Problem 99 - A Complete Resource Book in Chemistry for JEE Main

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) \(\mathrm{Zn}^{2+}\)

Step-by-Step Solution

Answer

(c) \( \mathrm{Mg}^{2+} \) will precipitate.

Step 1: Identifying the Chemical Reaction

When \( ext{NH}_4 ext{Cl}\) and aqueous \( ext{NH}_3\) are mixed, they create an ammonium buffer solution. This solution reacts with metal ions to form hydroxide precipitates if the solubility product ( ext{K}_{sp}) of the metal hydroxide is exceeded.

Step 2: Considering the Solubility of Metal Hydroxides

Calcium ( ext{Ca}^{2+}), aluminum ( ext{Al}^{3+}), magnesium ( ext{Mg}^{2+}), and zinc ( ext{Zn}^{2+}) can form different metal hydroxides. Evaluate their general solubility in basic conditions. \( ext{Ca(OH)}_2\) is slightly soluble, \( ext{Mg(OH)}_2\) is not very soluble, \( ext{Al(OH)}_3\) and \( ext{Zn(OH)}_2\) are somewhat insoluble but are more reactive than magnesium hydroxide.

Step 3: Applying Ammonium Buffer Concepts

In an ammonium buffer solution, which consists of \( ext{NH}_4^+\) and \( ext{NH}_3\), the pH is moderate, and favors the precipitation of metals that form insoluble hydroxides at relatively higher pH, such as \( ext{Mg(OH)}_2\).

Step 4: Conclusion about the Precipitating Metal

Given the pH maintained by the ammonium buffer, \( ext{Mg}^{2+}\) will precipitate as \( ext{Mg(OH)}_2\), whereas \( ext{Ca(OH)}_2\) remains slightly soluble, and other hydroxides are expected to either remain in solution or require greater alkalinity. Therefore, the metal that will precipitate in this solution is magnesium, \( ext{Mg}^{2+}\).

Key Concepts

Ammonium Buffer

An ammonium buffer comprises a mix of ammonium chloride \(\mathrm{NH}_4\mathrm{Cl}\) and ammonia \(\mathrm{NH}_3\). Together, they create a system that can resist changes in pH when acids or bases are introduced. This is crucial in reactions where pH must remain stable to achieve desired results, such as precipitating certain metal ions from a solution.
A buffer works by utilizing both components: the weak acid \(\mathrm{NH}_4^+\) and the weak base \(\mathrm{NH}_3\). When a small amount of acid is added to the solution, the \(\mathrm{NH}_3\) neutralizes it and forms more \(\mathrm{NH}_4^+\). Conversely, if a base is added, \(\mathrm{NH}_4^+\) converts to \(\mathrm{NH}_3\), thus maintaining a nearly constant pH level.

• Buffers make solutions resistant to drastic pH changes.
• They are pivotal in biochemical processes and laboratory experiments.
• An ammonium buffer provides a moderate pH, often around 9 to 11, depending on concentrations.

Understanding the role of ammonium buffers is essential when examining reactions involving metal ions in aqueous solutions.

Metal Hydroxides

Metal hydroxides are compounds formed when metal ions react with hydroxide ions (OH\(^{-}\)). They each have their own solubility characteristics that depend on the surrounding environment, especially the pH level. In the context of an ammonium buffer, the pH facilitates reactions where certain metal hydroxides become insoluble, leading to precipitation.
For instance, magnesium hydroxide \(\mathrm{Mg(OH)}_2\) is known to have low solubility in water. In a buffered solution with moderate pH, \(\mathrm{Mg}^{2+}\) ions will prefer to form \(\mathrm{Mg(OH)}_2\) and precipitate out. Contrast this with calcium hydroxide \(\mathrm{Ca(OH)}_2\), which is slightly soluble under similar conditions.

• Solubility is affected by the solution pH.
• Magnesium hydroxide is almost insoluble in neutral pH but precipitates at moderate pH.
• Aluminum and zinc hydroxides also tend to be less soluble but are influenced by more complex chemical equilibria.

Predicting whether a metal will precipitate as a hydroxide necessitates understanding its K\(_{sp}\) or solubility product values, which indicate how likely a compound is to remain in solution.

Precipitation Reactions

Precipitation reactions occur when ions in solution combine to form an insoluble compound, which appears as a solid, often referred to as a precipitate. The formation of a precipitate is heavily dependent on the conditions of the reaction, particularly the concentration of ions and the pH of the solution.
In an ammonium buffer system, the solution pH plays a significant role in whether a metal hydroxide becomes a precipitate. For example, when the \(\mathrm{Mg}^{2+}\) ions are added to an ammonium buffer, the moderate pH and the equilibrium position favor the formation of \(\mathrm{Mg(OH)}_2\) that falls out of the solution. This indicates a shift in equilibrium due to exceeding the solubility product, K\(_{sp}\).

• Precipitation occurs when the ionic product exceeds the K\(_{sp}\).
• The pH of the solution is crucial; higher pH can lead to more precipitation.
• Different metal ions require different conditions for precipitation.

Thus, controlling the pH with buffers is pivotal in designing experiments that exploit the precipitation of specific metal ions.