapter: 16 Problem: 16.26

Chemistry 12
Sketch titration curves for the following acid-base titrations: (a) HCl versus NaOH, (b) HCl versus CH3NH2, (c) CH3COOH versus NaOH. In each case, the base is added to the acid in an Erlenmeyer flask. Your graphs should show pH on the y-axis and volume of base added on the x-axis.

Chapter: 16 Problem: 16.27

Chemistry 12
A 0.2688-g sample of a monoprotic acid neutralizes 16.4 mL of 0.08133 M KOH solution. Calculate the molar mass of the acid

Chapter: 16 Problem: 16.28

Chemistry 12
A 5.00-g quantity of a diprotic acid was dissolved in water and made up to exactly 250 mL. Calculate the molar mass of the acid if 25.0 mL of this solution required 11.1 mL of 1.00 M KOH for neutralization. Assume that both protons of the acid were

Chapter: 16 Problem: 16.29

Chemistry 12
Very concentrated NaOH solutions should not be stored in Pyrex glassware. Why? (Hint: See Section 11.7.

Chapter: 16 Problem: 16.3b

Chemistry 12
In a titration experiment, 20.4 mL of 0.883 M HCOOH neutralize 19.3 mL of Ba(OH)2. What is the concentration of the Ba(OH)2 solution?

Chapter: 16 Problem: 16.31

Chemistry 12
A 0.1276-g sample of an unknown monoprotic acid was dissolved in 25.0 mL of water and titrated with 0.0633 M NaOH solution. The volume of base required to bring the solution to the equivalence point was 18.4 mL. (a) Calculate the molar mass of the acid. (b) After 10.0 mL of base had been added during

Chapter: 16 Problem: 16.32

Chemistry 12
A solution is made by mixing 5.00 3 102 mL of 0.167 M NaOH with 5.00 3 102 mL of 0.100 M CH3COOH. Calculate the equilibrium concentrations of H1, CH3COOH, CH3COO2, OH2, and Na1

Chapter: 16 Problem: 16.33

Chemistry 12
Calculate the pH at the equivalence point for the following titration: 0.20 M HCl versus 0.20 M methylamine (CH3NH2). (See Table 15.4.)

Chapter: 16 Problem: 16.34

Chemistry 12
Calculate the pH at the equivalence point for the follow ing titration: 0.10 M HCOOH versus 0.10 M NaOH.

Chapter: 16 Problem: 16.35

Chemistry 12
A 25.0-mL solution of 0.100 M CH3COOH is titrated with a 0.200 M KOH solution. Calculate the pH after the following additions of the KOH solution: (a) 0.0 mL, (b) 5.0 mL, (c) 10.0 mL, (d) 12.5 mL, (e) 15.0 mL

Chapter: 16 Problem: 16.36

Chemistry 12
A 10.0-mL solution of 0.300 M NH3 is titrated with a 0.100 M HCl solution. Calculate the pH after the following additions of the HCl solution: (a) 0.0 mL, (b) 10.0 mL, (c) 20.0 mL, (d) 30.0 mL, (e) 40.0 mL

Chapter: 16 Problem: 16.37

Chemistry 12
The diagrams shown here represent solutions at different stages in the titration of a weak acid HA with NaOH. Identify the solution that corresponds to (1) the initial stage before the addition of NaOH, (2) halfway to the equivalence point, (3) the equivalence point, (4) beyond the equivalence point.

Chapter: 16 Problem: 16.38

Chemistry 12
The diagrams shown here represent solutions at various stages in the titration of a weak base B (such as NH3) with HCl. Identify the solution that corresponds to (1) the initial stage before the addition of HCl, (2) halfway to the equivalence point, (3) the equivalence point, (4) beyond the equivalen

Chapter: 16 Problem: 16.39

Chemistry 12
A 0.054 M HNO2 solution is titrated with a KOH solution. What is [H1] at half way to the equivalence point?

Chapter: 16 Problem: 16.4b

Chemistry 12
A student titrates an unknown monoprotic acid with a NaOH solution from a buret. After the addition of 12.35 mL of NaOH, the pH of the solution read 5.22. The equivalence point is reached at 24.70 mL of NaOH. What is the Ka of the acid?

