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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

Calculate the concentrations of \(\mathrm{Cd}^{2+}, \mathrm{Cd}(\mathrm{CN})_{4}^{2-}\) and \(\mathrm{CN}^{-}\) at equilibrium when \(0.50 \mathrm{~g}\) of \(\mathrm{Cd}\left(\mathrm{NO}_{3}\right)_{2}\) dissolves in \(5.0 \times 10^{2} \mathrm{~mL}\) of \(0.50 \mathrm{M} \mathrm{NaCN}\).

How many grams of \(\mathrm{CaCO}_{3}\) will dissolve in \(3.0 \times\) \(10^{2} \mathrm{~mL}\) of \(0.050 \mathrm{M} \mathrm{Ca}\left(\mathrm{NO}_{3}\right)_{2} ?\)

A sample of \(0.1276 \mathrm{~g}\) of an unknown monoprotic acid was dissolved in \(25.0 \mathrm{~mL}\) of water and titrated with \(0.0633 \mathrm{M} \mathrm{NaOH}\) solution. The volume of base required to reach the equivalence point was \(18.4 \mathrm{~mL}\). (a) Calculate the molar mass of the acid. (b) After \(10.0 \mathrm{~mL}\) of base had been added to the titration, the \(\mathrm{pH}\) was determined to be 5.87 . What is the \(K_{\mathrm{a}}\) of the unknown acid?

Define buffer solution.

What is the \(\mathrm{pH}\) of the buffer \(0.10 \mathrm{M} \mathrm{Na}_{2} \mathrm{HPO}_{4} / 0.15 \mathrm{M}\) \(\mathrm{KH}_{2} \mathrm{PO}_{4} ?\)
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