How do I balance these oxidation reduction equations using the half cell method?

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Balancing oxidation-reduction equations can be complex, but using the half-cell method simplifies the process. To balance the given equations, begin by identifying the oxidation and reduction half-reactions. For example, in the reaction involving SO3^2- and MnO4-, the half-reaction for manganese is MnO4- being reduced to Mn2+, while SO3^2- is oxidized to SO4^2-. Similarly, for the reaction with Cl2 and OH-, Cl2 is reduced to Cl-, and Cl- is oxidized to ClO3-. In the final equation with SO4^2- and I-, SO4^2- is reduced to S2-, while I- is oxidized to I2. Each half-reaction must be balanced for mass and charge, ensuring that electrons lost in oxidation equal those gained in reduction. This methodical approach clarifies the balancing process and helps in understanding the underlying redox chemistry.
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I do not understand how to balance these oxidation reduction equations.
These equations seem really complicated, I don't understand them. The question that i have to answer is
Balance the following equations by the half cell method. Show both half cell reactions and identify them as oxidation or reduction.
a) SO3^2- + MnO4-+ H+ (arrows) Mn2+ + SO4^2- + H2O
b) Cl2 + OH- (arrows) Cl- + ClO3- + H2O
c) SO4^2- + I- + H+ (arrows) S2- + I2 + H2O
 
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You need to start by writing out the half reactions. For example the half reaction of iron +3 being reduced to iron +2 would be:

Fe+3 + e- ---> Fe+2

Try it from there.
 

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