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Chapter 6: Problem 52

Give an example how each of the following insoluble ionic compounds could be produced using a precipitation reaction. Write the balanced formula equation for each reaction. a. \(\mathrm{Fe}(\mathrm{OH})_{3}(s)\) c. \(\mathrm{PbSO}_{4}(s)\) b. \(\mathrm{Hg}_{2} \mathrm{Cl}_{2}(s)\) d. \(\mathrm{BaCrO}_{4}(s)\)

Short Answer

Expert verified
a. \(\mathrm{Fe(NO_3)_3 (aq) + 3NaOH (aq) \rightarrow Fe(OH)_3 (s) + 3NaNO_3(aq)}\) b. \(\mathrm{Hg}_2(NO_3)_2(aq) + 2NaCl(aq) \rightarrow Hg_2Cl_2(s) + 2NaNO_3(aq)\) c. \(\mathrm{Pb(NO_3)_2 (aq) + Na_2SO_4 (aq) \rightarrow PbSO_4 (s) + 2NaNO_3(aq)}\) d. \(\mathrm{Ba(NO_3)_2 (aq) + K_2CrO_4 (aq) \rightarrow BaCrO_4 (s) + 2KNO_3(aq)}\)

Step by step solution

01

a. Precipitation Reaction for Fe(OH)₃

To produce \(\mathrm{Fe}(\mathrm{OH})_{3}(s)\) as a precipitate, we need two soluble reactants that form this compound. One option is choosing a soluble iron compound (such as \(\mathrm{Fe}^{3+}\) salt) and a source of the hydroxide ion, \(\mathrm{OH}^{-}\) (like \(\mathrm{NaOH}\)). The balanced formula equation for this reaction is: \[\mathrm{Fe(NO_3)_3 (aq) + 3NaOH (aq) \rightarrow Fe(OH)_3 (s) + 3NaNO_3(aq)}\]
02

b. Precipitation Reaction for Hg₂Cl₂

To form \(\mathrm{Hg}_{2}\mathrm{Cl}_{2}(s)\) as the precipitate, we need a soluble mercury(I) compound and a soluble source of chloride ions. For instance, we can use \(\mathrm{Hg}_{2}(\mathrm{NO}_{3})_{2}\) as the mercury(I) compound and \(\mathrm{NaCl}\) as the chloride source. The balanced formula equation for this reaction is: \[\mathrm{Hg}_2(NO_3)_2(aq) + 2NaCl(aq) \rightarrow Hg_2Cl_2(s) + 2NaNO_3(aq)\]
03

c. Precipitation Reaction for PbSO₄

In order to form \(\mathrm{PbSO}_4(s)\) as a precipitate, we can choose a soluble lead(II) compound such as \(\mathrm{Pb(NO_3)_2}\) and a soluble sulfate source like \(\mathrm{Na_2SO_4}\). The balanced formula equation for this reaction is: \[\mathrm{Pb(NO_3)_2 (aq) + Na_2SO_4 (aq) \rightarrow PbSO_4 (s) + 2NaNO_3(aq)}\]
04

d. Precipitation Reaction for BaCrO₄

To produce \(\mathrm{BaCrO}_4(s)\) precipitate, we must select a soluble barium compound, for example, \(\mathrm{Ba(NO_3)_2}\), and a soluble source of chromate ions like \(\mathrm{K_2CrO_4}\). The balanced formula equation for this reaction is: \[\mathrm{Ba(NO_3)_2 (aq) + K_2CrO_4 (aq) \rightarrow BaCrO_4 (s) + 2KNO_3(aq)}\]

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Most popular questions from this chapter

Citric acid, which can he ohtained from lemon juice, has the molecular formula \(\mathrm{C}_{6} \mathrm{H}_{8} \mathrm{O}_{7}\). A 0.250 -g sample of citric acid dissolved in \(25.0 \mathrm{mL}\) of water requires \(37.2 \mathrm{mL}\) of \(0.105 \mathrm{M}\) NaOH for complete neutralization. What number of acidic hydrogens per molecule does citric acid have?

