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ELECTROCHEMICAL CELLS

Structure 13.1 Introduction Objective8 13.2 Electrochemical or Galvanic Cells 13.2 1 A Convenient Notation for UIC Representation of El& and Ceb 13.3 Types of 13.4 Electromotive and its Measurement 13.5 Free Change and Electrical 13.6 The 13.7 Standard Potentials 13.8 Using the Table of Standard Electrode Potentials 13.8.1 SigniT~canceof Positive and Negative Values 13.8.2 Elearode Potential and Stoichimetry 13.8.3 Displacement Reacriom 13.9 Calculation of the EMF of Galvanic Cells 13.10 Applications of EMF Measurements 13.1 1 Corrosion and its Prevention 13.12 Protective Measures Against Corrosion 13.13 Eec trochemical Energy Sources 13.3.1 Dry Cells 13.3.2 Nickel-CadmiumBatteries 13.3.3 Fuel Cells 13.14 Summary 13.15 Glossary 13.17 Answers to SAQs

13.1 INTRODUCTION can be used to bring about chemical reactions in electrolytic cells. The passage of through aqueous solutions of or molten electrolytes provides sufficient energy to cause an otherwise non-spontaneous reduction-oxidation or reaction to take place. This process, called is widely used in electroplating, extractive metallurgy, electrochemical machining, etc. The converse process of conversion of the of a spontaneous redox reaction into electrical energy takes palce in electrochemical or galvanic cells. These are also called voltaic cells in honour of Volta. He demostrated that by interposing sheets of paper soaked in salt water between different metals and connecting the metals, electricity could be produced. Batteries and fuel cells are such devices. The phenomenon of corrosion also belongs to this category. Both in electrolytic and galvanic cells, the redox reactions involve the transfer of charge between the electrode and the . Thus, electrochemisuy can be broadly defined as the science that deals with the consequences of transfer of electric charges from one phase to another. Application of such as the extraction and refining of metals, manufacture of important chemicals like chlorine, caustic soda etc., are quite well known. The increasing awareness to minimise the pollution by automobile exhausts and thermal power plants has rekindled an intrest in batteries and fuel cells. Non-conventional energy sources like the solar energy or wind power can be used to generate electricity which can be used to charge the batteries. The batteries can be used during periods of break down of power and also in running automobiles and mopeds. 0 bjectives After studying this unit, you should be able to : * understand how a redox reaction can be made to produce in a , * represent such a galvanic cell and to explain by means of electrode reactions how the cell produces an electric current, * identify the anode, cathode, psitive electrode and negative electrode of a' galvanic cell, * know how the standard electrode potentials are determined and how to use these to calculate the standard E.M.F. of a cell, * know how to use the relationship between free energy change and E.M.F. and to calculate the equilibrium constant from the standard E.M.F. of a cell, * use Nernst equation to calculate the of half cells or the E.M.F. of cells in which the reactants and products are not in their standard states. * Explain the working of a pH meter and use EMF measurements for measuring pH and solubility products, * appreciate the usefulness of potentiometric titrations, * Analyse the advantages and limitations of various electrochemical energy sources, and * explain what galvanic corrosion is and how it may be minimized.

These are devices in which the free energy change accompanying a redox reaction is converted into electrical energy. If a zinc rod is dipped into a solution of copper sulphate, a brownish-red deposit of copjgr is formed on the surface of zinc and the bluish-green colour of the copper sulphate solution disappears. The redox reactions that occur can be represented as oxidation : ~n(s)-2e+~x?+(aq) (13.1) reduction : cu2+(aq)+2e+Cu(s) (13.2) The overall reaction is represented as (13.3) Zn(s) +C$+(aq) +~x?+(aq) +Cu(s) (13.3) In reaction (13.3), only the active constituents taking part in the reaction are indicated. The sulphate ions do not take part in the reaction. At 298 K and 101.3 kPa (i.e., lam) the standard free energy change ( AGO ) corresponding to this spontaneous reaction is -212.9 kJ. Though AGO is a measure of the electrical work that can be derived from this reaction i.e., Eq.(13.3), the transfer from zinc to Cu2+( aq ) can be demonstrated only when the oxidation and reduction reaction are allowed to take place separately as in the Daniel1 cell, which is a Galvanic cell. [Figure 13.11. This cell consists of a zinc rod (electrode) dipping into a solution of zinc sulphate (one can use an aqueous solution of sodium chloride also) and a copper rod (a platinum electrode can also be used ) dipping into a solution Of capper sulphate. A porous barrier prevents the mixing of the two electrolytes [ Figurel3.l(a)] but allows the passage of the ions. The two electrolytes may also be kept in separate beakers and a enables the passage of ions from one compartment to another [Figure 13.1(b)]. When the two electrodes are connected by means ~f a wire (electronic conductor), oxidation (reaction-1) takes palce at the zinc electrode. The zinc ions dissolve in the Elwetron flow EMr~&emlerlCdls

Porous borrlrr

!(a) Solutlow equated by a porous bad=

so$-

-

(b) Sdutlons separated by a salt brldge Flpre 13.1 :Laboratory Verdons of the electrolyte and the left behind on the electrode push other electrons via the connecting wire to the copper electrode. These electrons are used by cu2+ions in solution and reduction (reaction-2) occurs. The electron flow from the zinc electrode to the copper electrode provides a source of electricity. The direction of the current flow, as measured by a current measuring device, is opposite to that of electron flow. Sirice oxidation takes place at the zinc electrode it is called the anode. In this galvanic cell, this electrode, being a source of a negative charge, is the negative electrode. The copper electrode at which reduction occurs is the cathode and since it is a sink for electrons (accepts electrons), it is the positive electrode. As the cell continues to produce current, the zn2+ions entering the electrically neutral solution make the solution positively charged. The zinc electrode acquires a nagative charge with respect to the solution. As a result of the build up of charges, an electrical double layer is established at the electrode-electrolyte interface. The resultant potential difference between the electrode and electrolyte is called the electrode potential. The negative electrode potential on the zinc electrode will prevent the zinc ions from the metal lattice leaving the electrode. At the other electrode, the cuZ+ions in solution are used up in the reduction (reaction-2). The remaining sulphate ions tend to make the solution Acquire a negative charge. Since electrons are removed from the copper electrode it acquires a positive potential and so prevents the approach of cu2+ions to the electrode (cathode). A continuous flow of current can be maintained if electrical neutrality is maintained at both the compartments A and C. The excess of sulphate ions from C move through the barrier into A so as to neutralize the excess positive charge of the solution in A. The sulphate ions do not undergo any chemical change. Instead of a porous partition, a salt bridge is often used. This consists of a glass tube containing a concentrated solution of KC1 or NH,NO, gelled by adding gelatin or Equilibria & Electroehmlstry agar. The gel confines the electrolyte to the tube, prevents the mixing of the electrolytes and allows the passage of ions. Here, K+or Neions move to C and the anions ( Cl- or NO; ) to A so as to maintain electrical neutrality. The high concentration ensures that most of the current is carried by the ions of the electrolyte used in the salt bridge. These ions also carry equal shares of the current and do not undergo any reaction with the electrode. 13.2.1 A Convenient Notation for the Representation of Electrodes and Cells Any like the Daniell cell consists of two electrodeelectrolyte assemblies and each of these is called a half cell or simply an electrode. Oxidation occurs at one half cell and reduction at the other. The cell reaction is simply the sum of the two half cell or electrode reactions. A convenient notation is generally followed to represent on paper the electrodes and cells. A single vertical line denotes the boundary between two phases. In the case of aqueous solutions, the concentrations of the ionic species are indicated in parenthesis. The Daniell cell [Fig.13.1] can be represented on paper as ~n 1 w+(1.0~1 II cu2+( 1.0~) I cu The double vertical line between the electrolyte solutions indicates that a barrier [Figure 13.l(a)] or a salt bridge [Figure 13.1 (b)] has been used. The half cell or electrode at which oxidation occurs, i.e., anode, is written on the left hand side, and the cathode on the right hand side. An easy way of remembering all that has been said about galvanic cells is given below, using the Daniell cell as an example.

