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Precipitation and Dissolution |
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The preservation of medical laboratory blood samples, mining of sea water for magnesium, formulation of over-the-counter medicines such as Milk of Magnesia and antacids, and treating the presence of hard water in your home’s water supply are just a few of the many tasks that involve controlling the equilibrium between a slightly soluble ionic solid and an aqueous solution of its ions.
In some cases, we want to prevent dissolution from occurring. Tooth decay, for example, occurs when the calcium hydroxylapatite, which has the formula Ca5(PO4)3(OH), in our teeth dissolves. The dissolution process is aided when bacteria in our mouths feast on the sugars in our diets to produce lactic acid, which reacts with the hydroxide ions in the calcium hydroxylapatite. Preventing the dissolution prevents the decay. On the other hand, sometimes we want a substance to dissolve. We want the calcium carbonate in a chewable antacid to dissolve because the CO32−
ions produced in this process help soothe an upset stomach.
In this section, we will find out how we can control the dissolution of a slightly soluble ionic solid by the application of Le Châtelier’s principle. We will also learn how to use the equilibrium constant of the reaction to determine the concentration of ions present in a saturated solution.
Silver chloride is what’s known as a sparingly soluble ionic solid ([link]). Recall from the solubility rules in an earlier chapter that halides of Ag+ are not normally soluble. However, when we add an excess of solid AgCl to water, it dissolves to a small extent and produces a mixture consisting of a very dilute solution of Ag+ and Cl– ions in equilibrium with undissolved silver chloride:
This equilibrium, like other equilibria, is dynamic; some of the solid AgCl continues to dissolve, but at the same time, Ag+ and Cl– ions in the solution combine to produce an equal amount of the solid. At equilibrium, the opposing processes have equal rates.
{: #CNX_Chem_15_01_AgCl}
The equilibrium constant for the equilibrium between a slightly soluble ionic solid and a solution of its ions is called the solubility product (Ksp){: data-type="term"} of the solid. Recall from the chapter on solutions and colloids that we use an ion’s concentration as an approximation of its activity in a dilute solution. For silver chloride, at equilibrium:
Note that the Ksp expression does not contain a term in the denominator for the concentration of the reactant, AgCl. According to the guidelines for deriving mass-action expressions described in an earlier chapter on equilibrium, only gases and solutes are represented. Solids and liquids are assigned term values of one and are thus do not appear in equilibrium constant expressions.
Some common solubility products are listed in [link] according to their Ksp values, whereas a more extensive compilation of solubility products appears in Appendix J{: .target-chapter}. Each of these equilibrium constants is much smaller than 1 because the compounds listed are only slightly soluble. A small Ksp represents a system in which the equilibrium lies to the left, so that relatively few hydrated ions would be present in a saturated solution.
| Common Solubility Products by Decreasing Equilibrium Constants | |
|---|---|
| Substance | Ksp at 25 °C |
| CuCl | 1.2 × 10–6 |
| CuBr | 6.27 × 10–9 |
| AgI | 1.5 × 10–16 |
| PbS | 7 × 10–29 |
| Al(OH)3 | 2 × 10–32 |
| Fe(OH)3 | 4 × 10–38 |
(a) AgI, silver iodide, a solid with antiseptic properties
(b) CaCO3, calcium carbonate, the active ingredient in many over-the-counter chewable antacids
(c) Mg(OH)2, magnesium hydroxide, the active ingredient in Milk of Magnesia
(d) Mg(NH4)PO4, magnesium ammonium phosphate, an essentially insoluble substance used in tests for magnesium
(e) Ca5(PO4)3OH, the mineral apatite, a source of phosphate for fertilizers
(Hint: When determining how to break (d) and (e) up into ions, refer to the list of polyatomic ions in the section on chemical nomenclature.)
Solution(a) AgI(s)⇌Ag+(aq)+I−(aq)Ksp=[Ag+][I−]
(b) CaCO3(s)⇌Ca2+(aq)+CO32−(aq)Ksp=[Ca2+][CO32−]
(c) Mg(OH)2(s)⇌Mg2+(aq)+2OH−(aq)Ksp=[Mg2+][OH−]2
(d) Mg(NH4)PO4(s)⇌Mg2+(aq)+NH4+(aq)+PO43−(aq)Ksp=[Mg2+][NH4+][PO43−]
(e) Ca5(PO4)3OH(s)⇌5Ca2+(aq)+3PO43−(aq)+OH−(aq)Ksp=[Ca2+]5[PO43−]3[OH−]
Check Your LearningWrite the ionic equation for the dissolution and the solubility product for each of the following slightly soluble compounds:
(a) BaSO4
(b) Ag2SO4
(c) Al(OH)3
(d) Pb(OH)Cl
(b) Ag2SO4(s)⇌2Ag+(aq)+SO42−(aq)Ksp=[Ag+]2[SO42−];
(c) Al(OH)3(s)⇌Al2+(aq)+3OH−(aq)Ksp=[Al3+][OH−]3;
(d) Pb(OH)Cl(s)⇌Pb2+(aq)+OH−(aq)+Cl−(aq)Ksp=[Pb2+][OH−][Cl−]
Now we will extend the discussion of Ksp and show how the solubility product is determined from the solubility of its ions, as well as how Ksp can be used to determine the molar solubility of a substance.
