I am having trouble with this question:

Identify the Lewis acid and Lewis base in each of the following reactions:

$\ce{SO2(g) + H2O(l) = H2SO3(aq)}$

What I tried:

A Lewis acid is an electron pair acceptor.

A Lewis base is an electron pair donor.

The other problems of this type involved cations or anions, so I was able to identify the cation as the Lewis acid and the anion as the Lewis base.

However, this question does not involve ions. So I tried to look at the lone pairs on the molecules. Since they both have lone pairs, I figured since $\ce{SO2}$ has more lone pairs, it would act as the Lewis base. Also in the final molecule, the hydrogen atoms would be bonded to two separate oxygen atoms, so I thought of them as migrating to the lone pairs on the $\ce{SO2}$ oxygen atoms.

However, the textbook says that $\ce{SO2}$ is the Lewis acid and $\ce{H2O}$ is the Lewis base.

Why would this be the case?

  • $\begingroup$ Clearly it is $\ce{H2O}$ that donates an electron pair to from the bond? So it is the Lewis base. $\endgroup$
    – Gert
    Commented Dec 9, 2017 at 21:25
  • $\begingroup$ Can you explain why it must be $\ce{H2O}$ that donates the electron pair(s) and not $\ce{SO2}$? $\endgroup$ Commented Dec 9, 2017 at 22:07
  • $\begingroup$ I'll try and answer that tomorrow. $\endgroup$
    – Gert
    Commented Dec 9, 2017 at 22:48

3 Answers 3


Molecules that contain polar multiple bonds can function as Lewis acid because the central atom is electron deficient with a vacant orbital that can accept an electron pair.

When $\ce{SO2}$ dissolves in water , it forms the weak diprotic acid $\ce{H2SO3}$ (sulfurous acid) from Lewis acid/Lewis base reaction:

$$\ce{O\bond{=}S\bond{=}O + H2O <=> H2SO3}$$

enter image description here

The $\ce{O}$ atom of $\ce{H2O}$ molecule donates a lone pair to the $\ce{S}$ of $\ce{SO2}$,forming a new $\ce{S\bond{-}O}$ $\sigma$-bond and breaking $\ce{S\bond{=}O}$ $\pi$-bond.


The oxygen in $\ce{H2O}$ already has its orbitals ($2s^2$,$2p^4$) filled and since oxygen is in the second row of the periodic table it can't form hypervalent molecules (That means that the oxygen would have more than 8 valence electrons).
Sulfur on the other hand is in the third row of the periodic table and can therefore form such molecules.

Because of this it is clear, that the $\ce{H2O}$ can't act as a Lewis acid as it can't have any more electrons. Therefore it has to be the Lewis-base which leaves $\ce{SO2}$ as the Lewis-acid.

  • $\begingroup$ I think what you say is misleading. Because you are saying that water cannot act as a Lewis acid. However, water is amphoteric and can act as a Lewis acid, as it has an electron deficient hydrogen atom capable of accepting an electron pair. $\endgroup$ Commented Dec 10, 2017 at 10:47
  • $\begingroup$ I think you are referring to hydrogen bonds aren't you? In that case water is indeed acting as a Lewis acid but I'm not sure if one would consider this as a real Lewis acid-vase-reaction... See quora.com/Can-water-be-classified-as-a-lewis-acid?share=1 $\endgroup$
    – Raven
    Commented Dec 10, 2017 at 11:10

Oxygen is more electronegative than sulfur, so it pulls the shared valence electrons closer to itself. This causes sulfur to have a partial positive charge and hence act as the electrophile. Furthermore, sulfur is able to have more than 8 valence electrons as it is able to expand its octet by using the d orbitals.

  • 2
    $\begingroup$ The funny thing is existence of H2SO3 in water has not been ever confirmed, only hydrated SO2 or HSO3-. $\endgroup$
    – Poutnik
    Commented Oct 27, 2021 at 13:11
  • $\begingroup$ You should say which oxygen atom you are refering to; both water and the SO2 contain oxygen atoms. $\endgroup$
    – Karsten
    Commented Oct 27, 2021 at 17:51

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