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Chapter, , 12, , Electrochemistry, Electrochemistry is the branch of physical chemistry which deals, with the relationship between electrical energy and chemical changes taking, place in redox reactions, , (v) The anions on reaching the anode give up their electrons and, converted into the neutral atoms., , At anode : A–  A  e  (Oxidation), , Electrolytes and Electrolysis, , (vi) On the other hand cations on reaching the cathode take up, electrons supplied by battery and converted to the neutral atoms., , (1) Definition : “The substances whose aqueous solution undergo, , decomposition into ions when electric current is passed through them are, known as electrolytes and the whole process is known as electrolysis or, electrolytic decomposition.”, , At cathode : B  e   B (Reduction), This overall change is known as primary change and products, formed is known as primary products., , Solutions of acids, bases, salts in water and fused salts etc. are the, examples of electrolytes. Electrolytes may be weak or strong. Solutions of, cane sugar, glycerine, alcohol etc., are examples of non-electrolytes., , The primary products may be collected as such or they undergo, further change to form molecules or compounds. These are called secondary, products and the change is known as secondary change., , (2) Electrolytic cell or Voltameter : The device in which the process, of electrolysis or electrolytic decomposition is carried out is known as, electrolytic cell or voltameter., , (3) Preferential discharge theory : According to this theory “If more, than one type of ion is attracted towards a particular electrode, then the ion, is discharged one which requires least energy or ions with lower discharge, potential or which occur low in the electrochemical series”., , (i) Voltameter convert electrical energy into chemical energy., (ii) The electrode on which oxidation takes place is called anode (or, +ve pole) and the electrode on which reduction takes place is called cathode, (or –ve pole), , The potential at which the ion is discharged or deposited on the, appropriate electrode is termed the discharge or deposition potential, (D.P.)., The values of discharge potential are different for different ions., , (iii) During electrolysis in voltameter cations are discharged on cathode, , The decreasing order of discharge potential or the increasing order, of deposition of some of the ions is given below,, , and anions on anode., (iv) In voltameter, outside the electrolyte electrons flow from anode, to cathode and current flow from cathode to anode., Anode, , Flow of electrons, , For cations : Li , K  , Na  , Ca 2  , Mg2  , Al3  , Zn 2  ,, Fe2 , Ni 2  , H  , Cu 2  , Hg 2  , Ag  , Au 3  ., , Cathode, , Flow of current, , For anions : SO 42  , NO 3 , OH  , Cl  , Br , I  ., , For voltameter, E cell  ve and ΔG  ve., , Table : 12.1 Products of electrolysis of some electrolytes, Electrolyte, Aqueous NaOH, , Electrode, Pt or Graphite, , Product at cathode, 2 H  2e  H 2, , 2OH  , , 1, O 2  H 2 O  2e , 2, , Fused NaOH, , Pt or Graphite, , Na   e   Na, , 2OH  , , 1, O2  H 2 O  2 e , 2, , , , , , Product at anode

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Aqueous NaCl, , Pt or Graphite, , 2 H   2e   H 2, , 2Cl   Cl 2  2e , , Fused NaCl, , Pt or Graphite, , Na   e   Na, , 2Cl   Cl 2  2e , , Aqueous CuSO, , Pt or Graphite, , Cu 2  2e   Cu, , 2OH  , , Aqueous CuSO, , Cu electrode, , Cu 2  2e   Cu, , Cu oxidised to Cu 2  ions, , Dilute H SO, , Pt electrode, , 2 H   2e   H 2, , 2OH  , , Conc. H SO, , Pt electrode, , 2 H   2e   H 2, , Peroxodisulphuric acid (H 2 S 2O8 ), , Aqueous AgNO, , Pt electrode, , Ag   e   Ag, , 2OH  , , Aqueous AgNO, , Ag electrode, , Ag   e   Ag, , Ag oxidised to Ag  ions, , 4, , 4, , 2, , 2, , 4, , 4, , 3, , 3, , (4) Application of electrolysis : Electrolysis has wide applications in, industries. Some of the important applications are, as follows,, (i) Production of hydrogen by electrolysis of water., (ii) Manufacture of heavy water (D2O) ., (iii) The metals like Na, K, Mg, Al, etc., are obtained by electrolysis, of fused electrolytes., (iv) Non-metals like hydrogen, fluorine, chlorine are obtained by, electrolysis., (v) In this method pure metal is deposited at cathode from a, solution containing the metal ions Ag, Cu etc., (vi) Compounds like NaOH, KOH, Na 2CO 3 , KClO3 , white lead,, KMnO4 etc. are synthesised by electrosynthesis method., , (vii) Electroplating : The process of coating an inferior metal with a, superior metal by electrolysis is known as electroplating. The aim of, electroplating is, to prevent the inferior metal from corrosion and to make, it more attractive in appearance. The object to be plated is made the, cathode of an electrolytic cell that contains a solution of ions of the metal to, be deposited., For electroplating, , Anode, , Cathode, , Electrolyte, , With copper, , Cu, , Object, , CuSO 4  dilute H 2 SO 4, , With silver, , Ag, , Object, , K[ Ag(CN )2 ], , With nickel, , Ni, , Object, , Nickel ammonium, sulphate, , With gold, , Au, , Object, , With zinc, , Zn, , Iron objects, , ZnSO 4, , With tin, , Sn, , Iron objects, , SnSO, , K[ Au(CN )2 ], , 4, , Thickness of coated layer : Let the dimensions of metal sheet to be, coated be (a cm  b cm)., Thickness of coated layer  c cm, Volume of coated layer  (a  b  c) cm 3, Mass, , of, , the, , deposited, , substance, ,  Volume  density, ,  (a  b  c)  dg, It E, 96500, Using above relation we may calculate the thickness of coated layer., ,  (a  b  c)  d , , 1, O2  H 2 O  2e , 2, , 1, O2  H 2 O  2e , 2, , 1, O2  H 2 O  2e , 2, , Faraday's laws of electrolysis, The laws, which govern the deposition of substances (In the form of, ions) on electrodes during the process of electrolysis, is called Faraday's laws, of electrolysis. These laws given by Michael Faraday in 1833., (1) Faraday's first law : It states that,, “The mass of any substance deposited or liberated at any electrode, is directly proportional to the quantity of electricity passed.”i.e., W  Q, Where,, W = Mass of ions liberated in gm,, Q  Quantity of electricity passed in Coulombs, = Current in Amperes (I) × Time in second (t),  W  I  t or W  Z  I  t, In case current efficiency () is given, then, , W  ZIt, , , 100, , where, Z  constant, known as electrochemical equivalent (ECE) of, the ion deposited., When a current of 1 Ampere is passed for 1 second (i.e., Q  1 ),, then, W  Z, Thus, electrochemical equivalent (ECE) may be defined as “the mass, , of the ion deposited by passing a current of one Ampere for one second, (i.e., by passing Coulomb of electricity)”. It's unit is gram per coulomb., Coulomb is the unit of electrical charge., 96500 Coulombs  6.023  10 23 electrons = 1 mole electrons., 1 Coulomb , , 6 .023  10 23,  6 .28  1018 electrons,, 96500, , or 1 electronic charge  1.6  10 19 Coulomb., (2) Faraday's second law : It states that,, “When the same quantity of electricity is passed through different, electrolytes, the masses of different ions liberated at the electrodes are, directly proportional to their chemical equivalents (Equivalent weights).” i.e.,, W1, E, Z It E, Z1, E,  1 or 1  1 or,  1, ( W  ZIt), W2 E2, Z 2 It E2, Z 2 E2, Thus the electrochemical equivalent (Z) of an element is directly, proportional to its equivalent weight (E), i.e.,, E  Z or E  FZ or E  96500  Z, , where, F  Faraday constant  96500 C mol 1, So, 1 Faraday = 1F =Electrical charge carried out by one mole of, electrons., 1F = Charge on an electron × Avogadro's number.

