15.1 Precipitation and dissolution  (Page 13/17)

Calculate the concentration of Sr ²⁺ when SrF ₂ starts to precipitate from a solution that is 0.0025 M in F ^– .

Calculate the concentration of ${\text{PO}}_4{}^{\text{3−}}$ when Ag ₃ PO ₄ starts to precipitate from a solution that is 0.0125 M in Ag ⁺ .

Calculate the concentration of F ^– required to begin precipitation of CaF ₂ in a solution that is 0.010 M in Ca ²⁺ .

Calculate the concentration of Ag ⁺ required to begin precipitation of Ag ₂ CO ₃ in a solution that is 2.50 $\times$ 10 ^–6 M in ${\text{CO}}_3{}^{\text{2−}}.$

What [Ag ⁺ ] is required to reduce $[{\text{CO}}_3{}^{\text{2−}}]$ to 8.2 $\times$ 10 ^–4 M by precipitation of Ag ₂ CO ₃ ?

What [F ^– ] is required to reduce [Ca ²⁺ ] to 1.0 $\times$ 10 ^–4 M by precipitation of CaF ₂ ?

A volume of 0.800 L of a 2 $\times$ 10 ^–4 - M Ba(NO ₃ ) ₂ solution is added to 0.200 L of 5 $\times$ 10 ^–4 M Li ₂ SO ₄ . Does BaSO ₄ precipitate? Explain your answer.

Perform these calculations for nickel(II) carbonate. (a) With what volume of water must a precipitate containing NiCO ₃ be washed to dissolve 0.100 g of this compound? Assume that the wash water becomes saturated with NiCO ₃ ( K _sp = 1.36 $\times$ 10 ^–7 ).

(b) If the NiCO ₃ were a contaminant in a sample of CoCO ₃ ( K _sp = 1.0 $\times$ 10 ^–12 ), what mass of CoCO ₃ would have been lost? Keep in mind that both NiCO ₃ and CoCO ₃ dissolve in the same solution.

Iron concentrations greater than 5.4 $\times$ 10 ^–6 M in water used for laundry purposes can cause staining. What [OH ^– ] is required to reduce [Fe ²⁺ ] to this level by precipitation of Fe(OH) ₂ ?

A solution is 0.010 M in both Cu ²⁺ and Cd ²⁺ . What percentage of Cd ²⁺ remains in the solution when 99.9% of the Cu ²⁺ has been precipitated as CuS by adding sulfide?

A solution is 0.15 M in both Pb ²⁺ and Ag ⁺ . If Cl ^– is added to this solution, what is [Ag ⁺ ] when PbCl ₂ begins to precipitate?

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 Cu ²⁺

(b) ${\text{SO}}_4{}^{\text{2−}}$ and Cl ^–

(c) Hg ²⁺ and Co ²⁺

(d) Zn ²⁺ and Sr ²⁺

(e) Ba ²⁺ and Mg ²⁺

(f) ${\text{CO}}_3{}^{\text{2−}}$ and OH ^–

A solution contains 1.0 $\times$ 10 ^–5 mol of KBr and 0.10 mol of KCl per liter. AgNO ₃ is gradually added to this solution. Which forms first, solid AgBr or solid AgCl?

A solution contains 1.0 $\times$ 10 ^–2 mol of KI and 0.10 mol of KCl per liter. AgNO ₃ is gradually added to this solution. Which forms first, solid AgI or solid AgCl?

The calcium ions in human blood serum are necessary for coagulation ( [link] ). Potassium oxalate, K ₂ C ₂ O ₄ , is used as an anticoagulant when a blood sample is drawn for laboratory tests because it removes the calcium as a precipitate of CaC ₂ O ₄ ·H ₂ O. It is necessary to remove all but 1.0% of the Ca ²⁺ in serum in order to prevent coagulation. If normal blood serum with a buffered pH of 7.40 contains 9.5 mg of Ca ²⁺ per 100 mL of serum, what mass of K ₂ C ₂ O ₄ is required to prevent the coagulation of a 10 mL blood sample that is 55% serum by volume? (All volumes are accurate to two significant figures. Note that the volume of serum in a 10-mL blood sample is 5.5 mL. Assume that the K _sp value for CaC ₂ O ₄ in serum is the same as in water.)

