BaSO 4 (s) Ba 2+ (aq) + SO 4 2– (aq). The Solubility Product Constant, K sp. Many important ionic compounds are only slightly soluble in water (we used to call them “insoluble” – Chapter 4).
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Ksp= [ Ba2+ ][ SO42–]
Write a solubility product constant expression for equilibrium in a saturated aqueous solution of the slightly soluble salts (a) iron(III) phosphate, FePO4, and (b) chromium(III) hydroxide, Cr(OH)3.
At 20 °C, a saturated aqueous solution of silver carbonate contains 32 mg of Ag2CO3 per liter of solution. Calculate Kspfor Ag2CO3 at 20 °C. The balanced equation is
Ag2CO3(s) 2 Ag+(aq) + CO32–(aq) Ksp= ?
From the Ksp value for silver sulfate, calculate its molar solubility at 25 °C.
Ag2SO4(s) 2 Ag+(aq) + SO42–(aq)
Ksp= 1.4 x 10–5 at 25 °C
Without doing detailed calculations, but using data from Table 16.1, establish the order of increasing solubility of these silver halides in water: AgCl, AgBr, AgI.
When Na2SO4(aq) is added to the saturated solution of Ag2SO4 …
… [Ag+] attains a new, lower equilibrium concentration as Ag+ reacts with SO42–to produce Ag2SO4.
Calculate the molar solubility of Ag2SO4 in 1.00 M Na2SO4(aq).
Qip andQc: new look, same great taste!
If 1.00 mg of Na2CrO4 is added to 225 mL of 0.00015 M AgNO3, will a precipitate form?
Ag2CrO4(s) 2 Ag+(aq) + CrO42–(aq)
Ksp= 1.1 x 10–12
If 0.100 L of 0.0015 M MgCl2 and 0.200 L of 0.025 M NaF are mixed, should a precipitate of MgF2 form?
MgF2(s) Mg2+(aq) + 2 F–(aq) Ksp= 3.7 x 10–8
Pictured here is the result of adding a few drops of concentrated KI(aq) to a dilute solution of Pb(NO3)2. What is the solid that first appears? Explain why it then disappears.
To a solution with [Ca2+] = 0.0050 M, we add sufficient solid ammonium oxalate, (NH4)2C2O4(s), to make the initial [C2O42–] = 0.0051 M. Will precipitation of Ca2+ as CaC2O4(s) be complete?
CaC2O4(s) Ca2+(aq) + C2O42–(aq) Ksp= 2.7 x 10–9
AgCl(s) Ag+(aq) + Cl–(aq)Effect of pH on Solubility—fly in the Ointment
Added H+ reacts with, and removes, F–; LeChâtelier’s principle says more F– forms.
H+ does not consume Cl– ; acid does not affect the equilibrium.
What is the molar solubility of Mg(OH)2(s) in a buffer solution having [OH–] = 1.0 x 10–5 M, that is, pH = 9.00?
Mg(OH)2(s) Mg2+(aq) + 2 OH–(aq) Ksp= 1.8 x 10–11
Another one of those “explain” problems
Without doing detailed calculations, determine in which of the following solutions Mg(OH)2(s) is most soluble:
(a) 1.00 M NH3
(b) 1.00 M NH3 /1.00 M NH4+
(c) 1.00 M NH4Cl.
Bromphenol blue, bromthymol blue, and phenolphthalein all change color at very nearly 20.0 mL
At about what volume would we see a color change if we used methyl violet as the indicator?
Calculate the pH at the following points in the titration of 20.00 mL of 0.500 M HCl with 0.500 M NaOH:
H3O+ + Cl– + Na+ + OH– Na+ + Cl– + 2 H2O
(a) before the addition of any NaOH
(b) after the addition of 10.00 mL of 0.500 M NaOH
(c) after the additionof 20.00 mL of 0.500 M NaOH
(d) after the addition of 20.20 mL of 0.500 M NaOH
The equivalence-point pH is NOT 7.00 here. Why not??
Bromphenol blue was ok for the strong acid/strong base titration, but it changes color far too early to be useful here.