Behaviorofelectrolyticcells

Behaviorofelectrolyticcells - BEHAVIOR OF ELECTROLYTIC...

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Unformatted text preview: BEHAVIOR OF ELECTROLYTIC CELLS (Excerpted from “Chemical Separations and Measurements” by D. G. Peters, G. M. Hieftje, and J. M. Hayes, Chapter 12, pp. 393—402, W. B. Saunders Co., Philadelphia, 1974.) A number of methods of chemical analysis are based upon the principle that one may cause a desired reaction to take place when an external voltage of appropriate magnitude and polarity is impressed across the electrodes ofan electrochemical cell. Some of these techniques are described in this chapter. Using electrogravimetry, one can accomplish a number ofusciill analyses by depositing a metal quantitatively upon a previously weighed electrode. and by reweighing the electrode to ohtaiu the amount of the metal. In controlled-potential eoulometry, the potential of an anode or cathode is maintained constant so that only a single reaction occurs at that electrode. By integrating the current which flows as a function of time, one can deter- mine the total quantity ol‘electricit).r produced by the desired reaction and can calcu- late the amount ol‘ the species of' interest according to Ii‘araday’s law. Another technique, coulometric titration, is a method in which a titrant, electrt'ichemicaIl)’ generated at constant Current, reacts with the substance to be determined. Since the magnitude of the constant current is analogous to the concentration ofa standard tilt'ant solution, and the time required to complete the titration is equivalent to the \‘ultin'h’: ol‘ titrant. solution, the ein‘rent-time product is directly related to the un- known quantity ol'suhstancc. in the present chapter, Wt.‘ shall investigate what happens when current passes through an electrmrlieniieal cell. With this knowledg- as background, clectrogravi- metric analysis, eontrolled—potential CI':tIlUII'tt’.lI'}', and coulorrietric titrimeti'y will I): discussed. BEHAVIOR OF ELECTROLYTIC CELLS At least three phenomena occur when current passes through an electroclarmical cell. first, changes in the concentrations ol‘ chmnical species other at one or both electrodes. Second‘ there is an ohmic potential drop which develops because current flows against the resistance ol'the cell. Third, elTects related to the kinetics ol‘clectron- transfer processes influence the lichavior of a. cell. In addition, if the cell contains dif‘lerent solutions which share a connnon interface, liquid—junction potentials exist, Current-Voltage Curve for a Reversible Cell Let us examine the behavior of the zinc-copper cell represented hy the short- hand notation Zn XIIiNTJulg {l I") {11.1501 {I F} l (111 If, for simplicity. the activity coellicients oi‘zinc and cupric ions are taken to be unity, we can calculate an elnl'ol' I.ltitl v for this cell under the condition of zero current- llow, the copper electrorle'hcing pesitivc with respect to the zinc electrode. Ohmic potential drop. Let us connect an external voltage source to this cell—the positive terminal oi‘the source to the copper electrode, the negative terminal to the zinc electrode—so that the external source opposes the zinc-copper cell. In addition, let to assume provisionally (flat in each carrifxartmcnt (if—(la? cell perfect lanaogeiteig' of (fit relation is meirttat'rtrrt’, tit-w: iii. the vicinity q‘zhe electrode, means of highly flficiertl stirring. If the magnitude of the external applied voltage is precisely equal to the emf of the zinc-copper cell—that is, 1.100 \'—no current passes through the cell and no net reaction occurs. “diet: the applied voltage is increased above LIOU v, the zinc- copper cell behaves as an electrolytic cell, the reaction Zn2+ + Cu h> Zn -I- [111“ is forced to proceed, and the relationship lJt_'.i.\\.'('_'.f_‘.i'l the applied voltage and the current which flows through the cell is depicted hv line .-l in the upper hall'ol' Figure 12—]. “the applied voltage is less than 1.100 v, the zinc-copper cell acts as a galvanic cell, the overall reaction becomes Zn + (1113+ —> 2112" + (111 and the current-voltage relation is shown by line .--'l in the letter hall'ol‘Figurc 12—l.* * Since ult‘t‘trochr‘uIical I'eaclions cause changes in the coneentratieus |,':1:3I.:|-.‘.'ililts_‘l of zinc and eupi'ic ions as Well as the and and resistant-{- ol' the cell‘ Figure l2--l is strictly applicable only durum,r the lirst hm: Iiioim-nts ol‘ current How. 8,, l___—————~_..—.______————.-.---.-.—.._.