Electrochemistry Chapter 12 PDF

Electrochemistry 491 Anode Flow of electrons Cathod e 60 Flow of current Chapter E3 12 Electrochemistry U Electrolytes and Electrolysis (iv) In voltameter, outside the electrolyte electrons flow from anode to cathode and current flow from cathode to anode. ID Electrochemistry is the branch of physic...

Electrochemistry 491 Anode Flow of electrons Cathod e 60 Flow of current Chapter E3 12 Electrochemistry U Electrolytes and Electrolysis (iv) In voltameter, outside the electrolyte electrons flow from anode to cathode and current flow from cathode to anode. ID Electrochemistry is the branch of physical chemistry which deals with the relationship between electrical energy and chemical changes taking place in redox reactions D YG (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.” 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 nonelectrolytes. U (2) Electrolytic cell or Voltameter : The device in which the process of electrolysis or electrolytic decomposition is carried out is known as electrolytic ST cell or voltameter. (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) (iii) During electrolysis in voltameter cations are discharged on cathode and anions on anode. For voltameter, E cell  ve and ΔG  ve. (v) The anions on reaching the anode give up their electrons and converted into the neutral atoms. 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. At cathode : B  e   B (Reduction) This overall change is known as primary change and products formed is known as primary products. 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 492 Electrochemistry discharge or deposition potential, (D.P.). The values of discharge potential are different for different ions. For cations : Li , K  , Na  , Ca 2  , Mg 2  , Al 3  , Zn2  , Fe 2 , Ni 2  , H  , Cu 2  , Hg 2  , Ag  , Au 3 . The decreasing order of discharge potential or the increasing order of deposition of some of the ions is For anions : SO 42  , NO 3 , OH  , Cl  , Br  , I . given below, Table : 12.1 Products of electrolysis of some electrolytes Electrode Product at cathode Product at anode Aqueous NaOH Pt or Graphite 2 H   2e  H 2 2OH  1 O 2  H 2 O  2e  2 Fused NaOH Pt or Graphite Na   e  Na 2 OH  1 O2  H 2 O  2e  2 Aqueous NaCl Pt or Graphite 2 H   2e  H 2 Fused NaCl Pt or Graphite Na   e  Na Aqueous CuSO4 Pt or Graphite Cu 2  2e  Cu 2OH  Aqueous CuSO4 Cu electrode Cu 2  2e  Cu Cu oxidised to Cu 2  ions Dilute H2SO4 Pt electrode 2 H   2e  H 2 2OH  Conc. H2SO4 Pt electrode 2 H   2e  H 2 Peroxodisulphuric acid (H 2 S 2O8 ) Aqueous AgNO3 Pt electrode Ag   e  Ag 2OH  Aqueous AgNO3 Ag electrode Ag   e  Ag Ag oxidised to Ag  ions E3 ID U ST (vi) Compounds like NaOH, KOH, Na 2CO 3 , KClO 3 , white lead, KMnO 4 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. Anode 2Cl  Cl 2  2e  1 O 2  H 2 O  2e  2 1 O2  H 2 O  2e  2 1 O2  H 2 O  2e  2 electroplati ng Cathode Electrolyte  dilute H 2 SO 4 With copper Cu Object With silver Ag Object K[ Ag (CN )2 ] With nickel Ni Object Nickel ammonium sulphate With gold Au Object K[ Au (CN )2 ] With zinc Zn Iron objects ZnSO 4 With tin Sn Iron objects SnSO 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. For 2Cl  Cl 2  2e  U D YG (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 60 Electrolyte CuSO 4 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  (a  b  c)  d  It E 96500 Using above relation we may calculate the thickness of coated layer. Electrochemistry 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, 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 (6.023  10 23 ) of electrons. = Current in Amperes (I) × Time in second (t)  W  I  t or W  Z  I  t In case current efficiency () is given, then  If in any reaction, n moles of electrons are involved, then the total electricity (Q ) involved in the reaction is given by, Q  nF  n  96,500 C Thus, the amount of electricity involved in any reaction is related to, ID 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 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. E3 Where, W = Mass of ions liberated in gm, Q  Quantity of electricity passed in Coulombs W  ZIt (3) Faraday's law for gaseous electrolytic product For the gases, we use It Ve V 96500 60 Faraday's laws of electrolysis 493 (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, U 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. 