Electrochemistry Ipe

Electrochemistry

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VISIT THE FOLLOWING SITE FOR OTHER FILES AND UPDATES IF ANY

http://www.adichemistry.com ELECTROCHEMISTRY * Electrochemistry deals with 1) The processes in which electrical energy is converted to chemical energy. 2) The processes in which chemical energy is converted to electrical energy. 3) Preparation of metals and alloys using electricity.

OM

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Based on electrical conductivity, substances are divided into two types. i) Conductors - which can conduct electricity. e.g. Al, Cu, Fe, Graphite etc... ii) Insulators - which resist the conduction of electricity. e.g. Diamond, Glass, Plastics etc....

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The electrical conductors are again divided into two types as follows; i) Metallic or electronic conductors:- The conductors which conduct the electricity through the electrons. e.g. All metals, Graphite etc... * No chemical reaction occurs during the conduction of electricity. * Conductivity decreases with increase in temperature due to vibrational disturbances.

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ii) Electrolytes :- The substances which furnish oppositely charged ions for the conduction of electricity. e.g :- NaCl, KCl, CH3COOH, HCl etc....... * There is flow of ions towards the oppositely charged electrodes. * During conduction of electricity through electrolytes, oxidation occurs at anode whereas reduction occurs at cathode i.e., a chemical reaction occurs. * The conductivity increases with increase in temperature as the extent of ionization increases. Non electrolytes :- The substances which do not furnish ions for electrical conduction are called non electrolytes. e.g : urea, glucose, sucrose etc..... Strong electrolytes : undergo complete ionization in water e.g. NaCl, KCl, K2SO4, HCl, H2SO4,NaOH, NaNO3 etc..... Weak electrolytes : under go partial ionization in water e.g. HF, CH3COOH, NH4OH, HCOOH etc.... Resistance and ohm's law :- According to Ohm's law, resistance (R) offered by an electrolyte in a solution is proportional to length (l) and inversely proportional to the cross sectional area (a) of electrodes. Resistance (R) 

length (l ) area (a) l a

i.e.,

R

or

R=s.

Where

l a s = specific resistance (resistivity)

Electrochemistry



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a s = R. l

When a = 1 cm2 and l = 1 cm Then s=R Specific resistance (s) : The resistance shown by a material of 1 cm length and 1 cm2 of cross sectional area is called specific resistance (or) resistivity. Units of 's' s=R.

a l

cm 2 cm  ohm . cm. (in CGS)  ohm . m (in SI)

1 a x s l

C=kx Where

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C=

a l

k=

DI TY A

1 R

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C=

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Conductance (C) :- It is the reciprocal of resistance.

OM

or

AN C

 ohm .

1 = specific conductance (conductivity) s l is called cell constant. a

V. A

Cell constant : The quantity,

distance between two electrodes l  area of cross section of electrodes a Specific conductance (k) :- The conductance of 1 cm3 solution is called specific conductance or conductivity. Units of 'k' = ohm-1 cm-1 or mho cm-1 (C.G.S) -1 -1 . = Siemen m or S.m (S I)

cell constant =

Equivalent conductance (  ) :- The conductance of a solution containing one equivalent weight of electrolyte present between two parallel electrodes separated by a unit distance of 1 cm (or) 1 m is called equivalent conductance (  ). k x 1000  N Where N = Normality k = Specific conductance Units of '  '

Ohm-1. cm -1  equivalents . cm -3

Electrochemistry

or

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= ohm-1 cm2 equivalents-1 = Siemen. m2. equivalent -1

(CGS) (S I)

Molar conductance (μ or Λ m ) :- The conductance shown by a solution containing 1 mole of electrolyte present between two parallel electrodes separated by a unit distance of 1 cm or 1m is called molar conductance k x 1000 M M = Molarity ohm-1. cm2. mol-1 Siemen . m2 . mol-1



