The existence and importance of the potential at a liquid junction have been noted in several places in this chapter. The detailed treatment is somewhat spe
cialized, however, and has therefore been placed in this section. We consider
first the potential of a cell having a liquid junction and then some aspects of the electrochemistry of the liquid junction itself.
A. Concentration Cells with a Liquid Junction
The least complicated type of cell which has a liquid junction is a concentration cell since the electrode reactions are then the same, except for a difference in concentration of the electrolyte across the junction. The general cell diagram is
/
anode/solution at mxj solution at ra2/cathode,
where the dashed line means that the two solutions are in electrolytic contact but are somehow prevented from mechanical mixing. This may be by means of a porous diaphragm, as in the Daniell cell of Fig. 13-2, or by a stiffening of one of the solutions at its point of contact with the other by agar-agar or gelatine, as is done in a salt bridge.
An example of a concentration cell with a liquid junction is the following:
Pt/H2(l atm)/HCl(m1)///HCl(m2)/H2(l atm)/Pt, (13-41) anode reaction ^Η2(1 atm) = Η+Ο^) + e~,
cathode reaction H+(m2) + e~ = | H2( 1 atm), (13-42) net reaction H+(m2) = H+(mx).
The emf of a cell corresponds to the free energy change associated with the sum of all processes occurring per faraday of electricity passed through the cell.
Current is carried by the ions in the solution, which means that t+ equiv of hydrogen ion must cross the junction from left to right and t_ equiv of chloride ion must do so from right to left, where t denotes transference number (see Sec
tion 12-6). The changes that occur in addition to the electrode reactions are then
Ζ+Η+ίΛΐ!) = /+H+(m2), i_Cl-(/w2) = tJZ\-{mx). (13-43) The sum of process (13-42) and processes (13-43) is
/_H+(m2) + t_C\-(m2) = r_H+(^i) + /_Cl-(/Wi), (13-44) bearing in mind that t+ + t_
δ
or
ê
If m1 an d m2 ar e 0. 1 an d 0.01 , respectively , the n th e γ± values are 0.796 and 0.904, respectively, and the t_ values are 0.1686 and 0.1749, respectively. The
1. The corresponding Nernst equation is
RT aH+,mlaCl~,ml
H+,m, C\-,m
(13-45)
^ a±,m9
SPECIAL TOPICS , SECTIO N 1 52 9
average valu e o f t_ i s 0.1717 , an d th e em f o f th e cel l a t 25° C i s calculate d t o b e
ë = - ( 2 X 0 . 1 7 1 7 X 0 . 0 2 5 6 9 ) I n | ^ $ g g = = - 0 . 1 9 1 9 .
The overal l cel l reactio n i s th e su m o f processe s (13-42 ) an d (13-43) , an d th e separate Nerns t expression s fo r thes e ar e
RT aH+ ~
« fE= I n - — - i - ( 1 3 - 4 6 )
and
RT aH+ m, RT aC\- m,
*, = - t + — I n - / _ I n . (13-47 )
^ aH + ,mi ^ aCl-,m2
Inspection confirm s tha t th e su m o f Eqs . (13-46 ) an d (13-47 ) give s Eq . (13-45) ; thus
g = (oE -\- ê} .
The em f fo r th e cel l ca n therefor e b e viewe d a s mad e u p o f tw o terms : tha t du e t o the electrod e reaction s SE , an d th e junctio n potentia l S}, whic h give s th e wor k associated wit h th e transpor t o f ion s acros s th e liqui d junction .
We ca n exten d th e precedin g numerica l exampl e t o th e calculatio n o f <fE an d
<f j onl y b y makin g som e assumptio n regardin g individua l ioni c activit y coefficients . If, fo r example , w e assum e tha t y H+ = 7 c i " = 7 ± ·> the n Eqs . (13-46 ) an d (13-47 ) give <?E = —0.0558 8 V , < f j = 0.0366 9 V . Th e junctio n potentia l represent s a majo r correction i n thi s case .
