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Solubility product constant, Ksp: the equilibrium constant expression for the dissolving of a slightly soluble solid.

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The Solubility Product Constant, Ksp Many important ionic compounds are only slightly soluble in water (we used to call them insoluble Chapter 4). – PowerPoint PPT presentation

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Title: Solubility product constant, Ksp: the equilibrium constant expression for the dissolving of a slightly soluble solid.


1
The Solubility Product Constant, Ksp
  • 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
  • Solubility product constant, Ksp the equilibrium
    constant expression for the dissolving of a
    slightly soluble solid.

Ksp Ba2 SO42
2
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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 doesnt 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 Ksp for 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 105 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
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.
Calculate the molar solubility of Ag2SO4 in 1.00
M Na2SO4(aq).
9
Will Precipitation Occur? Is It Complete?
  • Qip is the ion product reaction quotient and is
    based on initial conditions of the reaction.

Qip and Qc new look, same great taste!
  • Qip can then be compared to Ksp.
  • Precipitation should occur if Qip gt Ksp.
  • Precipitation cannot occur if Qip lt 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.

10
  • 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 1012
  • 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 108

11
  • Simple explanation problem
  • 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.

12
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.

13
  • 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 109
14
Effect of pH on Solubilityfly in the Ointment
  • 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
LeChâteliers 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.
15
  • Example 16.11
  • What is the molar solubility of Mg(OH)2(s) in a
    buffer solution having OH 1.0 x 105 M, that
    is, pH 9.00?

Mg(OH)2(s) Mg2(aq) 2 OH(aq) Ksp
1.8 x 1011
Another one of those explain problems 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.
16
Neutralization Reactions
  • At the equivalence point in an acidbase
    titration, the acid and base have been brought
    together in precise stoichiometric proportions.
  • The endpoint is the point in the titration at
    which the indicator changes color.
  • Ideally, the indicator is selected so that the
    endpoint and the equivalence point are very close
    together.
  • The endpoint and the equivalence point for a
    neutralization titration can be best matched by
    plotting a titration curve, a graph of pH versus
    volume of titrant.

17
Titration Curve, Strong Acid with Strong Base
Bromphenol blue, bromthymol blue, and
phenolphthalein all change color at very nearly
20.0 mL
At about what volume would we see a color change
if we used methyl violet as the indicator?
18
  • Example 15.20
  • Calculate the pH at the following points in the
    titration of 20.00 mL of 0.500 M HCl with 0.500 M
    NaOH
  • H3O Cl Na OH ? Na Cl
    2 H2O
  • (a) before the addition of any NaOH
  • (b) after the addition of 10.00 mL of 0.500 M
    NaOH
  • (c) after the addition of 20.00 mL of 0.500 M
    NaOH
  • (d) after the addition of 20.20 mL of 0.500 M
    NaOH

19
Titration Curve, Weak Acid with Strong Base
The equivalence-point pH is NOT 7.00 here. Why
not??
Bromphenol blue was ok for the strong acid/strong
base titration, but it changes color far too
early to be useful here.
20
Sample problem for WA/SB titration
  • The titration of 100.0 mL of 0.016 M HOCl (Ka
    3.5 x 10-8 with 0.0400 M NaOH. How many mL of
    0.040 M NaOH are required to reach the
    equivalence point?
  • Calc the pH after addition of 10.0 mL NaOH
  • Halfway to equivalence point
  • At equivalence point.
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