L H S electrode B H S electrode Qxidation occurs Reduction occurs Anode Cathode Uega tive Electrode Positive Electrode It is simpler to remember the letters "LOAN" which summarizes the conventions. It will be noticed that the underlined alphabets on the LHS appear earlier than the corresponding ones on the RHS in alphabetical order. In general, irrespective of how the electrodes are represented, a galvanic cell should be represented on paper as Anode 1 Anode solltion 1 1 Cathode solution 1 Cathode (C, (GI Thecell reaction in the case of the Daniell cell can be written as Cathode : cu2++2 e = cu Anode : a'-2 e = ~n2+ Therefore, Zn(s) +cu2+=Zn2++~(s) You ought to try and think of alternative conventions of representing cells, which may help you to remember the conventions better. 13.3 TYPES OF ELECTRODES There are three categories of electrodes (half cells). The electrodes of the first kind iqclude a metal electrode in contact with its ions or a gas (G) bubbled over the surface of an inert electrode like R,dipping in a solution containing G or G+ ions. In the case of gases, the pressure of the gas should be specified. The hydrogen electrode can be represented as

since the gas at a pressure of 1 atm and W = 1.OM conform to standard conditions, the hydrogen electrode is called the Standard Hydrogen Electrode (S.H.E.) [Figurel3.2]. More examples are given in Table 13.1 Table 13.1 : Representation and Electrode reactions in the case of Gas Electrodes

All electrode reactions have

Electrodes of the second kind consist of a metal coated with layer of an insoluble salt of the metal dipping in a solution containing the common anion. Some of these are given i0 Table 13.2

Table 13.2 : Representation of Electrodes of the Second Kind

Electrodes of the third kind are called redox electrodes. In these electrodes, an inert electronic conductor like Pt dips into a solution containing two different oxidation states of a species.

Glass tube -Hydrogen C 1 atn >

Pt electrode coated wi- Finely divided Pt

* Figure 13.2 :A diagrammatic representation of a Standard Hydrogen Elecvode (SHE)

Pt I Fe2+(c,) ,Fe3+(c2) ~t I sn4+(c',),sn2+(C',) The order of writing the two species is immaterial. Equlllbria & Electrochemistry SAQ 1 In the following reactions, identify the reactants that undergo oxidation. a) Zn(s) +2AgCl(s) ->ZnC&+2Ag(s) b) Fe(s) +2~+->~e~++~,(g) c) 2cuZ++4r->~cuI+& d) I,+s~+--->s~~++~I-

SAQ 2 Study the cell diagram of a galvanic cell and also the electrochemical terms arranged in an alphabetical order below the LHS compartment. Can you suggest a suitable alternative mnemonic (easy way of remembering) by concentrating on the underlined alphabets. Zn I ~2+(1.0 M) II cu2+(1.0 M) 1 CU Anode Cathode LHS electrode RHS electrode Negatively charged Eositively charged Qxidation occurs Reduction occurs If it is an will the same hold good ?

13.4 ELECTROMOTIVE FORCE AND ITS MEASUREMENT A flow of charge occurs only if there is a difference in the electrical potential. In the case of the Daniel1 cell, a flow of from anode (Zn) to the cathode (Cu) occurs, since the negative potential at the anode is more negative than that at the cathode. The potential difference which is respnsible for forcing the electrons to flow from the negative electrode to the positive electrode is called the electromotive force (EMF). Potential is energy per unit charge. If energy is expressed in (J) and charge in (C), both potential and potential difference are to be expressed in J/C or (V). A voltmeter connected across the terminals of the two electrodes will give the of the cell. This depends on the current drawn from the cell by the voltmeter. As the cell reaction progresses, the voltage will be found to decrease. The voltage will be a maximum only when the current drawn is zero. The voltage under zero current conditions is the EMF of the cell. For this purpose, a potentiometer lFigure13.31 is used. In this instrument, the EMF of the cell is balanced by an opposing EMF from within the potentiometer. When the two EMF'S are equal, no current flows, and thus the EMF is measured when the current drawn is nil. A storage battery is connected across the potentiometer wire NP Lhrough a variable resistance R. Both the standard cell (S) having a known and constant EMF ( E, ) and El~rtrochcmicalCells

Flpre133 :Measurement of EMF using a Patentlometer the cell X whose EMF ( Ex ) has to be determined are connected as shown in Figure 13.3. The negative pole of each cell is co~ectedto the negative end N of the potentiometer. The positive and of each cell is to be connected through a galvanometer (G)to the sliding contact, C. The EMF of the cell is measured by comparing its EMF with that of a standard. The standard cell used for this purpose is the Weston cadmium cell which can be represented as : Cd ( Hg ) 1 3 Cd SO, . 8H20( s ) 1 Cd SO, ( saturated solution) I Hg2S0, ( s ) I Hg It has an EMF of 1.0146V at 298 K and does not vary much with temperature. With the (S) included in the circuit by the DPDT switch and the sliding contact is moved along the wire till there is null deflection in G. If NC is the length of the potentiometer wire at which no deflection is observed in G, the potential drop across NC is balanced by the EMF, E,, of the cell S. The purpose of R is to adjust the resistance in such a way that the balance point occurs when NC is equal to 1014.6 cm. The potential drop will thus be 1 millivolt/cm of the potentiometer wire. Now the cell is brought into the circuit using the DPDT switch and without disturbing R, the sliding contact is moved till the galvanometer reads nil deflection. Let ND be the length of the potentiometer wire when there is nil deflection. The EMF of the unknown cell is given by Eqn. 13.1

EMF measurements using a potentiometer are also useful in hawing the polarity of the electrodes. In the case of the Daniel1 cell, only when the zinc electrode (negative electrode, anode) is connected to N and the copper electrode (positive electrode, cathode) to the sliding contact it will be possible to find a balance point. U the connections are reversed no balance point will be found. In a balanced potentiometer circuit, the electrode connected to N is the negative electrode (anode) and the other one is the positive electrode (cathode) of the galvanic cell. U the external EMF is decreased slightly, the cell reaction in the galvanic cell occurs in the forward direction so as to force a current through the potentiometer circuit. However if the external EMF is increased slightly, a current is forced through the galvanic cell and the reverse reaction will be favoured. Such a cell is said to be thermodynamically reversible.

