1. 18 GENERAL CHEMISTRY Solubility and Complex-Ion Equilibria
Chemistry 140 Fall 2002
TENTH EDITION
GENERAL CHEMISTRY
Principles and Modern Applications
PETRUCCI HERRING MADURA BISSONNETTE 18
Solubility and Complex-Ion Equilibria
In this chapter we will combine ideas about acid–base equilibria from Chapters 16
and 17 with ideas about the new types of equilibria.
PHILIP DUTTON
UNIVERSITY OF WINDSOR
DEPARTMENT OF CHEMISTRY AND BIOCHEMISTRY
General Chemistry: Chapter 18
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2. Solubility and Complex-Ion Equilibria
Chemistry 140 Fall 2002
Solubility and Complex-Ion Equilibria
Contents
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
Stalactite and stalagmite formations in a cavern in the Guangxi Province, People’s Republic of China. Stalactites and stalagmites are formed from calcium salts deposited as underground water seeps into the cavern and evaporates.
General Chemistry: Chapter 18
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3. Solubility and Complex-Ion Equilibria
Chemistry 140 Fall 2002
Solubility and Complex-Ion Equilibria
Contents
18-7
Solubility and pH
18-8
Equilibria Involving Complex Ions
18-9
Qualitative Cation Analysis
Stalactite and stalagmite formations in a cavern in the Guangxi Province, People’s Republic of China. Stalactites and stalagmites are formed from calcium salts deposited as underground water seeps into the cavern and evaporates.
General Chemistry: Chapter 18
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4. 18-1 The Solubility Product Constant, Ksp
Chemistry 140 Fall 2002
18-1 The Solubility Product Constant, Ksp
The equilibrium constant for the equilibrium established between a solid solute and its ions in a saturated solution.
CaSO4(s) Ca2+(aq) + SO42-(aq)
K =
aCaSO4
aCa2+aSO42-
The activity of a pure solid is 1. Because the concentrations of the ions are very small, we can set their activities equal to their molar concentrations. With these substitutions, we obtain expression shown in concentration terms.
Ksp = [Ca2+][SO42-] = 9.110-6 at 25°C
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5. General Chemistry: Chapter 18
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6. 18-2 The Relationship Between Solubility and Ksp
Chemistry 140 Fall 2002
18-2 The Relationship Between Solubility and Ksp
Molar solubility.
The molarity in a saturated aqueous solution.
Related to Ksp
g BaSO4/100 mL → mol BaSO4/L
→ [Ba2+] and [SO42-]
→ Ksp = 1.110-10
As a component of a “barium milk shake,” BaSO4 coats the intestinal tract so that this
soft tissue will show up when X-rayed. Even though Ba2+ is poisonous, BaSO4 is
harmless because its aqueous solubility is very low.
Barium sulfte, BaSO4, is a good absorber of X-rays.
General Chemistry: Chapter 18
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7. 18-3 The Common-Ion Effect in Solubility Equilibria
General Chemistry: Chapter 18
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8. The common-ion effect in solubility equilibrium
Chemistry 140 Fall 2002
(a) A clear saturated solution of lead(II) iodide from which excess undissolved solute has been filtered off. (b) When a small volume of a concentrated solution of KI (containing the common ion, I-) is added, a small quantity of PbI2(s) precipitates. A common ion reduces the solubility of a sparingly soluble solute.
FIGURE 18-1
The common-ion effect in solubility equilibrium
General Chemistry: Chapter 18
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9. 18-4 Limitations of the Ksp Concept
Ksp is usually limited to slightly soluble solutes. For more soluble solutes we must use ion activities Activities (effective concentrations) become smaller than the measured concentrations. The Diverse Noncommon Ion Effect: The Salt Effect Ionic interactions are important even when an ion is not apparently participating in the equilibrium. Uncommon ions tend to increase solubility.