In usin g Eq . (13-45) , a n averag e valu e o f r _ fo r th e tw o concentration s wa s used. I t turn s ou t tha t th e correc t averag e i s no t th e simpl e arithmeti c one , an d the wor k don e i n transportin g a n io n acros s th e junctio n actuall y involve s th e integral
w2
RT UdtfaOi), ( 1 3 - 4 8 )
so tha t th e correcte d versio n o f Eq . (13-45 ) i s RT r™2
S = - 2 ^ - \ t_ rf(ln a±). (13-49 )
A stud y o f cell s o f thi s typ e evidentl y permit s th e determinatio n o f transferenc e numbers a s a functio n o f concentration .
As a las t point , th e overal l cel l reactio n wil l depen d o n whethe r th e electrode s are reversibl e t o th e anio n o r t o th e cation . Thu s fo r th e cel l
Ag/AgCl/HCKiwo/HCK/WaVAgCl/Ag
Eqs. (13-45 ) an d (13-49 ) woul d hav e r + rathe r tha n r _ , an d <fE woul d hav e th e opposite sign . Th e numerica l exampl e woul d the n giv e
g = <fE + ^ , o r 0.0925 7 = 0.0558 8 + 0.03669 .
The sig n o f th e junctio n potentia l remain s th e same . Thus , althoug h w e analyz e
in detail each specific concentration cell to determine the correct overall cell reaction, the junction potential remains the same for any given junction.
6 . Junction Potentials
Equation (13-47) simplifies considerably if activity coefficient effects are neglected. We then obtain
This form brings out the important point that «f j depends on t+ — t_ . It is largely with this factor in mind that one expects the junction potential to be small in the case of a KCl bridge; t+ for 1 m KCl is 0.488, and so t+ — t_ is only 0.024. The same situation applies in the case of N H4N 03, an electrolyte that is used in a salt bridge if it is not desirable that Cl_-ion-containing solution have contact with the electrolyte of the half-cell being studied.
If various electrolytes are present on either side of a liquid junction, then the total work of transporting ions across the junction per faraday, which gives , depends on weighted averages or integrals such as that of Eq. (13-48). By using concentrated KCl in the salt bridge, one ensures that the important t{ will be those of K+ and Cl~", so that <f j will not be very sensitive to the nature of the more dilute electrolyte solution of the half-cell into which the salt bridge dips. The analysis makes it clear, however, that although we can reduce a junction potential to a fairly small value, cells with liquid junctions retain a residual uncertainty in the interpretation of their emf's and one which is difficult to analyze exactly.
Three types of polarization effects were noted in Section 13-9: ohmic, concentra
tion, and activation. We now discuss these last two in more detail.
A. Activation Polarization
The Tafel equation (13-38), which relates the activation polarization component of the overvoltage to the current /, may be accounted for in terms of the following analysis. Consider a general electrode reaction
where, as illustrated in Fig. 13-13, the oxidized state must pass through some high-energy intermediate I in the process of being reduced. We write an equation analogous to Eq. (8-65) for the rate of the forward process:
(13-50)
13-ST-2 Polarization at Electrodes. Polarography
e~ + O (oxidized state) = R (reduced state),
(O), (13-51)
where AG%% is the standard free energy change to go from Ο to I. It will be known
SPECIAL TOPICS, SECTION 2 531 Intermediate I
Reduced state R FIG. 13-13. Mechanism for activation polarization.
in the next chapter as the free energy of activation. Similarly, the rate of the back reaction is
*-[τΜττ4ΐ < >-
R (13-
52>
The presence of a potential difference φ between the electrode and the solution contributes to AG°* and we write this last as the sum of a chemical component and an electrical one:
AGV = AG?%hem - ^ φ , (13-53)
which allows that only some of this potential difference is effective; this fraction is called the transfer coefficient a. The current part due to the forward reaction is then
* =
[ t H = # ^ M ^ ) ]
<p). < • « . ) • where current is in faradays per square centimeter per second. At equilibriumthere will be equal and opposite currents in the two directions, or k° = z'b° = i°, where i° is known as the exchange current, and φ has the value φΌ and corresponds to ê°. The n
Combination o f Eqs . (13-54 ) an d (13-55 ) give s
Η = ι° exp K i FRT ψ 0) = i° exp -^J- , ( 1 3_5 6)
where η is the overvoltage. Equation (13-56) may be written in the form In / = In i°
RT
which rearranges to the Tafel equation.
