Laidler Physical Chemistry Chpt. 8

PREVIEW

[his chapter is concerned with galvanic cells, in which a chemical reaction produces an electric potential difference between two electrodes. An example is the I)aniell cell, in which zinc and copper electrodes are immersed in solutions ol' Zn2 and Cu2 ions separated by a porous membrane. which prevents hulk mixing, l'he chemical reactions occurring at the two electrodes cause a how of electrons in the outer circuit.

It is impossible to measure the electromotive force tenif) of a single electrode, and the convention is to use a standard hydrogen electrode as the left-hand electrode in a cell. With another standard electrode on the right-hand side, the cell emf is then taken to he the standard electrode potential for the right-hand electrode. Such standard electrode potentials are useful for predicting the direction of a chemical reaction. The thermodynamic relationship between the Gibbs energy change and the 'mnf of a galvanic cell leads to the determination of sjuilibrium constants. An equation due to Nernst relates lie cell emf to the activities of solutions in the cell. The Nernst potential is the potential difference established across a membrane when two diikrent solutions are on itpposite sides of it.

An important type of cell is a redox cell, in which oxidation-reduction processes occur in a special way. It is often convenient to relate the ernt's of such cells to a standard pH value, usually taken to he 7.

Practical applications of cml measurements include the determination of pH, activity coefficients, equilibriuni constants, soluhility products, and transport numbers. The technique of polarography has become an important analytical tool. Much valuable information has also been derived from studies of the kinetics of electrode processes, particularly with reference to the phenomenon of overvoltage.

Certain electrochemical cells are useful devices for the CeneI'ation of electric power. In a fuel cell, for example, the reacting substances are continuously fed into the system. Photogalvanic cells are electrochemical cells in which radiation induces a chemical process that gives rise to an electric current. Among batteries discussed, lithium ion batteries give a high energy per unit density. Consequently, they will he used in hybrid automobiles and other applications requiring high energy drain and light weight.

315

OBJECTIVES

After studying this chapter, the student should be able to: Obtain activity coefficients and equilibrium constants

Explain how electrochemical cells, such as the from emf measurements.

Daniell cell, function. Describe fuel cells and photogalvanic cells.

Understand the concept of the standard hydrogen Understand the use of the terms battery and cells, electrode and describe other standard electrodes. U Describe the different batteries discussed, especially

Apply the principles of thermodynamics to the chemistry of the lithium ion battery. electrochemical cells, including the derivation of the Nernst equation.

316

An electrochemical cell, also called a voltaic cell or a galvanic cell, is a device in Anode and cathode: Review which a chemical reaction occurs with the production of an electric potential

REDOX reactions and ionic difference between two electrodes. If the electrodes are connected to an external bonds circuit there results a flow of current, which can lead to the performance of

mechanical work so that the electrochemical cell transforms chemical energy into work. Besides being of considerable practical importance, electrochemical cells are valuable laboratory instruments, because they provide some extremely useful scientific data. For example, they lead to thermodynamic quantities such as enthalpies and Gibbs energies, and they allow us to determine transport numbers and activity coefficients for ions in solution. This chapter deals with the general principles of electrochemical cells and with some of their more important applications.

8.1 .' The Daniell Cell The cell originally developed by the English chemist John Frederic Daniell (1790-1845) consisted of a zinc electrode immersed in dilute sulfuric acid and a cop-per electrode immersed in a cupric sulfate solution. It was later found that the cell gave a more stable voltage if the sulfuric acid was replaced by zinc sulfate solution, and the expression "Daniell cell" today usually refers to such an arrangement, which is shown schematically in Figure 8.1. The voltage produced by the cell depends on the activities of the Zn2 and Cu 2+ ions in the two solutions. If the molalities of the two solutions are 1 mol kg' (1 m), the cell is called a standard Daniell cell.

When a Daniell cell is set up, there is a flow of electrons from the zinc to the copper electrode in the outer circuit. This means that positive current is moving

Anode and cathode examples; from left to right in the cell itself. By convention, a potential difference corre- Electrolytes. sponding to an external flow of electrons from the left-hand electrode to the

right-hand electrode is said to be a positive potential difference. The processes that occur when this cell operates are shown in Figure 8.1. Since

positive electricity in the form of positive ions moves from left to right within the cell, zinc metal dissolves to form Zn2 ions,

Zn - Zn2 + 2e

Some of these zinc ions pass through the membrane into the right-hand solution, and at the right-hand electrode cupric ions interact with electrons to form metallic copper:

Cu2 + 2e - Cu

Every time a zinc atom dissolves and a copper atom is deposited, two electrons travel round the outer circuit.

Cells of this kind can be made to behave in a reversible fashion, by balancing their potentials by an external potential so that no current flows. This can be done by means of a potentiometer, the principle of which is illustrated in Figure 8.2a. A current is passed through a uniform slide wire AB, along which there is a linear po-tential drop, the magnitude of which can be adjusted by a rheostat. The cell under investigation is connected through a tap-key switch and a galvanometer to one end (A) of the slide wire and to a movable contact C. The contact is moved along the slide wire until, when the switch is depressed, no current passes through the gal-vanometer. The potential difference produced by the cell is then exactly balanced

317

318 Chapter 8 Electrochemical Cells

Voltmeter

REDOX reactions: Zinc reduces copper; Copper oxidizes zinc.

FIGURE 8.1 The standard Daniell cell. 1 m solution of ZnSO4 1 m solution of CuSO4

by the potential difference between points A and C. The potentiometer wire can be calibrated by use of a standard cell, such as the Weston cell (Figure 8.2b),

Anode and cathode: See cations When the electric potential of a cell is exactly balanced in this way, the cell is migrate to the cathode and operating reversibly and its potential is then referred to as the electromotive force anions migrate to the anode.

(emf) of the cell. If the counter-potential in the slide wire is slightly less than the emf of the cell, there is a small flow of electrons from left to right in the external circuit. If the counter-potential is adjusted to be slightly greater than the cell emf, the cell is forced to operate in reverse; zinc is deposited at the left-hand electrode,

Zn2 + 2e 9 Zn

and copper dissolves at the right,

Cu - Cu2 + 2e

FIGURE 8.2 4- Rheostat - (a) A potentiometer circuit used for

cell Saturated determining the reversible emf of Potentiometer c

_: Crystals of a cell. (b) The Weston standard

B

solution

used standard cell. At 25 C its A JuuIuuuuuuuiuuuiiii7

.Galvanometer

J CdSO4H20 cell, which is the most widely Slide wire of CdSO4

emf is 1.018 32 V and it has a Mixture of

very small temperature coefficient mercury and* (see Problem 8.33).

Tap-key

Hg2SO4

Cell under Mercury investigation Cadmium amalgam

a. b.

12 gas at 1 ar (100 kPa) ressure

I solution of H ions

FIGURE 8.3 The standard hydrogen electrode.

8.2 Standard Electrode Potentials 319

The fact that the electrons normally flow from the zinc to the copper electrode indi-cates that the tendency for Zn -> Zn 2+ + 2e to occur is greater than for the reaction Cu * Cu 2+ + 2e, which is forced to occur in the reverse direction. The magnitude of the emf developed is a measure of the relative tendencies of the two processes. The emf varies with the activities of the Zn2 and Cu 2+ ions in the two solutions. Thus the tendency for Zn - Zn 2+ + 2e to occur is smaller when the concentration of Zn 2+ is large, while the tendency for 2e + Cu 2+ - Cu to occur increases when the concentration of Cu2 is increased. The relationship between emf and the activities will be considered in Section 8.5.

8.2

Standard Electrode Potentials It would be very convenient if we could measure the potential of a single electrode, such as the right-hand electrode in Figure 8. 1, which we will write as

Cu2 + Icu The potential of such an electrode would be a measure of the tendency of the process

Cu2 + 2e -* Cu

to occur. However, there is no way to measure the emf of a single electrode, since in order to obtain an emf there must be two electrodes, with an emf associated with each one. The procedure used is to choose one electrode as a standard and to mea-sure emf values of other electrodes with reference to that standard.

The Standard Hydrogen Electrode The electrode chosen as the ultimate standard is the standard hydrogen electrode, which is illustrated in Figure 8.3. It consists of a platinum electrode immersed in a I m solution of hydrogen ions maintained at 25 C and I bar pressure. Hydrogen gas is bubbled over the electrode and passes into solution, forming hydrogen ions and electrons:

H2 -2H + 2e

Standard Electrode Potentials

The .emf corresponding to this electrode is arbitrarily assigned the value of zero, and this electrode is used as a standard for other electrodes.

There are two conventions in common use, and the student should be aware of both methods of procedure. The standard hydrogen electrode can be either the left-hand or the right-hand electrode. In the convention adopted by the International Union of Pure and Applied Chemistry (IUPAC) the hydrogen electrode is placed on the left-hand side, and the emf of the other electrode is taken to be that of the cell. Such emf values, under standard conditions, are known as standard electrode po-tentials or standard reduction potentials and are given the symbol E. Alternatively, the standard hydrogen electrode may be placed on the right-hand side; the potential so obtained is known as the standard oxidation potential. The latter potentials are the standard electrode potentials with the signs reversed; the only difference is that the cells have been turned around.

To illustrate the standard electrode (reduction) potentials ([UPAC convention), consider the cell shown in Figure 8.4. The left-hand electrode is the standard


FIGURE 8.4 A voltaic cell in which a standard hydrogen electrode has been combined with a copper electrode immersed in a 1 m solution of cupric ions; the two solutions are connected by a potassium chlo-ride bridge.

H2 gas at 1 bai pressui

hydrogen electrode, with a hydrogen gas pressure of 1 bar, and the acid solution is 1 m in H ions. The right-hand electrode is the Cu2ICu electrode, and the concen-tration of Cu2 ions is 1 m. The two solutions are connected by a "salt bridge" such as a potassium chloride solution, which conducts electricity but does not allow bulk mixing of the two solutions. Alternatively, an agar gel containing KC1 is commonly used. This procedure is somewhat more reliable than separating the two solutions by a porous partition, which itself sets up a small emf; the salt bridge minimizes this effect.

