Electrochemistry Study Guide

Electrochemistry (ii) Electrolytic conductors :INTRODUCTION * Electrochemistry deals with the study of electrical Those substances whose water solution conducts properties of solutions of electrolytes and with the the electric current and which are decomposed by inter-relation of chemical phenomenon and electrical the passage of current are called electrolytes. In this energies. Electrical energy is carried through matter case, conduction takes place by movement of ions. in the form of electric curent with the help of suitable electrolytes also conduct ele ctricity in fused state source and charge carriers (ions or electrons) and undergo decomposition by passage the electric Conductors and non conductors current. Substances are devided into two classes * Substances whose aqueous solution does not (a) Non conductor or insulator (b) Conductor conduct electric current are called non-electrolytes. (a) Non - conductor :They do not conduct electricity in the fused state. Substances which do not allow electric current to Solutions of cane suger, glycerine, glucose, urea etc. pass through them are called non-conductors or insulators examples - pure water, ice, glass, rubber are the examples of non electrolytes. etc. Strong electrolyte :(b) Conductor :- Substances which allow electric Electrolytes which are completely ionized in aqueous current to flow through them are called conductors. solution or in their molten state; are called strong Examples - Metals, Aqueous solution of acids, bases and salts, fused salts and impure water etc. electrolyte. Conductors are of two types Example – all salts, strong acid and strong base (i) Metallic conductors Weak electrolyte :(ii) Electrolytic conductors or electrolytes. Electrolytes which are not completely ionized in aqueous (i) Metallic conductors :soluion or in their molten state; are called weak The conductors which conduct electric current by electrolytes. movement of electrons without undergoing any chemical change are known as metallic conductors. Examples :- All carbonic acids (except sulphonic Metals (Cu,Ag,Fe,Al etc), non metals (graphite) and acid), CH3COOH, HCN, NH3, amine, etc. various alloys and minerals are examples. 1 Electrochemistry Difference between metallic and electrolytic conduction (i) Metallic conduction Electrolytic conduction Electrons flow from negative end Ions flow towards the oppositely charged electrodes to positve end (ii) Flow of electricity takes place without the decomposition of the substance Flow of electricity takes place by chemical changes at electrodes. (iii) No transfer of matter takes place. Transfer of matter takes place in the form of ions (iv) The resistance to the flow of current The resistance to the flow of current decreases increases with the increase in with the increase in temperature and hence temperature and hence the increase in increase in temperature increases the conduction. temperature decreases the conduction. e At anode Cl  ELECTROLYSIS : Cl The process of decomposition of an electrolyte by Cl + Cl Cl2 the passage of electricity is called electrolysis or e At cathode Na+   Na electrolytic dissociation. It is carried out in electrolytic cell where electrical energy is converted (b) Electrolysis of aqueous solution of NaCl into chemical energy. For electrolysis to take place The solution of NaCl containts Na+ , Cl and small two suitable electrodes are immersed in the liquid. amounts of H+, OH (due to small dissociation of The solution of an electrolyte contains ions. water) When an electric potential is applied between the electrodes, the positively charged ions move towards * If more than one types of ions are present at a given the cathode and negatively ions move towards the electrode, then the one liberated is the ion which anode, when a cation reaches the cathode, its takes requires least energy. The energy required to liberate up electron(s) and thus gets its charge neutralised. an ion is provided by the applied potential between Thus the gain of electrons (decrease in oxidation electrodes. This potential is called discharge or number) means reduction