1 Coulomb  D YG 96500 Coulombs  6.023  10 23 electrons = 1 mole electrons. 6.023  10 23  6.28  10 18 electrons, 96500 ST U 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 Z It E Z1 E E  1  1 or 1  1 or ( 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. 1F = e   N  (1.602  10 19 c)  (6.023  10 23 mol 1 ). Number of Faraday  Number of electrons passed 6.023  10 23 (a) 1 mole of monovalent cation,(b) 1/2mole of divalent cation, (c) 1/3 mole of trivalent cation, valent cations. (d) 1/n mole of n 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 Electrolytic conduction (i) It is due to the flow of electrons. (i) It is due to the flow of ions. (ii) It is not accompanied by decomposition of the substance.(Only physical changes occurs) (ii) It is accompanied by decomposition of the substance. (Physical as well as chemical change occur) (iii) (iii) It involves transfer of It does not involve 494 Electrochemistry matter in the form of ions. (iv) Conductivity decreases with increase in temperature. (iv) Conductivity increases with increases in temperature and degree of hydration due to decreases in viscosity of medium. 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. Thus, resistivity is defined as the resistance of a conductor of 1 cm length and having area of crosssection equal to 1 cm 2. Units   R. The 2 a cm  Ohm l cm units of resistivity are  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., Electrolytic conduction 1 R E3 G 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 Siemens, S (i.e., 1S  1  1 ). ID 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. (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 I  or V  IR R : 60 transfer of matter. D YG U (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, conductivity is the conductance of one centimetre cube of a solution of an electrolyte. ST U 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 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 R l ; Now, if l  1 cm , a  1 cm 2 then R   a Thus,   1  Units : The units of conductivity are  1  Ohm 1 cm–1 or  1 cm 1 Ohm. cm In SI units, l is expressed in m area of crosssection 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 mole of an electrolyte in solution. It is denoted by  (lambda). Molar conductance is related to specific conductance (  ) as,   M where, M is the molar concentration. If M is in the units of molarity i.e., moles per litre 1 (mol L ), the  may be expressed as, Electrochemistry   1000 determination of the conductance of a solution involves the measurement of its resistance. 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, Conductanc e(G)  Conductivi ty  Molar conductivi ty() But if solution contains 1 gm mole of the electrolyte therefore, the measured conductance will be the molar conductivity. Thus, Molar conductivi ty()  100  Conductivi ty In other words, ()    V where V is the volume of the solution in If M is the concentration of the solution in mole per litre, then M mole of electrolyte is present in 1000 cm 3 or   1000 cm 3 of M     Volume in cm 3 containing 1 mole of   1000 M M D YG   1000 The units of  are S cm 1 and units of  are, cm 3 Λ  S cm 1   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. U (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 ST conductance as 1000   1000 e    (M is Molarity of the C M solution) 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. (8) Experimental measurement of conductance the a l 1l l   or   G   R a a where G is the conductance of the cell, l is the distance of separation of two electrodes having cross l The quantity   is called cell constant and is a expressed in 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., Conductivi ty  Conductanc e  Cell constant Factors affecting the electrolytic conductance In general, conductance of an electrolyte depends upon the following factors, Units of Molar Conductance : The units of molar conductance can be derived from the formula ,     and   R (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 Thus, electrolyte. 1  ID solution  section area a cm 2. cm 3 containing one gram mole of the electrolyte. 