The Ksp of a slightly soluble ionic compound may be simply related to its measured solubility provided the dissolution process involves only dissociation and solvation, for example:
For cases such as these, one may derive Ksp values from provided solubilities, or vice-versa. Calculations of this sort are most conveniently performed using a compound’s molar solubility, measured as moles of dissolved solute per liter of saturated solution.
10–4 M; therefore, that of F– is 4.30 ×
10–4 M, that is, twice the concentration of Ca2+. What is the solubility product of fluorite?
SolutionFirst, write out the Ksp expression, then substitute in concentrations and solve for Ksp:
Check Your LearningIn a saturated solution that is in contact with solid Mg(OH)2, the concentration of Mg2+ is 1.31 ×
10–4 M. What is the solubility product for Mg(OH)2?
10–12
10–9. Calculate the molar solubility of copper bromide.
SolutionThe solubility product of copper(I) bromide is 6.3 ×
10–9.
The reaction is:
10–5 M.
Check Your Learning The Ksp of AgI is 1.5 ×
10–16. Calculate the molar solubility of silver iodide.
10–8 M
10–6. Calculate the molar solubility of calcium hydroxide.
SolutionThe solubility product of calcium hydroxide is 1.3 ×
10–6.
The reaction is:
10–3 M.
Check Your Learning The Ksp of PbI2 is 1.4 ×
10–8. Calculate the molar solubility of lead(II) iodide.
10–3 M
Note that solubility is not always given as a molar value. When the solubility of a compound is given in some unit other than moles per liter, we must convert the solubility into moles per liter (i.e., molarity) in order to use it in the solubility product expression. [link] shows how to perform those unit conversions before determining the solubility product equilibrium.
10–6 g/L. Determine the solubility product for PbCrO4.
{: #CNX_Chem_15_01_OilPaints}
Solution We are given the solubility of PbCrO4 in grams per liter. If we convert this solubility into moles per liter, we can find the equilibrium concentrations of Pb2+ and CrO42−,
then Ksp:
-
Use the molar mass of PbCrO4 (323.2g1mol)
to convert the solubility of PbCrO4 in grams per liter into moles per liter:
[ PbCrO4 ]=4.6×10−6g PbCrO41L×1mol PbCrO4323.2g PbCrO4=1.4×10−8mol PbCrO41L=1.4×10−8M -
The chemical equation for the dissolution indicates that 1 mol of PbCrO4 gives 1 mol of Pb2+(aq) and 1 mol of CrO42−(aq):
PbCrO4(s)⇌Pb2+(aq)+CrO42−(aq)Thus, both [Pb2+] and [CrO42−]
are equal to the molar solubility of PbCrO4:
[Pb2+]=[CrO42−]=1.4×10−8M -
Solve. Ksp = [Pb2+][CrO42−]
= (1.4 ×
10–8)(1.4 ×
10–8) = 2.0 ×
10–16 {: data-number-style="arabic" .stepwise}
Check Your LearningThe solubility of TlCl [thallium(I) chloride], an intermediate formed when thallium is being isolated from ores, is 3.46 grams per liter at 20 °C. What is its solubility product?
10–4
and chloride ions, Cl–. Although most mercury compounds are now known to be poisonous, eighteenth-century physicians used calomel as a medication. Their patients rarely suffered any mercury poisoning from the treatments because calomel is quite insoluble:
SolutionThe molar solubility of Hg2Cl2 is equal to the concentration of Hg22+
ions because for each 1 mol of Hg2Cl2 that dissolves, 1 mol of Hg22+
forms:
-
Determine x and equilibrium concentrations. Concentrations and changes are given in the following ICE table:
Note that the change in the concentration of Cl– (2x) is twice as large as the change in the concentration of Hg22+
(x) because 2 mol of Cl– forms for each 1 mol of Hg22+
that forms. Hg2Cl2 is a pure solid, so it does not appear in the calculation.
