Acid–Base Titrations (Section)Assume that 30.0 mL of a

The following boxes represent aqueous solutions containing a weak acid, HX, and its conjugate base,X- . Water molecules, hydronium ions and cations are not shown. Which solution has the highest pH? Explain. [Section 17.1]

Problem 1PE Calculating the pH When a Common Ion Is Involved What is the pH of a solution made by adding 0.30 mol of acetic acid and 0.30 mol of sodium acetate to enough water to make 1.0 L of solution? Calculate the pH of a solution containing 0.085 M nitrous acid (HNO2, Ka = 4.5 × 10-4) and 0.10 M potassium nitrite (KNO2).

The beaker on the right contains 0.1 M acetic acid solution with methyl orange as an indicator. The beaker on the left contains a mixture of 0.1 M acetic acid and 0.1 M sodium acetate with methyl orange. (a) Using Figure 16.7, what can you say about the pH of each solution? (b) Which solution is better able to maintain its pH when small amounts of NaOH are added? Explain. [Sections 17.1 and 17.2]

Problem 2PE Calculating Ion Concentrations When a Common Ion Is Involved Calculate the fluoride ion concentration and pH of a solution that is 0.20 M in HF and 0.10 M in HCl. Calculate the formate ion concentration and pH of a solution that is 0.050 M in formic acid (HCOOH, Ka = 1.8 × 10-4) and 0.10 M in HNO3.

Problem 3PE Calculating the pH of a Buffer What is the pH of a buffer that is 0.12 M in lactic acid [CH3CH(OH)COOH, or HC3H5O3] and 0.10 M in sodium lactate [CH3CH(OH)COONa or NaC3H5O3]? For lactic acid, Ka = 1.4 × 10-4. Calculate the pH of a buffer composed of 0.12 M benzoic acid and 0.20 M sodium benzoate. (Refer to Appendix D.)

The drawing on the left represents a buffer composed of equal concentrations of a weak acid, HX, and its conjugate base, X- . The heights of the columns are proportional to the concentra- tions of the components of the buffer. (a) Which of the three drawings, (1), (2), or (3), represents the buffer after the addition of a strong acid? (b) Which of the three represents the buffer after the addition of a strong base? (c) Which of the three represents a situation that cannot arise from the addition of either an acid or a base? [Section 17.2]

Problem 4PE Preparing a Buffer How many moles of NH4Cl must be added to 2.0 L of 0.10 M NH3 to form a buffer whose pH is 9.00? (Assume that the addition of NH4Cl does not change the volume of the solution.) Calculate the concentration of sodium benzoate that must be present in a 0.20 M solution of benzoic acid (C6H5COOH) to produce a pH of 4.00. Refer to Appendix D.

The following drawings represent solutions at various stages of the titration of a weak acid, HA, with NaOH. (The ions and water molecules have been omitted for clarity.) To which of the following regions of the titration curve does each drawing correspond: (a) before addition of NaOH, (b) after addition of NaOH but before equivalence point, (c) at equivalence point, (d) after equivalence point? [Section 17.3]

Problem 5PE Calculating pH Changes in Buffers A buffer is made by adding 0.300 mol CH3COOH and 0.300 mol CH3COONa to enough water to make 1.000 L of solution. The pH of the buffer is 4.74 (Sample Exercise 17.1). (a) Calculate the pH of this solution after 5.0 mL of 4.0 M NaOH(aq) solution is added., (b) For comparison, calculate the pH of a solution made by adding 5.0 mL of 4.0 M NaOH(aq) solution to 1.000 L of pure water. Determine (a) the pH of the original buffer described in Sample Exercise 17.6 after the addition of 0.020 mol HCl and, (b) the pH of the solution that would result from the addition of 0.020 mol HCl to 1.000 L of pure water.

Problem 6PE Calculations for a Strong Acid–Strong Base Titration Calculate the pH when (a) 49.0 mL and, (b) 51.0 mL of 0.100 M NaOH solution have been added to 50.0 mL of 0.100M HCl solution. Calculate the pH when (a) 24.9 mL and, (b) 25.1 mL of 0.100 M HNO3 have been added to 25.0 mL of 0.100 M KOH solution.

Problem 7PE Calculations for a Weak Acid–Strong Base Titration Calculate the pH of the solution formed when 45.0 mL of 0.100 M NaOH is added to 50.0 mL of 0.100MCH3COOH (Ka = 1.8 × 10-5) (a) Calculate the pH in the solution formed by adding 10.0 mL of 0.050 M NaOH to 40.0 mL of 0.0250 M benzoic acid (C6H5COOH, Ka = 6.3 × 10-5). (b) Calculate the pH in the solution formed by adding 10.0 mL of 0.100 M HCl to 20.0 mL of 0.100 M NH3.

Equal volumes of two acids are titrated with 0.10 M NaOH resulting in the two titration curves shown in the following figure. (a) Which curve corresponds to the more concentrated acid solution? (b) Which corresponds to the acid with the larger Ka? Explain. [Section 17.3]

Problem 8PE Calculating the pH at the Equivalence Point Calculate the pH at the equivalence point in the titration of 50.0 mL of 0.100 M CH3COOH with 0.100 M NaOH. Calculate the pH at the equivalence point when (a) 40.0 mL of 0.025 M benzoic acid (C6H5COOH, Ka = 6.3 × 10-5) is titrated with 0.050 MNaOH and, (b) 40.0 mL of 0.100 M NH3 is titrated with 0.100 M HCl.

