Electrochemical cell: A device which converts chemical energy into electrical energy or electrical energy into chemical energy or both is called electrochemical cell. They are of three types. 1. Galvanic cell or voltaic cell A device used to convert chemical energy of a spontaneous redox reaction into electrical energy. Example When a Zinc rod is placed in copper sulfate solution then redox reaction begins resulted in dissolution of Zinc rod & precipitation of copper at bottom while solution becomes hot as chemical energy is converted into heat energy.

Zn0(s) + Cu2+SO42-(aq) Zn2+SO42-(aq) + Cu0(s) -H or Zn0(s) + Cu2+(aq) Zn2+(aq) + Cu0(s) This reaction may be split into two half reactions Zn0(s) Zn2+(aq) + 2e- 2e- + Cu2+(aq) Cu0(s) Oxidation Half reaction Reduction Half reaction So if two separate half cells are set up by placing a Zn rod in Zinc sulfate solution in a beaker and a copper rod in copper sulfate solution in another beaker while two rods are connected by a wire through an ammeter and two solutions are connected by a salt bridge than electrons given out by former half cell will be taken up by latter one hence current will flow.

Zn0(s) Zn2+(aq) + 2e- 2e- + Cu2+(aq) Cu0(s) A salt bridge is a U-shaped tube containing concentrated solution or solidified solution of an electrolyte in agar or gelatins matrix e.g. KCl, KNO3, K2SO4, NH4NO3 etc. while both ends of the U-tube plugged with porous material. Salt bridge allow anionic conductance from one solution to other without mixing of two solutions so when electrons flow in outer circuit wire, the inner circuit is completed by flow of anions from one solution to other through salt bridge. Therefore anions of both half cell solutions & of salt bridge should be identical. By convention, the electrode at which oxidation occurs is anode and the electrode where reduction occur is cathode. Thus Zinc rod serves as anode and copper rod as cathode. In galvanic cell, electrons are released at anode therefore electrons are in excess on it so it is ve electrode/pole/terminal while electrons are consumed at cathode therefore electrons are deficient there so it is +ve electrode/pole/terminal of cell.

As external circuit is closed, following changes take place Electrons flow from anode to cathode as indicated by ammeter. In oxidation Half cell, Zinc oxidized into Zn2+ ions so that Zn rod get dissolved & Zn2+

ions concentration increases in solution. To maintain electrical neutrality of solution, SO42- ions migrate from reduction to oxidation half cell through salt bridge hence solution becomes concentrate in terms of ZnSO4, in oxidation half cell.

In reduction half cell, Cu2+ ions reduced into copper metal so that copper metal cathode grow in size & Cu2+ ions concentration decreases in solution. As SO42- anions migrated to oxidation half cell hence solution becomes diluted in terms of CuSO4, in reduction half cell.

This type of cell called Daniel cell / gravity cell. So combination of any two-different electrodes constitutes a galvanic cell. The one with

greater tendency to go into the solution by liberating electrons, acts as the anode while other one with more tendency to get deposited acts as the cathode e.g. Cu-Ag galvanic cell.

2. Electrolytic Cell A device used to convert electrical energy into chemical energy, to drive reaction in non-spontaneous direction i.e decomposition of electrolyte solution through electricity. Example 1: Electrolysis of acidulated water in an electrolytic cell, using platinum or stainless steel electrodes, resulted in decomposition of water to yield H2 & O2 gases as follow At cathode 4H3O+ + 4e- 2H2(g) + 4H2O At Anode 6H2O O2(g) + 4H3O+ + 4e- Net electrolytic reaction 2H2O 2H2(g) + O2(g) Example 2: Electrolysis of CuCl2 aqueous solution in electrolytic cell while using inert electrodes (like platinum/ stainless steel) resulted in decomposition of CuCl2 into copper metal & chlorine gas so that CuCl2 becomes diluted.

At cathode Cu2+ + 2e- Cu At anode 2Cl- Cl2 + 2e- Overall electrolytic reaction CuCl2 Cu + Cl2 In contrast to galvanic cell, in electrolytic cell reduction takes place at ve electrode and oxidation at +ve electrodes hence they called as cathode & Anode respectively by convention.