-
Solve for x and the equilibrium concentrations. We substitute the equilibrium concentrations into the expression for Ksp and calculate the value of x:
Ksp=[Hg22+][Cl−]21.1×10−18=(x)(2x)24x3=1.1×10−18x=(1.1×10−184)3=6.5×10−7M[Hg22+]=6.5×10−7M=6.5×10−7M[Cl−]=2x=2(6.5×10−7)=1.3×10−6MThe molar solubility of Hg2Cl2 is equal to [Hg22+],
or 6.5 ×
10–7 M.
-
Check the work. At equilibrium, Q = Ksp:
Q=[Hg22+][Cl−]2=(6.5×10−7)(1.3×10−6)2=1.1×10−18The calculations check. {: data-number-style="arabic" .stepwise}
Check Your LearningDetermine the molar solubility of MgF2 from its solubility product: Ksp = 6.4 ×
10–9.
10–3 M
Tabulated Ksp values can also be compared to reaction quotients calculated from experimental data to tell whether a solid will precipitate in a reaction under specific conditions: Q equals Ksp at equilibrium; if Q is less than Ksp, the solid will dissolve until Q equals Ksp; if Q is greater than Ksp, precipitation will occur at a given temperature until Q equals Ksp.
10–10, very little of it dissolves as it coats the lining of the patient’s intestinal tract. Barium-coated areas of the digestive tract then appear on an X-ray as white, allowing for greater visual detail than a traditional X-ray ([link]).
{: #CNX_Chem_15_01_BariumXray}
Further diagnostic testing can be done using barium sulfate and fluoroscopy. In fluoroscopy, a continuous X-ray is passed through the body so the doctor can monitor, on a TV or computer screen, the barium sulfate’s movement as it passes through the digestive tract. Medical imaging using barium sulfate can be used to diagnose acid reflux disease, Crohn’s disease, and ulcers in addition to other conditions.
Visit this website for more information on how barium is used in medical diagnoses and which conditions it is used to diagnose.
The equation that describes the equilibrium between solid calcium carbonate and its solvated ions is:
We can establish this equilibrium either by adding solid calcium carbonate to water or by mixing a solution that contains calcium ions with a solution that contains carbonate ions. If we add calcium carbonate to water, the solid will dissolve until the concentrations are such that the value of the reaction quotient (Q=[Ca2+][CO32−])
is equal to the solubility product (Ksp = 8.7 ×
10–9). If we mix a solution of calcium nitrate, which contains Ca2+ ions, with a solution of sodium carbonate, which contains CO32−
ions, the slightly soluble ionic solid CaCO3 will precipitate, provided that the concentrations of Ca2+ and CO32−
ions are such that Q is greater than Ksp for the mixture. The reaction shifts to the left and the concentrations of the ions are reduced by formation of the solid until the value of Q equals Ksp. A saturated solution in equilibrium with the undissolved solid will result. If the concentrations are such that Q is less than Ksp, then the solution is not saturated and no precipitate will form.
We can compare numerical values of Q with Ksp to predict whether precipitation will occur, as [link] shows. (Note: Since all forms of equilibrium constants are temperature dependent, we will assume a room temperature environment going forward in this chapter unless a different temperature value is explicitly specified.)
Solution This problem asks whether the reaction:
10–8 is larger than Ksp = 8.9 ×
10–12), we can expect the reaction to shift to the left and form solid magnesium hydroxide. Mg(OH)2(s) forms until the concentrations of magnesium ion and hydroxide ion are reduced sufficiently so that the value of Q is equal to Ksp.
Check Your Learning Use the solubility product in Appendix J{: .target-chapter} to determine whether CaHPO4 will precipitate from a solution with [Ca2+] = 0.0001 M and [HPO42−]
= 0.001 M.
10–7, which is less than Ksp
10–4-M solution of AgNO3 and a 2.0 ×
10–4-M solution of NaCl are mixed?
(Note: The solution also contains Na+ and NO3−
ions, but when referring to solubility rules, one can see that sodium nitrate is very soluble and cannot form a precipitate.)
Solution The equation for the equilibrium between solid silver chloride, silver ion, and chloride ion is:
10–10 (see Appendix J{: .target-chapter}).
AgCl will precipitate if the reaction quotient calculated from the concentrations in the mixture of AgNO3 and NaCl is greater than Ksp. The volume doubles when we mix equal volumes of AgNO3 and NaCl solutions, so each concentration is reduced to half its initial value. Consequently, immediately upon mixing, [Ag+] and [Cl–] are both equal to:
Check Your Learning Will KClO4 precipitate when 20 mL of a 0.050-M solution of K+ is added to 80 mL of a 0.50-M solution of ClO4−?
(Remember to calculate the new concentration of each ion after mixing the solutions before plugging into the reaction quotient expression.)