A saturated solution of Cd(OH)2 is shown in the middle beaker. If hydrochloric acid solution is added, the solubility of Cd(OH)2 will increase, causing additional solid to dissolve. Which of the two choices, Beaker A or Beaker B, accurately represents the solution after equilibrium is reestablished? Explain. (The water molecules and CI- ions are omitted for clarity). [Sections 17.4 and 17.5]

The following graphs represent the behavior of BaCO3 under different circumstances. In each case the vertical axis indicates the solubility of the BaCO3 and the horizontal axis represents the concentration of some other reagent.(a)Which graph represents what happens to the solubility of BaCO3 as HNO3 is added? (b) Which graph represents what happens to the BaCO3 solubility as Na2CO3 is added? (c) Which represents what happens to the BaCO3 solubility as NaNO3 is added? [Section 17.5]

\(Ca(OH)_2\) has a \(K_{sp}\) of \(6.5 X 10^{-6}\). (a) If 0.370 g of \(Ca(OH)_2\) is added to 500 mL of water and the mixture is allowed to come to equilibrium, will the solution be saturated? (b) If 50 mL of the solution from part (a) is added to each of the beakers shown here, in which beakers, if any, will a precipitate form? In those cases where a precipitate forms, what is its identity? [Section 17.6]

Problem 9PE Writing Solubility-Product (Ksp) Expressions Write the expression for the solubility-product constant for CaF2, and look up the corresponding Ksp value in Appendix D. Give the solubility-product-constant expressions and Ksp values (from Appendix D) for (a) barium carbonate and, (b) silver sulfate.

Problem 10PE Calculating Kspfrom Solubility Solid silver chromate is added to pure water at 25 °C, and some of the solid remains undissolved. The mixture is stirred for several days to ensure that equilibrium is achieved between the undissolved Ag2CrO4(s) and the solution. Analysis of the equilibrated solution shows that its silver ion concentration is 1.3 × 10-4 M. Assuming that the Ag2CrO4 solution is saturated and that there are no other important equilibria involving Ag+ or CrO4 2-ions in the solution, calculate Ksp for this compound. A saturated solution of Mg(OH)2 in contact with undissolved Mg(OH)2(s) is prepared at 25 °C. The pH of the solution is found to be 10.17. Assuming that there are no other simultaneous equilibria involving the Mg2+ or OH- ions, calculate Ksp for this compound.

What is the name given to the kind of behavior demonstrated by a metal hydroxide in this graph? [Section 17.5]

Problem 11PE Calculating Solubility from Ksp The Ksp for CaF2 is 3.9 × 10-11 at 25 °C. Assuming equilibrium is established between solid and dissolved CaF2, and that there are no other important equilibria affecting its solubility, calculate the solubility of CaF2 in grams per liter. The Ksp for LaF3 is 2 × 10-19. What is the solubility of LaF3 in water in moles per liter?

Three cations, Ni2+, Cu2+, and Ag+, are separated using two different precipitating agents. Based on Figure 17.23, what two precipitating agents could be used? Using these agents, indicate which of the cations is A, which is B, and which is C. [Section 17.7]

Problem 12PE Practice Exercise 2 Calculating the Effect of a Common Ion on Solubility Calculate the molar solubility of CaF2 at 25 °C in a solution that is (a) 0.010 M in Ca(NO3)2 and, (b) 0.010 M in NaF. For manganese(II) hydroxide, Mn(OH)2, Ksp = 1.6 × 10-13. Calculate the molar solubility of Mn(OH)2 in a solution that contains 0.020 M NaOH.

Problem 13E (a) What is the common-ion effect? (b) Give an example of a salt that can decrease the ionization of HNO2 in solution.

Problem 13PE Predicting the Effect of Acid on Solubility Which of these substances are more soluble in acidic solution than in basic solution: (a) Ni(OH)2(s), (b) CaCO3(s), (c) BaF2(s), (d) AgCl(s)? Write the net ionic equation for the reaction between a strong acid and (a) CuS (b) Cu(N3)2.

(a) Consider the equilibrium B(aq) + H2O(I) = HB+(aq) + OH-(aq) Using Le Châtelier’s principle, explain the effect of the presence of a salt of HB+ on the ionization of B. (b) Give an example of a salt that can decrease the ionization of NH3 in solution.

(a) Consider the equilibrium \(\mathrm{B}(a q)+\mathrm{H}_{2} \mathrm{O}(l) \rightleftharpoons \mathrm{HB}^{+}(a q)+\mathrm{OH}^{-}(a q)\). Using Le Châtelier’s principle, explain the effect of the presence of a salt of \(\mathrm{HB}^{+}\) on the ionization of B. (b) Give an example of a salt that can decrease the ionization of \(\mathrm{NH}_{3}\) in solution.

Problem 15E The Common–Ion Effect (Section) Use information from Appendix D to calculate the pH of (a) a solution that is 0.060 M in potassium propionate (C2H5COOK or KC3H5O2) and 0.085 M in propionic acid (C2H5COOK or HC3H5O2); (b) a solution that is 0.075 M in trimethylamine, (CH3)3N, and 0.10 M in trimethylammonium chloride, (CH3)3NHCl; (c) a solution that is made by mixing 50.0 mL of 0.15M acetic acid and 50.0 mL of 0.20 M sodium acetate.

Problem 15PE Predicting Whether a Precipitate Forms Does a precipitate form when 0.10 L of 8.0 × 10-3 M Pb × (NO3)2 is added to 0.40 L of 5.0 × 10-3 M Na2SO4? Does a precipitate form when 0.050 L of 2.0 × 10-2 M NaF is mixed with 0.010 L of 1.0 × 10-2 MCa(NO3)2?

Problem 16E The Common–Ion Effect (Section) Use information from Appendix D to calculate the pH of (a) a solution that is 0.250 M in sodium formate (HCOONa) and 0.100 M in formic acid (HCOOH) (b) a solution that is 0.510 M in pyridine (C5H5N) and 0.450 M in pyridinium chloride (C5H5NHCl) (c) a solution that is made by combining 55 mL of 0.050 M hydrofluoric acid with 125 mL of 0.10 M sodium fluoride.