13.5 FREE ENERGY CHANGE AND ELECTRICAL .WORK At constant temperature and pressure, the change accompanying a process is equal to the reversible work, other than the work of expansion,(Unit 7). This is the electrical work in galvanic cells. When a charge of q coulomb is transported under the influence of a potential difference of E volt, the electrical work I.q~~ilil)ri.~S I~:lcc~rtrhernlstry done on the surroundings is qE volt- coulomb or qE . If a redox reaction involves n mol of electrons, q = nF where F is the Faraday. Hence the electrical work done by the cell on the surroundings ,W" WSur= n F E Joules (13.2) Since the cell is thermodynamically reversible, this is also the maximum electrical work, which is equal to -AG. Thus Eqn. 13.2 becomes AG-W"=-nFE (13.3) In Eqn. 13.3, n is the number of moles of electrons involved in the cell reaction, F is the Faraday and E is the EMF of the cell. The EMF of any galvanic cell depends on (i) the nature of half cells used to form the cell, (ii) the concentrations of the species involved in the redox reaction, and (iii) temperature. For the sake of comparison, rhe temperature is usually 298 K and the reactants and products are assumed to be in the respective standard states (Unit 6 and 7). All solids and liquids must be pure and in their stable forms. The gases should be at a pressure of 101.3 kPa (1 am) and the activities of all the species in solution must be unity. However, as an approximation, a concentration of 1.0 M can be used instead of unit activity. The EMF and the corresponding free energy change under the standard conditions are E0 and AGO respectively. Thus for standard conditions AGO=-nFEO (13.4) Eqn. 13.4 provides the link between thermodynamics and electrochemistry. Example 13.1 : For the galvanic cell, zn I zn2+ (1.0~1)I CI-(1.0~1I ~g~l(s)I ~g the EMF at 298 K is 0.985 V. (a) identify (i) the anode, (ii) the cathode (iii) the positive electrode (iv) the negative electrode. (b) Write down (i) the electrode reactions and (ii) the cell reactions and (c) calculate AGO at 298 for the cell reaction. Solution: (a) (i) LHS elecuode viz., Zn is the anode (ii) RHS electrode viz., Ag is the cathode (iii) The Ag electrode is the positive elecuode (iv) The Zn electrode is the negative electrode (b) (i) cathode : AgCl(s)+e'=Ag(s) +Cl- 1 anode: Zn - 2 e- = zn2+ 2 (ii) For the purpose of making the number of electrons gained equal to the number lost, eqn(1) is multiplied by 2 and added to Eqn (2) to get the cell reaction Zn(s) +2AgCl(s) =z~*++~cT+~A~(s)

(c) AGO=-nFEO Here n = 2, F = 96500 C and E0 = 0.985V

AGO=-2mole-~96500- C (0.985V) mol e' = -190105 J or -190.1 1 kJ 13.6 THE NERNST EQUATION Electrochemical Cclb The free energy change of a reaction depends on the concentrations (activities) of reactants and products. If the same reaction occurs reversibly in a galvanic cell, the EMF of the cell will also depend on the concentrations of the reactants and products. For a cell reaction aA+dBzmM+nN the free energy change is related to the standard free energy change as ( Eqn. 7 of Unit 7)

h Eqn 13.5, Q, the reaction quotient, has the same form as the expression for the equilibrium constant K. However, the concentration terms used in the expression for Q are arbitrary and do not correspond to those in the equilibrium state. From equations 13.3, 13.4 and 13.5,

Here EOis the standard EMF and E is the EMF under non-standard conditions. The standard EMF ( E0 ) is a constant for a given cell and at a pressure of 101.3 kPa, varies only with temperature. Substituting for R,T and F in Eqn. 13.7 we get Eqn. 13.8

Equations 13.6 to 13.8 are different forms of the Nernst equation. All forms of the Nernst equation are applicable not only to cell reactions but also to electrode reactions. When the concentrations of all the species are unity, Q=1, log 1 = 0 and so E = EO.When the concentrations are not unity (non-standard conditions) E is not equal to EOand the Nernst equation (Eqn. 13.6 to 13.8) should be used to calculate E. If all the species are at their equilibrium concentrations, Q = K, AG = 0 and E = 0. Eqn 13.6 can then be written as

In a similar fashion, Eqn. 13.8 can be written as K= 10(nlf/0.0591) Thus, it is possible to calculate the equilibrium constant of a reaction if it is allowed to take place in a galvanic cell under standard conditions. For example, the reduction reaction.cu2++ 2e+ -Cu, the Nernst's equation is given by

El = E: + 9log [cu2+] 2 for the reduction reaction zn2++ 2e- -- + Zn huiiibri8 EicctdclnYq The cell potential for the cell reaction a+cu2+ + ZnZ+ + cu, is given by the difference of the two reduction potentials expressed by Eqns (1) and (2) , 0.0591 Ed = (E! - @ + -2 [log (&+) - log (2h2+)]

For any cell reaction it is easy to write the Nernst eqn by keeping in mind that the Q term is similar to the expression for the equilibrium constant but set at actual concentrations and not equilibrium concentrations. Example f 3.2 Calculate the half cell potential at 298 K for cu2+(aq) + 2e- ----+ Cu (s), if (~$9= 10.0 M and E0 = + 0.336 V. Solution:

E=EO--0.0591 [Cu] 2 log [Cd+]

= 0.336 V 0.0591 log 10 + -2 = 0.3655 V Example 133 Calculate the equilibrium constant of the cell reaction given in Example 1. Solution: Cell reaction : Zn(s) +2AgCl(s) =Zn2++2~g(s)+2CT K= [zn2+] [CI-]~=~X~(~FEO/RT)

- exp [2 965m x 0q985 up ( 76-731 8.314 x 298 )=

= 2.14 x ld3 Example 13.4 For the galvanic cell ~d I C~~+(O.OSM)1) CT(O.~OM) I Cb(latm) I ~t calculate the EMF at 298 K, if E0 = 1.76 V Solution: The concentrations are arbitrary and also nonstandard. The Nernst Eqn. (Eqn. 13.8) can be applied.

The cell reaction can be written as ~d(s)+cI,(~) =~d++2~1-(n=2) Note that C%gas at 1 am pressure and Cd(s) are in their standard states and so their ~oehemkdCdb activities are unity.

SAQ 3 1. If the EMF of a cell under standard conditions is 1.00 volt, what will be the values of AGO and K if the cell reaction involves 2 moles electrons ?

2 Given that EzeI,is + 1.10 V. Calculate the EMF of the cell in which the cell reactionisZn(s) +cu2+(0.02~)-+Cu(s) +zn2+(0.40~)

13.7 STANDARD ELECTRODE POTENTIALS It is not possible to measure the potential of a single electrode since the solution can not be directly connected to the potential measuring device (electrode) in the solution. What is measured as EMF is the potential difference between the two electrodes. If single electrode potentials are available, it would be possible to predict the EMF of any combination of two half cells. In order to assign values to single electrode potentials we can use the approach that enabled us to evaluate AG and AH without knowing the absolute values of G and H respectively (using standard states as in Units 6 and 7). It is assumed that the electrode potential of a standard Hydrogen Electrode (SHE) (Sec 13.3) is exactly 0.000 V at all temperatures. What one has to do is to construct a cell consisting of a SHE and another half cell, whose electrode potential is required, and measure the EMF of the cell. The measured EMF (E) is a difference between the two electrode potentials (E) i.e.,

It follows that EhalPcellis equal to E in this case since EsM is zero. While determining the EMF, the SHE may have to be connected either to the negative end N or to the positive end P of the potentiometer circuit. For example, two cells in which SHE is a half cell, are given below and the polarities of each electrode are also indicated.