Chapter: 16 Problem: 16.41

Chemistry 12
Explain how an acid-base indicator works in a titration. What are the criteria for choosing an indicator for a particular acid-base titration?

Chapter: 16 Problem: 16.42

Chemistry 12
The amount of indicator used in an acid-base titration must be small. Why?

Chapter: 16 Problem: 16.43

Chemistry 12
Referring to Table 16.1, specify which indicator or indicators you would use for the following titrations: (a) HCOOH versus NaOH, (b) HCl versus KOH, (c) HNO3 versus CH3NH2.

Chapter: 16 Problem: 16.44

Chemistry 12
A student carried out an acid-base titration by adding NaOH solution from a buret to an Erlenmeyer flask containing HCl solution and using phenolphthalein as indicator. At the equivalence point, she observed a faint reddish-pink color. However, after a few minutes, the solution gradually turned color

Chapter: 16 Problem: 16.45

Chemistry 12
The ionization constant Ka of an indicator HIn is 1.0 3 1026 . The color of the nonionized form is red and that of the ionized form is yellow. What is the color of this indicator in a solution whose pH is 4.00?

Chapter: 16 Problem: 16.46

Chemistry 12
The Ka of a certain indicator is 2.0 3 1026 . The color of HIn is green and that of In2 is red. A few drops of the indicator are added to a HCl solution, which is then titrated against a NaOH solution. At what pH will the indicator change color?

Chapter: 16 Problem: 16.47

Chemistry 12
Use BaSO4 to distinguish between solubility, molar solubility, and solubility product

Chapter: 16 Problem: 16.48

Chemistry 12
Why do we usually not quote the Ksp values for soluble ionic compounds?

Chapter: 16 Problem: 16.49

Chemistry 12
Write balanced equations and solubility product expressions for the solubility equilibria of the following compounds: (a) CuBr, (b) ZnC2O4, (c) Ag2CrO4, (d) Hg2Cl2, (e) AuCl3, (f) Mn3(PO4)2.

Chapter: 16 Problem: 16.5b

Chemistry 12
Write the solubility product expression for the ionic compound AxBy

Chapter: 16 Problem: 16.51

Chemistry 12
How can we predict whether a precipitate will form when two solutions are mixed?

Chapter: 16 Problem: 16.52

Chemistry 12
Silver chloride has a larger Ksp than silver carbonate (see Table 16.2). Does this mean that AgCl also has a larger molar solubility than Ag2CO3?

Chapter: 16 Problem: 16.53

Chemistry 12
Calculate the concentration of ions in the following saturated solutions: (a) [I2] in AgI solution with [Ag1] 5 9.1 3 1029 M, (b) [Al31] in Al(OH)3 solution with [OH2] 5 2.9 3 1029 M.

Chapter: 16 Problem: 16.54

Chemistry 12
From the solubility data given, calculate the solubility products for the following compounds: (a) SrF2, 7.3 3 1022 g/L, (b) Ag3PO4, 6.7 3 1023 g/L

Chapter: 16 Problem: 16.55

Chemistry 12
The molar solubility of MnCO3 is 4.2 3 1026 M. What is Ksp for this compound?

Chapter: 16 Problem: 16.56

Chemistry 12
The solubility of an ionic compound MX (molar mass 5 346 g) is 4.63 3 1023 g/L. What is Ksp for the compound?

Chapter: 16 Problem: 16.57

Chemistry 12
The solubility of an ionic compound M2X3 (molar mass 5 288 g) is 3.6 3 10217 g/L. What is Ksp for the compound?

Chapter: 16 Problem: 16.58

Chemistry 12
Using data from Table 16.2, calculate the molar solubility of CaF2.

Chapter: 16 Problem: 16.59

Chemistry 12
What is the pH of a saturated zinc hydroxide solution?

Chapter: 16 Problem: 16.6b

Chemistry 12
The pH of a saturated solution of a metal hydroxide MOH is 9.68. Calculate the Ksp for the compound.

