Electrodes and Their Types

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Working of Electrochemical Cell:

  • Electrochemical (Voltaic or Galvanic) cell consists of two half cells. In one half-cell, an oxidation reaction takes place and electrons are generated in the process. While in the second half cell reduction reaction takes place and electrons are absorbed or consumed in this process.
  • The two half-cells are connected to each other internally by porous partition and connected externally by means of connecting wire.
  • Both the oxidation and reduction reaction takes place simultaneously and separately
  • The movement of electrons from oxidation half cell to reduction half-cell through external circuit constitutes the electrical current.

Convention for Representing the Voltaic Cell:

  • Every time it is not possible to draw neat diagrams of the cell.  So as to avoid this difficulty, following conventions (rules) are used.
  • The electrode with greater oxidation potential (negative electrode or anode) is written to the left.  Here oxidation takes place.
  • The electrode material is written first, followed by its electrolyte.
  • The electrode with lower oxidation potential ( positive electrode or cathode ) is written to the right.  Here reduction takes place.
  • The electrolyte is written first followed by the electrode material.
  • A single vertical line is drawn between the electrode and its electrolyte which represents direct contact but the separation of the two phases.
  • The double vertical line is drawn between two salt solutions which indicate indirect contact by means of a porous pot or a salt bridge.
  • Concentration or activity of solutions at two electrodes is written in brackets as (C = 1),    (C = 2) or (a = 1), (a  = 2).
  • In case of gas electrodes, inert metal conductors like platinum is used to establish electrical contact. It should be incorporated in the convention. e.g. a) Hydrogen electrode is represented as

Electrochemical Cells 01

  • Representation of Daniell cell by above convention is as follows.

Electrochemical Cells 02



By applying above conventions some cells are represented below:

A cell consisting of zinc and cadmium electrodes:

  • As zinc has higher oxidation potential, it acts as -ve electrode.  Cadmium acts as a +ve electrode.  Oxidation takes place at zinc while reduction takes place at Cadmium.

–   Z n   |   Z n++   ||   Cd++  |   Cd   +

( 1 M )       ( 1 M)



  • At anode (Zn) :  Zn  → Zn++  + 2 e   (oxidation)
  • At Cathode (Cd): Cd++  + 2 e →   Cd  (reduction)
  • Net cell reaction :  Zn   + Cd++ →  Zn++ + Cd     (redox)

Representation of a cell consisting of hydrogen and copper electrodes:

  • As hydrogen has greater oxidation potential than copper, acts as -ve electrode (anode) and copper act +ve electrode (cathode).  Oxidation takes place at hydrogen and reduction takes place at copper.

Electrochemical Cells 03

  • At anode (H2) :     H2(g) →  2 H+ +   2 (oxidation)
  • At cathode (Cu) :  Cu++ + 2 e–  Cu (reduction)
  • Net cell reaction :  H2(g) + Cu++  →  2H+ + Cu (redox)

Representation of a cell consisting of copper and silver electrodes:

  • As copper has greater oxidation potential than silver, acts as anode and silver act as the cathode.  Oxidation takes place at copper and reduction takes place at silver.

–    Cu  | Cu++ ||    Ag+  |   Ag+

(C =1 )   (C = 2)

  • At anode(Cu):  Cu  →  Cu++  + 2 e   (oxidation)
  • At cathode (Ag):  2 Ag+   +   2 e →  2 Ag     (reduction)
  • Net cell reaction: Cu + 2 Ag+   →  Cu++ + 2 Ag (redox)

Representation of cell from the cell reaction:

1/2 Cl2+   e →  Cl



2 OH–  →  1/2 O2 +  H2O   +  2e

  • As chlorine undergoes reduction hence chlorine is a cathode and it must be written to right.  In the second reaction OH ions undergo oxidation hence oxygen electrode is an anode and it must be written to the left.  The cell can be represented as,

Pt, O2(g) | OH (aq) || Cl(aq) | Cl2(g) ,Pt

Salt Bridge:

  • It consists of inverted U shaped glass tube containing the saturated solution of a strong electrolyte like KCl, KNO3, NH4NO3 immobilised by agar-agar gel with glass wool plugs at the two ends.

Electrochemical Cells 04

  • Working: The two electrolytic solutions in two half-cells are connected by dipping the arms of the tube in inverted position in the solutions.

Electrodes and its Types:

  • Electrodes are metallic or non-metallic rods immersed in the electrolyte. They conduct electric current through them. Carbon and platinum are mainly used electrodes because they are inert and do not get dissolved in the electrolytic solution.
  • Electrodes are of two types, a) indicator electrode b) Reference Electrode

Indicator Electrode:

  • The electrode whose potential depends upon the concentration of a particular ion in the solution in which it is dipped and is usually used to find out the concentration of ions in the solution is known as indicator electrode.
  • It usually consists of metal in the form of wire or rod kept in contact with its salt solution.

e.g. Ag(s) | Ag+ (aq)     (a = x M)



  • The electrode system is used for determination of the concentration of the solution used in that half cell. All electrodes except the reference electrode are called indicator electrodes.