At 298 K, the EMFs of cell - 1 and cell - 2 are found to be 0.763V and 0.340V mpectively. The cell reactions are spontaneous in both cases. The electrode potentials of the zinc electrode and copper electrode should thus be 0.763 and +0.340V respectively. Since the electrode potential is related to the free energy change, the sign and magnitude of the electrode potential should reflect the tendency of the electrode reaction to proceed spontaneously. For example in cell-1, at the zinc electrode the oxidation of Zn to i@ is spontaneous whereas at the copper electrode of cell-2, it is the reduction of cu2+to Cu that is spontaneous. This has led to the adoption of two sign conventions. In the American convention, all electrode reactions are written as oxidations and if the oxidation is spontaneous, a positive value is assigned to the electrode potential. If the oxidation is not spontaneous, a negative value is assigned to the electrode potential. For example, Zn(s)-2e-=&+;~~~=0.763V CU(S)-~~-=CI?+;&Dox =-0.340V Thus, electrode potentials, according to the American convention, are oxidation potentials. In the European convention, all electrode reactions are written as reductions. If the reduction is spontaneous, a positive value is assigned to the electrode potential and if it is not, a negative value is assigned. For example, zn2++2e-=Zn(s) ;cod =-0.763V Cu2++2e-=Cu(s) ;cod =+0.340V According to the IUPAC (The International Union of Pure and Applied Chemists) recommendation, half cell potentials corresponding to reduction potentials are to be called Standard Elecuode Potentials. A table of standard electrode potentials is given below (Table 13.3). I

Table 13.3 :Standard Electrode Potentials at 25" C

Equiiibria&EwochmMrg 13.8.1 Significance of Positive and Negative Values A positive value for E0 in this table indicates that the reaction Ox + ne = red, is spontaneous. Since oxidising agents undergo reduction (electronation), this would mean that the species "Ox" having a higher oxidation state than the " red " is an oxidising agent. Greater the positive value, greater is the strength of the oxidant. Thus among halogens, the oxidising ability decreases as F2(+2.870)>C12(+ 1.358)>Br2(+ 1.065)>$(+0.540) The E0 values are indicated in parenthesis. Another way of stating the same thing is that, in the half cell with a more positive potential, reduction occurs. Similarly a negative value for E0 indicates that the reduction reaction, Ox + ne = red is not spontaneous but the reverse reaction is. The species " red " in a lower oxidation state is a reducing agent. More negative the value of EO,greater is the strength of the reducing agent. Thus the reducing ability decreases as Na(-2.711) >Zn(-0.763) >Fe(-0.441) >H2(0.000) 13.8.2 Electrode Potential and Stoichiometry The potential of a half cell or an electrode is not related to the stoichiometric coefficients in the electrode reaction, since it is an intensive property. This means that the following half cell reactions cu2++2e- SCU(s)

have the same E0 value of +0.340V, at 298 K (refer to Example 13.4). If the half cell reaction is written in the Teverse direction, the sign of E0 changes. Thus

z$+ + 2e- Zn: EO = -0.763 V

but Zn = Zn2++ 2e-: EO = +0.763 V i.e., = - oxidation potential. 13.8.3 Displacement Reactions Metals like Mg, Al, Zn. Fe etc., tend to lose electrons readily to become the corresponding cations. Ions like cu2+,Ag+ ,~g~+ etc. exhibit a tendency to accept electrons. It is because of these opposing tendencies that the familiar displacement reactions such as . zn+cu2+ zn2++Cu

occur spontaneously. In general, a metal X having a more negative E0 value (higher up in the EMF series) will displace another metal Y having a positive or less negative E0 value from a solution of its salt. In fact, reaction of metals like Zn, Mg, Fe etc. with dilute acids to give hydrogen from acids can be considered as a displacement reaction. One can understand why metals like Cu and Ag do not liberate hydrogen from acids. These displacement reactions can also be made to take place in suitable galvanic cells and we shall now see how to calculate the EMF of such cells.

13.9 CALCULATION OF THE EMF OF GALVANIC CELLS As indicated earlier (13.2). the galvanic cell is always written as, electrode 1 ion I 1 ion 1 electrode, whatever be the manner in which the half cell is represented in Table 13.3. For such a galvanic cell, the EMF under standard conditions is Eleebochemkid Cells given by Eqn 13.1 1 E0 = qHS- EbS = ELe- Eonnode (13.1 1) In Eqn. 13.1 1, the E0 values are reduct ion potentials. Examples: Cells :E0 = Ebs - GS

(ii) Pt 1 s$+, sn4+ 11 Fe3+,FC?' I Pt EO= -770- ,150

i The cell reactions are spontaneous if the E0 values are positive and non-spontaneous i if the E0 values are negative. For writing the cell reaction, the reaction at the MS electrode is written as reduction and that at the LHS electrode as oxidation. The number of electrons gained is made equal to the number of electrons lost (see example 1). taking care to see that the corresponding electrode potential values are not multiplied or divided. The two half cell reactions are then added to give the cell reaction (Vide. example 1)

Example 13.5 : The standard electrode potential of Ag-AgC1 electrode on the hydrogen scale is +0.222 at 298 K. For the cell Fe 1 ~g'(l.0~)I I Cl-(1.OM) I AgCl(s) 1 Ag the EMF was found to be +0.663V at 298 K. What is the standard electrode potential of F$+/F~on the hydrogen scale'? Solution: EO=E;- EL + 0.663 V = 0.222 - Therefore et=-0.441 V Example 13.6 : I Forthecell Zn 1 Zn2+(1.0~)(1 6?+(1.0~)I Cr I (i) Calculate the EMF under standard conditions (ii) Write down the cell reaction and calculate the AGO value corresponding to the cell reaction. The standard electrode potentials of zn2+ 1 Zn and c?+I Cr are -0.763 and -0.740 V respectively. Solution : i. EO=Ehs-Gs = - 0.740 - ( - 0.763 ) = + 0.023 V ii. (i) RHS electrode : cPi+3e-=~r; (ii) LHS electrode : zn=zn2++2c;

Let AG: and AG; be the standard free energy changes for reactions 1 and 2 respectively. In order to balance the number of electrons, Eqn. (1) is multiplied by 2 and Eqn. 2 by 3

Adding,

3Zn(s)+2~?+$ 3Zn2++22(s); AGO = - ~FEO Since AGO = AG; + AG; -6FE0=- 6F(-0.740+0.763) = - 6F (-0.740 + 0.763)

EO = + 0.023 V A GO = - 6 x 96500 x 0.023 v =- 13317 J or - 13.32kJ It must be realised that the predictions regarding the spontaneity of a cell reaction are based on thermodynamic considerations. However, the rates at which the cell reactions take place can not be easily predicted. SAQ 4 Consider the cell I zn2+(l.0M) II cu2+(1.0~)I cu The standard reduction potentials ate + 0.35V for cu2++ 2,- +Cu and - 0.763 V for zn2+ + 2e- --+ Zn Write down the cell reaction. Calculate the E.M.F. of the cell.

SAQ 5 Oxidation potential ~f Zn (s) / ~n+~is 0.76 and the oxidation potential of Ag(s) / is -0.80 volts. What will be the emf of the Zn-Ag cell ? Predict whether the cell is feasible or not. SAQ 6 For the cell, Ag 1 Ag' 1 1 1 CU,find out the emf of the cell and state whether copper will displace silver from a solution containing silver ions.

SAQ 7 1. a) Will Cu react with dilute H2S04to give H2 ? Justify your answer. b) Mention the three metals which liberate Ag from AgNQ.

c) . The oxidation potential of the redox couple M I M~+ is + 1.50 V. Is M an oxidant or a reductant ?

2. The standard electrode potential of C& + I Cd on the hydrogen scale is - 0.403 V. If instead of SHE, SCE ( EO = 0.242) is used in combination with the Cd 1 C6L + couple, what will be the EG,, ?

