Question

Which of the following is the strongest Lewis acid?
(A) $B{I_3}$
(B) $BB{r_3}$
(C) $BC{l_3}$
(D) $B{F_3}$

Answer 1

Hint: Lewis acids are the compounds that have empty orbitals and so that they can accept electrons from Lewis bases. The strength of the B-X bond (X is the halogen) affects the strength of the Lewis acid.

Complete step by step solution:
- In $B{F_3}$ (Boron Trifluoride), boron atoms have half filled 2p orbital. A Fluorine atom has five electrons in its seven electrons in its valence orbital. So, fluorine donates two electrons from its 2p orbital to the vacant 2p orbitals of boron. As a result $p\pi - p\pi $ bond is formed. This bond is stronger and decreases the electron deficiency of boron. So, as a result the Lewis acid character of $B{F_3}$ diminishes.
- $BC{l_3}$ (Boron trichloride) is a poisonous gas. In $BC{l_3}$, the B-Cl bonds are less stronger than B-F bonds and as a result Lewis acid character of $BC{l_3}$ is slightly higher than $B{F_3}$.
- $BB{r_3}$ (Boron tribromide) is found in liquid form. B-Br bonds are weaker than B-Cl bonds in boron trihalides and as a result Lewis acid character of $BB{r_3}$ is higher than $BC{l_3}$.
- $B{I_3}$ (Boron triiodide) is found in liquid state. B-I bonds are the weakest amongst all the halogens. So, as a result $B{I_3}$ is the strongest Lewis acid because electron deficiency of boron is highest in $B{I_3}$ amongst other boron trihalides.
Therefore, we can say that the strongest lewis acid is $B{I_3}$.

So, correct answer is (A).

Note: The order of the electron deficiency of Boron atom in its trihalides is:
$B{F_3}$ < $BC{l_3}$ < $BB{r_3}$ < $B{I_3}$
The order of strength of Lewis acid amongst boron trihalides is:
$B{F_3}$ < $BC{l_3}$ < $BB{r_3}$ < $B{I_3}$