Chapter 16

Write the equilibrium equation and the \(K_{s p}\) expression for each of the following. (a) \(\mathrm{Co}_{2} \mathrm{~S}_{3}\) (b) \(\mathrm{PbCl}_{2}\) (c) \(\mathrm{Zn}_{2} \mathrm{P}_{2} \mathrm{O}_{7}\) (d) \(\mathrm{Sc}(\mathrm{OH})_{3}\)

Write the equilibrium equation and the \(K_{s p}\) expression for each of the following. (a) \(\mathrm{AgCl}\) (b) \(\mathrm{Al}_{2}\left(\mathrm{CO}_{3}\right)_{3}\) (c) \(\mathrm{MnS}_{2}\) (d) \(\mathrm{Mg}(\mathrm{OH})_{2}\)

Write the equilibrium equations on which the following \(K_{s p}\) expressions are based. (a) \(\left[\mathrm{Hg}_{2}{\underline{\phantom{xx}}}^{2+}\right]\left[\mathrm{Cl}^{-}\right]^{2}\) (b) \(\left[\mathrm{Pb}^{2+}\right]\left[\mathrm{CrO}_{4}^{2-}\right]\) (c) \(\left[\mathrm{Mn}^{4+}\right]\left[\mathrm{O}^{2-}\right]^{2}\) (d) \(\left[\mathrm{Al}^{3+}\right]^{2}\left[\mathrm{~S}^{2-}\right]^{3}\)

Write the equilibrium equations on which the following \(K_{s p}\) expressions are based. (a) \(\left[\mathrm{Ca}^{2+}\right]\left[\mathrm{CO}_{3}{\underline{\phantom{xx}}}^{2-}\right]\) (b) \(\left[\mathrm{Co}^{3+}\right]\left[\mathrm{OH}^{-}\right]^{3}\) (c) \(\left[\mathrm{Ag}^{+}\right]^{2}\left[\mathrm{~S}^{2-}\right]\) (d) \(\left[\mathrm{Pb}^{2+}\right]\left[\mathrm{Cl}^{-}\right]^{2}\)

Given \(K_{s p}\) and the equilibrium concentration of one ion, calculate the equilibrium concentration of the other ion. (a) cadmium(II) hydroxide: \(K_{\text {sp }}=2.5 \times 10^{-14} ;\left[\mathrm{Cd}^{2+}\right]=1.5 \times 10^{-6} M\) (b) copper(II) arsenate \(\left(\mathrm{Cu}_{3}\left(\mathrm{AsO}_{4}\right)_{2}\right): K_{\mathrm{sp}}=7.6 \times 10^{-36} ;\left[\mathrm{AsO}_{4}^{3-}\right]=\) \(2.4 \times 10^{-4} M\) (c) zinc oxalate: \(K_{s p}=2.7 \times 10^{-8} ;\left[\mathrm{C}_{2} \mathrm{O}_{4}{\underline{\phantom{xx}}}^{2-}\right]=8.8 \times 10^{-3} M\)

Fill in the blanks in the following table. (a) \(\mathrm{CoCO}_{3}\) (b) \(\mathrm{LaF}_{3}\) (c) \(\mathrm{Ba}_{3}\left(\mathrm{PO}_{4}\right)_{2}\)

Fill in the blanks in the following table. (a) \(\mathrm{BaC}_{2} \mathrm{O}_{4}\) (b) \(\mathrm{Cr}(\mathrm{OH})_{3}\) (c) \(\mathrm{Pb}_{3}\left(\mathrm{PO}_{4}\right)_{2}\)

Cadmium(II) chloride is added to a solution of potassium hydroxide with a \(\mathrm{pH}\) of \(9.62 .\left(K_{\mathrm{sp}} \mathrm{Cd}(\mathrm{OH})_{2}=2.5 \times 10^{-14}\right)\) (a) At what concentration of \(\mathrm{Cd}^{2+}\) does a precipitate first start to form? (b) Enough cadmium(II) chloride is added to make [Cd \(\left.^{2+}\right]=\) \(0.0013 M\). What is the \(\mathrm{pH}\) of the resulting solution? (c) What percentage of the original hydroxide ion is left in solution?