Chapter: 16 Problem: 16.61

Chemistry 12
If 20.0 mL of 0.10 M Ba(NO3)2 are added to 50.0 mL of 0.10 M Na2CO3, will BaCO3 precipitate?

Chapter: 16 Problem: 16.62

Chemistry 12
A volume of 75 mL of 0.060 M NaF is mixed with 25 mL of 0.15 M Sr(NO3)2. Calculate the concentrations in the final solution of NO3 2, Na1, Sr21, and F2. (Ksp for SrF2 5 2.0 3 10210.)

Chapter: 16 Problem: 16.63

Chemistry 12
Solid NaI is slowly added to a solution that is 0.010 M in Cu1 and 0.010 M in Ag1. (a) Which compound will begin to precipitate first? (b) Calculate [Ag1] when CuI just begins to precipitate. (c) What percent of Ag1 remains in solution at this point?

Chapter: 16 Problem: 16.64

Chemistry 12
Find the approximate pH range suitable for the separation of Fe31 and Zn21 ions by precipitation of Fe(OH)3 from a solution that is initially 0.010 M in both Fe31 and Zn21. Assume a 99 percent precipitation of Fe(OH)3.

Chapter: 16 Problem: 16.65

Chemistry 12
How does the common ion effect influence solubility equilibria? Use Le Chteliers principle to explain the decrease in solubility of CaCO3 in a Na2CO3 solution.

Chapter: 16 Problem: 16.66

Chemistry 12
The molar solubility of AgCl in 6.5 3 1023 M AgNO3 is 2.5 3 1028 M. In deriving Ksp from these data, which of the following assumptions are reasonable? (a) Ksp is the same as solubility. (b) Ksp of AgCl is the same in 6.5 3 1023 M AgNO3 as in pure water. (c) Solubility of AgCl is independent of the c

Chapter: 16 Problem: 16.67

Chemistry 12
How many grams of CaCO3 will dissolve in 3.0 3 102 mL of 0.050 M Ca(NO3)2?

Chapter: 16 Problem: 16.68

Chemistry 12
The solubility product of PbBr2 is 8.9 3 1026 . Determine the molar solubility (a) in pure water, (b) in 0.20 M KBr solution, (c) in 0.20 M Pb(NO3)2 solution.

Chapter: 16 Problem: 16.69

Chemistry 12
Calculate the molar solubility of AgCl in a 1.00-L solution containing 10.0 g of dissolved CaCl2

Chapter: 16 Problem: 16.7b

Chemistry 12
Calculate the molar solubility of BaSO4 (a) in water, (b) in a solution containing 1.0 M SO4 22 ions.

Chapter: 16 Problem: 16.71

Chemistry 12
Which of the following ionic compounds will be more soluble in acid solution than in water? (a) BaSO4, (b) PbCl2, (c) Fe(OH)3, (d) CaCO3.

Chapter: 16 Problem: 16.72

Chemistry 12
Which of the following will be more soluble in acid solution than in pure water? (a) CuI, (b) Ag2SO4, (c) Zn(OH)2, (d) BaC2O4, (e) Ca3(PO4)2.

Chapter: 16 Problem: 16.73

Chemistry 12
Compare the molar solubility of Mg(OH)2 in water and in a solution buffered at a pH of 9.0.

Chapter: 16 Problem: 16.74

Chemistry 12
Calculate the molar solubility of Fe(OH)2 in a solution buffered at (a) pH 8.00, (b) pH 10.00

Chapter: 16 Problem: 16.75

Chemistry 12
The solubility product of Mg(OH)2 is 1.2 3 10211. What minimum OH2 concentration must be attained (for example, by adding NaOH) to decrease the Mg21 concentration in a solution of Mg(NO3)2 to less than 1.0 3 10210 M?

Chapter: 16 Problem: 16.76

Chemistry 12
Calculate whether or not a precipitate will form if 2.00 mL of 0.60 M NH3 are added to 1.0 L of 1.0 3 1023 M FeSO4.

