Question

Suggest explanations for the following observations. (a) In aqueous solution, \(\mathrm{AgNO}_{3}\) and \(\mathrm{KCl}\) react to give a precipitate of AgCl, whereas in liquid \(\mathrm{NH}_{3}, \mathrm{KNO}_{3}\) and \(\mathrm{AgCl}\) react to produce a precipitate of \(\mathrm{KCl}\). (b) \(\mathrm{Mg}\) dissolves in a concentrated solution of \(\mathrm{NH}_{4}\) I in liquid \(\mathrm{NH}_{3}\) (c) Most common "acids' behave as bases in liquid \(\mathrm{H}_{2} \mathrm{SO}_{4}\) (d) \(\mathrm{HClO}_{4}\) is fully ionized in water and is strongly dissociated in pure (glacial) acetic acid; in liquid \(\mathrm{HSO}_{3} \mathrm{F},\) the following reaction occurs: \(\mathrm{KClO}_{4}+\mathrm{HSO}_{3} \mathrm{F}-\mathrm{KSO}_{3} \mathrm{F}+\mathrm{HClO}_{4}\)

Short answer

Different solvent environments influence precipitation, dissolution, and acid-base behaviors by altering solubility and ionization properties.

Step-by-step solution

Understanding the Reaction in Aqueous Solution

In aqueous solution, \(\mathrm{AgNO}_{3}\) and \(\mathrm{KCl}\) react to form \(\mathrm{AgCl}\) as a precipitate because \(\mathrm{AgCl}\) is insoluble in water. \(\mathrm{Ag}^+\) ions from \(\mathrm{AgNO}_{3}\) react with \(\mathrm{Cl}^-\) ions from \(\mathrm{KCl}\) to form \(\mathrm{AgCl}\), which precipitates out of the solution.

Understanding the Reaction in Liquid Ammonia

In liquid \(\mathrm{NH}_{3}\), the solubility principles are different. \(\mathrm{AgCl}\) is actually more soluble than \(\mathrm{KCl}\) in liquid ammonia, leading \(\mathrm{AgCl}\) to dissolve and \(\mathrm{KCl}\) to precipitate when \(\mathrm{AgCl}\) reacts with \(\mathrm{KNO}_{3}\). This supports understanding that solvent properties can significantly change reaction outcomes.

Role of NH4I with Mg in Liquid Ammonia

\(\mathrm{Mg}\) dissolves in a concentrated solution of \(\mathrm{NH}_{4}\mathrm{I}\) in liquid \(\mathrm{NH}_{3}\) because \(\mathrm{NH}_{3}\) acts as a complexing solvent. \(\mathrm{NH}_{3}\) helps stabilize ion pairs like \([\mathrm{Mg}(\mathrm{NH}_{3})_{4}]^{2+}\), making \(\mathrm{Mg}\) soluble.

Understanding Acids as Bases in Liquid H2SO4

In liquid \(\mathrm{H}_{2}\mathrm{SO}_{4}\), most common acids behave as bases because \(\mathrm{H}_{2}\mathrm{SO}_{4}\) is a strong proton donor. This medium donates protons more readily than many acids, compelling them to act as bases by accepting protons instead.

Dissociation and Reaction of HClO4 in Different Solvents

\(\mathrm{HClO}_{4}\) is fully ionized in water because water stabilizes ions. Meanwhile, in (glacial) acetic acid, it experiences strong dissociation due to the differing polarity. In \(\mathrm{HSO}_{3}\mathrm{F}\), the reaction given indicates ion exchange, where \(\mathrm{HClO}_{4}\) interacts with \(\mathrm{KSO}_{3}\mathrm{F}\) forming the corresponding compounds due to differing solvent characteristics.

