Solubility and Complex Ion Equilibria

1
Solubility and Complex Ion Equilibria
2

• Slightly soluble salts establish a
• dynamic equilibrium with the
• hydrated cations and anions in
• solution.

3

• When the solid is first added to
• water, no ions are initially present.

4

• As dissolution proceeds, the
• concentration of ions increases
• until equilibrium is established.
• This occurs when the solution is
• saturated.

5

• The equilibrium constant, the
• Ksp, is no more than the product of
• the ions in solution.
• (Remember, solids do not
• appear in equilibrium expressions.)

6

• For a saturated solution of
• AgCl, the equation would be
• AgCl (s) ? Ag (aq) Cl- (aq)

7

• The solubility product expression
• would be
• Ksp Ag Cl-

8

• The AgCl(s) is left out since
• solids are left out of equilibrium
• expressions (constant
• concentrations).

9
(No Transcript)
10
You can find loads of Ksps on tables.

• Find the Ksp values write the
• Ksp expression for the following

11

• CaF2(s) ? Ca2 2 F- Ksp
• Ag2SO4(s) ? 2 Ag SO4-2 Ksp
• Bi2S3(s) ? 2 Bi3 3 S-2 Ksp

12
Determining Ksp From Experimental Measurements

• In practice, Ksps are determined
• by careful laboratory
• measurements using various
• spectroscopic methods.
• Remember STOICHIOMETRY!!

13
Example

• Lead (II) chloride dissolves to a
• slight extent in water according
• to the equation
• PbCl2 ? Pb2 2Cl-

14

• Calculate the Ksp if the lead ion
• concentration has been found to
• be 1.62 x 10-2M.

15
Solution

• If leads concentration is x ,
• then chlorides concentration is
• 2x.
• So. . . .
• Ksp (1.62 x 10-2)(3.24 x 10-2)2
• 1.70 x 10-5

16
Exercise 12 Calculating Ksp from Solubility I

• Copper(I) bromide has a measured
• solubility of 2.0 X 10-4 mol/L at
• 25C.
• Calculate its Ksp value.

17
Solution

• Ksp 4.0 X 10-8

18
Exercise 13 Calculating Ksp from Solubility II

• Calculate the Ksp
• value for bismuth
• sulfide (Bi2S3), which
• has a solubility of
• 1.0 X 10-15 mol/L at
• 25C.

19
Solution

• Ksp 1.1 X 10-73

20

• ESTIMATING SALT
• SOLUBILITY FROM Ksp

21
Example

• The Ksp for CaCO3 is 3.8 x 10-9 _at_
• 25C.
• Calculate the solubility of calcium
• carbonate in pure water in
• a) moles per liter
• b) grams per liter

22

• The relative solubilities can be
• deduced by comparing values of Ksp.
• BUT, BE CAREFUL!
• These comparisons can only be
• made for salts having the same
• IONION ratio.

23

• Please dont forget solubility
• changes with temperature!
• Some substances become less
• soluble in cold while some become
• more soluble!
• Aragonite.

24
Exercise 14 Calculating Solubility from Ksp

• The Ksp value for copper(II) iodate,
• Cu(IO3)2, is 1.4 X 10-7 at 25C.
• Calculate its solubility at 25C.

25
Solution

• 3.3 X 10-3 mol/L

26
Exercise 15 Solubility and Common Ions

• Calculate the solubility of solid CaF2
• (Ksp 4.0 X 10-11)
• in a 0.025 M NaF solution.

27
Solution

• 6.4 X 10-8 mol/L

28
Ksp and the Reaction Quotient, Q

• With some knowledge of the
• reaction quotient, we can decide
• 1) whether a ppt will form, AND
• 2) what concentrations of ions
• are required to begin the ppt. of
• an insoluble salt.

29

• 1. Q lt Ksp, the system is not at equil.
  (unsaturated)
• 2. Q Ksp, the system is at equil. (saturated)
• 3. Q gt Ksp, the system is not at equil.
  (supersaturated)

30

• Precipitates form when the
• solution is supersaturated!!!

31
Precipitation of Insoluble Salts

• Metal-bearing ores often contain
• the metal in the form of an
• insoluble salt, and, to complicate
• matters, the ores often contain
• several such metal salts.

32

• Dissolve the metal salts to obtain
• the metal ion, concentrate in some
• manner, and ppt. selectively only
• one type of metal ion as an
• insoluble salt.