takes place at the cathode. deposition potential Similarly an anion when it reaches the anode, gives up electron(s) and thus gets discharged. Loss of e At cathode H+   H electrons(Increase in oxidation number) means H + H H2 oxidation takes place at anode. * The tendency of an electrode to loose electrons is e At anode Cl   Cl known as the oxidation potential. Cl + Cl Cl2 * The tendency of an electrode to gain electrons is + known as the reduction potential. Na ions move towards the cathode and combine with (a) Electrolysis of fused sodium chloride :OH ions furnished by feebly ionised water to form + When fused sodium chloride is electrolysed, Na NaOH. ions moves towards the cathode and Cl ions moves 2H2O + 2e  H2 + 2OH towards the anode. At cathode Na+ ions accept electrons to form sodium metal. At anode each Cl Na + + OH  NaOH ion loses an electron to form Cl2 gas. 2 Electrochemistry PRODUCTS OF ELECTROLYSIS OF SOME ELECTROLYTES S. Electrolyte Electrode Product obtained No. at anode cathode (i) Aqueous NaCl Pt or Graphite Cl2 H2 (ii) Fused NaCl Pt or Graphite Cl2 Na (iii) Aqueous NaOH Pt or Graphite O2 H2 (iv) Fused NaOH Pt or Graphite O2 Na (v) Aqueous CuSO4 Pt or Graphite O2 Cu (vi) Aqueous CuSO4 Copper Cu oxides to Cu+2 Ions - (vii) Aqueous CuCl2 Pt or Graphite Cu oxides to Cu+2 ions Cu (viii) Dilute HCl Pt or Graphite Cl2 H2 (ix) Dilute H2SO4 Pt or Graphite O2 H2 (x) Aqueous AgNO3 Pt of Graphite O2 Ag ELECTROLYTIC CONDUCTANCE :(a) Resistance (R) :- Metallic and electrolytic conductors obey ohm's law according to which the resistance of a conductor is the ratio of the applied potential difference (V) to the current following(I). R= V I * R is expressed in ohms. (b) Conductance (C) :- The conductance of a conductor is equal to reciprocal of resistance C= 1 R R  L A (d) Specific conductance () :- It is defined as the reciprocal of specific resistance 1 =  The above definitions apply to metallic conductors and electrolytes. * R  In the case of solution of electrolytes, the resistance offered by the solution to the flow of current is. (a) Directly proportional to the distance between the electrodes * C is expressed in mho. (c) Specific resistance () The resistance (R) of a conductor of uniform cross section is directly proportional to its length(L) and inversely proportional to its area of cross section (A) * Product obtained at R  L (b) Inversely proportional to the area of cross section of the electrodes R 1 A  The conductance C = L A 1 L A A L R= CL whre  is a constant and called resistivity or specific resistance.   =  so  = A If L = 1 cm , A = 1 cm2 then  = C When L = 1 , A = 1 , then  = R thus the specific resistance may be defined as the resistance of a conductor of unit length and unit area of cross section. Hence specific conductivity of a solution is defined as the conductance of one centimeter cube of the solution of the electrolyte. Unit of  ohm - cm * = CL A Where L A = cell constant 3 Electrochemistry *   100 0 Cell constant (x) is a fixed quantity for a particular 2 Sol. (i)  M  = 2 10 M cell and is defined as the distance between two parallel electrodes of a cell devided by the area of = 250 —1 cm2 mol—1 cross section of the electrodes.   100 0 (ii)  M   = C × cell constant 1 00 0  0 .0 8 M cell constant =  C = R = Specific conductivity x resistance * Unit of  is ohm-1 cm-1 (e) Equivalent Conductance (eq) :- It is defined as the product of specific conductance () and the volume (V in m) of the solution which contain one gram equivalent of the electrolyte.  eq =  × V If concentration of solution is C - gram equivalent per litre then eq = * Ex.1 The resistance of a 1N solution of salt is 50  Calculate the equivalent conductance of the solution, if the two platinum electrodes in solution are 2.1cm apart and each having an area of 4.2cm2. 1 1L   R  A    100 0 and  eq.  