Key concepts

Solubility in Different Solvents

Solubility is the ability of a substance to dissolve in a solvent. Different solvents can dramatically change the outcome of chemical reactions due to their chemical properties. In aqueous solutions, water serves as the solvent. For example, when silver nitrate (\(\mathrm{AgNO}_3\)) is mixed with potassium chloride (\(\mathrm{KCl}\)), silver chloride (\(\mathrm{AgCl}\)) is precipitated because it is insoluble in water. On the other hand, in liquid ammonia (\(\mathrm{NH}_3\)), the solubility of these compounds changes. Here, silver chloride is more soluble than potassium chloride, thus the reaction leads to the formation of a different precipitate – \(\mathrm{KCl}\). When studying solubility, consider these aspects:

• Solvent Interaction: Different solvents interact differently with solutes, affecting their dissolution or precipitation.
• Polarity: This is a key factor, influencing how well a solute can dissolve. Water is polar, supporting ionic compound dissolution while non-polar solvents, like liquid ammonia, alter outcomes.

Understanding these variables helps predict the behavior of substances in various chemical environments.

Acid-Base Behavior in Non-Aqueous Solvents

The behavior of acids and bases can change significantly outside water. This is due to the unique properties of non-aqueous solvents, like sulfuric acid (\(\mathrm{H}_2\mathrm{SO}_4\)). In this solvent, typical acids may act as bases. This phenomenon occurs because sulfuric acid is a strong proton donor. It readily donates protons, leading other acids to accept them, thus behaving like bases. Key points to understand acid-base behavior in non-aqueous solvents include:

• Proton Donation Ability: Some solvents are capable of donating protons more effectively than others, shifting the typical roles of acids and bases.
• Strength of Acids and Bases: The usual strength hierarchy can reverse, as observed with typical acids acting as bases in \(\mathrm{H}_2\mathrm{SO}_4\).

This understanding aids in predicting reactions and interactions in varied chemical environments.

Ionization in Various Solvents

Ionization is the process where a molecule or compound transforms into ions. Each solvent affects ionization differently depending on polarity and dielectric constant. For example, \(\mathrm{HClO}_4\) (perchloric acid) behaves distinctively in different solvents:

• In Water: \(\mathrm{HClO}_4\) fully ionizes, forming \(\mathrm{H}^+\) and \(\mathrm{ClO}_4^-\). Water stabilizes these ions efficiently.
• In Acetic Acid: Though it is less polar than water, \(\mathrm{HClO}_4\) still dissociates strongly due to acetic acid's partial polarity.
• In \(\mathrm{HSO}_3\mathrm{F}\): Perchloric acid undergoes ion exchange processes, resulting in the formation of new products.

This demonstrates the importance of solvent choice in influencing both ion stability and reaction feasibility.

Precipitation Reactions

Precipitation reactions occur when two soluble salts react in a solution to form one or more insoluble products, known as precipitates. These reactions are determined by solubility rules and depend heavily on the solvent environment:

• Aqueous Solutions: \(\mathrm{AgNO}_3\) reacts with \(\mathrm{KCl}\) in water, forming \(\mathrm{AgCl}\) as it is insoluble. Water does not dissolve \(\mathrm{AgCl}\), causing it to precipitate.
• In Liquid Ammonia: The solubility of compounds like \(\mathrm{AgCl}\) vs. \(\mathrm{KCl}\) switches, leading \(\mathrm{KCl}\) to precipitate instead.

Key rules governing these reactions are influenced by the ionic product and saturation level in the solvent. Recognizing these factors helps anticipate which substances will precipitate in different chemical contexts.

Complex Formation in Liquid Ammonia

Complex formation is a process where metal ions form coordination complexes with other atoms or molecules. Liquid ammonia (\(\mathrm{NH}_3\)) acts as an excellent complexing solvent because of its ability to stabilize metal ions and form new structures. For example, when magnesium metals, which are typically insoluble, interact with \(\mathrm{NH}_4\mathrm{I}\) in liquid ammonia, they form complexes like \([\mathrm{Mg}(\mathrm{NH}_3)_4]^{2+}\). This increases their solubility. Important aspects of complex formation include:

• Coordination Number: The number of \(\mathrm{NH}_3\) molecules attaching to a central ion impacts complex stability and solubility.
• Electronic Structure: Metal ions like \(\mathrm{Mg}^{2+}\) gain enhanced stability in these environments.

Understanding these interactions guides the use of liquid ammonia in synthetic and industrial chemical processes.