Which of the species $\ce{BF3}$ and $\ce{AlCl3}$ is more acidic, both are electron deficient.

In case of $\ce{BF3}$ one can argue that there is extensive p-π-p-π back bonding and as a result the electron deficiency is less.
That means that the positive charge on the overall molecule is less.

In case of $\ce{AlCl3}$ one can argue there is p-π-d-π back bonding and as a result the electron deficiency is more in case of aluminiuim.
Therefore the overall positive charge on the molecule is more, meaning it is more acidic.

So $\ce{AlCl3}$ is more acidic that $\ce{BF3}$ is conclusion of the above factors. Is this plausible?

I also faced the counter logic that $\ce{BF3}$ being a small compound tends to react more readily and therefore is a strong acid.

  • $\begingroup$ Back bonding would dominate $\endgroup$ Jan 24 '18 at 8:15
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    $\begingroup$ D-orbitals in aluminium or chlorine? No. $\endgroup$ Jan 24 '18 at 9:05
  • $\begingroup$ @Martin-マーチン why can you explain? $\endgroup$
    – Pole_Star
    Jan 24 '18 at 9:38
  • $\begingroup$ [AlCl6]3- exists what about that $\endgroup$
    – Pole_Star
    Jan 24 '18 at 9:40
  • 4
    $\begingroup$ @starunique2016 The molecular ion $\ce{[AlCl6]^3-}$ is predominantly bound by ionic interactions. The covalent part is best described with multi-centre bonds. In valence bond theory you have to invoke resonance. In none of these have the d-orbitals a role beyond polarisation effects. $\endgroup$ Jan 24 '18 at 10:18

Actually, there are three reasons why BF3 is more acidic than AlCl3.

  1. BF3 is smaller in size and can easily attract the incoming pair of electrons.
  2. Fluorine is more electronegative than chlorine. So boron has less ionisation potential as compare to aluminium thus forms anion easily.
  3. Boron is more electronegative than aluminium, therefore boron has more tendency to attract electrons. It means that BF3 behaves more like Lewis acid than AlCl3.

Scale of the strength of some Lewis acids

This photo has been taken from this article.

  • 1
    $\begingroup$ It would be preferable to include a human readable citation instead of just a link to a paper. $\endgroup$ Apr 4 at 17:53
  • $\begingroup$ I have inserted a photo for better understanding the matter. $\endgroup$
    – Exeplone
    Apr 4 at 20:25
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    $\begingroup$ The image certainly is a good addition. I kindly ask you to include a citation for it though. One that you would use in other articles. Thank you. $\endgroup$ Apr 5 at 2:39
  • $\begingroup$ really well explained, even i think that this is the correct answer...i was trying to convey the same in my answer $\endgroup$ Apr 5 at 17:21
  • $\begingroup$ @Exeplone According to the attached photo, Lewis acid strength order is: $\ce{BF3 < AlCl3 < BCl3}$ $\endgroup$
    – Apurvium
    Jul 6 at 4:32

Applying Pearson's HSAB concept-

Toward hard Lewis bases such as ethyl acetate, the Lewis acidities of the halides of group 13 elements would decrease as the softness of the acceptor element increases: $\ce{BX3 > AlX3 > GaX3 > InX3}$

Toward soft Lewis bases such as dimethyl sulphide, the Lewis acidities of the halides of group 13 elements would increase as the softness of the acceptor element increases: $\ce{BX3 < AlX3 < GaX3 < InX3}$

But for comparison of $\ce{BF3}$ and $\ce{AlCl3}$, experimental data is required as provided by @Exeplone


Lewis acids are those which can accept electrons easily. Now, in BF3 molecule, B is small in size and thus can easily attract the incoming pair of electrons. Also, the presence of three fluorine atoms makes the boron acquire a partial positive charge, due to which boron easily accept electrons acting as Lewis acid. But, in the case of AlCl3, the size of Al is large so incoming electrons are less attracted.

(Since Boron is more electronegative than Aluminium, it can attract electrons more easily and since Lewis acids are electron acceptors, BF3 is a stronger Lewis acid than AlCl3).

Thus BF3 is better Lewis acid than AlCl3.

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    $\begingroup$ Unfortunately that is not correct, and not thorough enough for such an involved problem. $\endgroup$ Apr 4 at 10:20

$\ce{BF3}$ is less acidic than $\ce{AlCl3}$ as there much stronger pi-backbonding between the Boron and Fluorine atoms while there is no backbonding in AlCl3 as well. The tendency of boron to accept electrons from the base, is way lesser than that of Al.

And, There is no backbonding in AlCl3, check out this link: https://byjus.com/questions/why-alcl3-does-not-show-back-bonding/

  • $\begingroup$ No. There can be backbonding in Al-Cl systems. $\endgroup$ Jan 25 '18 at 2:38
  • $\begingroup$ @KartikAnand That is one of the problems of this answer. You should have said that there is backbonding in $\ce{AlCl3}$, because there is. Also, nobody ever said that there is noc backbonding in $\ce{BF3}$. The whole reasoning with "octet completion" as a drop in for acidity is hand-wavy at best. $\endgroup$ Apr 4 at 8:41
  • $\begingroup$ isn't BF3 a stronger Lewis acid in comparison to AlCl3, considering that it is more electronegative and thus attracts electrons more easily? $\endgroup$ Apr 4 at 10:18
  • $\begingroup$ @Martin-マーチン I edited my answer. Let me know if that still doesn't sense. $\endgroup$ Apr 4 at 14:37
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    $\begingroup$ That source page you are referring to probably has so much tracking on it that they decided it's not going to be available in the EU. In any case, judging just from the title, the content will be completely wrong. $\endgroup$ Apr 4 at 17:47

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