About 50% of urinary calculi (kidney stones) consist of calcium phosphate, Ca ₃ (PO ₄ ) ₂ . The normal mid range calcium content excreted in the urine is 0.10 g of Ca ²⁺ 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?

The pH of normal urine is 6.30, and the total phosphate concentration $({[\text{PO}}_4{}^{\text{3−}}]$ + $[{\text{HPO}}_4{}^{\text{2−}}]$ + $[{\text{H}}_2{\text{PO}}_4{}^{\text{−}}]$ + [H ₃ PO ₄ ]) is 0.020 M . What is the minimum concentration of Ca ²⁺ necessary to induce kidney stone formation? (See [link] for additional information.)

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:

${\text{Mg}}^{\text{2+}}(aq)+{\text{Ca(OH)}}_2(aq)⟶{\text{Mg(OH)}}_2(s)+{\text{Ca}}^{\text{2+}}(aq)$

${\text{Mg(OH)}}_2(s)+\text{2HCl(}aq)⟶{\text{MgCl}}_2(s)+{\text{2H}}_2\text{O}(l)$

${\text{MgCl}}_2(l)\stackrel{\text{electrolysis}}{\to }\text{Mg}(s)+{\text{Cl}}_2(g)$

Sea water has a density of 1.026 g/cm ³ and contains 1272 parts per million of magnesium as Mg ²⁺ ( aq ) by mass. What mass, in kilograms, of Ca(OH) ₂ is required to precipitate 99.9% of the magnesium in 1.00 $\times$ 10 ³ L of sea water?

Hydrogen sulfide is bubbled into a solution that is 0.10 M in both Pb ²⁺ and Fe ²⁺ and 0.30 M in HCl. After the solution has come to equilibrium it is saturated with H ₂ S ([H ₂ S] = 0.10 M ). What concentrations of Pb ²⁺ and Fe ²⁺ remain in the solution? For a saturated solution of H ₂ S we can use the equilibrium:

(Hint: The $[{\text{H}}_3{\text{O}}^{\text{+}}]$ changes as metal sulfides precipitate.)

Perform the following calculations involving concentrations of iodate ions:

(a) The iodate ion concentration of a saturated solution of La(IO ₃ ) 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(IO ₃ ) ₂ ( K _sp = 7.4 $\times$ 10 ^–8 ).

Calculate the molar solubility of AgBr in 0.035 M NaBr ( K _sp = 5 $\times$ 10 ^–13 ).

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

Use the simulation from the earlier Link to Learning to complete the following exercise:. Using 0.01 g CaF ₂ , give the K _sp values found in a 0.2- M solution of each of the salts. Discuss why the values change as you change soluble salts.

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 = [H ₃ O ⁺ ][OH ^– ])

Two hypothetical salts, LM ₂ and LQ, have the same molar solubility in H ₂ O. If K _sp for LM ₂ is 3.20 $\times$ 10 ^–5 , what is the K _sp value for LQ?

Which of the following carbonates will form first? Which of the following will form last? Explain.

(a) ${\text{MgCO}}_3{K}_{\text{sp}}=3.5\times 10^{-8}$

(b) ${\text{CaCO}}_3{K}_{\text{sp}}=4.2\times 10^{-7}$

(c) ${\text{SrCO}}_3{K}_{\text{sp}}=3.9\times 10^{-9}$

(d) ${\text{BaCO}}_3{K}_{\text{sp}}=4.4\times 10^{-5}$

(e) ${\text{MnCO}}_3{K}_{\text{sp}}=5.1\times 10^{-9}$

How many grams of Zn(CN) ₂ ( s ) (117.44 g/mol) would be soluble in 100 mL of H ₂ O? Include the balanced reaction and the expression for K _sp in your answer. The K _sp value for Zn(CN) ₂ ( s ) is 3.0 $\times$ 10 ^–16 .