18. Aqueous Equilibrium

Ksp - Molar Solubility, Ice Tables, & Common Ion Effect

A 1.00-L solution saturated at 25 C with lead(II) iodide 
contains 0.54 g of PbI2. Calculate the solubility-product 
constant for this salt at 25 C.

Tooth enamel is composed of hydroxyapatite, whose simplest
formula is Ca51PO423OH, and whose corresponding
Ksp = 6.8 * 10-27. As discussed in the Chemistry and Life
box on page 746, fluoride in fluorinated water or in toothpaste
reacts with hydroxyapatite to form fluoroapatite,
Ca51PO423F, whose Ksp = 1.0 * 10-60. (a) Write the expression
for the solubility-constant for hydroxyapatite and
for fluoroapatite.

Ksp and solubility

(b) Write the expression for the solubility-product constant for each 
of the following ionic compounds: Cu31PO422.

Refer to the Ksp values in Table 17.2 to calculate the molar solubility
of each compound in pure water.
a. AgBr

The following pictures represent solutions of Ag2CrO4, which also may contain ions other than Ag+ and CrO42- that are not shown.
Solution 1 is in equilibrium with solid Ag2CrO4. Will a precipitate of solid Ag2CrO4 form in solutions 2-4? Explain. <img src="https://lightcat-files.s3.amazonaws.com/problem_images/27e3327f27717551-1673023496839.jpg" alt="" /> 
(1) <img src="" alt="" />
(2) <img src="" alt="" />
(3) <img src="" alt="" />
(4) <img src="" alt="" />

Use Le Châtelier’s principle to predict whether the solubility 
of BaF2 will increase, decrease, or remain the same on 
addition of each of the following substances. 
(a) HCl

Consider the equilibrium
B1aq2 + H2O1l2 Δ HB+1aq2 + OH-1aq2.
Suppose that a salt of HB+1aq2 is added to a solution of
B1aq2 at equilibrium. (c) Will the pH of the solution increase, decrease, or stay the same?

The precipitation of Al(OH)3 (Ksp) = 1.3x10-33) is sometimes used ot purify water. (a) Estimate the pH at which precipitation of Al(OH)3 will begin if 5.0 lb of Al2(SO4)3 is added to 2000 gal of water

Use the given molar solubilities in pure water to calculate Ksp for
each compound.
b. Ag2SO3; molar solubility = 1.55 * 10- 5 M

Calculate the molar solubility of PbCrO4 in: 
(a) Pure Water
(d) 1.0 x 10^-3 M K2CrO4

Calculate the molar solubility of MX (Ksp = 1.27 * 10 - 36) in
each liquid or solution.
b. 0.25 M MCl2

In which of these solutions will HNO2 ionize less than it does in 
pure water? 
a. 0.10 M NaCl 
b. 0.10 M KNO3 
c. 0.10 M NaOH 
d. 0.10 M NaNO2

17.4 Solubility and Ksp

Some progressive hair coloring products marketed to men, such as Grecian Formula 16, contain lead acetate Pb(CH3CO2)2. As the coloring solution is rubbed on the hhair, the Pb2+ ions react with the sulfur atoms in hair proteins to give lead(II) sulfide (PbS), which is black. A typical coloring solution contains 0.3 mass% Pb(CH3CO2)2, and about 2 mL of the solution is used per application. 
(b) Suppose the hair is washed with shampoo and water that has pH = 5.50. How many washings would be required to remove 50% of the black color? Assume that 3 gal of water is used per washing and that the water becomes saturated with PbS. 
(c) Does the calculated number of washings look reason-able, given that frequent application of the coloring solution is recommended? What process(es) in addition to dissolution might contribute to the loss of color?

(a) I f t he molar solubility of CaF2 at 35 C i s 
1.24 * 10-3 mol>L, what is Ksp at this temperature?

Use the following solubility data to calculate a value of Ksp for each compound. 
(d) Zn(CN)2; 4.95 x 10^-4 g/L

(b) If 0.0490 g of AgIO3 dissolves per liter of solution, calculate the solubility-product constant.

Calculate the solubility of LaF3 in grams per liter in (b) 0.010 M KF solution.

Use the given molar solubilities in pure water to calculate Ksp for
each compound.
a. BaCrO4; molar solubility = 1.08 * 10- 5 M

The solubility-product constant for barium permanganate, 
Ba1MnO422, is 2.5 * 10-10. Assume that solid Ba1MnO422 is 
in equilibrium with a solution of KMnO4. What concentration 
of KMnO4 is required to establish a concentration of 
2.0 * 10-8 M for the Ba2 + ion in solution?

