6th Edition Kotz
Study Guide:
Chapter 18.4 – 18.7
Principles of Reactivity:
Other
Aspects of Aqueous Equilibria: Hetergeneous Phase Equilibria and Ksp
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Twelve: Applications of Aquous
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O. Discussion Questions
Hetergenous Equilibria |
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By
Dr. Andrea
Wallace
CGCC
Edited by
John Taylor
FCCJ-North Campus
Chapter 18.4 – 18.7
Principles of Reactivity: Other Aspects of Aqueous Equilibria
18.4 Solubility
of Salts, p. 873
(CD-Rom, Screen 18.8)
Precipitation reactions are exchange reactions in which one of the products is a water-insoluble compound.
(CD-Rom Screen 18.9)
The Solubility Product Constant, Ksp - solubility product constant for a salt. It is really the equilibrium constant for the dissolution of a salt.
Consider a salt that is placed in water, it will dissolve until equilibrium is reached and a saturated solution is formed.
AgBr(s) ó Ag+1(aq) +
Br-1(aq)
At 25oC, a saturated solution will have a [Ag+1] = [Br-1] = 7.35 x 10-7 M. In other words, the compound has a solubility of 7.35 x 10-7 M.
Thus, Ksp = __________________________
And substituting in the concentrations, Ksp = ___________________ = _____________
General Formula for Ksp expressions:
AxBy(s) ó x Ay+(aq) + y Bx-(aq)
Ksp = [Ay+]x [Bx-]y
Ksp values can be found in Table 18.2, p. 874 and Appendix J.
Exercise 18.11, p. 875
Write the Ksp expressions for the following salts.
a) AgI
b) BaF2
c) Ag2CO3
Exercise 18.12, p. 763
The barium ion concentration, [Ba+2], in a saturated solution of barium fluoride is 3.6 x 10-3 M. Calculate the value of Ksp for BaF2.
BaF2(s) ó Ba+2(aq) + 2 F-1(aq)
(CD-Rom, Screen 18.11)
Exercise 18.13, p. 878
Using the value of Ksp in Appendix J, calculate the solubility of Ca(OH)2 in moles per liter and grams per liter. (Table 18.2, p. 874, Ksp [Ca(OH)2] = 5.5 x 10-5
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Ca(OH)2(s) ó |
Ca+2(aq) + |
OH-1(aq) |
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Relative solubilities can be deduced by comparing Ksp values if ion ratios are equal. In general, the smaller the Ksp, the ___________soluble the compound.
Exercise 18.14, page 879
Using Ksp values, tell which salt in each pair is more soluble in water.
1. AgCl (Ksp = 1.8 x 10-10) or AgCN (Ksp = 6.0 x 10-17)
2. Mg(OH)2 (Ksp = 5.6 x 10-12) or Ca(OH)2 (Ksp = 5.5 x 10-5)
3. MgCO3 (Ksp = 6.8 x 10-6) or CaCO3 (Ksp = 3.4 x 10-9)
Solubility and the Common Ion Effect
(CD-Rom, Screen 18.12)
See Figure 18.13, p. 879
Addition of aqueous Ag+ causes the equilibrium to shift to form more reactant (precipitation occurs).
AgCH3CO2(s) ó Ag+1(aq) + CH3CO2-1(aq)
Exercise 18.16, p. 882
Calculate the solubility of Zn(CN)2 at 25 oC
(a) in pure water and (b) in the presence of 0.10 M Zn(NO3)2. (Ksp for Zn(CN)2 is 8.0 x 10-12.)
Effect of Basic Anions on Salt Solubility
(CD-Rom, Screen 18.13)
In general, the solubility of a salt containing the _______________ of a _____________ is increased by the addition of a stronger acid.
Examples: Carbonate (CO3-2) salts from H2CO3
Sulfide (S-2) salts from H2S
Illustrate with reactions and describe how the equilibrium is shifted to products when a strong acid is added to
FeS(s) ó Fe+2(aq) + S-2(aq)
Illustrate with reactions and describe how the equilibrium is shifted to products when a strong acid is added to
CaCO3(s)
ó Ca+2(aq) +
CO3-2(aq)
In contrast, the salts are not soluble in strong acid if the anion is the _____________ of a ______________________.
Examples: Cl-1, Br-1, I-1 from HCl, HBr, and HI
The reactions below show that AgCl is not more soluble in strong acid. Explain.