Reference Electrode:

  • The electrode whose potential is arbitrarily fixed or is exactly known at a given constant temperature is known as a reference electrode.
  • Using reference electrode unknown potential of any other single electrode can be found out e.g. Two commonly used reference electrodes are
  • Standard hydrogen electrode (SHE) and Calomel electrode.
  • A cell is constructed using the given electrode and reference electrode. Using a potentiometer and standard cell (like Weston cell) the e.m.f. of the cell can be measured. By knowing the e.m.f. of the cell and potential of the reference electrode, the potential of the electrode in question can be easily determined.

Construction and Working of Standard Hydrogen Electrode (SHE).

  • SHE is defined as the electrode in which pure and dry hydrogen gas is bubbled at 1 atmospheric pressure and 298 K on a platinized platinum foil through a solution containing H+ ions at unit activity.

Construction:

Electrodes 02

  • SHE consists of a glass jacket which has a small inlet at the top and many outlets at the bottom.  Inside the glass jacket, there is a glass tube closed at both the ends.
  • It has a platinum wire sealed in it. At the lower end of the platinum wire, there is a platinized platinum plate.
  • At the bottom of the glass tube, there is little mercury which is meant for good electrical contact.
  • The glass jacket along with glass tube is dipped in a vessel containing 1 M HCI solution.

Working:

  • Pure and dry hydrogen gas is bubbled through HCI solution from the inlet at a constant pressure of 1 atm.  Hydrogen gas is adsorbed on the platinum plate and acts as hydrogen electrode.  An equilibrium between H2 gas and H+ ion is established across the metal.

Electrode reaction:

  • The electrode is reversible with respect to hydrogen ions. During working, hydrogen gas from platinum plate changes into hydrogen ions and electrons are set free. These electrons accumulate on the platinum plate.
  • If electrode is serving as anode, then the half-cell reaction is

H2(g) →   2H+(aq) + 2e (oxidation)

  • The electrons set free remains on the platinum plate and transferred to the other electrode through Pt. wire.  As the process is oxidation, a positive potential is developed.  It is comparatively very small, it is arbitrarily taken as a zero.
  • If electrode is serving as cathode, then the half-cell reaction is

2H+(aq) + 2e  →  H2(g) (reduction)



Representation of electrode:

  • When acting as anode,     Pt| H2(g) (1 atm.)| H+(aq) (1 M)
  • When acting as cathode,   H+(aq) (1 M) |  H2(g) (1 atm.) | Pt

Advantages 0f SHE:

  • SHE is used as reference electrode.  When it is coupled with any other electrode whose potential is to be determined.  Potential of the cell is then measured using a potentiometer.  Since the potential of SHE is zero, the potential of the cell is equal to the potential of another electrode or e.m.f. of the cell itself. Thus when SHE is used the correction for its own potential is not necessary.
  • It can be used over entire pH range.
  • It gives no salt error.
  • It consists pH scale with voltage measurement.

Difficulties in setting up of SHE:

  • It is difficult to obtain 100 % pure and dry hydrogen gas. Even traces of impurities in the hydrogen gas makes the electrode inactive and irreversible.
  • It is difficult to maintain exactly 1 atmospheric pressure on hydrogen gas for a longer time.
  • It is difficult to maintain the concentration of HCI solution as 1 M because due to bubbling of hydrogen gas through HCI solution, water is evaporated and hence the concentration of HCI solution may change.
  • Since it is made up of glass, it is not so handy.
  • Platinum used is rather expensive.
  • It is difficult to prepare ideal platinised platinum.

Arrangement to Find Oxidation Potential of another Electrode Using SHE:

Electrodes 06

Calomel Electrode:

Electrodes 03



Construction:

  • Calomel electrode consists of a broad glass tube having side arm as shown in the figure. The sidearm is used for dipping it any solution used for coupling the calomel electrode.
  • At the bottom of the glass tube, there is pure mercury and a platinum wire is sealed into it at the bottom for electrical connections. The wire runs through separator glass tube to the top of the tube for electrical contact.
  • Above pure mercury, there is a paste of mercurous chloride (calomel) (Hg2Cl2) in mercury.
  • Rest of the glass vessel and sidearm A is filled with saturated KCl solution.  KCI solution of 0.1 M or of 1 M can also be used.  Sidearm is plugged with glass wool.
  • The glass tube is closed from the top.