Water from a well is found to contain \(3.0 \mathrm{mg}\) of calcium ion per liter. If \(0.50 \mathrm{mg}\) of sodium sulfate is added to one liter of the well water without changing its volume, will a precipitate form? What should \(\left[\mathrm{SO}_{4}{\underline{\phantom{xx}}}^{2-}\right]\) be to just start precipitation?

Before lead in paint was discontinued, lead chromate was a common pigment in yellow paint. A \(1.0\) -L solution is prepared by mixing \(0.50 \mathrm{mg}\) of lead nitrate with \(0.020 \mathrm{mg}\) of potassium chromate. Will a precipitate form? What should \(\left[\mathrm{Pb}^{2+}\right]\) be to just start precipitation?

A solution is prepared by mixing \(13.00 \mathrm{~mL}\) of \(0.0021 M\) aqueous \(\mathrm{Hg}_{2}\left(\mathrm{NO}_{3}\right)_{2}\) with \(25.0 \mathrm{~mL}\) of \(0.015 \mathrm{M} \mathrm{HCl}\). Assume that volumes are additive. (a) Will precipitation occur? (b) Calculate \(\left[\mathrm{Hg}_{2}{\underline{\phantom{xx}}}^{2+}\right],\left[\mathrm{Cl}^{-}\right]\), and \(\left[\mathrm{NO}_{3}^{-}\right]\) after equilibrium is established.

A solution is prepared by mixing \(35.00 \mathrm{~mL}\) of a \(0.061 \mathrm{M}\) solution of zinc nitrate with \(20.0 \mathrm{~mL}\) of \(\mathrm{KOH}\) with a pH of \(9.00\). Assume that volumes are additive. (a) Will precipitation occur? (b) Calculate \(\left[\mathrm{Zn}^{2+}\right],\left[\mathrm{NO}_{3}^{-}\right],\left[\mathrm{K}^{+}\right]\), and the \(\mathrm{pH}\) after equilibrium is established.

Calculate the \(K_{s p}\) of the following compounds, given their molar solubilities. (a) \(\mathrm{AgCN}, 7.73 \times 10^{-9} \mathrm{M}\) (b) \(\mathrm{Ni}(\mathrm{OH})_{2}, 5.16 \times 10^{-6} \mathrm{M}\) (c) \(\mathrm{Cu}_{3}\left(\mathrm{PO}_{4}\right)_{2}, 1.67 \times 10^{-8} M\)

Calculate the solubility (in grams per liter) of silver chloride in the following. (a) pure water (b) \(0.025 \mathrm{M} \mathrm{BaCl}_{2}\) (c) \(0.17 \mathrm{M} \mathrm{AgNO}_{3}\)

Calculate the solubility (in grams per liter) of magnesium hydroxide in the following. (a) pure water (b) \(0.041 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\) (c) \(0.0050 \mathrm{MgCl}_{2}\)

Lead azide, \(\mathrm{Pb}\left(\mathrm{N}_{3}\right)_{2}\), is used as a detonator in car airbags. The impact of a collision causes \(\mathrm{Pb}\left(\mathrm{N}_{3}\right)_{2}\) to be converted into an enormous amount of gas that fills the airbag. At \(25^{\circ} \mathrm{C}\), a saturated solution of lead azide is prepared by dissolving \(25 \mathrm{mg}\) in water to make \(100.0 \mathrm{~mL}\) of solution. What is \(K_{\mathrm{sp}}\) for lead azide?

One gram of \(\mathrm{PbCl}_{2}\) is dissolved in \(1.0 \mathrm{~L}\) of hot water. When the solution is cooled to \(25^{\circ} \mathrm{C}\), will some of the \(\mathrm{PbCl}_{2}\) crystallize out? If so, how much?

\(\mathrm{K}_{\mathrm{sp}}\) for silver acetate \(\left(\mathrm{AgC}_{2} \mathrm{H}_{3} \mathrm{O}_{2}\right)\) at \(80^{\circ} \mathrm{C}\) is estimated to be \(2 \times 10^{-2}\). Ten grams of silver acetate are added to \(1.0 \mathrm{~L}\) of water at \(25^{\circ} \mathrm{C}\). (a) Will all the silver acetate dissolve at \(25^{\circ} \mathrm{C} ?\) (b) If the solution (assume the volume to be \(1.0 \mathrm{~L}\) ) is heated to \(80^{\circ} \mathrm{C}\), will all the silver acetate dissolve?