10–3, which is less than Ksp = 1.05 ×
10–2
In the previous two examples, we have seen that Mg(OH)2 or AgCl precipitate when Q is greater than Ksp. In general, when a solution of a soluble salt of the Mm+ ion is mixed with a solution of a soluble salt of the Xn– ion, the solid, MpXq precipitates if the value of Q for the mixture of Mm+ and Xn– is greater than Ksp for MpXq. Thus, if we know the concentration of one of the ions of a slightly soluble ionic solid and the value for the solubility product of the solid, then we can calculate the concentration that the other ion must exceed for precipitation to begin. To simplify the calculation, we will assume that precipitation begins when the reaction quotient becomes equal to the solubility product.
for this purpose ([link]). At sufficiently high concentrations, the calcium and oxalate ions form solid, CaC2O4·H2O (which also contains water bound in the solid). The concentration of Ca2+ in a sample of blood serum is 2.2 ×
10–3 M. What concentration of C2O42−
ion must be established before CaC2O4·H2O begins to precipitate?
{: #CNX_Chem_15_01_Blood}
SolutionThe equilibrium expression is:
CaC2O4 does not appear in this expression because it is a solid. Water does not appear because it is the solvent.
Solid CaC2O4 does not begin to form until Q equals Ksp. Because we know Ksp and [Ca2+], we can solve for the concentration of C2O42−
that is necessary to produce the first trace of solid:
= 8.9 ×
10–6 M is necessary to initiate the precipitation of CaC2O4 under these conditions.
Check Your Learning If a solution contains 0.0020 mol of CrO42−
per liter, what concentration of Ag+ ion must be reached by adding solid AgNO3 before Ag2CrO4 begins to precipitate? Neglect any increase in volume upon adding the solid silver nitrate.
10–9 M
It is sometimes useful to know the concentration of an ion that remains in solution after precipitation. We can use the solubility product for this calculation too: If we know the value of Ksp and the concentration of one ion in solution, we can calculate the concentration of the second ion remaining in solution. The calculation is of the same type as that in [link]—calculation of the concentration of a species in an equilibrium mixture from the concentrations of the other species and the equilibrium constant. However, the concentrations are different; we are calculating concentrations after precipitation is complete, rather than at the start of precipitation.
10–6 M) may be stained by the manganese upon oxidation, but the amount of Mn2+ in the water can be reduced by adding a base. If a person doing laundry wishes to add a buffer to keep the pH high enough to precipitate the manganese as the hydroxide, Mn(OH)2, what pH is required to keep [Mn2+] equal to 1.8 ×
10–6 M?
Solution The dissolution of Mn(OH)2 is described by the equation:
10–6 M. From that, we calculate the pH. At equilibrium:
10–6 M; at that concentration or less, the ion will not stain clothing.
Check Your Learning The first step in the preparation of magnesium metal is the precipitation of Mg(OH)2 from sea water by the addition of Ca(OH)2. The concentration of Mg2+(aq) in sea water is 5.37 ×
10–2 M. Calculate the pH at which [Mg2+] is diminished to 1.0 ×
10–5 M by the addition of Ca(OH)2.
Due to their light sensitivity, mixtures of silver halides are used in fiber optics for medical lasers, in photochromic eyeglass lenses (glass lenses that automatically darken when exposed to sunlight), and—before the advent of digital photography—in photographic film. Even though AgCl (Ksp = 1.6 ×
10–10), AgBr (Ksp = 5.0 ×
10–13), and AgI (Ksp = 1.5 ×
10–16) are each quite insoluble, we cannot prepare a homogeneous solid mixture of them by adding Ag+ to a solution of Cl–, Br–, and I–; essentially all of the AgI will precipitate before any of the other solid halides form because of its smaller value for Ksp. However, we can prepare a homogeneous mixture of the solids by slowly adding a solution of Cl–, Br–, and I– to a solution of Ag+.
When two anions form slightly soluble compounds with the same cation, or when two cations form slightly soluble compounds with the same anion, the less soluble compound (usually, the compound with the smaller Ksp) generally precipitates first when we add a precipitating agent to a solution containing both anions (or both cations). When the Ksp values of the two compounds differ by two orders of magnitude or more (e.g., 10–2 vs. 10–4), almost all of the less soluble compound precipitates before any of the more soluble one does. This is an example of selective precipitation{: data-type="term"}, where a reagent is added to a solution of dissolved ions causing one of the ions to precipitate out before the rest.
are often present in the water discharged from manufacturing facilities. An abundance of phosphate causes excess algae to grow, which impacts the amount of oxygen available for marine life as well as making water unsuitable for human consumption.
{: #CNX_Chem_15_01_Wastewater}
One common way to remove phosphates from water is by the addition of calcium hydroxide, known as lime, Ca(OH)2. The lime is converted into calcium carbonate, a strong base, in the water. As the water is made more basic, the calcium ions react with phosphate ions to produce hydroxylapatite, Ca5(PO4)3(OH), which then precipitates out of the solution:
View this site for more information on how phosphorus is removed from wastewater.