Problem 16PE Selective Precipitation A solution contains 1.0 × 10-2 M Ag+ and 2.0 × 10-2 M Pb2 +. When Cl- is added, both AgCl (Ksp = 1.8 × 10-10) and PbCl2(Ksp = 1.7 × 10-5) can precipitate. What concentration of Cl- is necessary to begin the precipitation of each salt? Which salt precipitates first? A solution consists of 0.050 M Mg2 + and Cu2 + . Which ion precipitates first as OH- is added? What concentration of OH- is necessary to begin the precipitation of each cation? [Ksp = 1.8 × 10-11 for Mg(OH)2, and Ksp = 4.8 × 10-20 for Cu(OH)2.]

Problem 17E The Common–Ion Effect (Section) (a) Calculate the percent ionization of 0.0075 M butanoic acid (Ka = 1.5 × 10–5) (b) Calculate the percent ionization of 0.0075 M butanoic acid in a solution containing 0.085 M sodium butanoate.

Problem 18E (a) Calculate the percent ionization of 0.125 M lactic acid (Ka = 1.4 × 10–4) (b) Calculate the percent ionization of 0.125 M lactic acid in a solution containing 0.0075 Msodium lactate.

Problem 19E Explain why a mixture of CH3COOH and CH3COONa can act as a buffer while a mixture of HCl and NaCl cannot.

Problem 20E Explain why a mixture formed by mixing 100 mL of 0.100 M CH3COOH and 50 mL of 0.100 MNaOH will act as a buffer.

Problem 21E Buffers (Section) (a) Calculate the pH of a buffer that is 0.12 M in lactic acid and 0.11 M in sodium lactate (b) Calculate the pH of a buffer formed by mixing 85 mL of 0.13 M lactic acid with 95 mL of 0.15M sodium lactate.

Problem 22E Buffers (Section) (a) Calculate the pH of a buffer that is 0.105 M in NaHCO3 and 0.125 M in Na2CO3 (b) Calculate the pH of a solution formed by mixing 65 mL of 0.20 M NaHCO3 with 75 mL of 0.15 MNa2CO3.

Problem 23E Buffers (Section) A buffer is prepared by adding 20.0 g of sodium acetate (CH3COONa) to 500 mL of a 0.150 Macetic acid (CH3COOH) solution (a) Determine the pH of the buffer (b) Write the complete ionic equation for the reaction that occurs when a few drops of hydrochloric acid are added to the buffer (c) Write the complete ionic equation for the reaction that occurs when a few drops of sodium hydroxide solution are added to the buffer.

Problem 25E Buffers (Section) You are asked to prepare a pH = 3.00 buffer solution starting from 1.25 L of a 1.00 M solution of hydrofluoric acid (HF) and any amount you need of sodium fluoride (NaF) (a) What is the pH of the hydrofluoric acid solution prior to adding sodium fluoride? (b) How many grams of sodium fluoride should be added to prepare the buffer solution? Neglect the small volume change that occurs when the sodium fluoride is added.

Problem 24E Buffers (Section) A buffer is prepared by adding 10.0 g of ammonium chloride (NH4Cl) to 250 mL of 1.00 M NH3 solution (a) What is the pH of this buffer? (b) Write the complete ionic equation for the reaction that occurs when a few drops of nitric acid are added to the buffer (c) Write the complete ionic equation for the reaction that occurs when a few drops of potassium hydroxide solution are added to the buffer.

Problem 27E Buffers (Section) A buffer contains 0.10 mol of acetic acid and 0.13 mol of sodium acetate in 1.00 L (a) What is the pH of this buffer? (b) What is the pH of the buffer after the addition of 0.02 mol of KOH? (c)What is the pH of the buffer after the addition of 0.02 mol of HNO3?

Problem 26E Buffers (Section) You are asked to prepare a pH = 4.00 buffer starting from 1.50 L of 0.0200 M solution of benzoic acid (C6H5COOH) and any amount you need of sodium benzoate (C6H5COONa) (a) What is the pH of the benzoic acid solution prior to adding sodium benzoate? (b) How many grams of sodium benzoate should be added to prepare the buffer? Neglect the small volume change that occurs when the sodium benzoate is added.

Problem 28E Buffers (Section) A buffer contains 0.15 mol of propionic acid (C2H5COOH) and 0.10 mol of sodium propionate (C2H5COONa) in 1.20 L (a) What is the pH of this buffer? (b) What is the pH of the buffer after the addition of 0.01 mol of NaOH? (c) What is the pH of the buffer after the addition of 0.01 mol of HI?

Problem 30E A buffer, consisting of H2PO4–and HPO42–, helps control the pH of physiological fluids. Many carbonated soft drinks also use this buffer system. What is the pH of a soft drink in which the major buffer ingredients are 6.5 g of NaH2PO4 and 8.0 g of Na2HPO4 per 355 mL of solution?

Problem 29E Buffers (Section) (a) What is the ratio of HCO3–to H2CO3 in blood of pH 7.4? (b) What is the ratio of HCO3–to H2CO3 in an exhausted marathon runner whose blood pH is 7.1?

You have to prepare a pH 3.50 buffer, and you have the following 0.10 M solutions available: HCOOH, \(\mathrm{CH}_{3} \mathrm{COOH}\), \(\mathrm{H}_{3} \mathrm{PO}_{4}\), HCOONa, \(\mathrm{CH}_{3} \mathrm{COONa}\), and \(\mathrm{NaH}_{2} \mathrm{PO}_{4}\)4. Which solutions would you use? How many milliliters of each solution would you use to make approximately a liter of the buffer?

