Electrochemical cell B.Sc. 2nd Year Notes

Electrochemical cell B.Sc. 2nd Year Notes

Electrochemical cell

Electrochemical Cell

An electrochemical cell is a device capable of either generating electrical energy from chemical reactions or using electrical energy to cause chemical reactions.
The electrochemical cells which generate an electric current are called Voltaic or Galvanic Cells and those that generate chemical reactions, via electrolysis are called Electrolytic Cells.
Electrochemical cells have two conductive electrodes, called the anode and the cathode. The anode is defined as the electrode where oxidation occurs and the cathode is the electrode where reduction takes place.

Electrmotive Force (emf)

emf stands for electromotive force. A galvanic cell consists of two electrodes each having its own potential. The difference of potential between these two electrodes of a cell causes a current to flow from electrode of higher potential to the electrode of lower potential is called emf of the cell.
emf is expressed in volts. Greater the emf (i.e. potential difference between two electrode) greater the electricity flow and greater is the tendency of the cell redox to occur.
Potential of a cell assembled of two electrodes can be determined from the two individual electrode potentials using this formula-
ΔVcell = Ered,cathode − Ered,anode
or, ΔVcell = Ered,cathode + Eoxy,anode

Potential Difference

It is the amount of work to be done to move a unit positive charge from one point to another point. It is also measured in volts. Potential difference does not remains constant, It is the result.

Q. The difference between the electrode potentials of two electrodes when no current is drawn through the cell is called
a. Cell Potential
b. Cell emf
c. Potential Difference
d. Cell Voltage

Electrode Potential or Half Cell Potential

Electrode potential is the electromotive force of a galvanic cell built from a standard reference electrode and another electrode to be characterized. The electrode potential has its origin in the potential difference developed at the interface between the electrode and the electrolyte.
When a metal(M) is dipped in a solution containing its own ion(M+), a potential is developed between them. This is called Single Electrode Potential(E). It is not possible to measure accurately the absolute value of single electrode potential directly. Only the difference in potential between two electrodes can be measured experimentally with the help of reference electrode whose potential is already known. Cell potential is measured by the following formula-
ECell = ECathode + EAnode
emf of single electrode potential is-
EM/M+n = EoM/M+n + (RT/nF)ln a
When a = 1 then E = Eo. So, Standard electrode potential is the single electrode potential at unit activity.
M ⇌ M+n + ne single elecrode potential for the above equilibrium-
EM/M+n = EoM/M+n + (RT/nF)ln [M]/[M+n]
or, EM/M+n = EoM/M+n + (RT/nF)ln 1/[M+n] (as [M] = 1)
or, EM/M+n = EoM/M+n − (RT/nF)ln [M+n]
or, EM/M+n = EoM/M+n − (0.0591/n)log [M+n]
This is the general expression for electrode potential.
Oxidation Potential and Reduction Potential
The tendency of an electrode to lose electrons is called oxidation potential while tendency of electrode to gain electrons is called reduction potential.

Standard Electrode Potential

The potential of a half-reaction measured against the Standard Hydrogen Electrode under standard conditions is called the standard electrode potential for that half-reaction.
Standard conditions are-
Temperature = 298K
Pessure = 1atm
Concentration of the electrolyte = 1M.
Consider Zn and Hydrogen electrode
Zn ⇌ Zn2+ + 2e
H2 ⇌ 2H+ + 2e
When these two electrodes (two equilibria) are brought into electrical contact using an external wire and a salt bridge, the electrons will be pushed from the zinc equilibrium (electrode) to the hydrogen equilbrium (electrode) with a force of - 0.76V (the negative sign simply indicates the direction of flow - from zinc to hydrogen ions). So, the standard electrode potential of Zn is −0.76 volts and the overall reaction is-
Zn + 2H+ → Zn2+ + H2

Uses of Standard Electrode Potentials

Uses of standard electrode potentials are given below –
1. It is used to measure relative strengths of various oxidants and reductants.
2. It is used to calculate standard cell potential.
3. It is used to predict possible reactions.
4. Prediction of equilibrium in the reaction.