AgCl(s) ó Ag+1(aq) + Cl-1(aq)
Cl-1(aq) + H3O+(aq) ó HCl(aq) + H2O(l)
18.5
Precipitation Reactions, p. 884
(CD-Rom, Screen 18.14)
Metal Ores contain metals in the form of insoluble metal salts. Pure metal is obtained by dissolving the metal ore. A precipitating agent is then added to selectively precipitate the metal ion from solution. The metal ion is then reduced to its zero state (pure metal).
Q = _____________________
Q = ___________ at equilibrium (when a ______________ solution exists and no more solute can be dissolved.)
Q < Ksp solution is _______________________.
1) if solid is present, solid will dissolve until a saturated solution develops and Q = Ksp.
2) if solid is not present, more ions can be added until a precipitate forms and a saturated solution develops and Q = Ksp. (Ppt does not form)
Q > Ksp solution is _______________________.
Concentration of ions is too high and solid will begin to precipitate until a saturated solution is formed and Q = Ksp. (Ppt does form)
Exercise 18.17, p. 771
Solid PbI2 (Ksp = 9.8 x 10-9) is placed in a beaker of water. After a period of time, the lead(II) concentration is measured and found to be 1.1 x 10-3 M. Has the system yet reached equilibrium? That is, is the solution saturated? If not, will more PbI2 dissolve?
Exercise 18.18, p. 886
Will SrSO4 precipitate from a solution containing 2.5 x 10-4 M strontium ion, Sr2+, if enough of the soluble salt Na2SO4 is added to make the solution 2.5 x 10-4 M in SO4-2? (Ksp for SrSO4 is 3.4 x 10-7.)
Exercise 18.19, p. 887
What is the minimum concentration of I-1 that can cause precipitation of PbI2 from a
0.050 M solution of Pb(NO3)2? Ksp for PbI2 is 9.8 x 10-9.
Exercise 18.20, p. 887
You have 100.0 mL of 0.0010 M silver nitrate. Will AgCl precipitate if you add 5.0 mL of 0.025 M HCl? (Ksp of AgCl = 1.8 x 10-10)
18.6 Solubility and Complex Ions, p. 887
(CD-Rom, Screen 18.16)
See p. 888 and Appendix K for Kformation.
See p. 889, Figure 18.16
(First Reaction taking place.)
Provide K for the following reaction:
AgCl(s) + 2 NH3(aq) ó Ag(NH3)2(aq)+1 + Cl-1(aq)
_______________________________________Ksp( AgCl) = 1.8 x 10-10
_______________________________________ Kformation ([Ag(NH3)2]+) = 1.6 x 107
Was the solubility of silver increased by adding ammonia?
Exercise 18.21, p. 890
Calculate the value of the equilibrium constant, Knet, for dissolving Cu(OH)2 in aqueous ammonia to form the complex ion [Cu(NH3)4]2+? (See Figure 17.7)
18.7 Solubility, Ion Separations, and Qualitative Analysis
(CD-Rom, Screen 18.17)
See p. 890, Figure 18.19
Ions in Solution
Ag+, Pb+2, Cu+2
(HCl is added, AgCl and PbCl2 are precipitated. Cu+2 is the blue color when dissolved in water.)
Separation scheme is determined by analyzing Ksp values for salts of cations with various anions.
Exercise 18.22, p. 892 (modified)
The cations of each of the following pairs appear together in one solution:
a) Ca+2 and Pb+2
b) Fe+2
and K+
You may add only one reagent to precipitate one cation and not the other. Consult the solubility product table in Appendix J (and Figure 5.3, p.179) and tell whether you would use Cl-1, S-2, or OH-1 as the precipitating ion in each case. (The precipitating ions are introduced in the form of HCl, (NH4)2S, or NaOH, for example.)
Equilibria in the Environment: Carbon Dioxide and Carbonates (Removed from 6th ed.)
CO2(g) + H2O(l) ó H2CO3(aq)
H2CO3(aq) + H2O(l) ó H3O+1(aq) + HCO3-1(aq) Ka = 4.2 x 10-7
These reactions explain why rain water is slightly ___________ with a pH of ________.
Oceans contain large quantities of calcium carbonate.
CaCO3(s) ó Ca+2(aq) + CO3-2(aq) Ksp = 3.4 x 10-9
CO3-2(aq) + H2O(l) ó OH-1(aq) +
HCO3-1(aq) Kb = 2.1 x 10-4
These reactions explain why sea water is slightly ___________ with a pH of ________.
One interesting aspect of the carbonate equilibrium system (HCO3-1/CO3-2) is that it acts as a____________________ system. Thus pH remains fairly constant even when an acid is added such as when there is undersea volcanic activity.