Problem 32E Buffers (Section) You have to prepare a pH 5.00 buffer, and you have the following 0.10 M solutions available: HCOOH, HCOONa, CH3COOH, CH3COONa, HCN, and NaCN. Which solutions would you use? How many milliliters of each solution would you use to make approximately 1 L of the buffer?

Problem 32PE Buffers (Section) You have to prepare a pH 5.00 buffer, and you have the following 0.10 M solutions available: HCOOH, HCOONa, CH3COOH, CH3COONa, HCN, and NaCN. Which solutions would you use? How many milliliters of each solution would you use to make approximately 1 L of the buffer?

The accompanying graph shows the titration curves for two monoprotic acids. (a) Which curve is that of a strong acid? (b) What is the approximate pH at the equivalence point of each titration? (c) 40.0 mL of each acid was titrated with 0.100 M base. Which acid is more concentrated?

Problem 34E How does titration of a strong, monoprotic acid with a strong base differ from titration of a weak, monoprotic acid with a strong base with respect to the following: (a) quantity of base required to reach the equivalence point, (b) pH at the beginning of the titration, (c) pH at the equivalence point, (d) pH after addition of a slight excess of base, (e) choice of indicator for determining the equivalence point?

The samples of nitric and acetic acid shown here are both titrated with a 0.100 M solution of NaOH(aq).

Determine whether each of the following statements concerning the titrations in Problem 17.35 is true or false. 1. The pH at the beginning of the two titrations will be the same. 2. The titration curves will both be essentially the same after passing the equivalence point. 3. Methyl red would be a suitable indicator for both titrations.

Predict whether the equivalence point of each of the following titrations is below, above, or at pH 7: (a) \(\mathrm{NaHCO}_{3}\) titrated with NaOH, (b) \(\mathrm{NH}_{3}\) titrated with HCl, (c) KOH titrated with HBr.

Problem 38E Acid–Base Titrations (Section) Predict whether the equivalence point of each of the following titrations is below, above, or at pH 7: (a) formic acid titrated with NaOH (b) calcium hydroxide titrated with perchloric acid (c) pyridine titrated with nitric acid.

As shown in Figure 16.7, the indicator thymol blue has two color changes. Which color change will generally be more suitable for titration of a weak acid with a strong base?

Problem 40E Acid–Base Titrations (Section) Assume that 30.0 mL of a 0.10 M solution of a weak base B that accepts one proton is titrated with a 0.10 M solution of the monoprotic strong acid HA (a) How many moles of HA have been added at the equivalence point? (b) What is the predominant form of B at the equivalence point? (c) Is the pH 7, less than 7, or more than 7 at the equivalence point? (d) Which indicator, phenolphthalein or methyl red, is likely to be the better choice for this titration?

Problem 41E Acid–Base Titrations (Section) How many milliliters of 0.0850 M NaOH are required to titrate each of the following solutions to the equivalence point: (a) 40.0 mL of 0.0900 M HNO3 (b) 35.0 mL of 0.0850 M CH3COOH (c)50.0 mL of a solution that contains 1.85 g of HCl per liter?

Problem 42E Acid–Base Titrations (Section) How many milliliters of 0.105 M HCl are needed to titrate each of the following solutions to the equivalence point: (a) 45.0 mL of 0.0950 M NaOH (b) 22.5 mL of 0.118 M NH3 (c) 125.0 mL of a solution that contains 1.35 g of NaOH per liter?

Problem 43E Acid–Base Titrations (Section) A 20.0–mL sample of 0.200 M HBr solution is titrated with 0.200 M NaOH solution. Calculate the pH of the solution after the following volumes of base have been added: (a) 15.0 mL, (b) 19.9 mL (c) 20.0 mL (d) 20.1 mL (e) 35.0 mL.

Problem 44E Acid–Base Titrations (Section) A 20.0–mL sample of 0.150 M KOH is titrated with 0.125 M HClO4 solution. Calculate the pH after the following volumes of acid have been added: (a) 20.0 mL (b) 23.0 mL (c) 24.0 mL, (d) 25.0 mL (e) 30.0 mL.

Problem 45E Acid–Base Titrations (Section) A 35.0–mL sample of 0.150 M acetic acid (CH3COOH) is titrated with 0.150 M NaOH solution. Calculate the pH after the following volumes of base have been added: (a) 0 mL, (b) 17.5 mL (c) 34.5 mL (d) 35.0 mL (e) 35.5 mL (f) 50.0 mL.

Problem 46E Acid–Base Titrations (Section) Consider the titration of 30.0 mL of 0.050 M NH3 with 0.025 M HCl. Calculate the pH after the following volumes of titrant have been added: (a) 0 mL (b) 20.0 mL (c) 59.0 mL, (d) 60.0 mL (e) 61.0 mL (f) 65.0 mL.

Calculate the pH at the equivalence point in titrating 0.100 M solutions of each of the following with 0.080 M NaOH: (a) hydrobromic acid (HBr), (b) chlorous acid (\(HClO_2\)), (c) benzoic acid (\(C_6 H_5 COOH\)).

Problem 49E Solubility Equilibria and Factors Affecting Solubility (Sections) (a) Why is the concentration of undissolved solid not explicitly included in the expression for the solubility–product constant? (b) Write the expression for the solubility–product constant for each of the following strong electrolytes: AgI, SrSO4, Fe1OH22, and Hg2Br2

Problem 50E (a) Explain the difference between solubility and solubility product constant. (b) Write the expression for the solubility-product constant for each of the following ionic compounds:MnCO3, Hg(OH)2, and Cu3(PO4)2.