1 mole of electrolyte is present in  (ii) Calculation of conductivity : We have seen that conductivity () is reciprocal of resistivity (  ) , i.e., 60 M E3  495 (i) The conductance of a solution is reciprocal of resistance, therefore, the experimental (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 496 Electrochemistry 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 c where  is the degree of dissociation, c is the molar conductance at concentration C Loss around anode Speed of cation  Loss around cathode Speed of anion 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 tc and ta or nc and na”. From this definition, Current carried by an anion ta  Total current passed through the solution 60 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 E3 and tc  0 is the molar conductance at infinite dilution. 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. ID 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 interionic 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. Current carried by a cation Total current passed through the solution (3) Factors affecting transport number D YG U A rise in temperature tends to bring the transport number of cation and anion more closer to 0.5 (3) Temperature : The conductivity of an electrolyte depends upon the temperature. With increase in temperature, the conductivity of an electrolyte increases. U Migration of ions ST Electricity is carried out through the solution of an electrolyte by migration of ions. Therefore, (1) Ions move toward electrodes at different speeds. oppositely charged (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 (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–1 cm2 or V–1S-1cm2 Ionic mobility and transport number are related as, a orc  ta or tc   Absolute ionic mobility is the mobility with which the ion moves under unit potential gradient. It's unit is cm sec 1. Absolute ionic 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 Electrochemistry HCl   H  H    Cl  Cl  ;  Cl   1. So, HCl  For HCl, (1  H  )  (1  Cl  ) ; H  1 and 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. (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 m [1  (cm cm and / m )]  is the molar conductivity of a weak electrolyte at any concentration C and,  m is the molar conductivity of a electrolyte at  cm ) ; Thus, at any concentration, the (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   water , 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) ; c cm by, c  m   m (     ) Cm  D YG infinite dilution. Then, the degree of ionisation is given Thus, knowing the value of cm , and m (From the Kohlrausch's equation), the degree of ionisation at any concentration ( c ) can be determined. U (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 ST 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 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 : Following are the important characteristics of electrochemical cell, e– e– Voltmeter Salt bridge Cu cathode Zn anode Then, the ionisation constant of AB is given by, C 2 [ A  ][B  ] C .C  K   C(1   ) (1   ) [ AB ] C(cm )2 m (m ionisation constant (K) of the electrolyte can be determined. U electrolyte at any concentration. If  cm knowing C(cm / m )2 ID CH 3 COOH  CH 3 COONa  HCl  NaCl K Then, 60 The molar conductivity of HCl at infinite dilution can be expressed as, We know, that at any concentration C, the degree of ionisation ( ) is given by,   cm / m E3 expressing the molar conductivity of an electrolyte is illustrated as, 497 Porous CuSO4 ZnSO4plug Fig. 12.1 498 Electrochemistry 60 (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 E3 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. 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  p2 , oxidation occurs 2 (p 2 ) ID (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. (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. D YG U (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, KNO3, NH4NO3 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. ST U (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, Left Bridge Right Anode Cathode Oxidation Reductio Negative n Positive at L. H. 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. M (Hg C1 ) | M n | Zn(Hg C 2 ) Electrochemistry M | M  n (C1 )| | M n  (C 2 )| M or Zn | Zn2  (C1 ) Zn2  (C2 )| Zn || Anode Cathode Ecell  Types of commercial cells : There are mainly two types of commercial cells, (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 (i) Voltaic cellCu rod C 0.0591 log 2(R. H.S ) e at 25o C n C1(L. H.S.) + D YG (7) Heat of reaction in an electrochemical cell : Let n Faraday charge flows out of a cell of emf E, then …….