Chapter: 16 Problem: 16.77

Chemistry 12
Explain the formation of complexes in Table 16.4 in terms of Lewis acid-base theory

Chapter: 16 Problem: 16.78

Chemistry 12
Give an example to illustrate the general effect of complex ion formation on solubility

Chapter: 16 Problem: 16.79

Chemistry 12
If 2.50 g of CuSO4 are dissolved in 9.0 3 102 mL of 0.30 M NH3, what are the concentrations of Cu21, Cu(NH3)4 21, and NH3 at equilibrium?

Chapter: 16 Problem: 16.8b

Chemistry 12
Calculate the concentrations of Cd21, Cd(CN)4 22, and CN2 at equilibrium when 0.50 g of Cd(NO3)2 dissolves in 5.0 3 102 mL of 0.50 M NaCN

Chapter: 16 Problem: 16.81

Chemistry 12
If NaOH is added to 0.010 M Al31, which will be the predominant species at equilibrium: Al(OH)3 or Al(OH)4 2? The pH of the solution is 14.00. [Kf for Al(OH)4 2 5 2.0 3 1033.]

Chapter: 16 Problem: 16.82

Chemistry 12
Calculate the molar solubility of AgI in a 1.0 M NH3 solution

Chapter: 16 Problem: 16.83

Chemistry 12
Both Ag1 and Zn21 form complex ions with NH3. Write balanced equations for the reactions. However, Zn(OH)2 is soluble in 6 M NaOH, and AgOH is not. Explain

Chapter: 16 Problem: 16.84

Chemistry 12
Explain, with balanced ionic equations, why (a) CuI2 dissolves in ammonia solution, (b) AgBr dissolves in NaCN solution, (c) HgCl2 dissolves in KCl solution.

Chapter: 16 Problem: 16.85

Chemistry 12
Outline the general procedure of qualitative analysis.

Chapter: 16 Problem: 16.86

Chemistry 12
Give two examples of metal ions in each group (1 through 5) in the qualitative analysis scheme.

Chapter: 16 Problem: 16.87

Chemistry 12
In a group 1 analysis, a student obtained a precipitate containing both AgCl and PbCl2. Suggest one reagent that would enable her to separate AgCl(s) from PbCl2(s)

Chapter: 16 Problem: 16.88

Chemistry 12
In a group 1 analysis, a student adds HCl acid to the unknown solution to make [Cl2] 5 0.15 M. Some PbCl2 precipitates. Calculate the concentration of Pb21 remaining in solution.

Chapter: 16 Problem: 16.89

Chemistry 12
Both KCl and NH4Cl are white solids. Suggest one reagent that would enable you to distinguish between these two compounds

Chapter: 16 Problem: 16.9b

Chemistry 12
Describe a simple test that would enable you to distinguish between AgNO3(s) and Cu(NO3)2(s).

Chapter: 16 Problem: 16.91

Chemistry 12
To act as an effective buffer, the concentrations of the acid and the conjugate base should not differ by more than a factor of 10, that is, 10 $ [conjugate base] [acid] $ 0.1 (a) Show that the buffer range, that is, the range of the concentration ratio over which the buffer is effective, is given by

Chapter: 16 Problem: 16.92

Chemistry 12
The pKa of the indicator methyl orange is 3.46. Over what pH range does this indicator change from 90 percent HIn to 90 percent In2?

Chapter: 16 Problem: 16.93

Chemistry 12
The iodide impurity in a 4.50-g sample of a metal nitrate is precipitated as silver iodide. If 5.54 mL of 0.186 M AgNO3 solution is needed for the precipitation, calculate the mass percent of iodide in the sample.

Chapter: 16 Problem: 16.94

Chemistry 12
A sodium acetate-acetic acid buffer solution was prepared by adding a 0.020 M HCl solution to 500 mL of 0.020 M CH3COONa and then diluting the mixed solution to 1.0 L. Calculate the original volume of the HCl solution needed to prepare a buffer solution of pH 5.00.

Chapter: 16 Problem: 16.95

Chemistry 12
Sketch the titration curve of a weak acid versus a strong base like the one shown in Figure 16.5. On your graph indicate the volume of base used at the equivalence point and also at the halfequivalence point, that is, the point at which half of the acid has been neutralized. Show how you can measure

Chapter: 16 Problem: 16.96

Chemistry 12
A 200-mL volume of NaOH solution was added to 400 mL of a 2.00 M HNO2 solution. The pH of the mixed solution was 1.50 units greater than that of the original acid solution. Calculate the molarity of the NaOH solution.