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1F = e   N  (1.602  10 19 c)  (6.023  10 23 mol 1 )., , Number of electrons passed, 6.023  10 23, (3) Faraday's law for gaseous electrolytic product For, the gases, we use, It Ve, V, 96500, where, V  Volume of gas evolved at S.T.P. at an electrode, Ve  Equivalent volume = Volume of gas evolved at an electrode at, S.T.P. by 1 Faraday charge, (4) Quantitative aspects of electrolysis : We know that,, one Faraday (1F) of electricity is equal to the charge carried by one mole, Number of Faraday , , (6.023  10 23 ) of electrons. So, in any reaction, if one mole of electrons, are involved, then that reaction would consume or produce 1F of electricity., Since 1F is equal to 96,500 Coulombs, hence 96,500 Coulombs of electricity, would cause a reaction involving one mole of electrons., If in any reaction, n moles of electrons are involved, then the total, electricity, involved in the reaction is given by,, (Q), Q  nF  n  96,500 C, , Thus, the amount of electricity involved in any reaction is related to,, (i) The number of moles of electrons involved in the reaction,, (ii) The amount of any substance involved in the reaction., Therefore, 1 Faraday or 96,500 C or 1 mole of electrons will reduce,, (a) 1 mole of monovalent cation,(b) 1/2mole of divalent cation,, (c) 1/3 mole of trivalent cation, (d) 1/n mole of n valent cations., , (1) Ohm's law : This law states that the current flowing through a, conductor is directly proportional to the potential difference across it, i.e.,, IV, where I is the current strength (In Amperes) and V is the potential, difference applied across the conductor (In Volts), , V, or V  IR, R, where R is the constant of proportionality and is known as, resistance of the conductor. It is expressed in Ohm's and is represented as, . The above equation is known as Ohm's law. Ohm's law may also be, stated as,, “the strength of current flowing through a conductor is directly, or I , , proportional to the potential difference applied across the conductor and, inversely proportional to the resistance of the conductor.”, (2) Resistance : It measures the obstruction to the flow of current ., The resistance of any conductor is directly proportional to the length (l), and inversely proportional to the area of cross-section (a) so that, l, l, R, or R  ρ, a, a, where  (rho) is the constant of proportionality and is called, specific resistance or resistivity. The resistance depends upon the nature of, the material., , Units : The unit of resistance is ohm (). In terms of SI, base unit, is equal to (kgm 2 ) / (s 3 A 2 )., (3) Resistivity or specific resistance : We know that resistance R is, , l, ; Now, if l  1 cm, a  1 cm 2 then R  , a, Thus, resistivity is defined as the resistance of a conductor of 1 cm, R, , Metallic and Electrolytic conductors, All substances do not conduct electrical current. The substances,, which allow the passage of electric current, are called conductors. The best, metal conductors are such as copper, silver, tin, etc. On the other hand, the, substances, which do not allow the passage of electric current through, them, are called non-conductors or insulators. Some common examples of, insulators are rubber, wood, wax, etc., The conductors are broadly classified into two types,, Metallic and electrolytic conductors., Metallic conduction, (i) It is due to the flow of electrons., (ii) It is not accompanied by decomposition, of the substance.(Only physical changes, occurs), (iii) It does not involve transfer of, matter., (iv) Conductivity decreases with increase, in temperature., , Electrolytic conduction, (i) It is due to the flow of ions., (ii), It, is, accompanied, by, decomposition of the substance., (Physical as well as chemical change, occur), (iii) It involves transfer of matter in, the form of ions., (iv) Conductivity increases with, increases in temperature and degree of, hydration due to decreases in viscosity, of medium., , length and having area of cross-section equal to 1 cm 2 ., , Units : The units of resistivity are, ,   R., , a, cm 2,  Ohm, l, cm, ,  Ohm. cm, , Its SI units are Ohm metre ( m ). But quite often Ohm centimetre, ( cm) is also used., , (4) Conductance : It is a measure of the ease with which current, flows through a conductor. It is an additive property. It is expressed as G., It is reciprocal of the resistance, i.e.,, , G, , 1, R, , Units : The units of conductance are reciprocal Ohm (ohm 1 ) or, mho. Ohm is also abbreviated as  so that Ohm 1 may be written as, ,  1 ., According to SI system, the units of electrical conductance is, , The electrolyte may, therefore, be defined as the substance whose, aqueous solution or fused state conduct electricity accompanied by chemical, decomposition. The conduction of current through electrolyte is due to the, movement of ions., On the contrary, substances, which in the form of their solutions or, in their molten state do not conduct electricity, are called non-electrolytes., , Siemens, S (i.e., 1S  1  1 )., , Electrolytic conduction, , conductivity is the conductance of one centimetre cube of a solution of an, electrolyte., , When a voltage is applied to the electrodes dipped into an, electrolytic solution, ions of the electrolyte move and, therefore, electric, current flows through the electrolytic solution. The power of the electrolytes, to conduct electric current is termed conductance or conductivity., , (5) Conductivity : The inverse of resistivity is called conductivity (or, specific conductance). It is represented by the symbol,  (Greek kappa)., The IUPAC has recommended the use of term conductivity over specific, conductance. It may be defined as, the conductance of a solution of 1 cm, length and having 1 sq. cm as the area of cross-section. In other words,, , Thus,  , , 1, , , , Units : The units of conductivity are

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, , 1,  Ohm 1 cm or  1 cm 1, Ohm. cm, –1, , In SI units, l is expressed in m area of cross-section in m 2 so that, the units of conductivity are S m 1 ., (6) Molar conductivity or molar conductance : Molar conductivity is, defined as the conducting power of all the ions produced by dissolving one, , (8) Experimental measurement of conductance, (i) The conductance of a solution is reciprocal of the resistance,, therefore, the experimental determination of the conductance of a solution, involves the measurement of its resistance., (ii) Calculation of conductivity : We have seen that conductivity (), is reciprocal of resistivity ( ) , i.e.,, ,  , , mole of an electrolyte in solution., It is denoted by  (lambda). Molar conductance is related to, specific conductance (  ) as,, , , , , ,   1000, , M, For the solution containing 1 gm mole of electrolyte placed between, two parallel electrodes of 1 sq. cm area of cross-section and one cm apart,, Conductance(G)  Conductivity  Molar conductivity(), , But if solution contains 1 gm mole of the electrolyte therefore, the, measured conductance will be the molar conductivity. Thus,, Molar conductivity()  100  Conductivity, , In other words, ()    V, where V is the volume of the solution in cm 3 containing one gram, mole of the electrolyte., If M is the concentration of the solution in mole per litre, then, , M mole of electrolyte is present in 1000 cm 3, 1 mole of electrolyte is present in , , 1000, cm 3 of solution, M, ,   1000, , Units of Molar Conductance : The units of molar conductance can, be derived from the formula ,, ,   1000, M, of, , , , are, , S cm, , 1, , and, , units, , of, ,  are,, , 3, , cm,  S cm 2 mol 1  S cm 2mol 1, mol, According to SI system, molar conductance is expressed as, S m 2mol 1 , if concentration is expressed as mol m 3 ., (7) Equivalent conductivity : It is defined as the conducting power of, all the ions produced by dissolving one gram equivalent of an electrolyte in, solution., It is expressed as  e and is related to specific conductance as, Λ  S cm 1 , ,   1000, , 1000,  , (M is Molarity of the solution), C, M, where C is the concentration in gram equivalent per litre (or, Normality). This term has earlier been quite frequently used. Now it is, replaced by molar conductance. The units of equivalent conductance are, Ohm 1 cm 2 (gm equiv)1 ., e , , is the conductance of the cell, l is the distance of, , l, The quantity   is called cell constant and is expressed in, a, cm 1 . Knowing the value of cell constant and conductance of the solution,, the specific conductance can be calculated as,, ,   G  Cell constant, i.e., Conductivity  Conductance  Cell constant, , Factors affecting the electrolytic conductance, In general, conductance of an electrolyte depends upon the following, factors,, (1) Nature of electrolyte : The conductance of an electrolyte depends, upon the number of ions present in the solution. Therefore, the greater the, number of ions in the solution the greater is the conductance. The number, of ions produced by an electrolyte depends upon its nature. The strong, electrolytes dissociate almost completely into ions in solutions and,, therefore, their solutions have high conductance. On the other hand, weak, electrolytes, dissociate to only small extents and give lesser number of ions., Therefore, the solutions of weak electrolytes have low conductance., , The molar conductance of strong electrolyte ( HCl, KCl , KNO 3 ) as, , M, , units, , 1l, l,   or   G , R a, a, , (2) Concentration of the solution : The molar conductance of, electrolytic solution varies with the concentration of the electrolyte. In, general, the molar conductance of an electrolyte increases with decrease in, concentration or increase in dilution., , Thus,     Volume in cm 3 containing 1 mole of electrolyte., , The, , a, l, , separation of two electrodes having cross section area a cm 2 ., ,  may be expressed as,, , , , , , where G, , If M is in the units of molarity i.e., moles per litre (mol L1 ), the, , or  , , and   R, ,   , , M, , where, M is the molar concentration., , , , 1, , well as weak electrolytes ( CH 3 COOH , NH 4 OH ) increase with decrease, in concentration or increase in dilution. The variation is however different, for strong and weak electrolytes., The variation of molar conductance with concentration can be, explained on the basis of conducting ability of ions for weak and strong, electrolytes., , For weak electrolytes the variation of  with dilution can be, explained on the bases of number of ions in solution. The number of ions, furnished by an electrolyte in solution depends upon the degree of, dissociation with dilution. With the increase in dilution, the degree of, dissociation increases and as a result molar conductance increases. The, limiting value of molar conductance (0 ) corresponds to degree of, dissociation equal to 1 i.e., the whole of the electrolyte dissociates., Thus, the degree of dissociation can be calculated at any, concentration as,, , , 0, where  is the degree of dissociation,, , , , c, , c is the molar conductance at concentration C and

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0 is the molar conductance at infinite dilution., For strong electrolytes, there is no increase in the number of ions, with dilution because strong electrolytes are completely ionised in solution, at all concentrations (By definition). However, in concentrated solutions of, strong electrolytes there are strong forces of attraction between the ions of, opposite charges called inter-ionic forces. Due to these inter-ionic forces the, conducting ability of the ions is less in concentrated solutions. With, dilution, the ions become far apart from one another and inter-ionic forces, decrease. As a result, molar conductivity increases with dilution. When the, concentration of the solution becomes very-very low, the inter-ionic, attractions become negligible and the molar conductance approaches the, limiting value called molar conductance at infinite dilution. This value is, characteristic of each electrolyte., (3) Temperature : The conductivity of an electrolyte depends upon, the temperature. With increase in temperature, the conductivity of an, electrolyte increases., , Absolute ionic mobility is the mobility with which the ion moves, under unit potential gradient. It's unit is cm sec 1 ., , Absoluteionic mobility , , Kohlrausch's law, (1) Kohlrausch law states that, “At time infinite dilution, the molar, conductivity of an electrolyte can be expressed as the sum of the, , contributions from its individual ions” i.e., m         , where,,   and   are the number of cations and anions per formula unit of, electrolyte respectively and,  and  are the molar conductivities of the, cation and anion at infinite dilution respectively. The use of above equation, in expressing the molar conductivity of an electrolyte is illustrated as,, The molar conductivity of HCl at infinite dilution can be expressed, , Migration of ions, Electricity is carried out through the solution of an electrolyte by, , as,, , HCl   H  H    Cl  Cl  ; For HCl,  H   1 and  Cl   1., , migration of ions. Therefore,, (1) Ions move toward oppositely charged electrodes at different, speeds., (2) During electrolysis, ions are discharged or liberated in equivalent, amounts at the two electrodes, no matter what their relative speed is., (3) Concentration of the electrolyte changes around the electrode, due to difference in the speed of the ions., (4) Loss of concentration around any electrode is proportional to, the speed of the ion that moves away from the electrode, so, , So, HCl  (1  H  )  (1  Cl  ) ; Hence, HCl  H   Cl , (2) Applications of Kohlrausch's law : Some typical applications of the, Kohlrausch's law are described below,, (i) Determination of m for weak electrolytes : The molar, conductivity of a weak electrolyte at infinite dilution (m ) cannot be, determined by extrapolation method. However, m values for weak, electrolytes can be determined by using the Kohlrausch's equation., , Loss around anode, Speed of cation, , Loss around cathode Speed of anion, , CH 3 COOH  CH 3 COONa  HCl  NaCl, , The relation is valid only when the discharged ions do not react with, atoms of the electrodes. But when the ions combine with the material of the, electrode, the concentration around the electrode shows an increase., , Transport number or Transference number, (1) Definition : “The fraction of the total current carried by an ion is, , known as transport number, transference number or Hittorf number may, be denoted by sets symbols like t and t or t and t or n and n ”., +, , –, , c, , a, , c, , a, , From this definition,, Current carried by an anion, ta , Total current passed through the solution, , tc , , Current carried by a cation, Total current passed through the solution, , evidently, ta  tc  1., (2) Determination of transport number : Transport number can be, determined by Hittorf's method, moving boundary method, emf method and, from ionic mobility., (3) Factors affecting transport number, A rise in temperature tends to bring the transport number of, cation and anion more closer to 0.5, (4) Transport number and Ionic mobility : Ionic mobility or Ionic, conductance is the conductivity of a solution containing 1 g ion, at infinite, dilution, when two sufficiently large electrodes are placed 1 cm apart., Ionic mobilities(a or c )  speeds of ions (uaor uc ), , (ii) Determination