13.10- APPLICATIONSOF EMF MEASUREMENTS By measuring the EMF of suitable galvanic cells, it is possible to calculate the changes in the thermodynamic functions like AH and AS and the solubility product and the pH of a solution. EMF measurements are also useful in analytical chemistry. 13.10.1 Evaluation of Thermodynamic Functions In this method both AG and AS are determined from EMF measurements and AH is . evaluated from the Gibbs - Helmohltz equation. EMF measurements are normally . made at constant temperature. Since EMF is affected by changes in temperature, the dE temperature coefficient of EMF viz., -dT can be determined by measuring the EMF at 2 or 3 different temperatures. It is related to AS as shown below. Differentiating Eqn. 13.3 w.r.t. temperature

since d(AG)/dT 7 -AS (Unit 7) we have

\ / AH at any given temperature can be calculated from the'aibbs-~elrnholtzequation AG=AH-TAS (13.12) Example 13.7: fie galvanic cell ~p'(s)I AgCl(s) I KCI(l.OM) I H&c~(s)I HgO) is an example of cell without a liquid junction since the two electrodes dip into the same electrolyte. The EMF of the cell is 0.058V at 298 K and 0.0614V at 308 K. (a) Write down the cell reaction. (b) Calculate (i) ASo (ii) AGO and (iii) AHoat 298 K. Solution: (1 RHS electrode : Hg2 C12 + 2e' ;f 2 Hg ( 1 ) + 2 Cl- LHSelectrode:2Ag(s)+2Cl' 2 2AgCl+2e The cell reaction is 2 Ag ( s ) + Hg2C12 (s ) 2 2 Hg ( 1 ) + 2 AgCl ( s ) (b) (i) -dE -AE dT AT - 0.0614 - 0.058 10 = 3.40 x 10-4 VK-' From Eqn. 13.11,

AS0 = nF -dE dT =2x965000C ~3.4~10-4 VK-' = 57.9 JK-' mo~' (ii) AGO = -nFEO = -2 x 96500 C x 0.058 V = -1 1194 J mol-' (iii) AGO = A* - T ASO AHO = AGO + T AS0 =-1 l194+ 298 x 57.9 = 6060 J mol-' Note : Had the cell reaction been written as A~(s)+~H~~cI,(s)2 Hg(l)+AgCl(s) 2 the values of AS', AGO and AW will-be 28.95 JK-I mol-1 - 5597 J mol-I and 3030 J moT1 respectively. In this case, the values are per mol of AgCl formed. In the former case, the values are per mol of Hg2CI, ( s ) reacted. 13.10.2 Determination of Solubility Products of Sparingly Soluble Salts The solubility product K,, (Sec. 14.15) is also an equilibrium constant and its magnitude can be predicted from a knowledge of the standard electrode potentials, so chosen, that the overall cell reaction is the solubility product equilibrium. One can also form a suitable galvanic cell and determine its EMF experimentally. Even if measurements @n not be made, it is possible to arrive at an approximate value of KBp from the appropriate half cell reaction potentials as shown below. Example 13.8: Calculate K, for Hg2C12 ( s ) at 298 K from data in Table 13.3 Solution : H&Cl,(s) 2 H&++~cI-;

K,=[H&+] [~rl2 One should look for a reaction in which Hg2C1, takes part and another reaction in which H$ takes part. These are Eloctradtemlcal Cells

Reaction (1) - Reaction (2) gives

From Eqn. 13.9 or 13.10, K, can be calculated

13.10.3 Determination of the pH of a Solution The electrode reaction in the case of a hydrogen electrode is H+ + e = -1 H,. Nernst 2 eqn. (Eqn. 13.8) can be used to calculate the electrode potential of a hydrogen electrode.

1 Since PH2= 1 atm and %,, = 0 ,this equation can be written as EH=O-0.0591 (-log[H+])=-0.0591pH (13.13) A hydrogen electrode dipping into a solution of unknown pH is combined with a reference electrode like the saturated calomel electrode (SCE) and its EMF is measured. Pt I H2(l aun) 1 [Ht]=x 1 IKCI I Hg2C12(s) 1 Hg E=ER-EL=0.242-E, (13.14) In Eqn. 13.14, since E is known and the EE ,, being + 0.242 V, E, can be calculated. This value is substituted in Eqn. 13.13 to obtain the pH of the solution. The commercial pH meters make use of a glass electrode-SCE assembly to measure the pH. The glass electrode consists of a thin glass bulb of a special quality glass containing a solution of a constant pH into which a reference electrode dips. The cell assembly can be represented as Ag 1 AgCl (s) I 0.1NHCl 1 Glass I experimentalsoln I I SatdKCI I Hg2CI,(s ) I Hg glass electrode SCE The potential of a glass electrode depends on the difference in pH on either side of the glass electrode. Since the solution in the glass electrode has a constant pH, the potential depends on the pH of the experimental solution. First of all, the glass electrode is immersed in a buffer solution and the pH meter adjusted so as to read the pH of this buffer solution. After washing, rinsing etc.. it is then immersed in the experimental solution and the pH of the solution is read off from the instrument dial. The glass electrode is an example of an ion- selective electrode. By making suitable changes in the composition of glass, a glass electrode can be made selective to other cations like Nat ,Kt, NH, +,etc. Compacted discs of AgCl, AgBr, AgI, Ag2S or Lq are found to be selective to the respective anions. 13.10.4 Potentiometric Titra tions One of the widely used applications of EMF measurements is to detect the end point of a titration by measuring the EMF of a cell consisting of an indicator electrode (electrode, whose potential depends on the concentration of the reactant ions) and a reference electrode (SCE) as the titration progresses. This is called a potentiometric titration. Since the electrode potential of the reference elecuode is constant, the observed change in EMF as the reaction progresses is due to the change in the electrode potential of the indicator electrode. During acid-base titfatiom, the pH of the solution changes and by monitoring the change in pH as the titration progresses. it is possible to detect the end point in acid-base titrations using a pH meter. Very near the equivalence point, the change in pH is quite%ge. In a redox titration between Fe?' and ~e~.the overall reaction ce'++ ~2'3 Ce* + F$+ can be discussed interns of the two half cell reactions cect+e- 2 Ce?' ; E&=+1.610 V and ~2++e-2 Fez+ ; &=+0.77lV Initially when no Cect is added to the system, the potential of the electrode. characteristic of the Fe?, Fe?+ couple. is given by Eqn. 13.15

When all the Fez+ has reacted. if a slight excess of ~e*is added, the potential is characteristic of the ce4+, ~2'. system and is given by Eqn. 13.16 E& = E& - 0.059 log uce" [ a*+I The variations in the reduction potential when 100 ml of 0.1N Fe?+ is titrated against de4+of the same normality are tabulated on next page : By plotting the potential against volume of &+added, Rgure 13.4 is obtained. Tim equivalence point corresponds to the point of inflection of the curve. A better procedure to obtain the equivalence point is to make use of the fact that WAVis large in the vicinity of the equivalance point. If AWAV is plotted against V (Figure 13.5). the sharp maximum corresponds to the equivalence point.