The voltaic cell shown in Figure 8.4 can be represented as follows:'

Pt, H2(1 bar)IH(1 m) Cu21(1 m)ICu where the double vertical dashed lines represent the salt bridge. The observed emf is +034 V; the sign is positive by convention since electrons flow from left to right in the outer circuit. There is therefore a greater tendency for the process

Cu 2+ + 2e- Cu

to occur than for

2W + 2e_ H2

to occur; the latter process is forced to go in the reverse direction. By the IUPAC convention, the Cu2ICu electrode is on the right, and the standard electrode poten-

Wheu a gas is a reactant in an electrochemical cell, its pressure must be known. Throughout this dis-cussion, the gas pressure, if not stated, is assumed to be I bar. Values of the emf in many tables relate to 1 arm pressure. The differences in values are generally within the margins of error of the measurement.

A single vertical line represents a phase separation. Double vertical dashed lines represent a salt bridge.

REDOX reactions: Galvanic cell; Watch the electrons flow through the external circuit and anions and cations pass through the salt bridge.


REDOX agents: Distinguish tial E of this electrode is 0.34 V. This is a measure of the tendency for the cupric between what is reduced and ions to be reduced by the process what is oxidized. 2+ Il Cu +2e-

-'Cu

It is for this reason that these electrode potentials are also known as standard reduc-tion potentials.

In this book, standard electrode (reduction) potentials will always be em-ployed, as recommended by the IUPAC. Table 8.1 lists such potentials; the reactions are written as reduction processes (e.g., Cu2 + 2e - Cu). A table of standard oxidation potentials, on the other hand, would have all the signs reversed and the corresponding reactions would be oxidations; for example,

Cu -* Cu2 + 2e

REDOX reactions: Electrolytic By combining the standard electrode potentials for two electrodes we can de- cells; Compare galvanic and duce the emf of a cell involving the two electrodes, neglecting the hydrogen electrolytic cells, electrode. Consider, for example, the following items in Table 8.1:

Cu2 +2e-+Cu E0.34V Zn2 +2e-Zn E-0.76V

These values are the emf values for the following cells, in which the electrode processes are shown:

Pt, H2 (1 bar) (H4(1 m) Cu2(l m)ICu E = 0.34 V H2 -2H + 2e Cu2 + 2e -*Cu

Standard reduction potentials: Pt, H2 (1 bar) H(1 m) Zn2(l m)IZn E = -0.76 V Example; Learn the strongest to H2 -4 2W + 2e Zn2 + 2e -* Zn weakest oxidizing and reducing agents. We could connect the two cells together as follows:

ZnIZn2(1 m) H(l m)(H2 (1 bar), PtPt, H2 (1 bar) jH(1 m) Cu2 (1 m)ICu and the emf would then be E = 0.34 - (-0.76) = 1.10 V (i.e., the right-hand electrode potential minus the left-hand potential). We could also eliminate the hydrogen electrodes altogether and set up the cell

ZnlZn2 (1 m)HCu2 (1 ,n)Cu The emf of this would be the same, 1.10 V, since we have merely eliminated two identical hydrogen electrodes working in opposition to each other. This cell is, of course, just the standard Daniel! cell (see Figure 8.1).

The potential for a half-reaction Note that in writing the individual cell reactions it makes no difference is an intensive property. whether they are written with one or with more electrons. Thus the hydrogen elec-

trode reaction can be written as either

2W +2e-*H2 or H +e-+--H2

Standard reduction potentials: However, in considering the overall process we must obviously balance out the Reactivity and stoichiometry; electrons. Thus, for the cell Take quiz on standard reduction + 2+ Pt, H2 H :: Cu Cu potentials.

the individual reactions can be written as

H2 -42H+ +2e- and 2e+Cu2-Cu


TABLE 8.1 Standard Electrode (Reduction) Potentials, T = 25 C Half-Reaction - Standard Electrode Potential, E/V

F2 + 2e -* 2F 2.866 H202 + 214+ + 2e - 2H2O 1.776 Ce" + e -, Ce3 1.72 Au4 + e -4 Au 1.692

Oxidation states; Summarize the Mn0 + 8W + 5e -* Mn 2+ + 4H20 1.52 oxidation states of elements by Ce" + e- Ce3 1.72 use of the periodic table. C12 + i - 2C1 1.35827

Cr2O + 14W + 6e - 2Cr + 7H20 1.232 Mn02 + 4W + 2e - Mn 2+ + 21420 1.224 02 + 4W + 4e -* 21420 1.229 Pt24 + 2e -* Pt 1.18 Br2 + 2e -+ 2Br 1.0873 Hg2 + 2e -* Hg 0.851 Ag + e -* Ag 0.7996 Hg+ 2e -4 2Hg 0.7973 Fe 3+ + e - Fe2 0.771 02 + 2W + 2e -* H202 0.695 12 + 2e - 21- 0.5355 Cu 2+ + 2e- Cu 0.3419 Hg2Cl2 + 2e -+ 2Hg + 20F, 0.1 m HC] 0,3337 H92Cl2 + 2e -+ 2Hg + 2C1, saturated KCI 0.2412 AgCI(s) + e- Ag + Cl- 0.22233

Oxidation states; Examples; Test Cu24 + e- Cu' 0.153 your ability to assign oxidation Sn" + 2e- Sn24 0.151 states. HgO + H20 + 2e -3 Hg + 20W 0.0977

2W -F 2e- H2 0.00 (by definition) Pb 2+ + 2e -9 Pb 0.1262 Sn24 + 2e_ Sn 0.1375 Ni24 + 2e - Ni 0.257 PbC12 + 2e -* Pb + 20F 0.2675 CO24 -F 2e -9 Co 0.280 Fe 2+ + 2e- Fe 0.447

Standard reduction potential; Cr 3+ + 3e -9 Cr 0.744 Cell EMF; Confirm what is Zn24 + 2e- Zn 0.7618 oxidized and what is reduced. A134 + 3e -9 Al -1,662

Mg2 + 2e -3 Mg 2.372 Na' + e -3 Na 2.71 Ca24 -I- 2e -4 Ca 2.868 K 4 + e' -3 K 2.931 Li 4

+ e -4 Li 3.0401 The standard oxidation potentials are the negatives of the values given here: The reactions are written in the opposite direction. Some of the values in this table were determined indirectly from other experimental results, since hypothetical electrodes (e.g., LiLi) are impossible to set up. The values are referenced to I bar pressure. See footnote on p. 320

Copper wire This process is accompanied by the passage of two electrons around the outer circuit. We could equally well write the reactions as

H2 -p W + e

+Cu2+ + e -4 +Cu

+ Cu2+ - --Cu + H

Sidearm for replacing This tells us that every time 0.5 mol of Cu2 disappears and 0.5 inol of Cu appears, KCI solution 1 mol of electrons passes from the left-hand electrode to the right-hand electrode.

Saturated KCI Other Standard Electrodes The standard hydrogen electrode is not the most convenient electrode because of the necessity of bubbling hydrogen over the platinum electrode. Several other elec-trodes are commonly used as secondary standard electrodes. One of these is the standard silver-silver chloride electrode, in which a silver electrode is in contact with solid silver chloride, which is a highly insoluble salt. The whole is immersed in potassium chloride solution in which the chloride-ion concentration is 1 m. This electrode can be represented as

Calomel and mercury paste

Mercury

Porous liquid junction a.


so that the overall process is

H, + Cu2 -4 Cu + 2H

Pt, H2 (1 bar) IH(1 m) Cl-(l m)AgCI(s)Ag with a salt bridge connecting the two solutions. The emf at 25 C is found to be 0.22233 V. The individual reactions are

* H + e e +AgC1-.Ag+Cl-

and the overall process is

-H2 +AgCl - H + Cl + Ag

The standard electrode potential for the silver-silver chloride electrode is thus 0.22233 V.

Another commonly used electrode is the calomel electrode, illustrated in Fig-ure 8.5a. In this, mercury is in contact with mercurous chloride (calomel, H92Cl2) immersed either in a 0.1 m solution of potassium chloride or in a saturated solution of potassium chloride. If the cell

Connecting wire

Buffer solution

Silver-silver chloride reference electrode

Thin glass membrane

AgAgC1(s)CF(l m) We can set up a cell involving this electrode and the hydrogen electrode,

FIGURE 8.5 The calomel electrode (a), and the glass electrode (b). The pH meter, commonly used in chemical and biological laboratories, often employs a glass electrode that is immersed in the unknown solution and is used with a reference calomel electrode.

Pt, 142(1 bar)tH(l m) C170.l m)Hg2Cl2(s)Hg is set up, the individual reactions are

-H2 H + e e + -H92Cl2 Hg + CF

and the overall process is

+ H92C12 -4 H + Cl-

+ Hg


The emf at 25 C is 0.3337 V, which is the standard electrode potential P. If a sat-urated solution of KC1 is used with the calomel electrode, the standard electrode potential is 0.2412 V.

Another electrode commonly used as a secondary standard is the glass elec-trode, illustrated in Figure 8.5b. In its simplest form this consists of a tube terminating in a thin-walled glass bulb; the glass is reasonably permeable to hy-drogen ions. The glass bulb contains a 0.1 m hydrochloric acid solution and a tiny silversilver chloride electrode. The theory of the glass electrode is some-what complicated, but when the bulb is inserted into an acid solution, it behaves like a hydrogen electrode. This electrode is particularly convenient for making pH determinations.

Ion-Selective Electrodes The glass electrode was devised in 1906 by the German biologist M. Cremer and was the prototype of a considerable number of membrane electrodes that have been developed subsequently. The importance of membrane electrodes is that some of them are highly selective to particular ions. For example, membrane electrodes have been constructed that are 104 times as responsive to Na as to K ions, and such electrodes are of great value for chemical analysis. Electrodes that are selective for more than 50 ions are now available, and most of them are membrane electrodes.