Describe how you would prepare 2.00 L of each of the following solutions. a. \(0.250 \mathrm{M}\) NaOH from solid \(\mathrm{NaOH}\) b. \(0.250 M\) NaOH from \(1.00 M\) NaOH stock solution c. \(0.100 M K_{2} C r O_{4}\) from solid \(K_{2} C r O_{4}\) d. \(0.100 M K_{2} C r O_{4}\) from \(1.75 M K_{2} C r O_{4}\) stock solution

Zinc and magnesium metal each react with hydrochloric acid according to the following equations: $$ \begin{array}{c} \mathrm{Zn}(s)+2 \mathrm{HCl}(a q) \longrightarrow \mathrm{ZnCl}_{2}(a q)+\mathrm{H}_{2}(g) \\ \mathrm{Mg}(s)+2 \mathrm{HCl}(a q) \longrightarrow \mathrm{MgCl}_{2}(a q)+\mathrm{H}_{2}(g) \end{array} $$ A \(10.00-\mathrm{g}\) mixture of zinc and magnesium is reacted with the stoichiometric amount of hydrochloric acid. The reaction mixture is then reacted with \(156 \mathrm{mL}\) of \(3.00 \mathrm{M}\) silver nitrate to produce the maximum possible amount of silver chloride. a. Determine the percent magnesium by mass in the original mixture. b. If \(78.0 \mathrm{mL}\) of HCl was added, what was the concentration of the HCl?

Specify which of the following equations represent oxidationreduction reactions, and indicate the oxidizing agent, the reducing agent, the species being oxidized, and the species being reduced. a. \(\mathrm{CH}_{4}(g)+\mathrm{H}_{2} \mathrm{O}(g) \rightarrow \mathrm{CO}(g)+3 \mathrm{H}_{2}(g)\) b. \(2 \mathrm{AgNO}_{3}(a q)+\mathrm{Cu}(s) \rightarrow \mathrm{Cu}\left(\mathrm{NO}_{3}\right)_{2}(a q)+2 \mathrm{Ag}(s)\) c. \(\mathrm{Zn}(s)+2 \mathrm{HCl}(a q) \rightarrow \mathrm{ZnCl}_{2}(a q)+\mathrm{H}_{2}(g)\) d. \(2 \mathrm{H}^{+}(a q)+2 \mathrm{CrO}_{4}^{2-}(a q) \rightarrow \mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}(a q)+\mathrm{H}_{2} \mathrm{O}(i)\)

A stream flows at a rate of \(5.00 \times 10^{4}\) liters per second (L/s) upstream of a manufacturing plant. The plant discharges \(3.50 \times 10^{3} \mathrm{L} / \mathrm{s}\) of water that contains \(65.0 \mathrm{ppm}\) HCl into the stream. (See Exercise 123 for definitions.) a. Calculate the stream's total flow rate downstream from this plant. b. Calculate the concentration of HCl in ppm downstream from this plant. c. Further downstream, another manufacturing plant diverts \(1.80 \times 10^{4} \mathrm{L} / \mathrm{s}\) of water from the stream for its own use. This plant must first neutralize the acid and does so by adding lime: $$ \mathrm{CaO}(s)+2 \mathrm{H}^{+}(a q) \longrightarrow \mathrm{Ca}^{2+}(a q)+\mathrm{H}_{2} \mathrm{O}(i) $$ What mass of \(\mathrm{CaO}\) is consumed in an 8.00 -h work day by this plant? d. The original stream water contained \(10.2 \mathrm{ppm} \mathrm{Ca}^{2+} .\) Although no calcium was in the waste water from the first plant, the waste water of the second plant contains \(\mathrm{Ca}^{2+}\) from the neutralization process. If \(90.0 \%\) of the water used by the second plant is returned to the stream, calculate the concentration of \(\mathrm{Ca}^{2+}\) in ppm downstream of the second plant.
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