Titrant added (V) Piprre 13.4: ApotentiawYic-anrn

v Fipn 135 :Dasmrinrbn of Uw squivhm point ia r pombnderilnrlm Vol of ce4' added . Excess ml of E in volt ~e~' ce4'

100.0 - - - - equivalence point 1.19

It will be seen that the change in potentid within s.l ml of the equivalence point is quite high, i.e., 0.485V (1.433 - 0.948) Equilibria& Elcet~ahemMry 13.11 CORROSION AND ITS PREVENTION Metals like iron are widely used as structural materials. Metals vary in their reactivity to oxygen, moisture and other substances in the environment. Metals like Au and Pt are not affected by the substances present in the environment and are called noble metals. It is partly because of this resistance to environmental attack that these metals are so sought after, and therefore expensive. On the other hand, when exposed to atmosphere, a piece of iron rusts, objects made of copper are covered with a green deposit and zinc becomes coated with a white deposit. These are the products resulting from their reaction with the environment. The term corrosion is generally used to describe the deterioration of a metal as a result of its interaction with the environment. The corrosion of iron is called rusting. Rust is hydrated ferric oxide and as such provides little protection to the metal and makes it useless as a structural material. It has been estimated that nearly 15% of the annual production of iron is mainly used to replace the rusted iron. The problem of rusting of iron has assumed serious proportions in view of the increasing pollution of the environment. The corrosion of iron will be discussed in this section. In order to combat the problem of corrosion, one should understand the factors that promote corrosion so that steps can be taken to avoid or minimise these. The factors that promote corrosion are (1) oxygen and moisture. (2) electrolytes dissolved in water, and (3) contact with metals like copper. The recognition of the redox nature of corrosion led to the electrochemical theory of corrosion. According to this theory, a potential difference exists on the surface of a metal because of small differences in the composition of the metal, such as lattice defects, impurities etc. This potential difference is responsible for the corrosion process, resembling the operation of a galvanic cell, to take place on the metal surface. At the anode of these galvanic cells, the oxidation reaction, Fe -2 e- = ~e~+occurs, and the ~2'goes into solution in the water formed by the condensation of water vapour from the air condensing on the surface. The electrons left behind at the anode flow to the cathode of the galvanic cell, the metal surface acting as the electronic conductor. At the cathode, a number of reactions are possible. Some of the important reduction half- reactions are

1) H* + e- 2 H~ ( ) ( acid solution) ( EO = 0.00 ) 2 v 2) 02+4H++4 e- 2 2 H20 ( acid solution) ( E'= 1.23 V ),

3) O2 + 2 H20 + 4 e- 2 4 OH(neutral or alkaline solutions) ( EO = 0.40 V ) In the galvanic cell, the circuit is completed by the ionic conduction through the aqueous solution af an electrolyte. This explains role of an electrolyte in promoting corrosion. The formation of anodic/and cathodic areas or sites on the surface of a metal may arise as a result of the following : (1) roughness on the metal surface (2) presence of impurities in the bulk metal, and (3) differences in the availability of oxygen along the metal surface (1) The head and the pointed end of a nail are rough surfaces. These uneven surfaces have crests and troughs. Since the metal atoms at the crest are more weakly bound compared to those at the trough, these are oxidised more readily. Thus trough areas act as anodic sites and the body of the nail acts as the cathodic sites. The formation of ~e~+at the anodic sites and OHat cathodic sites can be demonstrated using the ferroxyl indicator (a mixure of phenolphthelein and potassium ferricyanide.) If the iron nail (free from rust) is placed in a Petri dish containing a dilute solution of NaCl (3%) containing 1 or 2 drops of ferroxyl indicator, it will -be noticed i;i,t~ aker some time the solution near thc pointed end and the had regions of the nail would have become blue whereas that near the body pink. The blue colour is due to the formation of prussian blue, resulting from the reaction between ~e~+and ferrocyanide ions. The reduction of 0, at the cathode gives OKions, which react with phenolphthelein to n;.~nn Anb ~.rrln.lr (2) A galvanic cell is formed when two metals of different electrode potentials are in Eledro&embl Cdls contact with an electrolyte. This is exactly what happens when steel containing alloying elements or steel-copper combinations are in contact with an electrolyte. In a galvanic cell of this type, iron acting as the anode gets corroded. The alloying element or Cu acts as the cathode. The corrosion of a less noble member of a pair of metals is called galvanic corrosion. Graphite in grey cast iron and cementite in carbon steel are cathodic compared to iron, and so, it is iron that corrodes. It is therefore clear that if iron is made to act as a cathode in a galvanic cell, the oxidation of iron leading to corrosion can be avoided. Iron pipes are coated with zinc (galvanized iron or GI) or aluminium for this purpose. Even in a corrosive atmosphere, it is zinc. acting as the anode of the galvanic cell. that corrodes. Thus zinc is called the I sacrificial anode and iron is said to be cathodically protected from getting corroded. / Steel structures in marine environments and pipes buried under ground are I cathodically protected by connecting it to zinc rods. r (3) The standard reduction potentials of the half reactions, in which oxygen is , reduced, are more positive than that of ~e~+/~e,and so, when there is an uneven supply of oxygen on the metal surface in contact with an electrolyte. corrosion of iron becomes possible. The area exposed to higher concentration of oxygen acts as the cathode and at the region where the oxygen concentration is low, the oxidation of Fe leading to its corrosion occurs. This is referred to as the differential aeration corrosion. The corrosion of an automobile and of GI pipes buried in sand are familiar examples of this type of corrosion. When some. of the protective layer of paint or zinc peels off, corrosion does not occur at the site of the broken film but at a site nearby where the concentration of oxygen is less.

13.12 PROTECTIVE MEASURES AGAINST CORROSION Since water and oxygen are essential for rusting of iron, the best way would be to protect the metal surface from these two. The most widely used method is to protect the surface of iron with a metallic or a non-metallic coating. The metallic coating should be capable of forming a self- protecting layer (e.g.) Zn. Ni, Sn. Al, Cd, Cr etc. These coatings may act as physical barriers against moisture and oxygen. In the case of iron coated with zinc (galvanised iron), even if there are a few scratches on the surface, rusting does not take place since compounds of zinc resulting from the oxidation of zinc plug these scratches. The same thing can not be said about metals like tin and nickel, having less negative Eo values compared iron, used for the protection of iron. In these case. these metals are able to protect iron I because of their ability to form oxide layers on their surfaces. Once the protective oxide layer is exposed, Sn or Ni acts as the cathode and iron as the anode in the galvanic cell and corrosion occurs. This is in a way advantageous since tin plated iron vessels can be thrown off after use. The cans get rusted and are converted to ferric oxide powder. The non-metallic coatings may be inorganic or organic. Metals like A1 , Ti and Zr have a thin film of protective oxide layers. The oxide layer can thickened electrolytically by anodising (metal used as anode in electrolysis). Iron is protected by forming a coating of phosphates of Fe .Zn and Mn on its surface (phosphatizing). These phosphate films serve as excellent bases for paints. Paints are dispersions of finely ground inorganic or organic pigments in film-forming materials like linseed oil (medium) dissolved in suitable solvents (thinners). The paint may protect the iron surface by acting as a barrier against the attack by oxygen and moisture. There may also be a chemical interaction between the constituents of paint with the surface to prevent the anodic oxidation of iron. Another way of protection is to add corrosion inhibitors t~ the medium with which the iron objects are in contact. The corrosion reaction may be stifled if the anodic oxidation and / or cathodic reaction processes are not allowed to take place. Anodic inhibitors like sodium chromate, sodium nitrite, etc. ililiibit the anode reaction. These are believed to act by setting right the defects in oxide films, responsible for the reaction, Fe - 2 e- +J~e~', to occur. In neutral or slightly alkaline media, the cathode reaction, viz. ,the reduction of oxygen to hydroxide ion occurs. Certain salts like zinc sulphate and magnesium sulphate act as cathodic inhibitors by forming insoluble hydroxides on the metal surface, and thereby, preventing the occurrence of the cathode reaction. Organic inhibitors containing N or S (e.g.) are believed to be absorbed all over the metal surface and retard or suppress the corrosion processes. Organic inhibitors like dicyclohexyl ammonium nitrite (DCHAN) are called vapour phase inhibitors (VPI) and are largely used to protect iron and steel machine from corrosion during transit by sea or air. SAQ 8 1. Explain the following : a) Galvanic corrosion is more facile in salt water than in fresh water. b) When Ni and Zn form a galvanic couple, zinc is likely to corrode. C) Electrode potential data cannot be used to predict the rate at which cell reactions take place.

