In order to compare the electrode potentials of various electrodes, it is necessary to specify the concentration of the ions present in solution in which the electrode is dipped and the temperature of the half-cell.

The potential difference developed between metal electrode and the solution of its ions of unit molarity (1M) at 25°C (298 K) is called standard electrode potential.

According to the IUPAC convention, the reduction potential alone be called as the electrode potential `(E^O)` , i.e., the given value of electrode potential be regarded as reduction potential unless it is specifically mentioned that it is oxidation potential.

Standard reduction potential of an electrode means that reduc­tion reaction is taking place at the electrode. If the reaction is reversed and written as oxidation reaction, the numerical value of electrode potential will remain same but the sign of standard potential will have to be reversed. Thus

Standard reduction potential = - (Standard oxidation potential)

or

Standard oxidation potential = - (Standard reduction potential)

Standard Hydrogen Electrode, SHE

See fig.

`*` Hydrogen electrode is the primary standard electrode.

`*` It con­sists of a small platinum strip coated with platinum black as to adsorb hydrogen gas.

`*` A platinum wire is welded to the platinum strip and sealed in a glass tube as to make contact with the outer circuit through mercury.

`*` The platinum strip and glass tube is surrounded by an outer glass tube which has an inlet for hydrogen gas at the top and a number of holes at the base for the escape of excess of hydrogen gas.

`*` The platinum strip is placed in an acid solution which has `H^+` ion concentration 1 M.

`*` Pure hydrogen gas is circulated at one atmospheric pressure.

`*` A part of the gas is adsorbed and the rest escapes through holes. This gives an equilibrium between the adsorbed hydrogen and hydrogen ions in the solution.

`*` `H_2⇋ 2H^(+) + 2e^(-)`

The temperature of the cell is maintained at `25^oC`. By international agreement the standard hydrogen electrode is arbitrarily assigned a potential of exactly ± 0.000 volt.

The hydrogen electrode thus obtained forms one of two half-cells of a voltaic cell. When this half-cell is connected with any other half-cell, a voltaic cell is constituted. The hydrogen electrode can act as cathode or anode with respect to other electrode.

See Table.

The measurement of electrode potential of a given electrode is made by constituting a voltaic cell, i.e., by connecting it with a standard hydrogen electrode (SHE) through a salt bridge.

1 M solution is used in hydrogen half-cell and the temperature is maintained at `25^oC`

See fig.1.

The EMF of the cell is measured either by a calibrated potentiometer or by a high resistance voltmeter, i.e., a valve voltmeter. The reading of the voltmeter gives the electrode potential of the electrode in question with respect to the hydrogen electrode. The standard electrode potential of a metal may be determined as it is the potential difference in volt developed in a cell consisting of two electrodes: the pure metal is contact with a molar solution of one of its ions and the standard hydrogen electrode.

`text(Determination of standard electrode potential of)` `Zn//Zn^(2+)` `text(electrode:)`

`*` A zinc rod is dipped in 1 M zinc sulphate solution. This half-cell is combined with a standard hydrogen electrode through a salt bridge.

`*` Both the electrodes are con­nected with a voltmeter.

`*` The deflection of the voltmeter indicates that current is flowing from hydrogen electrode to metal electrode or the electrons are moving from zinc rod to hydrogen electrode.

`*` The zinc electrode acts as an anode and the hydrogen electrode as cathode and the cell can be represented as

See Table 1.

See fig.2.

`*` The EMF of the cell is 0.76 volt

`*` `E_(text(Cell)) = E_(text(Anode))^o + E_(text(Cathode))^o`

`0.76 = E_(text(Anode))^o + 0 text(or) E_(text(Anode))^o = +0.76 V`

`*` As the reaction on the anode is oxidation, i.e.,

`Zn → Zn^(2+) + 2e,`

`*` `E_(text(Anode))` is the standard oxidation potential of zinc. This potential is given the positive sign.

See fig.3.

