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Problem 46e,f
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: (e) 61.0 mL (f) 65.0 mL.
Problem 47b
Calculate the pH at the equivalence point for titrating 0.200 M solutions of each of the following bases with 0.200 M HBr: (b) hydroxylamine 1NH2OH2.
Problem 48b
Calculate the pH at the equivalence point in titrating 0.100 M solutions of each of the following with 0.080 M NaOH: (b) chlorous acid (HClO2).
Problem 48c
Calculate the pH at the equivalence point in titrating 0.100 M solutions of each of the following with 0.080 M NaOH: (c) benzoic acid (C6H5COOH).
Problem 50c
The solubility of two slightly soluble salts of M2 + , MA and MZ2, is the same, 4 * 10-4 mol/L. (c) If you added an equal volume of a solution saturated in MA to one saturated in MZ2, what would be the equilibrium concentration of the cation, M2+?
Problem 52a
(a) True or false: 'solubility' and 'solubility-product constant' are the same number for a given compound.
Problem 52b
(b) Write the expression for the solubility-product constant for each of the following ionic compounds: MnCO3, Hg(OH)2, and Cu3(PO4)3.
Problem 53a
(a) I f t he molar solubility of CaF2 at 35°C i s 1.24 × 10–3 mol/L, what is Ksp at this temperature?
Problem 53b
(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.
Problem 54b
(b) If 0.0490 g of AgIO3 dissolves per liter of solution, calculate the solubility-product constant.
Problem 55
A 1.00-L solution saturated at 25 C with calcium oxalate 1CaC2O42 contains 0.0061 g of CaC2O4. Calculate the solubility-product constant for this salt at 25 C.
Problem 56
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.
- Using Appendix D, calculate the molar solubility of AgBr in (b) 3.0 × 10^-2 M AgNO3 solution and (c) 0.10 M NaBr solution.
Problem 57
Problem 58a
Calculate the solubility of LaF3 in grams per liter in (a) pure water.
Problem 58b
Calculate the solubility of LaF3 in grams per liter in (b) 0.010 M KF solution.
Problem 58c
Calculate the solubility of LaF3 in grams per liter in (c) 0.050 M LaCl3 solution.
Problem 59a
Consider a beaker containing a saturated solution of CaF2 in equilibrium with undissolved CaF21s2. 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?
Problem 61
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 64
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 (e) CuBr (d) Hg2C2O4.
- From the value of Kf listed in Table 17.1, calculate the concentration of NH3 required to just dissolve 0.020 mol of NiC2O4 (Ksp = 4 * 10^-102) in 1.00 L of solution? (Hint: You can neglect the hydrolysis of C2O4^2- because the solution will be quite basic.)
Problem 66
- 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 67
- Using the value of Ksp for Ag2S, Ka1 and Ka2 for H2S, and Kf = 1.1 * 10^5 for AgCl2^-, calculate the equilibrium constant for the following reaction: Ag2S (s) + 4 Cl^- (aq) + 2 H^+ (aq) ⇌ 2 AgCl2^- (aq) + H2S (aq)
Problem 68
- (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 69
Problem 72
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 Pb1NO322. What is the minimum number of grams of I- that must be present for PbI21s2 to form?
- A solution contains 2.0 * 10^-4 M Ag^+ (aq) and 1.5 * 10^-3 M Pb^2+ (aq). If NaI is added, will AgI (Ksp = 8.3 * 10^-17) or PbI2 (Ksp = 7.9 * 10^-9) precipitate first? Specify the concentration of I^- (aq) needed to begin precipitation.
Problem 73
Problem 74a
A solution of Na2SO4 is added dropwise to a solution that is 0.010 M in Ba2+(aq) and 0.010 M in Sr2+(aq). (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.)
Problem 74c
A solution of Na2SO4 is added dropwise to a solution that is 0.010 M in Ba2+(aq) and 0.010 M in Sr2+(aq). (c) What is the concentration of SO42-(aq) when the second cation begins to precipitate?
- A solution contains three anions with the following concentrations: 0.20 M CrO4^2-, 0.10 M CO3^2-, 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 75
- A 1.0 M Na2SO4 solution is slowly added to 10.0 mL of a solution that is 0.20 M in Ca2+ and 0.30 M in Ag+. (a) Which compound will precipitate first: CaSO4 (Ksp = 2.4 * 10^-5) or Ag2SO4 (Ksp = 1.5 * 10^-5)?
Problem 76
Problem 77
A solution containing several 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 of these metal cations are either possibly present or definitely absent: Al3+, Na+, Ag+, Mg2+?