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More Equilibria in Aqueous Solutions: Slightly Soluble Salts and Complex Ions

Chapter Sixteen. More Equilibria in Aqueous Solutions: Slightly Soluble Salts and Complex Ions. BaSO 4 (s) Ba 2+ (aq) + SO 4 2– (aq). The Solubility Product Constant, K sp.

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More Equilibria in Aqueous Solutions: Slightly Soluble Salts and Complex Ions

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  1. Chapter Sixteen More Equilibria in Aqueous Solutions:Slightly Soluble Salts and Complex Ions

  2. BaSO4(s) Ba2+(aq) + SO42–(aq) The Solubility Product Constant, Ksp • Solubility product constant, Ksp: the equilibrium constant expression for the dissolving of a slightly soluble solid. • Many important ionic compounds are only slightly soluble in water (we used to call them “insoluble” – Chapter 4). • An equation can represent the equilibrium between the compound and the ions present in a saturated aqueous solution: Ksp= [ Ba2+ ][ SO42–]

  3. Example 16.1 Write a solubility product constant expression for equilibrium in a saturated aqueous solution of the slightly soluble salts (a) iron(III) phosphate, FePO4, and (b) chromium(III) hydroxide, Cr(OH)3.

  4. Ksp and Molar Solubility • Ksp is an equilibrium constant • Molar solubility is the number of moles of compound that will dissolve per liter of solution. • Molar solubility is related to the value of Ksp, but molar solubility and Ksp are not the same thing. • In fact, “smaller Ksp” doesn’t always mean “lower molar solubility.” • Solubility depends on both Ksp and the form of the equilibrium constant expression.

  5. Example 16.2 At 20 °C, a saturated aqueous solution of silver carbonate contains 32 mg of Ag2CO3 per liter of solution. Calculate Kspfor Ag2CO3 at 20 °C. The balanced equation is Ag2CO3(s) 2 Ag+(aq) + CO32–(aq) Ksp= ? Example 16.3 From the Ksp value for silver sulfate, calculate its molar solubility at 25 °C. Ag2SO4(s) 2 Ag+(aq) + SO42–(aq) Ksp= 1.4 x 10–5 at 25 °C

  6. Example 16.4 A Conceptual Example Without doing detailed calculations, but using data from Table 16.1, establish the order of increasing solubility of these silver halides in water: AgCl, AgBr, AgI.

  7. The Common Ion Effectin Solubility Equilibria • The common ion effect affects solubility equilibria as it does other aqueous equilibria. • The solubility of a slightly soluble ionic compound is lowered when a second solute that furnishes a common ion is added to the solution.

  8. Common Ion Effect Illustrated The added sulfate ion reduces the solubility of Ag2SO4. Na2SO4(aq) Saturated Ag2SO4(aq) Ag2SO4 precipitates

  9. Common Ion Effect Illustrated When Na2SO4(aq) is added to the saturated solution of Ag2SO4 … … [Ag+] attains a new, lower equilibrium concentration as Ag+ reacts with SO42–to produce Ag2SO4.

  10. Example 16.5 Calculate the molar solubility of Ag2SO4 in 1.00 M Na2SO4(aq).

  11. Solubility and Activities • Ions that are not common to the precipitate can also affect solubility. • CaF2 is more soluble in 0.010 M Na2SO4 than it is in water. • Increased solubility occurs because of interionic attractions. • Each Ca2+ and F– is surrounded by ions of opposite charge, which impede the reaction of Ca2+ with F–. • The effective concentrations, or activities, of Ca2+ and F– are lower than their actual concentrations.

  12. Will Precipitation Occur? Is It Complete? • Qip can then be compared to Ksp. • Precipitation should occur if Qip > Ksp. • Precipitation cannot occur if Qip < Ksp. • A solution is just saturated if Qip = Ksp. • In applying the precipitation criteria, the effect of dilution when solutions are mixed must be considered. • Qip is the ion product reaction quotient and is based on initial conditions of the reaction. Qip andQc: new look, same great taste!

  13. Example 16.6 If 1.00 mg of Na2CrO4 is added to 225 mL of 0.00015 M AgNO3, will a precipitate form? Ag2CrO4(s) 2 Ag+(aq) + CrO42–(aq) Ksp= 1.1 x 10–12

  14. Example 16.7 A Conceptual Example Pictured here is the result of adding a few drops of concentrated KI(aq) to a dilute solution of Pb(NO3)2. What is the solid that first appears? Explain why it then disappears. Example 16.8 If 0.100 L of 0.0015 M MgCl2 and 0.200 L of 0.025 M NaF are mixed, should a precipitate of MgF2 form? MgF2(s) Mg2+(aq) + 2 F–(aq) Ksp= 3.7 x 10–8

  15. To Determine Whether Precipitation Is Complete • A slightly soluble solid does not precipitate totally from solution … • … but we generally consider precipitation to be “complete” if about 99.9% of the target ion is precipitated (0.1% or less left in solution). • Three conditions generally favor completeness of precipitation: • A very small value of Ksp. • A high initial concentration of the target ion. • A concentration of common ion that greatly exceeds that of the target ion.