In a stricter analysis, an equation analogous to Eq. (13-54) is written for ή , ,
N e g a t i v e 0 — P o s i t i v e
F I G . 1 3 - 1 4 . A general current versus voltage diagram.
with i = it — ih ; the effect is to alter the coefficient of η in Eq. (13-56) to give
As the equation suggests, the situation is symmetric; η may be positive or negative, so that the electrode reaction is driven either forward or backward. Figure 13-14 shows the general behavior of Eq. (13-57).
This treatment, although somewhat sketchy, is designed to indicate how the detailed study of overvoltage effects can lead to information about the intrinsic rate of the electrode reaction, through i°, and the energy of the reaction barrier.
As an example, although the mechanism for the reduction of H+ ion at a metal electrode is still not fully elucidated, the evidence suggests that the rate-determining step is probably the reaction of H30+ with the metal surface to give adsorbed hydrogen atoms and water:
β . Concentration Polarization
The preceding analysis was based on the assumption that the rate-limiting step of the electrode process was some activated chemical reaction at the electrode surface. Another process must become rate-controlling at sufficiently high current densities, namely the rate of diffusion of reactant to and of product away from the electrode surface. If the reaction is one of the deposition of a metal from solution, only the rate of diffusion of the metal ion to the electrode is to be considered.
(13-57)
H30 + + M + e" = M - H + H20 .
SPECIAL TOPICS, SECTION 2 533
Recalling Eq. (2-65) and the discussion of Section 10-7B, we can write
J = "5) 8 ( 1 3
where J is now in moles per square centimeter per second. If there is an excess of inert electrolyte in the solution so that the ion being reduced carries little of the current, the potential term in Eq. (12-121) will not be important and Q) will be a constant, equal to the ion diffusion coefficient as given by Eq. (12-119). Further, a reasonable approximation to the physical situation is that the stirring conditions leave a thin stagnant layer of thickness δ within which C varies linearly from C, the bulk solution concentration, to Cs, the concentration at the electrode surface.
We expect the linear variation of C with χ since in a steady-state condition / is constant, and since S is constant, so must be dC/dx. Equation (13-58) then becomes
/ a = i / =
_
fW
zÇ )
(13-59)where id is the diffusion current density in faradays per square centimeter per second and ζ is the valence number of the metal ion being discharged. Under the condi
tions of the preceding subsection it is assumed that @/8 is large enough that Cs = C, and the observed current is determined by the rate of the chemical reduction reaction at the electrode, za . That is, under steady-state conditions there is no accumulation of material in the diffusion zone and ia = i& · As the over
voltage is increased, /a increases by Eq. (13-57), and Cs drops enough for zd to match the increase in fa · The maximum possible diffusion rate /a(max) is reached when Cs has dropped to zero. Further increase in the overvoltage cannot increase the current any further, and the plot of i versus V must level off at z'd(max) , or at
z@C
*d(max) = g · (13-60) Thus the effect of concentration polarization is to give a maximum current which
is proportional to the concentration of the ion being discharged.
If a mixture of potentially reducible ions is present, again with a supporting electrolyte, that is, an excess of inert electrolyte, present, then as the applied reducing potential is increased the first metal ion begins to deposit, reaches its i'max, and the second metal ion begins to deposit at some higher potential, and reaches its /m a x, so that the plot of / versus V is as shown in Fig. 13-15. The characteristic deposition potential for each ion could be identified as in Fig. 13-9 but when one is operating in the diffusion-controlled region Cs and hence <frev is changing with V so that one obtains a sinusoidal rather than a linear / versus V plot. Consequently, it is much more accurate to pick the potential at the half-way point of the step, or the half-wave potential V1/2, as the characteristic one.
Since the overvoltage η is given by
RT Λ Cs
^ Ι η ^ τ - , (13-61)
FIG. 13-15. Current versus voltage in the region of concentration polarization.
where ohmic polarization is neglected, as are activity coefficient terms, we can solve for Cs/C. Also, combination of Eqs. (13-59) and (13-60) gives
*d — Jd(max) The result is
- = 1 - exp - — { - .