33
Exercise 16 Determining Precipitation Conditions

• A solution is prepared by adding
• 750.0 mL of 4.00 X 10-3 M Ce(NO3)3
• to 300.0 mL of 2.00 X 10-2 M KIO3.
• Will Ce(IO3)3 (Ksp 1.9 X 10-10)
• precipitate from this solution?

34
Solution

• yes

35
Exercise 17 Precipitation

• A solution is prepared by mixing
• 150.0 mL of 1.00 X 10-2 M Mg(NO3)2
• and 250.0 mL of 1.00 X 10-1 M NaF.
• Calculate the concentrations of Mg2
• and F- at equilibrium with solid MgF2
• (Ksp 6.4 X 10-9).

36
Solution

• Mg2 2.1 X 10-6 M
• F- 5.50 X 10-2 M

37
SOLUBILITY AND THE COMMON ION EFFECT
38

• Experiment shows that the
• solubility of any salt is always less
• in the presence of a common
• ion.

39

• LeChateliers Principle, thats why!
• Be reasonable and use
• approximations when you can!!

40

• Just remember what happened
• earlier with acetic acid and
• sodium acetate.
• The same idea here!

41

• pH can also affect solubility.
• Evaluate the equation to see who
• would want to react with the
• addition of acid or base.

42

• Would magnesium hydroxide be
• more soluble in an acid or a base?
• Why?
• Mg(OH)2(s) ? Mg2(aq) 2 OH-(aq)
• (milk of magnesia)

43
Why Would I Ever Care About Ksp ???

• Keep reading to find out !
• Actually, very useful stuff!

44
Solubility, Ion Separations, and Qualitative
Analysis

• introduce you to some basic
• chemistry of various ions.
• illustrate how the principles of
• chemical equilibria can be applied.

45
Objective

• Separate the
• following
• metal ions
• silver,
• lead,
• cadmium and
• nickel

46

• From solubility rules, lead and silver
• chloride will ppt, so add dilute HCl.
• Nickel and cadmium will stay in
• solution.

47

• Separate by filtration
• Lead chloride will dissolve in HOT
• water
• filter while HOT and those two will
• be separate.

48

• Cadmium and nickel are more
• subtle.
• Use their Ksps with sulfide ion.
• Who ppts first???

49
Exercise 18 Selective Precipitation

• A solution contains 1.0 X 10-4 M Cu
• and 2.0 X 10-3 M Pb2.
• If a source of I- is added gradually to
• this solution, will PbI2 (Ksp 1.4 X
• 10-8) or CuI (Ksp 5.3 X 10-12)
• precipitate first?

50

• Specify the concentration of I-
• necessary to begin precipitation of
• each salt.

51
Solution

• CuI will precipitate first.
• Concentration in excess of
• 5.3 X 10-8 M required.

52

• If this gets you interested, lots
• more information on this topic in
• the chapter.
• Good bedtime reading for
• descriptive chemistry!

53

• THE EXTENT OF LEWIS ACID-BASE REACTIONS
• FORMATION CONSTANTS

54

• When a metal ion (a Lewis acid)
• reacts with a Lewis base, a
• complex ion can form.
• The formation of complex ions
• represents a reversible equilibria
• situation.

55

• A complex ion is a charged
• species consisting of a metal ion
• surrounded by ligands.

56

• A ligand is typically an anion or
• neutral molecule that has an
• unshared electron pair that can
• be shared with an empty metal
• ion orbital to form a metal-ligand
• bond.
• Some common ligands are
• H2O, NH3, Cl-, and CN-.

57

• The number of ligands attached to
• the metal ion is the coordination
• number.
• The most common coordination
• numbers are 6, 4, 2

58

• Metal ions add ligands one at a
• time in steps characterized by
• equilibrium constants called
• formation constants.
• Ag 2NH3 ? Ag(NH3)22
• acid base

59
Stepwise Reactions

• Ag(aq) NH3(aq) ? Ag(NH3)(aq)
• Kf1 2.1 x 103
• Ag(NH3) 2.1 x 103
• AgNH3

60

• Ag(NH3) NH3(aq) ? Ag(NH3)2(aq)
• Kf2 8.2 x 103
• Ag(NH3)2 8.2 x 103
• Ag(NH3)NH3

61

• In a solution containing Ag and
• NH3, all of the species NH3, Ag,
• Ag(NH3), and Ag(NH3)2 exist at
• equilibrium.
• Actually, metal ions in aqueous
• solution are hydrated.

62

• More accurate representations would
• be
• Ag(H2O)2 instead of Ag, and
• Ag(H2O)(NH3) instead of Ag(NH3).