FIGURE 8.6 (a) The principle of the membrane I Potentiometer ______ electrode. (b) A membrane elec- circuit trode in which the membrane is a External reference single crystal, or a mixed crystal, electrode i Internal reference electrode or a matrix impregnated with a precipitate. (c) A membrane elec-trode

in which the membrane is a ___ Internal solution liquid ion exchanger. (d) An External solution enzyme-substrate electrode, containing ion which makes use of the ability of to be measured an enzyme to react selectively Ion-selective with an organic ion.

a

membrane

Internal electrode

Ion-selective membrane

Immobilized

b. C. d.

8.3 Thermodynamics of Electrochemical Cells 325

The principle of the membrane electrode is illustrated in Figure 8.6a. The sam-ple solution is separated from an internal solution by an ion-selective membrane, and an internal reference electrode is placed within the internal solution. An exter-nal reference electrode, such as a silversilver chloride electrode, is also immersed in the sample solution, and a measurement is made of the reversible emf of the as-sembly. The glass electrode is illustrated in Figure 8.5b, and some other assemblies are shown in Figure 8.6bd.

The theory of membrane electrodes is quite complicated and is different in detail for each type of electrode. It is not necessary for the ion to which the electrode is sen-sitive actually to be transported through the membrane. What occurs at the membrane is a combination of an ion-exchange process at the solution-membrane in-terface and the movement of various cations at the interface. It is not necessary for the ion to be measured to be especially mobile or for the same ion to be present in the membrane. A complete theoretical treatment of a membrane electrode requires a con-sideration of the Donnan equilibrium that is established (Section 7.13), of the Nernst potential (Section 8.3), and of complications arising from deviations from equilib-rium. Since so many factors are involved, the development of ion-selective electrodes is necessarily done on the basis of a good deal of empiricism.

8.3 Thermodynamics of Electrochemical Cells During the last century studies were made of the relationship between the emi of a cell and the thermodynamics of the chemical reactions occurring in the cell. These studies made important contributions to the understanding of the basic principles of thermodynamics. An early contribution was made by Joule who, with very sim-ple apparatus but with accurate temperature and current measurements, found in 1840 that

The calorific effects of equal quantities of transmitted electricity are proportional to the resistance opposed to its passage, whatever may be the length, thickness, shape, or kind of metal which closes the circuit; and also that, caereris paribus, these effects are in the duplicate ratio of the quantities of transmitted electricity, and, consequently, also in the duplicate ratio of the velocity of transmission.

By "quantity of transmitted electricity," Joule meant the current; by "duplicate ra-tio," the square. His conclusion was therefore that the heat produced was proportional to the square of the current 12

and to the resistance R. Since it is also proportional to the time t, Joule had shown that the heat is proportional to

J 2Rt

Since the resistance R is the potential drop V divided by the current I (Ohm's law), it follows that the heat is proportional to

I Vt

These conclusions have been confirmed by many later investigations. The SI unit of heat is the joule, that of current is the ampere, and that of electric

potential the volt; in these units the proportionality factor relating heat to IVt is unity:

qIVr (8.!)


This is readily seen by expressing the joule and the volt in terms of the base SI units (see Appendix A, Table A-i): thus

J kg m2 s 2 V kg M2 S-3 A-]

The product lVt is

AXkgnz2 s 3 AXskgm2 s 2 J

Joule's conclusion that the heat generated in a wire is IVt is quite correct, but later he and others went wrong. In 1852 he concluded that there is a correspon-dence between the heat of reaction of a cell and the electrical work. This error was also made by Helmholtz and by William Thomson (later Lord Kelvin). Thomson's conclusion appeared to be supported by his calculation of the emf of the Daniell cell from the heat of the reaction; his value, 1.074 V, is practically the measured value, but this agreement is accidental.2

It remained for Willard Gibbs to draw the correct conclusion, in 1878, that the work done in an electrochemical cell is equal to the decrease in what is now known as the Gibbs energy. This is an example of the deduction we have already made, in Section 3,7, that non-PV work (i.e., available work) is equal to the decrease in Gibbs energy.

This may be illustrated for the standard cell

Pt, H2 (1 bar) I H(l m) Cu2(1 m)ICu for which the emf (E) is 0.3419 V. The overall reaction is

H2 + Cu2 - 2W + Cu

Every time 1 mol of H2 reacts with 1 mol of Cu2 , 2 mol of electrons pass through the outer circuit. According to Faraday's laws, this means the transfer of 2 X 96485 C of electricity. The emf developed is +0.3419 V, and the passage of 2 X 96 485 C across this potential drop means that

2 X 96485 X 0.3419 CV = 6.598 X l0J

of work has been done by the system. Thus, for this cell process,

= 6.598 x 104

In general, for any standard-cell reaction associated with the passage of z electrons and an emf of E, the change in Gibbs energy is

Gibbs Energy Change = zFE (8.2) in a Cell

Since this Gibbs energy change is calculated from the E value, which relates to a cell in which the molalities are unity, it is a standard Gibbs energy change, as indi-cated by the superscript .

The same argument applies to any cell; if the emf is E, the Gibbs energy change is

AG = zFE (8.3)

21t results from the fact that there is only a small entropy change in the Daniell cell (&E/T = 0, see Eq. 8.23), so that A H G.


The AG is the change in Gibbs energy when the reaction occurs with the concentra-tions having the values employed in the cell. Note that if E is positive, AG is negative; a positive E means that the cell is operating spontaneously with the reac-tions occurring in the forward direction (e.g., H2 + Cu2 * 2W + Cu), and this requires AG to be negative.

For any reaction

aA+bB+ ... ... +yY+zZ

the Gibbs energy change that occurs when a mol of A at a concentration [A] reacts with b mol of B at [B], etc., is given by3

AG = RT1In K' - in!_{Y]Y[Z]Z \u1

L \[A]0[B?...)j (8.4)

If the initial and final concentrations are unity, AG is AGO and is given by

AG* = RTIn K (8.5) For any cell involving standard electrodes, such as the standard Daniell cell

znpzn2 (1 m) Cu2 (l m)Cu Eq. 8.2 applies and therefore

E0=,1nK' . (8.6)

zF

At 25 C this becomes

E/V = 0.0257 In K (8.7)

These equations provide a very important method for calculating Gibbs energy changes and equilibrium constants. The extension of the method to the calculation of AN and AS' values is considered on p. 335.

EXAMPLE &1 Calculate the equilibrium constant at 25C for the reaction occurring in the Daniell cells

if the standard emf is 1.10 V.

Solution The reaction is Zn + 2+ Zn2 + Cu

and z = 2. From Eq. 8.7,

zE 2 X 1.10 in 85.6 K' 0.0257 0.0257

and thus K = 1.50 X 10

'This is the approximate relationship in which concentrations rather than activities are used. In Eq. 8.4 and subsequent equations we are again using the superscript to indicate the numerical value of the quantities such as the ratio [Y]YEZI/EA]0[B]b.


EXAMPLE 8.2 Using the data in Table 8.1, calculate the equilibrium constant for the reaction

H2 + 2Fe3 2H' + 2Fe2

Solution From Table 8.1, the standard electrode potentials are

2H"+2e+H2 E0

Fe' +eFe2 E0.771V

The E value for the process H2 + 2Fe3 - 2H++ 2Fe2

for which z = 2 is thus 0.771 V. Therefore

0.771 = 0.0257 K

and

K = 1.14 X 1026

It should be emphasized that the emf is an intensive property, so that in making this calculation we must not multiply the value of 0.771 V by 2; the emf of 0.771 V applies equally well to the process

2Fe3 + 2e - 2Fe2

If the problem had been to calculate the equilibrium constant K' for the process

+H2 + Fe' # W + Fe2

the reactions would have been written as

H+e*H2 E0

Fe' + e -9 Fe' E = 0.771 V

and again E = 0.771 V. In this case z = I and

0,771 = 0.0257 in K' or in K' = 0.771 = 30.01 0.0257

and therefore

K' = 1.1 X 1013 or K' = 1.1 X 1013 moldm 3

K' is, of course, the square root of K. In general, a more complicated procedure must be used to combine half-cell re-

actions in order to find an E for a desired half-cell reaction, as is shown in the

following example.


EXAMPLE 8.3 Calculate E for the process

Cu + e -* Cu

making use of the following E values:

1. Cu2 + e -* Cu E? = 0.153 V 2. Cu2 + 2e -4 Cu E2 = 0.337 V

Solution The AGO values for these two reactions are

Cu2" + e -4 Cu = zEF = 1 X 0.153 X 96485 J mot-' Cu2 + 2e - Cu tG2

= zE2F = 2 X 0.341 >< 96485 J mot-'

The reaction Cu' + e - Cu is obtained by subtracting reaction 1 from reaction 2 and the AGO value for Cu + e - Cu is therefore obtained by sub-tracting AGO from AGO:

AGO = [-2 X 0.341 )< 96485 - (-1 X 0.153 X 96 485)] J mot-' (0.153 - 0.682) 96485 J moE'

= 0.529 )< 96 485 1 mot-' Since, for Cu + e -+ Cu, z = 1, it follows that

E = 0.529 V

Note that it is incorrect, in working the previous example, which is for a single electrode half cell, simply to combine the E values directly.

In view of this the student may wonder why it is legitimate to calculate E val-ues for overall cell reactions by simply combining the E values for individual electrodes. Consider, for example, the following E values involving both a one-electron and a two-electron process as in the last example:

Fe3 + e -* Fe2 E = 0.771 V I1+2e-2I E=0.536V

The procedure we have been adopting is to combine these two E values:

2Fe3 + 21- -* 2Fe2' + 12 E = 0.771 - 0.536 = 0.235 V

The fact that this is justified can be seen by writing the AGO values: Fe 3+ + e -4 Fe 2+ G = 1 X 0.771 X 96485 J mot 12 + 2e -4 21- AGO = 2 X 0.536 X 96485 J mol'

We combine the two equations by multiplying the first by 2 and subtracting the second:

2Fe3 + 2F -* 2Fe2 + 12 AGO = [2(-1 X 0.771 X 96 485) (-2 X 0.536 X 96 485)] J moE'

= 2 X 0.235 )< 96485JmoF' Thus E = 0.235 V. and this is simply 0,771 - 0.536. We are justified in simply subtracting E values to find E for an overall reaction, in which there are no elec-trons left over. However, to obtain an E for a half-reaction (as in Example 8.3), we


can in general not siniply combine F values but must calculate the .G values as iusL done.