___——————— E E | I l E 5 i 30 l —l l l I o 0.2 0.4 0.5 0.8 1.0 1.2 1.4 / External applied voltage. volts Extra voltage needed m obtain current i' because . concentration gradients o . IR dr p most at electrode surfaces and because energy barriers for electron transfer must be overcome Figure l2—I. Current—applied Vt'iltagc curves for the reversible electrochemical cell Zn 121: (It‘i|13|3‘t._.;f_lil1‘(1I|[(tht.‘S('J_l [1 F} 1 Cu Clul't-‘C A: schen‘tatic variation 0" current. with applied Voltage. for ideal situation in which the concentrations {activities} ofzinclllj and copperlllt are uniform throughout the two half—cells and in which the clectron—trmtslcr processes are illlittitt’l}'l}-15l. Curve 3 : schematic variatith urcrn'rtrlil with applied voltage when concentration gradients exist at the electrode sttrlitees and when the t'ltrctrorIvtransfin' reactions occur at linite rates. At applied voltages less than 1.100 t'. the system behaves as a galvanic. cell; at applied voltages greater than [.100 v, the system behaves as an electrolytic Cell. By convention, when an electrochemical cell t'ipcrates ntins])ontaneously, the result- ing current is said to be positive, whereas the current is negative if an electrochemical cell functions sptmtaneonsly as a galvanic cell. For the remainder of' this discussion, we shall focus our attention mainly on the behavior olthe zinc-copper system as an electrolytic cell. ll'the zinc nitrate and cupric sttllitte solutions remained homogeneous as electrolysis began, the current passing through the cell would exhibit a linear dependence on the quantity (Emu, — Em“), where EM,“ is the magnitude ol‘ the external voltage impressed across the two elec— trodes and EL..." is the zero—current ernl‘ol‘ the cell calculated from the concentrations {activities} ol‘zinc and cupric ions. In other words, the electrochemieal cell would act :15 a simple resistance, and the "'lll'l'ltlll, i, would he governed by ('Jhm’s law, Lilli]: — ‘Et'ell R where R is the total cell resistance. Altt'n'nativcly: this expressit'nt may in: written as E :tpn '= EH!” {R I". in the present lt}'pr.itl1t:tical situation. we wished to begin operating tltt‘: Zinc- Cfipper cell clectrolytically at a rate corresponding to the passage of current i in 4. Figure 12—1, 15”“, must exceed EH.” by an amount equal to iii’ in order to overcome the ohmic potential drop or the sci-called iR drop ol'tllc cell. Since the if? drop shown in Figure 12—] is 0. ml) \-', the required value of EM”, would he 1.200 r. Actually, the proper relationship between current and applied Voltage for a reversible electrochemical cell is depicted by curve B in Figure. 12—1. Note that the current does not vary linearly with external applied voltage. Looking at curve B in the upper half of Figure 12—], which illustrates the hehavior of the zine-copper system as an electrolytic cell, we diseovtEr that I. passage of current i requires an applied voltage of nearly [.300 v, not 1.200 v. Why must the external applied voltage be larger than the value needed to overcome the. if? drop? To answer this question, we must con— sider two phenomena—first, the existence ol‘conccnlration gradients at the anode and cathode during electrolysis and, second, the energy barrier that must he overcome in the electron-transfer process at the surface of an electrode. Concentration gradients at electrode surfaces during electrolysis. As soon as the zinc-copper cell starts to function electrolytically, the reaction Zn“ + (In—v Zn —t— (1113+ occurs, and the concentrations of zinc ion and cupric ion at the clematis: unfores- undergo imlncdiate changes—the concentralitm of" zinc ion decreases below l A! at the surface of the zinc cathode and the concentration ol'cupric ion around the copper anode becomes larger than I .-lI—-:tllh011gh the bulk concentration ol‘ each ion re- mains at l M'. However, even electrolysis causes :1 Concentration gradient to develop at the surface ol‘ each electrode, Ll‘tCI‘f are three mass~transport processes — diffusion, electrical migration, and mechanical stirring—which act sin‘tullaneously to diminish thcse concentration gradients. Firsl, hecaus * a chemical species undergoes diffusion from a region of higher concentration to one of lower ctmcentration, zinc ions will dili‘usc from the bulk ol‘ the zinc nitrate solution toward the zinc cathode, whereas cupric ions will dil'l'use from the copper anode into the bulk of the cupric sulfate solution. Second, the direction of current flow through the cell is such that there is electrical migration of zinc ions toward the zinc cathode and of cupl‘ic ions