Problem 51E Solubility Equilibria and Factors Affecting Solubility (Sections) (a) If the molar solubility of CaF2 at 35 °C i s 1.24 × 10–3 mol/L, what is Ksp at this temperature? (b) It is found that 1.1 × 10–2 g SrF2 dissolves per 100 mL of aqueous solution at 25 °C. Calculate the solubility product for SrF2 (c) The Ksp of Ba1IO322 at 25 °C is 6.0 × 10–10. What is the molar solubility of Ba(IO3)2?

Problem 52E Solubility Equilibria and Factors Affecting Solubility (Sections) (a) The molar solubility of PbBr2 at 25 °C is 1.0 × 10–2 mol/L. Calculate Ksp(b) If 0.0490 g of AgIO3 dissolves per liter of solution, calculate the solubility–product constant (c) Using the appropriate Ksp value from Appendix D, calculate the pH of a saturated solution of Ca(OH)2.

Problem 53E Solubility Equilibria and Factors Affecting Solubility (Sections) A 1.00–L solution saturated at 25 °C with calcium oxalate (CaC2O4) contains 0.0061 g of CaC2O4. Calculate the solubility–product constant for this salt at 25 °C.

Problem 54E Solubility Equilibria and Factors Affecting Solubility (Sections) A 1.00–L solution saturated at 25 °C with lead(II) iodide contains 0.54 g of PbI2. Calculate the solubility–product constant for this salt at 25 °C.

Problem 56E Solubility Equilibria and Factors Affecting Solubility (Sections) Calculate the solubility of LaF3 in grams per liter in (a) pure water (b) 0.010 M KF solution (c) 0.050 M LaCl3 solution.

Problem 58E Solubility Equilibria and Factors Affecting Solubility (Sections) Consider a beaker containing a saturated solution of Pbl2 in equilibrium with undissolved Pbl2(s). Now solid KI is added to this solution (a) Will the amount of solid Pbl2 at the bottom of the beaker increase, decrease, or remain the same? (b) Will the concentration of Pb2+ ions in solution increase or decrease? (c) Will the concentration of I– ions in solution increase or decrease?

Problem 57E Solubility Equilibria and Factors Affecting Solubility (Sections) Consider a beaker containing a saturated solution of CaF2 in equilibrium with undissolved CaF2(s). Solid CaCl2 is then added to the solution (a) Will the amount of solid CaF2 at the bottom of the beaker increase, decrease, or remain the same? (b) Will the concentration of Ca2+ ions in solution increase or decrease? (c) Will the concentration of F– ions in solution increase or decrease?

Problem 59E Solubility Equilibria and Factors Affecting Solubility (Sections) Calculate the solubility of Mn(OH)2 in grams per liter when buffered at pH (a) 7.0 (b) 9.5 (c)11.8.

Problem 60E Solubility Equilibria and Factors Affecting Solubility (Sections) Calculate the molar solubility of Ni(OH)2 when buffered at pH (a) 8.0 (b) 10.0 (c) 12.0.

Problem 61E Solubility Equilibria and Factors Affecting Solubility (Sections) Which of the following salts will be substantially more soluble in acidic solution than in pure water: (a) ZnCO3 (b) ZnS (c) BiI3 (d) AgCN (e) Ba3(PO4)2?

Problem 63E Solubility Equilibria and Factors Affecting Solubility (Sections) From the value of Kf listed in Table calculate the concentration of Ni2+ in 1.0 L of a solution that contains a total of 1 × 10–3 mol of nickel(II) ion and that is 0.20 M in NH3. Table Formation Constants for Some Metal Complex Ions in Water at 25ºC

Problem 64E Solubility Equilibria and Factors Affecting Solubility (Sections) To what final concentration of NH3 must a solution be adjusted to just dissolve 0.020 mol of NiC2O4 (Ksp = 4 × 10–10) in 1.0 L of solution? (Hint: You can neglect the hydrolysis of C2O42–because the solution will be quite basic.)

Problem 62E Solubility Equilibria and Factors Affecting Solubility (Sections) For each of the following slightly soluble salts, write the net ionic equation, if any, for reaction with a strong acid: (a) MnS, (b) PbF2 (c) AuCl3 (d) Hg2C2O4 (e) CuBr.

Problem 65E Solubility Equilibria and Factors Affecting Solubility (Sections) Use values of Ksp for AgI and Kf for Ag(CN)2–to (a) calculate the molar solubility of AgI in pure water (b) calculate the equilibrium constant for the reaction AgI(s)+2 CN–(aq) ? Ag(CN)2–(aq) + I–(aq) (c) determine the molar solubility of AgI in a 0.100 M NaCN solution.

Problem 66E Solubility Equilibria and Factors Affecting Solubility (Sections) Using the value of Ksp for Ag2S, Ka1 and Ka2 for H2S, and Kf= 1.1 × 105 for AgCl2–, calculate the equilibrium constant for the following reaction: Ag2S(s) + 4 Cl–(aq) + 2 H+(aq) ?2 AgCl2–(aq) + H2S(aq)

Problem 67E Precipitation and Separation of Ions (Section) (a) Will Ca(OH)2 precipitate from solution if the pH of a 0.050 M solution of CaCl2 is adjusted to 8.0? (b) Will Ag2SO4 precipitate when 100 mL of 0.050 M AgNO3 is mixed with 10 mL of 5.0 × 10–2 M Na2SO4 solution?

Problem 68E Precipitation and Separation of Ions (Section) (a) Will Co(OH)2 precipitate from solution if the pH of a 0.020 M solution of Co(NO3)2 is adjusted to 8.5? (b) Will AgIO3 precipitate when 20 mL of 0.010 M AgIO3 is mixed with 10 mL of 0.015 M NaIO3? (Ksp of AgIO3 is 3.1 × 10–8).

Calculate the minimum pH needed to precipitate \(Mn(OH)_2\) so completely that the concentration of \(Mn^{2+}\) is less than 1 \(\mu g\) per liter [1 part per billion (ppb)].

