Generally, the direction of a redox reaction depends on the relative strength of oxidants and reductants in a solution. In simple situations, electrochemical series, as seen in Figure 1, can be very useful for determining the direction of the reaction.
However, some reactions cannot be "eyeballed" in this manner. The reactions require a more mathematical method to determine the direction. To achieve this, it is important to consider the standard electrode potential, which is a measure of the driving force behind a reaction. The larger the value of the standard electrode potential, the further the reaction is from equilibrium. The sign of the standard electrode potential indicates the direction in which the reaction must shift to reach equilibrium.
Consider the reaction between zinc and acid:
Eo = 0.76 V
The positive Eo value indicates that when this system is present at standard-state conditions, it has to shift to the right to reach equilibrium. Reactions for which Eo is positive, therefore, have equilibrium constants that favor the products of the reaction.
What happens to the standard electrode potential with a reversal of the direction in which a reaction is written? Turning the reaction around doesn't change the relative strengths of the oxidizing or reducing agents. The magnitude of the potential must remain the same. However, turning the equation around changes the sign of the standard electrode potential, and can therefore turn an unfavorable reaction into one that is spontaneous, or vice versa.
The relative reactivities of different half reactions can be compared to predict the direction of electron flow. Half reaction equations can be combined if one is reversed to an oxidation in a manner that cancels out the electrons.