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Chapter 16: Problem 51

How can we predict whether a precipitate will form when two solutions are mixed?

Short Answer

Expert verified
Precipitates form when the product of the concentrations of the ions in the solution exceed the solubility product constant (Ksp). This can be predicted by identifying the ions present in the initial solutions, writing all possible reactions that might occur, checking solubility rules to identify potential precipitates, and finally comparing the reaction quotient (Q) against the solubility product constant (Ksp).

Step by step solution

01

Identify the ions present in the solutions

Each solution contains ions, which are dissolved particles that carry a charge. Identify the ions present in the initial solutions.
02

Write down possible precipitation reactions

Write down all possible reactions that could occur when the ions from the two solutions come into contact with each other. Precipitation occurs when a solid forms from these reaction.
03

Check solubility rules

Check the solubility rules to identify if any of the possible products are insoluble. If a compound is insoluble, it will be the precipitate.
04

Use solubility product constant (Ksp) to predict precipitate formation

If a solid compound is insoluble, it may not necessarily form a precipitate. One must check the reaction quotient (Q) against the solubility product constant (Ksp). A precipitate will form if \(Q > Ksp\).

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Most popular questions from this chapter

The solubility product of \(\mathrm{PbBr}_{2}\) is \(8.9 \times 10^{-6}\). Determine the molar solubility (a) in pure water, (b) in \(0.20 M \mathrm{KBr}\) solution, (c) in \(0.20 \mathrm{M} \mathrm{Pb}\left(\mathrm{NO}_{3}\right)_{2}\) solution.

The solubility product of \(\mathrm{Mg}(\mathrm{OH})_{2}\) is \(1.2 \times 10^{-11}\) What minimum \(\mathrm{OH}^{-}\) concentration must be attained (for example, by adding \(\mathrm{NaOH}\) ) to decrease the \(\mathrm{Mg}^{2+}\) concentration in a solution of \(\mathrm{Mg}\left(\mathrm{NO}_{3}\right)_{2}\) to less than \(1.0 \times 10^{-10} M ?\)

Which of the following solutions has the highest \(\left[\mathrm{H}^{+}\right]:\) (a) \(0.10 \mathrm{M} \mathrm{HF},\) (b) \(0.10 \mathrm{M} \mathrm{HF}\) in \(0.10 \mathrm{M} \mathrm{NaF}\) (c) \(0.10 \mathrm{M}\) HF in \(0.10 \mathrm{M} \mathrm{SbF}_{5} ?\) (Hint: SbF\(_{5}\) reacts with \(\mathrm{F}^{-}\) to form the complex ion \(\mathrm{SbF}_{6}^{-}\).)

Find the approximate \(\mathrm{pH}\) range suitable for the separation of \(\mathrm{Fe}^{3+}\) and \(\mathrm{Zn}^{2+}\) ions by precipitation of \(\mathrm{Fe}(\mathrm{OH})_{3}\) from a solution that is initially \(0.010 \mathrm{M}\) in both \(\mathrm{Fe}^{3+}\) and \(\mathrm{Zn}^{2+}\). Assume a 99 percent precipitation of \(\mathrm{Fe}(\mathrm{OH})_{3}\)

\(\mathrm{CaSO}_{4}\left(K_{\mathrm{sp}}=2.4 \times 10^{-5}\right)\) has a larger \(K_{\mathrm{sp}}\) value than that of \(\mathrm{Ag}_{2} \mathrm{SO}_{4}\left(K_{\mathrm{sp}}=1.4 \times 10^{-5}\right) .\) Does it follow that \(\mathrm{CaSO}_{4}\) also has greater solubility \((\mathrm{g} / \mathrm{L}) ?\)
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