Units : or

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Factors affecting conductivity of electrolytes : * Strong electrolytes undergo complete ionization and hence show higher conductivities whereas weak electrolytes undergo partial ionization and hence show low conductivities in their solutions. * The ionic mobility decreases with increase in its size and hence conductivity also decreases. * In aqueous solutions the extent of hydration affect the mobility of the ion, which in turn affect the conductivity. Heavily hydrated ions show low conductance values. E.g.,In aqueous solutions Li+ ion (with high charge density) is heavily hydrated than Cs+ ion (with low charge density). Hence hydrated Li+ bigger than hydrated Cs+ . As a result, lithium salts show lower conductivities compared to those of cesium salts in water. * Specific conductance (k) decreases with decrease in concentration of solution as the number of ions per unit volume decreases. * Equivalent or molar conductances increase with decrease in concentration (upon dilution) as the extent of ionization increases * Weak electrolytes undergo complete dissociation at infinite dilution and show the maximum conductance. The equivalent conductance and molar conductance of solutions at infinite dilutions are denoted by   (or)  0 and µ  (or) o respectively..

V. A

Debye - Huckel - Onsagar equation It is possible to determine the equivalent conductances of electrolytes at given concentration by using Debye - Huckel - Onsagar equation.

Λc = Λo  A c Where

Λc = equivalent conductance at given concentration. Λ o = equivalent conductance at infinite dilution. c = concentration A = a constant =

82.4 1

 DT  2

+

8.2 x 105 3

Λo

 DT  2

D = Dipole moment of water A straight line with negative slope is obtained when equivalent conductance values are plotted against different concentrations. The equivalent conductance at infinite dilution can be determined by extending this straight line to zero concentration.

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0

ΛC c

Conductance ratio (  ) : The ratio of the equivalent conductance ( Λc ) at given concentration to that at infinite dilution ( Λ o ) is called conductance ratio (  )

 

Λc Λo

AN C

For weak electrolytes, ' ' is called degree of ionization.

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KOHLRAUSCH LAW OF INDEPENDENT MIGRATION OF IONS The equivalent conductance of an electrolyte at infinite dilution is the algebraic some of equivalent conductances of constituent ions (cations and anions) at infinite dilution. =   electrolyte 



VA R

+

 = equivalent conductance of cation at infinite dilution  = equivalent conductance of anion at infinite dilution

mobilities u o+ and u -o .

  u +o

i.e.,

DI TY A

 and  are also called ionic conductances at infinite dilution. These are proportional to ionic

  u -o

and

and λ +  k u +o λ - = k u -o 'k' is a proportionality constant and its value is equal to one Faraday (96,500 coulombs)



and λ + = F u o+ λ - = F u -o Hence Kohlrausch law can be written as

V. A

or

 electrolyte 

 F u +o  F u -o or

electrolyte 



F  u +o  u -o 

Applications : 1) It is not possible to experimentally determine the equivalent conductances of weak electrolytes at infinite dilution. But by using this law it is possible to calculate the equivalent conductance of weak electrolytes at infinite dilution. The equivalent conductance of acetic acid at infinite dilution can be calculated as follows e.g.

-

3COOH  CH  λ CH 3COO  λ H  -

+

3COONa  CH  λ CH3COO  λ Na  +

 HCl  λ H  λ Cl  +

-

-

Cl  NaCl  λ Na  λ 

+

Electrochemistry

Therefore

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  CH3 COOH   CH3 COONa   HCl   NaCl

2) The degree of ionization of a weak electrolyte (  ) can be calculated by using Kohlrausch law as follows. Degree of ionization ( ) =

equivalent conductance of solution of given concentration ( c ) equivalent conductance of solution at infinite dilution (  )

i.e.,

 

c 

Problems:1) At 250C, the specific conductance of acetic acid of 0.01N concentration is 0.000163 ohm-1 cm-1. What

are 91.0, 426.16 & 126.45 ohm-1. cm2. eq-1. What is  o of CH3COOH ?

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is  c at this concentration ? 2) At 250C, the equivalent conductances of CH3COONa, HCl and NaCl at infinite dilution respectively

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3)  c of acetic acid at 250C is 16.3 ohm-1 cm2 eq-1. The ionic conductances of H+ and CH3COO - are 349.83 & 40.89 ohm-1 cm2 eq-1. What is  of CH3COOH ?