The solubility of CaCO3 is pH dependent. (b) Use the Kb
expression for the CO32 - ion to determine the equilibrium
constant for the reaction
CaCO31s2 + H2O1l2 ΔCa2 + 1aq2 + HCO3-1aq2 + OH-1aq2

Ksp and pH

Write balanced equations and expressions for Ksp for the dissolution 
of each ionic compound. 
b. PbBr2

Consider a beaker containing a saturated solution of CaF2 
in equilibrium with undissolved CaF21s2. Solid CaCl2 is 
then added to the solution. (a) Will the amount of solid 
CaF2 at the bottom of the beaker increase, decrease, or remain 
the same?

Can Fe2+ be separated from Sn2+ by bubbling H2S through a 0.3 M HCl solution that contains 0.01 M Fe2+ and 0.01 M Sn2+? A saturated solution of H2S has [H2S] ≈ 0.10 M. Values of Kspa are 6 x 10^2 for FeS andd 1 x 10^-5 for SnS.

A particular saturated solution of Ca3(PO4)2 has [Ca2+] = [PO43-] = 2.9 x 10^-7 M. 
(c) What is [PO43-] in a saturated solution that has [Ca2+] = 0.010 M?

Consider the compounds with the generic formulas listed and 
their corresponding molar solubilities in pure water. Which 
compound has the smallest value of Ksp? 
a. AX; molar solubility = 1.35 * 10 - 4 M 
b. AX2; molar solubility = 2.25 * 10- 4 M 
c. A2X; molar solubility = 1.75 * 10 - 4 M

Use the following solubility data to calculate a value of Ksp for each compound. 
(a) SrF2: 1.03 x 10^-3 M

Determine the minimum concentration of the precipitating 
agent on the right to cause precipitation of the cation from the 
solution on the left. 
b. 0.085 M CaI2; K2SO4

Predict whether a precipitate will form if you mix 75.0 mL of a 
NaOH solution with pOH = 2.58 with 125.0 mL of a 0.018 M 
MgCl2 solution. Identify the precipitate, if any.

Use the given molar solubilities in pure water to calculate Ksp for
each compound.
a. MX; molar solubility = 3.27 * 10 - 11 M

Suppose that a 10-mL sample of a solution is to be tested for 
I- ion by addition of 1 drop (0.2 mL) of 0.10 M Pb1NO322. 
What is the minimum number of grams of I- that must be 
present for PbI21s2 to form?

(a) True or false: “solubility” and “solubility-product constant” 
are the same number for a given compound.

Will FeS precipitate in a solution that is 0.10 M in Fe(NO3)2, 0.4 M in HCl, and 0.10 M in H2S? Will FeS precipitate if the pH of the solution is adjusted to pH 8 with an NH4+ - NH3 buffer? Kspa = 6 x 10^2 for FeS.

Whenever the question refers to the solubility then that value represents our &#8220;x&#8221;&#160;variable.&#160;

Calculate the solubility of LaF3 in grams per liter in (c) 0.050 M LaCl3 solution.

Associated with any ionic compound is the solubility product constant, which&#160;measures how soluble the compound&#160;will be in a solvent.&#160;

A concentration of 10–100 parts per billion (by mass) of Ag+ 
is an effective disinfectant in swimming pools. However, if 
the concentration exceeds this range, the Ag+ can cause 
adverse health effects. One way to maintain an appropriate 
concentration of Ag+ is to add a slightly soluble salt to 
the pool. Using Ksp values from Appendix D, calculate the 
equilibrium concentration of Ag+ in parts per billion that 
would exist in equilibrium with (c) AgI.

The solubility of two slightly soluble salts of M2 + , MA
and MZ2, is the same, 4 * 10-4 mol/L. (c) If you added an equal
volume of a solution saturated in MA to one saturated in
MZ2, what would be the equilibrium concentration of the
cation, M2+?

A saturated solution of an ionic salt MX exhibits an osmotic pressure of 74.4 mm Hg at 25 °C. Assuming that MX is completely dissociated in solution, what is the value of its Ksp?

Use Le Châtelier’s principle to predict whether the solubility 
of BaF2 will increase, decrease, or remain the same on 
addition of each of the following substances. 
(d) Ba1NO322

Write balanced equations and expressions for Ksp for the dissolution 
of each ionic compound. 
c. Ag2CrO4

A 1.00-L solution saturated at 25 C with calcium oxalate 
1CaC2O42 contains 0.0061 g of CaC2O4. Calculate the 
solubility-product constant for this salt at 25 C.

Ca1OH22 has a Ksp of 6.5 * 10-6. (b) If 50 mL of the solution from part (a) is added to each of the beakers shown here, in which beakers, if any, will a precipitate form? In those cases where a precipitate forms, what is its identity? [Section 17.6]
<img src="https://lightcat-files.s3.amazonaws.com/problem_images/f286e1c0f4397451-1667321555336.jpg" alt="" />

Calculate the solubility of Mn1OH22 in grams per liter 
when buffered at pH (b) 9.5.