away front the copper anode. Third, the use of mechanical stirring helps decrease the differences between the concentrations of zinc ion and cupric ion at the electrode surfaced“- and in the hqu ol the solutions. How ellcctive arc dill‘usicin, electrical migration, and mechanical stirring in eliminating the concentration gradients caused by electrolysis? M'echanical stirring is most efficient for the kind of electrolysis experiment we are describing now. How- ever, even if the solution in contact with each electrode is very vigoroust stirred, the theory of hydrodynamics tells us that a thin layer of stationary liquid always sur- rounds an electrode immersed in a moving solution. Consequently, n'iovemem of zinc ions to the surface of the cathode proceeds in two steps; zinc ions are transported by stirring of the zinc nitrate solution up to the edge of the thin film of immobile liquid, and then di use and migrate: through the film to the electrode surface at which the}r are reduced. At the copper anode, the eupric ions 0? use and migrate away from the electrode through the thin film ol'stationar‘y liquid and, upon reaching the stirred solution, are swept into the main body of the copper sulfate medium."c In spite of difi’wion, migration, and stirring, {lie Zora-copper cell operate: at some usefiilt‘r large current level, it is inevitable that the concentration gradient produced at each electrode in! electralrrir will persist to some extent. “ Nitrate and sulfate ions also undergo mass transport by means of diffusion. cleclrical migrationI and mechanical stirrin . it would be instructive Ibr the reader to chart the movemdnt of tlmsg anions ‘ 3 during the electrolysis. 5 Just as the voltage impressed across the anode and cathode of an electrolytic cell niusL be increased to compensate for the ohmic potential drop, so the existence ofa concentration gradient at the surface ofeach electrode creates a similar need— the applied voltage must be increased still more to cause the desired rate of electrol- ysis. ‘Nith reference to Figure 12—1, suppose that we wish to obtain a current flow of ithrough the zincleopper cell. Let us assume arbitrarily that, in the attaimnent ofthis rate ofelectrolysis, the surface concentratit'in of zinc ion is lowered to (1.1 ill and the surface concentratit'nt ofcupric ion becomes 1.!) ill. By using the Nernst equation, we can calculate that the emf ofan electrochemical cell consisting:r of a zinc electrode in contact with (10.1 ill zinc ion solution and ofa copper electrode immersed in a 1.9 .-l.' cupric ion solution is 1.138 v, which is 38 mv larger than that of a similar cell with l 11-! solutions of the two metal ions in contact with the electrodes. ’l'hercfore, in this hypothetical situation, if we wish current i to flow initially through the zinc—copper cell, the applied voltage must not only be large enough to overcome the ohmic po- tential drop but rnus.‘ further exceed the emf of the cell at zero current llow by at least 38 mv.* Energy barriers for electron -rrarisfer reactions. Like other chemical processes, the transfer of electrons from a chemical species to an electrode {or vice versa) involves a reaction pathway that includes an energy barrier with an associated free energy of activation {530"}. In contrast to an ordinary chemical reaction, how- ever, the height of the energy barrier for such an electron-transfer process depends upon the potential of the electrode. We can use this simple picture to examine the behavior of the zinc electrode of the zinc-copper electrolytic cell. l'l’licn in; ma! otrrartfflows through the zoo-rapper rcl‘l, the reversible electron-transfer reaction 7mg" —|— 2 e {-—-‘- Zn exists in a state of dynamic equilibrium. This means that the ratc of the forward reaction (reduction ofzinc ion to zinc metal} equals the rate of the backward reaction (oxidation of zinc metal to zinc ion). In other words, equal cathodic and anoclic currents'l' are [lowing across the interft'ice between the. zinc electrode and the zinc nitrate solution. ll" the activities ofzincfll} and zinc metal are taken to be unit): at the electrode—solution interface, the equality of the cathodic and anodic currents requires that the energy barriers for the forward and backward electron-transfer reactions be identical. “filter: a net current dear flow titratith the zinc-rapper cell, the zinc electrode functions as the cathode and net reduction of zinctll) to zinc metal occurs. Consider the following hypothetical question: if the activities ofzincfll) and zinc metal at the electrode-solution