3. ELECTROPLATING CELL Its a kind of electrolytic cell in which instead of electrolyte anionic part oxidation at anode, anodic material itself undergoes oxidation. This alteration done by choosing anodic material metal which readily undergoes oxidation while using electrolyte, which consist of same metallic part (cations) as anode and have bulky anionic part to reduce anion mobility. For example electrolysis of CuSO4 solution while using copper metal as anode resulted in dissolution of copper anode and deposition of copper at cathode. So theoretically electrolyte solution concentration remains unchanged in electroplating cell while in electrolysis cell, electrolyte becomes dilute as process proceeds.

Difference b/w Galvanic cell & Electrolytic cell GALVANIC CELL ELECTROLYTIC CELL It converts chemical energy into electrical energy

It converts electrical energy into chemical energy

It drives Redox reaction in spontaneous direction

It drives Redox reaction in non spontaneous direction

Two electrodes may be set up in solution of the same or different electrolytes

Two electrodes are setup in the same electrolytic solution

+ve & -ve electrodes serve as cathode & anode respectively

+ve & -ve electrodes serve as anode & cathode respectively

A salt bridge or a porous membrane separator is used b/w both electrodes

No salt bridge or a porous membrane separator is used

Free energy supplied by cell reaction system is extracted to do work on surrounding i.e. G is -ve

Free energy is supplied in form of electrical work done on the system by the surrounding i.e. G is +ve

Faradays Law of Electrolysis

Cation Cu+1 Cu+2 Fe+2 Fe+3 Atomic mass 63.5 63.5 56 56 Case-I When Conc. is 1M 1M 1M 1M then, a) amount of metal deposited if 48,250 (96,500/2) coulombs charge i.e. half mole of electron

is passed.

31.7 g 15.8 g 14 g 9.3 g

b) amount of metal deposited if 96,500 coulombs charge i.e one mole of electron is passed.

63.5 g 31.7 g 28 g 18.6 g

Case-II When Conc. is 2M 2M 2M 2M then, a) amount of metal deposited if 48,250 (96,500/2) coulombs charge i.e half mole of electron

is passed.

31.7 g 15.8 g 14 g 9.3 g

b) amount of metal deposited if 96,500 coulombs charge i.e one mole of electron is passed.

63.5 g 31.7 g 28 g 18.6 g

Factors affecting amount of metal deposited, are as follow:

i- m Q ii- m Atomic mass iii- m 1 / Valency of cation

as z Atomic mass / Valency of cation iv- m z v- m Concentration of electrolyte First Law: The amount of given metal deposited at an electrode during electrolysis is directly proportional to the quantity of electricity (i.e. charge), which passes through the electrolyte solution". Second Law: "For a constant amount of electricity, the amount of metal deposited is directly proportional to their electrochemical equivalent. m Q

As I = Q / t so Q = I t

m I t m = z I t If I = 1 Ampere t = 1 Second Then m = z

Electrochemical equivalent (z) is that amount of metal, which deposited by 1 ampere current passing for 1 second (i.e one coulomb) through its electrolyte. Faraday : (F) 1 F = Charge on one mole of electron

1 F = Numbers in one mole Charge on an electron 1 F = (6.023 x 1023) (1.602 10-19 C) 1 F = 96,500 coulombs One gram equivalent is the amount of metal deposited by passing one Faraday i.e 96,500 coulombs charge One gram Equivalent = 96,500 Electrochemical Equivalent One gram equivalent weight = Atomic weight Valency of metal

One gram equivalent w.t of Ag+ = 108 = 108 g 1 One gram Equivalent w.t of Cu2+ = 63.5 = 31.7 g 2 One gram Equivalent w.t of Al3+ = 27 = 9 g 3

NUMERICAL 1Q. 0.1978 g of copper is deposited by a current of 0.2 ampere in 50 min. Calculate electrochemical equivalent of copper. Data: m = 0.1978 g I = 0.2 ampere t = 50 min 60 = 3,000 sec z = ? Solution: m = z I t z = m = 0.1978 I t 0.2 3,000 z = 0.003296 g/coulomb 2Q. How much current in amperes will be required to deposit 10g of silver from