Whenever we deal with the solubility product constant we use an ICE Chart and since&#160;the reactant will always be a&#160;solid then we can ignore it.&#160;

Use the values of Ksp in Appendix C to calculate the solubility of the following compounds (in g/L): 
(c) Cu3(PO4)2

Use Le Châtelier’s principle to predict whether the solubility
of BaF2 will increase, decrease, or remain the same on
addition of each of the following substances.
(c) NaNO3

Calculate the molar solubility of barium fluoride in each liquid 
or solution. 
a. pure water

Write balanced equations and expressions for Ksp for the dissolution
of each ionic compound.
a. BaSO4

The following graphs represent the behavior of BaCO3 under
different circumstances. In each case, the vertical axis
indicates the solubility of the BaCO3 and the horizontal
axis represents the concentration of some other reagent.
(a) Which graph represents what happens to the solubility
of BaCO3 as HNO3 is added? [Section 17.5]
<img src="https://lightcat-files.s3.amazonaws.com/problem_images/9ab13c3abfc1126e-1667321273878.jpg" alt="" />

A solution containing sodium fluoride is mixed with one containing 
calcium nitrate to form a solution that is 0.015 M in NaF 
and 0.010 M in Ca(NO3)2. Does a precipitate form in the mixed 
solution? If so, identify the precipitate.

Use the values of Ksp in Appendix C to calculate the molar solubility of the following compounds: 
(b) Mg(OH)2

Determine the minimum concentration of the precipitating 
agent on the right to cause precipitation of the cation from the 
solution on the left. 
a. 0.035 M Ba(NO3)2; NaF

Two compounds with general formulas AX and AX2 have 
Ksp = 1.5 * 10 - 5. Which of the two compounds has the higher 
molar solubility?

For each pair of compounds, use Ksp values to determine 
which has the greater molar solubility: (b) PbCO3 or BaCrO4.

The solubility of copper(I) chloride is 3.91 mg per 100.0 mL of 
solution. Calculate Ksp for CuCl.

What is Ksp? (Solubility Product Constant)

Is the solubility of Zn(OH)2 increased, decreased, or unchanged on addition of each of the following substances? Write a balanced net ionic equation for each dissolution reaction. (See Appendix C.6 for formulas of complex ions.) 
(a) HCl

Write the expression for the solubility product constant of MgF2 (see Problem 4.139). If [Mg2+] = 2.6 * 10-4 mol/L in a solution, what is the value of Ksp?

The solubility product constants of PbSO4 and SrSO4 are 
6.3 * 10-7 and 3.2 * 10-7, respectively. What are the values 
of 3SO4 
2 - 4, 3Pb2 + 4, and 3Sr2 + 4 in a solution at equilibrium 
with both substances?

Lead(II) carbonate, PbCO3, is one of the components of the
passivating layer that forms inside lead pipes. (d) The EPA threshold for acceptable levels of lead ions in water is 15 ppb. Does a saturated solution of lead(II) carbonate produce a solution that exceeds the EPA
limit?

Calculate the solubility of LaF3 in grams per liter in (a) pure
water.

A 0.0100 mol sample of solid Cd(OH)2 (Ksp = 5.3 x 10^-15) in 100.0 mL of water is titrated with 0.100 M HNO3. 
(c) How many milliliters of 0.100 M HNO3 must be added to completely neutralize the Cd(OH)2?

Does the pH increase, decrease, or remain the same when the substances are added to the solutions? 
(a) LiF to an HF solution
(b) KI to an HI solution 
(c) NH4Cl to an NH3 solution

Solubility Product Constant (Ksp)

Which of these statements about the common-ion effect is
most correct? (a) The solubility of a salt MA is decreased in
a solution that already contains either M+ or A-. (b) Common
ions alter the equilibrium constant for the reaction of
an ionic solid with water. (c) The common-ion effect does
not apply to unusual ions like SO32 - . (d) The solubility of a salt MA is affected equally by the addition of either A- or a noncommon ion.

What is the molar solubility of BaF2 in a solution containing 0.0750 M LiF (Ksp = 1.7 x 10^-6)
(a) 2.3 x 10^-5 M
(b) 3.0 x 10^-4 M
(c) 1.2 x 10^-2 M
(d) 1.3 x 10^-3 M

For each of the following compounds, write a balanced net ionic equation for the dissolution of the compound in water, and write the equilibrium-constant expression for Ksp. 
(d) Hg2Cl2

Calculate the molar solubility of calcium hydroxide in a solution
buffered at each pH.
a. pH = 4

Calculate the molar solubility of calcium hydroxide in a solution 
buffered at each pH. 
c. pH = 9

Calculate the solubility (in grams per 1.00 * 102 mL of solution) 
of magnesium hydroxide in a solution buffered at 
pH = 10. How does this compare to the solubility of Mg(OH)2 
in pure water?