Selective precipitation can also be used in qualitative analysis. In this method, reagents are added to an unknown chemical mixture in order to induce precipitation. Certain reagents cause specific ions to precipitate out; therefore, the addition of the reagent can be used to determine whether the ion is present in the solution.
Solution The two equilibria involved are:
For AgI: AgI precipitates when Q equals Ksp for AgI (1.5 ×
10–16). When [I–] = 0.0010 M:
10–13 M.
For AgCl: AgCl precipitates when Q equals Ksp for AgCl (1.6 ×
10–10). When [Cl–] = 0.10 M:
10–9 M.
AgI begins to precipitate at a lower [Ag+] than AgCl, so AgI begins to precipitate first.
Check Your Learning If silver nitrate solution is added to a solution which is 0.050 M in both Cl– and Br– ions, at what [Ag+] would precipitation begin, and what would be the formula of the precipitate?
10–11 M; AgBr precipitates first
As we saw when we discussed buffer solutions, the hydronium ion concentration of an aqueous solution of acetic acid decreases when the strong electrolyte sodium acetate, NaCH3CO2, is added. We can explain this effect using Le Châtelier’s principle. The addition of acetate ions causes the equilibrium to shift to the left, decreasing the concentration of H3O+
to compensate for the increased acetate ion concentration. This increases the concentration of CH3CO2H:
Because sodium acetate and acetic acid have the acetate ion in common, the influence on the equilibrium is called the common ion effect{: data-type="term"}.
The common ion effect can also have a direct effect on solubility equilibria. Suppose we are looking at the reaction where silver iodide is dissolved:
If we were to add potassium iodide (KI) to this solution, we would be adding a substance that shares a common ion with silver iodide. Le Châtelier’s principle tells us that when a change is made to a system at equilibrium, the reaction will shift to counteract that change. In this example, there would be an excess of iodide ions, so the reaction would shift toward the left, causing more silver iodide to precipitate out of solution.
10–28.
Solution The first thing you should notice is that the cadmium sulfide is dissolved in a solution that contains cadmium ions. We need to use an ICE table to set up this problem and include the CdBr2 concentration as a contributor of cadmium ions:
10–26 M.
Check Your Learning Calculate the molar solubility of aluminum hydroxide, Al(OH)3, in a 0.015-M solution of aluminum nitrate, Al(NO3)3. The Ksp of Al(OH)3 is 2 ×
10–32.
10–11 M
The equilibrium constant for an equilibrium involving the precipitation or dissolution of a slightly soluble ionic solid is called the solubility product, Ksp, of the solid. When we have a heterogeneous equilibrium involving the slightly soluble solid MpXq and its ions Mm+ and Xn–:
We write the solubility product expression as:
The solubility product of a slightly soluble electrolyte can be calculated from its solubility; conversely, its solubility can be calculated from its Ksp, provided the only significant reaction that occurs when the solid dissolves is the formation of its ions.
A slightly soluble electrolyte begins to precipitate when the magnitude of the reaction quotient for the dissolution reaction exceeds the magnitude of the solubility product. Precipitation continues until the reaction quotient equals the solubility product.
A reagent can be added to a solution of ions to allow one ion to selectively precipitate out of solution. The common ion effect can also play a role in precipitation reactions. In the presence of an ion in common with one of the ions in the solution, Le Châtelier’s principle applies and more precipitate comes out of solution so that the molar solubility is reduced.
- MpXq(s)⇌pMm+(aq)+qXn−(aq)Ksp=[Mm+]p[Xn−]q {: data-bullet-style="bullet"}
(a) AgI(s)⟶Ag+(aq)+I−(aq)x_____
(b) CaCO3(s)⟶Ca2+(aq)+CO32−(aq)____x
(c) Mg(OH)2(s)⟶Mg2+(aq)+2OH−(aq)x_____
(d) Mg3(PO4)2(s)⟶3Mg2+(aq)+2PO43−(aq)x_____
(e) Ca5(PO4)3OH(s)⟶5Ca2+(aq)+3PO43−(aq)+OH−(aq)__________x
{: data-type="newline"}
(b) CaCO3(s)⇌Ca2+(aq)+CO32−(aq)x_x
{: data-type="newline"}
(c) Mg(OH)2(s)⇌Mg2+(aq)+2OH−(aq)x2x_
{: data-type="newline"}
(d) Mg3(PO4)2(s)⇌3Mg2+(aq)+2PO43−(aq)3x_2x
{: data-type="newline"}
(e) Ca5(PO4)3OH(s)⇌5Ca2+(aq)+3PO43−(aq)+OH−(aq)5x_3x_x
(a) BaSO4(s)⟶Ba2+(aq)+SO42−(aq)x_____
(b) Ag2SO4(s)⟶2Ag+(aq)+SO42−(aq)_____x
(c) Al(OH)3(s)⟶Al3+(aq)+3OH−(aq)x_____
(d) Pb(OH)Cl(s)⟶Pb2+(aq)+OH−(aq)+Cl−(aq)_____x_____
(e) Ca3(AsO4)2(s)⟶3Ca2+(aq)+2AsO43−(aq)3x_____
in a saturated solution above 1.0 g of solid Ag2CrO4 change when 100 g of solid Ag2CrO4 is added to the system? Explain.