N = 1 2 .1 1   5 0 4 .2 1 00 = 1 1 00 0  100 1 = 10 S (f) Molar conductance :- (mor ) It is defined as the product of specific conductance () and the volume (V in m)) in which contains one mole of the electrolyte m  =  × v and m =   1000 M Unit  ohm-1 cm2 per mole Ex.2 Which of the following have maximum molar conductivity. (i) 0.08M solution and its specific conductivity is 2 × 10–2  (ii) 0.1M solution and its resistivity is 50 cm. 1 1 00 0  50 0 .1 200 —1 cm2 mol—1 So, the molar conductivity of 0.08M solution will be greater than 0.1M solution Relation between  eq. and M :M    100 0 M and  eq.    100 0 N We know that Normality = Valency Factor × Molarity or N = n × M  1 00 0 C Unit  ohm-1 cm2 per gram equivalent Sol.   1     M = so  eq.  M V.F. FACTORS INFLUENCING THE CONDUCTIVITY OF ELECTROLYTE (a) Inter ionic attraction :If inter ionic attraction between solute is more, then the conductivity will be less. (b) Polarity of solvent :If the solvant is greater polarized then the ionization and conductivity will be more (c) Viscosity of medium :By increasing the viscosity of medium, the conductivity decreases. (d) Temperature :As the temperature of electrolyte solution is increased, the conductivity increases because K.E. of ions increases and all type of attraction forces decreases and the viscosity of medium decreases. (e) Dilution :(i) The degree of ionisation increases with the increase of dilution of the solution the conductivity is increases due to increasing the no. of ions. (ii) Effect of dilution on specific conductance :Specific conductance decreases with the increase of dilution of the solution due to the presence of number of ions in 1cm3 solution decreases. 4 Electrochemistry (iii) Effect of dilution on equivalent/ molar conductivity :Current carried by cation The equivalent/molar conductivity increases with nc = Totalcurrent carried na + nc = 1 dilution. This is due to the fact that degree of ionisation * The transport number can be determined by Hittorf's increases with dilution. method, moving boundry method, EMF method and when the whole of the electrolyte has ionised, further from ionic mobility. addition of the water bring a small change in the value of equivalent/molar conductance. This stage is called KOHLRAUSCH LAW :infinite dilution. This law states that the equivalent conductivity of The ratio of equivalent conductivity at any dilution any electrolyte at infinite dilution ( eq. ) is the sum of to equivalent dilution at infinite dilution is called conductivity ratio or degree of dissociation of soluteionic conductances of the cation and anion given by the electrolytes at infinite dilution.     eq   a   C Current carried by anion na = Totalcurrent carried * * consider a salt NaCl, its equivalent conductivity at infinite dilution is the sum of two terms i.e. Ex.3 The equivalent conductivity of H2SO4 at infinite dilution is 384  cm2 eq–1. If 49 g H2SO4 per litre is present in solution and specific resistance is 184 then calculate the degree of dissociation. Sol. Equivalent of H2SO4 = 49 49 These terms are known as ionic mobilities of anion and cation . The ratio of ionic mobilities is equal to their transport numbers = = 1N 1 a c 1 specific conductance = specific resistance = 18.4  eq.    eq   Na     Cl  eq   Ceq.   eq.  55 38 4 = 0.14 na For AxBy type solution :- 1 00 0   1 0 00  1    55 N 1 8.4 Degree of dissociation (= na = n = (1  n ) c a 1  1  C   a Z Z here c = cation & a = anion and Z+, Z– are the charges on the cation and anion respectively. = 14%  eq  1  1  C  a y x (f) Effect temperature and pressure :Conductivity of all electrolytes increases with increase in temperature. Conductivity of solution is slightly changed with change in pressure due to Ex.4 At infinite dilution the equivalent conductance of Al+3 and SO4–2 ion are 189 and 160  cm2eq–1 change in viscosity of medium. respectively. Calculate the equivalent and molar The viscosity of dilute solution decreases with conductivity at infinite dilute of Al2(SO4)3. increase in pressure. i.e., equivalent conductivity 1  1  1 1  increases. Sol.  eq. Al SO    Al   SO =  1 89   16 0 2 TRANSPORT NUMBER :The fraction of the current carried by an anion is termed as its transport number. If na and nc are the transport numbers of anion and cation respectively then :- 4 3 3 3 2 2 4 3 2 = 143  cm2 eq–1 Molar conductivity = eq × V. F. = 143 × 6 = 858 cm2 mol–1 5 Application :- This law can be used to (a) Determine the equivalent conductivities of weak electrolytes at infinite dilution. (b) To determine the ionic mobilities and absolute velocity of ions. (c) Determination of degree of dissociation, since , the degree of dissociation of an electrolyte at the dilution V is given by  =   Electrochemistry Ex.6 The time required to coat a metal surface of 80 cm2 with 0.005mm thick layer of silver (density= 10.5 g cm–3) with the passage of 3A current through silver nitrate solution is – Sol.  