Chapter: 16 Problem: 16.97

Chemistry 12
The pKa of butyric acid (HBut) is 4.7. Calculate Kb for the butyrate ion (But2).

Chapter: 16 Problem: 16.98

Chemistry 12
A solution is made by mixing 5.00 3 102 mL of 0.167 M NaOH with 5.00 3 102 mL 0.100 M HCOOH. Calculate the equilibrium concentrations of H1, HCOOH, HCOO2, OH2, and Na1.

Chapter: 16 Problem: 16.99

Chemistry 12
Very concentrated NaOH solutions should not be stored in Pyrex glassware. Why? (Hint: See Section 11.7.

Chapter: 16 Problem: 16.1c

Chemistry 12
A student mixes 50.0 mL of 1.00 M Ba(OH)2 with 86.4 mL of 0.494 M H2SO4. Calculate the mass of BaSO4 formed and the pH of the mixed solution.

Chapter: 16 Problem: 16.101

Chemistry 12
For which of the following reactions is the equilibrium constant called a solubility product? (a) Zn(OH)2(s) 1 2OH2(aq) Zn(OH)4 22(aq) (b) 3Ca21(aq) 1 2PO32 4 (aq) Ca3(PO4)2(s) (c) CaCO3(s) 1 2H1(aq) Ca21(aq) 1 H2O(l) 1 CO2(g) (d) PbI2(s) Pb21(aq) 1 2I2(aq)

Chapter: 16 Problem: 16.102

Chemistry 12
A 2.0-L kettle contains 116 g of boiler scale (CaCO3). How many times would the kettle have to be completely filled with distilled water to remove all of the deposit?

Chapter: 16 Problem: 16.103

Chemistry 12
Equal volumes of 0.12 M AgNO3 and 0.14 M ZnCl2 solution are mixed. Calculate the equilibrium concentrations of Ag1, Cl2, Zn21, and NO3 2

Chapter: 16 Problem: 16.104

Chemistry 12
Calculate the solubility (in g/L) of Ag2CO3.

Chapter: 16 Problem: 16.105

Chemistry 12
Find the approximate pH range suitable for separating Mg21 and Zn21 by the precipitation of Zn(OH)2 from a solution that is initially 0.010 M in Mg21 and Zn21.

Chapter: 16 Problem: 16.106

Chemistry 12
A volume of 25.0 mL of 0.100 M HCl is titrated against a 0.100 M CH3NH2 solution added to it from a buret. Calculate the pH values of the solution (a) after 10.0 mL of CH3NH2 solution have been added, (b) after 25.0 mL of CH3NH2 solution have been added, (c) after 35.0 mL of CH3NH2 solution have been

Chapter: 16 Problem: 16.107

Chemistry 12
The molar solubility of Pb(IO3)2 in a 0.10 M NaIO3 solution is 2.4 3 10211 mol/L. What is Ksp for Pb(IO3)2?

Chapter: 16 Problem: 16.108

Chemistry 12
When a KI solution was added to a solution of mercury(II) chloride, a precipitate [mercury(II) iodide] formed. A student plotted the mass of the precipitate versus the volume of the KI solution added and obtained the following graph. Explain the appearance of the graph. Mass of HgI2 formed Volume of

Chapter: 16 Problem: 16.109

Chemistry 12
Barium is a toxic substance that can seriously impair heart function. For an X ray of the gastrointestinal tract, a patient drinks an aqueous suspension of 20 g BaSO4. If this substance were to equilibrate with the 5.0 L of the blood in the patients body, what would be [Ba21]? For a good estimate, we

Chapter: 16 Problem: 16.11c

Chemistry 12
The pKa of phenolphthalein is 9.10. Over what pH range does this indicator change from 95 percent HIn to 95 percent In2?