The Nernst fqual R)fl So far we ha\ e limited our discussion to standaid electrode potentials 1; and to I:- ali.ies ku cells in which the active species are present at I in concentrations. The

correspondtii standard Gibbs C n CljLiCs htvc been '. ni I en as Let us now remove this rcstriction and consider cells in which the eoiceiiiri-

110115 are (ilici- than Lnity.( for exaitiple, the cell

Tie Ncrnsl equation: Follow its derivation Pt, H (I hat) H (Lt(l (aq)Co

iJ in which it hvdroeen electrode has been coiiihineul with a copper cIcLirode rn- meised in a Cu

solution. the concentration of which is other thLtii unit\. The overall cell react I( in is

'I H -f Cu - 217I - CU and the Gibbs energy change is (see Eq. 8.4)'

IH 2\u L(, RI In I' In

\ I Cu Ho'cvcr. ..(; = - RI In K = - I I .wiJ iheieliiie

H il 0 = :L I RI In( ----- i89 i

I('tr

Since AG : El-. here F is the em I* of ibis ccli. we tilnaiti

/ 1-I I I RI In l I -- I - I (8.10) \ CLI ' j!

and therefore

Ri / IH I\' = !' - - In , ) 8.11

:1 'i (Cui /

In neiieral. we may consider ally cell for which the overall react ion his the general I orin

IIA I)B 4 - -- + iY * G is ivcii h' Eq. 83 and \G = Ri In K - EL. and ihcreInc

\

(-soi-1-Y111zi

+ K/'ln -I(8.12) JAI"IB ---1

Note that, as in an equi I ibriuni constant, products are in the nuiiicraiur and reaciants in the denominator. Since .G = -- :LF, Eq. 8.12 leads to the following general expression for the cmi':

R7 I Y 11Z 1 Nernst Equation F = E

- -j IAI'IEI'

H

L,.1111 liii simph, ;i. \\e tie eii1LenIrlii;i' IiStL,iE( Ill - iii 111c, scc Seeiiii 5.. icie ii1ii\ my :lklj\ Ilk.' alt t'iiipIii.c.J.


This relationship was first given in 1889 by Nernst and is known as the Nernst equation.

Suppose that we apply the Nernst equation to the cell

ZnIZn2 Ni2 f Ni

for which the overall reaction is

Zn + Ni2 - Zn2 + Ni

Nernst equations: Simplify and The standard electrode potentials (see Table 8.1) are restrict it to 298 K. Ni2 + 2e_ Ni E = 0.257 V i::

Zn2 +2e>Zn E= 0.762V

and E for the overall process is 0.257 - (-0.762) = 0.505 V. The Nernst equa-tion is thus

E = 0.505(V) - in [Zn 2+1

(8.14) zF [Ni2 J

(As always, the concentrations of solid species such as Zn and Ni are incorporated into the equilibrium constant and are therefore not included explicitly in the equa-tion.) At 25 C this equation becomes

E/V = 0.505 - 0.0257 in [Zn 2- ] (8.15)

2 [Ni2 ] since z = 2. We see from this equation that increasing the ratio [Zn 2 ]/[Ni2 ] de-creases the cell emf; this is understandable in view of the fact that a positive emf means that the cell is producing Zn2 and that Ni2

ions are being removed.

EXAMPLE 8.4 Calculate the emf of the cell

CoJCo2 IINi2 jNi if the concentrations are

a. [Ni2 ]=1.0m and [Co2 ]=0.10m b. [Ni21 = 0.010 m and [Co2 ] = 1.0 m Solution The cell reaction is

Co + Ni2 - CO2+ + Ni

and from Table 8.1 the standard electrode potentials are

Ni2 +2e*Ni E-0.257V CO2 + 2e -4 Co E = 0.280 V

The standard emf is thus 0.257 - (-0,280) = 0.023 V. and z = 2. The cell emf at the first concentrations specified (a) is

0,0257 [CO21 ] E=0.023 n

2 [Ni

= 0.023 - 0.0257 In 0.10 = 0.023 + 0.030 = 0.053 V 2

10)1

Initial state Final state: All concentrations equal: M = 0

a. Membrane permeable to both ions

p -

]NJ

Initial state: Final state:

> [K12

Hardly any concentration change:

F C2 b. Membrane permeable to KI only

1 d

1 dm 1 d 1 d

1 cm c. Two cubic cells separated

by a membrane

FIGURE 8.7 (a) Solutions of KCI separated by a membrane that is permeable to both ions. (b) Solutions of KCI separated by a membrane that is permeable only to K. (c) Two 1 -dm3

cubes separated by a membrane of area 1 dm2 and thickness 1 cm.


Lob,

E = 0.023 - ____ 0.0257 1.0 2 0.010

= 0.023 - 0.059 = 0.036 V

We see that the cell operates in opposite directions in the two cases.

Nernst Potentials If two electrolyte solutions of different concentrations are separated by a mem-brane, in general there will be an electric potential difference across the membrane. Various situations are possible. Suppose, for example, that solutions of potassium chloride were separated by a membrane, as shown in Figure 8.7. If both ions could cross the membrane, the concentrations would eventually become equal and there would be no potential difference (Figure 8.7a). If, however, only the potassium ions can cross, and the membrane is not permeable to the solvent," there will be very lit-tle change in the concentrations on the two sides of the membrane. Some K ions will cross from the more concentrated side (the left-hand side in Figure 8.7b), and as a result the left-hand side will have a negative potential with respect to the right-hand side. The effect of this will be to prevent more K ions from crossing. An equilibrium will therefore be established at which the electric potential will exactly balance the tendency of the concentrations to become equal. This potential is known as the Nernst potential.

The Gibbs energy difference AG, arising from the potential difference sI is

= zFM (8.16)

where z is the charge on the permeable ions (z = I in this case). In our particular example (Figure 8.7b), the potential is higher on the right-hand side ((I) > 0), because a few K ions have crossed, and there will thus be a higher Gibbs en-ergy, as far as K ions are concerned, on the right-hand side (AG, > 0). The Gibbs energy difference arising from the concentration difference is

AG, =RTln (8.17) C l

At equilibrium there is no net AG across the membrane and therefore

Ge + AG, =ZF+RTlfl=O (8.18) Cl

Thus

(8.19) z C2

It is important to realize that in situations of this kind the Nernst potential is due to the transfer of only an exceedingly small fraction of the diffusible ions, so

5This restriction is made so that there will be no complications due to osmotic effects.


By convention we take the ratio of the concentration of ions outside the cell to the concentration of ions inside the cell. The potential obtained is that inside the cell.

EXAMPLE 8.5 Mammalian muscle cells are freely permeable to K ions but much less permeable to Na and Cl_ ions. 1'pica1 concentrations of K ions are

Inside the cell: [K] = 155 mM Outside the cell: [K] = 4 mM

Calculate the Nernst potential at 310 K (37 C) on the assumption that the mem-brane is impermeable to Na and CF.

Solution From Eq. 8.19 withz =1,

-

- 8.3145(J K' moF') X 310(K) In

96 485(C moF')

155

=0.0267 In = 0.0977 V 97.7 mV 155

The potential. is negative inside the cell and positive outside. In reality, such potentials are more like 85 mV, because there is a certain amount of diffusion of Na and Cl-, and there is also a biological pumping mechanism.

that there is no detectable concentration change. Suppose, for example, that two so-lutions 1 dm3 in volume are separated by a membrane 1 dm2 in area (A) and 1 cm in thickness (1), and suppose that 0.1 M and 1.0 M solutions of KC1 are present on the two sides (Figure 8.7c), the membrane again being permeable only to the K ions. The capacitance C of the membrane is

(8.20)

where E, the dielectric constant of the membrane, will be taken to have a value of 3 (this is typical of biblogical and other organic membranes). The capacitance in this example is therefore

8.854 X 10 12(C2 N' m 2 ) X 3 X 10 2(m2 ) C = 102(m)

= 2.66 X 10" C2 N_' m = 2.66 X lO' F

where F is the farad, the unit of electric capacitance. In terms of base SI units: F CV' As(kgm2 s 3 A')' A2 s4 kg' m 2

C2 N' m' (As)2 (kg ms 2 ' m A2 s4 kg m 2 F The Nernst potential at 25 C is

= 8.3145 x 298.15 In 10 = 0.059 V 96485

With a capacitance of 2.66 X 10-" F, the net charge on each side of the wall, re-quired to maintain this potential, is

Q=2.66X 10"FX0.0591V = 1,57 X 10'2C


The number of K ions required to produce this charge is

Q 1.57 X 10 2 C = 9.80 X 10

e 1.602X 10'9 C However, I dm3 of the 0.1 M solution contains 0.1 X 6.022 X 1023 = 6.022 X 1022 ions. The fraction involved in establishing the potential difference is therefore ex-ceedingly small (

83 Thermodynamics of Electrochemical Cells 335

At equilibrium

(0.2 = 00. I -~- .v) = 0.08 ,ti

The final concentrations are therefore

Palmitate side: [Nai = 0.18 M: [CI I = 0.08 M Other side: ENaJ = [Cl = 0.12 M

The Nernst potential arising from the distribution of Na ions is, at 25 "C,

- 8.3145 x 298.15 In 0J8 = 0,0 104 V = 10.4 mV - 96485 0.12

The palmitate side, having the higher concentration of Na' , is the negative side, since Na ions will tend to cross from that side. The Nernst potential calculated from the distribution of Cl ions is exactly the same:

8.3145 x 298.15 0.12 ACD = 96 485 In 0.08

= 10.4 iTiV

Strictly speaking D values measured by placing platinum electrodes in two solutions separated by a membrane are not precisely the same as these calculated Nernst potentials. The measured values are For the potential drop between the Pt electrodes, but the potential drop between the electrode and the solution is not quite the same for the two solutions. However, the error is about the same as the error in the calculations, and in clectrophysiological and other studies it is usually assumed that the measured values can be equated to those obtained in the calculations.