interface are kept at unity. what must be done to cause net re— dttction of zincfll) to zinc metal? To obtain the desired result, we need to make. the rate of reduction of zincUl) to zinc metal greater than the rate of oxidation of zinc * As stated in Chapter 9 [page 2til)),the ernl'ofthc zinc—copper cell is governed by the course-titra— tions (activities) of species at the .1?thth oflr'm electrodes. However, when no current flows through the cell, it is permissible to substitute be“; crxrtcrntmt'ioas inlo the Nernst equation for calculation of the cell L‘l'flf, because there are no concentration gradients in the solution. \Vhen current does llow through the cell, the concentrations of species at the electronic surfaces deviate from the bulk concentrations, ' 50 1hr“ cell emfwill dillizr from the value based on the bulk concentral ions. In many books. this difference if» Inferred to as concentration overpolential, a term which implies sonic departure of the Cell from normal behavior. If the true concentrations of species at the electrode surfaces {luring current flow are substituted into the Ncrnsl equation, the calculated cell eml'sltould agree with the observed emf {if there are no other complications) and concentration overpotenlial becomes nonexistent. 'l‘herel'ore, WC prefer to avoid use of the artiliciai tel-in concentration overpotential and to focus attention instead an the real phenomena that occur in a cell during electrolysis. T This current, either anotlic or cathodic, is known as the exchange current. metal to zinClIlj. This can he accrnnplislnfl iii the. energy harrier for the Ii_'l'r'sr\vartl‘j reaction Zn“ —l— 2' t: —» Zn becomes smallcrtlnui the energy barrier For thc lhackward} reaction Zn ---> Zn“ -l,- ‘2 c It can be. shown. although we will not ollcr the I'H‘l')l.ll‘l‘.|l.'.l't‘_.* that ;t shill in the ,)o— tential oi the zinc clcctrot'le in the nrguiiru direction lowers the energy harrier for the iiirward reaction and, siinultanct'iusly, raises the energy barrier for the .jntcltward reaction. Therelin'c, a negative shill in the potential 01‘ the Ziltt; cathode from its Zi:t‘0-(.‘.III‘I‘CHI, equilibrium value [ta-m the reduction of ziucfll'} to zinc metal and causes a net cathodic current to How. This shift in the potential of" an electrode that produces a certain net. current (in the. absence of any concentration gradients} is called the activation overpotential. litre want the net current For the reduction of zinctll') to zinc. metal to increase, the potential ol‘ the cathode must he shifted to a greater extent in the negative clirt:t:tion——a bigger current requires a larger activation U\.'[‘.I'I)Ul(_‘.l'll,l£ll, Similar reasoning leads to the conclusion that {even in the absence ol'a concentra- tion gradient) the potential ol‘ the copper electrode of the zinc-copper cell must he Sllil'itftl in :1 {Mutt-it's direcl'it'tn from its Value when no current is llowing to cause net oxidation of copper metal to ctipric ion. Thus, there is an activation tn'erpotential asst'rciated with the electrtitn-ti'anslirr process at each electrode of‘an electrolytic cell. For some reactions, such as the reduction ol‘a tnetal cation at a mercury surface to fi'll‘tl] an amalgam, the activation ovel‘potential needed for rapid electron transfer is small. For other processes, such as those involving the formation 01' rupture of chemical bonds, the activation ovcrpotcntial required for fast electron transier is considerably larger, Although it is not I'iossiblc to predict the magnitude of the activation overpotential needed to cause an electrtnivtransier reaction to proceed at a specified rate, there are some quz'ilitative guides. First, activation m'erpotential in- creases, sometimes quite drastically, as the current density {current per unit electrode area) l)t‘.COl‘tltt$ larger. Second, an t‘.l(i\-'£tl.l011 Ul‘ telttl'n‘rature causes the activation ovcrpotential to decrease, because part of the free energy oi“ activation for the elec- tron—tl‘tlnsli‘n‘ process is provided thermally. Third, activation overpotcntials are cliaraeteristicalltr large for reactions resulting in liberation ol‘ gases, for the oxidation or reduction of organic molecules, and for processes involving multiple electron transfers. Fourth, the deposition of one metal upon the surface ol‘a diilerent metal nl‘tcn takes place with some activation UVCI‘I){'Jltfl‘ltiill until a e'nnpletc layer of the desired metal has been plated out. l’il‘th, electrode processes occurring at the so- called soft metals, includng lead, zinc, and mercury. usually