*d(max) Kl
(13-62)
(13-63) Equation (13-63) gives a symmetric curve such as shown in Fig. 13-15, for which
V1/2 occurs at the inflection point.
C. Polarography
The preceding analysis suggests that we could analyze for metal ions in solution by obtaining experimental i versus V curves, each limiting diffusion current giving the concentration of a particular ion, the ion being identified by its V1/2. This procedure is indeed used. To maintain δ more constant than is possible by stirring
Hg anode
/ Galvanometer
II-i — v w w w v w w v —1
•
FIG. 13-16. Schematic diagram of a polarographic cell.
SPECIAL TOPICS, SECTION 2 535
the solution, a common practice is to use a rotating electrode. There is still a problem with accumulation of reduction products, and a very ingenious alternative procedure was devised by J. Heyrovsky in 1922.
The basic experimental features of the Heyrovsky polarograph are shown in Fig. 13-16. The cathode is a mercury drop that is steadily growing at the tip of a capillary immersed in the electrolyte solution, and the anode is a large pool of mercury. By having a tiny drop as the anode surface, concentration polarization effects can be made to develop at relatively small currents; the large area of the anode pool of mercury essentially eliminates concentration polarization at its surface. As each drop grows and falls the anode surface is steadily kept fresh and reproducible, and with a supporting electrolyte the same general analysis applies as for a stationary electrode.
A typical polarogram is shown in Fig. 13-17. The oscillations are a result of the successive appearance of new drops at intervals of a few seconds; currents may be only microamperes in order of magnitude. Each step is called a polarographic wave and is characterized by its half-wave potential V1/2 and diffusion current
i'd(max) . The detailed algebraic analysis is complicated by the situation of an expanding cathode surface, and an equation derived by D. Ilkovic in 1938 will be given without the derivation:
where ζ is the valence number of the ion, w is the flow of the mercury in grams per second, / is the drop time in seconds, and C is the concentration in moles per liter. Equations have also been derived for the half-wave potential, but this remains essentially an empirically determined quantity for each species.
Polarography is capable of measuring ion concentrations as low as Ι Ο-4 M and is a rapid as well as a sensitive analytical tool. The physical chemist uses it to study the chemistry of reduction (or of oxidation) processes. One can determine from the diffusion current whether the reduction occurs as a one- or as a
two-Id(max) = 70.82ZVK§0(1^)(!C), (13-64)
1.0
<
a. 0.5
0
0.4 0 - 0 . 4 - 0 . 8 - 1 . 2 -1 . 6
V (relative to standard calomel electrode) F I G . 1 3 - 1 7 . A typical polarogram.
250 /aamps d i v-1
/ c
— —
0.125 V d i v1
1 1 I 1
(+)
^ \ /a
1 1
ο ( - )
v = 0.1 V s e c - i
F I G . 1 3 - 1 8 . Typical cyclic voltammogram for a reversible redox couple. [From D. T. Sawyer and J. L. Roberts, Jr., "Experimental Electrochemistry for Chemists:' Wiley (Interscience), New
York, 1974.]
GENERAL REFERENCES
ADAMSON, A . W. (1976). "The Physical Chemistry of Surfaces," 3rd ed. Wiley (Interscience), New York.
DANIELS, F., MATHEWS, J. H., WILLIAMS, J. W . , BENDER, P., AND ALBERTY, R. A. (1956). "Experi
mental Physical Chemistry," 5th ed. McGraw-Hill, New York.
DOUGLAS, Β . E., AND MCDANIEL, D . H. (1965). "Concepts and Models of Inorganic Chemistry."
Ginn (Blaisdell), Boston, Massachusetts.
KORTUM, G. (1965). "Treatise on Electrochemistry." Elsevier, Amsterdam.
electron step. If complexing ions are added to the solution, both the reduction potentials and the formation constants of complex ions can be found. With a commutator to interrupt the applied voltage, chemical rate processes can be followed.