Calculate the solubility (g/100 mL) of iron(II) hydroxide in buffered solutions with the following pH's. (a) 4 (b) 7 (c) 10

At \(25^{\circ} \mathrm{C}, 100.0 \mathrm{~mL}\) of a \(\mathrm{Ba}(\mathrm{OH})_{2}\) solution is prepared by dissolving \(\mathrm{Ba}(\mathrm{OH})_{2}\) in an alkaline solution. At equilibrium, the solution has \(2.37 \mathrm{~g}\) of \(\mathrm{Ba}(\mathrm{OH})_{2}\) and a \(\mathrm{pH}\) of \(13.28\). Estimate \(K_{s p}\) for \(\mathrm{Ba}(\mathrm{OH})_{2}\).

At \(25^{\circ} \mathrm{C}, 100.0 \mathrm{~mL}\) of a \(\mathrm{Cr}(\mathrm{OH})_{2}\) solution is prepared by dissolving \(\mathrm{Cr}(\mathrm{OH})_{2}\) in an alkaline solution. At equilibrium, the solution has \(8.65\) \(\mathrm{mg}\) of \(\mathrm{Cr}(\mathrm{OH})_{2}\) and a \(\mathrm{pH}\) of \(8.50\). Estimate \(K_{s p}\) for \(\mathrm{Cr}(\mathrm{OH})_{2}\).

A solution is \(0.035 M\) in \(\mathrm{Na}_{2} \mathrm{SO}_{4}\) and \(0.035 \mathrm{M}\) in \(\mathrm{Na}_{2} \mathrm{CrO}_{4}\). Solid \(\mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\) is added without changing the volume of the solution. (a) Which salt, \(\mathrm{PbSO}_{4}\) or \(\mathrm{PbCrO}_{4}\), will precipitate first? (b) What is \(\left[\mathrm{Pb}^{2+}\right]\) when the salt in (a) first begins to precipitate?

Solid lead nitrate is added to a solution that is \(0.020 \mathrm{M}\) in \(\mathrm{OH}^{-}\) and \(\mathrm{SO}_{4}^{2-}\). Addition of the lead nitrate does not change the volume of the solution. (a) Which compound, \(\mathrm{PbSO}_{4}\) or \(\mathrm{Pb}(\mathrm{OH})_{2}\left(K_{s p}=2.8 \times 10^{-16}\right)\), will precipitate first? (b) What is the \(\mathrm{pH}\) of the solution when \(\mathrm{PbSO}_{4}\) first starts to precipitate?

A 65-mL solution of \(0.40 \mathrm{M} \mathrm{Al}\left(\mathrm{NO}_{3}\right)_{3}\) is mixed with \(125 \mathrm{~mL}\) of \(0.17 \mathrm{M}\) iron(II) nitrate. Solid sodium hydroxide is then added without a change in volume. (a) Which will precipitate first, \(\mathrm{Al}(\mathrm{OH})_{3}\) or \(\mathrm{Fe}(\mathrm{OH})_{2} ?\) (b) What is \(\left[\mathrm{OH}^{-}\right]\) when the first compound begins to precipitate?

A solution is made up by adding \(0.925 \mathrm{~g}\) of silver(I) nitrate and \(6.25 \mathrm{~g}\) of magnesium nitrate to enough water to make \(375 \mathrm{~mL}\) of solution. Solid sodium carbonate is added without changing the volume of the solution. (a) Which salt, \(\mathrm{Ag}_{2} \mathrm{CO}_{3}\) or \(\mathrm{MgCO}_{3}\), will precipitate first? (b) What is the concentration of the carbonate ion when the first salt starts to precipitate?

A solution is made up by mixing \(125 \mathrm{~mL}\) of \(0.100 \mathrm{M} \mathrm{AuNO}_{3}\) and \(225 \mathrm{~mL}\) of \(0.049 \mathrm{M} \mathrm{AgNO}_{3}\). Twenty-five mL of a \(0.0100 \mathrm{M}\) solution of \(\mathrm{HCl}\) is then added. \(K_{\text {sp }}\) of \(\mathrm{AuCl}=2.0 \times 10^{-13} .\) When equilibrium is established, will there be- -no precipitate? -a precipitate of \(\mathrm{AuCl}\) only? -a precipitate of \(\mathrm{AgCl}\) only? -a precipitate of both \(\mathrm{AgCl}\) and \(\mathrm{AuCl}\) ?