Temperature Coefficients of Cell emfs Since it Gibbs energy change can he obtained From the standard emf of a reversible cell (Eq. 8.2), the S and AW values can he calculated if ernf measurements are made over a range of temperature.

The basic relationship is Eq. 3.119.-

S =

and tor an overall reaction

As =

_(li) (8.22) 11,

Introduction of Eq. 8.3 gives

iiE AS = (8.23)

The enthalpy change is thus

1I = G + TS (8.24)

AH = :F(E - T-) (8.25)


The measurement of emf values at various temperatures provides a very convenient method of obtaining thermodynamic values for chemical reactions and has fre-quently been employed. For the results to he reliable the temperature coefficients should be known to three significant figures, and this requires careful temperature and emf measurements.

EXAMPLE 8.7 The einf of the cell Pt. H7( I har)IHCI(0.() I n)AgCl(s)jAg

is 0.2002 Vat 25 C aE/iTis 8.665 X 10 V K . Write the cell reaction and calculate AG, AS, and AU at 25 C.

Solution The electrode reactions are

71-12, * H1 + e and e + AgCI(s) 'p Ag + Cl - The cell reaction is

+ AgCI(s) - AL, + H' + Cl The Gibbs energy change is

AG = 96 485 X 0.2002 = - 19 320 i mol

The entropy change is obtained by use of Eq. 8.23:

AS = 96485 X 8.665 X 10 5 = 8.360 J K 1 mol'

The enthalpy change can be calculated by use of Eq. 8.25 or more easily from the AG and AS values:

AH = AG + TAS = 1.932 X 10 + (-8.360 x 298.15) = 2.181 X l0 J niol

8.4 Types of Electrochemical Cells In the cells we have considered so far there is a net chemical change. Such electro chemical cells are classified as chemical cells. There are also cells in which the driving force, instead of being a chemical reaction taking place between electrodes and solute ions, is a dilution process. Such cells are known as concentration cells. The changes in concentration can occur either in the electrolyte or at the electrodes. Examples of concentration changes at electrodes are foutid with electrodes made of amalgams or consisting of alloys and with gas electrodes (e.g.. the Pt. H electrode) when there are different gas pressures at the two electrodes.

Figure 8.8 shows a classification of ckctrochemical cells. A subclassificatiu,, of chemical and concentration cells relates to whether or not there is a boundary he twecfl two solutions. If there is not, as in the cell

Pt, HHCl(aq)AgCI(s)Ag we have a cell without tm,isfre,,c'e. If there is a boundary, as in the Daniell cell (Figure 8.1). the cell is known as a cell with transference. In the latter case there

8.4 Types of Electrochemical Cells

33

Electrochemical cells

FIGURE 8.8 l,issfication of electrochemical

- ,[Is. I Pt, H2 HClAgCI(Ag

Zn)Zn2 Cu2 Cu Pt. H21H1, Fe24, Fe3IPt

.1 Pt, H2 HCl(m1)AgClAg - AgClHCl(m2)Pt, H2 Pt, H2H(m1) H(m2)IPt, H2

; Na in Hg at c1(Na4tNa in Hg at C2 Pt, H2(P,)IH

4 JPt, H2(P2).

Chemical cells

Without With transference (1) transference (2)

Concentration cells

Electrolyte Amalgam Gas concentration or alloy electrodes (7) cells electrodes (6)

Cells in which Redox Without With the chemical cells (3)

transference (4) transference (5) reaction involves the electrodes

(:)ncentration cells; Watch a i;iction in a concentration cell Ii yen by a concentration difference.

is a potential difference between the solutionswhich can be minimized by use of a salt bridgeand there are irreversible changes in the two solutions as the cell is operated.

Concentration Cells A simple example of a concentration cell is obtained by connecting two hydrogen electrodes by means of a salt bridge:

Pt. l-1I}-tCl(nhi) HCI(m1)IH1, Pt The salt bridge could be a tube containing saturated potassium chloride solution. The reaction at the left-hand electrode is

- H2 --H (in1 ) + e while that at the right-hand electrode is

H (in,) + e -

The net process is therefore

HTh(ni,) - H (in1 )

and is simply the transfer of hydrogen ions from a solution of molality rn to one of molality ni l . If m is greater than m. the process will actually occur in this direc-tion, and a positive emf is produced: if in, is less than in1 . the cmi is negative and electrons flow from the right-hand to the left-hand electrode.

The Gibbs energy change associated with the transfer of 1-1 ions from a molal-ity rn to a molality in1 is

nilAG = RTln - (8.26) 111


Since z = 1, the emf produced is

E=ln - - (8.27) F m1

and is positive when rn2 > ?n 1.

EXAMPLE 8.8 Calculate the emf at 25 C of a concentration cell of this type in which the molalities are 0.2 m and 3.0 m.

Solution The emf is given by

E = In = 0.0257 In = 0.0696 V F in 1 0.2

Redox Cells Since, when cells operate, there is an electron transfer at the electrodes, oxidations and reductions are occurring. In all the cells considered so far these oxidations and reductions have involved the electrodes themselvesfor example, the H2 gas in the hydrogen electrode. There is also an important class of oxidation-reduction cells, known as redox cells, in which both the oxidized and reduced species are in solu-tion; their interconversion is effected by an inert electrode such as one of platinum.

Consider, for example, the cell

Pt, H21H(l in) Fe2, Fe3IPt The left-hand electrode is the standard hydrogen electrode. The right-hand elec-trode consists simply of a platinum electrode immersed in a solution containing both Fe 2+ and Fe 3+

ions. The platinum electrode is able to catalyze the interconver-sion of these ions, and the reaction at this electrode is

e 3-1- +Fe Fe'+

Since the reaction at the hydrogen electrode is

+H2 - H'3 + e

the overall process is

Fe33- + 42 -* Fe23- + H

The emf of this cell represents the ease with which the Fe 3+ is reduced to Fe 2+ The emf of the cell is

RT E= E - In [Fe 2+

]{F31

(8.28)

(Note that the E relates to P, = 1 bar and [if3- ] = [Fe 21] = [Fe31-] = 1 in.) The E for this system is +0.771 V (Table 8.1),

The interconversion of oxidized and reduced forms frequently involves also the participation of hydrogen ions. Thus, the half-reaction for the reduction of fu-marate ions to succinate ions is

8.4 Types of Electrochemical Cells 339

CHCOO +2W + 2e * CH2COO 11 1 CHCOO CH2COO fumarate succinate

If we wished to study this system, we could set up the following cell:

Pt, H2IH(l in) H F2 , S2, H(mH+)IPt where F2 and S2 represent fumarate and succinate, respectively. The hydrogen-ion molality in the right-hand solution is not necessarily 1 in; it will here be denoted as mH

-1- . If we combine the standard hydrogen electrode,

H2 -2H(1 m) + 2e we obtain, for the overall cell reaction,

F2 + 2H(mH+) + H2 -> S2 + 2H4(l m) The equation for the emf is thus, since z = 2,

RT /[S21 \u E = - - In I 2F [F2im+) (8.29)

The E for this system is related to a standard Gibbs energy change G by the usual equation

iXG = zFE (8.30) This standard Gibbs energy change is related to the equilibrium constant:

K= (8.31) [F2im+

However, it is frequently convenient to deal with the modified equilibrium constant

K' [S21 (8.32)

[F2 1

at some specified hydrogen-ion concentration. Often this standard concentration is taken to be 10-7M, corresponding to a pH of 7. In that case

K' = (10 7)2K = 10 4K (8.33) In many cases the K' value corresponds to a fairly well-balanced equilibrium at pH 7, whereas K will be larger by the factor 1014; the K' value and the corresponding iXG' at pH 7 therefore give a clearer indication of the situation at that pH.

Equation 8.29 for the emf can be written as

RT[S21 RT E = - - In + - In (8.34) 2F [F2 ] F

= - RT [S21 + 0.0257 (V) In mH+

(8.35) 2F [F21 RT[S21

= in 0.05916(V) pH (8.36)

2F {F21


where pH is the pH of the solution in which the S2-:F2 system is maintained (pH -10910 mH+ = - 2.303 in mH+). Then, if we define a modified standard potential E' by

RT E=E' - [S21 1n (8.37) 2F [F2 ]

it follows that

E' = E - 0.05916 (V) pH (8.38) Different relationships apply to other reaction types.

Redox systems are particularly important in biological systems, as illustrated by the following example.

EXAMPLE 8.9 The enzyme glycollate oxidase is a catalyst for the reduction of cytochrome c in its oxidized formdenoted as cytochrome c (Fe 3+)--bygly- collate ions. The relevant standard electrode potentials E', relating to 25 C and pH 7, are as follows:

E'IV cytochrome c (Fe') + e + cytochrome c (Fe 2+)0.250 glyoxylate + 2H + 2e -4 glycollate -0.085

a. Calculate G' for the reduction, at pH 7 and 25 C. b. Then calculate the equilibrium ratio

[cytochrome c (Fe 2 )}2[g1yoxylat&i [cytochrome c (Fe3'1')}2(glycollate

at pH 7 and 25 C. c. What is the equilibrium ratio at pH 7.5 and 25 C?

Solution a. The balanced reaction is 2 cytochrome c (Fe 3+) + glycollate -. glyoxylate + 2 cytochrome c (Fe2 ) + 2W and this equation corresponds to z = 2.

= 0.250 - (-0.085) = 0.335 V = -2 X 0.335 )( 96485

-64 645 J moV' = -64,6 kJ moF'

b. The equilibrium ratio at pH 7 is

K' = exp [64 6451(8.3145 X 298.15)] = 2.115 X 10"

c. The equilibrium ratio at pH 7.5 is conveniently calculated in terms of the true (pH-independent) equilibrium constant K e. if K" is the equilibrium ratio at pH 7.5,

= K' X (10_7)2 = K" X (I o-7.5)2 K" = 2.115 X 1011 X (1005)2

= 2.115 X 1012

8.5 Applications of emf Measurements 341

8.5 Applications of emf Measurements A number of physical measurements are conveniently made by setting up appro-priate electrochemical cells. Since emf values can be determined very accurately, such techniques are frequently employed. A few of them will be mentioned briefly.

pH Determinations Since the emf of a cell such as that shown in Figure 8.4 depends on the hydrogen-ion concentration, p1-I values can be determined by dipping hydrogen electrodes into solutions and measuring the emf with reference to another electrode. In commercial pH meters the electrode immersed in the unknown solution is often a glass electrode, and the other electrode may be the silver-silver chloride or the calomel electrode. In accordance with the Nernst equation the emf varies loga-rithmically with the hydrogen-ion concentration and therefore varies linearly with the pH. Commercial instruments are calibrated so as to give a direct read-ing of the pH.