Volume of layer of silver = 0.005 × 10–1 × 80 = 0.04 cm3  Mass = Density × volume = 10.5 × 0.04 = 0.42 g E So w  9 65 00 × It where   = a + c or t = 0.42  96 5 00 108  3 108 or 0.42  965 00 × 3 × t = 125.09 seconds. FARADAY'S LAWS OF ELECTROLYSIS (b) Faraday's second law : This law states that the amounts of different substances deposited at electrodes by passage of the same quantity of (a) Faraday's first law of elecrolysis : This law states electricity are proportional to their chemical that "The amount of a substance deposited or equivalent(E). dissolved at an electrode is directly proportional to WE the charge passing through the electrolytes. If W 1 and W 2 be the amounts of two different If a currect of I amperes is passed for t seconds, the substances deposited at electrodes and E1 and E2 be quantity of charge Q in coulombs. If Wgram of W1 E1 substances is deposited by Q coulombs of electricity, the equivalent weights then - W = E combining 2 2 then ItE the two laws W  It E or W = E F It W  Q  It or W = z It = 9 65 00 1 Where is proportionality constant and F is called Where z is constant of proportionality and is known Michael Faraday deduced two important law:- F as electrochemical equivalent. * When Q = 1 coulomb , W = z thus electro chemical equivalent may be defined as the weight in grams of an element liberated by the passage of 1 coulomb of electricity. * Ex.5 How many cc of chlorine will be deposited by 100 A current flowing for 5 hours through melted NaCl. * 5 Sol. Q = It = 100 × 5 × 60 × 60 = 18 × 10 W = ZQ = E 9 65 0 0 1 8  3 5 .5  10 3 9 65 × 18 × 105 = 18E 9 65 0 0 zQ = 22.4 L  volume of 662.2 g Cl2 at NTP = 2 2.4 71 × 662.2 W Q W Q  w = zQ  M = z ZF = ZF  M = ZF = 662.2 g  Volume of 71 g Cl2 at NTP = 208.9 L × 103 = faraday. when It = F then W = E Hence faraday (F) is the quantity of charge in coulombs required to deposit one gram equivalent of any substance. The Faraday (F) is also the quantity of charge carried by one mole of electrones F = e x N = 1.6 x 10-19 x 6.023 x 1023 = 96500 coulombs. To deposit one mole of any metal of valency Z, requires Z moles of electrons. thus M = zZF (M  mass of one mole) M Z * The equivalent weight E = * When atoms of the element are deposited, M is called gm atomic weight and when molecules are deposited M is called gm molecular weight. 6 Electrochemistry Cell reaction – Zn + Cu+2  Zn+2 + Cu Ex.7 The same current if passed through solution of silver nitrate and cupric salt connected in series. If the weight of silver diposited is 1.08 g . Calculate the weight of copper deposited Representation of Galvanic cell. Zn/Zn+2 // Cu+2/ Cu Sol. According to faradays second law W1 E1  W2 E2  1.08 10 8  W2 3 1.7 5  w2 = 0.3175 g Example – Daniel Cell * The chemical reaction occuring in a galvanic cell is always a redox reaction. * Anode Cathode Electrode potential : When a strip of metal is brought in contact with the solution containing its own ions then the strip of metal gets positively charged or negatively charged and results into a potential being developed between the metallic strip and its solution which is known as electrode potential. GALVANAIC CELL / VOLTAIC CELL A cell in which the chemical energy is transformed into electrical energy. (C2) salt bridge ELECTRO CHEMICAL CELL / * (C1) At anode M M+n + ne— (Oxidation Potential) During the chemical process, the reduction in free energy will obtain as a result in the form of electrical * energy. At cathodeM+n ne— M (Reduction Potential) The value of electrode potential depends upon : (1) the nature of electrode (2) the concentration of solution (3) the temperature Standard electrode potnetial (E 0) : If the concentration of ions is unity, temperature is 250C and pressure is 1 atm (standard conditions), the potential of the electrode is called standard electrode potential. * Salt bridge :1. It allows to flow of current by completing the circuit. 