15.1 Precipitation and Dissolution

Complete the changes in concentrations for each of the following reactions:

(a)
$\begin{matrix}\text{AgI} \left(s\right) \rightarrow & \text{Ag}^{+} \left(a q\right) & + \text{I}^{−} \left(a q\right) \\ \\ & x & _____\end{matrix}$

(b)
$\begin{matrix}\text{CaCO}_{3} \left(s\right) \rightarrow & \text{Ca}^{\text{2}+} \left(a q\right) + & \text{CO}_{3} ^{\text{2}−} \left(a q\right) \\ \\ & ____ & x\end{matrix}$

(c)
$\begin{matrix}\text{Mg} \left(\text{OH}\right)_{2} \left(s\right) \rightarrow & \text{Mg}^{\text{2}+} \left(a q\right) + & 2 \text{OH}^{−} \left(a q\right) \\ \\ & x & _____\end{matrix}$

(d)
$\begin{matrix}\text{Mg}_{3} \left(\right. \text{PO}_{4} \left.\right)_{2} \left(s\right) \rightarrow & 3 \text{Mg}^{\text{2}+} \left(a q\right) + & 2 \text{PO}_{4} ^{\text{3}−} \left(a q\right) \\ \\ & x & _____\end{matrix}$

(e)
$\begin{matrix}\text{Ca}_{5} \left(\right. \text{PO}_{4} \left.\right)_{3} \text{OH} \left(s\right) \rightarrow & 5 \text{Ca}^{\text{2}+} \left(a q\right) + & 3 \text{PO}_{4} ^{\text{3}−} \left(a q\right) + & \text{OH}^{−} \left(a q\right) \\ \\ & _____ & _____ & x\end{matrix}$

How do the concentrations of Ag⁺ and $\text{CrO}_{4} ^{\text{2}−}$ in a saturated solution above 1.0 g of solid Ag₂CrO₄ change when 100 g of solid Ag₂CrO₄ is added to the system? Explain.

What additional information do we need to answer the following question: How is the equilibrium of solid silver bromide with a saturated solution of its ions affected when the temperature is raised?

Which of the following slightly soluble compounds has a solubility greater than that calculated from its solubility product because of hydrolysis of the anion present: AgCl, BaSO₄, CaF₂, Hg₂I₂, MnCO₃, and ZnS?

Write the ionic equation for the dissolution and the K_sp expression for each of the following slightly soluble ionic compounds:

(a) LaF₃

(b) CaCO₃

(c) Ag₂SO₄

(d) Pb(OH)₂

The Handbook of Chemistry and Physics gives solubilities of the following compounds in grams per 100 mL of water. Because these compounds are only slightly soluble, assume that the volume does not change on dissolution and calculate the solubility product for each.

(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

Assuming that no equilibria other than dissolution are involved, calculate the molar solubility of each of the following from its solubility product:

(a) KHC₄H₄O₆

(b) PbI₂

(c) Ag₄[Fe(CN)₆], a salt containing the $\text{Fe}(\text{CN})_{6} ^{\text{4}–}$ ion

(d) Hg₂I₂

Assuming that no equilibria other than dissolution are involved, calculate the concentration of all solute species in each of the following solutions of salts in contact with a solution containing a common ion. Show that changes in the initial concentrations of the common ions can be neglected.

(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 containin…

Assuming that no equilibria other than dissolution are involved, calculate the concentration of all solute species in each of the following solutions of salts in contact with a solution containing a common ion. State whether or not it is appropriate to neglect the changes in the initial concentrations of the common ions.

(a) TlCl(s) in 0.025 M TlNO₃

(b) BaF₂(s) in 0.313 M KF

(c) MgC₂O₄ in 2.250 L…

Explain why the changes in concentrations of the common ions in the exercise cannot be neglected.

Refer to Appendix J for solubility products for calcium salts. Determine which of the calcium salts listed is most soluble in moles per liter and which is most soluble in grams per liter.

Public Health Service standards for drinking water set a maximum of 250 mg/L (2.60 $\times$ 10^–3 M) of $\text{SO}_{4} ^{\text{2}−}$ 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 $\text{SO}_{4} ^{\text{2}−}$ in such water?

The solubility product of CaSO₄·2H₂O is 2.4 $\times$ 10^–5. What mass of this salt will dissolve in 1.0 L of 0.010 M $\text{SO}_{4} ^{\text{2}−} ?$

Assuming that no equilibria other than dissolution are involved, calculate the concentrations of ions in a saturated solution of each of the following (see Appendix J for solubility products):

(a) AgI

(b) Ag₂SO₄

(c) Mn(OH)₂

(d) Sr(OH)₂·8H₂O

(e) the mineral brucite, Mg(OH)₂

The following concentrations are found in mixtures of ions in equilibrium with slightly soluble solids. From the concentrations given, calculate K_sp for each of the slightly soluble solids indicated:

(a) TlCl: [Tl⁺] = 1.21 $\times$ 10^–2 M, [Cl^–] = 1.2 $\times$ 10^–2 M

(b) Ce(IO₃)₄: [Ce⁴⁺] = 1.8 $\times$ 10^–4 M, $\left[\text{IO}_{3} ^{−}\right]$ = 2.6 $\times$ 10^–13 M

(c) Gd₂(SO₄)₃: [Gd³⁺] = 0.132 M, $\left[\text{SO}_{\text{4}} ^{\text{2}−}\right]$ = 0.198 M

(d) Ag₂SO₄: [Ag⁺] = 2.40 $\times$ 10^–2 M, $\left[\text{SO}_{\text{4}} ^{\text{2}−}\right]$ = 2.05 $\times$…

Which of the following compounds precipitates from a solution that has the concentrations indicated? (See Appendix J for K_sp values.)