Chapter: 16 Problem: 16.111

Chemistry 12
Solid NaBr is slowly added to a solution that is 0.010 M in Cu1 and 0.010 M in Ag1. (a) Which compound will begin to precipitate first? (b) Calculate [Ag1] when CuBr just begins to precipitate. (c) What percent of Ag1 remains in solution at this point?

Chapter: 16 Problem: 16.112

Chemistry 12
Cacodylic acid is (CH3)2AsO2H. Its ionization constant is 6.4 3 1027 . (a) Calculate the pH of 50.0 mL of a 0.10 M solution of the acid. (b) Calculate the pH of 25.0 mL of 0.15 M (CH3)2AsO2Na. (c) Mix the solutions in part (a) and part (b). Calculate the pH of the resulting solution

Chapter: 16 Problem: 16.113

Chemistry 12
Radiochemical techniques are useful in estimating the solubility product of many compounds. In one experiment, 50.0 mL of a 0.010 M AgNO3 solution containing a silver isotope with a radioactivity of 74,025 counts per min per mL were mixed with 100 mL of a 0.030 M NaIO3 solution. The mixed solution wa

Chapter: 16 Problem: 16.114

Chemistry 12
The molar mass of a certain metal carbonate, MCO3, can be determined by adding an excess of HCl acid to react with all the carbonate and then back titrating the remaining acid with a NaOH solution. (a) Write equations for these reactions. (b) In a certain experiment, 18.68 mL of 5.653 M HCl were adde

Chapter: 16 Problem: 16.115

Chemistry 12
Acid-base reactions usually go to completion. Confirm this statement by calculating the equilibrium constant for each of the following cases: (a) A strong acid reacting with a strong base. (b) A strong acid reacting with a weak base (NH3). (c) A weak acid (CH3COOH) reacting with a strong base. (d) A

Chapter: 16 Problem: 16.116

Chemistry 12
Calculate x, which is the number of molecules of water in oxalic acid hydrate, H2C2O4 ? xH2O, from the following data: 5.00 g of the compound is made up to exactly 250 mL solution, and 25.0 mL of this solution requires 15.9 mL of 0.500 M NaOH solution for neutralization.

Chapter: 16 Problem: 16.117

Chemistry 12
Very concentrated NaOH solutions should not be stored in Pyrex glassware. Why? (Hint: See Section 11.7.

Chapter: 16 Problem: 16.118

Chemistry 12
Phenolphthalein is the common indicator for the titration of a strong acid with a strong base. (a) If the pKa of phenolphthalein is 9.10, what is the ratio of the nonionized form of the indicator (colorless) to the ionized form (reddish pink) at pH 8.00? (b) If 2 drops of 0.060 M phenolphthalein are

Chapter: 16 Problem: 16.119

Chemistry 12
Oil paintings containing lead(II) compounds as constituents of their pigments darken over the years. Suggest a chemical reason for the color change

Chapter: 16 Problem: 16.12c

Chemistry 12
What reagents would you employ to separate the following pairs of ions in solution? (a) Na1 and Ba21, (b) K1 and Pb21, (c) Zn21 and Hg21

Chapter: 16 Problem: 16.121

Chemistry 12
Look up the Ksp values for BaSO4 and SrSO4 in Table 16.2. Calculate the concentrations of Ba21, Sr21, and SO4 22 in a solution that is saturated with both compounds.

Chapter: 16 Problem: 16.122

Chemistry 12
In principle, amphoteric oxides, such as Al2O3 and BeO, can be used to prepare buffer solutions because they possess both acidic and basic properties (see Section 15.11). Explain why these compounds are of little practical use as buffer components

Chapter: 16 Problem: 16.123

Chemistry 12
CaSO4 (Ksp 5 2.4 3 1025 ) has a larger Ksp value than that of Ag2SO4 (Ksp 5 1.4 3 1025 ). Does it follow that CaSO4 also has greater solubility (g/L)?

Chapter: 16 Problem: 16.124

Chemistry 12
When lemon juice is squirted into tea, the color becomes lighter. In part, the color change is due to dilution, but the main reason for the change is an acid-base reaction. What is the reaction? (Hint: Tea contains polyphenols which are weak acids and lemon juice contains citric acid.)