(i) Gibbs – Helmholtz equation from thermodynamics may be given as  G  G  H  T    T  P …….(ii) Zn Fig. 12.2 Cathode : Cu rod Electrolyte : dil. H 2SO 4 Anode : Zn rod Emf : 1.08 V At cathode : Cu 2   2e  Cu At Anode : Zn Zn2   2e  Over all reaction : Zn  Cu 2  Zn2   Cu (ii) Daniel cell e– – Key Anode (Zn) Electrons flow Ammeter e– + Current Cathode (Cu) Salt bridge  (nFE )   E   nFE  H  T    H  nFT  T   T  P  P U Local action Cu Polarisatio n From equation (i) and (ii) we get, Cotton Plugs Cu2+ Zn2+ ST  E  H  nFE  nFT    T  P 1M ZnSO4 (aq)  E  where   = Temperature coefficient of cell  T  P Fig. 12.3 1M CuSO4 (aq) (Depolariser) Cathode : Cu rod Electrolyte : dil. H 2SO 4  E  Case I: When   = 0, then H  nFE  T  P  E  Case II: When   > 0, then nFE  H ,  T  Zn rod Dil. H2SO4 U 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. G  nFE – ID The emf of the cell is given by the following expression, 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 (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 C1 ). Some Commercial cell (Batteries) E3 The emf of the cell is given by the expression, 0.0591 C Ecell  log 1 at 25 o C n C2 499 Anode : Zn rod Emf : 1.1 V At cathode : Cu 2   2e  Cu At Anode : Zn Zn2   2e  i.e. process inside the cell is endothermic.  E  Case III: When   < 0, then nFE  H , i.e.,  T  Over all reaction : Zn  Cu 2  Zn2   Cu (iii) Lechlanche cell (Dry cell) + Seal Graphite (cathode) MnO2 +C (Depolariser) Paste of NH4Cl+ZnCl2 process inside the cell is exothermic. – Fig. 12.4 Zinc anode P 500 Electrochemistry Cathode : Mercury (II) oxide Zn rod Electrolyte : Paste of KOH  ZnO 1.35 V Cathode : Graphite rod Anode : Zn pot Electrolyte : Paste of NH 4 Cl  ZnCl2 in starch Anode : Emf : At cathode : HgO(s)  H 2 O(l)  2e  Hg (l)  2OH (aq) Emf : 1.2 V to 1.5 V Zn(s)  20 H (aq ) ZnO (s)  H 2 O(l)  2e  At cathode : NH 4  MnO 2  2e  MnO (OH )  NH 3 At Anode : At Anode : Zn Zn2   2e  Over all reaction : Zn  NH 4  MnO 2 Zn 2   MnO (OH )  NH 3 Over all reaction : Zn(s)  HgO(s) ZnO(s)  Hg(l) 60 (amalgam) (iv) Mercury cell Lead storage cell + – Alkali cell ID In charged E3 (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. + U Glass vessel PbO2 – Ni(OH)2 Fe(OH)2 Perforated steel grid Pb Positive electrode Negative electrode Electrolyte D YG dil. H2SO4 Perforated lead plates coated with PbO2 Perforated steel plate coated with Ni(OH)4 Perforated lead plates coated with pure lead Perforated steel plate coated with Fe dil. H2SO4 20% solution of KOH + 1% LiOH Chemical reaction At anode : PbSO4 + 2H+ + 2e– Pb + H2SO4 At cathode : PbSO4 + SO4– – + 2H2O – 2e– PbO2 + 2H2SO4 Specific gravity of H2SO4 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 + SO4– – – 2e– PbSO4 At cathode : PbO2 + 2H+ + 2e– + H2SO4 PbSO4 + 2H2O Specific gravity of H2SO4 decreases and when specific gravity falls below 1.18 the cell requires recharging. Chemical reaction At anode : Ni (OH)2 + 2OH+ – 2e– Ni(OH)4 At cathode : Fe(OH)2 + 2K+ + 2e– Fe + 2KOH Emf of cell : When cell is fully charged then E = 1.36 volt ST U During charging During discharging KOH 20% + Li(OH), 1% Chemical reaction At anode : Fe + 2OH– – 2e– Fe(OH)2 At cathode : Ni(OH)4 + 2K+ + 2e– Ni(OH)2 + 2KOH Emf of cell : When emf of cell falls below 1.1 V it requires charging. Electrochemistry Emf of cell : When emf of cell falls below 1.9 volt the cell requires recharging. 80% 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 At anode : 2[H 2 (g)  2OH  ](aq)  2 H 2 O(l)  2e  At cathode : O2 (g)  2 H 2 O(l)  4 e   4 OH  (aq) Overall reaction 2 H 2 (g)  O2 (g)  2 H 2 O(l) 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. Electrode Potential (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 n+ collides with the electrode, and undergoes no change. (ii) A metal ion M n+ collides with the electrode, gains n electrons and gets converted into a metal atom M, (i.e. the metal ion is reduced). ID 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 battery. The fuel can be fed continuously to produce power. For this reason, H 2  O 2 fuel cells have been used 60 Fuel cells : 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 U 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 D YG the reactants are fed. Fuel cells convert the energy of the fuel directly into electricity EMF of fuel cell is 1.23 V. This cell has been used for electric power in the H2OThe important advantages of Apollo space programme. fuel cells are + Cathode Anode – Porous carbon electrode U OH– H2 60% E3 Efficiency 501 O2 Electrolyte ST Fig. 12.6 (1) High efficiency : The 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 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. (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. 