`*` `E_(text(ox))^o (Zn//Zn^(2+)) = +0.76 text(volt)`

`*` So standard reduction potential of Zn, i.e., `E^o (Zn//Zn^(2+))`

`= E_(text(ox))^o = -(+0.76)`

`= -0.76 text(volt)`

`*` The EMF of such a cell gives the positive value of standard oxidation potential of metal M. The standard reduction potential `(E^o)` is obtained by reversing the sign of standard oxidation potential.

`text(Determination of standard electrode potential of)` `Cu^(2+)//Cu,` `text(electrode:)`

A copper rod is dipped in 1 M solution of `CuSO_4` . It is combined with hydrogen electrode through a salt bridge. Both the electrodes joined through a voltmeter. The deflection of the voltmeter indicates that current is flowing from copper electrode towards hydrogen electrode, i.e., the electrons ate moving from hydrogen to copper electrode. The hydrogen electrode acts as an anode and the copper electrode as a cathode. The cell can be represented as

See Table 2.

`*` The EMF of the cell is 0.34 volt.

`*` `E_(text(Cell))^o = E_(text(Anode))^o + E_(text(Cathode))^o`

`0.34 = 0 + E_(text(Cathode))^o`

`*` Since the reaction on the cathode is reduction, i.e.,

`*` `Cu^(2+) + 2e^(-) → Cu, E_(text(Cathode))^o` is the standard reduction potential of copper. This is given the +ve sign.

`*` `E°`, i.e., standard reduction potential of `Cu^(2+)//Cu = 0.34` volt

So `E°_x` (standard oxidation potential of copper) `= -0.34 text(volt)`

`*` The EMF of such a cell gives positive value of reduction potential of metal electrode. The standard oxidation potential of this electrode is obtained by reversing the sign of standard reduction potential.

`*` The EMF of such a cell gives the positive value of standard oxidation potential of metal M. The standard reduction potential (E°) is obtained by reversing the sign of standard oxidation potential.

See fig.4.



`*` It is thus concluded that at the metal electrode which acts as anode with respect to hydrogen electrode (cathode), the reduc­tion potential is given the minus sign and at the metal electrode which acts as cathode with respect to hydrogen electrode (anode), the reduction potential is given the positive sign.

`*` The standard electrode potentials (oxidation or reduction) of various elements can be measured by combining the electrode in question with a standard hydrogen electrode and measuring the emf of the cell constituted.

Since a standard hydrogen electrode is difficult to prepare and maintain, it is usually replaced by other reference electrodes, which are known as secondary reference electrodes. These are convenient to handle and are prepared easily. Two important secondary reference electrodes are described here.

`*` `text(Calomel electrode:)`

It consists of mercury at the bot­tom over which a paste of mercury-mercurous chloride is placed. A solution of potassium chloride is then placed over the paste. A platinum wire sealed in a glass tube helps in making the electrical contact. The electrode is connected with the help of the side tube on the left through a salt bridge with the other electrode to make a complete cell.

The potential of the calomel electrode depends upon the concentration of the potassium chloride solution. If potassium chloride solution is saturated, the electrode is known as saturated calomel electrode (SCE) and if the potassium chloride solution is 1 N, the electrode is known as normal calomel electrode (NCE) while for 0.1 N potassium chloride solution, the electrode is referred to as decinormal calomel electrode (DNCE). The electrode reaction when the electrode acts as cathode is:

`1//2 Hg_2Cl_2 + e^(-) ⇋ Hg + Cl^(-)`

The reduction potentials of the calomel electrodes on hydrogen scale at 298K are as follows:

See Table.

See fig.1.

The electrode potential of any other electrode on hydrogen scale can be measured when it is combined with calomel electrode. The emf of such a cell is measured. From the value of electrode potential of calomel electrode, the electrode poten­tial of the other electrode can be evaluated.


`*` `text(Ferrous – Ferric Electrode :)`

An example is a Pt wire dipping in a solution containing ferrous and ferric ions. Such a cell is described as: `Pt | Fe^(2+) (C_1), Fe^(3+) (C_2)`
The comma is used to separate two chemical species in the same solution. The electrode reaction is

`Fe^(3+) + e^(–) → Fe^(2+)`

The function of a Platinum wire is to pick up the electrons and provide electrical contact to the electrode.