Activity Coefficients Up to now we have expressed the Gibbs energy changes and emf values of cells in terms of molalities. This is an approximation and the errors become more serious as concentrations are increased. For a correct formulation, activities must be em-ployed, and the emf measurements over a range of concentrations lead to values for the activity coefficients.

Consider the cell

Pt, 1-120 bar) jHCI(aq) jAgCl(s)Ag The overall process is

+ AgCl - Ag + H + Cl-

and the Gibbs energy change is

AG = G + RT1n[a,a_]'

(8.39) where a+ and a_ are the activities of the H and Cl- ions, andAG' is the standard Gibbs energy change, when the activities are unity. The emf is, since z = 1.

E = E - RT ln[a+a_]

(8,40)

Since each of the activities a and a_ is the molality m multiplied by the activity coefficient y or y, we can write

RT a2 RT E=E --1 --ln[m} - ---lny+ y_ (8.41)

RT 2RT =E0__4 _lnm

- -- -lny (8.42)


,// where y, equal to V'ii, is the mean activity coefficient. This equation can be written as

(8.43)

-E - In and if E is measured over a range of molalities of HC1, the quantity on the left-hand side can be calculated at various molalities. If this quantity is plotted against m, as

+ ,.' shown schematically in Figure 8.9, the value extrapolated to zero m gives E, since ---------- at zero in the activity coefficient -/ is unity so that the final term vanishes. At any

molality the ordinate minus E then yields

2RT - j---

Molality/mol kg-1 from which the activity coefficient y, can be calculated.

FIGURE 8.9 A plot that will provide the value of E for a cell and the mean Equilibrium Constants activity coefficients at various concentrations.

The emf determination of equilibrium constants may be illustrated with reference to the measurement of the dissociation constant of an electrolyte HA such as acetic acid. Suppose that the following cell is set up:

Pt, H20 bar)IHA(mi), NaA(m2), NaCl(m1)AgCl(s)Ag The essential feature of this cell is that the solution contains NaA, which being a salt is essentially completely dissociated and therefore provides a known concentra-tion of A ions. The emf of the cell provides a measure of the hydrogen-ion concentration (since the hydrogen electrode is used); since A and H are known, and the total amount of HA is known, the amount of undissociated HA is known, and hence the dissociation constant can be obtained.

In more detail, the result is obtained as follows. The reactions at the two elec-trodes and the overall reaction are

-4 W + e e +AgCl - Ag+CF -H2 + AgC1 - H + Cl + Ag

The emf of the cell is

E = RT - -i-- ln[aH+acI_]u (8.44)

= E - RT -i- ln[mH+rncI_]u - RT in yH+ycl- (8.45)

E is the standard electrode potential of the silver-silver chloride electrode (0.2224 V at 25 C).

The dissociation constant for the acid HA is

mHmA- YH+YA- Ka = (8.46)

rnHA YHA and Eq. 8.45 can be written as

RT E = E - Ifl1 1 YHAY - - In - In K (8.47)

F L inA j F YA- F

cc

8.5 Applications of emf Measurements 343

or as

(EE)---- F +ln I mHAmCI- ' in YHAYCI - in K (8.48) RT

mA- ] YA- The molality mci- is equal to m, the molality of the NaCl solution. The value of mHA is equal to m1

and the value of m- is m 2 + mH+. The molality mH+ can be sufficiently well estimated from an approximate value of the dissociation constant; it is usually much less than m1 and rn2 so that not much error arises from a rough estimate. The quantities on the left-hand side of Eq. 8.48 are therefore known at various values of the molalities m 1 , tn2, and m3, and the left-hand side can be plotted against the ionic strength 1, as shown schematically in Figure 8.10. At zero ionic strength the activity coefficients become unity, and the first term on the right-hand side of Eq. 8.48 is therefore zero; extrapolation to I = 0 thus gives In Ka.

Similar methods can be employed for the determination of the dissociation constants of bases and of polybasic acids.

Solubility Products The determination of solubility products by emf measurements may be exempli- fied by the use of the following cell, which gives the solubility product of silver chloride:

Cl20 bar)fHCl(aq)AgCl(s)Ag The electrode processes are

Cl * C12 + e and e + AgCl(s) Ag + CF and the overall process is

AgC1(s) -4 Ag +--Cl2 However, the AgCl(s) is in equilibrium with Ag and CF ions present in solution, and we can write the overall process as

Ag + CF * Ag +

The emf corresponding to this process is6

E = E + RT-i

- ln[aAg+ac]-IM (8.49)

= E + I (8.50)

where E is given by

E = E(Ag + e -4 Ag) E0(fCl2 + e CF) (8.51) (see Table 8.1). At 25 C the value of E is 0.7996 1.35827 = 0.5587 V, and the measured emf of the cell is 1.140 V; thus

1.140= 0.5587 + RT

Ionic strength

FIGURE 8.10 Plot of the function shown on the left-hand side of Eq. 8.48 against the ionic strength.

It is of interest that E does not depend on the concentration of HCl used in the cell, since if a- is large, is correspondingly small.

800

700

600

500 DEIDV

400

300 0 1, Lu

200

11.2 IT'

11.3 11.4 11.5 V/cm3

b.

0 2 4 6 8 10 12 14 16 Volume V, of AgNO3 solution added/cm3

a.

FIGURE 6.11 0.40 Typical curves for a potentiometric titration. The example shown is for 0.38 the titration of a solution of sodium chloride with one of silver nitrate. At the end point there is a 0.34 sharp change in the electromotive force E (a), and a maximum in 0.30 & Fla V(b).

E/V 0.26

0.22

0.18

0.14


and

= 1.57 x 10-10 m012 dm -6

Similar cells can be devised for other electrochemical reactions. The general principle, for a salt AB, is to use the following type of cell:

Bisoluble salt of B ionslAB(s)IA The emf method is a valuable one for measuring solubility products for salts of very low solubility, for which direct solubility measurements cannot be made with high accuracy. A practical difficulty with the emf method is that sometimes the electrodes do not operate reversibly.

Potentiometric Titrations One of the most important practical applications of electrode potentials is to titrations. The procedure will be briefly explained with reference to a titration involving a pre-cipitation, but it can also be applied to acid-base and oxidation-reduction titrations.

Suppose that a solution of sodium chloride is titrated with a solution of silver nitrate; silver chloride, which is only slightly soluble, is precipitated, and at the end point of the titration the concentration of the silver ions in the solution rises sharply. The titration can therefore be followed by inserting a clean silver sheet or wire into the solution, and connecting it, through a salt bridge, with a reference electrode, such as a calomel electrode. Typical results of such a potentiometric titration are shown in Figure 8.11 a, in which the electromotive force, E, is plotted against the volume V of silver nitrate solution added to a solution of sodium chlo-ride. To obtain a precise value of the end point of the titration, 3EThV can be determined and plotted against V; as shown in Figure 8.11b, the end point is given by the maximum value. Modern instrumentation allows this differentiation to be obtained electronically.

H2 - 02

Anode Cathode

Anode reaction: H2 -4 21-1 + 2e

3.6 Fuel Cells 345

To carry out an acid-base titration potentiometrically, one uses any convenient form of hydrogen electrode, such as a glass electrode. At the end point of the titra-tion there is a sharp change in the concentration of hydrogen ions, which is reflected in a sharp change in the electromotive force. To carry out a titration of an oxidizing agent against a reducing agent one measures the potential of the redox system, which again varies sharply at the end point. In this case an inert electrode, such as a platinum electrode, is employed.

8.6 C Fuel Cells The first fuel cell was constructed in 1839 by the British physicist and lawyer Sir William Robert Grove (1811-1896); it had platinum electrodes with hydrogen bub-bled over one electrode and oxygen over the other. For many years little was done to develop fuel cells for commercial purposes, but since the 1960s there has been a considerable revival of interest in this problem, particularly in view of present en-ergy shortages. Fuel cells have been used as sources of auxiliary power in spacecraft, and major research efforts are under way to develop their use in auto-mobiles in order to minimize air pollution and noise.

Fuel cells employ the same electrochemical principles as conventional cells. Their distinguishing feature is that the reacting substances are continuously fed into the system, so that fuel cells, unlike conventional cells, do not have to be discarded when the chemicals are consumed. The simplest type of fuel cell uses hydrogen and oxygen as fuel. Figure 8.12 is a schematic representation of a hydrogen-oxygen fuel

FIGURE 8.12 Diagrammatic representation of a hydrogen-oxygen fuel cell, show-ing the reactions occurring at the two electrodes and the overall reaction.

Cathode reaction: 02 + H20 + 2e -, 20H Overall reaction: H2 + 2102 - H20 AG' = -237.2 kJ mol 1


cell and shows the reactions occurring at the two electrodes. Various electrolyte so-lutions, such as sulfuric acid, phosphoric acid, and potassium hydroxide solutions, have been employed in such cells.

The overall reaction in this type of fuel cell is

H2(g) + 02(9) - H200) This process corresponds to the transfer of two electrons. The standard Gibbs en-ergy change in this reaction [i.e., the standard Gibbs energy of formation of H20(l)] is 237.18 kJ moF' at 25 C, and the theoretical reversible emf, from Eq. 8.2, is

237180 =l.23V 2 x 96485

In practice voltages are less than this, because of deviations from reversible be-havior. The extent of these deviations depends on the materials used as electrodes. Much present research is devoted to the development of improved electrodes. If the difficulties can be overcome, fuel cells will offer a considerable advantage over other methods of obtaining energy from fuels. As discussed in Section 3. 10, fur-naces in which fuels are burned suffer from second-law limitations to their efficiency. There are no such Carnot limitations to the efficiencies of fuel cells.

