Give reasons for the following: (a) Acetamide is basic in water but acidic in liquid \(\mathrm{NH}_{3}\). (b) \(\mathrm{SbF}_{6}\) behaves as an acid in \(\mathrm{HF}\). (c) Alkali metals dissolve in liquid \(\mathrm{NH}_{3}\) to give reducing solutions. (d) Solution of \(\mathrm{KNH}_{2}\) in liquid \(\mathrm{NH}_{3}\) gives pink colour to phenolphthalein. (e) Solution of \(\mathrm{Hg}\left(\mathrm{ClO}_{A}\right)_{2}\) in \(\mathrm{HgCl}_{2}\) is acidic.

Multiple factors determine the acidity or basicity of a compound, such as the solvent it's dissolved in, the structure of the compound, and the presence of functional groups. In aqueous solutions, acid-base behavior is often discussed in terms of the Brønsted-Lowry theory, which defines acids as proton donors and bases as proton acceptors. For instance, acetamide is basic in water as it can accept protons, forming the ammonium ion.
However, when the same compound is placed in a different solvent, such as liquid ammonia, its behavior can change. In liquid ammonia, acetamide acts as an acid.

• Basic Environment: Acetamide accepts a proton and increases pH.
• Acidic Environment: Acetamide donates a proton and decreases pH.

Understanding these behaviors helps to predict reactions and properties of chemical substances in various environments.

Lewis acid-base theory broadens the definition of acids and bases beyond proton transfer. Lewis acids are electron pair acceptors, while Lewis bases are electron pair donors. This concept explains reactions without an obvious proton transfer, such as the reaction of SbF6 with HF. Here, the fluoride ion from HF (the Lewis base) donates an electron pair to SbF6 (the Lewis acid), forming a stable complex ion, SbF7−.

• Lewis Acid: Electron pair acceptor (e.g., SbF6)
• Lewis Base: Electron pair donor (e.g., fluoride ion from HF)

These concepts are crucial for understanding a wide variety of chemical reactions, including coordination chemistry and catalysis.

Solvation involves interactions between a solute and solvent that lead to a stable solution. Liquid ammonia is a non-aqueous solvent that exhibits interesting solvation properties, particularly with metals. When alkali metals dissolve in liquid ammonia, the metal atoms donate electrons to the solvent, forming metal cations and solvated electrons.
This results in a blue to bronze-colored solution that can act as a powerful reducing agent, capable of donating electrons to other substances. The reaction showcases the distinct solvation effects in liquid ammonia due to its capacity to solvate and stabilize electrons. It is these solylation effects that also impart the reducing nature to the solutions.

• Alkali Metals: Donate electrons to liquid ammonia.
• Solvated Electrons: Impart color and reducing properties to the solution.

Comprehending these solvation effects is fundamental in understanding reactions in non-aqueous solvents.

pH indicators are substances that change color depending on the acidity or basicity of the environment they are in. They are typically weak acids or bases that exist in different color forms depending on the pH. For example, phenolphthalein is commonly used as an indicator which is colorless in acidic conditions but turns pink in basic solutions.
The solution of KNH2 in liquid ammonia changes the color of phenolphthalein to pink, indicating a basic medium. This occurs because KNH2 acts as a strong base in the presence of liquid ammonia, altering the equilibrium of the phenolphthalein molecule towards its basic, colored form.

• Acidic Solution: Phenolphthalein is colorless.
• Basic Solution: Phenolphthalein turns pink.

Identifying these color changes allows for a visual understanding of the pH level in a solution and is a vital concept in acid-base chemistry.