  16. Example 16.9 To a solution with [Ca2+] = 0.0050 M, we add sufficient solid ammonium oxalate, (NH4)2C2O4(s), to make the initial [C2O42–] = 0.0051 M. Will precipitation of Ca2+ as CaC2O4(s) be complete? CaC2O4(s) Ca2+(aq) + C2O42–(aq) Ksp= 2.7 x 10–9

  17. Selective Precipitation AgNO3 added to a mixture containing Cl– and I–

  18. Example 16.10 An aqueous solution that is 2.00 M in AgNO3 is slowly added from a buret to an aqueous solution that is 0.0100 M in Cl– and also 0.0100 M in I–. • Which ion, Cl– or I–, is the first to precipitate from solution? • When the second ion begins to precipitate, what is the remaining concentration of the first ion? • Is separation of the two ions by selective precipitation feasible? AgCl(s) Ag+(aq) + Cl–(aq) Ksp= 1.8 x 10–10 AgI(s) Ag+(aq) + I–(aq) Ksp= 8.5 x 10–17

  19. CaF2(s) Ca2+(aq) + 2 F–(aq) AgCl(s) Ag+(aq) + Cl–(aq) Effect of pH on Solubility • If the anion of a precipitate is that of a weak acid, the precipitate will dissolve somewhat when the pH is lowered: Added H+ reacts with, and removes, F–; LeChâtelier’s principle says more F– forms. • If, however, the anion of the precipitate is that of a strong acid, lowering the pH will have no effect on the precipitate. H+ does not consume Cl– ; acid does not affect the equilibrium.

  20. Example 16.11 What is the molar solubility of Mg(OH)2(s) in a buffer solution having [OH–] = 1.0 x 10–5 M, that is, pH = 9.00? Mg(OH)2(s) Mg2+(aq) + 2 OH–(aq) Ksp= 1.8 x 10–11 Example 16.12 A Conceptual Example Without doing detailed calculations, determine in which of the following solutions Mg(OH)2(s) is most soluble: (a) 1.00 M NH3 (b) 1.00 M NH3 /1.00 M NH4+ (c) 1.00 M NH4Cl.

  21. Equilibria Involving Complex Ions Silver chloride becomes more soluble, not less soluble, in high concentrations of chloride ion.

  22. Ag+(aq) + 2 Cl–(aq) [AgCl2]–(aq) Complex Ion Formation • A complex ion consists of a central metal atom or ion, with other groups called ligands bonded to it. • The metal ion acts as a Lewis acid (accepts electron pairs). • Ligands act as Lewis bases (donate electron pairs). • The equilibrium involving a complex ion, the metal ion, and the ligands may be described through a formation constant, Kf: [AgCl2]– Kf = –––––––––– = 1.2 x 108 [Ag+][Cl–]2

  23. Complex Ion Formation Concentrated NH3 added to a solution of pale-blue Cu2+ … … forms deep-blue Cu(NH3)42+.

  24. Complex Ion Formationand Solubilities But if the concentration of NH3 is made high enough … … the AgCl forms the soluble [Ag(NH3)2]+ ion. AgCl is insoluble in water.

  25. Ag+(aq) + 2 NH3(aq) [Ag(NH3)2]+(aq) Kf= 1.6 x 107 AgBr(s) Ag+(aq) + Br–(aq) Example 16.13 Calculate the concentration of free silver ion, [Ag+], in an aqueous solution prepared as 0.10 M AgNO3 and 3.0 M NH3. Example 16.14 If 1.00 g KBr is added to 1.00 L of the solution described in Example 16.13, should any AgBr(s) precipitate from the solution? Ksp= 5.0 x 10–13

  26. Example 16.15 What is the molar solubility of AgBr(s) in 3.0 M NH3? AgBr(s) + 2 NH3(aq) [Ag(NH3)2]+(aq) + Br–(aq) Kc= 8.0 x 10–6 Example 16.16 A Conceptual Example Figure 16.10 shows that a precipitate forms when HNO3(aq) is added to the solution in the beaker on the right in Figure 16.9. Write the equation(s) to show what happens.

