NEET Chemistry Notes - Chapter 12 Electrochemistry PDF

60 Chapter E3 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 ID (v) The anions on reaching the anode give up their electrons and converted into the neutral atoms...

60 Chapter E3 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 ID (v) The anions on reaching the anode give up their electrons and converted into the neutral atoms. Electrolytes and Electrolysis At anode : A–  A  e  (Oxidation) (vi) On the other hand cations on reaching the cathode take up electrons supplied by battery and converted to the neutral atoms. U (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.” D YG At cathode : B  e   B (Reduction) 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. (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. (i) Voltameter convert electrical energy into chemical energy. U (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) (iii) During electrolysis in voltameter cations are discharged on cathode ST (iv) In voltameter, outside the electrolyte electrons flow from anode to cathode and current flow from cathode to anode. Flow of electrons 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. (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”. 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. The decreasing order of discharge potential or the increasing order of deposition of some of the ions is given below, and anions on anode. Anode This overall change is known as primary change and products formed is known as primary products. 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 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 Aqueous AgNO Ag electrode Ag   e  Ag 60 4 2 4 2 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). D YG (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 Ag Object K[ Ag(CN )2 ] Ni Object Nickel ammonium sulphate With silver With nickel Au Object U With gold K[ Au(CN )2 ] With zinc Zn Iron objects ZnSO 4 With tin Sn Iron objects SnSO 4 ST 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 Ag oxidised to Ag  ions 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 U (vi) Compounds like NaOH, KOH, Na 2CO 3 , KClO3 , white lead, KMnO4 etc. are synthesised by electrosynthesis method. 1 O2  H 2 O  2e  2 ID (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. 2OH  1 O2  H 2 O  2e  2 E3 4 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. (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 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 ID 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. 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  60 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  (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) E3 1F = e   N  (1.602  10 19 c)  (6.023  10 23 mol 1 ). D YG 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. 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. U 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) ST (iii) It does not involve transfer of matter. (iv) Conductivity decreases with increase in temperature. (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 U Metallic and Electrolytic conductors is equal to (kgm 2 ) / (s 3 A 2 ). 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 –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,  and   R     where G M a l 1l l   or   G  R a a is the conductance of the cell, l is the distance of separation of two electrodes having cross section area a cm 2. where, M is the molar concentration. If M is in the units of molarity i.e., moles per litre (mol L1 ), the l The quantity   is called cell constant and is expressed in a   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, 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 ID Conductance(G)  Conductivity  Molar conductivity() E3  may be expressed as,  1  60 1  Ohm 1 cm or  1 cm 1 Ohm. cm  But if solution contains 1 gm mole of the electrolyte therefore, the measured conductance will be the molar conductivity. Thus, Molar conductivity()  100  Conductivity (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. U In other words, ()    V In general, conductance of an electrolyte depends upon the following factors, D YG 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 Thus,     Volume in cm 3 containing 1 mole of electrolyte. or     1000 M Units of Molar Conductance : The units of molar conductance can be derived from the formula , The   1000 U  M units of  are S cm 1 and units of  are, 3 ST cm Λ  S cm 1   S cm 2 mol 1  S cm 2mol 1 mol According to SI system, molar conductance is expressed as 2 S m mol 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   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  (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. The molar conductance of strong electrolyte ( HCl, KCl , KNO 3 ) as 