Chapter 19 PRECIPITATION REACTIONS

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Chapter 19PRECIPITATION REACTIONS
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Solubility of Ionic Solids

• Depends on the balance of two forces
• Attraction between H2O molecules and ions of
  solid.
• Force of attraction between oppositely charged
  ions within solid.

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Solubility Rules

• Use in predicting results of precipitation
  reactions. MEMORIZE THE SOLUBILITY RULES!!!!!
• Determine ions present and possible products.
• Use solubility rules to determine if any are
  insoluble.

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Example 1

• Ba(NO3)2(aq) Na2CO3(aq)

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Solubility Rules

• 1. Salts containing Group I elements are soluble
  (Li, Na, K, Cs, Rb). Exceptions to this rule
  are rare. Salts containing the ammonium ion
  (NH4) are also soluble.
• 2. Salts containing nitrate ion (NO3-) are
  generally soluble.
• 3. Salts containing Cl-, Br-, I- are generally
  soluble. Important exceptions to this rule are
  halide salts of Ag, Pb2, and (Hg2)2. Thus,
  AgCl, PbBr2, and Hg2Cl2 are all insoluble.

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Solubility Rules Continued

• 4. Most silver salts are insoluble. AgNO3 and
  Ag(C2H3O2) are common soluble salts of silver
  virtually anything else is insoluble.
• 5. Most sulfate salts are soluble. Important
  exceptions to this rule include BaSO4, PbSO4,
  Ag2SO4 and SrSO4.
• 6. Most hydroxide salts are only slightly
  soluble. Hydroxide salts of Group I elements are
  soluble. Hydroxide salts of Group II elements
  (Ca, Sr, and Ba) are slightly soluble. Hydroxide
  salts of transition metals and Al3 are
  insoluble. Thus, Fe(OH)3, Al(OH)3, Co(OH)2 are
  not soluble.

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Solubility Rules Continued

• 7. Most sulfides of transition metals are highly
  insoluble. Thus, CdS, FeS, ZnS, Ag2S are all
  insoluble. Arsenic, antimony, bismuth, and lead
  sulfides are also insoluble.
• 8. Carbonates are frequently insoluble. Group II
  carbonates (Ca, Sr, and Ba) are insoluble. Some
  other insoluble carbonates include FeCO3 and
  PbCO3.
• 9. Chromates are frequently insoluble. Examples
  PbCrO4, BaCrO4
• 10. Phosphates are frequently insoluble.
  Examples Ca3(PO4)2, Ag2PO4
• 11. Fluorides are frequently insoluble. Examples
  BaF2, MgF2 PbF2.

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Example 1 Continued

• Ba(NO3)2(aq) Na2CO3(aq)

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Stoichiometry

• Mole Relations
• Coefficients in the net ionic equation can be
  used in the usual way to relate the moles of
  reactants and products.
• Moles of ions can be deduced from solute
  concentrations.

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Example 2

• What is the molar concentration of Ba2 and F-
  in a solution containing 0.0075 M BaF2

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Precipitation Titrations

• Used to determine the concentration of species in
  solution or in a solid mixture.
• Indicator shows, usually by color change, when
  the species being analyzed for has been consumed

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General Principles

• Involves formation of a precipitate
• Must determine the volume of a
• standardized titrant needed to just
• precipitate all of the ion.
• Need an indicator or electrode to
• determine when the precipitation is
• complete

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Solubility Equilibria

• Solubility Product Constant, Ksp
• Precipitation reactions like all reactions, reach
  a position of equilibrium.
• Expression for Ksp
• MaXb lt----------gt aMb bX-a
• Ksp Mba X-ab

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Solubility Product Principle

• In any water solution in equilibrium with a
  slightly soluble ionic compound, the product of
  the concentrations of its ions, each raised to a
  power equal to its coefficient in the solubility
  equation is a constant. This constant, Ksp, has
  a fixed value at a given temperature, independent
  of the concentrations of the individual ions.

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Two Ion Compound

• AgCl

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Three Ion Compound

• PbCl2

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Four Ion Compound

• Al(OH)3

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Calculation of Ksp

• Calculated from measured solubility
• AgCl Ksp (s) (s) s2
• PbCl2 Ksp (s) (2s)2 4s3
• Al(OH)3 Ksp (s) (3s)3 27s4

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Example 3

• At 20 oC, a saturated aqueous solution of silver
  acetate, AgC2H3O2, contains 1.0 g dissolved in
  100.0 mL of solution. Calculate the Ksp for
  AgC2H3O2.