Problem 70E Precipitation and Separation of Ions (Section) Suppose that a 10–mL sample of a solution is to be tested for I–ion by addition of 1 drop (0.2 mL) of 0.10 M Pb(NO3)2. What is the minimum number of grams of I– that must be present for Pbl2(s) to form?

A solution contains \(2.0 \times 10^{-4} \mathrm{M} \mathrm{Ag}^{+}\) and \(1.5 \times 10^{-3} \mathrm{M}\) \(\mathrm{Pb}^{2+}\). If \(\mathrm{NaI}\) is added, will \(\mathrm{AgI}\left(K_{s p}=8.3 \times 10^{-17}\right)\) or \(\mathrm{PbI}_2\) \(\left(K_{s p}=7.9 \times 10^{-9}\right)\) precipitate first? Specify the concentration of \(\mathrm{l}^{-}\) needed to begin precipitation.

Problem 72E Precipitation and Separation of Ions (Section) A solution of Na2SO4 is added dropwise to a solution that is 0.010 M in Ba2+ and 0.010 M in Sr2+ (a) What concentration of SO42– is necessary to begin precipitation? (Neglect volume changes. BaSO4: Ksp = 1.1 × 10–10; SrSO4: Ksp = 3.2 * 10–7.) (b) Which cation precipitates first? (c) What is the concentration of SO42–when the second cation begins to precipitate?

A 1.0 M \(Na2SO_4\) solution is slowly added to 10.0 mL of a solution that is 0.20 M in \(Ca^{2+}\) and 0.30 M in \(Ag^+\) (a) Which compound will precipitate first: \(CaSO_4\) \((K_{sp} = 2.4 \times 10^{–5})\) or \(Ag_{2} SO_{4}\) \((K_{sp} = 1.5 \times 10^{–5})\)? (b) How much \(Na2SO_4\) solution must be added to initiate the precipitation?

Problem 73E Precipitation and Separation of Ions (Section) A solution contains three anions with the following concentrations: 0.20 M CrO42–, 0.10 M CO32–, and 0.010 M Cl–. If a dilute AgNO3 solution is slowly added to the solution, what is the first compound to precipitate: Ag2CrO4 (Ksp = 1.2 × 10–12), Ag2CO3(Ksp = 8.1 × 10–12), or AgCl (Ksp = 1.8 × 10–10)?

Problem 75E Qualitative Analysis for Metallic Elements (Section) A solution containing an unknown number of metal ions is treated with dilute HCl; no precipitate forms. The pH is adjusted to about 1, and H2S is bubbled through. Again, no precipitate forms. The pH of the solution is then adjusted to about 8. Again, H2S is bubbled through. This time a precipitate forms. The filtrate from this solution is treated with (NH4)2HPO4. No precipitate forms. Which metal ions discussed in Section are possibly present? Which are definitely absent within the limits of these tests?

Problem 77E Qualitative Analysis for Metallic Elements (Section) In the course of various qualitative analysis procedures, the following mixtures are encountered:(a) Zn2+ and Cd2+ , (b) Cr(OH)3 and Fe(OH)3 (c) Mg2+ and K+ (d) Ag+ and Mn2+ . Suggest how each mixture might be separated.

An unknown solid is entirely soluble in water. On addition of dilute HCl, a precipitate forms. After the precipitate is filtered off, the is adjusted to about 1 and is bubbled in; a precipitate again forms. After filtering off this precipitate, the is adjusted to 8 and is again added; no precipitate forms. No precipitate forms upon addition of . The remaining solution shows a yellow color in a flame test (see Figure 7.21). Based on these observations, which of the following compounds might be present, which are definitely present, and which are definitely absent: , , and

Problem 78E Qualitative Analysis for Metallic Elements (Section) Suggest how the cations in each of the following solution mixtures can be separated: (a) Na+and Cd2+ (b) Cu2+ and Mg2+ (c) Pb2+ and Al3+ (d) Ag+ and Hg2+

(a) Precipitation of the group 4 cations of Figure 17.23 requires a basic medium. Why is this so? (b) What is the most significant difference between the sulfides precipitated in group 2 and those precipitated in group 3? (c) Suggest a procedure that would serve to redissolve the group 3 cations following their precipitation.

A student who is in a great hurry to finish his laboratory work decides that his qualitative analysis unknown contains a metalion from group 4 of Figure 17.23. He therefore tests his sample directly with , skipping earlier tests for the metalions in groups 1, 2, and 3. He observes a precipitate and concludes that a metal ion from group 4 is indeed present. Why is this possibly an erroneous conclusion?

Problem 81AE Derive an equation similar to the Henderson–Hasselbalch equation relating the pOH of a buffer to the pKb of its base component.

Problem 82AE Benzenesulfonic acid is a monoprotic acid with pKa = 2.25. Calculate the pH of a buffer composed of 0.150 Mbenzenesulfonic acid and 0.125 M sodium benzenesulfonate.

Problem 83AE Furoic acid (HC5H3O3) has a Ka value of 6.76 × 10–4 at 25 °C. Calculate the pH at 25 °C of (a) a solution formed by adding 25.0 g of furoic acid and 30.0 g of sodium furoate (NaC5H3O3) to enough water to form 0.250 L of solution (b) a solution formed by mixing 30.0 mL of 0.250 MHC5H3O3 and 20.0 mL of 0.22 M NaC5H3O3 and diluting the total volume to 125 mL (c) a solution prepared by adding 50.0 mL of 1.65 M NaOH solution to 0.500 L of 0.0850 MHC5H3O3.

Problem 84AE The acid–base indicator bromcresol green is a weak acid. The yellow acid and blue base forms of the indicator are present in equal concentrations in a solution when the pH is 4.68. What is the pKa for bromcresol green?

