Chapter 19: Solubility and Simultaneous Equilibria

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Chapter 19 Solubility and Simultaneous
Equilibria

• None of the salts describe as insoluble in
  Chapter 5 are totally insoluble
• For example, if solid AgCl is added to water a
  small amount dissolves
• The ion product is the product of the molar
  concentrations of the dissolved ions of the solute

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• The solubility product constants for a number of
  ionic compounds are given in Table 19.1 (and
  Appendix C)
• Ksp can be calculated by determining the molar
  solubility or the number of moles of salt
  dissolved in one liter of the saturated solution
• We can assume that the salt that dissolves is
  100 dissociated (see Facets of Chemistry 19.1
  for a discussion of the validity of this
  assumption)

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• Example The molar solubility of PbF2 in a 0.10 M
  Pb(NO3)2 solution at 25oC is 3.1x10-4 mol L-1.
  What is Ksp for PbF2?
• ANALYSIS Write the solubility equation, setup
  and evaluate the ICE table, and calculated the
  requested quantity.
• SOLUTION

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• Often, a value for Ksp is available and the molar
  solubility needs to be calculated
• This can be done with or without the presence of
  a common ion
• An ion in solution that has been supplied from
  more than one solute is called a common ion
• When a common ion is present, the molar
  solubility of an ionic compound decreases
• This is called the common ion effect
• Care must be taken to use all the information
  correctly

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• Example What is the molar solubility of PbI2 in
  0.10 M NaI?
• ANALYSIS This is a common ion problem. NaI is
  soluble in water and provides the common ion
  (I-).
• SOLUTION

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• This is a common ion problem, and Ksp is small
• Ksp can be used to determine if a precipitation
  will occur
• If the ion product gt Ksp, the solution is
  supersaturated and a precipitate will form
• If the ion product Ksp, the solution is
  saturated and no precipitate will form
• If the ion product lt Ksp, the solution is
  unsaturated and no precipitate will form

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• The pH dependence of metal carbonate solubility
  is similar to the pH dependence of metal sulfide
  solubility (see Example 19.10) and can provide a
  basis for selective precipitation
• Metal ions are Lewis acids and so can become
  covalently bonded with ions or neutral molecules
  in solution
• The resulting species is called a complex ion

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• Copper(II) ion is a typical example
• The Lewis base that attaches itself to the metal
  ion is called a ligand (H2O is the ligand)
• The atom in the ligand that actually provides the
  electron pair is called the donor atom (O is the
  donor atom)
• The metal ion is called the acceptor (Cu2 is the
  acceptor)
• Compounds that contain complex ions are called
  coordination compounds and the complex itself is
  sometimes called a coordination complex

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• Formation or stability constants, Kform, can be
  used to describe complex ion formation
• For example, for copper(II) ions in solutions of
  ammonia
• The inverse of the formation constant is called
  the instability constant, Kinst, which are
  preferred by some

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• Table 19.3 lists both formation and instability
  constants for some ions
• Appendix C lists only formation constants
• Complex ion formation can affect the solubility
  of salts

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• The solubility of a slightly soluble salt
  increases when one of its ions can be changed
  into soluble complex ion
• Example Calculate the solubility of silver
  chloride in 0.10 M NH3.
• ANALYSIS The equations describing the solubility
  and complex ion formation must be combined into a
  single equation and solved.

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• SOLUTION

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• Setup the ICE table and solve for the molar
  solubility x

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