V. A

DI TY A

VA R

ELECTROLYSIS The decomposition of chemical compound in the molten state or in solution state into it's constituent elements under the influence of an applied E.M.F is called electrolysis. Electrolysis is carried out in an electrolytic cell, provided with two electrodes. The electrode connected to negative end of a battery is called cathode and which is connected to positive end is called anode. An electrolyte either in molten state or in solution state is taken into this cell and electrolyzed by applying E.M.F. The cations are reduced at cathode and anions are oxidized at anode.

Battery +

e-

Anode

+

e-

-

cathode

electrolyte

Examples: Electrolysis of molten KCl Following reactions will occur at electrodes when molten KCl is electrolyzed using platinum electrodes. KCl   K + + Clmolten At cathode K + +1e-   K

Reduction

2Cl-   Cl 2  2e -

oxidation

At anode

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Complete redox reaction 2KCl

  2K + Cl 2

Electrolysis of aqueous solution of KCl In this case, hydrogen gas is evolved at cathode as K+ ions can not undergo reduction in presence of water. The electrode reactions are shown below.

KCl(aq)   K + (aq) + Cl-(aq) At cathode 2H 2 O + 2e-   H 2 + 2OH -

At anode 2Cl-   Cl 2  2e  Overall reaction is

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2KCl + 2H 2 O   H 2 + 2KOH + Cl 2 Electrolysis of water Following reactions will occur during the electrolysis of water. At cathode At anode 

 O2 + 4H+ + 4e2H2O 

Net reaction

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 2H2 + 4OH4H2O + 4e- 

DI TY A

 2H2 + O2 2H2O  Hence 4 Faradays of electricity is required to electrolyze 2 moles of water.

V. A

Some important aspects in electrolysis 1) The ions of highly reactive metals like alkaline and alkali earth metals do not undergo reduction in presence of water. These metals have low reduction potentials than water and hence, water molecules undergo reduction by liberating H2 gas at cathode. 2) The metals with low reactivity, like transition metals, can undergo reduction even in presence of water. These metals have higher reduction potentials than water. 3) Oxy anions like SO-24 , NO3- , CO32- , PO3-4 etc. are very stable and can not undergo oxidation at anode. Instead water molecule is oxidized by liberating oxygen gas. E.g., In the electrolysis of aqueous solution of AgNO3 by using platinum electrodes, silver can be deposited at cathode as it is less reactive metal. But NO3- ion cannot be oxidized. Instead, water is oxidized by liberating oxygen gas. AgNO3 (aq)   Ag + (aq)  NO3- (aq)

At cathode:

Ag + (aq)  1e -   Ag

At anode: 2H 2 O   4H + + O 2 + 4e 4) Electrodes which do not participate in the electrochemical reaction are called inert or passived electrodes. Usually platinum and graphite electrodes are used as inert electrodes. But some electrodes may participate in the electrochemical reaction. These are said to be active electrodes. e.g. When aqueous solution of AgNO3 is electrolysed by using copper electrodes, silver is deposited at cathode and dissolved at anode.

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At cathode:

Ag + (aq)  1e -   Ag

Deposition of silver on cathode

At anode:

Ag   Ag + (aq)  1e-

Dissolution of silver anode

In this case, silver electrode is acting as active electrode. FARADAY'S LAWS OF ELECTROLYSIS First law : When an electrolyte, either in molten state or solution state is electrolyzed, the amount of substance deposited or dissolved at electrodes is directly proportional to the quantity of electricity passed through the electrolyte. Mathematically amount of substance (m)  quantity of electricity in coulombs (q) mq

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m = ect

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

e = electrochemical equivalent q = ct c = current in amperes t = time in seconds

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Where But Where

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m  eq

V. A

DI TY A

When q = 1 Coulomb or when c=1 ampere and t = 1 sec Then e=m Hence Electrochemical equivalent :- The amount of substance deposited or liberated or dissolved at electrodes when 1 coulomb of electricity is passed through the electrolyte in molten or solution state is called electrochemical equivalent of that substance. or The amount of substance deposited or liberated or dissolved at electrodes when 1 ampere current is passed through the electrolyte in molten or solution state for one second is called electrochemical equivalent of that substance. Faraday :- The amount of charge transported during the migration of 1 mole of electrons is called Faraday. Faraday (F) = 1.602 x 10-19C x 6.022 x 1023 = 96,500 coulombs Equivalent weight (E) :- The amount of substance formed at electrodes when one Faraday of electricity is passed through the electrolyte. Hence E=Fxe or