63
The equations would be

• Ag(H2O)2(aq) NH3(aq) ?
• Ag(H2O)(NH3)(aq) H2O(l)
• Kf1 2.1 x 103
• Ag(H2O)(NH3) 2.1 x 103
• Ag(H2O)2NH3

64

• Ag(H2O)(NH3)(aq) NH3(aq) ?
• Ag(NH3)2(aq) 2H2O(l)
• Kf2 8.2 x 103
• Ag(NH3)2 8.2 x 103
  Ag(H2O)(NH3)NH3

65

• The sum of the equations gives the
• overall equation, so the product of
• the individual formation constants
• gives the overall formation constant

66

• Ag 2NH3 ? Ag(NH3)2
• or
• Ag(H2O)2 2NH3 ? Ag(NH3)2 2H2O
• Kf1 x Kf2 Kf
  
• (2.1 x 103) x (8.2 x 103) 1.7 x 107

67
Exercise 19

• Calculate the equilibrium
• concentrations of Cu2, NH3, and
• Cu(NH3)42 when 500. mL of 3.00 M
• NH3 are mixed with 500. mL of 2.00 x
• 10-3 M Cu(NO3)2.
• Kformation 6.8 x 1012.

68
Solubility and Complex Ions
69

• Complex ions are often insoluble in
• water.
• Their formation can be used to
• dissolve otherwise insoluble salts.
• Often as the complex ion forms,
• the equilibrium shifts to the right
• and causes the insoluble salt to
• become more soluble.

70

• If sufficient aqueous ammonia is
• added to silver chloride, the latter
• can be dissolved in the form of
• Ag(NH3)2.

71

• AgCl(s) ? Ag(aq) Cl-(aq)
• Ksp 1.8 x 10-10
• Ag(aq) 2 NH3(aq) ? Ag(NH3)2(aq)
• Kformation 1.6 x 107

72
Sum

• K Ksp x Kformation 2.0 x 10-3
• Ag(NH3)2Cl-
• NH32

73

• The equilibrium constant for
• dissolving silver chloride in ammonia
• is not large however, if the
• concentration of ammonia is
• sufficiently high, the complex ion
• and chloride ion must also be high,
• and silver chloride will dissolve.

74
Exercise 20 Complex Ions

• Calculate the concentrations of
• Ag, Ag(S2O3)-, and Ag(S2O3)23- in a
• solution prepared by mixing 150.0
• mL of 1.00 X 10-3 M AgNO3 with
• 200.0 mL of 5.00 M Na2S2O3.

75
The stepwise formation equilibria are

• Ag S2O32- ? Ag(S2O3)-
• K1 7.4 X 108
• Ag(S2O3)- S2O32- ? Ag(S2O3)23-
• K2 3.9 X 104

76
Solution

• Ag 1.8 X 10-18 M
• Ag(S2O3)- 3.8 X 10-9 M

77
ACID-BASE AND PPT EQUILIBRIA OF PRACTICAL
SIGNIFICANCE

• SOLUBILITY OF SALTS IN WATER
• AND ACIDS

78
The solubility of PbS in water

• PbS (s) ? Pb2 S-2
• Ksp 8.4 x 10-28

79
The Hydrolysis of the S-2 ion in Water

• S-2 H2O ? HS- OH-
• Kb 0.077

80
Overall Process

• PbS H2O ? Pb2 HS- OH-
• Ktotal Ksp x Kb 6.5 x 10-29

81

• May not seem like much, but it can
• increase the environmental lead
• concentration by a factor of about
• 10,000 over the solubility of PbS
• calculated from simply Ksp!

82

• Any salt containing an anion that is
• the conjugate base of a weak acid
• will dissolve in water to a greater
• extent than given by the Ksp.

83

• This means salts of sulfate,
• phosphate, acetate, carbonate, and
• cyanide, as well as sulfide can be
• affected.

84

• If a strong acid is added to
• water-insoluble salts such as ZnS
• or CaCO3, then hydroxide ions from
• the anion hydrolysis is removed by
• the formation of water.
• This shifts the anion hydrolysis
• further to the right the weak acid is
• formed and the salt dissolves.

85

• Carbonates and many metal
• sulfides along with metal
• hydroxides are generally soluble
• in strong acids.
• The only exceptions are sulfides
• of mercury, copper, cadmium and
• a few others.

86

• Insoluble inorganic salts containing
• anions derived from weak acids
• tend to be soluble in solutions of
• strong acids.
• Salts are not soluble in strong acid
• if the anion is the conjugate base of
• a strong acid!!