Problem 86AE Two buffers are prepared by adding an equal number of moles of formic acid (HCOOH) and sodium formate (HCOONa) to enough water to make 1.00 L of solution. Buffer A is prepared using 1.00 mol each of formic acid and sodium formate. Buffer B is prepared by using 0.010 mol of each. (a) Calculate the pH of each buffer, and explain why they are equal. (b) Which buffer will have the greater buffer capacity? Explain. (c) Calculate the change in pH for each buffer upon the addition of 1.0 mL of 1.00 M HCl. (d) Calculate the change in pH for each buffer upon the addition of 10 mL of 1.00 M HCl. (e) Discuss your answers for parts (c) and (d) in light of your response to part (b).

Problem 87AE A biochemist needs 750 mL of an acetic acid–sodium acetate buffer with pH 4.50. Solid sodium acetate (CH3COONa) and glacial acetic acid (CH3COOH) are available. Glacial acetic acid is 99% CH3COOH by mass and has a density of 1.05 g/mL. If the buffer is to be 0.15 M in CH3COOH, how many grams of CH3COONa and how many milliliters of glacial acetic acid must be used?

A sample of 0.2140 g of an unknown monoprotic acid was dissolved in 25.0 mL of water and titrated with 0.0950 M NaOH. The acid required 27.4 mL of base to reach the equivalence point (a) What is the molar mass of the acid? (b) After 15.0 mL of base had been added in the titration, the pH was found to be 6.50. What is the \(K_a\) for the unknown acid?

Problem 89AE A sample of 0.1687 g of an unknown monoprotic acid was dissolved in 25.0 mL of water and titrated with 0.1150 M NaOH. The acid required 15.5 mL of base to reach the equivalence point (a) What is the molecular weight of the acid? (b) After 7.25 mL of base had been added in the titration, the pH was found to be 2.85. What is the Ka for the unknown acid?

Problem 90AE Show that the pH at the halfway point of a titration of a weak acid with a strong base (where the volume of added base is half of that needed to reach the equivalence point) is equal to pKa for the acid.

Problem 91AE A hypothetical weak acid, HA, was combined with NaOH in the following proportions: 0.20 mol of HA, 0.080 mol of NaOH. The mixture was diluted to a total volume of 1.0 L and the pH measured. (a) If pH = 4.80, what is the pKa of the acid? (b) How many additional moles of NaOH should be added to the solution to increase the pH to 5.00?

Problem 92AE What is the pH of a solution made by mixing 0.30 mol NaOH, 0.25 mol Na2HPO4, and 0.20 mol H3PO4 with water and diluting to 1.00 L?

Problem 93AE Suppose you want to do a physiological experiment that calls for a pH 6.50 buffer. You find that the organism with which you are working is not sensitive to the weak acid H2A (Ka1 = 2 × 10-2;Ka2 = 5.0 × 10-7) or its sodium salts. You have available a 1.0 M solution of this acid and a 1.0M solution of NaOH. How much of the NaOH solution should be added to 1.0 L of the acid to give a buffer at pH 6.50? (Ignore any volume change.)

Problem 94AE How many microliters of 1.000 M NaOH solution must be added to 25.00 mL of a 0.1000 Msolution of lactic acid [CH3CH(OH)COOH or HC3H5O3] to produce a buffer with pH = 3.75?

Problem 95AE A person suffering from anxiety begins breathing rapidly and as a result suffers alkalosis, an increase in blood pH (a) Using Equation 17.10, explain how rapid breathing can cause the pH of blood to increase (b) One cure for this problem is breathing in a paper bag. Why does this procedure lower blood pH?

Problem 96AE For each pair of compounds, use Ksp values to determine which has the greater molar solubility: (a) CdS or CuS (b) PbCO3 or BaCrO4 (c) Ni(OH)2 or NiCO3 (d) AgI or Ag2SO4

The solubility of \(\mathrm{CaCO}_{3}\) is pH dependent. (a) Calculate the molar solubility of \(\mathrm{CaCO}_3\ \left(K_{sp}=4.5\times10^{-9}\right)\) neglecting the acid–base character of the carbonate ion. (b) Use the \(K_{b}\) expression for the \(\mathrm{CO}_3^{\ 2-}\) ion to determine the equilibrium constant for the reaction \(\mathrm{CaCO}_3(s)+\mathrm{H}_2\mathrm{O}(l)\rightleftharpoons\mathrm{Ca}^{2+}(aq)+\mathrm{HCO}_3^{\ -}(aq)+\mathrm{OH}^-(aq)\) (c) If we assume that the only sources of \(\mathrm{Ca}^{2+}\), \(\mathrm{HCO}_3^{\ -}\), and \(\mathrm{OH}^{-}\) ions are from the dissolution of \(\mathrm{CaCO}_{3}\), what is the molar solubility of \(\mathrm{CaCO}_{3}\) using the preceding expression? What is the pH? (d) If the pH is buffered at 8.2 (as is historically typical for the ocean), what is the molar solubility of \(\mathrm{CaCO}_{3}\)? (e) If the pH is buffered at 7.5, what is the molar solubility of \(\mathrm{CaCO}_{3}\)? How much does this drop in pH increase solubility? Equation Transcription: ? Text Transcription: CaCO_3 CaCO_3 (Ksp = 4.5 x 10^-9) K_b CO_3 ^2- CaCO_3(s) + H_2 O(l) ? Ca^2+ (aq) + HCO_3 ^- (aq) + OH^- (aq) Ca^2+ HCO_3 ^- OH^-

Problem 98AE Tooth enamel is composed of hydroxyapatite, whose simplest formula is Ca5(PO4)3OH, and whose corresponding Ksp = 6.8 × 10–27. As discussed in the “Chemistry and Life” box on page 755, fluoride in fluorinated water or in toothpaste reacts with hydroxyapatite to form fluoroapatite, Ca5(PO4)3F, whose Ksp = 1.0 × 10–60 (a) Write the expression for the solubility-constant for hydroxyapatite and for fluoroapatite (b) Calculate the molar solubility of each of these compounds

Problem 99AE Use the solubility-product constant for Cr(OH)3(Ksp = 6.7 × 10–31) and the formation constant for Cr(OH)4– from Table to determine the concentration of Cr(OH)4– in a solution that is buffered at pH = 10.0 and is in equilibrium with solid Cr(OH)3. Table Formation Constants for Some Metal Complex Ions in Water at 25ºC

Problem 100AE Calculate the solubility of Mg(OH)2 in 0.50 M NH4Cl.