Where

e=

E F

m

E=

Ect Eq or m  F F

Atomic weight Valency

Faraday's Second law :- If the same quantity of electricity is passed through different electrolytic cells

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connected in series containing different electrolytic solutions or melts, the amounts of substances deposited or liberated or dissolved are directly proportional to their equivalent weights. mE or

m1 E  1 m2 E2

Illustration :- When three electrolytic cells containing different electrolytes i.e., CuSO4 , AgNO3 and H2SO4 in aqueous solutions are connected in series and same quantity of electricity is passed through them, then from Faraday's second law the ratio of amounts of Cu, Ag, H2 and O2 formed at different electrodes can be given as follows. m Cu : m Ag : m H 2 : m O 2 = E Cu : E Ag : E H 2 : E O 2

EH2 EO 2

mCu E = Cu m O2 E O2

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mO2



etc.......

VA R

or

mH2

DH

or

mCu E  Cu mAg E Ag

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or

+

DI TY A

Battery

-

-

AgNO3

+

-

H2SO4

V. A

CuSO4

+

Problems : 1) A current of 10 amperes is passed through molted AlCl3 for 96.5 sec. Calculate the mass of aluminium deposited at cathode. 2) What is the ratio of weights of Ag and Al deposited at respective cathodes when the same current is passed for same period through aqueous AgNO3 and Al2(SO4)3 solutions. GALVANIC CELLS (OR) VOLTAIC CELLS The device which makes use of a spontaneous redox reaction for the generation of electrical energy is called galvanic cell or voltaic cell or electrochemical cell. Construction of a galvanic cell : A galvanic cell consists of two half cells called single electrodes which are connected to each other. Single electrode: A single electrode consists of a metal or a non metal in contact with their ions. While representing a single electrode, the metal or non metal and its ion are written by separating with a vertical line. e.g: Metal electrodes Zn2+ / Zn Cu2+ / Cu

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Ag+ / Ag Non Metal electrodes 1 H 2 , Pt 2 Cl2, Pt / ClBr2, Pt / BrH+

Daniel cell : This cell contains two half cells divided into two compartments by a porous diaphragm. A zinc rod immersed in ZnSO4 solution in one half cell acts as anode. Another half cell is provided with a copper rod immersed in CuSO4 solution. It acts as cathode. These are connected by a metallic wire externally. When connected, following reactions occur in the two half cells and thus by producing electricity. In Zn half cell (anode) In Cu half cell (cathode)

Zn  Zn+2 + 2eCu2+ + 2e-  Cu

Total reaction

Zn + Cu+2  Zn2+ + Cu

OM

DH

AN C

oxidation Reduction

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Note: In galvanic cell, anode is indicated by negative sign whereas cathode is indicated by positive sign. Reactive resistance

VA R

Ammiter

e-

e-

DI TY A

e-

Anode

Zn rod

Zn-2

Cu+2

Cathode

Cu rod

CuSO4 solution

ions ions

Porous diaphragm

V. A

ZnSO4 solution

Voltaic cell : Daniel cell is modified by connecting two half cells internally by a salt bridge. One half cell contains a zinc rod immersed in ZnSO4 solution and another half cell consist of a copper rod immersed in CuSO4 solution. These two half cells are connected externally by a metallic wire whereas a salt bridge is used to connect them internally. Salt bridge is a U-shaped tube filled with Agar - Agar solution of KCl or NH4NO3. This is used to avoid junction potential. Following reactions occur in this cell. In Zn half cell (anode) Zn  Zn+2 + 2eoxidation In Cu half cell (cathode) Cu2+ + 2e-  Cu Reduction Total reaction

Zn + Cu+2  Zn2+ + Cu

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-

e

Ammeter Reactive resistance e-

ee-

cathode

Anode

salt bridge Electrolyte - I

Electrolyte - II

Anode half cell

Cathode half cell

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Representation of galvanic cell :- According to IUPAC convention the oxidation half cell is written on the left hand side and reduction half cell is written on the right hand side. Two vertical lines are used to indicate the salt bridge. Negative sign is used to indicate the oxidation half cell (anode) and a positive sign is used to indicate the reduction half cell (cathode). e.g : Above galvanic cell can be represented as follows. Zn / Zn2+ // Cu2+ / Cu  Where Zn / Zn2+ = Zn half cell 2+ Cu / Cu = Cu half cell // = salt bridge

DI TY A

Differences between electrolytic and galvanic cells

V. A

Electrolytic cell 1) Electrical energy is converted to chemical energy 2) A non spontaneous reaction is carried out by using electrical energy. 3) Anode is indicated by positive sign whereas cathode is indicated by negative sign.