General Chemistry: Chapter 18
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10. General Chemistry: Chapter 18
Chemistry 140 Fall 2002
The presence of ions, derived CrO42- from reduces the solubility
of K2CrO4 by a factor of about 35 over the concentration range shown (from
0 to 0.10 M added salt). Over the same
concentration range, the solubility of
is increased by the presence
of the diverse ions from but only
by about 25%
FIGURE 18-2
Comparison of the common-ion effect and the salt effect on the molar solubility of Ag2CrO4
General Chemistry: Chapter 18
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11. Incomplete Dissociation of Solute into Ions
Chemistry 140 Fall 2002
Incomplete Dissociation of Solute into Ions
We have assumed that all the dissolved solute appears in solution as separated cations
and anions. This assumption is often not valid.
The solute might not be 100% ionic, and some of the solute might enter solution in molecular form. Alternatively, some ions in solution might join together into ion pairs. An ion pair is two oppositely charged ions that are held together by the electrostatic attraction between the ions. For example, in a saturated solution of magnesium fluoride, although most of the solute exists as Mg2+ and F- ions, some
exists as the ion pair MgF+.
An ion pair, in a magnesium fluoride aqueous solution. The water molecules surrounding the ion pair form what is referred to as a solvent cage.
General Chemistry: Chapter 18
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12. Ion Pair formation. Some solute “molecules” are present in solution.
Chemistry 140 Fall 2002
Ion Pair formation.
Some solute “molecules” are present in solution.
Increasingly likely as charges on ions increase.
Ksp (CaSO4) = 2.310-4
when considering solubility in g/100 mL
Table 19: Ksp = 9.110-6
Activities take into account ion pair formation and must be used.
Assumption that all ions in solution are completely dissociated is not valid.
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13. Simultaneous Equilibria
Other equilibria are usually present in a solution. Kw for example. Reactions between solute ions and other species. Acid base reactions. Complex-ion formation These must be taken into account if they affect the equilibrium in question.
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14. Simultaneous Equilibria
The dissolution of PbI2(s) is more complex than we have shown:
PbI+(aq) + 2 I-(aq)
-I-
-I-
PbI42-(aq)
PbI3-(aq)
PbI2 (s)
Pb2+(aq) + 2 I-(aq)
+I-
+I-
PbI2 (aq)
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15. 18-5 Criteria for Precipitation and Its Completeness
Chemistry 140 Fall 2002
18-5 Criteria for Precipitation and Its Completeness
AgI(s) Ag+(aq) + I-(aq)
Ksp = [Ag+][Cl-] = 8.510-17
Mix AgNO3(aq) and KI(aq) to obtain a solution that is M in Ag+ and M in I-.
Saturated, supersaturated or unsaturated?
Q = [Ag+][Cl-] = (0.010)(0.015) = 1.10-4 > Ksp
General Chemistry: Chapter 18
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16. General Chemistry: Chapter 18
The Ion Product
Q is generally called the ion product.
Q > Ksp Precipitation should occur.
Q = Ksp The solution is just saturated.
Q < Ksp Precipitation cannot occur.
General Chemistry: Chapter 18
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17. General Chemistry: Chapter 18
Chemistry 140 Fall 2002 a b
(a) When three drops of 0.20 M KI are first added to mL of M Pb(NO3)2
a precipitate forms because Ksp is exceeded in the immediate vicinity of the drops.
(b) When the KI becomes uniformly mixed in the Pb(NO3)2, Ksp is no longer
exceeded and the precipitate re-dissolves. The criteria for precipitation must be
applied after dilution has occurred/
FIGURE 18-3
Applying the criteria for precipitation from solution – Example 18-5 illustrated
General Chemistry: Chapter 18
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18. 18-6 Fractional Precipitation
Chemistry 140 Fall 2002
18-6 Fractional Precipitation
Fractional precipitation is 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: One ion is precipitated, while the other(s) remains in solution. The primary condition for a successful fractional precipitation is that there be a significant difference in the solubilities of the substances being separated. (Usually this means a significant difference in their Ksp values.) The key to the technique is the slow addition of a concentrated solution of the precipitating reagent to the solution from which precipitation is to occur, as from a buret.
