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Chapter 18 Solubility and Complex-Ion Equilibria

Chapter 18 Solubility and Complex-Ion Equilibria. Contents in Chapter 18. 18-1 Solubility Product Constant, K sp 18-2 Relationship Between Solubility and K sp 18-3 Common-Ion Effect in Solubility Equilibria 18-4 Limitations of the K sp Concept

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Chapter 18 Solubility and Complex-Ion Equilibria

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  1. Chapter 18Solubility and Complex-Ion Equilibria

  2. Contents in Chapter 18 18-1 Solubility Product Constant, Ksp 18-2 Relationship Between Solubility and Ksp 18-3 Common-Ion Effect in Solubility Equilibria 18-4 Limitations of the Ksp Concept 18-5 Criteria for Precipitation and Its Completeness 18-6 Fractional Precipitation 18-7 Solubility and pH 18-8 Equilibria Involving Complex Ions 18-9 Qualitative Cation Analysis

  3. 18-1 Solubility Product Constant, Ksp • Solubility: The number of grams of solute in one liter of a saturated solution (g/L). • Molar solubility (s): The number of moles of solute in one liter of a saturated solution (M). • Solubility product constant (Ksp): The product of the molar concentrations of the constituent ions: • AmBn(s) mAn+(aq) + nBm–(aq) • m·s n·s s: molar solubility

  4. 18-2 Relationship Between Solubility and Ksp

  5. (Continuous)

  6. 18-3 Common-Ion Effect in SolubilityEquilibria • Effect in solubility by common ion: Followed Le Châtelier’s principle, i.e., the solubility of a slightly soluble ionic compound is lowered in the presence of a second solute that furnishes a common ion. Example: For equilibrated PbI2(s) Pb2+(aq) + 2I–(aq) • What is happened once adding KI(s). • What is happened once adding Pb(NO3)2(s). • What is happened once adding PbI2(s). Ans: • Shift left • Shift left • No effect

  7. 18-4 Limitations of the Ksp Concept • The Diverse Noncommon Ion Effect: The Salt Effect • Noncommon ions (diverse ions, inert ions): The ions different from those directly involved in a solution equilibrium. • Salt effect for the solubility of ionic compound, PbI2(s) in KNO3(aq) for example, The activities (effective concentrations) of Pb2+ and I– are lower than their actual concentrations, resulting the higher solubility than PbI2(s) in pure water.

  8. (Continuous) • Incomplete Dissociation of Solute into Ions • Ion pair: An association of a cation and an anion in solution. An example of ion pair MgF+: Therefore, the solubility of MgF2(s) increased because of the ion pair.

  9. (Continuous) • Simultaneous Equilibria: Two or more reactions that occur and reach equilibria at the same time. • Assessing the Limitations of Ksp • The value listed in Table is based on ion activities. • The value calculated from the experimentally determined solubility is based on ion concentrations, assuming complete dissociation of the solute into ions and no ion-pair formation.

  10. 18-5 Criteria for Precipitation and Its Completeness • Whether Precipitation Occurs • Calculating Qip, the ion product: Qip > Ksp supersaturated, precipitation occur Qip = Ksp saturated, at equilibrium Qip < Ksp unsaturated, precipitation cannot occur * The effect of dilution when solutions are mixed must be considered.

  11. (Continuous) • Whether Precipitation Is Complete • A slightly soluble solid never totally precipitates from solution. • Generally, complete precipitation is considered as 99.9% of the target ion is precipitated, i.e., 0.1% left in solution. • Complete precipitation is favored when: • A very small value of Ksp. • A high initial concentration of the target ion. • A concentration of common ion (precipitating reagent) that greatly exceeds that of the target ion.

  12. 18-6 Fractional Precipitation • Fractional precipitation: A technique in which two or more ions in solution, each capable of being precipitated by the same reagent, are separated by the proper use of that reagent. Usually this means a significant difference in their Kspvalues.