Because of the hazards of using hydrogen, and the cost of preparing it from other fuels, attention is being given to the development of fuel cells using other materials, such as hydrocarbons and alcohols. Much modem research has been done along these lines, some of it focusing on specially designed solid oxide elec-trodes that operate at 1000 C. The focus is now on cells that operate at much lower temperatures of 600 C and below. These cells consist of a number of units, each one of which is a ceramic device consisting of an electrode which uses air as the oxidant, the electrolyte, and a fuel electrode. These sets are arranged in stacked layers to provide the desired voltage, the oxidant and fuel being able to flow between the electrodes. Such solid oxide fuel cells are capable of operating at high energy conversion efficiency and with low pollutant emissions, but a num-ber of practical problems still remain. Other types of fuel cells that operate at much lower temperatures include those based on polymer electrolytes, phosphoric acid, or alkaline systems, and are well under development.

Attempts are also being made to develop biochemical fuel cells that would use as fuels the products of biological processes. For example, at the bottom of the Black Sea, bacteria obtain oxygen from sulfates and produce large amounts of hy-drogen sulfide that could be used as anodic fuel. It will probably be many years before these speculative possibilities are turned into practical devices.

8.7 Photogalvanic Cells Certain materials, such as selenium, produce electricity directly when they are irra-diated, and devices that employ this effect are known as photovoltaic cells. Such cells are to be distinguished from photogalvanic cells, in which irradiation induces a chemical process that in turn gives rise to an electric current. The generation of electricity by the chemical system follows the same principles as in an ordinary electrochemical cell; the special feature of a photogalvanic cell is that the reaction is brought about photochemically.

8.7 Photogalvanic Cells 347

A number of devices of this kind have been set up on the laboratory scale, but more remains to he achieved in the direction of producing photogalvanic cells that will be of practical use in the utilization of solar energy. One reaction that has been studied in the laboratory is the light-induced reaction between the purple dye thio-nine and Fe2

ions in aqueous solution. If we represent the thionine ion as T. and its reduced colorless form as TI-I ' the overall reaction is

Ti-Fe21 +H TH - Fe

pLilpic

In the dark the equilibrium for this reaction lies over to the left. Irradiation with visible light caises this equilibrium to shift considerably to the right, and the solu-tion becomes colorless. The most effective wavelength for bringing about this equilibrium shift is 478 nm, corresponding to the yellow-green region of the spec-[rum. In the dark the reaction has a positive standard Gibbs energy change, but absorption of photons by the thionine molecules provides energy for displacement of the equilibrium.

A cell that utilizes this equilibrium shift is shown schematically in Figure .l 3. There are two compartments separated by a membrane that is impermeable

to thionine but permeable to Fe and Fe3 . The left-hand compartment, which

contains thionine as well as Fe2 and Fc is irradiated. The right-hand compartment contains no thionine and is kept in darkness. The irradiation of the

Voltmeter

e f I Light t

TH T - -

Fe2 Fe3 Fe2 J&lI

Pt Pt Anode Cathode

I IGURE 8.13 Membrane

liomatic representation of a impermeable togalvanic cell. to thionine

348 Chapter 8 Electrochemical Ce115

left-hand compartment causes the ratio [TH]/[T] to be abnormally large, and the process

TH -4 T + H + e

therefore occurs at the electrode. in the dark right-hand solution, the process - 3+

e +Fe *Fe

occurs at the electrode. The overall reaction giving rise to the emf is thus TH + Fe 3+ -4 T + Fe2 + H

and is the reverse of the reaction that is occurring as a result of the irradiation. As long as the light shines on the thionine solution, the current flows; but when it is turned off, the reaction rapidly reverts to equilibrium and the current ceases. The theoretical standard emf for the system is 0.47 V. The reversible potential obtained depends on the concentrations, in accordance with the Nernst equation, and these concentrations depend on the initial concentrations of the materials and also on the intensity of the radiation. In practice the voltages obtained are only 2 to 3% of the theoretical values because of deviations from reversibility.

There are formidable problems that make it difficult to make effective use of solar radiation with photogalvanic cells. The most important of these relates to the small proportion of light that is absorbed. With thionine the maximum absorption is at a wavelength of 478 am, and even when monochromatic light of this wavelength is used only about 0.1% of the light is absorbed in a typical experiment. With sun-light, covering a wide range of wavelengths, the fraction absorbed is reduced by a further factor of 10 or more, since most of the visible and near-ultraviolet radiation lies outside the absorption region of the dye.

Photogalvanic cells thus represent an interesting laboratory device, but only if completely new systems are used will they become of practical importance.

8.8 ' Batteries, Old and New In this section we discuss details of several commercial batteries, but first we define some terms. The expression electrochemical cell, or more simply, cell, is a general term that applies to a device designed either to convert chemical energy into elec-tricity or electricity into chemical energy. A device that is designed specifically to

Voltaic Cell

convert chemical energy into electricity is called a voltaic cell, a galvanic cell, or a battery. In this book we will often use the word battery for any device that gener-ates electricity, even though this word is sometimes regarded as applying only to a single cell; we will follow modern dictionary usage and use the word battery for a single cell that generates electricity, as well as for a battery of such cells.

Batteries can also be classified according to whether or not they are capable of being recharged. One that cannot be recharged is known as a primary cell or bat-tery. A cell or battery that can be recharged is commonly referred to as a secondary; storage, or rechargeable battery. In the United Kingdom a storage battery is com-monly called an accumulator

Electrolytic Cell We will use the term electrolytic cell to refer to a cell through which an exter- nally generated electric current is passed in order to produce chemical action, such as electrolysis. This term could be applied to a battery that is being recharged.

The convention is that when a battery is in the discharge mode the electrode of negative sign, where oxidation occurs with the electrode providing electrons to the

8.8 Batteries, Old and New

349

FIGURE 8.14 Relationship between voltaic and electrolytic cells. The voltaic cell on the left provides the electrical current to operate the cell on the right. The electrode signs are dif-ferent for the two types of cell, but oxidation occurs at the anode and reduction occurs at the cathode.

Anode (oxidation) Zn -i Zn2 + 2e negative pole

Zinc j---,-

electrode c Zn2 -v'

Copper - electrode

Cathode (reduction) + Cu2 +2e-* Cu positive pole

Voltaic cell (specifically, a

Daniell cell)

ZnSO4 solution

Porous barrier

CuSO4 solution

Cathode (reduction) 2H-i-2e-* H2

+ negative pole

-H 'f Inert electrodes

Dilute HI

- Anode (oxidation) 21-*12+2e

.J negative pole

Electrolytic cell

external circuit, is called the anode. In Section 8.1 we discussed the Daniell cell or battery, where the anode is the left-hand electrode. The potential in the external cir-cuit, as measured with a voltmeter, is positive. Figure 8.14 shows the relationship between batteries and electrolytic cells.

The reason a primary battery cannot be recharged is that the electrodes and electrolytes are consumed and cannot be restored to their original state by reversing the flow of electrical current. The Daniell cell (Figure 8.1) is an example of a pri-mary battery. A secondary battery on the other hand may be assembled in a discharged state and be charged or recharged when its original reactants have be-come depleted. Passing current from an outside source through the cell in the direction opposite to the discharge process can reverse the spontaneous discharge process. The battery therefore functions as an electrolytic cell in the recharge process.

Most modern primary batteries are based on aqueous electrolyte systems, but free liquid in the battery introduces the possibility of leaks. Thus almost all com-mercial primary batteries have the electrolyte immobilized in a gel or incorporated into a microporous separator that prevents direct interaction between the electrodes. Although some water is present in the gel such batteries are referred to as dry cells, where the word cell now refers to the smallest physical unit having the components necessary to produce a voltage. Several cells may be linked in series (anode of 1 to cathode of 2, and so forth) to make a battery capable of producing a higher voltage. Thus six cells each with a potential of 1.5 V are needed to form a 9 V battery. We now consider some primary batteries.

The Original and Modified Leclanch Cell The modern Leclanchd cell, first proposed by George Leclanch in 1866, has evolved into the modern "heavy duty" battery consisting of a zinc anode in the form of a cylindrical vessel that forms the outside container. Inside it is placed a central inert cylindrical carbon (graphite) rod, C(gr), acting as the electron collec-tor, with a paste of manganese dioxide as cathode surrounding the graphite rod. This latter combination is often referred to as the electrode, although only the ac-tual cathode (or anode) material undergoes change when the discharge occurs. The


electrolyte consists of an aqueous solution of zinc chloride, in gel form or incorpo-rated into an inert filter that occupies the space between anode and cathode. A porous liner separates the zinc can from the cathode, thereby preventing internal discharge (an electrical short). The cell is represented as follows:

Zn(s)ZnC12(aq)MnO2(s), C(gr) The half-reaction at the anode where oxidation occurs is

Zn(s) Zn2 (aq) + 2e and the overall cell reaction can be summarized as

4Zn(s) + 8Mn02(s) + ZnCl2(aq) + 8H200) 8MnOOH(s) + ZnC12 4Zn(OH)2(s)

The potential of this cell is about 1.5 V. This cell was modified from the origi-nal Leclanch cell, which contained an ammonium and zinc chloride electrolyte. One advantage over the original cell is that it has less polarization (caused by for-mation of a layer of ammonia molecules on the cathode) and thus avoids the subsequent reduction of voltage under heavy drain. This change allows an almost twofold increase in battery life over that of the original battery, and better service capacity under high current drain and continuous discharge.

When naming batteries it is common to place the anode name first, followed by the cathode name. However, deviations from this naming method will sometimes be found. Thus we have the zinc-carbon cell with the bottom of the Zn case in a typical "D"-cell battery being designated negative, and with the cathode center top being positive.

