The applications of electrochemical series (or electromotive series) are as follows :
(1) Relative strength of oxidising agents in terms of `E_("red")^(0)` values : The `E_("red")^(0)` value is a measure of the tendency of the species to be reduced i.e. to accept electrons and act as an oxidising agent. The species mentioned on left hand side of the half reactions are oxidising agents.
The substances in the upper positions in the series and hence in the upper left side of the half reactions have large positive Ered values hence are stronger oxidising agents. For example, `F_(2+), Ce^(4+) , Au^(3+)`, etc. As we move down the series, the oxidising power decreases. Hence from the position of the elements in the electrochemical series,oxidising agents can be selected.
(2) Relative strength of reducing agents in terms of `E_("red")^(0)` values: The lower `E_("red")^(0)` value means lower tendency to accpet electrons but higher tendency to lose electrons. The tendency for reverse reaction or oxidation increases as `E_("red")^(0)` becomes more negative and we move towrads the lower side of the series. For example, Li, K , Al etc. are good reducing agent .
(3) Identifying the spontaneous direction of reaction : From the standard reduction potentials, Eed the spontaneity of a redox reaction can be determined. The difference between `E_("red")^(0)` value for any tow electrodes represenets cell potenital `E_("cell")^(0)` , constituted by them.
If `E_("cell")^(0)` is positive then the reaction is spontaneous while if `E_("cell")^(0)` is negative the reaction is non-spontaneous . For example `E_(Mg^(2+)//Mg)^(2)` and `E_(Ag^(+)//Ag)^(0)` have values - 2.37V and 0.8 V respectively. The Mg will be a better reducing agent than Ag. Therefore Mg can reduce `Ag^(+)` to Ag.
The corresponding reaction will be .
`Mg_((s)) + 2Ag_((aq))^(+) + 2Ag_((s))`
`E_("cell")^(0) =E_(Ag^(+)//Mg)^(0) = 0.8 -(-2.37)`
= 3.17V
Therefore above reaction in the forward direction will be spontaneous while in the reverse direction will be non-spontaneous. since for it `E_("cell")^(0) = - 3.17` V.
(4) Calculation of standard cell potential `E_("cell")^(0)` From the electrochemical series, the standard cell potential, `E_("cell")^(0) ` form the `E_("red")^(0) ` value for the half reactions gives can be calculate.
For example
`Zn^(2+) + 2e^(-) to Zn_((s)) " "E_(Zn^(2+)//Zn)^(0) = - 0.76V`
`Cu^(2+) + 2e^(-) to Cu_((s)) " " E_(Cu^(2+)//Cu)^(0) = - 0.34 V`
For the cell,
`Zn|Zn_((aq))^(2+) (1M) || Cu_((aq))^(2+) (1M) |Cu`
`E_("cell")^(0) = E_(underset("(cathode)")(Cu^(2+)//Cu))^(0) -E_(underset("(andoe)")(Zn^(2+)//Zn))^(0)`
`= 0.34-(-0.76)`
= 1.1