Galvanic cell 1) Chemical energy is converted to electrical energy. 2) A spontaneous reaction occurs. 3) Anode is indicated by negative sign whereas cathode is indicated by positive sign.

4) Both oxidation and reduction reactions occur 4) Oxidation and reduction reactions occur in in one cell different half cells.

Note: 1) Oxidation occurs at anode and reduction occurs at cathode in both the cells. 2) The flow of electrons is always from anode to cathode in both the cells. Single electrode potential (E) :- The potential difference existing at the surface of contact between metal (or) non metal and its ionic solution is called single electrode potential. Standard electrode potential (E0) :- The potential difference at the surface of contact between metal (or) non metal and its ionic solution at unit concentration and at 250C is called standard electrode potential (E0). Single electrode potential can be written for oxidation half cell or reduction half cell. Both oxidation electrode potential and reduction electrode potential have same magnitude but have opposite signs. e.g : Zn / Zn2+ is oxidation half cell and its E0 value is + 0.762V. Zn2+ / Zn is reduction half cell and its E0 value is - 0.762V. Standard hydrogen electrode (SHE) (or) Normal hydrogen electrode (NHE) In this cell, hydrogen gas at 1atm pressure is continuously bubbled into an acid solution of unit activity ( 1M HCl ) along a platinum rod, which is coated by platinum black.

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The platinum rod is surrounded by a glass tube consisting of two circular holes. The glass tube is immersed in the solution such that these holes are half exposed to air. 1 H  H+ + 1e2 2(g)  (aq) The electrode potential of this cell is arbitrarily taken as Zero.

Eo

H + (1M)

1 H 2 (1atm) 2

= 0 volts

This electrode is taken as primary reference electrode.

H2 (1 Atm)

Pt rod

OM

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Black Pt foil

DH

1M HCl

AN C

hole

half cell

DI TY A

half cell

VA R

EMF of galvanic cell : The potential difference between two different electrodes in the galvanic cell is called EMF of that cell. E Cell = E Reduction - E Oxidation

V. A

E Cell = ER - EL According to IUPAC convention, EMF of a galvanic cell is calculated by substraction of the reduction electrode potential of left hand cell from that of right hand cell. E Cell = E Right E Left Electrochemical series :- The standard reduction electrode potentials of various electrodes are determined relative to that of SHE. These electrodes are arranged in their increasing order of standard reduction potentials in a series called as electrochemical series. In this series, the electrodes with negative Eo values are placed above hydrogen electrode whereas the electrodes with positive Eo values are placed below hydrogen electrode. Applications of electrochemical series 1) The elements with low reduction potential values are good reducing agents and those with high reduction potentials are good oxidizing agents. e.g. 'Li' with very low reduction potential is a reducing agent 'F2' with high reduction potential is an oxidizing agent 2) Metals with high negative potentials can displace metals with low negative potentials or high positive potentials e.g. 'Zn' can displace metals beneath it in the series from their ionic solutions. Zn + CuSO4   Cu + ZnSO4 3) Nonmetals with high reduction potentials are good oxidizing agents and can displace non metals with low reduction potential values from their compounds. e.g. F2 can displace Cl2 from NaCl

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F2 + 2 NaCl   2 NaF + Cl2 4) A proper short hand notation for the galvanic cell can be written. While writing the galvanic cell, the electrode with higher reduction potential is written on the right hand side and considered as cathode (reduction half cell) and another electrode with lower reduction potential is written on the left hand side and considered as anode (oxidation half cell) Problems 1) Construct a galvanic cell containing two electrodes Ag+ / Ag and Zn2+ / Zn calculate the EMF of cell. Write the cell reactions. Eo for Ag+ / Ag = + 0 .799 volts Eo for Zn2+ / Zn = - 0 .762 volts

