Redox reaction between hydrogen peroxide and permanganate balanced neutral

acid ==> Mn²⁺
neutral ==> MnO₂
basic ==> MnO₄^2-

Medium for reaction means H3O+, OH- or H2O (or none) on the left side of the reaction.

Agree, ,but for this reason I asked to develop an equation with sulfite, ascorbic acid sodium salt, hypochlorite or others in alkaline envirement, additional use chromate as oxidizer. Hydroxide will be fine as product.

Thank you! 

I don't agree.

Even Hydroxide is developed the reaction can take place in alkaline envirement.

Or please show, the right equation with OH- on educt side. You will see there is no.

To get rid of Peroxide, use sulfite instead and develop the equation.

3 SO₃^2- + 2 MnO₄^-  + H₂O => 2 MnO₂ + 3 SO₄^2- + 2 OH^-

Quote

And in a basic medium:

2MnO4- + 5H2O2 --> 2 Mn2+ + 2H2O + 5O2 + 6OH-

wrong balancing - basic medium means - ions OH- are consumed

In a neutral medium for manganates(VII) - ions OH- are formed

Moreover, manganese compounds catalytically decompose H2O2 then reactions can be balanced in an infinite number of ways.

And so in an acid medium H⁺ in a balanced redox equation would have to show up only in the reactants? I've never come across that condition.

In Chang's book, that I like to use to check things, there are many examples of balanced redox equations in acid medium with H⁺ in the products and balanced redox equations in basic medium with OH^- in the products.

If more base is used, the green manganate (VI) color can be seen for a moment. In a weakly alkaline medium, this cannot be seen.

Neutral and alkaline take the same calculation because pH change to alkaline

But not valid in redox envirement MnO₄ ^2- will also reduced to manganese dioxide and develop OH^-.

Manganate(VI) is not very stable and it is not known whether it is its spontaneous decay to MnO2 or reduction or both. All manganese salts also decompose hydrogen peroxide catalytically, so these reactions have virtually no stoichiometry. There is so many manganate(VII) reactions in various mediums that are relatively straightforward; those with hydrogen peroxide without any additional conditions are better left out.

This is correct, but at the end  the redoxreaction will not stop at Mn- VI. Violet, green and finally brown sludge.
You can also use Chromate CrO₄^2- and sulfite, or other reducer. In alkaline envirement it creates more hydroxide.