Working:

  • Since calomel electrode is reversible, two types of reaction are possible depending upon the nature of another electrode with which it is coupled.
  • When acting as  negative electrode:

2 Hg(l) → 2 Hg+ + 2 e

2 Hg+ + 2 Cl – → Hg2Cl2(s)

The net oxidation reaction is

2Hg(l) + 2Cl(sat) → Hg2Cl2(s)+  2e

Thus oxidation takes place when it is coupled with other electrode having lower oxidation potential.



  • When acting as positive electrode:

Hg2Cl2(s)      →   2Hg(l) + 2Cl

 2 Hg+ + 2 e    → 2 Hg

The net reduction reaction is

Hg2Cl2(s)   +  2 e   →  2 Hg+ + 2Cl



Thus reduction takes place when it is coupled with other electrode having greater oxidation potential.

Representation of Electrode:

  • When acting as anode:   Pt | Hg(l) | Hg2Cl2(s)  | KCl(sat)
  • When acting as anode:  KCl(sat) |  Hg2Cl2(s) | Hg(l) |Pt

Oxidation Potential of Calomel Electrode:

  • The oxidation potential of calomel electrode depends upon the concentration of KCI solution used. The negative potentials indicate that when combined with SHE reduction takes place at calomel electrode.

Electrodes 03

Advantages of calomel electrode:

  • It is easy to set up and easily reproducible.
  • It is convenient and easy to transport.
  • It is very compact and smaller in size requires little space.
  • No separate salt bridge is required as it has already a side tube containing KCl solution.
  • Potential does not change appreciably with time and a slight change in temperature.

Disadvantages of Calomel Electrode:

  • When half-cell potentials are to be measured, compensation for potential is necessary.
  • Calomel electrode cannot be used in the measurement of potentials of the cell where K+ and Cl – ions interfere in the electrochemical reactions of the cell.
  • The oxidation potential of the electrode depends on the concentration of KCl. If the concentration of KCl changes, the oxidation potential of electrode changes.


Different types of electrodes:

  • There are four types of electrodes
    • Gas electrodes
    • Metal–sparingly soluble metal salt electrodes
    • Metal – metal ion electrodes
    • Redox Electrodes

Gas Electrodes:

  • Gas electrode consists of a gas ( e.g. H2, Cl2, O2) in contact with a solution containing the ions derivable from the gas e.g. H+, Cl ,OH.
  • The potential of gas electrode depends upon the concentration of its ions in the solution and the pressure of a gas.
  • A gas electrode consists of a gas , bubbled about inert metal wire (platinized platinum electrode) immersed in a solution containing ions with which gas is irreversible. Platinum is used as conductor and to adsorb the gas.
  • e.g. Standard hydrogen electrode.

Examples of gas Electrodes:

  • Standard Hydrogen Electrode (SHE):

  • SHE is represented as,

 Pt| H2(g) (1 atm.)| H+(aq) (1 M)

  • The half cell reactions are

H2(g) →   2H+(aq) + 2e (oxidation) (L.H.S.)

2H+(aq) + 2e → H2(g)  (reduction) (R.H.S.)

The electrode potential is arbitrarily assigned zero. This electrode is cation electrode.



  • Chlorine gas electrode:

  • This electrode is anion electrode.  Chlorine gas electrode is represented as,

Pt| Cl2(g) (1 atm.)| Cl(aq) (1 M)

  • The half cell reactions are

2Cl(aq) → Cl2(g) + 2e   (oxidation) (L.H.S.)

Cl2(g) + 2e →  2Cl(aq)   (reduction) (R.H.S.)

  • Oxygen gas electrode:

  • Oxygen gas electrode is represented as,

Pt | O2(g) (1 atm)| OH (aq) (1M)



  • The half cell reaction is

4OH → 2H2O+ O2(g) + 4e   (oxidation) (L.H.S.)

2H2O + O2(g) + 4e– →  4OH (reduction) (R.H.S.)

Reversible Anion Electrode or Metal-Sparingly Soluble Metal Salt Electrode:

  • Reversible anion electrode is also called as metal- sparingly soluble metal salt electrode.
  • In this electrode a metal, a sparingly soluble salt of the metal in equilibrium with a solution containing the same anion as the sparingly soluble salt.
  • E.g. Calomel electrode.

Metal – Metal Ion Electrodes:

  • In this case, the metal strip is kept in contact with the solution of a water-soluble salt containing cation of the same metal.
  • e.g. Zn(s) | Zn++(aq)
  • In the electrochemical cell, the electrode having higher oxidation potential undergoes oxidation and acts as the anode or negative electrode and the electrode having lower oxidation potential undergoes reduction and acts as the cathode or positive electrode.