DH

VA R

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OM

G = - nFEcell n = no. of electrons transferred in redox reaction F = Faraday E = EMF of the cell Under standard conditions  Go = - nFEocell Hence Where

AN C

NERNST EQUATION The electrical energy (nFE cell ) produced in a reversible galvanic cell is equal to the decrease in Gibbs free energy ( G)

Consider the following equilibrium for a redox reaction occurring in a Galvanic cell;

DI TY A

 cC + dD aA + bB 

For above reaction, the equilibrium constant from law of mass action can be written as c

(Kc) =

d

 C  D a b  A   B

V. A

From laws of thermodynamics, the relation between Gibbs free energy and equilibrium constant (Kc) is given as G = G o + RTlnK c

nFE   nFE o + RTlnK c

-E = -E o +

RT ln K c nF

E  Eo 

RT ln K c nF

E  Eo 

2.303RT log K c nF c

d

 C  D  2.303RT EE  . log a b nF  A   B o

E  Eo 

 product of equilibrium conc. of products  2.303RT . log   nF  product of equilibrium conc. of reactants 

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This equation is also applicable to single electrodes. For metal or anode electrodes represented by Mn+ / M Mn+ + ne-  M

Kc =

E = Eo -

2.303RT 1 . log n+ nF [M ]

E = Eo -

0.059 1 . log n+ n [M ]

0.059 . log[M n+ ] n For non metal or cathode electrodes represented by A / An-

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E = Eo +

[A n- ] [A,Pt]

0.059 .log[A n- ] n

E = Eo +

0.059 1 . log nn [A ]

[  (A, Pt)=1]

DI TY A

E = Eo -

VA R

 AnA, Pt + ne- 

Kc =

OM



2.303RT [M] . log n+ nF [M ] But [M] = 1 E = Eo -

AN C



[M] [M n+ ]

V. A

Problems : 1) Calculate EMF of the cell constructed by using electrodes Cl-(0.1 M) / Cl2,Pt and Ni / Ni2+ (0.01 M) write the cell reactions 2) Calculate the potential of single electrode Zn++(aq) / Zn ? (0.1 M) (Eo =- 0.762V) 3) Calculate the electrode potential of the single electrode ? Cu2+ (C = 0.01 M) /Cu ? Eo = 0.337V) 4) Calculate the electrode potential of the single electrode. Ag+ (0.01 M) / Ag ? (Eo = 0.799V) Battery or cell : It is a galvanic cell which is used as source of electrical energy. These are of two types viz., 1) Primary cells, 2) Secondary cells 1) Primary cell It is an electrochemical cell which acts as a source of electrical energy without being previously charged up by an electric current from an external source. In this cell, the reaction occurs only once and it becomes dead after using over a period of time. This cell cannot be reused again as the reaction cannot be reversed. E.g., 1) LACLANCHE CELL It consists of a carbon rod placed in a porous pot containing a mixture of MnO2 and carbon powder. It acts as cathode whereas an amalgamated zinc acts as anode. The porous pot and zinc rod are placed in 20% NH4Cl solution.

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Following reaction occur in this cell At cathode :

2MnO 2  2H 2O  2e    2MnO  OH   2OH  MnO 2  NH 4  e   MnO  OH   NH3 At anode : Zn   Zn 2   2e 

Zn  2MnO2  2H 2 O   Zn 2  2OH  + 2MnO  OH  Some secondary reactions are 2NH 4 Cl  2OH    2NH 3  2Cl   2H 2O

DH

+

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C

OM

-

AN C

Zn 2  2NH 3  2Cl     Zn  NH 3  2  Cl2 It does not maintain steady current for long periods. Its voltage is 1.5 V. This cell cannot be recharged after its use.