A modern variant of polarography uses a fixed (not a dropping mercury) electrode and current is recorded as the voltage is changed at a fixed rate. In cyclic voltammetry the voltage sweep is first in one direction and then in the other. Sweep rates of up to 100 V s e c-1 may be used. An idealized voltammogram is shown in Fig. 13-18. On the forward sweep (upper curve) reduction products are formed and on the backward sweep (lower curve) these are oxidized. About the same informa
tion is obtained as in ordinary polarography but, in addition, much can be learned about the nature and the reaction kinetics of the reduction product.
EXERCISES 537 LATIMER, W. M. (1952). "The Oxidation States of the Elements and Their Potentials in Aqueous
Solution," 2nd ed. Prentice-Hall, Englewood Cliffs, New Jersey.
LEWIS, G. N., AND RANDALL, M. (1961). 'Thermodynamics," 2nd ed. (revised by K. S. Pitzer and L . Brewer). McGraw-Hill, New York.
CITED REFERENCES
BATES, R. G. (1954). "Electrometric pH Determinations." Wiley, New York.
DOLE, M. (1941). 'The Glass Electrode." Wiley, New York.
HARNED, H. S., AND EHLERS, R. W. (1932). / . Amer. Chem. Soc. 5 4 , 1350.
LATIMER, W. M. (1952). "The Oxidation States of the Elements and Their Potentials in Aqueous Solution," 2nd ed. Prentice-Hall, Englewood Cliffs, New Jersey.
EXERCISES
Activity coefficient effects are to be neglected in Exercises and Problems unless specifically noted otherwise. Assume 25°C unless otherwise specified. Take as exact numbers given to one significant figure.
1 3 - 1 The emf of the cell Pb/PbCl2(5)/KCl(m)/Hg2Cl2(s)/Hg is 0.5357 V at 25°C and increases with temperature by 1.45 χ 1 0 "4 V Κ"1. Write the electrode and overall cell reactions and calculate AG\ AH0 JS°, and qTev .
Ans. Pb + Hg2Cl2 = PbCl2 + 2Hg(/), AG0 = - 1 0 3 . 4 kJ, AS0 = 28.0 J K - \ AH0 = - 9 5 . 1 kJ, qvev = 8.34 kJ.
1 3 - 2 The emf for the cell Cd/0.1 m C d ( N 03)2, 0.01 m AgN03/Ag is - 0 . 4 4 6 V at 25°C.
Calculate <?° for this cell.
Ans. - 0 . 3 5 7 V.
1 3 - 3 Calculate «?2 98 for the cell Ag/AgI(j)/0.1 m HI/H2(1 atm)/Pt.
Ans. 0.038 V.
1 3 - 4 Calculate <f° and Κ at 25°C for the reaction Tl + A g+ = T1+ + Ag.
Ans. £° = 1.135 V, Κ = 1.56 x 1 01 9. 1 3 - 5 Calculate Κ for the reaction \Cd + T1+ = ^Cd2+ + Tl and the ratio (Cd2 +)/(Tl+) if
excess Cd is added to a solution which is 0.1 m in Tl+.
Ans. K= 13.4 M ~1 / 2; 2.76.
1 3 - 6 Calculate «f<>98 for the half-cell reaction Tl = T l3+ + 3e~.
Ans. - 0 . 7 2 1 V.
13-7 Obtain «?<>98 for the half-cell Mg + C2Ol~ - Mg(C204) + 2e~ given that KBp = 9.0 χ 10" 5
for Mg(C204).
Ans. 2.49 V.
1 3 - 8 What fraction of Ag(CN)2- is dissociated into Ag+ in 1 χ IO-3 M CN~ if *J98 = 0.289 V for the half-cell Ag + 2 CN" = Ag(CN)2- + e"? (Neglect the hydrolysis of CN~.)
Ans. 4.0 χ 1 0 - " ( ! ) .
13-9 Calculate <^298 for the following concentration cells: (a) Na(Hg, xNa = 0.1)/NaCl(m = 0.1)/Na(Hg, jcNa - 0.01); (b) Pt/H2(/> = 0.1 atm)/HCl(m = 0.1)/H2CP - 0.01 atm)/Pt;
(c) Ag/AgBr(5)/HBr(A7z = 0.1)/H2(1 atm)/Pt—Pt/H2(l atm)/HBr(m = 0.01)/AgBr(j)/Ag.