show much larger activation ovcrpotentiuls than at Noble metals such as platinum. pnlladitnn, and gold. Table 12- -l lists some values of activation over-potentials for the evolution ofhvdro— gen and oxygen gases at different electrodes and current densities. Notice that: the data nicely illustrate several ol. the qualitative rulesjust I]']('.'Jllit.)l'](i(.l. Let us recall the ’arious plient'm'iernt that occur during operation of an electrt'i- lvtic cell. It should now he evident that, to make a definite current flow through a cell, the. external applied voltage must he greater than thr- zero—current enil‘ of the cell by the amounts needed to overcome the ohmic potential drop, to t:rnn}'urns:ttc l'or 3‘ A Illut'otlglt presentation ol' the subject of (:lt‘fl‘lI'DCIII‘II‘Iil'JIl kinetics is given It}: ll. [3. Conway: Tin-err our! Principles of Elan-trod:- I’rm'rste'ri Ronald Press, New York, 196:3, 111:. Elf—I33. 7 Table 12-1. Activation Overpotentinls for Hydrogen and Oxygen Evolution at Various Electrodes and Current Densities at 25°C* Oval-potential, Volts Electrode Material 0:001 sump-fern2 0.01 Straggle-4:112 0.1 arnplern:1 1.0 Simplex-n“ H2 02 H2 02 H2 02 H3 (:12 Pt, smooth 0.024 0.72 0.000 0.85 0.20 1.3 0.03 1.5 1’t,platinizcd 0.015 0.35 0.030 0.52 0.041 0.54 0.048 0.75 At] 0.24 0.57 0.30 0.95 —- — 0.80 1.5 Ctb 0.48 0.4-2 0.53 ‘0.58 0.00 0.00 1.3 0.79 Ni 0.50 0.35 0.75 0.52 1.1 0.54 1.2 0.35 Hgtt 0.9 — 1.0. v— 1.1 — 1.2 H Zn 0.72 — 0.75 —- 1.1 — 1.2 —- Ft) 0.40 —- 0.55 — 0.82 -—- 1.3 -— Pb 0.52 —- 1.0 H—‘ 1.2 -——— 1.3 — * Data from International Critical Tables 6339,1929. the existence of‘coneentrrttimt gradients, and to supply the activation ot-‘erpotentia'tls for rapid electron transfer at each electrode. We. ham: not discussed the detuiled behavior of the zinc-copper system :15 a galvanic cell. It is sullleient to say here that, since the cell behaves revet‘sil:113.-'—thc overall cell reaction. can he made to go in either direction—the upper and lower l.‘11‘:‘1nches til—curve 13’ in Figure 12—--| are symmetrical rthout the point at which the external applied voltage is 1.100 t-'. “then the cell acts gztlvztnically, the phenomena resptmsilfle for the ohmic potential drop. the concentration gradients, and the activat- lion overpotentiztls are still operative. Current-Voltage Curve for an Irreversible Cell Some. electrocltett'ticztl cells behave irreversihly. As an example, consider the cell Pt Ztttrttm. {IF} House, {11:} | cu in which the zinc electrode 01‘ the. previous ziitc-comjer cell has been replaced by :1 platinum electrode. Assignment of :m emf to the present cell is impossible; since the platinum electrode initially has no deposit of zinc 111etal on its surface. the potential of this electrode. is undefined. li'urthermnre, the ahsence of zinc metal precludes the SP‘JnlimF-UIIS discharge. I'll this electrochemical t’t'fll. What happens when at variable Voltage is in'tpressed across the two electrodes :11 this 'eell, the negative side of the source being motioned to the platinum electrode E'tlltl the positive side of the source to the copper electrode? Figure 12—2 shows the quulitatit'e dependence ol‘ the current Upon the value 01‘ the applied voltage. As long as the. applied voltage remains below 1.100 v. the current is relatively small, altht'tugh it does inert-ruse more or less linearly 21s the. applied \-'t_'11t'.1ge becomes larger. T-Iowe'rer, when the applied voltage reaches 1.111019, the current rises :thruptly. 11nd the current—voltztgc curve mimics curve 15’ in Figure 12—1. Residual current. What is the origin of" the. small current --the residual current—that flows through the cell at applied \‘rtltztges less than 1.100 v 1’ At the titrpper :motle, ()deélllt'nl of :1 tiny amount of metallic copper to euprir. ion occurs. 0.2 0.4 0.6 0.8 1.0 1.2 1.4 External applied voltage. volts Current v Figure {2—2. Current-applied voltage curve for the irreversible electrochemical cell Pt] Zn{N03)3 (l |CuSCII4 {l FllCu At an}r applied voltage, this system can function only as an electrolytic cell. Extrapolation of the rising portion of the curve down to the zero-current line gives the so-callml decomposition potential. On the other hand, at least two current-producing reactions take place at the plati- num cathode. First, traces ol‘t‘lissolvcd oxygen are reduced to hydrogen peroxide and water; other impurities mayr undergo reduction as well. Second, there can he dep- osition of a small quantity of zinc metal upon the platinum electrode. 