Standard Hydrogen Electrode (SHE)

Standard Hydrogen Electrode is a reference electrode consists of a container, containing solution kept at 298K.
A wire containing Platinum electrode coated with platinum black is immersed in the solution.
Pure hydrogen gas is bubbled in the solution at 1bar pressure.
The potential of standard hydrogen electrode is taken as zero volt at all temperatures.
Standard hydrogen electrode may act as anode or cathode depending upon the nature of the other electrode.
If its acts as anode, the oxidation reaction taking place is
H2(gas) → 2H+(aq) + 2e
If it acts as cathode then the reduction half reaction occurring is
2H+(aq) + 2e → H2(gas)
Standard Hydrogen Electrode

Calomel Electrode

It is a secondary electrode. It has a glass tube fitted with two side tubes. One side tube is used to fill saturated KCl solution and other is connected to another electrode. At the bottom of the main tube, an extra pure Hg is kept which is an contact with Hg2Cl2. A pt-wire is sealed into a glass tube for making electrical contact with the external circuit.
Reduction occurs when it is combined with hydrogen electrode. Its potential depends on the Concentration of KCl solution. Its standard electrode potential is 0.2415V. It can act as an anode or cathode depending on the electrode potential of the coupled electrode. The calomel electrode has following cell reaction-
2Hg + 2Cl → Hg2Cl2 + 2e (Oxidation)
Hg2Cl2 + 2e → 2Hg + 2Cl (Reduction)
Calomel Electrode

Quinhydrone Electrode

The foloowing half cell reaction takes place between quinone and hydroquinone in the presence of H+ ion-
quinhydrone electrode
This electrode is easy to set up and small amount of solution required for measurement. It gives accurate pH between 0 to 9. In alkaline solution, the pH values are not accurate due to ionization and aerial oxidation of QH2.
Determination of pH
For pH determination, it is kept in the solution whose pH is to be determined and connected to a calomel electrode. Hence the cell is represented as-
(−),Pt,Hg|Hg2Cl2(s),KCl sat. solution || Solution of unknown pH,Q|QH2Pt(+)
Therefore, the emf of cell is given by-
Ecell = − 0.2415 − (−0.69994 + 0.0591 pH)
or, Ecell = − 0.2415 + 0.69994 − 0.0591 pH
or, 0.0591 pH = 0.4579 − Ecell
or, pH = (0.4579 − Ecell)/0.0591
By Knowing the value of emf of the cell, pH of the solution can easily be calculated.

Q. The calomel and quinhydrone electrodes are reversible with respect to which ions, respectively ?
a. Cl,H+
b. H+,Cl
c. Hg2+2,OH
d. Hg2+2,H+

Reversible and Irreversible Cells

A cell in which the driving and opposing force differ infinitesimally small amount from each other and the the chemical change taking place in it can be reversed by applying an external force infinitesimally greater than the emf of the cell. A reversible cell should satisfy the following conditions-
1. When the external emf of the cell is infinitesimally greater than the emf of the cell, then current should flow through the cell and the cell reaction of the cell should get reversed.
2. When the external emf of the cell is infinitesimally less than the emf of the cell, then current should flow from the cell.
3. When the external emf of the cell is exactly equal to the emf of the cell, then no current should flow through the cell.
Daniel cell is a reversible cell.
Zn(s) + Cu+2 ⇌ Zn+2 + Cu(s)
The cell reaction can be adjusted to forward or reverse direction by adjusting the external emf of the cell.
Any cell which does not satisfy the condition of reversibility is an Irreversible Cell.
Let us consider the following cell set up by dipping copper and zinc electrodes in a solution of sulphuric acid. When the cell operates, the following reaction take place
Zn + 2H+ → Zn+2 + H2
When this cell is connected to an external emf which is slightly greater than the emf of the cell, not reverse reaction occur but the following reaction occurs
Cu + 2H+ → Cu+2 + H2
So, it is not a reversible cell.

Glvanic Cell or Voltaic Cell

An electrochemical cell that converts the chemical energy of spontaneous redox reactions into electrical energy is known as a galvanic cell or a voltaic cell.
galvanic cell

Important Points of the Galvanic Cell

Oxidation occurs at Anode (−ve Electrode)
Reduction occurs at Cathode (+ve Electrode)
Current flows from Cathode to Anode
Electrone flows from Anode to Cathode
Salt bridge(Contains electrolytes) complete the circuit
Electromotive Force (emf) of the Cell is 1.1 Volt
Oxidation Half Cell Reaction- (at Anode)
Zn(s) → Zn+2 + 2e
Reduction Half Cell Reaction- (at Cathode)
Cu+2 + 2e → Cu(s)
Complete Cell Reaction-
Zn(s) + Cu+2 → Zn+2 + Cu(s)
Cell Representation-
Zn(s)|ZnSO4(aq) || CuSO4|Cu(s)

Reversible Electrodes

The electrodes of a reversible cell are called reversible electrodes.The various types of reversible electrodes are given below-
1. Gas electrode
2. Metal-Metal ion electrode
3. Metal –Metal sparingly salt electrode
4. Redox electrodes

1. Gas electrode

This is developed by bubbling pure and dry gas around a platinised platinum foil dipped in the solution containing ions (of the gas) reversible with respect to the gas bubbled. The gas is adsorbed on the surface of platinum foil and establishes an equilibrium with its ions in the solution. Pt electrode provides electrical contact and also acts as a catalyst. The potential of the gas electrode depends on the pressure of the gas, the concentration of the solution and the pressure.