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 (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 502 Electrochemistry contact with a molar solution of one of its ions and the normal hydrogen electrode (NHE)”. 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, Example Half – cell reaction Pt(H 2 )| H  (aq) Gas ion half - cell 1 H 2 (g) H  (aq)  e  2 Pt(Cl 2 )| Cl  (aq) 1 Cl (aq) Cl 2 (g)  e  2  Ag | Ag  (aq) Ag(s) Ag  (aq)  e  Ag, AgCl | Cl  (aq) Ag(s)  Cl  (aq) AgCl(s)  e  Hg, Hg2Cl 2 | Cl  (aq) 2 Hg(l)  2Cl  (aq) Reversibl e to Electrode Potential (oxidn), E = [H  ] H E0  0.0591log[H  ] 1 [Cl  ] Cl  E0  0.0591log[Cl  ] [ Ag  ] Ag  E0  0.0591log[Ag  ] 1 [Cl  ] Cl  E0  0.0591log[Cl  ] 1 [Cl  ]2 Cl  E0  0.0591log[Cl  ] 1 [OH  ]2 OH  E0  0.0591log[OH  ] [ Fe 3  ] [ Fe 2  ] Fe2  , Fe3  ID Metal – metal ion half – cell Metal insoluble salt anion half – cell Calomel electrode Q= E3 Type Hg 2Cl 2 (s)  2e  Hg(l)  2OH  (aq) U Hg, HgO | OH  (aq) HgO(s)  H 2O(l)  2e Pt | Fe 2  , Fe 3  (aq) (a q) 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 U to cathode and flow of current from cathode to anode. Flow of electrons Anode Cathode ST Flow of current (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 0 0 (anode )  E red (cathode ) potential of cathode  E ox (ii) When reduction potentials of both electrodes are taken into account  Fe2  (aq) Fe3  (aq)  e  D YG Metal – metal oxide hydroxide half - cell Oxidation – reduction half – dell 60 Table : 12.2 Various Types of Half – cells E 0  0.0591 log [Fe 3  ] [Fe 2  ] 0 Ecell  Reduction potential of cathode – Reduction potential of anode 0 o 0 o  E Cathode  Eright  Eleft  E Anode (iii) When oxidation potentials of both electrodes are taken into account o Ecell  Oxidation potential of anode – Oxidation 0 0 (anode )  Eox (cathode ) potential of cathode  Eox (3) Difference between emf and potential difference 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. It is the maximum voltage that the cell can It is always less then the maximum value of deliver. voltage which the cell can deliver. Electrochemistry It is responsible for the steady flow of current in It is not responsible for the steady flow of the cell. current in the cell. Nature of o ΔG(or ΔG ) E cell (or E ocell At 298 K, the Nernst equation for the ) Oxidised form  ne  Reduced form – + Equilibrium 0 0 Non – spontaneous + – The Nernst equation for the electrode is written 60 Spontaneous as, 0 Ehalf  cell  Ehalf  cell  Nernst's equation reaction, M n  (aq)  ne  M(s) ID The Nernst equation is written as, 0 Ecell  Ecell  U 2.303 RT [M (s)]  log nF [M n  (aq)] D YG 0 EM = the standard electrode potential n /M R = the universal gas constant, 8.31 J K 1 mol 1 T= the temperature on the absolute scale, U n = the number of electrons involved in the electrode reaction, ST [M (s)] = the concentration of the deposited metal, [M n (aq)] = the molar concentration of the metal ion in the solution, o  or Ecell  Ecell 2. 303 RT [C ]c [D]d log nF [ A]a [B]b At 298 K, above eq. can be written as, o  or Ecell  Ecell 0. 0592 [C ]c [D]d log n [ A]a [B]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. Zn(s) Zn2  (aq)  2e  At anode : At cathode : Net The concentration of pure metal M(s) is taken as 2.303 RT 1 log n nF [M (aq)] o The E cell is called the standard cell potential. (3) Nernst’s equation for Daniells cell : Daniell’s cell consists of zinc and copper electrodes. The electrode reactions in Daniell’s cell are, F = the Faraday constant : (96500 C), unity. So, the Nernst equation for the RT [C]c [D]d ln nF [ A]a [B]b 0 0 0  E cathode  E anode Where, E cell. E M n  / M = the potential of the electrode at a given electrode is written as, 0.0591 [Reduced form ] log n [Oxidised form] For a cell in which the net cell reaction involving n electrons is, aA  bB cC  dD above eq. is called the Nernst equation. 