2. It maintains electrical neutrality of electrolytes in two half cells. Characteristics of electrolyte used in salt bridge: The given value of electrode potential be regarded as reduction potential unless it is specifically mentioned that it is oxidation potential. Electro motive force of cell or cell voltage : The difference in the electrode potentials of the two electrodes of the cell is termed as electro motive force [EMF] or cell voltage. Ecell = Ered (cathode ) – Ered (anode) 1. The electrolyte should be inert or Ecell = Eoxi.(anode) – Eoxi.(Cathode) 2. The cations and anions of the electrolyte used should be of the same ionic mobality or Ecell = Eoxi.(anode) + Ered(cathode) 7 3. ELECTRO CHEMICAL SERIES The arrangement of various elements in order of increasing values of standard reduction potentials is called electrochemical series. Standard electrode Potential 2. for example : 2AgNO3 + Cu  Cu(NO3)2 + 2Ag E°red (Volt) CuSO4 + Ag  Reaction is not observed Li K –3.05 –2.925 Ba –2.90 For example : Ca –2.87 Na + H2SO4  Na2SO4 + H2 Na –2.714 Cu + H2SO4  Reaction is not observed Mg –2.37 Al –1.66 Mn –1.305 Zn –0.7628 Cr –0.74 Fe –0.44 Cd –0.403 Co –0.28 Ni –0.25 Sn –0.14 With the help of Nernst equation, we can calculate the electrode potential of electrode or EMF of cell. Pb –0.13 E cell  E 0  H 0.00 Cu +0.337 I2 +0.535 Ag +0.799 Hg +0.885 Br2 +1.08 Cl2 +1.36 Let, in the cell : Pt +1.20 at Anode : M1 M1 +n + ne– Au +1.50 for this reaction : F2 +2.87 Important points about series :1. Electrochemistry Metals near the top of the series can displace more electronegative metal below them from their salt. Electrode whose standard reduction potential is less, act as anode, and other one which has high reduction potential acts as cathode. Metals near the top of the series are strongly electropositive. 4. Metal above hydrogen can displace H2 from dilute acid 5. Hydro oxides of metal in the upper part of series are strongly basic while hydro oxides of a metal in lower part are weakly basic. 6. The activity of non metals increases from top to bottom. 7. The metals which come below copper form unstable oxides, i.e., these are decomposed on heating. NERNST EQUATION RT [Product] log e nf [Reactant] Where - E0 = standard electrode potential, R = gas constant, T = temperature (in K) F = Faraday (96500 coulomb mol–1), n = number of e– gained or loosed in balanced equation. 0 or Ecell  E  E oxi  0 E oxi 0.059 [Product] loge n [Reactant] 0 .05 9  log1 0 n M1 n    M1  -------(1) at Cathode : M2+n + ne— M2 for this reaction : 0 E red.  E red.  0 .05 9 log10 n M2  M2n    -------(2) 8 Electrochemistry Note : Concentration of solid taken as unity. 2 so [M1] = [M2] = 1  Cr 3   = 0.30  0.059 log  3 6  2 Fe    We know that EMF of cell is EMF = Eoxi + (Anode) Ered. 2 (Cathode) By adding equation (1) & (2) we get – EMF  E 0oxi 0.05 9  log M1 n    n EMF  E cell   M1 n  0.05 9   log   n n  M2    Ex.8 The 0.1M copper sulphae solution in which copper electrode is dipped at 250C. Calculate the electrode potential of copper electrode [Given E 0Cu+2 / Cu = 0.34V] Sol. Cu+2 + 2e— Cu E red  E red 0 Calculate = 0.34 – 0.03 = 0.31 Volts Note : 1. The value of electrode potential does not depend on sticheometry of half reactions. 2. In electrolytic cell, cathode is negative terminal whereas in Galvanic cell cathode is positive terminal. 