Question

The reagents, \(\mathrm{NH}_{4} \mathrm{Cl}\) and aqueous \(\mathrm{NH}_{3}\) will precipitate (a) \(\mathrm{Ca}^{2+}\) (b) \(\mathrm{Al}^{3+}\) (c) \(\mathrm{Mg}^{2+}\) (d) \(\mathrm{Zn}^{2+}\)

Short answer

(b) \(\mathrm{Al}^{3+}\) will precipitate.

Step-by-step solution

Understanding the Reaction

The reagents \(\mathrm{NH}_{4} \mathrm{Cl}\) and aqueous \(\mathrm{NH}_{3}\) together form an ammonium buffer solution that adjusts the pH of the solution. This buffer solution gives a weak basic condition that allows certain metal cations to precipitate as hydroxides.

Analyzing Options for Precipitation

Not all metallic cations will precipitate in the presence of \(\mathrm{NH}_{4} \mathrm{Cl}\) and \(\mathrm{NH}_{3}\). We need to determine which metals are capable of forming insoluble hydroxides under basic conditions. The metals in question are \(\mathrm{Ca}^{2+}\), \(\mathrm{Al}^{3+}\), \(\mathrm{Mg}^{2+}\), and \(\mathrm{Zn}^{2+}\).

Key concepts

Ammonium Buffer

Ammonium buffers play a crucial role in controlling the pH of a solution. When ammonium chloride (\(\mathrm{NH}_4\mathrm{Cl}\)) is added to aqueous ammonia (\(\mathrm{NH}_3\)), an ammonium buffer system forms. This combination stabilizes the pH around a mildly basic range. This buffer helps maintain an environment where pH-sensitive chemical reactions can take place smoothly.

• Ammonium chloride dissolves in water, releasing \(\mathrm{NH}_4^+\) ions, which slightly acidify the solution by donating protons (H\(^+\)).
• Aqueous ammonia contributes ammonia molecules, which can accept protons, resulting in the formation of \(\mathrm{NH}_4^+\) ions and hydroxide ions (\(\mathrm{OH}^-\)).
• The balance between \(\mathrm{NH}_4^+\) and \(\mathrm{OH}^-\) ions lends the buffer its ability to resist drastic changes in pH.

This stable environment is essential when forming metal hydroxides, as too much acidity or basicity might inhibit or falsely initiate precipitation.

Metal Hydroxides

Metal hydroxides are compounds formed when metal cations react with hydroxide ions (\(\mathrm{OH}^-\)). These compounds play a significant role in various chemical reactions, and their solubility often dictates whether a precipitate will form in a given reaction. The extent to which metal hydroxides are soluble or insoluble directly affects their ability to precipitate in a solution.

• In the context of precipitation reactions, the formation of metal hydroxides depends significantly on the solution's pH.
• The buffer system provided by ammonium chloride and ammonia helps achieve the appropriate pH range for the precipitation of some metal hydroxides.
• Different metals have different solubility profiles; therefore, only certain metals will precipitate as hydroxides under slightly basic conditions induced by ammonium buffers.

Understanding these properties allows chemists to predict which metal cations will likely form a precipitate in a given reaction.

Insoluble Hydroxides

Insoluble hydroxides are a subset of metal hydroxides characterized by their limited solubility in water. When dealing with precipitation reactions, it's crucial to identify which metal hydroxides are insoluble, as this determines their ability to form a solid precipitate when a buffer influences the pH.

• For a metal cation to form an insoluble hydroxide and thus a precipitate, it must react with hydroxide ions to form a compound that does not dissolve easily in water.
• The metals (\(\mathrm{Ca}^{2+}\), \(\mathrm{Al}^{3+}\), \(\mathrm{Mg}^{2+}\), and \(\mathrm{Zn}^{2+}\)) each have different propensities to form such hydroxides based on the solution's conditions.
• An ammonium buffer creates the requisite environment for this, but only \(\mathrm{Al}^{3+}\) and \(\mathrm{Zn}^{2+}\) will typically form insoluble hydroxides under these conditions, resulting in their precipitation from the solution.

Grasping which combinations lead to insoluble products is vital for tasks such as metal ion separation and water treatment processes.