Examples of metal – metal ions electrodes:

  • Zn(s) | Zn++(aq)

Zn(s) →  Zn++(aq) +   2e–   (Oxidation)

Zn++(aq) +   2e–  → Zn(s)    (Reduction)

  • Cu(s) | Cu++(aq)

Cu(s) →  Cu++(aq) +   2e–   (Oxidation)

Cu++(aq) +   2e–  → Cu(s)    (Reduction)



Redox Electrode:

  • In these electrodes, an inert metal like Pt is dipped in a solution containing ions of an active metal in two different oxidation states.
  • Pt | Fe2+, Fe3+

Fe2+      →     Fe3+  e (Oxidation)

Fe+++ +   e      →   Fe++ (Reduction)

  • Pt | Sn2+, Sn4+

Sn2+  →  Sn4+  +     2e (Oxidation)

Sn4+ +  2e    →   Sn2+ (Reduction)

Redox Potential:

  • The potential developed due to the ability of ions to lose or gain electrons forming higher or lower stable oxidation state is called redox potential.
  • The redox potential depends upon the ratio of concentrations of two types of ions.

Pt | Fe2+(aq)(1M),  Fe3+ (aq) (1M)        E2ox= – 0.771 V

Representation of cells containing standard and reference electrodes:

  • A cell composed of zinc rod contact with 1 molar zinc ion solution and saturated calomel electrode.

– Zn(s)| Zn2+(1M) || KCl(aq) (saturated) | Hg2Cl2(s)|Hg(l), Pt +

  •    Cell composed of SHE and saturated calomel electrode

– Pt | H2(g) (1 atm)| H+(aq) (1M) || KCl(aq)(saturated)|Hg2Cl2(s)| Hg(l) ,Pt  +

Cell Reactions:

Steps to Write Cell Reaction of Galvanic Cell:

  • Represent the given galvanic cell with standard convention.
  • The electrode on left side of the representation shows that it is anode and oxidation takes place at this electrode. Write half cell oxidation reaction half-cell reaction for it.
  • The electrode on right side of the representation shows that it is cathode and reduction takes place at this electrode. Write half-cell reduction reaction half-cell reaction for it.
  • Balance above two reactions for electrons for oxidation and reduction reaction.
  • Add the two reactions and obtain net (overall)  cell reaction.


Example – 1: To Write cell reaction of  Pb-Ag cell:

  • Step – 1: Represent the cell conventionally:

– Pb(s) | Pb2+(aq) (1M) || Ag+(aq) (1M)| Ag(s) +

  • Step – 2: Write left hand side half-cell reaction: Pb(s) is on the left side of the representation shows that it is anode and oxidation takes place at Pb(s) electrode.

Pb(s) →  Pb2+(aq) +   2e (Oxidation) … (1)

  • Step – 3: Write right hand side half-cell reaction: Ag(s) is on right side of the representation shows that it is cathode and reduction takes place at Ag(s) electrode.

Ag+(aq)   +   e → Ag(s)    (Reduction)  … (2)

  • Step – 4: Balance the Electrons of above two half cell reactions:

Multiply equation (2) by 2 to balance electrons.

2Ag+(aq)   +   2e → 2Ag(s)    (Reduction)  … (2)

  • Step – 5: Adding equations (1) and (3) we get overall reaction.

Pb(s) +  Ag+(aq) →  Pb2+(aq)    +  Ag(s)

Steps to Find E.M.F. of Galvanic Cell:

  • Represent the given galvanic cell with standard convention.
  • The electrode on left side of the representation shows that it is anode and oxidation takes place at this electrode.
  • The electrode on right side of the representation shows that it is cathode and reduction takes place at this electrode.
  • Obtain standard oxidation potential values from electromotive series for the material of cathode and anode.
  • Use following formula for calculation of e.m.f. of a cell.

EoCell =  Eo(ox/cathode) –   Eo(ox/anode)

OR

EoCell =  Eo(ox/cathode)+   Eo(red/anode)

Example – 1: To find e.m.f. of Daniel Cell :

  • Step – 1: Represent the cell conventionally

 

Electrochemical Cells 02

  • Step – 2: Decide anode and cathode: Pb(s) is on the left side of the representation shows that it is anode and oxidation takes place at Pb(s) electrode. Ag(s) is on the right side of the representation shows that it is cathode and reduction takes place at Ag(s) electrode.
  • Step – 3: Get values of oxidation potential or reduction potential for electrodes from From electrochemical series

Eo(ox/Zn) = 0.76 V and EEo(ox/Cu) =-0.34 V

  • Step – 4: calculate e.m.f of cell:

EoCell = Eo(ox/cathode) –   E(ox/anode)

EoCell =  Eo(ox/Zn) –  Eo(ox/Cu)

EoCell   =      0.76    –  (- 0.34)

EoCell  =  0.76    +   0.34

EoCell   =      1.1 V

 

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