of the degree of ionisation of a weak electrolyte :, The Kohlrausch's law can be used for determining the degree of ionisation, of a weak electrolyte at any concentration. If  cm is the molar conductivity, of a weak electrolyte at any concentration C and,  m is the molar, conductivity of a electrolyte at infinite dilution. Then, the degree of, ionisation is given by,  c , , –1, , 2, , –1, , -1, , equation), the degree of ionisation at any concentration ( c ) can be, determined., (iii) Determination of the ionisation constant of a weak electrolyte :, Weak electrolytes in aqueous solutions ionise to a very small extent. The, extent of ionisation is described in terms of the degree of ionisation ( ). In, solution, the ions are in dynamic equilibrium with the unionised molecules., Such an equilibrium can be described by a constant called ionisation, constant. For example, for a weak electrolyte AB, the ionisation equilibrium, is, AB ⇌ A   B  ; If C is the initial concentration of the electrolyte AB, in solution, then the equilibrium concentrations of various species in the, solution are, [ AB]  C(1   ), [ A  ]  C and [B  ]  C, Then,, , 2, , Ionic mobility and transport number are related as,, , cm, cm, , , , m (      ), , Thus, knowing the value of cm , and m (From the Kohlrausch's, , Unit of ionic mobility is Ohm cm or V S cm, , a orc  ta or tc  , , Ionic mobility, 96,500, , K, , , , the, , , ionisation, , constant, , [ A ][B ] C  .C , C, , , [ AB], C(1   ) (1   ), 2, , of, , AB, , is, , given, , by,

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We know, that at any concentration C, the degree of ionisation ( ), is given by,  , , cm, , K, , Then,, , / m, C(cm / m )2, [1, ,  (cm, , / m )], , , , C(cm ) 2, m (m, ,  cm ), , ; Thus, knowing, , m and cm at any concentration, the ionisation constant ( K) of the, electrolyte can be determined., , (vii) Like electrolytic cell, in electrochemical cell, from outside the, electrolytes electrons flow from anode to cathode and current flow from, cathode to anode., (viii) For electrochemical cell, Ecell  ve , G  ve ., (ix) In a electrochemical cell, cell reaction is exothermic., (2) Salt bridge and its significance, (i) Salt bridge is U – shaped glass tube filled with a gelly like, substance, agar – agar (plant gel) mixed with an electrolyte like KCl, KNO ,, NH NO etc., (ii) The electrolytes of the two half-cells should be inert and should, not react chemically with each other., (iii) The cation as well as anion of the electrolyte should have same, ionic mobility and almost same transport number, viz., KCl , KNO 3 , NH 4 NO 3 etc., (iv) The following are the functions of the salt bridge,, (a) It connects the solutions of two half - cells and completes the, cell circuit., (b) It prevent transference or diffusion of the solutions from one, half cell to the other., (c) It keeps the solution of two half - cells electrically neutral., (d) It prevents liquid – liquid junction potential i.e. the potential, difference which arises between two solutions when they contact with each, other., (3) Representation of an electrochemical cell, The cell may be written by arranging each of the pair left – right,, anode – cathode, oxidation – reduction, negative and positive in the, alphabetical order as,, Right, Bridge, Left, 3, , 4, , (iv) Determination of the solubility of a sparingly soluble salt : The, solubility of a sparingly soluble salt in a solvent is quite low. Even a, saturated solution of such a salt is so dilute that it can be assumed to be at, infinite dilution. Then, the molar conductivity of a sparingly soluble salt at, infinite dilution (m ) can be obtained from the relationship,, m      , , ........(i), , The conductivity of the saturated solution of the sparingly soluble, salt is measured. From this, the conductivity of the salt ( salt ) can be, obtained by using the relationship,  salt   sol   wate r , where,  water is, the conductivity of the water used in the preparation of the saturated, solution of the salt., , salt , , 1000  salt, Cm, , ........(ii), , From equation (i) and (ii) ;, , Cm , , 1000  salt, ,, (       ), , Cm is the molar concentration of the, , sparingly soluble salt in its saturated solution. Thus, Cm is equal to the, solubility of the sparingly soluble salt in the mole per litre units. The, solubility of the salt in gram per litre units can be obtained by multiplying, Cm with the molar mass of the salt., , Electrochemical or Galvanic cell, “Electrochemical cell or Galvanic cell is a device in which a, spontaneous redox reaction is used to convert chemical energy into, electrical energy i.e. electricity can be obtained with the help of oxidation, and reduction reaction”., (1) Characteristics of electrochemical cell :, important characteristics, of electrochemical cell,, –, , e, , Voltmeter, , Following are the, , e–, , Salt bridge, , Cu cathode, , Zn anode, , Porous plug, , ZnSO4, , CuSO4, , Fig. 12.1, , (i) Electrochemical cell consists of two vessels, two electrodes, two, electrolytic solutions and a salt bridge., (ii) The two electrodes taken are made of different materials and, usually set up in two separate vessels., (iii) The electrolytes are taken in the two different vessels called as, half - cells., (iv) The two vessels are connected by a salt bridge/porous pot., (v) The electrode on which oxidation takes place is called the anode, (or – ve pole) and the electrode on which reduction takes place is called the, cathode (or + ve pole)., (vi) In electrochemical cell, ions are discharged only on the cathode., , 3, , Anode, , Cathode, , Oxidation, , Reductio, n, Positive, , Negative, , (4) Reversible and irreversible cells : A cell is said to be reversible if the, following two conditions are fulfilled, (i) The chemical reaction of the cell stops when an exactly equal, external emf is applied., (ii) The chemical reaction of the cell is reversed and the current, flows in opposite direction when the external emf is slightly higher than, that of the cell. Any other cell, which does not obey the above two, conditions, is termed as irreversible. Daniell cell is reversible but, Zn| H 2 SO 4 | Ag cell is irreversible in nature, (5) Types of electrochemical cells : Two main types of, electrochemical cells have been reported, these are,, (i) Chemical cells : The cells in which electrical energy is produced, from the energy change accompanying a chemical reaction or a physical, process are known as chemical cells. Chemical cells are of two types,, (a) Chemical cells without transference : In this type of chemical, cells, the liquid junction potential is neglected or the transference number is, not taken into consideration. In these cells, one electrode is reversible to, cations while the other is reversible to the anions of the electrolyte., (b) Chemical cells with transference : In this type of chemical cells,, the liquid-liquid junction potential or diffusion potential is developed across, the boundary between the two solutions. This potential develops due to the, difference in mobilities of  ve and  ve ions of the electrolytes., (6) Concentration cells : “A cell in which electrical energy is, , produced by the transference of a substance from a system of high, concentration to one at low concentration is known as concentration cells”., Concentration cells are of two types., (i) Electrode concentration cells : In these cells, the potential, difference is developed between two electrodes at different concentrations, dipped in the same solution of the electrolyte. For example, two hydrogen, electrodes at different gaseous pressures in the same solution of hydrogen, ions constitute a cell of this type.