`*` `text(Quin – Hydrone Electrode :)`

Oxidation reduction electrodes can also be made with organic molecules that can exist in two different oxidation states.

A generally used material of this type is related to important biochemical oxidation – reduction reactions, is the system of hydroquinone, which can form the oxidation reduction system, between quinone (Q) and hydroquinone (QH2).

See fig.2.

The presence of a Pt electrode in a solution containing these two species again clearly provides an electrode that can donate or accept electrons.

If hydro – quinone is represented by `QH_2` and quinone by `Q`, the electrode is abbreviated as

`Pt|QH_2,Q,H+ (C)`

This electrode is called a quin-hydrone electrode, because of the charged complex that is formed between `QH_2` and `Q`.

`*` `text(Metal – Metal ion Electrode :)`

An example of this is a metallic silver electrode in an `AgNO_3` solution. The electrode is represented as `Ag | Ag^(+)` (C) and the electrode reaction is `Ag^(+) + e^(–) → Ag`.

Note: Very active metals react directly with water itself and cannot be used for such electrodes.

Another example is Silver-silver chloride electrode

This is another widely used reference electrode. It is reversible and stable and can be combined with cells containing chlorides without insetting liquid junctions.

Silver chloride is deposited electrolytically on a silver or platinum wire and it is then immersed in a solution containing chloride ions. Its standard electrode potential with respect to the standard hydrogen electrode is 0.2224 V at 298 K. the electrode is represented as: `Ag|AgCl|Cl^(-)` and the electrode reaction is, `AgCl^(+)e^(-)→ Ag + Cl^(-)`


`*` `text(Amalgam Electrode :)`

In a variation of the previous electrode, the metal is in the form of an amalgam, i.e., it is dissolved in mercury, rather than in the pure form. Electrical contact is made by a Pt wire dipping into the amalgam pool. The reaction is the same as in the metal–metal ion electrode, with the Hg playing no role. A sodium amalgam electrode is represented as `Na [ text(in) Hg text(at) C_1] | Na^(+) (C_2)`.

By measuring the potentials of various electrodes versus stand­ard hydrogen electrode (SHE), a series of standard electrode potentials has been established.

When the electrodes (metals and non-metals) in contact with their ions are arranged on the basis of the values of their standard reduction potentials or standard oxidation potentials, the resulting series is called the electrochemical or electromotive or activity series of the elements.

By international convention, the standard potentials of electrodes are tabulated for reduction half reactions, indicating the tendencies of the electrodes to behave as cathodes towards SHE.

Electrodes with positive E° values for reduction half reactions do in fact act as cathodes versus SHE, while those with negative E° values of reduction half reactions behave instead as anodes versus SHE. The electrochemical series is shown in the follow­ing table.

Standard Aqueous Electrode Potentials at 25°C 'The Electrochemical Series'

See Table.

The negative sign of standard reduction potential indicates that an electrode when joined with SHE acts as anode and oxidation occurs on this electrode.

For example, standard reduction potential of zinc is -0.76 volt.

When zinc electrode is joined with SHE, it acts as anode (-ve electrode) i.e., oxidation occurs on this electrode. Similarly, the +ve sign of standard reduction potential indicates that the electrode when joined with SHE acts as cathode and reduction occurs on this electrode.

`*` The substances which are stronger reducing agents than hydrogen are placed above hydrogen in the series and have negative values of standard reduction potentials.

`*` All those substances which have positive values of reduction potentials and placed below hydrogen in the series are weaker reducing agents than hydrogen.

`*` The substances which are stronger oxidising agents than H+ion are placed below hydrogen in the series.

`*` The metals on the top (having high negative values of standard reduction potentials) have the tendency to lose electrons readily. These are active metals.

`*` The activity of metals decreases from top to bottom.

`*` The non-metals on the bottom (having high positive values of standard reduction potentials)

`*` have the tendency to accept electrons readily. These are active non-metals.

`*` The activity of non-metals increases from top to bottom.

`text(Reactivity of metals:)`

The activity of the metal depends on its tendency to lose electron or electrons, i.e., tendency to form cation (M"+). This tendency depends on the magnitude of standard reduction potential.