0 EM n  / M  EM  n /M 0 Ehalf  cell  Ehalf  cell  (2) Nernst’s equation for cell EMF takes place is described by the equation, RT [M (s)] 0 EM n  / M  EM  ln n /M nF [M n  (aq.)] concentration, E3 The potential of the electrode at which the 0 EM n /M 2.303 RT [Reduced form ] log nF [Oxidised form] At 298 K, the Nernst equation can be written as, (1) Nernst’s equation for electrode potential Where, 0.0591 1 log n n [M (aq)] For an electrode (half - cell) corresponding to the electrode reaction, reaction or EM n  / M  M n / M electrode can be written as, 0 EM n  / M  EM  n /M (4) Cell EMF and the spontaneity of the reaction : We know, G  nFEcell 503 M n / M Cu 2  (aq)  2e  Cu(s) cell 2 reaction : 2 Zn(s)  Cu (aq) Cu(s)  Zn (aq) Therefore, the Nernst equation for the Daniell’s cell is, 0 Ecdll  Ecell  2.303 RT [Cu (s)][Zn 2  (aq)] log 2F [Zn(s)][Cu 2  (aq)] 504 Electrochemistry 2.303 RT [ Zn 2  (aq] log 2F [Cu 2  (aq)] The above eq. at 298 K is, o Ecdll  Ecell  0.0591 [Zn2  (aq] log V 2 [Cu 2  (aq)] 0  1.1 V For Daniells cell, E cell (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.....(i) ID 0 ECell  Ecell  At equilibrium, the cell cannot perform any useful work. So at equilibrium, ECell is zero. Also at U equilibrium, the ratio [C]c [D]d  [C]c [D]d    Kc  [ A]a [B]b  [ A]a [B]b  equilibriu m (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 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 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. 60 0 Ecdll  Ecell  (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. E3 Since, the activities of pure copper and zinc metals are taken as unity, hence the Nernst equation for the Daniell’s cell is, D YG 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.,......(i) Wmax  nFEcell According to thermodynamics the free energy change (G) is equal to the maximum work. In the cell U work is done on the surroundings by which electrical energy flows through the external circuit, So Wmax,  G (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 non-metals 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. Table : 12.3 Standard reduction electrode potentials at 298K......(ii) Element ST from eq. (i) and (ii) G  nFEcell Electrode Reaction (Reduction) Standard Electrode Reduction potential E0, volt 0 In standard conditions G 0   nFE cell Where G 0  standard free energy change 2.303 0  RT log K c But E cell nF Li Li++ e– = Li –3.05 K K++ e– = K –2.925 Ba Ba+++ 2e = Ba –2.90  G  nF  2.303 RT log K c nF Sr Sr+++ 2e = Sr –2.89 G   2.303 RT log K c or G  G  2.303 RT log Q Na G 0   RT ln K c Mg (2.303 log X  ln X ) Electrochemical series Al Ca 2  –2.87  2e  Ca Na   e   Na Mg 2   2e   Mg Al 3   3 e   Al rength as reducing agent Increasing ndency to lose electrons 0 Ca tendency to accept electrons Increasing strength as oxidising agent 0 –2.714 –2.37 –1.66 Electrochemistry Mn+++ 2e = Mn Zn Zn2+ +2e–=Zn Cr Cr3++3 e– = Cr –0.74 Fe Fe2++ 2e– = Fe –0.44 Cd Cd2++2e– = Cd –0.403 Co Co+++ 2e = Co –0.27 Ni Ni2++2e– = Ni –0.25 Sn Sn2++2e– = Sn –0.14 Pb Pb+++ 2e = Pb –0.12 H2 2H++2e– = H2 0.00 Cu Cu2++ 2e– = Cu +0.337 I2 I2+2e– = 2I– +0.535 – Hg Hg +2e = Hg +0.885 Ag Ag++ e– = Ag +0.799 Br2 Br2+2e– = 2Br– +1.08 Pt Pt+++ 2e = Pt +1.20 Cl2 Cl2+2e– = 2Cl– +1.36 Au Au 3++3e– = Au +1.50 F2 F2+2e–= 2F– +2.87 (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. 60 2+ negative like alkali metals, alkaline earth metals are strongly electropositive in nature. –1.18 –0.7628 (iii) Displacement reactions (3) Application of Electrochemical series ID (i) Reactivity of metals: The activity of the metal depends on its tendency to lose electron or electrons, (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. E3 Mn 505 i.e., tendency to form cation (M n  ). This tendency D YG U 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, (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 U (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 corresponding salts and combine with those substances which accept electrons. ST (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 and iodine from bromides and iodides. Cl2  2 KI 2 KCl  I2 2 I  I2  2e  Cl 2  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 Mn n  ne  .....(Oxidation)  2 H  2e H 2.....(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. 506 Electrochemistry The metals which are below hydrogen in electrochemical series like Cu, Hg, Au, Pt, etc., do not evolve hydrogen from dilute acids. 