(a) PbCl2
(b) Ag2S
(c) Sr3(PO4)2
(d) SrSO4
(a) LaF3
(b) CaCO3
(c) Ag2SO4
(d) Pb(OH)2
{: data-type="newline"}
(b) CaCO3(s)⇌Ca2+(aq)+CO32−(aq)Ksp=[Ca2+][CO32−];
{: data-type="newline"}
(c) Ag2SO4(s)⇌2Ag+(aq)+SO42−(aq)Ksp=[Ag+]2[SO42−];
{: data-type="newline"}
(d) Pb(OH)2(s)⇌Pb2+(aq)+2OH−(aq)Ksp=[Pb2+][OH−]2
(a) BaSiF6, 0.026 g/100 mL (contains SiF62−
ions)
(b) Ce(IO3)4, 1.5 ×
10–2 g/100 mL
(c) Gd2(SO4)3, 3.98 g/100 mL
(d) (NH4)2PtBr6, 0.59 g/100 mL (contains PtBr62−
ions)
(a) BaSeO4, 0.0118 g/100 mL
(b) Ba(BrO3)2·H2O, 0.30 g/100 mL
(c) NH4MgAsO4·6H2O, 0.038 g/100 mL
(d) La2(MoO4)3, 0.00179 g/100 mL
10–7; (b) 1.6 ×
10–6; (c) 2.2 ×
10–9; (d) 7.91 ×
10–22
(a) KHC4H4O6
(b) PbI2
(c) Ag4[Fe(CN)6], a salt containing the Fe(CN)4−
ion
(d) Hg2I2
10–2 M; (b) 1.5 ×
10–3 M; (c) 2.27 ×
10–9 M; (d) 2.2 ×
10–10 M
(a) Ag2SO4
(b) PbBr2
(c) AgI
(d) CaC2O4·H2O
(a) AgCl(s) in 0.025 M NaCl
(b) CaF2(s) in 0.00133 M KF
(c) Ag2SO4(s) in 0.500 L of a solution containing 19.50 g of K2SO4
(d) Zn(OH)2(s) in a solution buffered at a pH of 11.45
10−9 M = [Ag+], [Cl−] = 0.025 M* * * {: data-type="newline"}
Check: 6.4×10−9M0.025M×100%=2.6×10−5%,
an insignificant change;* * * {: data-type="newline"}
(b) 2.2 ×
10−5 M = [Ca2+], [F−] = 0.0013 M* * * {: data-type="newline"}
Check: 2.26×10−5M0.00133M×100%=1.70%.
This value is less than 5% and can be ignored.* * * {: data-type="newline"}
(c) 0.2238 M = [SO42−];
[Ag+] = 7.4 ×
10–3 M* * * {: data-type="newline"}
Check: 3.7×10−30.2238×100%=1.64×10−2;
the condition is satisfied.* * * {: data-type="newline"}
(d) [OH–] = 2.8 ×
10–3 M; 5.7 ×
10−12 M = [Zn2+]* * * {: data-type="newline"}
Check: 5.7×10−122.8×10−3×100%=2.0×10−7%;
x is less than 5% of [OH–] and is, therefore, negligible.