  27. Complex Ions in Acid–Base Reactions • Water molecules are commonly found as ligands in complex ions (H2O is a Lewis base). [Na(H2O)4]+ [Al(H2O)6]3+ [Fe(H2O)6]3+ • The electron-withdrawing power of a small, highlycharged metal ion can weaken an O—H bond in one of the ligand water molecules. • The weakened O—H bond can then give up its proton to another water molecule in the solution. • The complex ion acts as an acid.

  28. Ionization of a Complex Ion The highly-charged iron(III) ion withdraws electron density from the O—H bonds. [Fe(H2O)6]3+ + H2O [Fe(H2O)5OH]2+ + H3O+ Ka= 1 x 10–7

  29. Amphoteric Species • Certain metal hydroxides, insoluble in water, are amphoteric; they will react with both strong acids and strong bases. • Al(OH)3, Zn(OH)2, and Cr(OH)3 are amphoteric.

  30. Qualitative Inorganic Analysis • Acid–base chemistry, precipitation reactions, oxidation–reduction, and complex ion formation all apply to an area of analytical chemistry called classical qualitative inorganic analysis. • “Qualitative” signifies that the interest is in determining what is present. • Quantitative analyses are those that determine how much of a particular substance or species is present. • Although classical qualitative analysis is not used as widely today as are instrumental methods, it is still a good vehicle for applying all the basic concepts of equilibria in aqueous solutions.

  31. Qualitative Analysis Outline In acid, H2S produces very little S2–, so only the most-insoluble sulfides precipitate. In base, there is more S2–, and the less-insoluble sulfides also precipitate. Some hydroxides also precipitate here.

  32. Cation Group 1 • If aqueous HCl is added to an unknown solution of cations, and a precipitate forms, then the unknown contains one or more of these cations: Pb2+, Hg22+, or Ag+. • These are the only ions to form insoluble chlorides. • Any precipitate is separated from the mixture and further tests are performed to determine which of the three Group 1 cations are present. • The supernatant liquid is also saved for further analysis (it contains the rest of the cations). • If there is no precipitate, then Group 1 ions must be absent from the mixture.

  33. Cation Group 1 (cont’d)Analyzing for Pb2+ • Precipitated PbCl2 is slightly soluble in hot water. • The precipitate is washed with hot water, then aqueous K2CrO4 is added to the washings. • If Pb2+ is present, a precipitate of yellow lead chromate forms, which is less soluble than PbCl2. • (If all of the precipitate dissolves in the hot water, what does that mean?)

  34. Cation Group 1 (cont’d)Analyzing for Ag+ and Hg22+ • Next, any undissolved precipitate is treated with aqueous ammonia. • If AgCl is present, it will dissolve, forming Ag(NH3)2+ (the dissolution may not be visually apparent). • If Hg22+ is present, the precipitate will turn dark gray/ black, due to a disproportionation reaction that forms Hg metal and HgNH2Cl. • The supernatant liquid (which contains the Ag+, if present) is then treated with aqueous nitric acid. • If a precipitate reforms, then Ag+ was present in the solution.

  35. Group 1 Cation Precipitates PbCl2 precipitates when HCl is added. The presence of lead is confirmed by adding chromate ion; yellow PbCrO4 precipitates. Hg2Cl2 reacts with NH3 to form black Hg metal and HgNH2Cl.

  36. Hydrogen Sulfide in theQualitative Analysis Scheme • Once the Group 1 cations have been precipitated, hydrogen sulfide is used as the next reagent in the qualitative analysis scheme. • H2S is a weak diprotic acid; there is very little ionization of the HS– ion and it is the precipitating agent. • Hydrogen sulfide has the familiar rotten egg odor that is very noticeable around volcanic areas. • Because of its toxicity, H2S is generally produced only in small quantities and directly in the solution where it is to be used.

  37. Cation Groups 2, 3, 4, and 5 • The concentration of HS– is so low in a strongly acidic solution, that only the most insoluble sulfides precipitate. • These include the eight metal sulfides of Group 2. • Five of the Group 3 cations form sulfides that are soluble in acidic solution but insoluble in alkaline NH3/NH4+. • The other three Group 3 cations form insoluble hydroxides in the alkaline solution. • The cations of Groups 4 and 5 are soluble. • Group 4 ions are precipitated as carbonates. • Group 5 does not precipitate; these must be determined by flame test.

  38. Cumulative Example A solid mixture containing 1.00 g of ammonium chloride and 2.00 g of barium hydroxide is heated to expel ammonia. The liberated NH3 is then dissolved in 0.500 L of water containing 225 ppm Ca2+ as calcium chloride. Will a precipitate form in this water?

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