Determine the minimum concentration of the precipitating 
agent on the right to cause precipitation of the cation from the 
solution on the left. 
c. 0.0018 M AgNO3; RbCl

Calculate the molar solubility of calcium hydroxide in a solution
buffered at each pH.
b. pH = 7

Use Le Châtelier’s principle to predict whether the solubility 
of BaF2 will increase, decrease, or remain the same on 
addition of each of the following substances. 
(b) KF

Calculate the solubility of Mg1OH22 in 0.50 M NH4Cl.

Ksp Determination

Calculate the molar solubility of Ni1OH22 when buffered at 
pH (b) 10.0.

Calculate the minimum pH needed to precipitate Mn1OH22 
so completely that the concentration of Mn2 +1aq2 is less 
than 1 mg per liter [1 part per billion (ppb)].

Fluoridation of drinking water is employed in many places 
to aid in the prevention of tooth decay. Typically. the Fion 
concentration is adjusted to about 1 ppm. Some water 
supplies are also “hard”; that is, they contain certain cations 
such as Ca2 + that interfere with the action of soap. 
Consider a case where the concentration of Ca2 + is 8 ppm. 
Could a precipitate of CaF2 form under these conditions? 
(Make any necessary approximations.)

How To Solve Ksp (Solubility & Precipitation) Problems

Selective Precipitation

Calculate the molar solubility of barium fluoride in each liquid 
or solution. 
c. 0.15 M NaF

Calculate the molar solubility of Ni1OH22 when buffered at 
pH (c) 12.0.

What is the solubility-constant expression for Ca3(PO4)2 (s)? 
(a) [Ca2+]3 [PO43-]2
(b) [Ca2+]3 [PO43-]2 / [Ca3(PO4)2]
(c) [Ca(PO4)2] / [Ca2+]3 [PO43-]2
(d) [Ca2+]2 [PO43-]3

Calculate the molar solubility of Ni1OH22 when buffered at 
pH (a) 8.0.

Calculate the molar solubility of barium fluoride in each liquid 
or solution. 
b. 0.10 M Ba(NO3)2

Molar Solubility of Ions

A solution of Na2SO4 is added dropwise to a solution that
is 0.010 M in Ba2 +1aq2 and 0.010 M in Sr2 +1aq2. (b) Which cation precipitates first?

The general&#160;rule is that the greater the solubility product constant for a compound then&#160;the more soluble the&#160;compound.&#160;

If our compound dissolves in a solution with similar ions then the Common Ion Effect&#160;takes place.&#160;

The Reaction Quotient & Ksp

Q vs Ksp

Calculate the molar solubility of MX (Ksp = 1.27 * 10 - 36) in 
each liquid or solution. 
c. 0.20 M Na2 X

The osmotic pressure of a saturated solution of strontium 
sulfate at 25 C is 21 torr. What is the solubility product of 
this salt at 25 C?

Introduction to solubility equilibria | Equilibrium | AP Chemistry | Khan Academy

Pseudogout, a condition with symptoms similar to those of 
gout (see Problem 126), is caused by the formation of calcium 
diphosphate (Ca2P2O7) crystals within tendons, cartilage, and 
ligaments. Calcium diphosphate will precipitate out of blood 
plasma when diphosphate levels become abnormally high. If 
the calcium concentration in blood plasma is 9.2 mg>dL, and 
Ksp for calcium diphosphate is 8.64 * 10 - 13, what minimum 
concentration of diphosphate results in precipitation?

Use the given molar solubilities in pure water to calculate Ksp for 
each compound. 
b. PbF2; molar solubility = 5.63 * 10- 3 M

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12. Molecular Shapes & Valence Bond Theory

13. Liquids, Solids & Intermolecular Forces

14. Solutions

15. Chemical Kinetics

16. Chemical Equilibrium

17. Acid and Base Equilibrium

19. Chemical Thermodynamics

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21. Nuclear Chemistry

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24. Transition Metals and Coordination Compounds

[BOOK CHEM] Clutch Version

Weak Acid Strong Base Titrations

Whenever a WEAK ACID reacts with a STRONG BASE we use an ICF Chart to determine the pH of the solution.&#160;

Strong Acid Strong Base Titrations

In a Strong Acid-Strong Base Titration we do not use an ICF Chart or ICE Chart to determine the pH of the solution.&nbsp;

Titrations of Diprotic and Polyprotic Acids

A diprotic or polyprotic buffer can be approached in a way similar to monoprotic buffers. The key difference is that multiple pKa values will be involved.&#160;

Buffer

Weak Base Strong Acid Titrations

Whenever a&nbsp;WEAK BASE reacts with a&nbsp;STRONG ACID we use an&nbsp;ICF Chart to determine the pH of the solution.&nbsp;