(a) TlCl(s) in 1.250 M HCl
(b) PbI2(s) in 0.0355 M CaI2
(c) Ag2CrO4(s) in 0.225 L of a solution containing 0.856 g of K2CrO4
(d) Cd(OH)2(s) in a solution buffered at a pH of 10.995
(a) TlCl(s) in 0.025 M TlNO3
(b) BaF2(s) in 0.0313 M KF
(c) MgC2O4 in 2.250 L of a solution containing 8.156 g of Mg(NO3)2
(d) Ca(OH)2(s) in an unbuffered solution initially with a pH of 12.700
10−3 M* * * {: data-type="newline"}
Check: 7.6×10−30.025×100%=30%
{: data-type="newline"}
This value is too large to drop x. Therefore solve by using the quadratic equation:* * * {: data-type="newline"}
[Ti+] = 3.1 ×
10–2 M* * * {: data-type="newline"}
[Cl–] = 6.1 ×
10–3* * * {: data-type="newline"}
(b) [Ba2+] = 7.7 ×
10–4 M* * * {: data-type="newline"}
Check: 7.7×10−40.0313×100%=2.4%
{: data-type="newline"}
Therefore, the condition is satisfied.* * * {: data-type="newline"}
[Ba2+] = 7.7 ×
10–4 M* * * {: data-type="newline"}
[F–] = 0.0321 M;* * * {: data-type="newline"}
(c) Mg(NO3)2 = 0.02444 M* * * {: data-type="newline"}
[C2O42−]=2.9×10−5
{: data-type="newline"}
Check: 2.9×10−50.02444×100%=0.12%
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The condition is satisfied; the above value is less than 5%.* * * {: data-type="newline"}
[C2O42−]=2.9×10−5M
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[Mg2+] = 0.0244 M* * * {: data-type="newline"}
(d) [OH–] = 0.0501 M* * * {: data-type="newline"}
[Ca2+] = 3.15 ×
10–3* * * {: data-type="newline"}
Check: 3.15×10−30.050×100%=6.28%
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This value is greater than 5%, so a more exact method, such as successive approximations, must be used.* * * {: data-type="newline"}
[Ca2+] = 2.8 ×
10–3 M* * * {: data-type="newline"}
[OH–] = 0.053 ×
10–2 M
10–3 M) of SO42−
because of its cathartic action (it is a laxative). Does natural water that is saturated with CaSO4 (“gyp” water) as a result or passing through soil containing gypsum, CaSO4·2H2O, meet these standards? What is the concentration of SO42−
in such water?
10–3 M = [SO42−]
= [Ca2+]; Since this concentration is higher than 2.60 ×
10–3 M, “gyp” water does not meet the standards.
(a) Calculate [Ag+] in a saturated aqueous solution of AgBr.
(b) What will [Ag+] be when enough KBr has been added to make [Br–] = 0.050 M?
(c) What will [Br–] be when enough AgNO3 has been added to make [Ag+] = 0.020 M?
10–5. What mass of this salt will dissolve in 1.0 L of 0.010 M SO42−?
(a) TlCl
(b) BaF2
(c) Ag2CrO4
(d) CaC2O4·H2O
(e) the mineral anglesite, PbSO4
(a) AgI
(b) Ag2SO4
(c) Mn(OH)2
(d) Sr(OH)2·8H2O
(e) the mineral brucite, Mg(OH)2
10–5 M; (b) [Ag+] = 2.88 ×
10–2 M, [SO42−]
= 1.44 ×
10–2 M; (c) [Mn2+] = 3.7 ×
10–5 M, [OH–] = 7.4 ×
10–5 M; (d) [Sr2+] = 4.3 ×
10–2 M, [OH–] = 8.6 ×
10–2 M; (e) [Mg2+] = 1.3 ×
10–4 M, [OH–] = 2.6 ×
10–4 M.
(a) AgBr: [Ag+] = 5.7 ×
10–7 M, [Br–] = 5.7 ×
10–7 M
(b) CaCO3: [Ca2+] = 5.3 ×
10–3 M, [CO32−]
= 9.0 ×
10–7 M
(c) PbF2: [Pb2+] = 2.1 ×
10–3 M, [F–] = 4.2 ×
10–3 M
(d) Ag2CrO4: [Ag+] = 5.3 ×
10–5 M, 3.2 ×
10–3 M
(e) InF3: [In3+] = 2.3 ×
10–3 M, [F–] = 7.0 ×
10–3 M
(a) TlCl: [Tl+] = 1.21 ×
10–2 M, [Cl–] = 1.2 ×
10–2 M
(b) Ce(IO3)4: [Ce4+] = 1.8 ×
10–4 M, [IO3−]
= 2.6 ×
10–13 M
(c) Gd2(SO4)3: [Gd3+] = 0.132 M, [SO42−]
= 0.198 M
(d) Ag2SO4: [Ag+] = 2.40 ×
10–2 M, [SO42−]
= 2.05 ×
10–2 M
(e) BaSO4: [Ba2+] = 0.500 M, [SO42−]
= 2.16 ×
10–10 M
10–4; (b) 8.2 ×
10–55; (c) 1.35 ×
10–4; (d) 1.18 ×
10–5; (e) 1.08 ×
10–10
(a) KClO4: [K+] = 0.01 M, [ClO4−]
= 0.01 M
(b) K2PtCl6: [K+] = 0.01 M, [PtCl62−]
= 0.01 M
(c) PbI2: [Pb2+] = 0.003 M, [I–] = 1.3 ×
10–3 M
(d) Ag2S: [Ag+] = 1 ×
10–10 M, [S2–] = 1 ×
10–13 M
(a) CaCO3: [Ca2+] = 0.003 M, [CO32−]
= 0.003 M
(b) Co(OH)2: [Co2+] = 0.01 M, [OH–] = 1 ×
10–7 M
(c) CaHPO4: [Ca2+] = 0.01 M, [HPO42−]
= 2 ×
10–6 M
(d) Pb3(PO4)2: [Pb2+] = 0.01 M, [PO43−]
= 1 ×
10–13 M
(b) The compound does not precipitate.* * * {: data-type="newline"}
(c) The compound does not precipitate.* * * {: data-type="newline"}
(d) The compound precipitates.
