Electrochemistry - Reduction Potential in Feasibility of Redox Reactions

In an earlier entry on the topic of Electrochemistry -Reduction Potential, I spoke about my interpretation of reduction potential, whose main tenet is for us to decide which substance is more susceptible to reduction and which is less. Therefore, having the positive and negative sign is IMPORTANT.

In this entry I would be writing about the application (or uses) of the Standard Reduction Potential. There are three main uses, (a) to determine the Ecell of an electrochemical cell. (b) to determine whether a redox reaction will occur and (c) deciding which substance will be electrolysed. These will be discussed in a series of three posts. The second one is:

(b) To determine feasibility of a redox reaction

Firstly, you need to identify which is the species that will be oxidised and which will be reduced. For example, in the question of asking whether I^- will cause the reduction of Br₂. In this situation I^- will be oxidised to I₂, while Br₂ will be reduced to Br^-.

Therefore, by using logical deduction, for the redox reaction between Br₂ and I^- to feasible, it must be that Br₂ tends to reduce more favourably than I₂. (If the converse is true, wouldn't the formation of I₂ and Br^- then cause I₂ to oxidise Br^-?)

Thus, if we substract the reduction potential of I₂ from the reduction potential of Br₂, we will get a positive number. This positive number is the Ecell of this redox reaction and it informs us that the redox reaction is spontaneous.

Therefore, we will have the equation Ecell = Ered - Eoxd . (Since reduction potential of Br₂ is greater than I₂ (+1.07V against +0.54V), thus Br₂ will oxidise I^-)

This equation can also be used to predict whether a reaction is non-spontaneous. For example, to predict whether a reaction between F^- and I₂ will occur, we are suggesting that F₂ and I^- are formed.

However, the reduction potential of I₂ is smaller than the reduction of F₂. Therefore the redox reaction between F^- and I₂ wouldn't occur. To prove it, we calculate Ecell of the reaction (Ecell = +0.54V - (+2.87)) which is smaller than zero.

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Article written by Kwok YL 2007.

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