Write net ionic equations for the reaction of \(\mathrm{H}^{+}\) with (a) \(\mathrm{Cu}_{2} \mathrm{~S}\) (b) \(\mathrm{Hg}_{2} \mathrm{Cl}_{2}\) (c) \(\mathrm{SrCO}_{3}\) (d) \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}{\underline{\phantom{xx}}}^{2+}\) (e) \(\mathrm{Ca}(\mathrm{OH})_{2}\)

Write net ionic equations for the reactions of each of the following with strong acid. (a) \(\mathrm{CaCO}_{3}\) (b) NiS (c) \(\mathrm{Al}(\mathrm{OH})_{3}\) (d) \(\mathrm{Sb}(\mathrm{OH})_{4}^{-}\) (e) \(\mathrm{AgCl}\)

Write a net ionic equation for the reaction with ammonia by which (a) silver chloride dissolves. (b) aluminum ion forms a precipitate. (c) copper(II) forms a complex ion.

Write a net ionic equation for the reaction with ammonia by which (a) \(\mathrm{Cu}(\mathrm{OH})_{2}\) dissolves. (b) \(\mathrm{Cd}^{2+}\) forms a complex ion. (c) \(\mathrm{Pb}^{2+}\) forms a precipitate.

Write a net ionic equation for the reaction with \(\mathrm{OH}^{-}\) by which (a) \(\mathrm{Sb}^{3+}\) forms a precipitate. (b) antimony(III) hydroxide dissolves when more \(\mathrm{OH}^{-}\) is added. (c) \(\mathrm{Sb}^{3+}\) forms a complex ion.

Write a net ionic equation for the reaction with \(\mathrm{OH}^{-}\) by which (a) \(\mathrm{Ni}^{2+}\) forms a precipitate. (b) \(\mathrm{Sn}^{4+}\) forms a complex ion. (c) \(\mathrm{Al}(\mathrm{OH})_{3}\) dissolves.

Write an overall net ionic equation and calculate \(K\) for the reaction where \(\mathrm{CuCl}\left(K_{s p}=1.9 \times 10^{-7}\right)\) is dissolved by \(\mathrm{NaCN}\) to form \(\left[\mathrm{Cu}(\mathrm{CN})_{2}\right]^{-}\) \(\left(K_{f}=1.0 \times 10^{16}\right)\).

Consider the reaction \(\mathrm{Cu}(\mathrm{OH})_{2}(s)+4 \mathrm{NH}_{3}(a q) \rightleftharpoons \mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}^{2+}(a q)+2 \mathrm{OH}^{-}(a q)\) (a) Calculate \(K\) given that for \(\mathrm{Cu}(\mathrm{OH})_{2} K_{s p}=2 \times 10^{-19}\) and for \(\mathrm{Cu}\left(\mathrm{NH}_{3}\right)_{4}^{2+} K_{\mathrm{f}}=2 \times 10^{12}\) (b) Determine the solubility of \(\mathrm{Cu}(\mathrm{OH})_{2}\) (in \(\mathrm{mol} / \mathrm{L}\) ) in \(4.5 \mathrm{M} \mathrm{NH}_{2}\).

Calculate the molar solubility of gold(I) chloride \(\left(K_{s p}=2.0 \times 10^{-13}\right)\) in \(0.10 \mathrm{M} \mathrm{NaCN}\). The complex ion formed is \(\left[\mathrm{Au}(\mathrm{CN})_{2}\right]^{-}\) with \(K_{\mathrm{f}}=2 \times 10^{38}\). Ignore any other competing equilibrium systems.

For the reaction $$ \mathrm{CdC}_{2} \mathrm{O}_{4}(s)+4 \mathrm{NH}_{3}(a q) \rightleftharpoons \mathrm{Cd}\left(\mathrm{NH}_{3}\right)_{4}^{2+}(a q)+\mathrm{C}_{2} \mathrm{O}_{4}^{2-}(a q) $$ (a) calculate \(K .\left(K_{s p}\right.\) for \(\mathrm{CdC}_{2} \mathrm{O}_{4}\) is \(\left.1.5 \times 10^{-8} .\right)\) (b) calculate \(\left[\mathrm{NH}_{3}\right]\) at equilibrium when \(2.00 \mathrm{~g}\) of \(\mathrm{CdC}_{2} \mathrm{O}_{4}\) is dissolved in \(1.00 \mathrm{~L}\) of solution.