10−7 M
when Ag3PO4 starts to precipitate from a solution that is 0.0125 M in Ag+.
10−13 M
10–6 M in CO32−.
10–3 M
to 8.2 ×
10–4 M by precipitation of Ag2CO3?
10–4 M by precipitation of CaF2?
10–4
10–4-M Ba(NO3)2 solution is added to 0.200 L of 5 ×
10–4 M Li2SO4. Does BaSO4 precipitate? Explain your answer.
10–7).
(b) If the NiCO3 were a contaminant in a sample of CoCO3 (Ksp = 1.0 ×
10–12), what mass of CoCO3 would have been lost? Keep in mind that both NiCO3 and CoCO3 dissolve in the same solution.
10–7 g
10–6 M in water used for laundry purposes can cause staining. What [OH–] is required to reduce [Fe2+] to this level by precipitation of Fe(OH)2?
(a) Hg22+
and Cu2+
(b) SO42−
and Cl–
(c) Hg2+ and Co2+
(d) Zn2+ and Sr2+
(e) Ba2+ and Mg2+
(f) CO32−
and OH–
and Cu2+: Add SO42−.
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(b) SO42−
and Cl–: Add Ba2+.* * * {: data-type="newline"}
(c) Hg2+ and Co2+: Add S2–.* * * {: data-type="newline"}
(d) Zn2+ and Sr2+: Add OH– until [OH–] = 0.050 M.* * * {: data-type="newline"}
(e) Ba2+ and Mg2+: Add SO42−.
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(f) CO32−
and OH–: Add Ba2+.
10–5 mol of KBr and 0.10 mol of KCl per liter. AgNO3 is gradually added to this solution. Which forms first, solid AgBr or solid AgCl?
10–2 mol of KI and 0.10 mol of KCl per liter. AgNO3 is gradually added to this solution. Which forms first, solid AgI or solid AgCl?
10−12 M
-
[HPO42−]
-
[H2PO4−]
-
[H3PO4]) is 0.020 M. What is the minimum concentration of Ca2+ necessary to induce kidney stone formation? (See [link] for additional information.)
Mg2+(aq)+Ca(OH)2(aq)⟶Mg(OH)2(s)+Ca2+(aq)
Mg(OH)2(s)+2HCl(aq)⟶MgCl2(s)+2H2O(l)
MgCl2(l)→electrolysisMg(s)+Cl2(g)
Sea water has a density of 1.026 g/cm3 and contains 1272 parts per million of magnesium as Mg2+(aq) by mass. What mass, in kilograms, of Ca(OH)2 is required to precipitate 99.9% of the magnesium in 1.00 ×
103 L of sea water?
changes as metal sulfides precipitate.)
(a) The iodate ion concentration of a saturated solution of La(IO3)3 was found to be 3.1 ×
10–3 mol/L. Find the Ksp.
(b) Find the concentration of iodate ions in a saturated solution of Cu(IO3)2 (Ksp = 7.4 ×
10–8).
10–11; (b) [Cu2+] = 2.6 ×
10–3; [IO3−]
= 5.3 ×
10–3
10–13).
10–15)?
10–5 g Pb(OH)2
10–12. Include the ionic reaction and the expression for Ksp in your answer. (Kw = 1 ×
10–14 = [H3O+][OH–])
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1.23 ×
10−3 g Mg(OH)2
10–5, what is the Ksp value for LQ?
(a) MgCO3Ksp=3.5×10−8
(b) CaCO3Ksp=4.2×10−7
(c) SrCO3Ksp=3.9×10−9
(d) BaCO3Ksp=4.4×10−5
(e) MnCO3Ksp=5.1×10−9
10–16.
common ion effect : effect on equilibrium when a substance with an ion in common with the dissolved species is added to the solution; causes a decrease in the solubility of an ionic species, or a decrease in the ionization of a weak acid or base ^
molar solubility : solubility of a compound expressed in units of moles per liter (mol/L) ^
selective precipitation : process in which ions are separated using differences in their solubility with a given precipitating reagent ^
solubility product (Ksp) : equilibrium constant for the dissolution of a slightly soluble electrolyte