AgNO3 solution in 30 min. Data: I = ? m = 10g t = 30 min 60 = 1800 sec Solution: 108g of Ag i.e (1g eq. w.t) is deposited by = 96,500 Coulombs 1g of Ag deposited by = 96,500 Coulombs 108 10g of Ag deposited by = 96,500 10 = 8935.18 coulombs 108 I = Q t I = 8935.18 coulombs = 4.96 Ampere 1800 sec

Electrode potential When a metal or a non-metal comes in contact with its electrolyte solution than two types of reaction are possible for each case. Case I: 1. Metal shows tendency to go into solution as metallic cations by loosing electrons. M Mn+ + ne- 2. Simultaneously metal ions in solution show tendency to get deposited as metal

atoms by gaining electrons. Mn+ + ne- M

Case II: 1. Non metal shows tendency to go into solution as non metal anion by gaining

electrons X + ne- Xn-

2. Simultaneously non metal anions in solution show tendency to oxidize into non metal by loosing electrons.

Xn- X + ne- In both cases, competition exist b/w two reactions lead to dominance of more tend reaction. This net tendency or potential is called electrode potential. Standard electrode potential (E0) The potential (tendency) developed at the interface between electrode and solution to loose or gain electrons when comes in contact with its electrolyte solution of unit activity at 298 K i.e. 25 oC. Standard Reduction potential (EoRed) The tendency of an electrode metal cation to reduce itself by gain of electrons from electrode metal, developed when electrode comes in contact with its 1 M electrolyte solution at 25 0C. In case of gas electrode conditions further specified as gas at 1 atm pressure comes in contact with its 1M electrolyte solution at 250C.

Cu|Cu+

(1M)+e-

+0.521V

0V

2H+(1M) | H2(1atm)+2e-

Or The tendency of an electrode (usually non metal) to reduce itself by gain of electrons to form its anion, developed when this electrode (at 1atm pressure if a gas) comes in contact with its 1 M electrolyte solution at 25 0C.

F2(1atm) | 2F-(1M)

+2e-+2.87V

Cl2(1atm) |2Cl-(1M)

+2e-+1.359V

+ve sign of E0Red shows that reduction is favorable at that electrode system & vice versa

STANDARD OXIDATION POTENTIAL (EOOxd) Tendency or potential of a metal to go into solution as metal cations by loosing electrons, developed when comes in contact with its 1 M electrolyte solution at 25 oC.

-2e-

Ni Ni2+(1M)|

+0.23V

K K+

(1M)|-e-

+2.295V

Or Potential of a non metal anion in solution to oxidize into non metal by loosing electrons, developed when non metal electrode (at 1 atm pressure if a gas) comes in contact with its 1 M electrolyte solution at 25oC.

Cl2(1atm)|2Cl-(1M)-2e-

-1.359V

Br2(1atm)|2Br-(1M)-2e-

-1.08V

+ve sign of Eooxd shows that oxidation is spontaneous at that electrode system & vice versa.

STANDARD HYDROGEN ELECTRODE SCALE FOR ELECTRODE POTENTIAL The absolute value of individual electrode potential can not be measured experimentally b/c no electrode itself can oxidize or reduce until connected with another reference electrode system to maintain electrical neutrality of system. The reference electrode used is the Standard Hydrogen Electrode (SHE) in which H2 gas at 1 atm pressure is bubbled around a platinum electrode in 1M HCl solutions at 25oC. SHE arbitrarily fixed at zero. Hence resulted numerical emf value of cell is assigned as electrode potential of that electrode with respect to SHE. Direction of current flow would determine whether measured potential value is for oxidation or reduction potential. If subject electrode acts as anode than current would flow away from that electrodes & resulted value is for oxidation potential and vice versa. Voltmeters indicator tilled opposite to direction of current flow. According to convention, all electrode potential value should be reported as reduction tendency of electrodes i.e reduction potential. Since the reduction half reaction is just the reverse of oxidation half reaction. Therefore reduction potential is obtained from oxidation potential by simply changing the sign of value.

