Question

Why $BBr_{3}$ is a stronger Lewis acid as compared to $BF_{3}$ though fluorine is more electronegative than bromine.

Answer 1

Boron tribromide is the STRONGEST Lewis acid known…and this is the experimental fact despite the reduced electronegativity of bromine as compared to fluorine, and chlorine. So why so? For boron trifluoride, and the trichloride, donation from the lone pairs of the halogen to the empty p-orbital on boron is fairly effective. In the case of bromine, the lone pairs are larger and more diffuse and LESS capable of this donation….and as a result the boron centre is more electropositive, and Lewis-acidic…. Of course, we might also finger boron triiodide as a strong Lewis acid by this same argument, however, (and I have NEVER used this reagent!), it tends not to figure in these scenarios perhaps because of the expense, and its particularly high molar mass.

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Answer 2

   Boron tribromide $BBr_{3}$ is a stronger Lewis acid than Boron trifluoride $BF_{3}$ due to the reason mentioned below.

    There is a certain pi-bonding which involves lone pairs of halogen and also empty 2p orbitals present in the Boron halides. A Boron halide receives an electron from the donar molecule breaking the strong bond in the  $BF_{3}$.

     Because of this reason, Boron trifluoride $BF_{3}$ loses the strength and becomes a weaker lewis acid as compared to the Boron tribromide $BBr_{3}$.