3. Work obtained by electro chemical cell = free energy of cell G = –nFEcells G° = –nFE°cells 4. Cell reaction is spontaneous, if G = – ve or Ecell = +ve 5. Cell reaction will be in equilibrium e.g.G = 0 or –nFEcell; = 0 or Ecell = 0 6. In nernst equation : 0.05 9 log1 0 K n Cr / Cr+3 (0.1M) // Fe+2 (0.01M) / Fe = – 0.45 V) Fe Sol. Half cell reactions are – At Anode & [Cr Cr+3 + 3e-— ] × 2 At Cathode & [Fe+2 + 2e— Fe ] × 3 ——————————— 2Cr + 3Fe+2 2Cr+3 + 3Fe E°cell = oxidation potantial + Reduction potantial = 0.70 + (–0.45) = 0.30 or E0cell  RT log e K nF put the value of E0Cell in : G° = – nFE°cell , we get G 0   nF Ex.9 The EMF of the cell over all reaction = 0 .3 0  0 .2 4 = 0.26 Volt. or E0cell  log10 (Given ECr  3/ Cr = – 0.75V, E Fe 2 / 0.01  0.059 [Product] log10 n [Reactant] 0.059 [Product] 0 If Ecell = 0, then Ecell  n log10 [Reactant] here n = 2 so E = 0.34 – 6 E cell  E0cell  0.059 [Product]  log n [Reactant] 0.0 5 9 2 0.1  = 0.30  0.05 9 log   3 6 0.0 59 1  E0red.  log n  M2 n    0.05 9    n  0 0 log M1  log  M2 n   = E oxi  Ered.     n   E 0cell 0.059 [Product] log n [Reactant] E cell  E0  RT log e K nF or G° = – RT loge K Applications of Electrolysis :(a) Electrometallurgy :- The metals (Na, K, Mg, Al etc.) are obtained by the electrolysis of fused electrolytes. (b) Electro refining of metals :- A large number of metals (Cu , Ag, Au, Al, Sn) are purified by electrolysis. (c) Electroplating :- Deposition of precious metals (Au, Ag, Pt) on less expensive metals (Cu) is carried out by electrolysis. 9 (d) Manufacture of non - metals :- Electrolysis has been employed in the manufacture of H2, Cl2, F2 etc. (e) Manufacture of compounds :Electrolysis has been used in the large scale production of various compounds such as NaOH, KOH, KCl, Na2CO3, H2O2 etc. Reversible cell :- A cell is said to work reversible (a) If the external emf is equal to that of cell , no current flows and no chemical reaction occurs in the cell. Electrochemistry There are mainly two types of cells : (i) primary cells and (ii) secondary cells. In primary cells, the reaction occurs only once and the battery then becomes dead over a period of time and cannot be used again. (For example, dry cell, mercury cell.) Contrary to this, secondary cells can be recharged by passing a current through them so that they can be used again and again.(For example, lead storage battery, nickelcadmium storage cell.) PRIMARY CELLS The most familiar type of battery is the dry cell which (b) If the external emf is slightly smaller than the emf of the cell, a small current flows out the cell and a small amount of chemical reaction takes place in the cell. (c) If the external e.m.f. is slightly greater than the e.m.f. of the cell current flows into the cell and cell reaction gets reversed. The Daniell cell is a example of reversible cell. * The e.m.f of the Daniell cell is 1.10 volt * If an external e.m.f. equal to 1.10 volt is applied to this cell no current flow through the cell and no reaction will occure. * If external emf is slightly less than 1.10 volts a small current flow out of the cell and the reactions are Zn(s)  Zn+2 + 2e * Cu+2 + 2e  Cu(s) If the external emf is slightly more than 1.10 volts, the curent will flow in to cell and reverse reactions occur :+2 - Zn + 2e  Zn(s) +2 Cu(s)  Cu + 2e is acompact of Leclanche cell known after its discover Leclanche (fig.). In this cell, the anode consists of a zinc container and the cathode is a graphite rod surrounded by powdered MnO2 and carbon. The space between the electrodes is filled with a moist paste of NH4Cl and ZnCl2. The electrode reactions are complex, but they can be written approximately as follows. Anode : Zn(s) Zn+2 + 2e– - SOME COMMERCIAL BATTERIES Any battery or cell that we use as a source of electrical energy is basically an electrochemical cell where oxidising and reducing agents are made to react by using a suitable device. In principle, any redox reaction can be used as the basis of an electrochemical cell, but there are limitations to the use of most reactions as the basis of practical batteries. A battery should be reasonably right and compact and its voltage should not vary appreciably during the use. Cathode : MnO2+NH4+ + e– MnO(OH) + NH3 In the cathode reaction, manganese is reduced from the 4+ oxidation state to the 3+ state. Ammonia is not liberated as a gas but combines with Zn2+ to form Zn(NH3)42+ ion. Dry cells do not have an indenfinite life as acidic NH4Cl corrodes the zinc container even when not in use. Dry cells have a potential of approximately 1.25 to 1.5V. A new type of dry cell which has found use in small electrical circuits (such as hearing aids, watches and camera) is the mercury cell. Here zinc-mercury amalgam is the anode; a paste of HgO and carbon is the cathode. 