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 (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. not stable towards heat. The metals which come below copper form unstable oxides, i.e., these are decomposed on heating. Element : Na  Zn  Fe Reduction potential :  2.71  0.76  0.44 alkaline earth metals are 60 BaO    2 HgO  2 Hg  O2 ; Na 2 O   No decomposit ion  Al 2 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 ID Alkali and reducing agents. 1 O2 2 E3 (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)  Ag 2 O  2 Ag  strong (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 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. D YG U (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. Corrosion F2 (Fluorine) is a stronger oxidant than Cl 2 , Br2 and I 2 , Cl 2 (Chlorine) is a stronger oxidant than Thus corrosion is a process of deterioration of a metal as a result of its reaction with air or water (environment) surrounding it. I2 Br2 Cl 2 F2  0.53  1.06  1.36  2.85  Oxidising nature increases 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 U Br 2 and I 2 Element : Reduction potential : ST Thus, in electrochemical series Highest negative reduction potential or (Minimum reduction potential) nature (Strongest reducing agent) Reducing nature Oxidising Top (Strongest oxidising agent) Botto Highest positive value of reduction potential m (vi) Thermal stability of metallic oxides : The thermal stability of the metal oxide depends on its electropositive nature. As the electropositivity 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 Electrochemistry Presence of CO 2 in natural water increase rusting of iron. 507 H 2O  H   OH  or CO 2  H 2 O  H   HCO 3 At cathode : H   e   H (v) When the iron surface is coated with layers of metals more active than iron, then the rate of corrosion is retarded. dissolved oxygen. 4 H  O2  2 H 2 O A rise in temperature (with in a reasonable limit) increases the rate of corrosion. Therefore, the overall reaction at cathode of different electrochemical cells may be written as, (a) Reactive gases come in contact with metals at high temperatures e.g., corrosion in chemical industry. Oxi. half reaction : Fe(s) Fe 2  (aq)  2e  ]  2 (E  0. 44 V ) Red. half reaction : 4 H   O2  4 e  2 H 2 O (E  1. 23 V ) Overall cell reaction : 2 Fe (s)  4 H   O2 2 Fe 2  (aq)  2 H 2 O ID (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. (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. 60 (i) Chemical corrosion : Such corrosion, generally takes place when 4 H   O2  4 e   2 H 2 O E3 (3) Classification of corrosion process : Depending upon the nature of corrosion, and the factors affecting it, the corrosion may be classified as follows. The hydrogen atoms on the iron surface reduce of iron : D YG 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)  Fe 2  (aq)  2e . U 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. ST Rust Fe anode Drop of moisture 4H++O2+4e–2H2O (Cathode) Fe2+ + 2e– Iron Flow of electron s 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 further by 4 Fe 2  (aq)  O2 (g)  4 H 2 O  2 Fe 2 O3  8 H  and U (4) Mechanism of rusting Electrochemical theory of rusting. The ferrous ions are oxidised atmospheric oxygen to form rust. (ECell  1.67 V ) Fe 2 O3  xH 2 O  Fe 2 O3. xH 2 O 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 508 Electrochemistry 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. E3 ID  When two or more ions compete at the electrodes 60 (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. then the ion with higher reduction potential gets liberated at the cathode while the one with lower reduction potential at the anode.  Cell U constant is determined with the help of conductivity bridge, where a standard solution of KCl is used.  If the external EMF is slightly more than the D YG 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 / NH4Cl etc., can not be used in the salt bridge of a cell 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. U  In Appolo moon flights, H2 – O2 fuel cell was the source of energy and drinking water.  Conductivity water is the highly purified water ST whose on conductance is very small. It is prepared by the demineralisation of ordinary water by passing through cation and anion exchange resins.