Calcium ions in blood trigger clotting. To prevent that in donated blood, sodium oxalate, \(\mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\), is added to remove calcium ions according to the following equation. $$ \mathrm{C}_{2} \mathrm{O}_{4}^{2-}(a q)+\mathrm{Ca}^{2+}(a q) \longrightarrow \mathrm{CaC}_{2} \mathrm{O}_{4}(s) $$ Blood contains about \(0.10 \mathrm{mg} \mathrm{Ca}^{2+} / \mathrm{mL}\). If a 250.0-mL sample of donated blood is treated with an equal volume of \(0.160 \mathrm{M} \mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\), estimate \(\left[\mathrm{Ca}^{2+}\right]\) after precipitation. \(\left(K_{s p} \mathrm{CaC}_{2} \mathrm{O}_{4}=4 \times 10^{-9}\right)\)

A town adds \(2.0\) ppm of \(F^{-}\) ion to fluoridate its water supply (fluoridation of water reduces the incidence of dental caries). If the concentration of \(\mathrm{Ca}^{2+}\) in the water is \(3.5 \times 10^{-4} M\), will a precipitate of \(\mathrm{CaF}_{2}\) form when the water is fluoridated?

Predict what effect each of the following has on the position of the equilibrium $$ \mathrm{PbCl}_{2}(s) \rightleftharpoons \mathrm{Pb}^{2+}(a q)+2 \mathrm{Cl}^{-}(a q) \quad \Delta H=23.4 \mathrm{~kJ} $$ (a) addition of \(1 \mathrm{M} \mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\) solution (b) increase in temperature (c) addition of \(\mathrm{Ag}^{+}\), forming \(\mathrm{AgCl}\) (d) addition of \(1 M\) hydrochloric acid

When \(25.0 \mathrm{~mL}\) of \(0.500 \mathrm{M}\) iron(II) sulfate is combined with \(35.0 \mathrm{~mL}\) of \(0.332 M\) barium hydroxide, two different precipitates are formed. (a) Write a net ionic equation for the reaction that takes place. (b) Estimate the mass of the precipitates formed. (c) What are the equilibrium concentrations of the ions in solution?

Consider a \(1.50-\mathrm{L}\) aqueous solution of \(3.75 \mathrm{M} \mathrm{NH}_{3}\), where \(17.5 \mathrm{~g}\) of \(\mathrm{NH}_{4} \mathrm{Cl}\) is dissolved. To this solution, \(5.00 \mathrm{~g}\) of \(\mathrm{MgCl}_{2}\) is added. (a) What is \(\left[\mathrm{OH}^{-}\right]\) before \(\mathrm{MgCl}_{2}\) is added? (b) Will a precipitate form? (c) What is \(\left[\mathrm{Mg}^{2+}\right]\) after equilibrium is established?

Consider the following solubility data for calcium oxalate \(\left(\mathrm{Ca} \mathrm{C}_{2} \mathrm{O}_{4}\right)\) : $$ \begin{aligned} &K_{s p} \text { at } 25^{\circ} \mathrm{C}=4 \times 10^{-9} \\ &K_{s p} \text { at } 95^{\circ} \mathrm{C}=1 \times 10^{-8} \end{aligned} $$ Five hundred mL of a saturated solution is prepared at \(95^{\circ} \mathrm{C}\). How many milligrams of \(\mathrm{CaC}_{2} \mathrm{O}_{4}\) will precipitate when the solution is cooled to \(25^{\circ} \mathrm{C}\) ? (Assume that supersaturation does not take place.)

Shown below is a representation of the ionic solid \(\mathrm{MX}\), where \(\mathrm{M}\) cations are represented by squares and \(X\) anions are represented by circles. Fill in the box after the arrow to represent what happens to the solid after it has been completely dissolved in water. For simplicity, do not represent the water molecules.

