South Pasadena • AP Chemistry
Name
3 ▪ Chemical Equilibrium
Period
3.3
PROBLEMS
1. Addition of a strong acid would increase the
solubility of which of the following salts?
AgCl
CaSO4
CdS
CaCO3
PbBr2
CaHPO4
Cd(OH)2
AuCl
2. What is the [Mg2+] in a saturated solution of
magnesium fluoride, MgF2 if its solubility product
constant is 6.4 × 10−9? What is the [Mg2+] if the
solution also contains 0.30 M NaF?
(a) [Mg2+][F−]2 = Ksp
(s)(2s)2 = 6.4 × 10−9
s = 0.0012 M
[Mg2+] = 0.0012 M
(b) [Mg2+][F−]2 = Ksp
(s)(0.30 + 2s)2 = 6.4 × 10−9
Assume 2s << 0.30
2
−9
(s)(0.30) = 6.4 × 10
s = 7.1 × 10−8 M
[Mg2+] = 7.1 × 10−8 M
3. Calculate the molar solubility of cobalt(II) sulfide
in a solution that contains 0.030 M cobalt(II)
chloride. Ksp for cobalt(II) sulfide CoS is
5 × 10−22.
[Co2+][S2−] = Ksp
(0.030 + s)(s) = 5 × 10−22
(0.030)(s) = 5 × 10−22
s = 1.7 × 10−20 M
Assume s < 0.030
–
Date
COMMON IONS
4. Calculate the solubility of silver phosphate,
Ag3PO4, in (a) pure water, and (b) in a solution
with 0.020 M AgNO3. Ksp for Ag3PO4 is
1.8 × 10−18. (Ch. 19, #52)
(a) [Ag+]3[PO43−] = Ksp
(3s)3(s) = 1.8 × 10−18
27s4 = 1.8 × 10−18
s = 1.6 × 10−5 M
(b) [Ag+]3[PO43−] = Ksp
(0.020 + 3s)3(s) = 1.8 × 10−18
Assume 3s << 0.020
3
(0.020) (s) = 1.8 × 10−18
s = 2.3 × 10−13 M
5. A solution contains 0.10 M iodide ion, I−, and
0.10 M carbonate ion, CO32−. Ksp for PbI2 and
PbCO3 are 1.4 × 10−8 and 1.5 × 10−15, respectively.
a. If solid Pb(NO3)2 is slowly added to the
solution, which salt precipitates first, PbI2 or
PbCO3?
b. What is the concentration of the first ion that
precipitates (CO32− or I−) when the second or
more soluble salt begins to precipitate?
(Ch. 19, #56)
(a) Find [Pb2+] when salt precipitates.
Ksp = [Pb2+][I−]2
Ksp 1.4 × 10−8
[Pb2+] = − 2 =
= 1.4 × 10−6 M
[I ]
(0.10)2
Ksp = [Pb2+][CO32−]
Ksp
1.5 × 10−15
[Pb2+] =
= 1.5 × 10−14 M
2− =
[CO3 ]
0.10
PbCO3 will precipitate first.
(b) When PbI2 precipitates, [Pb2+] =
1.4 × 10−6 M
Find [CO32−].
Ksp
1.5 × 10−15
[CO32−] =
=
[Pb2+] 1.4 × 10−6
−9
= 1.1 × 10 M
AP Chemistry 2001 #1
Answer the following questions relating to the solubility of the chlorides of silver and lead.
(a) At 10°C, 8.9 × 10–5 g of AgCl(s) will dissolve in 100. mL of water.
(i) Write the equation for the dissociation of AgCl(s) in water.
AgCl (s)  Ag+ (aq) + Cl− (aq)
(ii) Calculate the solubility, in mol L–1, of AgCl(s) in water at 10°C.
s=
8.9 × 10−5 g AgCl  1 mol AgCl  1000 mL
= 6.2 × 10−6 M
100 mL solution 143.32 g AgCl  1 L 
(iii) Calculate the value of the solubility-product constant, Ksp, for AgCl(s) at 10°C.
Ksp = [Ag+][Cl−] = (s)(s) = (6.2 × 10−6)2 = 3.9 × 10−11
(b) At 25°C, the value of Ksp for PbCl2(s) is 1.6 × 10–5 and the value of Ksp for AgCl(s) is 1.8 × 10–10.
(i) If 60.0 mL of 0.0400 M NaCl(aq) is added to 60.0 mL of 0.0300 M Pb(NO3)2(aq), will a precipitate form?
Assume that volumes are additive. Show calculations to support your answer.
[Cl−] =
(0.0400 M)(60.0 mL)
= 0.0200 M
(120.0 mL)
[Pb2+] =
(0.0300 M)(60.0 mL)
= 0.0150 M
(120.0 mL)
Q = [Pb2+][Cl−]2 = (0.0150)(0.0200)2 = 6.0 × 10−6
No precipitate forms because Q < Ksp.
(ii) Calculate the equilibrium value of [Pb2+(aq)] in 1.00 L of saturated PbCl2 solution to which 0.250 mole of
NaCl(s) has been added. Assume that no volume change occurs.
Ksp = [Pb2+][Cl−]2
1.6 × 10−5 = (s)(0.250 + 2s)2
Assume 2s << 0.250
s = 2.6 × 10−4 M
[Pb2+] = 2.6 × 10−4 M
(iii) If 0.100 M NaCl(aq) is added slowly to a beaker containing both 0.120 M AgNO3(aq) and 0.150 M
Pb(NO3)2(aq) at 25°C, which will precipitate first, AgCl(s) or PbCl2(s)? Show calculations to support
your answer.
Find [Cl−] to precipitate AgCl:
Ksp 1.8 × 10−10
[Cl−] =
=
= 1.5 × 10−9 M
[Ag+]
0.120
Find [Cl−] to precipitate PbCl2:
Ksp
1.6 × 10−5
[Cl−] =
= 0.010 M
2+ =
[Pb ]
0.150
AgCl will precipitate first.