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Pt, H 2 (pressure p1 ), H 2 (pressure p 2 ) Pt, ;, | H |, Anode, Cathode, 0 .0591, (p ), Ecell , log 1 at 25 o C If p1  p 2 , oxidation occurs at L. H., 2, (p2 ), S. electrode and reduction occurs at R. H. S. electrode., In the amalgam cells, two amalgams of the same metal at two, different concentrations are immersed in the same electrolytic solution., , Some Commercial cell (Batteries), , M (Hg C1 ) | M n | Zn(Hg C 2 ), The emf of the cell is given by the expression,, 0 .0591, C, Ecell , log 1 at 25 o C, n, C2, (ii) Electrolyte concentration cells : In these cells, electrodes are, identical but these are immersed in solutions of the same electrolyte of, different concentrations. The source of electrical energy in the cell is the, tendency of the electrolyte to diffuse from a solution of higher, concentration to that of lower concentration. With the expiry of time, the, two concentrations tend to become equal. Thus, at the start the emf of the, cell is maximum and it gradually falls to zero. Such a cell is represented in, the following manner ( C 2 is greater then C 1 )., , Types of commercial cells : There are mainly two types of, commercial cells,, , One of the main use of galvanic cells is the generation of portable, electrical energy. These cells are also popularly known as batteries. The term, battery is generally used for two or more Galvanic cells connected in series., Thus, a battery is an arrangement of electrochemical cells used as an energy, source. The basis of an electrochemical cell is an oxidation – reduction, reaction., , (1) Primary cells : In these cells, the electrode reactions cannot be, reversed by an external electric energy source. In these cells, reactions occur, only once and after use they become dead. Therefore, they are not, chargeable. Some common example are, dry cell, mercury cell, Daniell cell, and alkaline dry cell, – +, Cu rod, Zn rod, (i) Voltaic cell, , Local action, , Cu, , Polarisation, , Fig. 12.2, , Cathode : Cu rod, Electrolyte : dil. H 2 SO 4, , The emf of the cell is given by the following expression,, Ecell , , At Anode : Zn  Zn 2   2e , , o, , The concentration cells are used to determine the solubility of, sparingly soluble salts, valency of the cation of the electrolyte and transition, point of the two allotropic forms of metal used as electrodes, etc., (7) Heat of reaction in an electrochemical cell :, charge flows out of a cell of emf E, then, , Over all reaction : Zn  Cu 2   Zn 2   Cu, Electrons flow, (ii) Daniel ecell, –, –, , Key, , Ammeter, , Anode (Zn), , e–, +, , Current, , Cathode (Cu), , Salt bridge, , Let n Faraday, , …….(i), G  nFE, Gibbs – Helmholtz equation from thermodynamics may be given as,  G , G  H  T , ,  T  P, , Anode : Zn rod, Emf : 1.08 V, , At cathode : Cu 2   2e   Cu, , 0 .0591, log, e at 25 C, n, C1( L. H .S .), , Cotton Plugs, , Cu2+, , Zn2+, , …….(ii), 1M ZnSO (aq), , Fig. 12.3, , 4, , From equation (i) and (ii) we get,,  (nFE) ,  E ,  nFE  H  T ,   H  nFT T ,  T  P, , P,  E , H  nFE  nFT, ,  T  P, , Cathode : Cu rod, Electrolyte : dil. H 2 SO 4, , Anode : Zn rod, Emf : 1.1 V, , cell is endothermic., ,  E , Case III: When   < 0, then nFE  H , i.e., process inside the,  T , , (Depolariser), , At Anode : Zn  Zn 2   2e , Over all reaction : Zn  Cu 2   Zn 2   Cu, (iii) Lechlanche cell (Dry cell), , +, , Seal, Graphite (cathode), , MnO +C (Depolariser), 2, ,  E , Case I: When   = 0, then H  nFE,  T  P, ,  E , Case II: When   > 0, then nFE  H , i.e. process inside the,  T , , 1M CuSO4 (aq), , At cathode : Cu 2   2e   Cu, ,  E , where ,  = Temperature coefficient of cell,  T  P, , cell is exothermic., , Paste of KOH+ Z, (Electrolyte), , Zn, , Zn | Zn 2  (C1 ) Zn 2  (C2 )| Zn, ||, Anode, Cathode, C 2( R . H . S ), , Glass, , Dil. H2SO4, , M | M n (C1 )| | M n  (C2 )| M, , or, , Zn, , Paste of NH Cl+ZnCl, 4, , –, Zinc anode, , Fig. 12.4, , Cathode : Graphite rod Anode : Zn pot, Electrolyte : Paste of NH 4 Cl  ZnCl2 in starch, Emf : 1.2 V to 1.5 V, At cathode : NH 4  MnO2  2e   MnO(OH )  NH 3, At Anode : Zn  Zn 2   2e , Over all reaction :, , 2

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At cathode : HgO(s)  H 2 O(l)  2e   Hg(l)  2OH (aq), , Zn  NH 4  MnO2  Zn 2   MnO(OH )  NH 3, (iv) Mercury cell, , At Anode :, , Zn (s) , (amalgam), , 20 H (aq )  ZnO(s)  H 2 O(l)  2e , , Over all reaction : Zn(s)  HgO(s)  ZnO(s)  Hg(l), (2) Secondary cells : In the secondary cells, the reactions can be, reversed by an external electrical energy source. Therefore, these cells can, be recharged by passing electric current and used again and again. These, are also celled storage cells. Examples of secondary cells are, lead storage, battery and nickel – cadmium storage cell., , Anode : Zn rod, Emf : 1.35 V, , Cathode : Mercury (II) oxide, Electrolyte : Paste of KOH  ZnO, In charged, , Lead storage cell, , Alkali cell, , –, , +, , Ni(OH), , 2, , –, , +, Glass vessel, , Fe(OH), , 2, , Perforated steel, grid, , PbO, , 2, , Pb, , KOH 20%, + Li(OH), 1%, , dil. H SO, 2, , Positive electrode, Negative electrode, Electrolyte, During charging, , 4, , Perforated lead plates coated with PbO, Perforated lead plates coated with pure lead, dil. H SO, Chemical reaction, At anode : PbSO + 2H + 2e  Pb + H SO, At cathode : PbSO + SO + 2H O – 2e  PbO, 2, , 2, , 4, , 4, , +, , –, , 4, , 2, , ––, , 4, , +, , 4, , 2, , 2, , 4, , 4, , ––, , –, , 4, , 4, , +, , –, , 2, , 2, , 4, , 4, , 2, , 2, , Efficiency, , 4, , Fuel cells, These are Voltaic cells in which the reactants are continuously, supplied to the electrodes. These are designed to convert the energy from, the combustion of fuels such as H 2 , CO, CH 4 , etc. directly into electrical, energy. The common example is hydrogen-oxygen fuel cell as described, below,, In this cell, hydrogen and oxygen are bubbled through a porous, carbon electrode into concentrated aqueous sodium hydroxide or potassium, hydroxide. Hydrogen (the fuel) is fed into the anode compartment where it, is oxidised. The oxygen is fed into cathode compartment where it is, reduced. The diffusion rates of the gases into the cell are carefully regulated, to get maximum efficiency. The net reaction is the same as burning of, hydrogen and oxygen to form water. The reactions are, , At anode : 2[H 2 (g)  2OH  ](aq)  2 H 2O(l)  2e , At cathode : O2 (g)  2 H 2O(l)  4 e   4 OH  (aq), 2 H 2 (g)  O2 (g)  2 H 2O(l), Each electrode is made of porous compressed carbon containing a, small amount of catalyst (Pt, Ag or CoO ) . This cell runs continuously as, long as the reactants are fed. Fuel cells convert the energy of the fuel, , Overall reaction :, , –, , 2, , 2, , During discharging, , 4, , +, , + 2H SO, Specific gravity of H SO increases and when specific gravity, becomes 1.25 the cell is fully charged., Emf of cell: When cell is fully charged then E = 2.2 volt, Chemical reaction, At anode : Pb + SO – 2e  PbSO, At cathode : PbO + 2H + 2e + H SO  PbSO +, 2H O, Specific gravity of H SO decreases and when specific gravity falls, below 1.18 the cell requires recharging., Emf of cell : When emf of cell falls below 1.9 volt the cell, requires recharging., 80%, 2, , –, , 2, , –, , 4, , Perforated steel plate coated with Ni(OH), Perforated steel plate coated with Fe, 20% solution of KOH + 1% LiOH, Chemical reaction, At anode : Ni (OH) + 2OH – 2e  Ni(OH), At cathode : Fe(OH) + 2K + 2e  Fe + 2KOH, Emf of cell : When cell is fully charged then, , E = 1.36, , volt, , Chemical reaction, At anode : Fe + 2OH – 2e  Fe(OH), At cathode : Ni(OH) + 2K + 2e  Ni(OH) +, –, , –, , 2, , +, , –, , 4, , 2, , 2KOH, Emf of cell : When emf of cell falls below 1.1 V it requires, charging., , 60%, , directly into electricity EMF of fuel cell is 1.23 V. This cell has been used for, electric power in the Apollo space programme. The important advantages of, H2 O, fuel cells are, + Cathode, , Anode –, , Porous carbon electrode, , OH–, H2, , O2, Electrolyte, , 12.6, (1) High efficiency : TheFig.fuel, cells convert the energy of a fuel, directly into electricity and therefore, they are more efficient than the, conventional methods of generating electricity on a large scale by burning, hydrogen, carbon fuels. Though we expect 100 % efficiency in fuel cells, so, far 60 – 70% efficiency has been attained. The conventional methods of, production of electrical energy involve combustion of a fuel to liberate heat