AgNO3(aq) is slowly added to a solution that is M in Br- and M in CrO42-
Essentially all the Br- has precipitated as pale yellow AgBr(s), with [Br-] in solution = 5 x 10-8 M. Red-brown Ag2CrO4(s) is just about to precipitate.
FIGURE 18-4
Fractional precipitation-Example 18-7 illustrated
General Chemistry: Chapter 18
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19. General Chemistry: Chapter 18
18-7 Solubility and pH
pH can affect the solubility of a salt. Especially when the anion of the salt is the conjugate base of a weak acid. Mg(OH)2 Milk of Magnesia.
General Chemistry: Chapter 18
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20. General Chemistry: Chapter 18
Chemistry 140 Fall 2002
18-7 Solubility and pH
Mg(OH)2 (s) Mg2+(aq) + 2 OH-(aq)
Ksp = 1.810-11 x2
OH-(aq) + H3O+(aq) H2O(aq)
K = 1/Kw = 1.01014
2 OH-(aq) + 2 H3O+(aq) H2O(aq)
K' = (1/Kw)2 = 1.01028
The large value of K for reaction (18.4) indicates that a greater amount of Mg(OH)2 will react (dissolve) in acidic solution than in pure water. Other slightly soluble solutes having basic anions (such as ZnCO3, MgF2, and CaC2O4) also dissolve to a greater extent in acidic solutions. For these solutes, we can write overall net ionic equations for solubility equilibria and corresponding K values based on Ksp for the solutes and Ka for the conjugate acids of the anions.
Mg(OH)2(s) + 2 H3O+(aq) Mg2+(aq) + 4 H2O(l)
K = Ksp(1/Kw)2 = (1.810-11)(1.01028) = 1.81017
General Chemistry: Chapter 18
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21. 18-8 Equilibria Involving Complex Ions
Chemistry 140 Fall 2002
18-8 Equilibria Involving Complex Ions
A saturated solution of silver chloride in contact with excess AgCl(s).
When NH3(aq) is added, the excess AgCl(s) dissolves through the formation of the complex ion [Ag(NH3)2]+.
AgCl(s) + 2 NH3(aq) → [Ag(NH3)2]+(aq) + Cl-(aq)
FIGURE 18-5
Complex-ion formation: dissolution of AgCl(s) in NH3
General Chemistry: Chapter 18
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22. General Chemistry: Chapter 18
Complex ion:
A polyatomic cation or anion composed of:
A central metal ion.
Ligands
Coordination compounds:
Are substance which contain complex ions.
General Chemistry: Chapter 18
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23. Formation Constant of Complex Ions
Chemistry 140 Fall 2002
Formation Constant of Complex Ions
AgCl(s) + 2 NH3(aq) → [Ag(NH3)2]+(aq) + Cl-(aq)
Think of this reaction as two simultaneous equilbria.
AgCl(s) Ag+(aq) + Cl-(aq)
Ksp = 1.810-11
Ag+(aq) + 2 NH3(aq) [Ag(NH3)2]+(aq)
To describe the ionization of a weak acid, we use the ionization constant Ka. For a solubility equilibrium, we use the solubility product constant, Ksp. The equilibrium constant that is used to deal with a complex-ion equilibrium is called the formation constant. The formation constant, Kf, of a complex ion isthe equilibrium constant describing the formation of a complex ion from a central ion and its attached groups. For reaction (18.7) this equilibrium constant expression is given on this slide.
Kf =
= 1.6107
[Ag(NH3)2]+
[Ag+]
[NH3]2
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24. General Chemistry: Chapter 18
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25. 18-9 Qualitative Cation Analysis
An analysis that aims at identifying the cations present in a mixture but not their quantities.
Think of cations in solubility groups according to the conditions that causes precipitation.
chloride group hydrogen sulfide group
ammonium sulfide group carbonate group
soluble group
Selectively precipitate the first group of cations then move on to the next.