2. In order to protect steel structures from corrosion, which of the following will be useful? Ni ,Na, Pb .Cd .Zn ,Mg ,Al.

13.13 ELECTROCHEMICAL ENERGY SOURCES - Commercial cells which are used as a source of electrical energy are of three types: Primary cells, Secondary cells and Fuel cells. Primary cells are based on cell reactions which are not reversible. Once the cell reaction is complete, the cell is discharged and cannot be charged again. Examples: Weston Cd Cell. Leclanche' cell (dry cell) etc. Secondary cells (storage cells or Accumulators) are galvanic cells in which the cell reactions that produce current can be reversed by applying an extemal source of current. These can be discharged and recharged many times until the electrode materials last. Examples: Lead-acid battery, Nickel-Cadmium battery, NiFe cells, etc. The term 'battery' is normally used to denote a number of galvanic cells connected in series. Fuel cells are also galvanic cells in which the reactants, to be oxidised at the anode (fuels) and reduced at the cathode (oxidants), are provided continuously from an external source and the products removed as they are formed. In a conventional cell, the reactants form a part of the cell. These are not replenished in primary cells but replenished in secondary cells. 13.13.1 Dry Cells These are modified Leclanchc's Cells in which an aqueous solution of the electrolyte is mixed with enough flour or starch to prevent spillage of the electrolyte. It can be represented as

- ' + Zn 1 NH,CI ( 25% ).ZnC12 ( 10% ) + MnO, ( s ) + C ( s ) 1 Graphite Figure 13.6 :Schematic Rep~sentaLionof a Dry Cd 1. Outer Cardboard Cover 2. Zinc Cup mode (Negative Eldmde) 3. Plastic or Pitch seal 4. Gra- rod (Positive Elearode) 5. Electmlyte (20 % NH4 CI + lOaD Zn (32 + Mna (s) + carbon .thickened with starch :

The dry cell is capable of delivering 1.5 V and the cell reactions can be represented as Anode : Zn - 2e- = zn2+ Cathode : 2Mn02+2H,0+2e-=2MnO(OH)+2OH Cell reaction: z~+~M~o~+~H~o=z~~++~M~o(OH)+~OH Mn02 is called the 'depolariser' since it prevents the formation of H2 at the cathode by preferentially getting reduced at the cathode. A secondary reaction (local action) results in the consumption of anode material and the electrolyte during the discharging of the cell. Once discharged the dry cell cannot be charged for reuse. Dry cells deteriorate on storage due to local action and also due to the evaporation of water from the electrolyte.

Miniature flat and round cells, capable of delivering 1.5 V, are used in calculators, hearing aids etc. In the alkaline manganese cell, zinc is the anode, a mixture of Mn$ I and graphite is the cathode and the elecuolyte is a solution of KOH contained in an absorbent material. The overall cell reaction is

The silver oxide-Zn cell, though costlier, lasts for a longer time. Here the n~deis Zn, cathode is Ag20 and the electrolyte is the KOH aq. The cell reaction is, Zn + Ago + H20 --> ZnO + H20 + Ag Secondary Cells The widely used automobile battery is a storage battery capable of delivering either 6 V or 12 V depending on the number of cells connected in series. They are used as stationary power sources in telephone exchanges, switching systems, emergency lighting etc. The cell can be represented as

Electrode reacrhns Anode : Equilibria & Electrochemistry To thpr TO other 0noJes

Figuie 13.7 :Schematic reprerentation of r lead accumulator 1. Glass a Plastic antainer 2. Gdl eleamdes made of Pb - Sb day 3. Spargy hdPadred in Mode 4. Pb& padred in cathode 5. Poms non-conductive Plastic Separator 6. Elecmlyce (38% by weight of Hs) sp. pvity = 1.30

Cathode :

Cell reaction for the passage of 2F discharging Pb02 + Pb + 2 H2S04 2 Pb SO4 + 2 H20 (1) charging The normal voltage of a single lead storage cell is about 2.OV. As the cell discharges electricity, PbS04 is deposited on both the electrodes and sulphuric acid is consumed, resulting in a decrease in the specific gravity of the electrolyte. With the aid of a hydrometer, the specific gravity can be checked and if it is equal to or below 1.20 (approx 28% by weight of &SO4), the battery is charged. The charging operation is performed in such a way that the negative pole of the battery is connected to the negative pole and the positive pole to the positive pole of the external charging device. The charging is done in an automobile by its electrical generator or . During the charging operation, water is converted into sulphuric acid (refer Eqn.l). The charging is done till the specific gravity increases to the required value. 13.13.2 Nickel-Cadmium Batteries The electrolyte is 21% by weight aqueous solution of KOH. The cell can be represented as .

anode cathode discharge Cd+2NiO(OH)+2H20 2 Cd(OH),+2Ni(OH), charging The voltage of single Ni-Cd cell is 1.3 to 1.4 V. Compared ta a lead storage battery, it has a longer life. It is also available as a sealed unit for use in electronic flash units (photography) and in calculators. The search is on for more efficient batteries for (i) storing electric power generated during the hours of low consumption for use during the periods of peak consumption (ii! storing electricity form solar and wind-powered generators, and (iii) use in developing quiet-running, non-polluting electric automobiles and mopeds. Though the lead-acid and the Ni-Cd batteries have been used to replace the internal combustion engines of automobiles, their performance is not as good as the gasoline-powered vehicles. Lithium or sodium-sulphur batteries have been developed for this purpose. Their drawbacks are (i) higher temperatures required for efficient operation, (ii) use of metals capable of violently reacting with water, and (iii) the necessity of using corrosion resistant materials. 13.13.3 Fuel Cells A can be represented as - + Fuel 1 Electrode 1 Electrolyte I Electrode I Oxidant Fuel can be eg. H2. N2H4 .Hydrocarbons. Oxidants can be eg. O2 ,H202 , HN03 etc. A fuel cell is a galvanic cell in which the chemical energy associated with the oxidation of reducing agents (fuel) is directly converted into electrical energy. The conventional method of utilising the chemical energy of the fuel to produce electrical energy and the direct conversion process can be represented as given below. (dl Chemical energy --+ Electrical energy 1 (c) I (b) Thermal energy -----., Any losses in energy in steps (a) and (c) can be minimised. However the efficiency of the process (b) is limited by the second law of thermodynamics. Hence the process (d) might be expected to have a higher efficiency. If AH is the enthalpy change of the reaction, the amount of useful work that can be obtained out of this is AG. The rest of it, i.e., AH - AG = T AS is unavailable for work. Work obtained from the cell reaction A G Efficiency = e = Heat change accompanyingthe reaction - A PI