The metal which has high negative value (or smaller positive value) of standard reduction potential readily loses the electron or electrons and is converted into cation. Such a metal is said to be chemically active.

The chemical reactivity of metals decreases from top to bottom in the series. The metal higher in the series is more active than the metal lower in the series. For example,

`*` Alkali metals and alkaline earth metals having high negative values of standard reduction potentials are chemically active. These react with cold water and evolve hydrogen. These readily dissolve in acids forming corresponding salts and combine with those substances which accept electrons.

`*` Metals like `Fe`, `Pb`, `Sn`, `Ni`, `Co`, etc., which lie a little down in the series do not react with cold water but react with steam to evolve hydrogen.

`*` Metals like `Cu`, `Ag` and `Au` which lie below hydrogen are less reactive and do not evolve hydrogen from water.

`text(Electropositive character of metals:)`

The electropositive character also depends on the tendency to lose electron or electrons. Like reactivity, the electropositive character of metals decreases from top to bottom in the electrochemical series. On the basis of standard reduction potential values, metals are divided into three groups:

See fig.1.

`*` Strongly electropositive metals: Metals having standard reduction potential near about -2.0 volt or more negative like alkali metals, alkaline earth metals are strongly electropositive in nature.

`*` Moderately electropositive metals: Metals having values of reduction potentials between 0.0 and about -2.0 volt are moderately electropositive. `Al`, `Zn`, `Fe`, `Ni`, `Co`, etc., belong to this group.

`*` Weakly electropositive metals: The metals which are below hydrogen and possess positive values of reduction potentials are weakly electropositive metals. `Cu`, `Hg`, `Ag`, etc., belong to this group.

`text(Displacement reactions :)`

See fig.2.



`*` To predict whether a given metal will displace another, from its salt solution: A metal higher in the series will displace the metal from its solution which is lower in the series, i.e., the metal having low standard reduction poten­tial will displace the metal from its salt's solution which has higher value of standard reduction potential. A metal higher in the series has greater tendency to provide electrons to the cations of the metal to be precipitated.

`*` Displacement of one nonmetal from its salt solution by another nonmetal: A nonmetal higher in the series (towards bottom side), i.e., having high value of reduction potential will displace another nonmetal with lower reduction potential i.e., occupying position above in the series. The nonmetal's which possess high positive reduction potentials have the tendency to accept electrons readily. These electrons are provided by the ions of the nonmetal having low value of reduction potential. Thus, `Cl_2` can displace bromine and iodine from bromides and iodides.

See fig.3.

[The activity or electronegative character or oxidising nature of the nonmetal increases as the value of reduction potential increases.]

`*` Displacement of hydrogen from dilute acids by metals: The metal which can provide electrons to H+ ions present in dilute acids for reduction, evolve hydrogen from dilute acids.

See fig.4.

The metal having negative values of reduction potential possess the property of losing electron or electrons.

Thus, the metals occupying top positions in the electrochemical series readily liberate hydrogen from dilute acids and on descending in the series tendency to liberate hydrogen gas from dilute acids decreases.

The metals which are below hydrogen in electrochemical series like Cu, Hg, Au, Pt, etc., do not evolve hydrogen from dilute acids.

`*` Displacement of hydrogen from water: Iron and the metals above iron are capable of liberating hydrogen from water. The tendency decreases from top to bottom in electrochemical series.Alkali and alkaline earth metals liberate hydrogen from cold water but Mg, Zn and Fe liberate hydrogen from hot water or steam.

`text(Reducing power of metals:)`

Reducing nature depends on the tendency of losing electron or electrons. More the negative reduction potential, more is the tendency to lose electron or electrons. Thus, reducing nature decreases from top to bottom in the electrochemical series. The power of the reducing agent increases as the standard reduction potential becomes more and more negative.

Sodium is a stronger reducing agent than zinc and zinc is a stronger reducing agent than iron.

See fig.5.

Alkali and alkaline earth metals are strong reducing agents.

Oxidising nature of nonmetals:

Oxidising nature depends on the tendency to accept electron or electrons. More the value of reduction potential, higher is the tendency to accept electron or electrons. Thus, oxidising nature increases from top to bottom in the electrochemical series. The strength of an oxidising agent increases as the value of reduction potential becomes more and more positive.

