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) \(\mathrm{Fe}^{2+}(a q)\) and \(\mathrm{Pb}^{2+}(a q)\) (b) \(\mathrm{K}^{+}(a q)\) and \(\mathrm{Ca}^{2+}(a q)\) (c) \(\mathrm{Ag}^{+}(a q)\) and \(\mathrm{Ba}^{2+}(a q)\) (d) \(\mathrm{Cu}^{2+}(a q)\) and \(\mathrm{Hg}_{2}{ }^{2+}(a q)\)

Understanding solubility rules is essential for predicting whether a compound will dissolve in water or form a precipitate. These rules are a set of guidelines that help us determine the solubility of various ionic compounds.

For instance, most sulfate (SO₄²⁻) salts are soluble, with exceptions like barium sulfate (BaSO₄) and lead sulfate (PbSO₄), which are insoluble. Similarly, most compounds of alkali metal ions and ammonium are soluble, which is why compounds containing potassium (K⁺) and sodium (Na⁺) ions dissolve readily in water. Conversely, carbonates (CO₃²⁻) are typically insoluble except when paired with alkali metal ions, as in the case of potassium carbonate (K₂CO₃).

These rules are crucial when selecting a reagent to precipitate a specific cation out of a mixture, as demonstrated in the original exercise. By adding the appropriate reagent that forms an insoluble compound with the targeted cation, we can selectively precipitate it out from the solution.

A net ionic equation represents the chemical reaction that occurs in solution at the ionic level. It focuses on the species that participate directly in the reaction, omitting the spectator ions that do not change during the process.

For example, in the original exercise's step 2, to precipitate lead(II) sulfate, we combine lead (II) ions with sulfate ions. The net ionic equation is simplified to \(\mathrm{Pb}^{2+}(aq) + \mathrm{SO}_4^{2-}(aq) \rightarrow \mathrm{PbSO}_4(s)\), showing only the ions that form the precipitate and leaving out the sodium ions which are spectator ions.

Net ionic equations provide a clearer picture of the underlying chemistry in precipitation reactions, helping students understand the process without the complexity of the complete ionic reaction.

The concept of selective precipitation involves adding a reagent to a mixture of ions that will cause one of them to precipitate, while the others remain in solution. This technique is used for separating ions in a mixture based on the differences in their solubility.

As seen in the original exercise, by choosing an appropriate reagent, we can selectively remove a specific ion. For instance, adding sodium sulfate to a mixture of Fe²⁺ and Pb²⁺ will precipitate out the Pb²⁺ as lead sulfate, while Fe²⁺ stays dissolved as a sulfate. This method is not only vital in chemical analysis but also in metal purification and wastewater treatment processes.

Executing selective precipitation effectively requires a thorough understanding of the solubility rules to predict the reagent needed for the desired separation.

A double displacement reaction, also known as a metathesis reaction, occurs when ions from two compounds exchange places in aqueous solution to form two new compounds. One of the products is typically a precipitate, a gas, or a weakly dissociated molecule like water.

For example, when sodium carbonate (Na₂CO₃) is added to a solution of calcium chloride (CaCl₂), calcium carbonate precipitates, and sodium chloride remains in solution. The reaction can be represented by the ionic equation: \(\mathrm{Na}_2\mathrm{CO}_3(aq) + \mathrm{CaCl}_2(aq) \rightarrow \mathrm{CaCO}_3(s) + 2\mathrm{NaCl}(aq)\). In the exercise provided, the reactions described are all double displacement reactions leading to the formation of an insoluble salt, demonstrating the kind of transformation characteristic of this reaction type.

Recognizing double displacement reactions is critical for predicting the outcomes of mixing various ionic compounds in solution, especially in the context of forming a precipitate.