10 The electrolyte is a paste of KOH and ZnO. The reaction of the cell is as follows : Anode Zn(amalgam) + 2OH–ZnO(s) + H2O +2e– Cathode HgO(s) + H2O + 2e–  Hg(I) + 2OH– ——————————————————— Overall reaction : Zn(amalgam) + HgO(s)  ZnO(s) + Hg(I) Since the overall cell reaction does not involve any ion in solution whose concentration can change, the cell shows a constancy in potential throughout its life. The cell potential is approximately 1.35V. SECONDARY CELLS The most important secondary cell is the lead Electrochemistry Over all cell reaction add equation (i), (ii), (iii) & (iv) Pb + PbO2 + 4H+ + 2SO42– 2PbSO4 + 2H2O Ecell is 2.041 volt (A 12 V lead storage battary is generally used which consist of 6 cells each producing 2 volt) When a potential slightly greater than the potential of battery is applied, the battery can be reacharged. During recharging the cell is operated like an electrolytic cell ie. now electrical energy is supplied to it from an external source. The electrodes reactions are reverse of those that occur during discharge : 2PbSO4 + 2H2O Pb + PbO2 + 2H2SO4 It is clear that H2SO4 is used up during the discharge & produced during reachrging reaction of cell. Another secondary cell is the nickel-cadmium storage cell which has a longer life than the lead storage cell but it is more expensive to manufacture. We shall not go into details of these cells here. FUEL CELLS It is possible to make batteries in which the reactants storage battery. It consists of a lead anode and a grid of lead packed with lead dioxide as cathode. A solution of sulphuric acid (38 percent by mass or having density 1.30 gm cm –3) is used as an electrolyte. When the cell discharges, it operates as a voltaic cell. The lead is oxidised to Pb+2 ion and lead plates acquire a negative charge. At Anode : Pb Pb+2 + 2e– .....(i) Pb+2 ions combine with sulphate ions of solution to form insoluble lead sulphate PbSO4 , when begins to coat lead electrode Pb+2 + SO4 2– PbSo4 (Precipitation) .....(ii) The electrons given by Pb are used at cathode (PbO2 electrode) At Cathode : PbO2 + 4H+ + 2e– Pb2+ + 2H2O .....(iii) Pb+2 + SO4–2 PbSO4 (Precipitation) .....(iv) are fed continuously to the electrodes. Electrical cells that are designated to convert the free energy from the combustion of fuels such as hydrogen, carbon monoxide or methane directly into electrical energy are called FUEL CELLS. One of the most successful fuel cells uses the reaction of hydrogen with oxygen to form water (fig.). This cell has been used for electric power in the Appollo space programme. The water vapour produced was condensed and added to the drinking water supply for the astronauts. 11 In the cell (fig.) hydrogen and oxygen are bubbled through a porous carbon electrode into concentrated aqueous sodium hydroxide. Catalysts are incorporated in the electrode. The electrode reactions are : Anode 2[H2(g) + 2OH– (aq)  2H2O() + 2e– ] Cathode O2(g) + 2H2O() + 4e– 4OH– (aq) ——————————————————— Overall reaction : 2H2(g) + O2(g)  2H2O() This cells runs continuously as long as the reactants are supplied. Since fuel cells convert the energy of a fuel directly to electricity, they are potentially more efficient than the conventional method of generating electricity on a large scale by burning hydrogen, carbon fuels or by using nuclear reactor. Although theoretically we expect 100 percent efficiency in fuel cells, so far an efficiency of only 60-70 per cent has been attained. Since fuel cells are efficient and free from pollution, efforts are being made to produce better commercially practical fuel cells. CORROSION Corrosion is basically an electrochemical phenomenon. The rusting of iron, the tarnishing of silver, development of a green coating on copper and bronze are some of the examples of corrosion. Corrosion causes enormous damage to building, bridges, ships and to many other articles made of iron. We spend crores of rupees every year on this account. Electrochemistry Electrons released at anode move through the metal and go to