Thus & depends on AS and AH The combustion (oxidation) reaction 2 H2 ( g ) + O2 ( g ) = 2 H,O, can be made tcr occur in a galvanic cell of the type (Fig. 13.8) - + Inert electrode I H, ( g ) 1 H+ I O2 ( g ) 1 Inert electrode Anode: 2%+4~,0=4~,0++4e(~O=0) Cathode: 02+4H,0++4e=6H,0(E0=1.23V) In the H2-02 fuel cell, the electrodes are Ti coated with porous R,and a water-soaked cation exchange resin in the acid form is used as the electrolyte. If the pressure of the gases are 1 aan, and water in the resin is pure, the EMF of this cell should be + 1.23 V, corresponding to the cathode reaction O2 + 4 H,O+ + 4e = 6 %O. This is because the potential of the hydrogen electrode under these conditions is zero. In practice the EMF is about 0.8 to 1.0 V. The only product discharged by the cell is water. The electrolyte in a % - O2 fuel cell may also be.alkaline. The electrode reactions in this case are Figurn 13.8 :Schematic reprearnuion of a H2- 9 fuel cell using H3O+ u electrolyta 1. Ti electrodcr canted with pomus Pt 2. Fuel (Hz) inlet 3. Oxidiser (02) inlet 4. Cation exchange membrane in acid form ~rourccofH3

Figure 13.9 :Schematic representation of a fuel cell ( H2 - 02) using an alkaline elemlyce 1. Pomus Ni anode 2. Pow Ni - NiO cathode 3. Hz inla 4. 02 mlet s. n-mlyte. KOH ( 4. )

In spite of higher efficiencies, fuel cells are rather expensive. However they are used in space-crafts because of their light weight and also because the product of oxidation, water, can be used by astronauts. SAQ 9 1. Distinguish between primary and secondary cells. 2. Write down the electrode reactions and the cell reactions when a lead-acid battery is discharged.

13.14 SUMMARY Many reduction-oxidation (redox) reactions involve electron transfers from reductants to oxidants. If the electron transfer is allowed to take place through an external circuit of a galvanic cell, rather than directly, electrical energy can be produced. The electromotive force (EMF) of the galvanic cell is a direct measure of free energy change of the cell reaction. One can therefore calculate the equilibrium constants bf redox reactions. The standard electrode potentials ( E' ) of the redox couples, derived on the basis of the assumption that the potential of a standard hydrogen electrode (SHE) is zero, are useful in predicting whether a redox reaction is I feasible or not, under standard conditions. The Nernst equation is used to calculate 1 the electrode potential or EMF under non-standard conditions. Galvanic cells in which the electrode potential of an electrode depends on [ H+] form the basis of pH measurements. Commercial pH meters make use of a glass electrode and a reference electrode like the saturated calomel electrode (SCE) to measure the pH. Potentiometrically, acid-base and redox titrations can be carried out more accurately than by using indicators. Corrosion, which is an electrochemical reaction, can take place when i) two metals of different electrode potentials are in contact with an electrolyte (Galvanic corrosion) or ii) the metal is exposed to different concentrations of oxygen. Iron can be protected from corroding using Zn or Mg as sacrificial anodes and also with the help of metallic or non-metallic coatings. Primary cells, secondary cells and fuel cells are the sources of electrical energy. In the case of secondary cells the cell reaction producing electrical energy (discharge process) can be reversed by applying an external source of current ( the charging process). In the case of primary cells, the discharge process cannot be reversed. In fuel cells, the chemical energy associated with the oxidation of fuels (reducing agents) is directly converted into electrical energy more efficiently than in other conventional processes.

13.15 GLOSSARY Gslvanic cells : A device in which the free energy of a redox reaction is converted into electrical energy. Anode : In a galvanic cell, it is the negative electrpe wherein oxidation is occuring and which is placed on the left in the cell representation. Equilibria & Electroehemlstrg Cathode In a galvanic cell, it is the positive electrode wherein reduction is occuring and which is placed'on the right in the cell representation.

electromotive force : the driving force arising due to the difference in the between the electrode and the solution.

standard electrode The potential resulting when the electrode is placed in a potential , solution of unit concentration (activity), with the electrode potential of the standard hydrogen electrode taken to be zero.

Nernst equation : E=EO-ElnB=EO-RTlnQ nF [Reactants] Potentiometric Titrations in which the EMF is measured as a function Titration of the volume of the titrant added.

Batteries and fuel Devices for converting chemical energy in a redox cells reaction or a combustion reaction into electrical energy.

Galvanic Corrosion : The electrochemical reaction occuring when two metals of different electrode potentials come in contact with an electrolyte. Sacrificial anodes : Electrodes (such as those of Zn or Mg) used for protecting a metal (iron) wherein they act as the anode in the place of iron, since they have a move negative E0 value than that of Fe.

13.16 ANSWERS TO SAQs SAQ 1 a) Zn(s) b). Fe (s) c) r

SAQ 2 a The underlined letkrs on the left hand side are alphabetically arranged (A. L, N, 0) b. Each alphabet on the LHS (with a single underline) comes earlier in the alphabetical order compared to the corresponding underlined alphabet on the RHS. Note that when the opposite electrochemical terms are arranged alphabetically, this mnemonic works. You may remember the word LOAN for left, oxidation anode, negative or you can not think of something better, e.g., the set RRCP for the right hand side electrode. c. In the case of an electrolytic cell, anode is positively charged electrode. The rest of the electrochemical terms are applicable to electrolyte cells also.

SAQ 3 1) AGO = -193KJ;K = 6.72x]d3 2) E,,, = 1.062 V SAQ 4 Cell reaction : zn+ZnZ++W cu2+ + 2e- -+ Cu Zn + cu2+-+ Cu + ZI?+ E.M.F. of the cell Eo (oxidatiori) - EO(oxidation) Left (Right), or E0 (Reductiod - E0 (Reduction) Right (Left) 1 = +0.35 - (-0.763) I = 1.113V i SAQ 5 Zn (s) -+ + X,Oxidation, E0 = + 0.76 volts 2Ag+ + 2e- +2Ag (s), Oxidation, E0 = - 0.80 volt 3ecell is Zn I ~n+~I Ag+ I Ag E.M.F. of the cell = E& - Ei, = 0.76 - (- 0.80) = 1.56 volts Since the value of emf is positive, the cell is feasible Rewrite the solutions of SAQ 5 and SAQ 6 using the reduction on potentials. Using reduction potentials has become the common practice and should be adopted. Of course, both conventions lead to the same result. SAQ 6 j . The oxidation potentials of the electrodes : EO Cu 1 CU+~= - 0.337 volts Eo Ag I Ag+ = - 0.799volts The cell reaction : 2Ag + 2Ag+ + 2e-, oxidation, EO = -0.799volt I + 2e- + Cu, reduction, E0 = - 0.337 volt EOcell = EOoxiaatim - Eooxidatim = - 0.799 - (- 0.337) = - 0.462 volts Since the value of e.m.f. is negative, the forward reaction is not spontaneous. The reverse reaction would be spontaneous, that is

will have +ve value of emf and hence it is possible to displace Ag from the solution of Ag+ ions with Cu cell in Cu I cu2+I Ag+ I Ag

SAQ 7 ' 1) a) No. The reduction Potential of cu2+I Cu is more positive. b) Any metal above Ag in the electrochemical series, eg., Cu, Zn, Fe c) M is a reductant. 2) E0 cell = 0.645 V Equillhrla LQ Electrochemistry SAQ 8 1) a) Ionic conduction is facile. b) Reduction potential of zn2+ I Zn is more negative. c) Predictions are them~odynamicallybased. 2) Zn, Mg, Al.

sac) 9 1) Ref :Text 2) Ref :Text