Chapter: 16 Problem: 16.125

Chemistry 12
How many millilitres of 0.1 M NaOH must be added to a 200 mL of 0.10 M to make a buffer solution with a pH of 7.50?

Chapter: 16 Problem: 16.126

Chemistry 12
The maximum allowable concentration of Pb21 ions in drinking water is 0.05 ppm (that is, 0.05 g of Pb21 in 1 million g of water). Is this guideline exceeded if an underground water supply is at equilibrium with the mineral anglesite, PbSO4 (Ksp 5 1.6 3 1028 )?

Chapter: 16 Problem: 16.127

Chemistry 12
One of the most common antibiotics is penicillin G (benzylpenicillinic acid), which has the structure shown next: O B H COOH G D O H3C C G D NOC J H C AAA D G COCONOCOCH2O H3C S A A B H H O It is a weak monoprotic acid: HP H1 1 P2 Ka 5 1.64 3 1023 where HP denotes the parent acid and P2 the conjugat

Chapter: 16 Problem: 16.128

Chemistry 12
Which of the following solutions has the highest [H1]? (a) 0.10 M HF, (b) 0.10 M HF in 0.10 M NaF, (c) 0.10 M HF in 0.10 M SbF5. (Hint: SbF5 reacts with F2 to form the complex ion SbF6 2.)

Chapter: 16 Problem: 16.129

Chemistry 12
Distribution curves show how the fractions of nonionized acid and its conjugate base vary as a function of pH of the medium. Plot distribution curves for CH3COOH and its conjugate base CH3COO2 in solution. Your graph should show fraction as the y axis and pH as the x axis. What are the fractions and

Chapter: 16 Problem: 16.13c

Chemistry 12
Water containing Ca21 and Mg21 ions is called hard water and is unsuitable for some household and industrial use because these ions react with soap to form insoluble salts, or curds. One way to remove the Ca21 ions from hard water is by adding washing soda (Na2CO3 ? 10H2O). (a) The molar solubility o

Chapter: 16 Problem: 16.131

Chemistry 12
Consider the ionization of the following acid-base indicator: HIn(aq) H1(aq) 1 In2(aq) The indicator changes color according to the ratios of the concentrations of the acid to its conjugate base as described on p. 740. Show that the pH range over which the indicator changes from the acid color to th

Chapter: 16 Problem: 16.132

Chemistry 12
Amino acids are building blocks of proteins. These compounds contain at least one amino group (NH2) and one carboxyl group (COOH). Consider glycine (NH2CH2COOH). Depending on the pH of the solution, glycine can exist in one of three possible forms: Fully ionized: NH2CH2COO Dipolar ion: NH3CH2COO Fu

Chapter: 16 Problem: 16.133

Chemistry 12
(a) Referring to Figure 16.6, describe how you would determine the pKb of the base. (b) Derive an analogous Henderson-Hasselbalch equation relating pOH to pKb of a weak base B and its conjugate acid HB1. Sketch a titration curve showing the variation of the pOH of the base solution versus the volume

Chapter: 16 Problem: 16.134

Chemistry 12
A 25.0-mL of 0.20 M HF solution is titrated with a 0.20 M NaOH solution. Calculate the volume of NaOH solution added when the pH of the solution is (a) 2.85, (b) 3.15, (c) 11.89. Ignore salt hydrolysis.

Chapter: 16 Problem: 16.135

Chemistry 12
Draw distribution curves for an aqueous carbonic acid solution. Your graph should show fraction of species present as the y axis and pH as the x axis. Note that at any pH, only two of the three species (H2CO3, HCO3 2, and CO3 22) are present in appreciable concentrations. Use the pKa values in Table

Chapter: 16 Problem: 16.136

Chemistry 12
One way to distinguish a buffer solution with an acid solution is by dilution. (a) Consider a buffer solution made of 0.500 M CH3COOH and 0.500 M CH3COONa. Calculate its pH and the pH after it has been diluted 10-fold. (b) Compare the result in (a) with the pHs of a 0.500 M CH3COOH solution before an