Reduction potential = -Oxidation potential Oxidation potential = -Reduction potential

Pushing Force

Pulling Force

Oxidation potential Reduction potential Tendency to loose es- Tendency to gain es-

e-

Electromotive force (EMF) of cell Net force which motivate electrons flow from one (oxidizing) electrode to other (reducing) electrode in galvanic cell. As motivation of electron flow cause by difference of potential b/w both electrodes so it is difference of reduction potential, which causes the current to flow from low reduction potential electrode to high reduction potential electrode. Its units are volts & represented by EMF. Net force = Pushing force + Pulling force Ecell = Eoxd of oxidation half reaction + ERed of reduction half reaction but Eoxd of oxidation half reaction = - ERed of oxidation half reaction so Ecell = ERed of reduction half reaction - ERed of oxidation half reaction

Thus Ecell = EMF

Thus emf of a cell is the driving force of cell reaction which thermodynamically correlated as G = -FE Where G = Free energy change accompanying in cell reaction n = Number of electrons transfer per cell reaction F = Faraday constant = 96,500 coulombs/ mole E = EMF If G is ve then cell reaction would be spontaneous & if G is +ve then cell reaction would be non spontaneous. As n & F are always positive, sign of G now depends upon EMF. So if EMF is +ve then cell reaction would be feasible and spontaneous but if its ve then cell reaction would be non spontaneous and would not be feasible. EMF of galvanic cell when the reactants and products of cell reaction are taken as unit concentration at 298 K temperature and 1 atmospheric pressure is called standard EMF of cell. The EMF of a cell can not be measured accurately by using a voltmeter b/c 1. Voltmeter draw some current from the cell thereby alters the concentration of

electrolyte which cause change in emf of cell. 2. A part of the EMF will be utilized to overcome the internal resistance of the cell during

flow of current. so potentiometer method is used in which emf of a test is opposed by impressed emf from an external source of emf, and measurements are made when there is no net flow of current in the circuit.

Weight

Anode

Cathode

ELECTROCHEMICAL SERIES OR EMF SERIES A series comprises different electrode systems, arranged in descending order of their standard reduction potential at SHE scale.

Electrochemical series with standard reduction potential of redox systems at 25 0C Reduction half-reaction E0 cell Reduction half-reaction E0 cell F2 (g) + 2e- 2F- +2.87 Cu+ + e- Cu + 0.521 O3 (g) + 2H+ + 2e- O2 + H2O +2.07 Fe(CN)63- + e- Fe(CN)64- +0.356 CO3+ + e- CO2+ +1.82 Cu2+ + 2e- Cu +0.337 H2O2 + 2H+ + 2e- 2H2O +1.77 Hg2Cl2 + 2e- 2Hg + 2Cl- +0.2680 Ce4+ + e- Ce3+ (In HClO4 solution) +1.70 AgCl + e- Ag+ + Cl- +0.2224 MnO4- + 4H+ + 3e- MnO2 + 2H2O +1.69 Cu2+ + e- Cu+ +0.153 2HClO + 2H+ + 2e- Cl2(g) + 2H2O +1.63 Sn4+ + 2e- Sn2+ (in HCl solution) +0.14 Ce4+ + e- Ce3+ (in HNO3 solution) +1.60 S + 2H+ + 2e- H2S +0.14 2HBrO + 2H+ + 2e- Br2 + 2H2O +1.60 S4O62- + 2e- 2S2O32- +0.09 MnO4- + 8H+ + 5e- Mn2+ + 4H2O +1.51 2H+ + 2e- H2 0.0000 2BrO3 + 12H+ + 10e- Br2 + 6H2O +1.50 Pb2+ + 2e- Pb -0.126 Mn3+ + e- Mn2+ +1.49 Sn2+ + 2e- Sn -0.140 Ce4+ + e- Ce3+ (in H2SO4 Solution) +1.44 Ni2+ + 2e- Ni -0.23 Cl2 + 2e- 2Cl- +1.359 Co2+ + 2e- Co -0.28 Cr2O72- + 14H+ + 6e- 2Cr3+ + 7H2O +1.33 Cr3+ + e- Cr2+ (in HCl solution) -0.38 Ce4+ + e- Ce3+ (in HCl solution) +1.28 Cd2+ + 2e- Cd -0.402 MnO2 + 4H+ + 2e- Mn2+ + 2H2O +1.23 Fe2+ + 2e- Fe -0.440 O2 + 4H+ + 2e- 2H2O +1.229 2CO2 + 2H+ + 2e- H2C2O4 -0.49 ClO4- + 2H+ + 2e- ClO3- + H2O +1.19 Cr3+ + 3e- Cr -0.74 2IO3- + 12H+ + 10e- I2 + 6H2O +1.19 Zn2+ + 2e- Zn -0.7628 Br2 (g) + 2e- 2Br- +1.087 SO42- + H2O + 2e- SO32- + 2OH+ -0.93 N2O4 + 2H+ + 2e- 2HNO2 +1.07 Mn2+ + 2e- Mn -0.190 NO3- + 3H+ + 2e- HNO2 + H2O +0.94 Al3+ + 3e- Al -1.66 2Hg2+ + 2e- Hg22+ +0.907 H2 + 2e- 2H- -2.25 2NO3 + 4H+ + 2e- N2O4 + H2O +0.80 Mg2+ + 2e- Mg -2.37 Ag+ + e- Ag +0.7994 Na+ + e- Na -2.7 Hg22++ 2e- 2Hg +0.792 Ca2+ + 2e- Ca -2.87 Fe3+ + e- Fe2+ +0.771 Sr2+ + 2e- Sr -2.89 O2 + 2H+ + 2e- H2O2 +0.69 Ba2+ + 2e- Ba -2.90 H3AsO4 + 2H+ + 2e- HAsO2 + 2H2O +0.56 K+ + e- K -2.925 I3- + 2e- 3I- +0.545 Rb+ + e- Rb -2.93 I2- + 2e- 3I- +0.536 Li+ + e- Li -3.03

