Question

To convert \(\mathrm{NH}_{4}^{+}(\text {aq })\) to \(\mathrm{NH}_{3}(\mathrm{aq}),\) (a) add \(\mathrm{H}_{3} \mathrm{O}^{+}\) (b) raise the \(\mathrm{pH} ;\) (c) add \(\mathrm{KNO}_{3}(\mathrm{aq}) ;\) (d) add \(\mathrm{NaCl}\).

Short answer

The correct method to convert ammonium ion to ammonia is (b), which is to raise the pH of the solution.

Step-by-step solution

Evaluate the Conversion

Look at the given options and think about how they will affect the solution. Keep in mind that the goal is to remove a H+ ion from ammonium. (a) Adding \( \mathrm{H}_{3} \mathrm{O}^{+} \) would increase the H+ concentration, making the solution more acidic, which is the opposite of what's required. (c) Addition of \( \mathrm{KNO}_{3} \) and (d) \( \mathrm{NaCl} \) neither supply nor remove H+ ions, so they would not affect the conversion. (b) Raising the pH of the solution actually means reducing the concentration of H+ ions and making the solution more alkaline. This would promote the equilibrium to shift to ammonia side (Le Chatelier’s Principle) and facilitate the conversion of ammonium to ammonia.

Choose the correct option

According to the analysis of each option, it’s clear that the correct option is (b), which is to raise the pH of the solution. By increasing the alkalinity, one promotes the removal of a H+ ion from the ammonium ion, which triggers its conversion to ammonia.

Key concepts

pH Adjustment

The concept of pH adjustment refers to changing the pH level of a solution to favor a certain chemical reaction or equilibrium. pH is a measure of the acidity or basicity of a solution, with lower pH values indicating acidity and higher pH values indicating basicity.
In chemical reactions involving acids and bases, pH can influence the direction in which the reaction proceeds.
For example, in the conversion of ammonium ions \(\mathrm{NH}_{4}^{+}\), increasing the pH by adding a base reduces the concentration of hydrogen ions, \(\mathrm{H}^{+}\), in the solution. This reduction makes the environment less acidic and more alkaline.

• By raising the pH, you shift the equilibrium towards the production of ammonia \(\mathrm{NH}_3\).
• This is in line with Le Chatelier’s Principle, which predicts how changes in concentration, temperature, or pressure affect chemical equilibria.

To adjust pH effectively, it is crucial to choose the appropriate substances that will alter the H+ ion concentration without introducing other interfering ions into the reaction.

Ammonium to Ammonia Conversion

The conversion of ammonium \(\mathrm{NH}_{4}^{+}\) to ammonia \(\mathrm{NH}_3\) involves a deprotonation process, which is essentially removing a hydrogen ion \(\mathrm{H}^{+}\) from ammonium.
In an aqueous solution, ammonium exists in equilibrium with ammonia and hydronium ions. \\[\mathrm{NH}_{4}^{+}(aq) \rightleftharpoons \mathrm{NH}_3(aq) + \mathrm{H}^{+}(aq)\]Adjusting the pH to favor ammonia formation is essential in this reaction. When you increase the pH, you're effectively removing \(\mathrm{H}^{+}\) ions, which shifts the equilibrium towards ammonia production. This is a direct consequence of Le Chatelier’s Principle.

• Le Chatelier's Principle states that a system at equilibrium will adjust to minimize the effect of any changes, such as pH shifts.
• For instance, raising pH means reducing the concentration of \(\mathrm{H}^{+}\), driving more \(\mathrm{NH}_{4}^{+}\) to convert to \(\mathrm{NH}_3\).

Achieving this shift allows the ammonia to become the dominant species in the solution.

Acid-Base Equilibrium

Acid-base equilibrium explains how acids and bases interact in a solution, maintaining a balance between their different forms.
The equilibrium principle in chemistry governs how reversible reactions respond to external changes, such as pH variation.
In the equation \[ \mathrm{NH}_{4}^{+} + \mathrm{H}_2\mathrm{O} \rightleftharpoons \mathrm{NH}_3 + \mathrm{H}_3\mathrm{O}^{+}\] you can see how liquid water interacts with ammonium ions to create ammonia and hydronium ions, establishing an equilibrium.

• When the concentration of hydronium ions \(\mathrm{H}_3\mathrm{O}^{+}\) is increased, the equilibrium shifts towards \(\mathrm{NH}_{4}^{+}\), according to Le Chatelier's Principle, resulting in less ammonia.
• Conversely, reducing hydronium ions shifts the equilibrium towards \(\mathrm{NH}_3\), producing more ammonia.

Understanding this balance is important in manipulating chemical reactions for industrial or laboratory purposes, ensuring that the desired products are formed efficiently and in higher yield.