Amalgamated Zn

20% NH4Cl

C+MnO 2

mixture

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solution

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porous pot

Glass Jar

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2) DRY CELL It is a modified Laclanche cell. It consists of a cylindrical zinc vessel acting as anode (negative electrode). It is covered with a cardboard and sealed with pitch. A carbon rod acting as cathode (positive electrode) is introduced at the centre of Zn vessel. It is surrounded by a paste of (C+MnO2). The remaining empty space is filled with NH4Cl + ZnCl2 paste. These two pastes are separated by a porous sheet. Following reaction occur in this cell At cathode :

2MnO 2  2H 2O  2e    2MnO  OH   2OH  MnO 2  NH 4  e   MnO  OH   NH3 At anode : Zn   Zn 2   2e 

Zn  2MnO2  2H 2 O   Zn 2  2OH  + 2MnO  OH  Some secondary reactions are 2NH 4 Cl  2OH    2NH 3  2Cl   2H 2O Zn 2  2NH 3  2Cl     Zn  NH 3  2  Cl2 It does not maintain steady current for long periods. Its voltage is 1.5 V. This cell cannot be recharged after its use.

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Carbon rod (cathode)

MnO2+

Zinc cup carbon black (anode) +NH4 Cl paste

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SECONDARY CELL The cell in which electrical energy from an external source is first converted to chemical energy and then made to operate in opposite direction by removing the external source. In this cell, the reaction can be reversed practically. It can be recharged after its discharging. E.g., 1) Acid storage cell (lead accumulator) 2) Alkali storage cell ( Edison battery)

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LEAD ACCUMULATOR Lead accumulator is an acid storage secondary cell which can be charged and discharged for several times. It consists of two lead electrodes Anode ------- made by sponge lead Cathode ----- Lead coated with PbO2 It can be represented as

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Pb H2SO4 PbO2

or

Pb PbSO4(s), H2SO4 (aq), PbSO4 (s), PbO2 (s) Pb

Double Sulphation Theory : According to this theory, proposed by Glasstone and Traube, following reactions occur at electrodes. At Left Hand side electrode (Anode)

Net Reaction

Pb

 

Pb 2+

+

Pb  s 

+

Pb 2+

SO 2-4 SO 2-4

+ 2e  (Oxidation)

 

PbSO 4  s 

 

PbSO 4  s  

2e 

At Right Hand side electrode (cathode) PbO 2

+

2H 2 O

Pb 4+

+

2e 

Pb 2+

+

SO 24 

     

Pb 4+

+

Pb 2+  reduction  PbSO 4  s 

4OH 

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4OH -

Net Reaction

+

4H +

 

4H 2O

 

PbO 2 + 4H + + SO 42- + 2e 

PbSO 4 s  + 2H 2 O

Complete cell reaction for 2 Faradays is

Pb

+

discharge     charge

PbO2 + 2H 2SO4

2PbSO4

+ 2H2 O

In the discharging process, equal amount of water is formed in place of sulfuric acid. After charging the lead accumulator, sulfuric acid is formed again. Hence the specific gravity of H2SO4 is changed during charging and discharging processes. The voltage of the cell is changed from 2.15 V (40% H2SO4) to 1.88 V (5% H2SO4) during the discharging process.

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FUEL CELLS A fuel cell is a galvanic cell in which the chemical energy of fuel-oxidant system is converted directly into electrical energy. In fuel cells, the fuel and oxidant are supplied continuously on to the two different electrodes. The fuel is oxidized at the anode and the oxidant is reduced at cathode. Theoretical efficiency of fuel cells is 100%. Fuels used may be gases or liquids. E.g., H2, alkanes, carbon monoxide, Methanol, Ethanol, hydrazine, formaldehyde etc., Oxygen, air, hydrogen peroxide, nitric acid etc., are used as oxidants. Platinum, porous PVC or PTFE (Teflon) coated with silver, nickel boride, raney nickel etc., can be used as electrodes.

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Examples H2 - O2 Fuel cell In this fuel cell, hydrogen and oxygen gases are continuously bubbled over porous carbon electrodes suspended in concentrated NaOH solution. These electrodes are embedded with finely divided Pt or Pd catalysts. Electricity is produced due to the following reactions occurring at electrodes. -

O 2 (g) + 2H 2O (l) + 4e

At anode:

2 x [ H2 (g) + 2OH (aq)

Overall reaction:

2H2 (g) + O2 (g)

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At cathode:

-

4OH

-

-

2H2O (aq) + 2e ]

2H2O (l)

The heat of formation of water is directly converted to electrical energy. Hydrocarbon - O2 fuel cell In these cells, hydrocarbons are used as fuels. Pure oxygen is used as oxidant. The electrodes are made up of platinum. The electrolyte used is phosphoric acid instead of KOH. Advantages of fuel cells 1) High efficiency of energy conversion process (nearly 100%) 2) Absence of moving parts in the cell eliminate wear & tear problems. 3) Silent operation. 4) Absence of harmful waste products. 5) Pure H2O is formed in the fuel cells used in space crafts and can be used for drinking purposes by space travellers. Gemini and Apollo used these successfully.

