Question

The best way to ensure complete precipitation from saturated \(\mathrm{H}_{2} \mathrm{S}(\mathrm{aq})\) of a metal ion, \(\mathrm{M}^{2+}\), as its sulfide, \(\mathrm{MS}(\mathrm{s}),\) is to \((\mathrm{a})\) add an acid; \((\mathrm{b})\) increase \(\left[\mathrm{H}_{2} \mathrm{S}\right]\) in the solution; (c) raise the \(\mathrm{pH} ;\) (d) heat the solution.

Short answer

Increasing the concentration of H2S in the solution (b) and raising the pH (c) of the solution will ensure complete precipitation of a metal ion as its sulfide from a saturated H2S solution.

Step-by-step solution

Adding an Acid

Adding an acid to the saturated solution of H2S will increase the concentration of H+ ions in the solution. The increase in H+ will consume more H2S according to the equilibrium: \(H2S_{(aq)} \leftrightarrow 2H^+ + S^{2-}\). This is due to Le Chatelier's principle where an increase in reactant will be followed by the shift of the reaction to the right, in the direction that consumes the added substance. However, we're interested in formation of MS, which means we don't want to use up H2S. This option may not be most suitable for the formation of more precipitate.

Increasing [H2S]

Increasing the concentration of H2S in the solution will shift the reaction to produce more metal sulfide precipitate. According to Le Chatelier's principle, increasing the concentration of a reactant results in the reaction shifting in the direction that leads to the consumption of the reactant. This involves formation of MS. Therefore, this option will definitely increase the precipitation of a metal ion as its sulfide.

Raising the pH

By raising the pH of the solution, we essentially decrease the concentration of H+ ions. As per Le Chatelier's principle, this decrease in the concentration of H+ ions shifts the equilibrium towards the reactant side to compensate for the loss of H+. It thereby leads to the consumption of more H2S, which, in turn, will lead to the formation of more metal sulfide precipitate. This option is favorable to increase the precipitation of the metal ion as its sulfide.

Heating the Solution

Typically, heating a solution involves providing heat energy that can affect the position of equilibrium. However, the direction of shift depends on whether the reaction is endothermic or exothermic. In this case, precipitation reactions are usually exothermic, thus heating the solution would cause the equilibrium to shift to the left in order to absorb the provided heat, resulting in a decrease in precipitate formation. Therefore, this option would not favor the formation of more precipitate.

Key concepts

Le Chatelier's Principle

Imagine you're at a dance party where everyone is trying to maintain some personal space. Suddenly, more people enter the room, making it crowded. To adapt, you and others might move around to reestablish a comfortable space. In chemistry, Le Chatelier's principle is somewhat similar; it describes how a chemical system at equilibrium reacts to changes, like our partygoers adjusting to the crowd.

When a chemical system experiences a change in concentration, temperature, or pressure, the principle predicts that the system will adjust, or shift, to counteract the change and regain a new state of equilibrium. For instance, if more reactants are added to the system, the equilibrium will shift in such a way to use up those reactants, potentially forming more products. Conversely, removing a product will generally cause the system to produce more of that product. This adaptability helps chemists control the direction and extent of chemical reactions, crucial for processes such as precipitating a metal ion from a solution.

Saturated Solutions

Just as a sponge can only soak up a certain amount of water, a saturated solution has dissolved the maximum amount of solute that can be dissolved in the solvent at a given temperature and pressure. It's a delicate balance, really. Add a bit more solute into a saturated solution and it won't dissolve; instead, it remains undissolved and can form a precipitate.

In the context of our exercise, when metal ions in a saturated solution begin to form a solid, or precipitate, they leave the solution phase and form a more solid phase – the substance's solubility limit has been reached. It's important for students and chemists to understand this saturation point, because beyond this limit, any added solute will lead to precipitation, which is central to processes such as purifying compounds or, as in our textbook scenario, isolating a metal ion as its sulfide.

Equilibrium Shift

A seesaw in a playground is a great visual for equilibrium. When two children of equal weight sit on either end, the seesaw balances perfectly. But what happens if one child hops off? The seesaw will tilt, and the other child will hit the ground. Likewise, an equilibrium can be disturbed, or shifted, by changes in a chemical reaction. This is known as an equilibrium shift.

In the context of chemical reactions and precipitations, when a reaction is at equilibrium and you alter a condition like concentration or temperature, the equilibrium will shift to counter this change. This shift is crucial in the formation of precipitates; for example, increasing the concentration of a certain reactant may favor the production of a precipitate, as the system seeks to reduce the impact of that change, thereby 'shifting' towards product formation. Understanding this concept helps us manipulate conditions to control precipitation, and identify the best methods for inducing or preventing it, depending on our needs.