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which is then used to produce electricity. The efficiency of these methods is, only about 40%., (2) Continuous source of energy : There is no electrode material to, be replaced as in ordinary battery. The fuel can be fed continuously to, produce power. For this reason, H 2  O 2 fuel cells have been used in space, crafts., (3) Pollution free working : There are no objectionable byproducts, and, therefore, they do not cause pollution problems. Since fuel cells are, efficient and free from pollution, attempts are being made to get better, commercially practical fuel cells., , (3) Types of electrode potential : Depending on the nature of the, metal electrode to lose or gain electrons, the electrode potential may be of, two types,, (i) Oxidation potential : When electrode is negatively charged with, respect to solution, i.e., it acts as anode. Oxidation occurs., , Electrode Potential, , (M) is suspended in a solution of one molar concentration, and the, temperature is kept at 298 K, the electrode potential is called standard, electrode potential, represented usually by E o ”. ‘or’, The standard electrode potential of a metal may be defined as “the, potential difference in volts developed in a cell consisting of two electrodes,, the pure metal in contact with a molar solution of one of its ions and the, normal hydrogen electrode (NHE)”., , M  M n   ne , (ii) Reduction potential : When electrode is positively charged with, respect to solution, i.e. it acts as cathode. Reduction occurs., M n   ne   M, (4) Standard electrode potential : “If in the half cell, the metal rod, , (1) When a metal (M) is placed in a solution of its ions (M ), either, of the following three possibilities can occurs, according to the electrode, ++, , potential solution pressure theory of Nernst., (i) A metal ion M collides with the electrode, and undergoes no, n+, , change., (ii) A metal ion M collides with the electrode, gains n electrons and, gets converted into a metal atom M, (i.e. the metal ion is reduced)., n+, , Standard oxidation potential for any half - cell  – (Standard reduction potential), Standard reduction potential for any half - cell  – (Standard reduction potential), , M n  (aq)  ne   M (s), (iii) A metal atom on the electrode M may lose an electrons to the, , (5) Reference electrode or reference half - cells, It is not possible to measure the absolute value of the single, electrode potential directly. Only the difference in potential between two, electrodes can be measured experimentally. It is, therefore, necessary to, couple the electrode with another electrode whose potential is known. This, electrode is termed as reference electrode or reference half - cells. Various, types of half – cells have been used to make complete cell with spontaneous, reaction in forward direction. These half – cells have been summarised in, following table,, , electrode, and enter to the solution as M n  , (i.e. the metal atom is, oxidised). M (s)  M n  (aq)  ne  ., Thus, “the electrode potential is the tendency of an electrode to lose, or gain electrons when it is in contact with solution of its own ions.”, (2) The magnitude of electrode potential depends on the following, factors,, (i) Nature of the electrode, (ii) Concentration of the ions in solution,, (iii) Temperature., , Table : 12.2 Various Types of Half – cells, Type, , Example, , Q=, , Half – cell reaction, , Reversible to, , Electrode Potential (oxidn), E, , =, Gas ion half - cell, , , , , , 1, H 2 (g)  H  (aq)  e , 2, , Pt(H 2 ) | H (aq), , , Pt(Cl 2 ) | Cl (aq), , [H ], , 1, Cl (aq)  Cl 2 (g)  e , 2, , , Metal – metal ion, half – cell, Metal insoluble salt anion, half – cell, Calomel electrode, , Oxidation – reduction, half – dell, , 1, [Cl  ], , Cl, , E  0.0591 log[ H  ], , , , E 0  0.0591 log[Cl  ], , 0, , Ag | Ag  (aq), , Ag(s)  Ag  (aq)  e , , [ Ag  ], , Ag , , E 0  0.0591 log[ Ag  ], , Ag, AgCl | Cl  (aq), , Ag(s)  Cl  (aq)  AgCl (s)  e , , 1, [Cl  ], , Cl , , E 0  0.0591 log[Cl  ], , Hg, Hg 2 Cl 2 | Cl  (aq), , 2 Hg(l)  2Cl  (aq) , , 1, [Cl  ]2, , Cl , , E 0  0.0591 log[Cl  ], , 1, [OH  ]2, , OH , , E 0  0.0591 log[OH  ], , [Fe 3  ], [Fe 2  ], , Fe 2  , Fe 3 , , Hg 2 Cl 2 (s)  2e, , Metal – metal oxide, hydroxide half - cell, , H, , , , , , Hg, HgO | OH (aq), , , , Hg(l)  2OH  (aq) , , HgO(s)  H 2 O(l)  2e , , Pt | Fe, , 2, , (aq), , , Fe, , 3, , (aq), , 2, , Fe (aq)  Fe 3  (aq)  e , , Cell potential or EMF of the cell, (1) “The difference in potentials of the two half – cells of a cell known as, electromotive force (emf) of the cell or cell potential.”, The difference in potentials of the two half – cells of a cell arises, due to the flow of electrons from anode to cathode and flow of current, from cathode to anode., , Flow of electrons, Anode, , Cathode, Flow of current, , E 0  0 .0591 log, , [Fe 3  ], [Fe 2  ], , (2) The emf of the cell or cell potential can be calculated from the, values of electrode potentials of two the half – cells constituting the cell., The following three methods are in use :, (i) When oxidation potential of anode and reduction potential of, cathode are taken into account, 0, E cell,  Oxidation potential of anode + Reduction potential of, 0, 0, cathode  Eox, (anode )  E red, (cathode)

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(ii) When reduction potentials of both electrodes are taken into, account, 0, Ecell,  Reduction potential of cathode – Reduction potential of anode, , 0, 0, o, o,  E Cathode,  E Anode,  Eright,  Eleft, , cathode , ,  Oxidation potential of anode – Oxidation potential of, , 0, Eox, (anode ) , , [M n (aq)] = the molar concentration of the metal ion in the, solution,, The concentration of pure metal M(s) is taken as unity. So, the, , (iii) When oxidation potentials of both electrodes are taken into, account, o, Ecell, , [M (s)] = the concentration of the deposited metal,, , 0, Eox, (cathode), , (3) Difference between emf and potential difference, , Nernst equation for the, , M n  / M electrode is written as,, , 0, EM n  / M  EM, , n, /M, , 2 .303 RT, 1, log n , nF, [M (aq)], , At 298 K, the Nernst equation for the M n  / M electrode can be, written as,, 0, EM n  / M  EM, , n, /M, , 0.0591, 1, log n , n, [M (aq)], , Emf, , Potential difference, , It is the potential difference, between two electrodes when no, current is flowing in the circuit., , It is the difference of the electrode, potentials of the two electrodes, when the cell is under operation., , For an electrode (half - cell) corresponding to the electrode reaction,, , It is the maximum voltage that the, cell can deliver., , It is always less then the maximum, value of voltage which the cell can, deliver., , The Nernst equation for the electrode is written as,, , It is responsible for the steady, flow of current in the cell., , It is not responsible for the steady, flow of current in the cell., , (4) Cell EMF and the spontaneity of the reaction :, We know, G  nFEcell, Nature of reaction, , ΔG(or ΔG o ), , Ecell (or Eocell ), , Spontaneous, , –, , +, , Equilibrium, , 0, , 0, , Non – spontaneous, , +, , –, , (1) Nernst’s equation for electrode potential, The potential of the electrode at which the reaction,, , At 298 K, the Nernst equation can