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 Absoluteionic mobility  Ionic mobility 96,500 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, (1) Ions move toward oppositely charged electrodes at different speeds. D YG 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 Current carried by a cation Total current passed through the solution U evidently, ta  tc  1. ST (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 ) Unit of ionic mobility is Ohm cm or V S cm –1 2 –1 -1 (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 U 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 – So, HCl  (1  H  )  (1  Cl  ) ; Hence, HCl  H   Cl  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 + HCl   H  H    Cl  Cl  ; For HCl,  H   1 and  Cl   1. ID (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 as, E3 Electricity is carried out through the solution of an electrolyte by migration of ions. Therefore, 2 Ionic mobility and transport number are related as, a orc  ta or tc   under unit potential gradient. It's unit is cm sec 1. The molar conductivity of HCl at infinite dilution can be expressed Migration of ions tc  Absolute ionic mobility is the mobility with which the ion moves 60 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. CH 3 COOH  CH 3 COONa  HCl  NaCl (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  cm cm    m (      ) Thus, knowing the value of cm , and m (From the Kohlrausch's 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, the ionisation constant [ A  ][B  ] C .C  C 2 K   [ AB] C(1   ) (1   ) of AB is given by, We know, that at any concentration C, the degree of ionisation ( ) is given by,   cm / m K Then, 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 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 salt  1000  salt Cm........(ii) From equation (i) and (ii) ; 1000  salt , (       ) Cm is the molar concentration of the U Cm  ID 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. D YG 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 ST U Voltmeter Zn anode E3 (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 3 60 4 Following are the e– Salt bridge Cu cathode 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. 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. M | M n (C1 )| | M n  (C2 )| M (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 Dil. H2SO4 Local action Cu Polarisation Zn Fig. 12.2 Zn | Zn 2  (C1 ) Zn 2  (C2 )| Zn || Anode Cathode Cathode : Cu rod Electrolyte : dil. H 2 SO 4 The emf of the cell is given by the following expression, C 0.0591 log 2(R. H.S ) e at 25 C n C1( L. H.S.) Anode : Zn rod Emf : 1.08 V At cathode : Cu 2   2e  Cu At Anode : Zn Zn 2   2e  o Over all reaction : Zn  Cu 2  Zn 2   Cu Electrons flow (ii) Daniel ecell – U Ecell  Types of commercial cells : There are mainly two types of commercial cells, ID or 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. 60 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 ). Some Commercial cell (Batteries) E3 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. D YG 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 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 – Cotton Plugs Cu2+ Zn2+ …….(ii) From equation (i) and (ii) we get, U  (nFE)   E   nFE  H  T    H  nFT T   T  P  P ST  E  H  nFE  nFT   T  P 1M ZnSO (aq) Fig. 12.3 4 Cathode : Cu rod Electrolyte : dil. H 2 SO 4 Anode : Zn rod Emf : 1.1 V 1M CuSO4 (aq) (Depolariser) At cathode : Cu 2   2e  Cu At Anode : Zn Zn 2   2e  Over all reaction : Zn  Cu 2  Zn 2   Cu (iii) Lechlanche cell (Dry cell) +  E  where   = Temperature coefficient of cell  T  P Seal Graphite (cathode) MnO +C (Depolariser) 2  E   = 0, then H  nFE  T  P Paste of NH Cl+ZnCl Case I: When   E  Case II: When   > 0, then nFE  H , i.e. process inside the  T  cell is endothermic.  E  Case III: When   < 0, then nFE  H , i.e., process inside the  T  cell is exothermic. 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 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 Lead storage cell Alkali cell 60 In charged – + Glass vessel dil. H SO 2 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 4 + – 4 2 –– 4 – 4 2 4 2 U + 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 During discharging 4 D YG 2 –– – 4 4 + – 2 2 2 Efficiency Fuel cells 4 4 4 ST U 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 4 KOH 20% + Li(OH), 1% 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 ID 2 4 + 2 Perforated steel grid E3 2 Pb 2 Fe(OH) PbO Positive electrode Negative electrode Electrolyte During charging Ni(OH) – + 4 – 2 4 + – 2 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 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)”. 