[ussername]

In the simplest approach, permanganate in acidic solutions (pH < 3.5) is generally reduced to Mn²⁺ salts while in neutral and base solutions it's reduced to MnO₂ resp. MnO₄^2-.
Why does this happen?  I can only speculate that some kinds of metal complexes are created in base solutions with sufficient concentration of OH^- and those are less oxidizing. What is the reality?

[mjc123]

Consider the general case of a redox couple which involves the participation of H⁺ or OH^-:
A: Ox + ne^- + nH⁺  Red + nH₂O
B: Ox + ne^-  Red + nOH^-
What is the difference in ΔG between these reactions? (Hint: set up a Hess's law cycle involving the reaction H⁺ + OH^-  H₂O) You will find that ΔG is more negative for A than B, and thus E° is more positive for A.
Thus for redox reactions involving the ionisation of water, the oxidising agent is more strongly oxidising in acidic than in basic solution, and reducing agents are more strongly reducing in base.
Now sometimes (especially with transition metals) there are several possible lower oxidation states, and the species formed can vary with pH. This complicates the picture, but the situation is essentially that outlined above. You might express it by saying that the oxidation state diagram is rotated anticlockwise in acid compared to base, and the minimum may come in a different place; see attached diagram.

[unsu]

It is all about the thermodynamic stabilities of different forms of Mn at different pH. You can look at the Frost diagram for Mn (see the above post by mjc123) to see which species are most stable at the specified conditions and also pay attention to the slopes (the line connecting the two species - the greater the slope the greater the reduction potential, the stronger the oxidizing abilities of the element in the high oxidation state in this particular half-reaction)

[mjc123]

Didn't know it was called a Frost diagram, you learn something every day!