Problem 101AE The solubility-product constant for barium permanganate, Ba(MnO4)2, is 2.5 × 10–10. Assume that solid Ba(MnO4)2 is in equilibrium with a solution of KMnO4. What concentration of KMnO4is required to establish a concentration of 2.0 × 10–8 M for the Ba2+ion in solution?

Problem 102AE Calculate the ratio of [Ca2+] to [Fe2+] in a lake in which the water is in equilibrium with deposits of both CaCO3 and FeCO3. Assume that the water is slightly basic and that the hydrolysis of the carbonate ion can therefore be ignored.

The solubility product constants of \(PbSO_4\) and \(SrSO_4\) are \(6.3 \times 10^{–7}\) and \(3.2 \times 10^{–7}\), respectively. What are the values of [\(SO_{4}^{2–}\)], [\(Pb^{2+}\)], and [\(Sr^{2+}\)] in a solution at equilibrium with both substances?

Problem 104AE A buffer of what pH is needed to give a Mg2+ concentration of 3.0 × 10–2 M in equilibrium with solid magnesium oxalate?

Problem 105AE The value of Ksp for Mg3(AsO4)2 is 2.1 × 10–20. The AsO4 3–ion is derived from the weak acid H3AsO4 (pKa1 = 2.22; pKa2 = 6.98; pKa3 = 11.50). When asked to calculate the molar solubility of Mg3(AsO4)2 in water, a student used the Ksp expression and assumed that [Mg2+] = 1.5[AsO4 3–]. Why was this a mistake?

Problem 106AE The solubility product for Zn(OH)2 is 3.0 × 10–16. The formation constant for the hydroxo complex, Zn1OH242–, is 4.6 × 1017. What concentration of OH– is required to dissolve 0.015 mol of Zn(OH)2 in a liter of solution?

Problem 107AE The value of Ksp for Cd(OH)2 is 2.5 × 10–14. (a) What is the molar solubility of Cd(OH)2? (b) The solubility of Cd(OH)2 can be increased through formation of the complex ion CdBr42–(Kf= 5 × 103). If solid Cd(OH)2 is added to a NaBr solution, what is the initial concentration of NaBr needed to increase the molar solubility of Cd(OH)2 to 1.0 × 10–3 mol/L?

(a) Write the net ionic equation for the reaction that occurs when a solution of hydrochloric acid (HCl) is mixed with a solution of sodium formate (NaCHO2) (b) Calculate the equilibrium constant for this reaction (c) Calculate the equilibrium concentrations of Na+, Cl–, H+, CHO2–, and HCHO2 when 50.0 mL of 0.15 M HCl is mixed with 50.0 mL of 0.15 M NaCHO2.

Problem 110IE A sample of 7.5 L of NH3 gas at 22 °C and 735 torr is bubbled into a 0.50-L solution of 0.40 MHCl. Assuming that all the NH3 dissolves and that the volume of the solution remains 0.50 L, calculate the pH of the resulting solution.

Problem 109IE (a) A 0.1044-g sample of an unknown monoprotic acid requires 22.10 mL of 0.0500 M NaOH to reach the end point. What is the molecular weight of the unknown? (b) As the acid is titrated, the pH of the solution after the addition of 11.05 mL of the base is 4.89. What is the Ka for the acid? (c) Using Appendix D, suggest the identity of the acid.

Aspirin has the structural formula At body temperature (37°C), Ka for aspirin equals 3x10-5. If two aspirin tablets, each having a mass of 325 mg, are dissolved in a full stomach whose volume is 1 L and whose pH is 2, what percent of the aspirin is in the form of neutral molecules?

What is the at of water saturated with at a partial pressure of ? The Henry's law constant for at is . The is an acidic oxide, reacting with to form

Excess is shaken with water to produce a saturated solution. The solution is filtered, and a sample titrated with requires of to your result with that in Appendix D. Do you think the solution was kept at ?

Problem 114IE The osmotic pressure of a saturated solution of strontium sulfate at 25 °C is 21 torr. What is the solubility product of this salt at 25 °C?

Problem 115IE A concentration of 10–100 parts per billion (by mass) of Ag+ is an effective disinfectant in swimming pools. However, if the concentration exceeds this range, the Ag+ can cause adverse health effects. One way to maintain an appropriate concentration of Ag+ is to add a slightly soluble salt to the pool. Using Ksp values from Appendix D, calculate the equilibrium concentration of Ag+ in parts per billion that would exist in equilibrium with (a) AgCl (b) AgBr (c) AgI.

Problem 116IE Fluoridation of drinking water is employed in many places to aid in the prevention of tooth decay. Typically the F- ion concentration is adjusted to about 1 ppb. Some water supplies are also “hard”; that is, they contain certain cations such as Ca2+ that interfere with the action of soap. Consider a case where the concentration of Ca2+ is 8 ppb. Could a precipitate of CaF2form under these conditions? (Make any necessary approximations.)