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CORROSION The natural tendency of conversion of a metal into its mineral upon interaction with the environment is called as corrosion. E.g. 1) Rusting: Conversion of iron into its oxide (Fe2O3 - Heamatite) 2) Tarnishing: Silver is converted to its sulphide (Ag2S- silver glance) 3) Conversion of copper to its green colored carbonate (malachite) Mechanism of corrosion Corrosion may be a chemical or an electrochemical process. It is considered as anodic dissolution of metal due to oxidation. M   M n+ + neHence the corrosion occurs under the conditions which favor the formation of voltaic cell with the metal acting as anode. Electrochemical corrosion is of two types as follows.

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1. Hydrogen evolution type Metals with more negative reduction potentials than that of hydrogen can undergo oxidation by liberating hydrogen gas from aqueous solutions. The electrode potentials of metal and hydrogen depend on PH, nature of metal and impurity. Hence the corrosion also depends on these conditions. E.g. Pure zinc does not undergo corrosion in salt solution or in neutral conditions. But it undergoes corrosion in presence of Cu as impurity or in 2M acid solutions.

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2. Differential oxygenation type A metal can undergo corrosion if concentration of O2 over the metal surface is not uniform. The portion at which O2 concentration is more acts as cathode and the portion at which the O2 concentration is less acts as anode. Hence the metal undergoes corrosion at points where O2 concentration is less. E.g. When half portion of iron rod is immersed in NaCl solution, the immersed portion undergoes corrosion due to less oxygenation. Factors promoting corrosion 1) Nature of impurity: Corrosion is favored when the impurity is a more cathodic metal. E.g. Cu is more cathodic than Zn and favors the corrosion of Zinc. 2) Concentration of O2: The portion of metal with less oxygen concentration favors corrosion. E.g. When a half portion of iron rod is immersed in NaCl solution, the immersed portion undergoes corrosion due to less oxygenation. 3) Highly conducting solutions favor rapid corrosion. Prevention of corrosion 1) By avoiding the contact with surroundings by painting the surface of metal. 2) By adding another metal which is more anodic than the metal to prevent corrosion. e.g. Zinc is more anodic than iron. Hence zinc is added to iron to prevent corrosion. This is called galvanization. 3) By avoiding the contact of metal with other materials which are good electical conductors. PASSIVITY The state of non reactivity reached after an initial state of reactivity is called passivity. Passivity of metals can be classified into 1) Chemical passivity

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2) Mechanical passivity 3) Electrochemical passivity 1) Chemical passivity : This state is achieved by chemical reaction over metals. E.g. Iron is attacked initially by conc.HNO3 and dissolves in it. But after sometime, the reactions is stopped and iron becomes passive. At this stage iron cannot displace Ag from AgNO3 solution. 2) Mechanical passivity: An invisible thin film of oxide formed over metal surface prevent it from reacting with acids. This is called mechanical passivity. E.g. An invisible layer of PbO2 formed over the surface of Pb makes it passive and hence Pb is not soluble in acids. This type of passivity is also shown by Fe, Co, Ni, Mn etc.,

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3) Electrochemical passivity: During electrochemical processes, the dissolution of metal at anode is stopped after some time. This is called electrochemical passivity. E.g. Zn, Fe, Ni metals acting as anodes dissolve in the solution due to oxidation as follows. Zn --------> Zn2+ + 2eNi --------> Ni2+ + 2eBut they stop dissolving after some time during electrochemical processes.

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Theory of passivity A metal becomes passive due to the formation of invisible thin film of its oxide over the surface which prevents further reaction. Usually a thick layer of oxide is also formed beneath the thin film.