(a) CaCO₃: [Ca²⁺] = 0.003 M, $\left[\text{CO}_{\text{3}} ^{\text{2}−}\right]$ = 0.003 M

(b) Co(OH)₂: [Co²⁺] = 0.01 M, [OH^–] = 1 $\times$ 10^–7 M

(c) CaHPO₄: [Ca²⁺] = 0.01 M, $\left[\text{HPO}_{\text{4}} ^{\text{2}−}\right]$ = 2 $\times$ 10^–6 M

(d) Pb₃(PO₄)₂: [Pb²⁺] = 0.01 M, $\left[\text{PO}_{\text{4}} ^{\text{3}−}\right]$= 1 $\times$ 10^–13 M

Calculate the concentration of sulfate ion when BaSO₄ just begins to precipitate from a solution that is 0.0758 M in Ba²⁺.

Calculate the concentration of $\text{PO}_{4} ^{\text{3}−}$ when Ag₃PO₄ starts to precipitate from a solution that is 0.0125 M in Ag⁺.

Calculate the concentration of Ag⁺ required to begin precipitation of Ag₂CO₃ in a solution that is 2.50 $\times$ 10^–6 M in $\text{CO}_{3} ^{\text{2}−} .$

What [F^–] is required to reduce [Ca²⁺] to 1.0 $\times$ 10^–4 M by precipitation of CaF₂?

Perform these calculations for nickel(II) carbonate. (a) With what volume of water must a precipitate containing NiCO₃ be washed to dissolve 0.100 g of this compound? Assume that the wash water becomes saturated with NiCO₃ (K_sp = 1.36 $\times$ 10^–7).

(b) If the NiCO₃ were a contaminant in a sample of CoCO₃ (K_sp = 1.0 $\times$ 10^–12), what mass of CoCO₃ would have been lost? Keep in mind that both NiCO₃ and CoCO_…

A solution is 0.010 M in both Cu²⁺ and Cd²⁺. What percentage of Cd²⁺ remains in the solution when 99.9% of the Cu²⁺ has been precipitated as CuS by adding sulfide?

What reagent might be used to separate the ions in each of the following mixtures, which are 0.1 M with respect to each ion? In some cases it may be necessary to control the pH. (Hint: Consider the K_sp values given in Appendix J.)

(a) $\text{Hg}_{2} ^{\text{2}+}$ and Cu²⁺

(b) $\text{SO}_{4} ^{\text{2}−}$ and Cl^–

(c) Hg²⁺ and Co²⁺

(d) Zn²⁺ and Sr²⁺

(e) Ba²⁺ and Mg²⁺

(f) $\text{CO}_{3} ^{\text{2}−}$ and OH^–

A solution contains 1.0 $\times$ 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?

About 50% of urinary calculi (kidney stones) consist of calcium phosphate, Ca₃(PO₄)₂. The normal mid range calcium content excreted in the urine is 0.10 g of Ca²⁺ per day. The normal mid range amount of urine passed may be taken as 1.4 L per day. What is the maximum concentration of phosphate ion that urine can contain before a calculus begins to form?

Magnesium metal (a component of alloys used in aircraft and a reducing agent used in the production of uranium, titanium, and other active metals) is isolated from sea water by the following sequence of reactions:

$\text{Mg}^{\text{2}+} \left(a q\right) + \text{Ca}(\text{OH})_{2} \left(a q\right) \rightarrow \text{Mg}(\text{OH})_{2} \left(s\right) + \text{Ca}^{\text{2}+} \left(a q\right)$

$\text{Mg}(\text{OH})_{2} \left(s\right) + \text{2HCl}( a q \left.\right) \rightarrow \text{MgCl}_{2} \left(s\right) + \text{2H}_{2} \text{O} \left(l\right)$

$\text{MgCl}_{2} \left(l\right) \overset{ \text{electrolysis} }{\rightarrow} \text{Mg} \left(s\right) + \text{Cl}_{2} \left(g\right)$

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…

Perform the following calculations involving concentrations of iodate ions:

(a) The iodate ion concentration of a saturated solution of La(IO₃)₃ was found to be 3.1 $\times$ 10^–3 mol/L. Find the K_sp.

(b) Find the concentration of iodate ions in a saturated solution of Cu(IO₃)₂ (K_sp = 7.4 $\times$ 10^–8).

How many grams of Pb(OH)₂ will dissolve in 500 mL of a 0.050-M PbCl₂ solution (K_sp = 1.2 $\times$ 10^–15)?

How many grams of Milk of Magnesia, Mg(OH)₂ (s) (58.3 g/mol), would be soluble in 200 mL of water. K_sp = 7.1 $\times$ 10^–12. Include the ionic reaction and the expression for K_sp in your answer. (K_w = 1 $\times$ 10^–14 = [H₃O⁺][OH^–])

The carbonate ion concentration is gradually increased in a solution containing equal concentrations of the divalent cations of magnesium, calcium, strontium, barium, and manganese. Which of the following carbonates will precipitate first? Which will precipitate last? Explain.

(a) $\text{MgCO}_{3} \cdot 3 \text{H}_{2} \text{O} K_{\text{sp}} = 1 \times 1 0^{- 5}$

(b) $\text{CaCO}_{3} K_{\text{sp}} = 8.7 \times 1 0^{- 9}$

(c) $\text{SrCO}_{3} K_{\text{sp}} = 7 \times 1 0^{- 10}$

(d) $\text{BaCO}_{3} K_{\text{sp}} = 1.6 \times 1 0^{- 9}$

(e) $\text{MnCO}_{3} K_{\text{sp}} = 8.8 \times 1 0^{- 11}$