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Determination of Solubility

• In pure water
• Ksp s2 s (Ksp)1/2
• Ksp 4s3 s (Ksp/4)1/3
• Ksp 27s4 s (Ksp/27)1/4

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Example 4

• Estimate the solubility of lead (II) bromide in
  (a) moles per liter and (b) grams per liter of
  pure water. Ksp 6.3 x 10-6

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Uses of Ksp

• Calculation of concentration of one ion, knowing
  that of the other

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Example 5

• You have a solution that has a lead (II)
  concentration of 0.0012 M. What is the maximum
  concentration of chloride ions that would be
  present? Ksp 1.7 x 10-5

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Uses of Ksp

• Determination of whether a precipitate will form
• Compare original concentration product, P, to Ksp
• if P lt Ksp, no precipitate will form
• if P gt Ksp, precipitate forms until P becomes
  equal to Ksp

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Example 6

• You have 100.0 m of a solution that has a lead
  (II) concentration of 0.0012 M. Does PbCl2
  precipitate when 1.20 g of solid NaCl is added?

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Determination of Solubility

• In solution containing a common ion
• Solubility is much less than in pure water

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Example 7

• Calculate the solubility of silver carbonate,
  Ag2CO3, in moles per liter, in pure water.
  Compare this with the molar solublity of Ag2CO3
  in 225 mL of water to which 0.15 g of Na2CO3 has
  been added.

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Simultaneous Equilibria

• Two or more reactions occur at the same time is a
  solution, all of them being described as
  equilibrium processes.
• The equilibrium for the overall reaction is the
  product of the equilibrium constants for the
  summed reactions.
• That is Knet K1 x K2

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Solubility and pH

• Any salt containing an anion that is the
  conjugate base of a weak acid dissolves in water
  to a greater extent than that given by Ksp
  because the ions undergo a hydrolysis reaction.

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Example 8

• PbS
• PbS(s) lt----------gt Pb2(aq) S2-(aq) Ksp
  8.4 x 10-28
• S2-(aq) H2O(l) lt----------gt HS-(aq) OH-(aq)
  Kb 1 x 105
• Overall
• PbS(s) H2O(l) lt---gt Pb2(aq) HS-(aq)
  OH-(aq) Knet 8.4 x 10-23

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• In general, the solubility of a salt containing
  the conjugate base of a weak acid is increased by
  addition of a stronger acid to the solution. In
  contrast, the salts are not soluble in strong
  acid if the anion is the conjugate base of a
  strong acid.

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• CaCO3 lt--------gt Ca2(aq) CO32-(aq) K Ksp
  3.8 x 10-9
• CO32-(aq) H2O(l) lt--------gt HCO3-(aq)
  OH-(aq) K Kb 2.1 x 10-4
• OH-(aq) H3O(aq) lt--------gt 2H2O(l) K
  1/Kw 1 x 1014

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• NET
• CaCO3(s) H3O(aq) lt------gt Ca2(aq)
  HCO3-(aq) H2O(l)
• Knet (Ksp) (Kb)/(Kw) 79.8

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Solubility and Complex Ions

• Examples of complex ions AgCl2-, Ag(S2O3)23-,
  Ag(CN)2-
• The solubility of certain insoluble compounds
  can be increased when a complex ion is formed.
  Complex ions usually refer to cations in which
  surrounding water molecules have been replaced by
  some other electron pair donor. The equilibrium
  constant will equal the solubility constant times
  the formation constant for the complex ions.
  (Note - Chapter 23 in your textbook covers
  complex ions - their formation and nomenclature)

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Solubility and Complex ions

• AgCl(s) 2NH3 ?? Ag(NH3)2 Cl-
• K (Ksp) (Kf)
• (1.8 x 10-10) (1.6 x 107) 0.00288 0.0029

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Example 9

• Solid gold (I) chloride AuCl, is dissolved when
  excess cyanide ions, CN-, are added to give a
  water soluble complex ion.
• AuCl(s) 2CN- (aq) ??Au(CN)-(aq) Cl-
• Show that this equation is the sum of two other
  equations and calculate the Knet.

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Solubility, Ion Separations, and Qualitative
Analysis