Lewis Acid / Base reaction for SO2 and H2O 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?
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:

$$\ce{O\bond{=}S\bond{=}O + H2O &lt;=&gt; H2SO3}$$



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.