Chapter 18 Solubility. Equilibria of Slightly Soluble Ionic Compounds Explore the aqueous equilibria of slightly soluble ionic compounds. Chapter 5. Precipitation.
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Equilibria of Slightly Soluble Ionic Compounds
• Explore the aqueous equilibria of slightly soluble ionic compounds.
• Chapter 5. Precipitation Reactions:
AgNO3(aq) + NaCl(aq) ??
• Want to consider quantitative predictions
The Formation & Dissolution of Precipitates
• Solubility: maximum amount of solute that will dissolve in a given amount of solvent (depends on solvent, temperature and pressure)
No compound is infinitely soluble andno compound is perfectly insoluble.
Solute Solubility
(g solute/100 g solvent)
Qualitative Solubility
Description
Less than 0.1 Insoluble
0.1 – 1 Slightly soluble
1 – 10 Soluble
Greater than 10 Very soluble
The Formation & Dissolution of Precipitates
• Saturated solution: contains maximum concentration of solute– Equilibrium between undissolved and dissolved solute.
• Solutes (even those called “soluble”) have a limited solubility in a particular solvent.
• Slightly soluble (often called “insoluble”) ionic compounds have a relatively low solubility – Reach equilibrium with little solute dissolved– Heterogeneous equilibrium
Why is this important?
Dissolving and Precipitation occurs around us:– Tooth enamel dissolves in acidic soln (tooth decay)
– Ppt of certain salts in kidneys causes kidney stones
– Waters of Earth contains dissolved salts as water passes over and through the ground
– Ppt of CaCO3 from groundwater is responsible for cave formation.
Let’s look at the factors that affect solubility!
Solubility-Product Constant (Ksp)
• Solubility-product constant (Ksp): equilibrium constant for equilibrium between slightly soluble ionic solid and a solution of its ions– Indicates how soluble the solid is in water
• Solubility: quantity that dissolves to form a saturated solution (g/L)
• Molar solubility: number of moles of solute that dissolves in forming a liter of saturated solution of solute (mol/L)
• Solubility depends on concentrations of other ions and pH but Ksp is a constant.
Solubility-Product Constant (Ksp)
Practice: Write an ionic equation for the dissolution, and the equation for the solubility product for:
(a) Calcium carbonate
(b) Magnesium hydroxide
(c) Ag3PO4
• Magnitude of Ksp is measure of how far to the right dissolution proceeds at equilibrium (saturation).
– Used to compare solubilities if same total number of ions
Ksp of Selected Ionic Compounds (25 °C)
Name, Formula Ksp
Aluminum hydroxide, Al(OH)3 3 x 10-34
Cobalt(II) carbonate, CoCO3 1.0 x 10-10
Iron(II) hydroxide, Fe(OH)2 4.1 x 10-15
Lead(II) fluoride, PbF2 3.6 x 10-8
Lead(II) sulfate, PbSO4 1.6 x 10-8
Mercury(I) iodide, Hg2I2 4.7 x 10-29
Silver sulfide, Ag2S 8 x 10-48
Zinc iodate, Zn(IO3)2 3.9 x 10-6
See Appendix D in your book for a much more extensive list.
Example
Predict which of the following compounds will
have the greatest molar solubility in water
A) AgCl Ksp = 1.8 x 10-10
B) AgBr Ksp = 5.0 x 10-13
C) AgI Ksp = 8.3 x 10-17
D) all have the same molar solubility
Solubilities and Solubility Products
• Ksp for a slightly soluble solid can be determined from its solubility – as long as there is no other reaction
• Example 1: Lead(II) sulfate (PbSO4) is a key component in lead-acid car batteries. Its solubility in water at 25 °C is 4.25 x 10-3 g/100 mL solution. What is the Ksp of PbSO4?
• Example 2: Determine the molar solubility of MgF2 from its solubility product (Ksp = 6.4 x 10-9).
Factors that Affect Solubility: Common Ion Effect
•The presence of a common ion decreases the solubility of a slightly soluble ionic compound.
•The shift in equilibrium that occurs because of the addition of an ion already involved in the equilibrium reaction.
AgCl(s) ⇌ Ag+(aq) + Cl(aq)
adding NaCl( ) shifts equilibrium positionaq
The effect of a common ion on solubility
PbCrO4(s) Pb2+(aq) + CrO42-(aq) PbCrO4(s) Pb2+(aq) + CrO4
2-(aq)
CrO42- added
Example
The solubility of Ca(OH)2 in water is 0.012 M.
What is its solubility in 0.10 M Ca(NO3)2?
Ksp of Ca(OH)2 is 6.5 x 10-6
Le Châtelier’s Principle
Determine the effects of solubility when each of
the following is added to a mixture of the
slightly soluble solid NiCO3 and water at
equilibrium:
(a)Ni(NO3)2 (c) K2CO3
(b)KClO4 (d) HNO3
Effect of pH on solubility
• [H3O+] can have a profound effect on the solubility of an ionic compound.
• Solubility of slightly soluble salts containing basic anions increases as [H+] increases– More basic anion…more solubility is influenced by pH
• Predict the effect on solubility of adding a strong acid
CaCO3(s) ⇌ Ca2+(aq) + CO3
2-(aq)
AgCl(s) ⇌ Ag+(aq) + Cl-
(aq)
Ca(OH)2(s) ⇌ Ca2+(aq) + 2OH-
(aq)
If limestone (CaCO3) deposit is well below the surface….caves
A view inside Carlsbad Caverns, New Mexico
Predicting the Formation of a Precipitate
Compare Qsp to Ksp to predict if a precipitate will form
and, if not, what concentrations of ions will cause it to
do so.
Qsp = Ksp soln is saturated & no changes occur
Qsp > Ksp ppt forms until soln is saturated
Qsp < Ksp soln is unsaturated & no ppt forms
Practice
• Example 1: Determine whether CaHPO4 will precipitate from a solution with [Ca2+] = 0.0001 M and [HPO4
2-] = 0.001 M.
• Example 2: Does silver chloride precipitate when equal volumes of a 2 x 10-4 M solution of AgNO3 and a 2 x 10-4 M solution of NaCl are mixed.
AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq)
Practice
Will a precipitate form when 0.10 L of
8.0 x 10-3 M Pb(NO3)2 is added to 0.40 L of
5.0 x 10-3 M Na2SO4?
Ksp for PbSO4 = 6.3 x 10-7
Concentration Necessary to Form a Ppt
• We can also determine the concentration of an ion necessary for precipitation to begin.
• Assume that precipitation begins when Qsp = Ksp
• Example: If a solution contains 0.0020 mol CrO42-
per liter, what concentration of Ag+ ion must be added as AgNO3 before Ag2CrO4 begins to precipitate. (Neglect any increase in volume upon adding the solid silver nitrate.)
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