Problem 40 - A Complete Resource Book in Chemistry for JEE Main

Question

The compound that is not a Lewis acid is (a) \(\mathrm{SnCl}_{4}\) (b) \(\mathrm{BeCl}_{2}\) (c) \(\mathrm{BF}_{3}\) (d) \(\mathrm{AlCl}_{3}\)

Step-by-Step Solution

Answer

The compound that is not a Lewis acid is \(\mathrm{SnCl}_{4}\).

Step 1: Understanding Lewis Acids

A Lewis acid is defined as a species that can accept a pair of electrons. This typically means that the compound has an incomplete octet or can form additional bonds by accepting electron pairs.

Step 2: Review Each Compound's Structure

Evaluate each compound to determine if it can accept electron pairs. - **\(\mathrm{SnCl}_{4}\)**: Tin (IV) chloride has a central tin atom bonded to four chlorine atoms. Since tin can expand its octet, it has the capacity to accept electron pairs.- **\(\mathrm{BeCl}_{2}\)**: Beryllium chloride has a beryllium atom bonded to only two chlorine atoms. Beryllium is an electron-deficient atom with only four electrons around it, making it a Lewis acid capable of accepting electron pairs.- **\(\mathrm{BF}_{3}\)**: Boron trifluoride has a boron atom with only six electrons, making it electron-deficient and capable of accepting electron pairs.- **\(\mathrm{AlCl}_{3}\)**: Aluminum chloride can exist as a dimer, meaning it still has a central aluminum atom that can accept electron pairs, making it a Lewis acid.

Step 3: Identifying the Non-Lewis Acid

All the compounds except \(\mathrm{SnCl}_{4}\) are classical Lewis acids due to their electron-deficient nature. Conversely, \(\mathrm{SnCl}_{4}\) is not considered a traditional Lewis acid as it does not readily accept additional electron pairs in typical conditions owing to its electron configuration.

Key Concepts

Electron-deficient Compounds

Electron-deficient compounds are unique because they have fewer electrons than needed to achieve a stable electronic structure. This concept primarily involves elements that do not satisfy the octet rule. Such compounds are often short of fulfilling their electron demand and can act as Lewis acids. They seek additional electrons to stabilize their structures by accepting electron pairs from electron-rich species.

Examples of electron-deficient compounds include:

• **Boron trifluoride (\(\mathrm{BF}_{3}\))**: It has only six electrons around the central boron atom, which is why it is eager to accept electron pairs to complete its octet.
• **Beryllium chloride (\(\mathrm{BeCl}_{2}\))**: Beryllium also forms electron-deficient compounds as it bonds with only two chlorine atoms, having just four electrons in its valence shell, leaving it short of the eight-electron rule.

Electron-deficient compounds are interesting in chemistry, as their bonding and reactions can often defy more commonly understood rules. They illustrate important concepts such as electron sharing and transfer in chemical bonding.

Chemical Bonding

Chemical bonding explains how atoms combine to form compounds. It involves interactions that hold these atoms together in a stable configuration. One of the most essential forms of chemical bonding is covalent bonding, where atoms share electron pairs to achieve stability. However, not all compounds strictly adhere to these classic bonding scenarios.

In the case of **Lewis acids**, we observe situations where one atom can adopt electron pairs, forming coordinate covalent bonds. For instance, in compounds like \(\mathrm{BF}_{3}\), boron needs more electrons to complete its orbit, making it capable of accepting electrons from donors. This ability to form bonds by electron pair acceptance is a hallmark of Lewis acids, which diversifies the typical binary perspective of acids and bases.

Understanding these interactions helps clarify more complex bonding scenarios encountered in advanced chemistry.

Acid-base Theory

Acid-base theory provides a foundation for understanding how acids and bases react. Lewis acid-base theory broadens this further by focusing not on protons like the classical definitions but on electronic interactions. It defines a **Lewis acid** as a species capable of accepting an electron pair, contrasting with **Lewis bases**, which donate electron pairs.

This theory is essential because it explains reactions in terms of electron pair exchanges rather than focusing strictly on hydrogen ions. It offers a more generalized view of acid-base interactions, applicable to gas-phase, non-aqueous, and solid-state scenarios.

• In Lewis terms, electron-deficient molecules like \(\mathrm{BeCl}_{2}\) or \(\mathrm{BF}_{3}\) readily accept electron pairs, functioning as acids.
• This concept can resolve compounds like \(\mathrm{SnCl}_{4}\), which is less typically recognized as a Lewis acid due to its propensity to remain electron-neutral under usual conditions.

Understanding acid-base reactions through this lens influences how chemists approach reaction mechanisms and the synthesis of new compounds.