another spot on the metal and reduce the oxygen in presence of H+ ions (which is believed to be available from H2CO3 formed due to dissolution of CO2 in water). This spot works as a cathode : Cathode O2(g) + 4H+ (aq) + 4e–  2H2O (); (E0 = 1.23 V) Here Fe2+ ions move through water on the surface of the iron object. (If water present is saline, it will help more in carrying the current in the miniature cell thus formed and will enhance corrosion). The overall reaction of the miniature cell is the sum of the cathode and anode reactions as follows : 2Fe(s) + O 2 (g) + 4H +(aq)  2Fe 2+ (aq) + 2H2O(); E0Cell = 1.67 V) The Fe2+ ions are further oxidised by atmospheric oxygen to Fe3+ (as Fe2O3) and comes out as rust in the form of a hydrated iron (III) oxide expressed as Fe2O3, xH2O. 4Fe2+(aq) + O2(g) + 4H2O()  2Fe2O3(s) + 8H+ H+ ions produced in the above reaction help further in rusting. Impurities present in iron also enhance rusting by setting a number of miniature cells. Very pure iron does not rust quickly. In corrosion, a metal is oxidised by loss of electrons to oxygen and forms metal oxide. Corrosion of iron(which is commonly known as rusting) occurs in presence of water and oxygen (air).Although the chemistry of corrosion is complex but it is understood that at one spot of an iron object oxidation occurs and that spot behaves as an anode. Anode Fe(s) Fe2+ (aq) + 2e– ; ( E 0 Fe 2 / Fe  0.44 V ) 12 PREVENTION FROM CORROSION Corrosion of a metal is prevented by applying protective coating (such as greese, paint or metal coatings) on its (metals) surface. In case of iron this coating is done in two ways : (i) electrolysis (Cr, Ni and Cd coating) (ii) dipping iron objects in a molten metal (Zn and Sn coating). Use of zinc coating to protect iron is called Galvanisation. In galvanising iron, zinc being more reactive than irons serves as an anode and is oxidised ( E0 Zn 2  / Zn  0.76 V and E0 Fe 2  / Fe  0.44 V ). Here it is important to note that iron does not rust even after zinc coating is broken which is not true in case of tin coating over iron ( E0 Sn 2  / Sn  0.14 V ). Now if the coating is broken, iron is exposed and iron being more active than both copper and tin, is corroded. Here iron corrodes more rapidly than it does in the absance of tin. But tin protects copper in the same way as zinc protects iron ( E0 Cu2  / Cu  0.3 4 V ). Electrochemistry Instead of coating more reactive metals on iron, the use of such metals is made as sacrificial anode. This method of preventing iron from corrosion is called CATHODIC PROTECTION. In this method, a plate of reactive metal (Zn or Mg) is burried beside the iron pipe or tank and connected to it by wires. Here iron becomes cathode and more reactive metal becomes anode. The reactive metal anode is sacrificed to protect the iron. Since these reactive metal plates are oxidised quickly, they are replaced from time to time which is easy to do. IMPORTANT POINTS 1. Electrolytes also conduct electricity in fused state and undergo decomposition by the passage of electric current. 2. The theory of ionisation was proposed by Arrhenius. 3. Galvanic cells are used to convert chemical energy to electrical energy. 4. The amount of current which produce one gm equivalent of any substance is called faraday. 5. Mg can librate H2 gas from HCl 6. When zinc piece is added to CuSO4 solution, copper gets precipitated because standard reduction potential of zinc is less than that of copper. 7. The electrode at which oxidation occurs is called positive electrode and at which reduction occurs is known as negative electrode. 8. In a galvanic cell cathode acts as positive electrode and anode as negative electrode. 9. In electrolytic cell cathode acts as negative electrode and anode as positive electrode. 10. More negative the standard reduction potential of a metal, the greater is its ability to displace hydrogen from acids. 11. More the reduction potential less is the reducing power. 12. Cell voltage is independent of the size of cell or electrodes. 13. Standard potential is the potential measured at 250C, 1 atm. pressure and one molar concentration of the ions. 13