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I asked this question to my teacher a few days ago. In a reaction such as

$$\ce{NH3 + H2O -> NH4+ + OH-},$$

$\ce{NH3}$ is a Lewis base, but is $\ce{H2O}$ a Lewis acid?

He said that it isn't. Now I fail to see the reason why. Can someone explain?

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  • $\begingroup$ I would think that H2O can act as a Lewis acid. The bound hydrogens are electron-deficient and can act as lone pair acceptors. $\endgroup$ – Tan Yong Boon Sep 23 '17 at 11:57
  • $\begingroup$ Perhaps, your teacher thought that it could not be a Lewis acid because all the hydrogens had complete octets. $\endgroup$ – Tan Yong Boon Sep 23 '17 at 11:59
  • $\begingroup$ So the bound hydrogen ions that get separated then bond with the NH3 ? Does that make the hydrogen ion the lewis acid? Since it's the one accepting the electron pair $\endgroup$ – Dahen Sep 23 '17 at 11:59
  • $\begingroup$ Sorry, I meant complete duets. $\endgroup$ – Tan Yong Boon Sep 23 '17 at 12:05
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    $\begingroup$ Not true... The hydrogen ion may not even form in the first place. A hydrogen bond can be formed between the ammonia and the water molecule. Then there is a transition state, with the hydrogen having partial covalent bonds with the nitrogen atom and the oxygen atom. Finally, the bond between N and H forms fully and the bond between O and H breaks fully. $\endgroup$ – Tan Yong Boon Sep 23 '17 at 12:08
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If you agree that ammonia is a Lewis base (electron pair donor) and that this is an Lewis acid-base reaction, (donor-acceptor interaction) then water is a Lewis acid (electron pair acceptor), by definition and the fact that there is a reaction. Now, identifying the mechanism of acidity or the acidic site is more complicated, but you can't argue that it is acidic, by definition.

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  • $\begingroup$ I asked my teacher again and he said that it was indeed acidic. Thanks! $\endgroup$ – Dahen Oct 5 '17 at 22:00
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Strictly speaking, a Lewis acid is something that forms a bond by accepting an electron pair form another molecule (Lewis base). Water as such does not do that, rather it is a hydrogen ion from the water that does so. When ammonia acts as a proton acceptor in water, the Lewis theory calls it a displacement reaction where the acid, $\ce{H+}$, is initially combined with one base ($\ce{OH-}$) and ends up combined with another base ($\ce{NH3}$).

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  • $\begingroup$ Yes i thought so too, but the issue is that same logic can be applied to all the other acids like HCL, HF ect, and my teacher called those lewis acids. Though I did ask him the other day and he said that water was a lewis acid in that reaction all technical details aside ( or at least it'll be for our book) $\endgroup$ – Dahen Oct 5 '17 at 22:25
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    $\begingroup$ Not even HCl. No matter how strong, protic acids are sources of the actual acid, hydrogen ion, in the Lewis theory. $\endgroup$ – Oscar Lanzi Oct 5 '17 at 23:30
  • $\begingroup$ Yup he said that as well. It's just weird that he didnt say it from the start tbh $\endgroup$ – Dahen Oct 6 '17 at 8:31
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I have finally found a satisfactory answer to your question after much research. Jensen (1978) provides an excellent review of the Lewis acid-base theory and p. 4 of the article provides the answer that you seek$^1$. I shall quote from that page the relevant insights:

So rapid and striking were many of these neutralizations that Lewis went on to propose that that criterion 1 (i.e., rapid kinetics) was the salient feature of acid-base behavior, suggesting further that a fundamental subdivision of acids and bases be made on this basis...

Lewis classified those acids and bases which underwent acid-base reactions which showed "essentially zero activation energy" as primary, while those which had measurable activation energies were termed secondary. He further broke down this secondary class into two types (ref 1, p. 4):

The first of these involved species, such as $\ce {CO2}$, in which the slow kinetic behavior was apparently due to the necessity of the species undergoing some sort of internal activation before its primary acid or base properties became apparent.

The second class involved those species in which the finite activation energy was due to the breaking of one or more auxiliary bonds upon neutralization, causing the initial $\ce {AB}$ complex to dissociate into several smaller fragments. Hence, Bronsted acids like $\ce {HCl}$ and $\ce {HNO3}$ were still acids, though now of the secondary variety, and their neutralizations could be thought of as initially resulting in an unstable hydrogen-bridged adduct which then underwent further decomposition.

To clarify, the Lewis acid-base reaction defined by Lewis is as such:

$\ce {A + :B -> AB}$

Back to your question... Essentially, $\ce {H2O}$ can be seen as the second class of secondary acids proposed by Lewis. During a reaction with Lewis base $\ce {:B}$, there is essentially some sort of complex formed, that looks like this $\ce {[B -- H -- OH]}$. The dotted lines indicate partial covalent bonds. This complex can be seen as sort of a "transition state". However, note that this was not in the original formulation of Lewis. The image below shows the reaction between pyridine and $\ce {HCl}$ viewed from Lewis' perspective. Lewis called this complex an "unstable adduct".

enter image description here

Consolidation

Based on the above, we can say that acids of the type $\ce {HA}$ (where $\ce {A}$ is an electronegative atom or group of atoms) are secondary Lewis acids, which participate in acid-base reactions with simultaneous bond breaking of auxiliary bonds. This is because the idea of complexation to form an adduct is still present.

Reference

  1. Jensen, W. B. The Lewis acid-base definitions: a status report. Chem. Rev., 1978, 78(1), 1-22. doi:10.1021/cr60311a002
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My friend, here is how i understand it: H2O is a neutral medium, does not like to be acid or base, but if we mix acid with it, then the water will try to fight back by turning to Lewis base. Same thing goes with adding base to H2O, H2o will fight back by turning acid to naturalize the solution. So yes, water can be both Lewis acid or base.

I hope that makes sense

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    $\begingroup$ It would be better if you used the definitions for Lewis acidity in your answer... $\endgroup$ – Zhe Oct 5 '17 at 21:54

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