Application of ECS 1. To predict the relative ease of oxidation or reduction of various elements. In ECS, a positive value indicates reduction tendency and a negative value indicates oxidation tendency with respect to hydrogen. So a substance (element or ion) with greater electrode potential values in ECS means it will easily reduce or it will be stronger oxidation agent or weaker reducing agent & vice versa. Example # 01 In ECS F2 has +2.87 value and Li+ has 3.03 value, So F2 reduced more easily compared to Li+ ion which reduced with greater difficultly or F2 is strong oxidizing agent compared to Li+ ion which is weak oxidizing agent or F2 oxidize more difficulty compared to Li+ ion which oxidize more feasibly or F2 is weak reducing agent compared to Li+ ion which is strong reducing agent. 2. To predict the spontaneity of a redox reaction. A reaction is spontaneous if free energy is decreased in that reaction i.e. G is ve but

G = -FE where E = EMF As & F has +ve value

So G is negative only if EMF is positive. Hence reaction is spontaneous if EMF is positive & vice versa 3. To calculate standard EMF of any electrochemical cell EMF = EoRed of reduction half cell + Eooxd of oxidation half cell As Eooxd = -EoRed EMF = EoRed of reduction half cell - EoRed of oxidation half cell 4. To predict whether a metal reacts with acid to give hydrogen gas or not. A metal would react with an acid to liberate hydrogen gas if emf of that reaction has +ve value.

M + H+ M+ + H2

[oxidation half reaction] M M+ + e- E ooxd = x EoRed = x [Reduction half reaction] H+ + e- H2 EoRed = 0 EMF = EoH - EoM EMF = - EoM

So EMF is only +ve if EoM is ve. Thus all metals laying below hydrogen in ECS can react with acid to give hydrogen gas. Moreover metal with highly ve value, would have highly +ve reaction emf thus possesses high reactivity & liberates H2 gas vigorously & vice versa.

5. To predict whether metal is able to replace another metal from its solution. If a metal Ma is dipped in a solution of another metal Mb then if Ma replaces Mb from solution then reaction will be

Ma + Mbn+ Ma n+ + Mb

[Oxidation half reaction] i) Ma Man+ + ne- [Reduction half reaction] ii) Mbn+ + ne- Mb

So EMF = EoRed of Mb - EoRed of Ma

Therefore if EoRed of Mb > EoRed of Ma then EMF would have +ve sign and above reaction would occur feasibly that is lower electrode potential metal Ma would replace high electrode potential metal Mb. But if EoRed of Mb < EoRed of Ma then EMF would have ve sign & above reaction would not feasible that is high electrode potential metal Ma would not replace low electrode potential metal Mb e.g. Electrode potential value for Ag = +0.79 V, Cu = +0.15 V & Zn = -0.76 V so Zn can displace Cu2+ from CuSO4 solution and Cu can displace Ag+ from AgNO3 solution.