Question

What solution can you add to each cation mixture to precipitate one cation while keeping the other cation in solution? Write a net ionic equation for the precipitation reaction that occurs. (a) \(\mathrm{Sr}^{2+}(a q)\) and \(\mathrm{Hg}_{2}{ }^{2+}(a q)\) (b) \(\mathrm{NH}_{4}^{+}(a q)\) and \(\mathrm{Ca}^{2+}(a q)\) (c) \(\mathrm{Ba}^{2+}(a q)\) and \(\mathrm{Mg}^{2+}(a q)\) (d) \(\mathrm{Ag}^{+}(a q)\) and \(\mathrm{Zn}^{2+}(a q)\)

Short answer

For Sr2+ and Hg22+, add Na2SO4 to precipitate SrSO4. For NH4+ and Ca2+, add Na2C2O4 to precipitate CaC2O4. For Ba2+ and Mg2+, add Na2CO3 to precipitate BaCO3. For Ag+ and Zn2+, add NaCl to precipitate AgCl.

Step-by-step solution

- Determine the Precipitation Reagent for Sr2+ and Hg22+

To precipitate only one cation from the mixture of Sr2+ and Hg22+, we choose a reagent that forms an insoluble compound with one of them while keeping the other in solution. We can use a sulfate ion (SO4^2-) source such as sodium sulfate (Na2SO4), as SrSO4 is insoluble but Hg2SO4 is soluble.

- Write the Net Ionic Equation for Sr2+ Precipitation

The net ionic equation for the precipitation of Sr2+ as strontium sulfate is: \[\mathrm{Sr}^{2+}(aq) + \mathrm{SO}_4^{2-}(aq) \rightarrow \mathrm{SrSO}_4(s)\]

- Determine the Precipitation Reagent for NH4+ and Ca2+

To precipitate only one cation from the mixture of NH4+ and Ca2+, we choose a reagent that forms an insoluble compound with one of them while keeping the other in solution. We can use oxalate ion (C2O4^2-) source such as sodium oxalate (Na2C2O4), as CaC2O4 is insoluble but NH4C2O4 is soluble.

- Write the Net Ionic Equation for Ca2+ Precipitation

The net ionic equation for the precipitation of Ca2+ as calcium oxalate is: \[\mathrm{Ca}^{2+}(aq) + \mathrm{C}_2\mathrm{O}_4^{2-}(aq) \rightarrow \mathrm{CaC}_2\mathrm{O}_4(s)\]

- Determine the Precipitation Reagent for Ba2+ and Mg2+

To precipitate only one cation from the mixture of Ba2+ and Mg2+, we choose a reagent that forms an insoluble compound with one of them while keeping the other in solution. We can use a carbonate ion (CO3^2-) source such as sodium carbonate (Na2CO3), as BaCO3 is insoluble but MgCO3 is soluble.

- Write the Net Ionic Equation for Ba2+ Precipitation

The net ionic equation for the precipitation of Ba2+ as barium carbonate is: \[\mathrm{Ba}^{2+}(aq) + \mathrm{CO}_3^{2-}(aq) \rightarrow \mathrm{BaCO}_3(s)\]

- Determine the Precipitation Reagent for Ag+ and Zn2+

To precipitate only one cation from the mixture of Ag+ and Zn2+, we choose a reagent that forms an insoluble compound with one of them while keeping the other in solution. We can use chloride ion (Cl-) source such as sodium chloride (NaCl), as AgCl is insoluble but ZnCl2 is soluble.

- Write the Net Ionic Equation for Ag+ Precipitation

The net ionic equation for the precipitation of Ag+ as silver chloride is: \[\mathrm{Ag}^{+}(aq) + \mathrm{Cl}^{-}(aq) \rightarrow \mathrm{AgCl}(s)\]

Key concepts

Net Ionic Equations

A net ionic equation is a simplified chemical equation that shows only the reacting ions in a precipitation reaction. When two aqueous solutions containing different ions are mixed, a precipitation reaction can occur if any combination of ions forms an insoluble compound, known as a precipitate. The net ionic equation represents the actual chemical change by only including the ions that participate in the formation of the precipitate and omitting the spectator ions that do not change during the reaction.

For example, consider the precipitation of strontium sulfate from a mixture of aqueous strontium and sulfate ions:
\[\begin{equation}\mathrm{Sr}^{2+}(aq) + \mathrm{SO}_4^{2-}(aq) \rightarrow \mathrm{SrSO}_4(s)\end{equation}\]In this equation, strontium ions (\(\mathrm{Sr}^{2+}\)) combine with sulfate ions (\(\mathrm{SO}_4^{2-}\)) to form strontium sulfate, a solid precipitate (\(\mathrm{SrSO}_4(s)\)). Creating a net ionic equation involves several steps which include writing a balanced molecular equation, separating the strong electrolytes into ions (complete ionic equation), and then identifying and removing the spectator ions, resulting in the net ionic equation.

Solubility Rules

Understanding which compounds are soluble or insoluble in water is essential for predicting the outcomes of precipitation reactions. Solubility rules are a set of guidelines that help in determining the solubility of different ionic compounds in water. Certain ions tend to form soluble compounds, such as those containing alkali metal ions and the ammonium ion (\(\mathrm{NH}_4^+\)), while others often form precipitates, like carbonate (\(\mathrm{CO}_3^{2-}\)) and phosphate ions (\(\mathrm{PO}_4^{3-}\)).

For example, all nitrates (\(\mathrm{NO}_3^-\)) are soluble in water, which means a nitrate compound would not precipitate no matter what cation it is paired with. In contrast, most sulfates are soluble except for those of barium, strontium, and lead. This explains why strontium sulfate would precipitate out of a mixture with sulfate ions while other sulfates would remain in solution. Knowing these rules is vital for selecting appropriate reagents in precipitation methods for separating cations.

Cation Separation

Cation separation relies on the selective precipitation of specific cations from a mixture based on their varied solubility with certain anions. It is a common analytical procedure in chemistry used to identify the presence of various cations in a solution. By adding a reagent that forms an insoluble compound with one cation but not with others, chemists can separate and identify the cations through stepwise precipitation.

For instance, to separate \(\mathrm{Sr}^{2+}\) from \(\mathrm{Hg}_2^{2+}\), a source of sulfate ions can be used since \(\mathrm{SrSO}_4\) is insoluble while \(\mathrm{Hg}_2\mathrm{SO}_4\) remains soluble. Similarly, to separate \(\mathrm{NH}_4^+\) from \(\mathrm{Ca}^{2+}\), using an oxalate ion source will precipitate calcium as calcium oxalate, leaving ammonium in solution. The separated cation can then be further analyzed, if necessary, with additional techniques such as flame tests or spectroscopy. This process is crucial for the analysis of complex mixtures in both qualitative and quantitative analyses.