Chapter: 16 Problem: 16.137

Chemistry 12
Histidine is one of the 20 amino acids found in proteins. Shown here is a fully protonated histidine molecule where the numbers denote the pKa values of the acidic groups. CH2 A A H3NOCHOCOOH O B HN NH 9.17 1.82 6.00 (a) Show stepwise ionization of histidine in solution. (Hint: The H1 ion will fi

Chapter: 16 Problem: 16.138

Chemistry 12
A sample of 0.96 L of HCl at 372 mmHg and 22C is bubbled into 0.034 L of 0.57 M NH3. What is the pH of the resulting solution? Assume the volume of solution remains constant and that the HCl is totally dissolved in the solution.

Chapter: 16 Problem: 16.139

Chemistry 12
(a) Assuming complete dissociation and no ionpair formation, calculate the freezing point of a 0.50 m NaI solution. (b) What is the freezing point after the addition of sufficient HgI2, an insoluble compound, to the solution to react with all the free I2 ions in solution? Assume volume to remain cons

Chapter: 16 Problem: 16.14c

Chemistry 12
Calculate the maximum mass (in grams) of each of the following soluble salts that can be added to 200 mL of 0.100 M MgCl2 without causing a precipitate to form: (a) Na2CO3, (b) AgNO3, (c) KOH. Assume volume to remain constant.

Chapter: 16 Problem: 16.141

Chemistry 12
A 1.0-L saturated silver carbonate solution at 5C is treated with enough hydrochloric acid to decompose the compound. The carbon dioxide generated is collected in a 19-mL vial and exerts a pressure of 114 mmHg at 25C. What is the Ksp of Ag2CO3 at 5C?

Chapter: 16 Problem: 16.142

Chemistry 12
The two curves shown on p. 772 represent the titration of two weak acids of the same concentration with a strong base such as NaOH. Use three observations to determine which of the two acids is stronger. and Solubility Equilibria 0 0 10 20 30 pH Volume of NaOH added (mL) (a) 40 2 4 6 8 10 12 14 0 0 1

Chapter: 16 Problem: 16.143

Chemistry 12
The two curves shown here represent the titration of two weak bases of the same concentration with a strong acid such as HCl. Use three observations to determine which of the two bases is stronger. 0 0 10 20 30 pH Volume of HCl added (mL) (a) 40 2 4 6 8 10 12 14 0 0 10 20 30 pH Volume of HCl added (m

Chapter: 16 Problem: 16.144

Chemistry 12
A 100-mL 0.100 M CuSO4 solution is mixed with a 100-mL 0.100 M Ba(OH)2 solution. Calculate the concentrations of the ions in the combined solution.

Chapter: 16 Problem: 16.145

Chemistry 12
The titration curve shown here represents the titration of a weak diprotic acid (H2A) versus NaOH. Identify the major species present at the marked points and estimate the pKa1 and pKa2 values of the acid. 0 pH 0 10 20 30 Volume of NaOH added (mL) 40 50 60 70 2 4 6 8 10 12 14

Chapter: 16 Problem: 16.146

Chemistry 12
The titration curve shown here represents the titration of a weak dibasic base (for example, a compound that contains two NH2 groups) versus HCl. Identify the major species present at the marked points and estimate the pKb1 and pKb2 values of the base. 0 pH 0 10 20 30 Volume of HCl added (mL) 40 50 6

Chapter: 16 Problem: 16.147

Chemistry 12
Use appropriate equations to account for the solubility of the amphoteric aluminum hydroxide [Al(OH)3)] at low and high pHs. Molar solubility pH 9876543 10 11 12 13

Chapter: 16 Problem: 16.148

Chemistry 12
From Table 16.2 we see that silver bromide (AgBr) has a larger solubility product than iron(II) hydroxide [Fe(OH)2]. Does this mean that AgBr is more soluble than Fe(OH)2?

Chapter: 16 Problem: 16.149

Chemistry 12
Aspirin is a weak acid with pKa 5 3.5. What is the ratio of neutral (protonated) aspirin to deprotonated aspirin in the following body fluids: (a) saliva, (b) gastric juices in the stomach, and (c) blood?

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