Formation of MnO2 vs [Mn(H2O)6]4+ in aqueous solution
Solid $\ce{MnO2}$ can be hydrated, but it's the solid which precipitates from solutions as a product of redox reactions of other $\ce{Mn}$ compounds. There is also for example crystalline $\ce{[MnF4(py)(H2O)]}$ What about aquo cations you mentioned?

One thing is that $\ce{[Mn(H2O)_n]^4+}$ cations should be indeed quite strongly acidic, but that's no problem as another strong acid in solution would prevent precipitation (similar thing like with $\ce{Fe(OH)3}$). The problem is oxidative power, what may be surprising knowing that Mn has higher oxidation states. 

Why do I think so? Well, let's see what happens if we try this trick with acid to make a solution. Concentrated $\ce{H2SO4}$ and $\ce{HCl}$ react with $\ce{MnO2}$, but lets see the reactions:

$$\ce{2 MnO2 + 2 H2SO4 → 2 MnSO4 + O2 + 2 H2O}$$
$$\ce{MnO2 + 4 HCl → MnCl2 + Cl2 + 2 H2O}$$

In first case oxygen gets oxidised and in second chlorine!

Now what if we used $\ce{HF}$? Concentrated solutions are really powerful stuff! Unfortunately the problem persists, as you could see earlier $\ce{Mn^{4+}}$ can oxidise even water if acidity is high enough. With difficulty $\ce{MnF4}$ was synthesised, but guess what, Mn can gradually oxidise fluorine in this compound, making $\ce{F2}$ gas! That was first chemical reaction known to do that.

So, generally not much chance for simple aquo complexes.

Check out these articles on Wikipedia for more info and citations to articles: about $\ce{MnF4}$ and about $\ce{MnO2}$.