01" course, the total anodic residual current is equal to the total cathodic residual current. It might be expected that no metallic zinc would be deposited until the external applied voltage exceeds 1.100 v, as is true for the reversible zine-copper cell de— scribed earlier [Figure 12—1). However, this statement is correct only il‘zinc metal is deposited in a state of unit activity (or ii‘a pure zinc electrode is initially in contact with the zinc nitrate solution). “then only a minute quantity of zincmetal is plated upon a platinum surface, the activity of zinc metal will be Far less than unity—a value such as 10—10 Mr is not unreasonable. Through use of the Nernst equation, it can be shown For the present cell that deposition ot‘zinc 'metal at an activity of 10"” £11! upon the platinum electrode occurs at an external applied voltage of approxi— mately 0.800 v. Furthermore, deposition of zinc metal at even smaller activities may proceed at lower applied voltages. Thurs, part ofthe cathodic residual current can be attributed to deposition of zinc metal at less than unit activity upon the platinum surface. Decomposition potential. Commencement of rapid deporsition of metallic zinc upon the platinum electrode accounts for the sudden rise in current at approxi- mately 1.100 V, because sufficient zinc is plated on the platinum cathode to transform C? the latter into a zinc electrode. Some electrochemists designate the applied voltage at which zinc ion begins to he reduced as the decomposition potential. However, the conCcpt ofa decomposition potential lacks theoretical significance For the irreversfl iblc cell Pt | 2:19:03): {I 1-"; House),1 (I 1-‘)I(.‘u because, as we have seen1 some zinc metal is deposited at all values of the applied voltage. Thus, Wt: may only' regard the decomposition potential as the approximate applied voltage at which the cell current exhibits a rapid increase and above which reduction ofzinc ion occurs at a convenient rate for practical applications. On the other hand, the decomposition potential for the reversible cell an | almosta (J F)|]C111804 (t smelt is identical to the. emi‘ot' the cell at zero current flow; deposition ofzinc metal begins as soon as the applied voltage exceeds 1.100 v. Relationship between Applied Voltage, Individual Electrode PotentiaIS, and Ohmic Potential Drop Ifwe consider a practical electrolysis for the reversible zinc-copper cell described earlier, Zn Zntwoaja (1 F) H (11180,, (I F} | cu the total voltage {Haw} impressed across the anode and cathode is the sum of two terms Eapp : Eur-II + where Emu is the cell emf based on the surfi-tce concentrations of zinc ion and cupric ion, plus the sum of the activation overpotentials for the reduction of' zincfll} and the oxidation of copper metal, and where it? represents the ohmic potential drop. In turn, the cm? of the cell consists of two contributions Ecell : ILCuzflL‘u + £211.th2+ if liquid-junction potentials are ignored. Another form of the last equation is EN” = ECtii'l'fiu _ isZn." ,Zn in which ECughC“ denotes the actual potential of the copper electrode and Ewfln the actual potential of the zinc electrode, each measured with respect to the same reference electrode. For each electrode the actual potential is governed by the surface Concentration of the pertinent metal cation and by the activation overpotential for the electron-transfer reaction occurring at that electrode. Combining the preceding ExpreSsions; we obtain F ram: :13 — F ’Znit‘Zn + SR Cttz'l .L.‘u Noting that the copper electrode is the anode of the zinc-copper electrolytic cell and that the zinc electrode is the cathode: we can write a general equation for the applied IO voltage in any electrolysis experiment Iafllrll : j"‘u __ be + ER where EEl and Eu are actual anode and cathode potentials, respectiver. During an electrolysis, the various terms in the preceding equation [nay Cllitllgu with time. For example, when Ea“, is heltl constant, the cell resistance (R) probably changes as electrolysis proceeds, whereas the current {i} decreases as the reacting species are consumed. IF the latter tWo quantities change unequally, the ohmic potential drop (it?) will vary. and so must the potentials of the anode and cathodg Similarly, if one performs 2.1 constant-current electrolysis, it is necessary to readjust the applied voltage continuously as the resistance varies; but any change in the applied voltage alters the anode and cathode potentials. ...
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Behaviorofelectrolyticcells - BEHAVIOR OF ELECTROLYTIC...

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