`F_2` (Fluorine) is a stronger oxidant than `Cl_2, Br_2` and `I_2`.

`Cl_2` (Chlorine) is a stronger oxidant than `Br_2` and `I_2` .

See fig.6.

`text(Thermal stability of metallic oxides:)`

`*` The thermal stability of the metal oxide depends on its electropositive nature.

`*` As the electropositivity decreases from top to bottom, the thermal stability of the oxide also decreases from top to bottom.

`*` The oxides of metals having high positive reduction potentials are not stable towards heat.

`*` The metals which come below copper form unstable oxides, i.e., these are decomposed on heating.

`Ag_2O oversettext(Heat)-> 2 Ag + O_2`

`2HgO oversettext(Heat)-> 1/2 O_2 + 2Hg`

`text(Products of electrolysis:)`

In case two or more types of positive and negative ions are present in solution, during electrolysis certain ions are discharged or liberated at the electrodes in preference to others. In general, in such com­petition the ion which is stronger oxidising agent (high value of standard reduction potential) is discharged first at the cathode.

The increasing order of deposition of few cations is:

`underset(->)(K^+, Ca^(2+), Na^+, Mg^(2+), Al^(3+), Zn^(2)+, Fe^(2+), H^+, Cu^(2+), Ag^+, Au^(3+))`

`text(Increasing order of deposition`

Similarly, the anion which is stronger reducing agent (low value of standard reduction potential) is liberated first at the anode.

The increasing order of discharge of few anions is:

`underset(->)(SO_4^(2-), NO_3^(-), OH^(-), Cl^(-), Br^(-), I^(-))`

`text(Increasing order of discharge)`

Thus, when an aqueous solution of NaCl containing `Na^(+), Cl^(-), H^(+)` and `OH"` ions is electrolysed, `H^(+)` ions are discharged at cathode and CF ions at the anode, i.e., `H_2` is liberated at cathode and chlorine at anode.

When an aqueous solution of `CuSO_4` containing `Cu^(2+), , H^+` and `OH^(-)` ions is electrolysed, `Cu^(2+)` ions are dis­charged at cathode and `OH^(-)` ions at the anode.

See fig.7.

Cu is deposited on cathode while `O_2` is liberated at anode.

`text(Corrosion of metals:)`

Corrosion is defined as the deterioration of a substance because of its reaction with its environment. This is also defined as the process by which metals have the tendency to go back to their combined state, i.e., reverse of extraction of metals.

Ordinary corrosion is a redox reaction by which metals are oxidised by oxygen in presence of moisture. Oxidation of metals occurs more readily at points of strain. Thus, a steel nail first corrodes at the tip and head. The end of a steel nail acts as an anode where iron is oxidised to `Fe^(2+)` ions.

`Fe → Fe^2 + 2e^(-)` (Anode reaction)

The electrons flow along the nail to areas containing im­purities which act as cathodes where oxygen is reduced to hydroxyl ions.

`O_2 + 2H_2O + 4e^(-) → 4OH^(-)` (Cathode reaction)

See fig.8.

The overall reaction is

`2Fe + Oz + 2H_2O = 2Fe(OH)_2`

`Fe(OH)_2` may be dehydrated to iron oxide, `FeO`, or further oxidised to `Fe(OH)_3` and then dehydrated to iron rust, `Fe_2O_3` .

Several methods for protection of metals against corrosion have been developed. The most widely used are (i) plating the metal with a thin layer of a less easily oxidised metal (ii) allowing a protective film such as metal oxide (iii) galvanis­ing-steel is coated with zinc (a more active metal).

`text(Extraction of metals:)`

A more electropositive metal can displace a less electropositive metal from its salt's solution. This principle is applied for the extraction of Ag and Au by cyanide process.

Silver from the solution containing sodium argento cyanide, `NaAg(CN)_2`, can be obtained by the addition of zinc as it is more electro-positive than Ag.

`2NaAg(CN)_2 + Zn → Na_2Zn(CN)_4 + 2Ag`