The box below represents one liter of a saturated solution of the species \(\square\), where squares represent the cation and circles represent the anion. Water molecules, though present, are not shown. Complete the next three figures below by filling one-liter boxes to the right of the arrow, showing the state of the ions after water is added to form saturated solutions. The species represented to the left of the arrow is the solid form of the ions represented above. Do not show the water molecules.

Which of the following statements are true? (a) For an insoluble metallic salt, \(K_{\text {sp }}\) is always less than 1 . (b) More \(\mathrm{PbCl}_{2}\) can be dissolved at \(100^{\circ} \mathrm{C}\) than at \(25^{\circ} \mathrm{C}\). One can conclude that dissolving \(\mathrm{PbCl}_{2}\) is an exothermic process. (c) When strips of copper metal are added to a saturated solution of \(\mathrm{Cu}(\mathrm{OH})_{2}\), a precipitate of \(\mathrm{Cu}(\mathrm{OH})_{2}\) can be expected to form because of the common ion effect.

Consider the insoluble salts \(\mathrm{JQ} \mathrm{K}_{2} \mathrm{R}, \mathrm{L}_{2} \mathrm{~S}_{3}, \mathrm{MT}_{2}\), and \(\mathrm{NU}_{3}\). They are formed from the metal ions \(\mathrm{J}^{+}, \mathrm{K}^{+}, \mathrm{L}^{3+}, \mathrm{M}^{2+}\), and \(\mathrm{N}^{3+}\) and the nonmetal ions \(\mathrm{Q}^{-}, \mathrm{R}^{2-}, \mathrm{S}^{2-}, \mathrm{T}^{-}\), and \(\mathrm{U}^{-}\). All the salts have the same \(K_{\text {sp }}, 1 \times 10^{-10}\), at \(25^{\circ} \mathrm{C}\). (a) Which salt has the highest molar solubility? (b) Does the salt with the highest molar solubility have the highest solubility in g salt/100 g water? (c) Can the solubility of each salt in \(\mathrm{g} / 100 \mathrm{~g}\) water be determined from the information given? If yes, calculate the solubility of each salt in \(\mathrm{g} / 100 \mathrm{~g}\) water. If no, why not?

A plot of the solubility of a certain compound \(\left(\mathrm{g} / 100 \mathrm{~g} \mathrm{H}_{2} \mathrm{O}\right)\) against temperature \(\left({ }^{\circ} \mathrm{C}\right)\) is a straight line with a positive slope. Is dissolving that compound an exothermic process?

What is the solubility of \(\mathrm{CaF}_{2}\) in a buffer solution containing \(0.30 \mathrm{M}\) \(\mathrm{HCHO}_{2}\) and \(0.20 \mathrm{M} \mathrm{NaCHO}_{2}\) ? Hint: Consider the equation $$ \mathrm{CaF}_{2}(s)+2 \mathrm{H}^{+}(a q) \longrightarrow \mathrm{Ca}^{2+}(a q)+2 \mathrm{HF}(a q) $$ and solve the equilibrium problem.

What is the \(\mathrm{I}^{-}\) concentration just as \(\mathrm{AgCl}\) begins to precipitate when \(1.0 \mathrm{M} \mathrm{AgNO}_{3}\) is slowly added to a solution containing \(0.020 \mathrm{M} \mathrm{Cl}^{-}\) and \(0.020 M \mathrm{I}^{-}\)

The concentrations of various cations in seawater, in moles per liter, are $$ \begin{array}{llllll} \hline \text { Ion } & \mathrm{Na}^{+} & \mathrm{Mg}^{2+} & \mathrm{Ca}^{2+} & \mathrm{A}^{3+} & \mathrm{Fe}^{3+} \\ \text { Molarity }(M) & 0.46 & 0.056 & 0.01 & 4 \times 10^{-7} & 2 \times 10^{-7} \\ \hline \end{array} $$ (a) At what \(\left[\mathrm{OH}^{-}\right]\) does \(\mathrm{Mg}(\mathrm{OH})_{2}\) start to precipitate? (b) At this concentration, will any of the other ions precipitate? (c) If enough \(\mathrm{OH}^{-}\) is added to precipitate \(50 \%\) of the \(\mathrm{Mg}^{2+}\), what percentage of each of the other ions will precipitate? (d) Under the conditions in (c), what mass of precipitate will be ob- tained from one liter of seawater?