Problem 29 - A Complete Resource Book in Chemistry for JEE Main

Question

Which of the following is the strongest Lewis acid? (a) \(\mathrm{BBr}_{3}\) (b) \(\mathrm{BCl}_{3}\) (c) \(\mathrm{BI}_{3}\) (d) \(\mathrm{BF}_{3}\)

Step-by-Step Solution

Answer

The strongest Lewis acid is \( \mathrm{BI}_3 \).

Step 1: Understanding Lewis Acids

A Lewis acid is a chemical species that can accept an electron pair. In this context, the strength of a Lewis acid is determined by its ability to accept electrons.

Step 2: Identifying Electron-Withdrawing Tendencies

In boron halides like \ \( \mathrm{BX}_3 \ \), the strongest Lewis acids are those in which the boron atom is least capable of stabilizing its negative charge after accepting an electron pair. This generally occurs when the terminal halogen atoms increase the electron deficiency of the boron atom through inductive or resonance effects.

Step 3: Analyzing Electron Withdrawal

The halogen's electronegativity and the ability to form back-bonding will influence the Lewis acidity. \ \( \mathrm{BF}_3 \ \) has strong \ \( p_\pi \rightarrow p_\pi \) back-bonding which stabilizes boron's electron deficiency, making it a weaker Lewis acid compared to others where such back-bonding is less.

Step 4: Comparing Electronegativity and Back-Bonding

Among the options, \ \( \mathrm{BBr}_3 \ \), \ \( \mathrm{BCl}_3 \ \), \ \( \mathrm{BI}_3 \ \), and \ \( \mathrm{BF}_3 \ \): \ \( \mathrm{BF}_3 \ \) has the highest electronegativity, but significant \ \( p_\pi \rightarrow p_\pi \) back-bonding reducing its acidity. \ \( \mathrm{BI}_3 \ \) has the least effective electronegative withdrawal effects and almost negligible back-bonding, making the electron deficiency on boron the greatest.

Step 5: Conclusion

Since \ \( \mathrm{BI}_3 \ \) lacks the stabilizing back-bonding and strong electronegative electron-withdrawing effects, boron remains a strong candidate to accept electron pairs, making it the strongest Lewis acid in the group.

Key Concepts

Electron-Withdrawing Effects

In chemistry, electron-withdrawing effects play a crucial role in determining the properties of a molecule, especially its acidity. When a molecule "withdraws" electrons, it essentially pulls electron density away from a central atom, such as boron in a boron halide compound (\( \mathrm{BX}_3 \)). This results in an increased positive charge or electron deficiency around that central atom, making it more eager, or acidic, to accept electrons from a donor, like in a Lewis acid-base reaction.

Now, understanding electron-withdrawing effects requires us to consider both inductive and resonance influences:

• Inductive effects: This occurs when electronegative atoms, such as fluorine, chlorine, or bromine, pull electron density through sigma bonds. It is mostly a straightforward, distance-dependent effect.
• Resonance effects: Unlike inductive, resonance can involve the delocalization of electrons across bonds, allowing for more complex stabilization, such as back-bonding.

In the context of Lewis acids like \( \mathrm{BF}_3 \) and \( \mathrm{BI}_3 \), although fluorine and iodine both impact boron's electron density, the negligible resonance through iodine results in much less electron density stabilization compared to fluorine. Thus, \( \mathrm{BI}_3 \) becomes a stronger Lewis acid as its electron deficiency on boron is not countered by significant electron-stabilizing effects.

Electronegativity

Electronegativity is a chemical property that describes an atom's ability to attract and hold on to electrons. In the context of Lewis acids, electronegativity directly impacts how electron-deficient a central atom, like boron, can become when bonded to halogens. This, in turn, affects the strength of the acid.

When we look at boron halides such as \( \mathrm{BX}_3 \), each halogen atom's electronegativity heavily influences the electron-withdrawing ability:

• Fluorine (\( \mathrm{BF}_3 \)): It is the most electronegative and effectively pulls electron density away. However, its strong back-bonding abilities counterbalance this effect, making \( \mathrm{BF}_3 \) a weaker Lewis acid.
• Iodine (\( \mathrm{BI}_3 \)): Having the least electronegativity among halogens discussed, it withdraws electrons less effectively but lacks significant back-bonding capabilities, reinforcing boron’s electron deficiency and hence, making \( \mathrm{BI}_3 \) the strongest Lewis acid.

Therefore, the interplay between electronegativity and back-bonding is pivotal; a high electronegativity combined with minimal compensating back-bonding enhances a compound's Lewis acid character.

Back-Bonding

Back-bonding is a fascinating concept that adds depth to our understanding of molecular structure and reactivity. It occurs when electrons in a filled p orbital of a more electronegative atom, like fluorine in \( \mathrm{BF}_3 \), are "donated back" into an empty orbital of the less electronegative central atom, such as boron. This can effectively stabilize electron-deficient centers by decreasing susceptibility to additional electron attraction.

The implications of back-bonding in Lewis acids can be summarized as:

• Back-bonding stabilizes the electron deficiency at the central atom, reducing its ability to behave as a strong Lewis acid.
• Fluorine in \( \mathrm{BF}_3 \), due to its high electronegativity and effective p-p orbital overlap, exhibits significant back-bonding, which stabilizes boron's electron-deficient character.
• In contrast, \( \mathrm{BI}_3 \), with its larger iodine atoms that feature less effective orbital overlap with boron, exhibits minimal back-bonding, thereby leaving boron's electron deficiency mostly unaddressed.

In summary, when back-bonding is strong, like in \( \mathrm{BF}_3 \), the Lewis acidity decreases. But in cases like \( \mathrm{BI}_3 \), where back-bonding is weak or non-existent, the molecule exhibits stronger acidic behavior.