Question

Which of the following would be most effective, and which would be least effective, in reducing the concentration of the complex ion \(\left[\mathrm{Zn}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}\) in a solution: \(\mathrm{HCl}, \mathrm{NH}_{3},\) or \(\mathrm{NH}_{4} \mathrm{Cl} ?\) Explain your choices.

Short answer

\(\mathrm{HCl}\) would be the most effective in reducing the concentration of the complex ion, while \(\mathrm{NH}_{3}\) would be the least effective, likely increasing its concentration. \(\mathrm{NH}_{4}\mathrm{Cl}\) would likely fall in between the two in its effectiveness.

Step-by-step solution

Introduction to the Problem

First, consider how different ions will affect the equilibrium of the complex ion \( \left[\mathrm{Zn}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+} \). The possible clues are: \n 1. \( \mathrm{HCl} \) will provide \( \mathrm{H}^{+} \) ions, which can react with \( \mathrm{NH}_{3} \) to form \( \mathrm{NH}_{4}^{+} \), and thus remove \( \mathrm{NH}_{3} \) from the solution. This will surely disturb the equilibrium and likely reduce the concentration of the complex ion.\n 2. \( \mathrm{NH}_{3} \) will increase the concentration of \( \mathrm{NH}_{3} \) in the solution, which would logically push the equilibrium towards more formation of the complex ion, hence increasing its concentration.\n 3. \( \mathrm{NH}_{4}\mathrm{Cl} \) provides \( \mathrm{NH}_{4}^{+} \) ions, which can produce \( \mathrm{NH}_{3} \) and \( \mathrm{H}^{+} \), so this could have a dual effect.

Interaction with HCl

\(\mathrm{HCl}\) would be the most effective in reducing the complex ion concentration since it provides \(\mathrm{H}^{+}\) ions which react with \(\mathrm{NH}_{3}\) to form \(\mathrm{NH}_{4}^{+}\), thus removing \(\mathrm{NH}_{3}\) from the solution. This would cause more of the \(\left[\mathrm{Zn}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}\) ion to disassociate in order to replace the \(\mathrm{NH}_{3}\) that was lost, leading to a decrease in the concentration of the complex ion.

Interaction with NH3

\(\mathrm{NH}_{3}\) would be least effective at reducing the concentration, in fact it would likely increase it, since adding more \(\mathrm{NH}_{3}\) to the solution would push the equilibrium towards the formation of more complex ion, hence raising its concentration.

Interaction with NH4Cl

\(\mathrm{NH}_{4}\mathrm{Cl}\) falls somewhere in between \(\mathrm{NH}_{3}\) and \(\mathrm{HCl}\). It could possibly reduce some concentration of the complex ion by providing \(\mathrm{NH}_{4}^{+}\) ions, which can produce \(\mathrm{NH}_{3}\) and \(\mathrm{H}^{+}\), removing some \(\mathrm{NH}_{3}\) from the solution. However, this effect would not be as intense as that of \(\mathrm{HCl}\), due to the availability of both \( \mathrm{NH}_{3} \) and \( \mathrm{H}^{+} \) ions.

Key concepts

Chemical Equilibrium

Understanding chemical equilibrium is crucial when studying reactions, such as those involving complex ions in the presence of other substances. A chemical reaction is said to have reached equilibrium when the rate of the forward reaction equals the rate of the reverse reaction, resulting in no net change in the concentrations of the reactants and products over time.

For the complex ion \(\left[\mathrm{Zn}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}\), the equilibrium can be expressed as:\[\mathrm{Zn}^{2+} + 4\mathrm{NH}_{3} \leftrightarrow \left[\mathrm{Zn}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}\]
Any change that alters the concentration of either reactants or products shifts the position of equilibrium. This can be adding reactants like \(\mathrm{NH}_{3}\), or products, or by removing them through chemical reactions with other added substances, such as \(\mathrm{HCl}\) or \(\mathrm{NH}_{4}\mathrm{Cl}\).

Le Chatelier's Principle

Le Chatelier’s Principle offers a predictive tool about how a system at equilibrium responds to disturbances. It states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium moves to counteract the change.

In context of the complex ion \(\left[\mathrm{Zn}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}\), adding \(\mathrm{HCl}\) introduces \(\mathrm{H}^{+}\) ions, and these react with \(\mathrm{NH}_{3}\), forming \(\mathrm{NH}_{4}^{+}\) and decreasing \(\mathrm{NH}_{3}\) concentration. According to Le Chatelier’s Principle, the equilibrium will shift to increase the concentration of \(\mathrm{NH}_{3}\) by dissociating some of the complex ion, thus reducing its concentration in the solution. On the flip side, adding \(\mathrm{NH}_{3}\) would shift the equilibrium towards more complex ion formation.

Ammonia Ligand Reactions

In coordination chemistry, ammonia ligand reactions are crucial for the formation of complex ions like \(\left[\mathrm{Zn}\left(\mathrm{NH}_{3}\right)_{4}\right]^{2+}\). Ammonia (\(\mathrm{NH}_{3}\)) acts as a ligand, a molecule that donates a pair of electrons to a central metal ion, resulting in a coordination complex.

When \(\mathrm{NH}_{3}\) is plentiful in the solution, it readily binds to metal ions like \(\mathrm{Zn}^{2+}\) and increases the stability of the complex ion. If the concentration of free \(\mathrm{NH}_{3}\) is depleted, say, by reaction with \(\mathrm{H}^{+}\) to form \(\mathrm{NH}_{4}^{+}\), then the complex will tend to break down to release \(\mathrm{NH}_{3}\), restoring some of the lost ligands.

Equilibrium Disturbances

Equilibrium disturbances refer to any change in the conditions of a system at equilibrium that impacts the concentrations of either reactants or products. Such disturbances include changes in concentration, pressure, volume, or temperature.

In the given problem, adding \(\mathrm{HCl}\) leads to a significant disturbance as it reacts with one of the constituents of the complex, \(\mathrm{NH}_{3}\), to form \(\mathrm{NH}_{4}^{+}\). This effectively pulls \(\mathrm{NH}_{3}\) away from the complex ion, thus shifting the equilibrium per Le Chatelier's Principle to the left, reducing the complex ion concentration. Alternatively, the addition of \(\mathrm{NH}_{3}\) would not disturb the equilibrium negatively; in fact, it would promote the formation of more complex ion.