Problem 31 - A Complete Resource Book in Chemistry for JEE Main

Question

Which one of the following is not a Lewis acid? (a) \(\mathrm{C}_{2} \mathrm{H}_{4}\) (b) \(\mathrm{FeCl}_{3}\) (c) \(\mathrm{BF}_{3}^{2}\) (d) \(\mathrm{SiF}_{4}\)

Step-by-Step Solution

Answer

The compound that is not a Lewis acid is \\( \mathrm{C}_{2} \mathrm{H}_{4} \\).

Step 1: Understanding Lewis Acids

A Lewis acid is a substance that can accept a pair of electrons from a Lewis base to form a Lewis adduct. They are typically electron-deficient atoms or molecules that can accommodate additional electrons.

Step 2: Evaluating Each Option

Now, we'll assess whether each given compound can act as a Lewis acid: (a) \(\mathrm{C}_{2} \mathrm{H}_{4}\) - Ethylene is a stable molecule with no electron deficiency. (b) \(\mathrm{FeCl}_{3}\) - Iron (III) chloride can accept electron pairs due to the vacant d orbitals on the iron.(c) \(\mathrm{BF}_{3}^{2-}\) - This species contains the borate anion, which typically does not act as a Lewis acid because of its negative charge.(d) \(\mathrm{SiF}_{4}\) - Silicon tetrafluoride can accept electron pairs through silicon's empty d orbitals.

Step 3: Identifying which Compound is Not a Lewis Acid

Given the analysis, \(\mathrm{C}_{2} \mathrm{H}_{4}\) is a stable molecule without empty orbitals or positive charge, and \(\mathrm{BF}_{3}^{2-}\) is unlikely to accept additional electrons due to its negative charge. However, since \(\mathrm{BF}_{3}\) itself is a known Lewis acid, the confusion arises from the notation error, so \(\mathrm{C}_{2} \mathrm{H}_{4}\) is indeed non-acidic by structure.

Key Concepts

Electron Deficiency

Electron deficiency is a core concept in understanding Lewis acids. A molecule or atom is considered electron deficient if it does not have a complete octet, meaning it lacks two electrons to complete its outer shell. This deficiency creates a desire for electrons, making electron-deficient species good candidates for Lewis acids.
One example is boron trifluoride (\(\mathrm{BF}_{3}\)). Boron has only six electrons in its outer shell, falling short of the stable octet. This inherent **electron deficiency** allows \(\mathrm{BF}_{3}\) to act as a Lewis acid by seeking additional electrons from Lewis bases.
It's important to distinguish between molecules that inherently lack electrons, like boron compounds, and those that are stable due to having a full electron count, such as ethylene (\(\mathrm{C}_{2} \mathrm{H}_{4}\)), which does not act as a Lewis acid due to its electron stability.

Lewis Adduct

A Lewis adduct forms when a Lewis acid accepts a pair of electrons from a Lewis base. This electron pair donation results in a stable complex through a dative bond.
In general:

• The **Lewis acid** is the electron pair acceptor.
• The **Lewis base** is the electron pair donor.

A classic example of a Lewis adduct is the reaction between ammonia (\(\mathrm{NH}_{3}\)) and boron trifluoride (\(\mathrm{BF}_{3}\)). Here, **ammonia acts as a Lewis base**, donating its lone pair to boron, thus forming a dative bond and completing the boron octet.
This new compound is referred to as a Lewis adduct, showcasing the transformative capability of electron pair sharing between substances with differing electron needs.

Vacant d Orbitals

The concept of vacant d orbitals is crucial in explaining how some transition metal compounds act as Lewis acids. Transition metals have d orbitals available that can accept electrons from Lewis bases.
Consider the case of iron(III) chloride (\(\mathrm{FeCl}_{3}\)). Iron, as a transition metal, possesses **vacant d orbitals** that allow it to accept electron pairs. This capability enhances its effectiveness as a Lewis acid.
The **vacant d orbitals** enable **transition metals** to exhibit variable coordination numbers and oxidation states. Transition metals' versatility is significant in various chemical reactions where electron acceptance is essential. In \(\mathrm{FeCl}_{3}\), the availability of these orbitals allows interactions and bonding with electron-rich species, modifying reactivity and compound stability.