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26. General Chemistry: Chapter 18
Chemistry 140 Fall 2002
Figure 18-7
Outline of a qualitative cation analysis
Various aspects of this scheme are described in the text. A sample containing all 25 cations can be separated into five groups by the indicated reagents.
General Chemistry: Chapter 18
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27. Cation Group 1: The Chloride Group
Chemistry 140 Fall 2002
Cation Group 1: The Chloride Group
Group 1 precipitate
Wash ppt with hot water. PbCl2 is slightly soluble. Test aqueous solution with CrO42-.
(c) Pb2+(aq) + CrO42- → PbCrO4(s)
Test remaining precipitate with ammonia.
(b) AgCl(s) + 2 NH3(aq) → Ag(NH3)2 (aq) + Cl-(aq)
Group 1 precipitate: a mixture of PbCl2 (white), HgCl2 (white), and AgCl (white) precipitated using HCl.
Test for Hg22+: a mixture of Hg (black) and HgNH2Cl (white).
Test for Pb2+: a yellow precipitate of PbCrO4(s).
(b) Hg2Cl2(a) + 2 NH3 → Hg(l) + HgNH2Cl(s) NH4+(aq) + Cl-(aq)
FIGURE 18-8
Chloride group precipitates
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28. Cation Groups 2 and 3: Equilibria Involving Hydrogen Sulfide.
H2S causes the familiar smell of rotten eggs.
It is detectable at 1 ppm, and can shut down your
respiratory system at 100 ppm. It is particularly
hazardous around “sour gas” wells.
H2S(aq) + H2O(l) HS-(aq) + H3O+(aq) Ka1 = 1.010-7
HS-(aq) + H2O(l) S2-(aq) + H3O+(aq) Ka2 = 1.010-19
S2- is an extremely strong base and is unlikely to be the precipitating agent for the sulfide groups.
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29. General Chemistry: Chapter 18
PbS(s) + H2O(l) Pb2+(aq) + HS-(aq) + OH-(aq) Ksp = 310-28
H3O+(aq) + HS-(aq) H2S(aq) + H2O(aq) 1/Ka1 = 1.0/1.010-7
H3O+(aq) + OH-(aq) H2O(l) + H2O(l) /Kw = 1.0/1.010-14
PbS(s) + 2 H3O+(l) Pb2+(aq) + H2S(aq) + 2 H2O(l)
Kspa =
= 310-7
Ksp
Ka1 Kw
310-28
1.010-7 1.010-14 =
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30. Dissolving Metal Sulfides
Several methods exist to re-dissolve precipitated metal sulfides.
React with an acid.
FeS readily soluble in strong acid but PbS and HgS are not because their Ksp values are too low.
React with an oxidizing acid.
3 CuS(aq) + 8 H+(aq) + 2 NO3-(aq) →
3 Cu2+(aq) + 3 S(s) + 2 NO(g) + 4 H2O(l)
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31. Complex-ion formation: A test for Cu2+
Chemistry 140 Fall 2002
Dilute CuSO4(aq) (left) derives its pale blue color from the complex ions [Cu(H2O)4]2+
When is added (here labeled “conc ammonium hydroxide”), the color changes to a deep violet, signaling the presence of [Cu(NH3)4]2+ (right). The deep violet color is detectable at much lower concentrations than the pale blue; the formation of [Cu(NH3)4]2+ is a sensitive test for the presence of Cu2+.
[Cu(H2O)4]2+(aq) + 4 NH3(aq) → [Cu(NH3)4]2+(aq) + 4 H2O(l)
FIGURE 18-9
Complex-ion formation: A test for Cu2+
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32. End of Chapter Questions
Look at problem solving as a skill.
This is not a memory exercise.
Transfer the skills learned in one type of problem to other types.
You are building a tool kit that will help you in any problem you have.
General Chemistry: Chapter 18
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