  13. 18-7 Solubility and pH Example 1: Mg(OH)2(s) Mg2+(aq) + 2OH–(aq) Ksp=1.8x10–11 2H3O+(aq) + 2OH–(aq) 4H2O(l) K’ =1/Kw2=1.0x1028 Mg(OH)2(s) + 2H3O+(aq) Mg2+(aq) + 4H2O(l) K=1.8x1017 increasing [H3O+], increasing the solubility of Mg(OH)2 Example 2: AgCl(s) Ag+(aq) + Cl–(aq) increasing [H3O+], no effect on the solubility of AgCl

  14. Example 3: CaCO3(s) Ca2+(aq) + CO32–(aq) 2HCl + 2H2O → 2H3O+(aq) + Cl–(aq) CO32–(aq) + H3O+(aq) HCO3–(aq) + H2O(l) HCO32–(aq) + H3O+(aq) H2CO3(aq) + H2O(l) H2CO3(aq) → CO2(g) + H2O(l) CaCO3(s) + 2HCl → Ca2+(aq) + 2Cl–(aq) + H2O(l) + CO2(g) increasing [H3O+], increasing the solubility of CaCO3

  15. 18-8 Equilibria Involving Complex Ions • Complex-Ion: Consists of a central metal atom or ion bonded with ligand(s): • Metal center: Lewis acid, electron pair acceptor • Ligands: Lewis base, electron pair donor • At least one lone pair of electron in the Lewis structures of ligand • The ligand and metal center are jointed by coordinate covalent (dative) bonds. • Coordination compound: The neutral complex or compound containing complex ion(s). • The constant of formation reaction of a complex ion is called a formation constant (Kf), or stability constant.

  16. 18-8 (Continuous) • Complexation effect on Solubility AgCl(s) in NH3(aq) solution for example: AgCl(s) Ag+(aq) + Cl–(aq) Ksp = 1.8 x 10–10 Ag+(aq) + 2NH3(aq) [Ag(NH3)2]+(aq) Kf = 1.6 x 107 __________________________________________________________________________ AgCl(s) + 2NH3(aq) [Ag(NH3)2]+(aq) + Cl–(aq) K=KspxKf=2.9x10–3 Increasing the concentration of NH3, increasing the solubility of AgCl.

  17. EXAMPLE 18-10 Predicting Reactions Involving Complex Ions Predict what will happen if nitric acid is added to a solution of [Ag(NH3)2]Cl in NH3(aq). • Solve: • 1. Adding HNO3, shift right, NH3 decreasing [NH3]: • H3O+(aq) + NH3(aq)→ NH4+(aq) + H2O(l) • 2. decreasing [NH3], shift left, increasing [Ag+]: • Ag+(aq) + 2NH3(aq) [Ag(NH3)2] • Increasing [Ag+], shift left: • AgCl(s) Ag+(aq) + Cl–(aq) • 4. Ans: AgCl(s)precipitation occurred

  18. 18-9 Qualitative Cation Analysis • Classical qualitative cation analysis • Identify inorganic ions by acid-base chemistry, precipitation reactions, oxidation-reduction, and complex-ion formation etc. • “Qualitative” interest in what is present, not how much is present. • Classical qualitative analysis relevant to all the basic concepts of equilibria in aqueous solutions.

  19. Outline of qualitative cation analysis The group numbers used in the qualitative analysis scheme are unrelated to group numbers in the periodic table

  20. Group 1 Cations (Chloride group: Pb2+, Hg22+, Ag+) • Adding HCl resulted in PbCl2(s), Hg2Cl2(s), AgCl(s) if the sample contains those cations. • Separation by filtration or centrifugation: – Filtrate or supernatant for other groups’ cations analysis. – Precipitates for analyzing Pb2+, Hg22+, Ag+. • Group 1 cations analysis i) Analyzing for Pb2+(aq): • Treat the precipitate with hot water fordissolving PbCl2(s), i.e., releasing Pb2+(aq). • Treat the washings with K2CrO4(aq). If Pb2+(aq) present, forming a yellow precipitate, PbCrO4(s).

  21. ii) Analyzing for Ag+: • Treat the undissolved precipitate (in step i)) with NH3(aq) for dissolving AgCl(s), i.e., releasing Ag+(aq). • (Save precipitate for iii) Hg22+ analysis). • Treat the supernatant with HNO3(aq). If Ag+(aq) was present, reforming a white, AgCl(s). iii) Analyzing for Hg22+: • Treat the undissolved precipitate (in step ii)) with NH3(aq). If Hg22+ present, forming dark gray mixture, Hgo(l) and HgNH2Cl(s).

  22. End of Chapter 18

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