Identifying Half Reactions

[Hussain Chharawalla 1G]

In our discussion, we did an example of balancing redox reactions for:
MnO4- (aq) + H2SO3 (aq) --> Mn2+ (aq) + HSO4-

In section, it was said that MnO4- was reduced and H2SO3 was oxidized. I didn't quite understand this because to me it seemed like Mn2+ means that two electrons were lost meaning oxidized, and the negative charge on HSO4- means it gained an electron and reduced. Can someone explain this please

Thanks

[Christine Honda 2I]

MnO4- is being reduced because when we calculate the oxidation of state of Mn in this compound it is +7 (do this knowing that O is -2, multiplying it by 4 and knowing that the entire molecule has a charge of -1). But in the products, Mn has a charge of +2. This shows that Mn has gained 5 electrons and has been reduced.

[Riya Sood 4G]

The best way to recognize which element is oxidized and reduced is to look at the oxidation number.
The general rule is that:
1. If oxidation number increases= element is oxidized (reducing agent)
2. If oxidation number decreases= element is reduced (oxidizing agent).

[Matt F]

A little bit of cheat too is that MnO₄^- lost its Oxygen, while H₂SO₃ gained an Oxygen. Therefore, the MnO₄^- is reduced while the H₂SO₃ is oxidized. I'm not sure if this is true in every reaction, but it makes sense

[Emil Velasco 1H]

Gaining electrons is reduction (going for negative) and losing electrons is oxidation (going for positive)