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+ M n  (aq)  ne   M (s) (iii) A metal atom on the electrode M may lose an electrons to the electrode, and enter to the solution as M n , (i.e. the metal atom is Standard oxidation potential for any half - cell  – (Standard reduction potential) Standard reduction potential for any half - cell  – (Standard reduction potential) (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, U ID 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. 60 ++ M n   ne   M (4) Standard electrode potential : “If in the half cell, the metal rod E3 (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 M  M n   ne  (ii) Reduction potential : When electrode is positively charged with respect to solution, i.e. it acts as cathode. Reduction occurs. Table : 12.2 Various Types of Half – cells Gas ion half - cell Example Half – cell reaction D YG Type Pt(H 2 ) | H  (aq) Metal – metal ion half – cell Metal insoluble salt anion half – cell Oxidation – reduction half – dell 1 Cl  (aq) Cl 2 (g)  e  2 [H  ] H  E 0  0.0591 log[ H  ] 1 [Cl  ] Cl  E 0  0.0591 log[Cl  ] 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, HgO | OH  (aq) ST Metal – metal oxide hydroxide half - cell Electrode Potential (oxidn), E Ag | Ag  (aq) U Calomel electrode Reversible to = 1 H 2 (g) H  (aq)  e  2 Pt(Cl 2 ) | Cl  (aq) Q= Pt | Fe 2  , Fe 3  (aq) (aq) Hg 2 Cl 2 (s)  2e  Hg(l)  2OH  (aq) HgO(s)  H 2 O(l)  2e  Fe 2  (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  Oxidation potential of anode + Reduction potential of E cell 0 0 cathode  Eox (anode )  E red (cathode) (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  Eright  Eleft  E Cathode  E Anode [M (s)] = the concentration of the deposited metal, [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  Oxidation potential of anode – Oxidation potential of 0 0 cathode  Eox (anode )  Eox (cathode) 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)] (3) Difference between emf and potential difference 60 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)] 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. Nature of reaction Ecell (or Eocell ) Spontaneous – + Equilibrium 0 0 D YG Non – spontaneous + – Nernst's equation (1) Nernst’s equation for electrode potential The potential of the electrode at which the reaction, M n  (aq)  ne  M(s) takes place is described by the equation, RT [M (s)] ln n  nF [M (aq.)] U 0  EM n  / M  EM n /M 2.303 RT [M (s)]  log n  nF [M (aq)] ST or EM n  / M  0 EM n /M 0 Ehalf  cell  Ehalf  cell  above eq. is called the Nernst equation. Where, E M n  / M = the potential of the electrode at a given concentration, At 298 K, the Nernst equation can be written as, 0 Ehalf  cell  Ehalf  cell  0.0591 [Reduced form ] log n [Oxidised form] 1 R = the universal gas constant, 8.31 J K mol (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, 0  Ecell  Ecell RT [C]c [D]d ln nF [ A]a [B]b 0 0 0 Where, Ecell.  Ecathode  Eanode o The E cell is called the standard cell potential. o or Ecell  Ecell  1 T= the temperature on the absolute scale, 2.303 RT [C]c [D]d log a b nF [ A] [B] At 298 K, above eq. can be written as, o or Ecell  Ecell  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 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), 2.303 RT [Reduced form ] log nF [Oxidised form] U ΔG(or ΔG o ) Oxidised form  ne  Reduced form ID (4) Cell EMF and the spontaneity of the reaction : We know, G  nFEcell E3 Emf 0 Ecdll  Ecell  2.303 RT [Cu (s)][Zn 2  (aq)] log 2F [ Zn(s)][Cu 2  (aq)] 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 RT [C]c [D]d ln nF [ A]a [B]b Table : 12.3 Standard reduction electrode potentials at 298K.....(i) Element [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 Li K Ba Sr Ca Li++ e– = Li K++ e– = K Ba+++ 2e = Ba Sr+++ 2e = Sr ID At equilibrium, the cell cannot perform any useful work. So at equilibrium, E Cell is zero. Also at equilibrium, the ratio Na Mg Al D YG 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 U 2.303 RT log K c nF  G 0  nF  2.303 RT log K c nF Mn Zn Cr Fe Cd Co Ni Sn Pb H2 Cu I2 Hg Ag Br2 Pt Cl2 Au F2 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 Ca 2   2e   Ca U 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.,......(i) Wmax  nFEcell 0  But E cell Electrode Reaction (Reduction) –2.714 Increasing strength as reducing agent Increasing tendency to lose electrons 0 ECell  Ecell  E3 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 60 Since, the activities of pure copper and zinc metals are taken as unity, hence the Nernst equation for the Daniell’s cell is, –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 G  RT ln K c (i) Reactivity of metals: The activity of the metal depends on its ST G 0   2.303 RT log Kc or G  G  2.303 RT log Q 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 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 Element : Na  Zn  Fe Reduction potential :  2.71  0.76  0.44 60 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