(a) Suggest a solvent in which aniline acts as strong base. (b) Write equation for the auto ionisation of (i) \(\mathrm{HCOOH}\), (ii) \(\mathrm{NH}_{3}\). (c) \(\left[\mathrm{Al}\left(\mathrm{H}_{2} \mathrm{O}\right)_{6}\right]^{3+}\) is acid or base and write its conjugate partner and reaction. (d) Write the order of acidic nature of \(\mathrm{HCl}, \mathrm{HCOOH}\) and \(\mathrm{CH}_{3} \mathrm{COOH}\) in (i) \(\mathrm{H}_{2} \mathrm{O}\), (ii) liq. \(\mathrm{NH}_{3}\).

Understanding acid-base reactions is fundamental in physical chemistry, especially for students preparing for competitive exams like IIT-JEE. In these reactions, acids donate protons (H⁺ ions), and bases accept them. The Bronsted-Lowry theory is often applied, where an acid-base reaction is seen as a proton transfer process.

For example, in the case of aniline acting as a base, we look for a solvent where it can readily accept a proton. Aniline, or C₆H₅NH₂, is a weak base in water, but in a non-polar solvent like toluene, it can behave as a strong base due to reduced competition for protonation. This is essential information for students to determine the strength of an acid or base depending on the solvent.

In the context of IIT-JEE, recognizing how solvents influence acid-base strength aids students in predicting reaction outcomes and creating stronger analytical skills for problem-solving.

Proton Transfer and Solvent Effects

To contextualize this with an example for deeper understanding: in water, hydrochloric acid (HCl) releases a proton to form hydronium ions (H₃O⁺), demonstrating its strong acidic nature. But in ammonia, which is a weaker base than water, acetic acid (CH₃COOH) releases protons more readily than HCl, inverting the acidic order observed in water.

Autoionization, sometimes referred to as self-ionization, is a process where solvent molecules react with each other to produce ions. This concept is pivotal for students to grasp as it lays the foundation for understanding the solvent's role in acid-base reactions.

In autoionization, a solvent acts as both an acid and a base; one molecule donates a proton, while another accepts it. For instance, the autoionization of formic acid (HCOOH) leads to the formation of hydronium ions (H₃O⁺) and formate ions (HCOO^-). Similarly, with ammonia (NH₃), the process generates ammonium ions (NH₄⁺) and amide ions (NH₂^-).

For IIT-JEE aspirants, understanding these processes is crucial in predicting the extent of ionization, pH level of solutions, and even the stability of ions formed from various solvents.

Comparative Strengths of Solvents

The autoionization in different solvents can vary significantly. Water's high dielectric constant makes it a generally strong solvent for ionization. In contrast, autonization is much less pronounced in ammonia because it's a weaker solvent. Recognizing these differences allows students to better appreciate the nuances in various chemical environments and reaction conditions.

A more generalized concept in acid-base chemistry, beyond the Bronsted-Lowry definition, is that of Lewis acids and bases. Lewis acids are electron pair acceptors, and Lewis bases are electron pair donors. This concept is particularly helpful when dealing with non-proton reactions and is a valuable addition to the IIT-JEE chemistry syllabus.

An example from the textbook exercise is the [Al(H₂O)₆]³⁺ complex. It operates as a Lewis acid because it can accept a pair of electrons from a water molecule, leading to the formation of its conjugate base, [Al(H₂O)₅(OH)]²⁺. Here, it's important for students to understand that acidity and basicity also involve the ability to accept or donate electron pairs, not just protons.

Lewis theory extends the concept of acids and bases to a broader range of reactions, which is essential for students facing the diverse array of chemistry problems in IIT-JEE.

Electron Pair Interaction

The interaction of electron pairs is at the heart of Lewis theory. In metal complexes, such as the hydrated aluminum ion, electron pair donation from water leads to coordination bond formation. These interactions are key to the structure and reactivity of compounds and underscore the importance of Lewis acid-base theory in predicting and explaining chemical behavior.