2. Metal-Metal ion electrode

It consist of a metal rod in contact with a solution of its own ion or ions(cations), with which the electrode is reversible. The potential of this electrode depends on the concentration of the metal ions in the solution and the temperature.
Zn rod dipped into ZnSO4 solution containing Zn+2 ions of concentration C. It is represented as-
Zn(aq)+2 | Zn(s)
The reduction reaction at the electrode is-
Zn(aq)+2 + 2e → Zn(s)

3. Metal –Metal sparingly salt electrode

This electrode consists of a metal coated with one of its sparingly soluble salts and immersed in a solution containing an electrolyte having a common anion as that of the salt. The electrode involves a reversible reaction between the metal and the negative ion to form sparingly soluble salt with the liberation of electrons.
Silver electrode coated with sparingly soluble AgCl dipped in KCl solution with common anion Cl-. This electrode is represented as-
Cl-(aq) | AgCl(s) | Ag(s)
and the reduction reaction is-
AgCl(s) + e → Ag(s) + Cl-(aq)

4. Redox electrodes

This type of electrodes consists of an inert metal like platinum in contact with an aqueous solution of the salt of an element in different oxidation state. The potential of this electrode depends on the ratio of the activities of the metal into in different oxidation state.
Fe+2(aq), Fe+3(aq) | Pt
Reduction Reaction-
Fe+3(aq) + e → Fe+2(aq)

Nernst's Equation or Relation between emf and Activity of a Cell

A galvanic cell is the cell which generates electricity from energy change accompanying a chemical reaction. The emf of such cell arises due to the reaction occuring in the cell. If the following reaction occurs in a reversible chemical cell- Nernst's Equation

Applications of EMF Measurement

1. Determination of pH by using Hydrogen Electrode

The potential of hydrogen electrode in contact with a solution of H+ ions involving a cell reaction-
H+ + e ⇌ 1/2H2 (1 atm) is given by Nernst equation, viz-
E(H+,H2) = Eo(H+,H2) + 2.303RT/F log[H+]
Stansard electrode potential of hydrogen is zero i.e. Eo(H+,H2) = 0
E(H+,H2) = 2.303RT/F log[H+]
or, E(H+,H2) = −0.0591pH at 25oC
Thus, the potential of a hydrogen electrode depends upon the pH of the solution with which it is in contact. and thus, pH can be determined by combining hydrogen electrode with a reference electrode (say calomel electrode).
The cell reaction will be-
Pt, H2(1atm), H+(conc. unknown) || KCl(salt solution),Hg2Cl2(s),Hg
Thus, Ecell = Ecell(Right) − Ecell(Left)
or, Ecell = − 0.2422 −(−0.0591pH)
or, 0.0591pH = 0.2422 − Ecell
or, pH = (0.2422 − Ecell)/0.0591
Thus knowing the value of Ecell, we can easily calculate the pH of the solution.

2. Determination of Solubility and Solubility Product of a Sparingly Soluble Salt

The solubility and solubility product of sparingly soluble salt(AgCl) can be determined from the emf of suitably constructed cells. The sparingly soluble salt is considered as completely ionized even in its saturated solution. So its solubiliy is proportional to its ionic concentration. A few drops of AgNO3 solution is added to 0.1N KCl solution so that the solution is saturated with AgCl and a wire is dipped in this solution. This function as a Ag electrode Ag|Ag+. This electrode is connected to a 0.1N calomel electrode. The experimental emf of the cell comes 0.0494 volts at 298K.
Ag|0.1N KCl || Hg2Cl2|Hg,Pt
The emf of this cell-
E = EAg − Ecalomel
or, 0.0494 = EAg + 0.3338
or, EAg = − 0.3338 + 0.0494
or, EAg = − 0.2844volts
But from the Nernst equation-
EAg = − EoAg − 0.0591 log aAg+
or, log aAg+ = (− 0.799 − 0.2844)/0.0591
or, aAg+ = 2 x 10−9
or [Ag+] = 2 x 10−9
and the activity or concentration of chloride ions in the solution = 0.1 x 0.77 = 0.077
Hence, activity (or solubility) product = aAg+ x aCl
2 x 10−9 x 0.077 = 1.54 x 10−10