11/30/2009
Solubility
•
Molar solubility (s) (mol/L) is the number of moles
of solute dissolved in 1 L of a saturated solution
•
Solubility (g/L) is the number of grams of solute
dissolved in 1 L of a saturated solution
The Solubility Product Constant
•
Chapter 7.6
General Solubility Rules:
Ionic Compounds in Water
• Single + and – ions are usually water soluble
– Na+, K+, Cl-, NO3-
– Ca2+, Fe3+, CO32-, PO43– Major exception = Mg2+
Solubility is a relative concept: depends on the ratio
of solute to solvent
–
Soluble: > 1.0 g/100 mL of solution
–
Insoluble: < 1.0 g/100 mL of solution
–
Solubility depends on formula weight therefore
quantitative expression related to equilibrium is useful
Specific Solubility Rules:
Ionic Compounds in Water
• Alkali metal salts are always soluble
• Nitrate, acetate, ammonium salts are always soluble
– Ca(NO3)2 Fe(NO3)3 (NH4)2S
• Halide salts are mostly soluble (XF, XCl, XBr, XI)
– Exceptions: AgX, PbX2, HgX, some fluorides
• When ions that can have strong attractions are
present, precipitates will form
• Precipitates can be dissolved if other
attractions come into play
• Sulfates salts are mostly soluble (XySO4)
– Exceptions: SrSO4 BaSO4 PbSO4
• Most carbonates (CO32-), hydroxides (OH-) , sulfides
(S2-) and phosphates (PO43-) are insoluble
– Except ones of alkali metals and NH4+ and some alkaline
earth metals
Solubility Equilibria
Solubility Equilibria
•
Solute can be a molecular substance, dissolved salt (both
ions) or a simple cation or anion or complex cation or
anion
– Li3PO4, Na2O, K2SO4, Rb2S, Cs2CO3,
• Double and triple + and – ions are usually
water insoluble
•
–
Ag+ (aq) + Cl- (aq)
Solubility product constant, Ksp, is the way of
expressing solubility in terms of an equilibrium
constant
AgCl (s)
Rules for Ksp are same as for all equilibrium
constants
MgF2 (s)
Ksp = [Ag+][Cl-]
Mg2+ (aq) + 2F- (aq)
Ksp = [Mg2+][F-]2
– No solids! (Therefore no denominator!!)
– Molar solubility of a salt is not always the same
as the molar solubility of the ions in the salt!
–
[Products]/ [Reactants]  Keq
[X+][Y-] Ksp
Ag2CO3 (s)
2Ag+ (aq) + CO32- (aq)
Ksp = [Ag+]2[CO32-]
Ca3(PO4)2 (s)
3Ca2+ (aq) + 2PO43- (aq)
Ksp = [Ca2+]3[PO43-]2
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11/30/2009
The molar solubility of Pb3(PO4)2 is 6.2 x 10-12 mol/L.
Calculate the Ksp value
s = 6.2 x 10-12 M
6.2 x 10-12 M Pb3(PO4)2 x 3 moles Pb 2+/mole Pb3(PO4)2
= 18.6 x 10-12 M Pb 2+
• Molar solubility of “entire” salt will always be less
than or equal to the “molar solubilities” of the
component ions
• For substances with multiple ions per formula unit
have to multiply AND raise by same power!
• Have to know how to take apart ionic substances
What is the solubility of silver chloride in g/L if
Ksp = 1.6 x 10-10?
AgCl (s)
Initial (M)
Change (M)
Equilibrium (M)
[Ag+] = 1.3 x 10-5 M
Solubility of AgCl =
Ag+ (aq) + Cl- (aq)
0.00
0.00
+s
+s
s
s
[Cl-] = 1.3 x 10-5 M
Ksp = 1.6 x 10-10
Ksp = [Ag+][Cl-]
Ksp = s2
s = Ksp
s = 1.3 x 10-5
1.3 x 10-5 mol AgCl 143.35 g AgCl
x
= 1.9 x 10-3 g/L
1 L soln
1 mol AgCl
6.2 x 10-12 M Pb3(PO4)2 x 2 moles PO43-/mole Pb3(PO4)2
= 12.4 x 10-12 M PO4 3Ksp = [Pb2+]3[PO43-]2 = [18.6 x 10-12]3[12.4 x 10-12]2
= 9.9 x 105 x 10-60 = 9.9 x 10-55
Predicting Precipitation from Solution
• Reaction Quotient (Q) review
– Q gives status of reaction in terms of
concentrations present
– If K > Q  equilibrium has not been reached and
reaction will go forward to products
– If K < Q  reaction is past equilibrium and
reaction will go from products to reactants
Q < Ksp
Unsaturated solution
No precipitate
Q = Ksp
Saturated solution
Q > Ksp
Supersaturated solution
Precipitate
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11/30/2009
If 2.00 mL of 0.200 M NaOH are added to 1.00 L of
0.100 M CaCl2, will a precipitate of Ca(OH)2form?
Ksp of Ca(OH)2 = 8.0 x 10-6
Is Q > Ksp for Ca(OH)2?
Change in volume negligible
(1.000 L to 1.002 L)
[Ca2+]0 = 0.100 M
[OH-]0 = 4.0 x 10-4 M
(0.002 L x .200 moles/L in
1.00L)
Q = [Ca2+]0[OH-]02 = 0.10 x (4.0 x 10-4)2 = 1.6 x 10-8
Q < Ksp
No precipitate will form
What is the molar solubility of AgBr in (a)
pure water and (b) 0.0010 M NaBr?
NaBr (s)
Na+ (aq) + Br- (aq)
AgBr (s)
Ag+ (aq) + Br- (aq)
Ksp = [Ag+][Br-] = s2
[Br-] = 0.0010 M
[Ag+] = s
Ksp =
Ksp = 7.7 x 10-13
s = 8.8 x 10-7 M (solubility in water)
AgBr (s)
Ag+ (aq) + Br- (aq)
[Br-] = 0.0010 + s
[Ag+][Br-]
The Common Ion Effect
• The presence of a common ion decreases the
solubility of the salt
XY (s)  X+ (aq) + Y- (aq)
• Increasing either ion suppresses the amount of
dissolved salt (in terms of paired ions in solution)
• Increasing concentration of one ion decreases
concentration of the other ion
– Same idea as with pH
– As H+ goes down, OH- goes up and vice versa
• Hard wire the Principle of LeChatelier!!
Summary of Pressure, Temperature
Affect on Solubility
H (s, l or g)
• (+) Endothermic
• (+) Endothermic
Temp


Direction
 Prod
 React
Solubility
 increase
 decrease
• (-) Exothermic
• (-) Exothermic


 React
 Prod
 decrease
 increase
• Gas solute
• Gas solute
Pressure


Direction
 Prod
 React
Solubility
 increase
 decrease
0.0010
= [s][0.0010]
s = 7.7 x 10-10M (solubility with common ion)
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