I read in my book this: $$\ce{H2SO3 + H2O -> HSO3- +H3O+}$$ while I rather would write (not balanced): $$\ce{H2SO3 + H2O -> 2 H+ + SO3^2- + 2 H2O -> SO3^2- + 2H3O+}$$ I mean I don't understand why the book writes only a partial deprotonated acid. Someone could explain this?

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    $\begingroup$ What is justification to ever write not balanced chemical reaction equation ? $\endgroup$ – Poutnik Oct 31 '19 at 16:36
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    $\begingroup$ The middle equation would presumably show intermediate forms. There is no free H+ ion floating freely in the water as an intermediate. $\endgroup$ – MaxW Oct 31 '19 at 16:55

$\ce{H2A}$ and $\ce{A^2-}$ forms of a biprotic acid seldom coexist as major components, expecially if the acid is weak, unless $\mathrm{p}K_\mathrm{a1}$ and $\mathrm{p}K_\mathrm{a2}$ are similar. Instead, coexistence of major components $\ce{H2A / HA-}$ or $\ce{HA- / A^2-}$ occurs.

$\ce{SO3^2-}$ starts to form in about neutral or alkalic solutions, while $\ce{SO2}$ solution is acidic.
As $\mathrm{p}K_\mathrm{a2}=7.18$ - see polyprotic_acids

The wikipedia article about Sulfurous_acid says:

Raman spectra of solutions of sulfur dioxide in water show only signals due to the $\ce{SO2}$ molecule and the bisulfite ion, $\ce{HSO3−}$. The intensities of the signals are consistent with the following equilibrium:

$\ce{SO2 + H2O <=> HSO3- + H+}$
$K_\mathrm{a} = 1.54 \cdot 10^{−2}$
$\mathrm{p}K_\mathrm{a} = 1.81$

Note that $\ce{H+}$ is just convenience for $\ce{H3O+}$, so it is rather $\ce{SO2 + 2 H2O <=> HSO3- + H3O+}$

See hydronium for complex chemistry of proton hydration.

Similarly, the strong acid $\ce{H2SO4}$ in not too diluted solution disociates just to $\ce{HSO4-}$. $\ce{SO4^2-}$ is a minor ion.

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  • $\begingroup$ Kudos for pointing out the "real" equilibrium between SO2 and bisulfite. Problems ought to reflect reality so that the student is learning chemistry at the same time as how the mathematics of chemistry works. $\endgroup$ – MaxW Oct 31 '19 at 18:54

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