be written as,, 0, Ehalf  cell  Ehalf,  cell , , RT, [M (s)], ln, nF [M n  (aq.)], , RT [C]c [D]d, ln, nF [ A]a [B]b, , 0, 0, 0, Where, Ecell, .,  Ecathode,  Eanode, , above eq. is called the Nernst equation., Where,, E M n  / M = the potential of the electrode at a given concentration,, , R = the universal gas constant, 8.31 J K mol, , o, or Ecell  Ecell, , , 1, , T= the temperature on the absolute scale,, , 0 .0592, [C]c [D]d, log a b, n, [ A] [B], , It may be noted here, that the concentrations of A, B, C and D, referred in the eqs. are the concentrations at the time the cell emf is, measured., (3) Nernst’s equation for Daniells cell : Daniell’s cell consists of zinc, and copper electrodes. The electrode reactions in Daniell’s cell are,, , Zn(s)  Zn 2  (aq)  2e , , At anode :, , 0, EM, = the standard electrode potential, n, /M, 1, , 2 .303 RT, [C]c [D]d, log a b, nF, [ A] [B], , At 298 K, above eq. can be written as,, , 2 .303 RT, [M (s)], , log n , nF, [M (aq)], , At cathode : Cu 2  (aq)  2e   Cu(s), Net cell reaction :, , Zn(s)  Cu 2  (aq)  Cu(s)  Zn 2  (aq), , Therefore, the Nernst equation for the Daniell’s cell is,, , n = the number of electrons involved in the electrode reaction,, F = the Faraday constant : (96500 C),, , 0 .0591, [Reduced form ], log, n, [Oxidised form], , (2) Nernst’s equation for cell EMF, For a cell in which the net cell reaction involving n electrons is,, aA  bB  cC  dD, The Nernst equation is written as,, , o, or Ecell  Ecell, , , takes place is described by the equation,, , or EM n  / M , , 2 .303 RT, [Reduced form ], log, nF, [Oxidised form], , o, The E cell, is called the standard cell potential., , M n  (aq)  ne   M(s), , 0, EM, n, /M, , 0, Ehalf  cell  Ehalf,  cell , , 0, Ecell  Ecell, , , Nernst's equation, , 0, EM n  / M  EM, , n, /M, , Oxidised form  ne   Reduced form, , 0, Ecdll  Ecell, , , 2 .303 RT, [Cu (s)][Zn 2  (aq)], log, 2F, [ Zn(s)][Cu 2  (aq)]

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Since, the activities of pure copper and zinc metals are taken as, unity, hence the Nernst equation for the Daniell’s cell is,, , 2 .303 RT, [ Zn 2  (aq], log, 2F, [Cu 2  (aq)], , The above eq. at 298 K is,, , 0 .0591, [ Zn 2  (aq], log, V, 2, [Cu 2  (aq)], , 0, For Daniells cell, Ecell,  1.1 V, , (4) Nernst's equation and equilibrium constant, For a cell, in which the net cell reaction involving n electrons is,, aA  bB  cC  dD, The Nernst equation is, 0, ECell  Ecell, , , RT [C]c [D]d, ln, nF [ A]a [B]b, , Table : 12.3 Standard reduction electrode potentials at 298K, .....(i), , At equilibrium, the cell cannot perform any useful work. So at, equilibrium, E Cell is zero. Also at equilibrium, the ratio, , [C]c [D]d  [C]c [D]d , ,  Kc, , [ A]a [B]b  [ A]a [B]b  equilibrium, , Relationship between potential, Gibbs energy, and equilibrium constant, The electrical work (electrical energy) is equal to the product of the, , EMF of the cell and electrical charge that flows through the external circuit, i.e.,, Wmax  nFEcell, , ......(i), , According to thermodynamics the free energy change (G) is, equal to the maximum work. In the cell work is done on the surroundings, by which electrical energy flows through the external circuit, So, Wmax,  G, ......(ii), from eq. (i) and (ii) G  nFEcell, 0, In standard conditions G 0   nFEcell, , Where G 0  standard free energy change, 0, , But E cell, , 2 .303, RT log K c, nF, ,  G 0  nF , , 2.303, RT log K c, nF, , G 0   2.303 RT log Kc or G  G  2.303 RT log Q, G  RT ln K c, 0, , (2.303 log X  ln X ), , Electrochemical series, (1) The standard reduction potentials of a large number of electrodes, have been measured using standard hydrogen electrode as the reference, electrode. These various electrodes can be arranged in increasing or, decreasing order of their reduction potentials. The arrangement of elements, in order of increasing reduction potential values is called electrochemical, series.It is also called activity series, of some typical electrodes., (2) Characteristics of Electrochemical series, (i) 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, , Element, , Li, K, Ba, Sr, Ca, Na, Mg, Al, Mn, Zn, Cr, Fe, Cd, Co, Ni, Sn, Pb, H2, Cu, I2, Hg, Ag, Br2, Pt, Cl2, Au, F2, , Electrode Reaction (Reduction), , Li++ e– = Li, K++ e– = K, Ba+++ 2e = Ba, Sr+++ 2e = Sr, Ca 2   2e   Ca, Na   e   Na, , Mg 2   2e   Mg, Al 3   3e   Al, , Mn+++ 2e = Mn, Zn2+ +2e–=Zn, Cr3++3 e– = Cr, Fe2++ 2e– = Fe, Cd2++2e– = Cd, Co+++ 2e = Co, Ni2++2e– = Ni, Sn2++2e– = Sn, Pb+++ 2e = Pb, 2H++2e– = H2, Cu2++ 2e– = Cu, I2+2e– = 2I–, Hg2++2e– = Hg, Ag++ e– = Ag, Br2+2e– = 2Br–, Pt+++ 2e = Pt, Cl2+2e– = 2Cl–, Au 3++3e– = Au, F2+2e–= 2F–, , Standard Electrode, Reduction, potential E0, volt, –3.05, –2.925, –2.90, –2.89, –2.87, –2.714, , Increasing strength as reducing agent Increasing tendency to, lose electrons, , o, Ecdll  Ecell, , , standard reduction potential indicates that the electrode when joined with, SHE acts as cathode and reduction occurs on this electrode., (ii) 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., (iii) The substances, which are stronger oxidising agents than, H  ion are placed below hydrogen in the series., (iv) The metals on the top (having high negative value 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 nonmetals 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., , Increasing tendency to accept electrons Increasing strength as, oxidising agent, , 0, Ecdll  Ecell, , , 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, , –2.37, –1.66, –1.18, –0.7628, –0.74, –0.44, –0.403, –0.27, –0.25, –0.14, –0.12, 0.00, +0.337, +0.535, +0.885, +0.799, +1.08, +1.20, +1.36, +1.50, +2.87, , (3) Application of Electrochemical series, (i) Reactivity of metals: The activity of the metal depends on its, tendency to lose electron or electrons, i.e., tendency to form cation (M n  ) ., 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,, (a) 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

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Cl2  2 KI  2 KCl  I2, 2 I   I2  2e , , , Cl2  2e  2Cl, , .....(Oxidation), , , .....(Reduction), , [The activity or electronegative character or oxidising nature of the, nonmetal increases as the value of reduction potential increases.], (c) 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., Mn  Mnn  ne , , , , , 2 H  2e  H 2, , .....(Oxidation), .....(Reduction), , 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., (d) 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., , Element :, Na , Zn , Fe, Reduction potential :  2.71,  0.76  0.44, , Alkali and alkaline earth metals are strong reducing agents., (v) Oxidising nature of non-metals : 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., F2 (Fluorine) is a stronger oxidant than Cl 2 , Br2 and I 2 , Cl 2, , (Chlorine) is a stronger oxidant than Br2 and I 2, <-Element :, I2 Br2 Cl 2, F2, <-Reduction potential :,  0.53  1.06  1.36  2.85, , Oxidisingnature increases, Thus, in electrochemical series, , Highest negative reduction potential, or, (Minimum reduction potential), , nature, , (Strongest reducing agent), Top, , Reducing, , (b) 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., (c) Metals like Cu, Ag and Au which lie below hydrogen are less, reactive and do not evolve hydrogen from water., (ii) 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, (a) 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., (b) 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., (c) Weakly electropositive : 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., (iii) Displacement reactions, (a) 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., (b) Displacement of one nonmetal from its salt solution by another, nonmetal: A non-metal higher in the series (towards bottom side), i.e.,, having high value of reduction potential will displace another non-metal, with lower reduction potential, i.e., occupying position above in the series., The non-metal'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., , (iv) 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. (decreasing order of, reducing nature), , Oxidising, nature, , corresponding salts and combine with those substances which accept, electrons., , (Strongest oxidising agent), , Highest positive value of reduction potential, Bottom, (vi) 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., , Ag2 O , 2 Ag , , 1, O2, 2, , BaO , , , No decomposit ion, 2 HgO , 2 Hg  O2 ; Na 2 O  , , Al2 O3 , (vii) 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., 2 NaAg(CN )2  Zn  Na2 Zn(CN )4  2 Ag, , Corrosion, (1) When metals are exposed to atmospheric conditions, they react, with air or water in the environment to form undesirable compounds, (usually oxides). This process is called corrosion. Almost all metals except, the least active metals such as gold, platinum and palladium are attacked by, environment i.e., undergo corrosion. For example, silver tarnishes, copper

