Here's a way (at least that I use) to keep the signs (+/-) straight when working with reduction potentials.
Remember the formula for the Gibbs Free Energy of an electrochemical cell:
ΔG°= -nFE°
where n is the number of electrons transferred per mole of reactant and E° is the standard reduction potential. n, and F are positive.
When ΔG° is negative, the reaction is thermodynamically favourable (K>1), so the reactant will 'want' to be reduced (though a high energy barrier can prevent this, like not hooking up the wire in your beakers, for example, but this is a kinetic issue, not a thermodynamic one).
Anyway, for ΔG° to be negative, E° must be positive, and the more positive it is the more negative ΔG° is, and thus the more thermodynamically favourable the reduction half-reaction is.
One other thing that I used to mess up all the time, which isn't that relevant but worth saying - when you're adding two half-reactions together to get an overall oxidation/reduction reaction, the standard reduction potentials simply add together, regardless of whether or not you have to multiply one reaction by 2 or 3 so the electrons on both sides cancel out. Basically, multiplying a half-reaction by a number N doesn't change its reduction potential by a factor of N; it's still just E°.
