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Chapter 16: Problem 63

When aqueous KI is added gradually to mercury(II) nitrate, an orange precipitate forms. Continued addition of KI causes the precipitate to dissolve. Write balanced equations to explain these observations. (Hint: \(\mathrm{Hg}^{2+}\) reacts with \(\mathrm{I}^{-}\) to form \(\mathrm{HgI}_{4}{ }^{2-}\).)

Short Answer

Expert verified
The balanced chemical equations representing the formation of the orange precipitate and its dissolution when aqueous KI is added gradually to mercury(II) nitrate are: 1. \(Hg(NO_3)_2(aq) + 2KI(aq) \rightarrow HgI2(s) + 2KNO3(aq)\) 2. \(HgI2(s) + 2I^-(aq) \rightarrow HgI_4^{2-}(aq)\)

Step by step solution

01

Write the chemical equations for the formation of orange precipitate and its dissolution.

For the formation of the orange precipitate, the reaction between mercury(II) nitrate (Hg(NO_3)_2) and potassium iodide (KI) will produce mercury(II) iodide (HgI2) and potassium nitrate (KNO3): \[Hg(NO_3)_2(aq) + 2KI(aq) \rightarrow HgI2(s) + 2KNO3(aq)\] The dissolution reaction is the formation of HgI_{4}{ }^{2-} complex ion, which occurs after continued addition of KI: \[HgI2(s) + 2I^-(aq) \rightarrow HgI_4^{2-}(aq)\]
02

Balance the equations

Both equations are already balanced: 1. Formation of the orange precipitate: \[Hg(NO_3)_2(aq) + 2KI(aq) \rightarrow HgI2(s) + 2KNO3(aq)\] 2. Dissolution of the orange precipitate: \[HgI2(s) + 2I^-(aq) \rightarrow HgI_4^{2-}(aq)\] In conclusion, the balanced chemical equations representing the formation of the orange precipitate and its dissolution when aqueous KI is added gradually to mercury(II) nitrate are: 1. \(Hg(NO_3)_2(aq) + 2KI(aq) \rightarrow HgI2(s) + 2KNO3(aq)\) 2. \(HgI2(s) + 2I^-(aq) \rightarrow HgI_4^{2-}(aq)\)

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Most popular questions from this chapter

The concentration of \(\mathrm{Ag}^{+}\) in a solution saturated with \(\mathrm{Ag}_{2} \mathrm{C}_{2} \mathrm{O}_{4}(s)\) is \(2.2 \times 10^{-4}\) M. Calculate \(K_{\text {sp }}\) for \(\mathrm{Ag}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\).

a. Calculate the molar solubility of AgBr in pure water. \(K_{\text {sp }}\) for \(\mathrm{AgBr}\) is \(5.0 \times 10^{-13}\) b. Calculate the molar solubility of AgBr in \(3.0 M \mathrm{NH}_{3}\). The overall formation constant for \(\mathrm{Ag}\left(\mathrm{NH}_{3}\right)_{2}^{+}\) is \(1.7 \times 10^{7}\), that is, \(\mathrm{Ag}^{+}(a q)+2 \mathrm{NH}_{3}(a q) \longrightarrow \mathrm{Ag}\left(\mathrm{NH}_{3}\right)_{2}^{+}(a q) \quad K=1.7 \times 10^{7}\) c. Compare the calculated solubilities from parts a and b. Explain any differences. d. What mass of AgBr will dissolve in \(250.0 \mathrm{~mL}\) of \(3.0 \mathrm{M} \mathrm{NH}_{3}\) ? e. What effect does adding \(\mathrm{HNO}_{3}\) have on the solubilities calculated in parts a and b?

Which of the following will affect the total amount of solute that can dissolve in a given amount of solvent? a. The solution is stirred. b. The solute is ground to fine particles before dissolving. c. The temperature changes.

Calculate the molar solubility of \(\mathrm{Mg}(\mathrm{OH})_{-}, K_{\mathrm{sp}}=8.9 \times 10^{-12}\).

A solution contains \(2.0 \times 10^{-3} \mathrm{M} \mathrm{Ce}^{3+}\) and \(1.0 \times 10^{-2} \mathrm{M} \mathrm{IO}_{3}^{3-}\). Will \(\mathrm{Ce}\left(\mathrm{IO}_{3}\right)_{3}(s)\) precipitate? \(\left[K_{\text {sp }}\right.\) for \(\mathrm{Ce}\left(\mathrm{IO}_{3}\right)_{3}\) is \(\left.3.2 \times 10^{-10} .\right]\)
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