Question

Suppose you have a solution that might contain any or all of the following cations: \(\mathrm{Ni}^{2+}, \mathrm{Ag}^{+}, \mathrm{Sr}^{2+},\) and \(\mathrm{Mn}^{2+} .\) Addition of HCl solution causes a precipitate to form. After filtering off the precipitate, \(\mathrm{H}_{2} \mathrm{SO}_{4}\) solution is added to the resulting solution and another precipitate forms. This is filtered off, and a solution of \(\mathrm{NaOH}\) is added to the resulting solution. No precipitate is observed. Which ions are present in each of the precipitates? Which of the four ions listed above must be absent from the original solution?

Short answer

The precipitate formed after adding HCl consists of \(\mathrm{AgCl}\), and the precipitate formed after adding \(\mathrm{H}_{2} \mathrm{SO}_{4}\) consists of \(\mathrm{SrSO}_{4}\). No precipitate is observed after adding \(\mathrm{NaOH}\), indicating the absence of \(\mathrm{Ni}^{2+}\) and \(\mathrm{Mn}^{2+}\) ions in the original solution. Thus, the original solution contains \(\mathrm{Ag}^{+}\) and \(\mathrm{Sr}^{2+}\) cations.

Step-by-step solution

1. Precipitate with HCl

After the addition of HCl, a precipitate forms. This indicates the presence of cations that form insoluble chloride salts. We need to check the solubility rules for each possible cation: - \(\mathrm{Ni}^{2+}\) and \(\mathrm{Mn}^{2+}\): Both nickel(II) and manganese(II) chlorides are soluble in water, so they are not responsible for the precipitate. - \(\mathrm{Ag}^{+}\): Silver chloride (AgCl) is insoluble in water, so the silver ions would form a precipitate with the chloride ions. - \(\mathrm{Sr}^{2+}\): Strontium chloride (SrCl2) is also soluble in water, so it would not form a precipitate. Thus, the precipitate formed after adding HCl consists of \(\mathrm{AgCl}\).

2. Precipitate with \(\mathrm{H}_{2} \mathrm{SO}_{4}\)

After the first precipitate has been filtered, we add H2SO4 to the resulting solution and observe another precipitate formation. We need to examine the solubility rules for sulfate salts containing the remaining possible cations: - \(\mathrm{Ni}^{2+}\): Nickel(II) sulfate (NiSO4) is soluble in water, so it would not form a precipitate. - \(\mathrm{Sr}^{2+}\): Strontium sulfate (SrSO4) is slightly insoluble in water, so the strontium ions would form a precipitate with the sulfate ions. - \(\mathrm{Mn}^{2+}\): Manganese(II) sulfate (MnSO4) is soluble in water, so it would not form a precipitate. Thus, the precipitate formed after adding H2SO4 consists of \(\mathrm{SrSO}_{4}\).

3. Precipitate with \(\mathrm{NaOH}\)

After the second precipitate has been filtered, we add NaOH to the resulting solution and observe no precipitate formation. This indicates that none of the remaining possible cations in the solution forms insoluble hydroxide salts. - \(\mathrm{Ni}^{2+}\): Nickel(II) hydroxide (Ni(OH)2) is insoluble in water. However, since no precipitate is observed, it indicates that nickel ions are not present in the original solution. - \(\mathrm{Mn}^{2+}\): Manganese(II) hydroxide (Mn(OH)2) is also insoluble in water, but as there is no precipitate formed, manganese ions are not present in the original solution. Based on the results, we conclude that the original solution contains \(\mathrm{Ag}^{+}\) and \(\mathrm{Sr}^{2+}\) cations, while \(\mathrm{Ni}^{2+}\) and \(\mathrm{Mn}^{2+}\) ions are absent.

Key concepts

Solubility Rules

Solubility rules are essential for understanding what substances will dissolve in water to form aqueous solutions. They are handy for predicting whether a precipitate will form when two solutions are mixed. Here are some simple guidelines:

• Most nitrate (\( ext{NO}_3^- \)) and acetate (\( ext{C}_2 ext{H}_3 ext{O}_2^- \)) salts are soluble.
• Most chloride (\( ext{Cl}^- \)), bromide (\( ext{Br}^- \)), and iodide (\( ext{I}^- \)) salts are soluble, except those of lead (\( ext{Pb}^{2+} \)), silver (\( ext{Ag}^+ \)), and mercury (\( ext{Hg}_2^{2+} \)).
• Most sulfate salts (\( ext{SO}_4^{2-} \)) are soluble, except for those of barium (\( ext{Ba}^{2+} \)), lead (\( ext{Pb}^{2+} \)), calcium (\( ext{Ca}^{2+} \)), and strontium (\( ext{Sr}^{2+} \)).
• Hydroxides (\( ext{OH}^- \)) are generally insoluble, except for those of the alkali metals and barium (\( ext{Ba}(OH)_2 \)).
• Carbonates (\( ext{CO}_3^{2-} \)) and phosphates (\( ext{PO}_4^{3-} \)) are generally insoluble, except for those of the alkali metals and ammonium (\( ext{NH}_4^+ \)).

By applying these rules to specific cations and anions, one can determine the likelihood of a precipitate forming.

Insoluble Salts

Insoluble salts are compounds that do not dissolve well in water. They remain as solid particles or precipitates when mixed with water or other solutions. Understanding why some salts are insoluble while others are not is important.
Insoluble salts form due to strong ionic bonds within the compound that the surrounding water molecules can't easily disrupt. The concept is connected with lattice energy, which is the energy required to separate a mole of an ionic solid into gaseous ions.
Examples include:

• Silver chloride (\( ext{AgCl} \)) does not dissolve well, making it an insoluble salt.
• Calcium carbonate (\( ext{CaCO}_3 \)) is also insoluble in water.

When such salts are mixed with solutions containing their respective ions,
a precipitate forms because the ionic compound reassembles faster than it dissolves.

Precipitation Reactions

Precipitation reactions occur when two aqueous solutions are mixed, resulting in the formation of an insoluble compound. This insoluble compound, or precipitate, comes out of the solution as a solid. These reactions are depicted by double displacement reactions where cations and anions swap partners.
For instance, when a solution containing silver ions (\( ext{Ag}^+ \)) is mixed with chloride ions (\( ext{Cl}^- \)), silver chloride (\( ext{AgCl} \)) forms as a precipitate.
We can write this reaction as follows:y(\[ ext{Ag}^+ (aq) + ext{Cl}^- (aq) \rightarrow ext{AgCl} (s) \])This example illustrates the simple rules of precipitation:

• The reaction should lead to the formation of an insoluble compound based on solubility guidelines.
• Only the ions involved in forming the solid are noted in the net ionic equation.

Precipitation reactions are crucial in many chemical processes, helping in the identification of ions and in various applications such as water treatment.