Alkaline Manganese Cells In alkaline batteries the electrolyte is a concentrated aqueous solution of ( 30 wt. %) potassium hydroxide that has been partially converted to potassium zincate by the addition of zinc oxide. The cathode is a compressed mixture of MnO, and graphite, moistened with electrolyte in contact with a carbon rod that acts as a cur-rent collector. The anode is a cylindrical zinc vessel in contact with the other current collector, an external steel-can casing being used to prevent leakage through the case. For most batteries, increasing the rate of current drawn from the battery decreases the voltage; however, this system gives a relatively constant voltage even at high discharge rates. The cell reaction in this battery is

Zn(s) + 2MnO2(s) + H200) 2MnOOH(s) + ZnO(s) Reactions involving Mn02 are much more complex than given in this equation be-cause of possible further reduction of the manganese. Also, because of impurities in the raw material, the source of Mn02 plays a role in the reactions that are observed. The cell potential for the alkaline cell is 1.55 V at room temperature.

ZincMercuric Oxide and ZincSilver Oxide Batteries Both of these batteries are familiarly known as 'button" cells. Both cells have a high volumetric capacity (that is, they can supply a high amount of current from a

8.8 Batteries, Old and New 351

particular volume), and the zincsilver oxide cell has the added advantage of hav-ing an almost constant voltage even under conditions of high discharge. The zincmercuric oxide cell or "mercury cell" used in watches is of environmental concern because, in its discharged state, it can release toxic elemental mercury. Only the chemistry of the zincsilver oxide battery will be discussed here because the chemical reactions in both are quite similar. The center cap of the battery is the negative pole formed by a steel electron collector in contact with the zinc anode. This is insulated from the rest of the positive steel case that contains the silver ox-ide cathode in a KOH electrolyte. The internal contact is made through a porous separator. The zincsilver oxide cell may be written as

Zn(s) jZnO(s)KOH(aq)Ag2O(s), C(s) The overall cell reaction is

Zn(s) + A920(s) ZnO(s) + 2Ag(s) where the anode (oxidation) reaction is

Zn(s) + 20H(aq) ZnO(s) + H200) + 2e and the cathode (reduction) reaction is

A920(s) + H200) + 2e - 2Ag(s) + 20H(aq) The for this battery is 1.60 V, whereas the overall reaction for the mercury battery is

Zn(s) + HgO(s) ZnO(s) + Hg(l) and its Ece11 is 1.357 V.

Metal-Air Batteries The primary objective of miniature battery design is to maximize the energy density or cell capacity in a small container. What better way is there to accomplish this than to cause one of the necessary components to flow continuously into the cell? In metal-air batteries the cathodic reactant is the oxygen in the air, which need not be stored within the confines of the battery. In essence, this is similar to one-half of a fuel cell. (See Section 8.6.) Aluminum, lithium, magnesium, and zinc have been ex-amined extensively for use as the anode in these systems. The first three suffer from severe corrosion problems; only zinc is satisfactory and is used in hearing-aid batter-ies. The zinc anode is separated from the cathode by a separator disc that allows oxygen to pass but prevents liquid from escaping. Many of the techniques used are proprietary, that is, trade secrets. The oxygen-electrode reaction is complex but in ba-sic solution it may be considered as a two-stage process, written as

02(aq) + H200) + 2e HO(aq) + 0H(aq) followed by

HO(aq) + H2O(l) + 2e - 30H(aq) The hydroperoxide ion may form a metal-oxygen bond that may be reduced or the hydroperoxide ion may decompose to re-form oxygen. The anode reaction may be written as

Zn(s) + 40H(aq) - 2e Zn022 (aq) + 211200)


When the solution is saturated with zincate ion, zinc oxide is formed

Zn(s) + 20H(aq) - 2e ZnO(s) + H200) Preventing zinc-oxide passivation of the cell (in which the anode fails to function properly) is a major design problem. The cell voltage is about 1.4 V. which drops to about 1.25 V when half of the anode material is consumed. The volumetric energy density, however, is about 30% better than that of the closest competitor. the zincsilver oxide cell.

Nickel-Cadmium (Ni-cad) Secondary Battery The Nicad battery should be called the cadmium-nickel battery in keeping with the convention that the anode should be named first. Because of its rechargeability, this battery is popular in portable hand tools, but it presents a disposal problem because of the toxicity of cadmium. In this battery the anode reaction is

Cd(s) + 20H(aq) Cd(OH)2(s) + 2e The cathode reaction is

2NiO(OH)(s) + 21-120(1) + 2e 2M(OHMS) + 20W(aq) Combining these gives the overall cell reaction

discharge Cd(s) + 2NiO(OH)(s) + 2H20(1) = 2M(OHMS) + Cd(OH)2(s)

charge

The Eceii = 1.4 V and the reaction can be reversed upon recharging as indicated by the double arrow.

The Lead-Acid Storage Battery The familiar lead-acid battery was invented by Gaston Plantd in 1859 it should be named the leadlead oxide battery to be consistent with practice. Today more than a third of the world's output of lead goes into constructing these batteries. The typical lead-acid battery anode has a negative lead grid support filled with porous lead (known as lead sponge) where oxidation occurs, and a positive lead grid support cathode with a filler of lead dioxide, Pb02. Both anode and cathode electrodes are submerged into a solution of sulfuric acid with a porous separator sheet between them to keep the plates from touching. The cell is represented by

Pb(s)IPbSO4(s), H2SO4(aq)PbSO4(s)IPb02(s)Pb(s) The overall process may be written as

Pb(s) + Pb01(s) + 2H1SO4(aq) = 2PbSO4(s) + 2H20(1) The anode (oxidation) negative electrode is

discharge Pb(s) + S042 (aq) PbS0(s) + 2e

charge

In the discharge process lead sulfate coats the lead electrode, and charging causes regeneration of the lead. The cathode (reduction) electrode reaction is

e_j

Active graphite

I

:1 Li

ions

Electrode (electron collector)

C(gr)

Charged anode

Li Intercalated ions Li metal

FIGURE 8.15 Intercalation. In the charging process electrons from an outside source are supplied along the conductor to the electrode, which will be the anode in the discharge process. In the process of interca-lation shown, ions of Li are Inserted into the structure of the graphite and, in the process, each gains an electron from the con-ductor without damage to the orig-inal structure of the host. The process is reversed in the dis-charge mode.

8.8 Batteries, Old and New 353

discharge Pb02(s) + 4H(aq) + S042 (aq) + 2e PbSO4(s) + 2H20(I)

charge

Again, lead sulfate forms on the lead dioxide plates. As the cell is discharged, sul-furic acid is consumed and water is formed so that the state of charge of the battery may be determined from the density of the electrolyte. At full charge, each cell produces 2.15 V at 25 C. This value is different from that calculated at standard-state conditions and is typical for all batteries of this type. At full discharge the voltage drops to 1.98 V. Thus, the charged state of the battery may be followed by measuring the specific gravity of the sulfuric acid. When the specific gravity of the sulfuric acid falls below 1.2, the battery needs to be recharged. The open-circuit voltage may be predicted thermodynamically from the sulfuric acid and water ac-tivities and the temperature. For automotive use, six cells are placed in series, giving a voltage of 12.6 V. Complex chemistry and careful design engineering are involved in making such batteries function for five or more years, and the reader is referred to the references.

Lithium Ion Batteries Familiar in cellular phones, camcorders, and computers, lithium ion batteries offer distinct advantages. They are rechargeable and have a high volumetric energy den-sity of 300 W h/dm3, where W h is the watt hour, and with a gravimetric energy density of 125 W h/kg, far surpassing most batteries. However, their development required understanding of a more complicated chemistry than that previously dis-cussed. The materials used as electrodes in lithium ion batteries were first patented in 1981. Curiously, these batteries do not contain lithium as isolated metal. The chemistry involves a process called intercalation, in which a lithium ion is inserted as Li metal into a host without damaging its structure. See Figure 8.15, where, in the charging process, graphite intercalates lithium. The process is

Li + e + C6 LiC6


Anode M Electrolyte M Cathode

FIGURE 8.16 C (gr) LiPF6 salt and LICoO2 Discharge process for the lithium

organic carbonates (LNiO2 or LiMn204) ion battery. Electrons leave the anode conductor via the external Active I circuit. Li ions move through the graphite electrolyte (LiPF6 salt and organic carbonate in the case shown) to ' Ll the cathode in the Li0 5CoO2 state, where electrons from the external circuit enter the cathode from the

LiC6 Li + e + 2Li05C0O2 conductor and the Li fully loads

-

the cathode with Li metal. C6+Lr+e- 2LiC002

where external electrons are supplied via a carrier to the active material, graphite. On discharge, the process is reversed and the Li-loaded graphite anode de-intercalates lithium as L11 with the electrons going into the external circuit. Several materials

FIGURE(LiCo02, LiNi02, or LiMn204) may be used as the cathode material. The selection

(a) LiCO02 in its discharged state of the actual material depends upon the particular characteristics desired from the is fully loaded with Li. (b) Charg-

cell. We will use LiCo02 as an example throughout this discussion. At the cathode, ing of LiCoO2 removes Li from the LiCoO2 de-intercalates lithium during charging. Thus the processes at both elec- structure, leaving Li*

now in the trodes are somewhat similar. electrolyte. (c) Overcharging the

, + - LiCoO2 removes too much lithium,

2LiC002 2Li0 CoO2 + Li + e causing its structure to collapse which releases oxygen to the During discharge, Li0 5CoO2 intercalates lithium ions. The overall reaction for atmosphere and thereby shortens

the discharge process is shown in Figure 8.16, where the salt LiPF6 and an or- battery life. ganic carbonate serve as the electrolyte. In one design of the battery, an organic

ON* (D'SIS4 0000 111011104W

.

: '.'SSSYSSSS00 .... : GCQQo

Discharged Li Charged LI Overcharged graphite graphite graphite

Key Equations 355

FIGURE 818 fluoropolymer is used as the electrolyte to transfer ions, but not the electrons. Discharged graphite is loaded with Since charging involves just the reverse processes, it is instructive to consider LI in the charging process. Over- what happens when LiCo02 and graphite are overcharged. In Figure 8.17 LiCoO2 charging of the graphite allows the exce

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Laidler Physical Chemistry Chpt. 8