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Chapter 17: Problem 31

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

Short Answer

Expert verified
To predict whether a precipitate will form when two solutions are mixed, one can use the solubility rules and the complete ionic equation for the reaction. If any of the potential product combinations is insoluble according to the solubility rules, a precipitate will form.

Step by step solution

01

Understand the Solubility Rules

The first step is to familiarize yourself with the solubility rules. They give insights on whether certain compounds will dissolve in water. For example, most nitrates (NO3-) are soluble, while most sulfates (SO4 2-) are soluble, except those of barium (Ba2+), calcium (Ca2+), and lead(Pb2+) which are insoluble.
02

Write out the Complete Ionic Equation

With the knowledge of the solubility rules, the next step is to write out the complete ionic equation for the potential reaction. This includes all ions present in the reaction. For example, if copper sulfate (CuSO4) - a soluble compound, is mixed with sodium phosphate (Na3PO4) - also a soluble compound, the ionic equation would be: \(Cu^{2+}(aq) + SO4^{2-}(aq) + 3Na^{+}(aq) + PO4^{3-}(aq) \rightarrow\) products
03

Use the Solubility Rules to Predict the Precipitate

The final step involves using the solubility rules to predict which, if any, of the potential products are insoluble. If there are any insoluble products, a precipitate will form. Using our example, the potential products from the combination of these ions are: copper phosphate (Cu3(PO4)2) and sodium sulfate (Na2SO4). Copper Phosphate is insoluble according to the solubility rules, meaning this combination will form a precipitate. The balanced total ionic equation showing the formation of the precipitate is: \(3Cu^{2+}(aq) + 2PO4^{3-}(aq) \rightarrow Cu3(PO4)2(s)\). The Na+ and SO4 2- ions remain in solution and do not form a precipitate.

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Most popular questions from this chapter

The buffer range is defined by the equation \(\mathrm{pH}=\) \(\mathrm{p} K_{\mathrm{a}} \pm 1 .\) Calculate the range of the ratio [conjugate base \(]\) / [acid] that corresponds to this equation.

The molar mass of a certain metal carbonate, \(\mathrm{MCO}_{3}\), can be determined by adding an excess of \(\mathrm{HCl}\) acid to react with the carbonate and then "back-titrating" the remaining acid with \(\mathrm{NaOH}\). (a) Write an equation for these reactions. (b) In a certain experiment, \(20.00 \mathrm{~mL}\) of \(0.0800 \mathrm{M} \mathrm{HCl}\) were added to a \(0.1022-\mathrm{g}\) sample of \(\mathrm{MCO}_{3}\). The excess HCl required \(5.64 \mathrm{~mL}\) of \(0.1000 \mathrm{M} \mathrm{NaOH}\) for neutralization. Calculate the molar mass of the carbonate and identify \(\mathrm{M}\).

Define solubility, molar solubility, and solubility product. Explain the difference between solubility and the solubility product of a slightly soluble substance such as \(\mathrm{BaSO}_{4}\) .

Both \(\mathrm{KCl}\) and \(\mathrm{NH}_{4} \mathrm{Cl}\) are white solids. Suggest one reagent that would enable you to distinguish between these two compounds.

A student wishes to prepare a buffer solution at \(\mathrm{pH}=\) \(8.60 .\) Which of these weak acids should she choose and why: HA \(\left(K_{\mathrm{a}}=2.7 \times 10^{-3}\right),\) HB \(\left(K_{\mathrm{a}}=4.4 \times\right.\) \(10^{-6}\) ), or HC \(\left(K_{\mathrm{a}}=2.6 \times 10^{-9}\right) ?\)
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