Question

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

Short answer

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

Step-by-step solution

Identifying a Lewis Acid

A Lewis acid is a compound that can accept a pair of electrons. Typically, atoms that can accept electrons to complete their outer electron shell, like those with vacant p or d orbitals, are good Lewis acids.

Review the Compounds

Examine each compound: (a) \(\mathrm{SnCl}_{4} \) - Tin (IV) chloride has a full octet but can expand its octet due to d-orbitals, making it a potential Lewis acid.(b) \(\mathrm{BeCl}_{2} \) - Beryllium chloride is electron-deficient and can accept electrons, functioning as a Lewis acid.(c) \(\mathrm{BF}_{3} \) - Boron trifluoride lacks a full octet and can accept an electron pair, therefore, a Lewis acid.(d) \(\mathrm{AlCl}_{3} \) - Aluminum chloride is electron-deficient with a vacant p-orbital, acting as a Lewis acid.

Determining the Non-Lewis Acid

Since a Lewis acid must be able to accept electrons, identify the compound that does not fulfill this criterion. \(\mathrm{SnCl}_{4} \) has a full octet and while d-orbital expansion is possible, it is less typical under normal conditions compared to the electron deficiency seen in the other compounds.

Key concepts

Electron Pair Acceptor

A Lewis acid is known for its ability to accept an electron pair. Just as a host welcomes a guest, a Lewis acid invites an electron duo into its midst. It acts almost like a chemical magnet, attracting electrons. This is due to its incomplete electron configuration.
A classic example of an electron pair acceptor is a compound with an atom that doesn't have a full octet. Think of it as needing to fill seats at a dinner table. There are not enough electrons to complete the atom's outer shell, so it accepts more.

• Atoms like Boron in BF_{3} or Aluminum in AlCl_{3} are prime examples. They have space in their p-orbitals waiting to be filled by electrons from other molecules.
• This is a key characteristic that defines them as Lewis acids.
• Accepting an electron pair means forming a coordinate covalent bond, where both electrons come from the donor molecule, unlike typical covalent bonds where each atom provides one electron.

Vacant Orbitals

Vacant orbitals are like empty parking spots in an atom's structure, ready to hold and stabilize incoming electrons. Atoms with these orbitals can act as Lewis acids, given their availability to accept electrons from other molecules.
It's all about electron deficiency. This doesn’t mean the atom is incomplete, but it has room to spare at the top!

• In the periodic table, atoms in groups 13, like Boron (B) and Aluminum (Al), naturally have three electrons in their outer shell, but they aim for a more stable configuration with eight.
• So, compounds such as BF_{3} and AlCl_{3} come into play, offering to fill those vacant spots with electron donors.
• SnCl_{4}, on the other hand, typically has its octet filled, leading it to not act as strongly as a Lewis acid unless external conditions force d-orbital expansion.

Boron's situation is a classic case where the absence of electrons in its p-orbital urges it to accept electron pairs, making it actively seek electron donors.

Compound Analysis

When analyzing a compound to determine if it's a Lewis acid, the focus is on its electron configuration and shape. It's like detective work, investigating the electron status and willingness to accept more.
Here’s how you analyze:

• Begin by looking at the central atom of the compound. Does it have less than eight electrons? If so, it’s primed to be a Lewis acid.
• Consider the plant-like Beryllium in BeCl_{2} . It naturally seeks stability by filling its electron deficiency, behaving as a Lewis acid by accepting electron pairs.
• Check for exceptions. Even though some compounds like SnCl_{4} have a full octet, their ability to expand their coordination number via d-orbitals, makes them potential Lewis acids but not under standard conditions.

Hence, understanding the electron configuration and the potential for bond formation provides insight into whether a compound acts as a Lewis acid, showcasing its chemical flexibility.