Question

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How do acidic character vary among the trihalides of Boron? Give reason in support of your answer.
Kinda confused, give any 5 points in support of your answer. :D
Wrong answers will be reported.
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Answer 1

Answer:

The Lewis acidity of BF3, BF2Cl, BFCl2, and BCl3 in acid−base orbital interactions has been studied. We have derived the unoccupied reactive orbitals that show the maximum localization on the boron pπ atomic orbital overlapping with the lone-pair orbital of an electron donor and have evaluated the electrophilicity of the boron center in these compounds. The Lewis acidity of boron is shown to be controlled by two factors:  localizability of the unoccupied reactive orbital on the boron pπ atomic orbital and the polarizability of the boron center. The former has been shown to be similar in magnitude in these boron halide compounds. Contrary to common belief, the conjugation between the boron atom and the attached halogen atoms is not necessarily stronger in BF3 relative to others. The trend observed in experiments and in theoretical calculations for BF3, BF2Cl, BFCl2, and BCl3 is interpreted in terms of these factors.

Explanation:

Answer 2

Answer:

Explanation:

All three lighter boron trihalides, BX3 (X = F, Cl, Br), form stable adducts with common Lewis bases. Their relative Lewis acidities can be evaluated in terms of the relative exothermicities of the adduct-forming reaction. Such measurements have revealed the following sequence for the Lewis acidity: BF3 < BCl3 < BBr3 (in other words, BBr3 is the strongest Lewis acid).

This trend is commonly attributed to the degree of π-bonding in the planar boron trihalide that would be lost upon pyramidalization (the conversion of the trigonal planar geometry to a tetrahedral one) of the BX3 molecule, which follows this trend: BF3 > BCl3 > BBr3 (that is, BBr3 is the most easily pyramidalized). The criteria for evaluating the relative strength of π-bonding are not clear, however. One suggestion is that the F atom is small compared to the larger Cl and Br atoms, and the lone pair electron in the 2pzorbital of F is readily and easily donated, and overlaps with the empty 2pz orbital of boron. As a result, the [latex]\pi[/latex] donation of F is greater than that of Cl or Br. In an alternative explanation, the low Lewis acidity for BF3 is attributed to the relative weakness of the bond in the adducts F3B-L.