Electrochemical series (Electromotive series)

The standard potentials of a number of electrodes have been determined using the standard hydrogen electrode. These electrodes, with their half reactions, are arranged according to their decreasing standard potentials, as shown in Table 5.1. This arrangement is called the electrochemical series.

Key points of electrochemical series

i. The half reactions are written as reductions. The oxidizing agents and electrons appear on the left side of the half reactions, while the reducing agents are shown on the right side of the half reactions.

ii. Below the hydrogen electrode the negative standard potential increases, and above the hydrogen electrode the positive standard potential increases.

iii. $E^0$ values apply to the reduction half reactions that occur in the forward direction as written.

iv. A higher (more positive) $E^0$ value for a half reaction indicates its greater tendency to occur in the forward direction, and in turn a greater tendency for the substance to reduce. Conversely, a low (more negative) $E^0$ value of a half reaction corresponds to its greater tendency to occur in the reverse direction, or for the substance to oxidise. The half reactions are listed in order of their decreasing tendency in the forward direction.

Applications of electrochemical series

i. Relative strength of oxidising agents : The species on the left side of the half reactions are oxidizing agents. The $E^0$ value is a measure of the tendency of the species to accept electrons and get reduced — in other words, the $E^0$ value measures the strength of a substance as an oxidising agent. The larger the $E^0$ value, the greater is the oxidising strength. The species in the top left side of the half reactions are strong oxidising agents; as we move down the table, the $E^0$ value and the strength of the oxidising agents decrease from top to bottom.

ii. Relative strength of reducing agents : The species on the right side of the half reactions are reducing agents. The half reactions at the bottom of the table, with large negative $E^0$ values, have little or no tendency to occur in the forward direction as written; they tend to favour the reverse direction. It follows that the species appearing at the bottom right side of the half reactions, associated with large negative $E^0$ values, are effective electron donors — they serve as strong reducing agents. The strength of the reducing agents increases from top to bottom as the $E^0$ values decrease.

iii. Spontaneity of redox reactions : A redox reaction in a galvanic cell is spontaneous only if the species with the higher $E^0$ value is reduced (accepts electrons) and that with the lower $E^0$ value is oxidised (donates electrons). The standard cell potential must be positive for a cell reaction to be spontaneous under standard conditions. A noteworthy application of the electromotive series is predicting the spontaneity of redox reactions from the knowledge of standard potentials.

Suppose we ask a question : at standard conditions, would $\mathrm{Ag^+}$ ions oxidise metallic magnesium ? To answer this question, first we write the oxidation of Mg by $\mathrm{Ag^+}$:

From Table 5.1 of the electrochemical series we have $E^0_{Mg} = -2.37$ V and $E^0_{Ag} = 0.8$ V. For the cell having Mg as anode and Ag as cathode,

The EMF being positive, the cell reaction is spontaneous: $\mathrm{Ag^+}$ ions oxidise metallic Mg.

General rules

i. An oxidizing agent can oxidize any reducing agent that appears below it, and cannot oxidize the reducing agent appearing above it in the electrochemical series.

ii. A reducing agent can reduce the oxidising agent located above it in the electrochemical series.

Table 5.1 : The standard aqueous electrode potentials at 298 K (Electrochemical series)

In the printed table, the oxidising agents on the left grow stronger from bottom to top (strongest at the $\mathrm{F_2}$ row) and the reducing agents on the right grow stronger from top to bottom (strongest at the Li row).