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develops a green coating, lead or stainless steel lose their lusture due to, corrosion. Corrosion causes enormous damage to building, bridges, ships, and many other articles made of iron., Thus corrosion is a process of deterioration of a metal as a result of, its reaction with air or water (environment) surrounding it., In case of iron, corrosion is called rusting. Chemically, rust is, hydrated form of ferric oxide, Fe 2 O 3 . xH 2 O . Rusting of iron is, generally caused by moisture, carbon dioxide and oxygen present in air. It, has been observed that rusting takes place only when iron is in contact with, moist air. Iron does not rust in dry air and in vacuum., (2) Factors which affect corrosion : The main factors which affect, corrosion are, More the reactivity of metal, the more will be the possibility of the, metal getting corroded., The impurities help in setting up voltaic cells, which increase the, speed of corrosion, Presence of electrolytes in water also increases the rate of corrosion, , At cathode : H   e   H, The hydrogen atoms on the iron surface reduce dissolved oxygen., 4 H  O2  2 H 2O, , Therefore, the overall reaction, electrochemical cells may be written as,, , The overall rusting involves the following steps,, (i) Oxidation occurs at the anodes of each electrochemical cell., Therefore, at each anode neutral iron atoms are oxidised to ferrous ions., , At anode : Fe(s)  Fe2  (aq)  2e  ., Thus, the metal atoms in the lattice pass into the solution as ions,, leaving electrons on the metal itself. These electrons move towards the, cathode region through the metal., Drop of moisture, 4H +O +4e 2H O, (Cathode), +, , –, , 2, , Rust, , 2, , Fe, +, , 2+, , Fe anode, , 2e, , cathode, , of, , different, , 4 H   O2  4 e   2 H 2O, , (iii) The overall redox reaction may be written by multiplying, reaction at anode by 2 and adding reaction at cathode to equalise number, of electrons lost and gained i.e., , Oxi. half reaction : Fe(s)  Fe2  (aq)  2e  ]  2, , , , , Red. half reaction : 4 H  O2  4 e  2 H 2O, , (E  0.44 V ), (E  1.23V ), , Overall cell reaction : 2 Fe(s)  4 H   O2  2 Fe2  (aq)  2 H 2O, (ECell  1.67 V ), , Presence of CO 2 in natural water increase rusting of iron., (v) When the iron surface is coated with layers of metals more, active than iron, then the rate of corrosion is retarded., A rise in temperature (with in a reasonable limit) increases the rate, of corrosion., (3) Classification of corrosion process : Depending upon the nature, of corrosion, and the factors affecting it, the corrosion may be classified as, follows., (i) Chemical corrosion : Such corrosion, generally takes place when, (a) Reactive gases come in contact with metals at high temperatures, e.g., corrosion in chemical industry., (b) Slow dissolution of metal takes place when kept in contact with, non conducting media containing organic acids., (ii) Bio-chemical corrosion or Bio-corrosion: This is caused by the, action of microorganisms. Soils of definite composition, stagnant water and, certain organic products greatly favour the bio-corrosion., (iii) Electrochemical corrosion : It occurs in a gaseous atmosphere in, the presence of moisture, in soils and in solutions., (4) Mechanism of rusting of iron : Electrochemical theory of rusting., , at, , The ferrous ions are oxidised further by atmospheric oxygen to form, rust., 4 Fe2  (aq)  O2 (g)  4 H 2O  2 Fe2O3  8 H  and, , Fe2O3  xH 2O  Fe2O3 . xH 2O, Rust, , It may be noted that salt water accelerates corrosion. This is mainly, due to the fact that salt water increases the electrical conduction of, electrolyte solution formed on the metal surface. Therefore, rusting becomes, more serious problem where salt water is present., (5) Corrosion protection : Corrosion of metals can be prevented in, many ways. Some commonly used methods are, (i) By surface coating, (a) By applying, oil, grease, paint or varnish on the surface., (b) By coating/depositing a thin layer of any other metal which does, not corrode. For example, iron surface can be protected from corrosion by, depositing a thin layer of zinc, nickel or chromium on it. Copper/brass can, be protected by coating it with a thin layer of tin. Tinning of brass utensils, is a very common practice in our country., (c) By Galvanization : Prevention of corrosion of iron by Zn coating., (ii) By connecting metal to a more electropositive metal : As long as, the more electropositive metal is there, the given metal does not get, corroded. For example, iron can be protected from corrosion by connecting, it to a block/plate of zinc or magnesium. This method of corrosion, protection is called cathodic protection., (iii) By forming insoluble phosphate or chromate coating : Metal, surfaces are treated with phosphoric acid to form an insoluble phosphate., Formation of a thin chromate layer also prevents the corrosion of metals., (iv) Using anti – rust solutions : Solutions of alkaline phosphates, and alkaline chromates are generally used as anti – rust solutions. For, example, iron articles are dipped in boiling alkaline sodium phosphate, solutions, when a protective insoluble sticking film of iron phosphate is, formed., , –, , Flow of, electrons, , Iron, Schematic representation of mechanism of rusting of iron, , Fig. 12.7, , (ii) At the cathodes of each cell, the electrons are taken up by, hydrogen ions (reduction takes place). The H  ions are obtained either, from water or from acidic substances (e.g. CO 2 ) in water, , H 2 O  H   OH  or CO 2  H 2O  H   HCO3, ,  When two or more ions compete at the electrodes then the ion, with higher reduction potential gets liberated at the cathode while, the one with lower reduction potential at the anode., ,  Cell constant is determined with the help of conductivity bridge,

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where a standard solution of KCl is used., ,  If the external EMF is slightly more than the actual EMF, the, current will flow into the cell and reverse reaction takes place., ,  Identification of cathode and anode is done by the use of, galvanometer., ,  KCl / NaCl / NH Cl etc., can not be used in the salt bridge of a cell, 4, , containing silver salt as one of the electrodes as Cl ions form a ppt., of AgCl with silver ion., –, ,  Weston cell is a common example of standard cell. The emf of a, standard cell does not change with temperature., ,  In Appolo moon flights, H2 – O2 fuel cell was the source of energy and, drinking water., ,  Conductivity water is the highly purified water whose on, conductance is very small. It is prepared by the demineralisation of, ordinary water by passing through cation and anion exchange, resins.