Ans. (a) 0.0592 V, (b) 0.0296 V, (c) 0.1183 V.
13-10 Calculate ^3 50 for the cell in Exercise 13-9(b).
Ans. 0.0348 V.
13-11 Calculate <f298 for the cell Ag/AgCIO)/0.1 m HC1/H2(1 atm)/Pt, using activity coefficient data from Chapter 12.
Ans. - 0 . 3 5 2 V.
13-12 A pH meter gives a reading of 200 mV when measuring a solution whose hydrogen ion activity is considered to be 10~4; what is the hydrogen ion activity of a solution for which the meter reads 100mV? (Assume 25°C.)
13-2 The cell Cu/CuCl(i)/KCl(0.100/w)/Cla(l atm)/C has a potential of 1.234 V at 25°C.
(a) Calculate the solubility product for CuCl in water at 25°C.
(b) What is the concentration of cuprous ion at the anode of this cell?
13-3 Given that *° = 0.152 for Ag + I" = Agi + e~ at 25°C and is - 0 . 8 0 0 for Ag = Ag+ + e~, calculate the solubility product for Agi.
13-4 The voltage of the cell Ag/Ag2S04(.y)/saturated solution of A g2S 04 and H g2S 04/ H g2S 04/ Hg is 0.140 V at 25°C and its temperature coefficient is 0.00015 V X "1.
(a) Give the cell reaction.
(b) Calculate the free energy change for the cell reaction.
(c) Calculate the enthalpy change for the cell reaction.
(d) Calculate the entropy change for the cell reaction.
(e) Does the cell absorb or emit heat as the cell reaction occurs ? Calculate the number of calories per mole of cell reaction.
(f) One mole each of Hg, Ag2S04(j), Hg2S04(.y), and some saturated solution of the two salts are mixed. What solid phases finally will be present and in what amounts?
13-5 & for the cell Ag/AgCl/ZnCl2(m)/Zn is - 1 . 2 4 0 V at 25°C and - 1 . 2 6 0 V at 35°C, if m — 1 x 1 0 ~3m . Write the cell reaction and calculate AG, AH, and AS for this reaction at 30°C.
13-6 Write the cell diagram for a cell whose # could be used to determine the solubility product for Ag2S04(.s). What additional information would be needed besides the measured emf?
13-7 Given the cell
Ag/AgNO3(0.00100 m)////ÀgNOa(0.00100 m), KCN (0.004 m)/Ag,
PROBLEMS 53 9 whose em f i s - 0 . 7 6 7 V a t 25°C :
(a) Writ e th e electrod e an d ne t cel l reactions .
(b) Calculat e th e equilibriu m constan t fo r A g+- f 2 C N- = Ag(CN)2~ (neglec t activit y coefficient effects) . Th e doubl e dashe d diagona l denote s a sal t bridg e whic h (i t i s hoped) make s th e junctio n potentia l negligible .
13-8 A t 25°C , S fo r th e cel l Pb/PbSO4(5)/H2SO4(0.00200 m)/H2(l atm)/P t i s 0.12 5 V . Calculat e the solubilit y produc t o f P b S 04 a t 25°C .
13-9 Calculat e th e percentag e o f mercur y i n th e mercuri c stat e i n a solutio n o f mercuri c nitrate tha t i s i n equilibriu m wit h liqui d mercury .
13-10 Calculat e th e solubilit y produc t o f ferrou s hydroxid e fro m th e cel l Fe(s)/Fe(OH)2(.y)/
Ba(OH)2(0.05 m)/HgO(.s)/Hg . <T2 98 = 0.97 3 V .
13-11 Molecula r weight s ar e give n i n k g i n th e S I system ; thus , th e molecula r weigh t o f 0 2
is 0.03 2 kg . I n goin g fro m th e cg s t o th e S I system , stat e whethe r th e numerica l valu e of eac h o f th e followin g i s changed , an d i f i t is , calculat e th e ne w value , (a ) Avogadro' s number, (b ) Faraday' s number , (c ) £ ° fo r th e cel l Pt/H2(l atm)/HCl/AgCl/Ag . (d ) AG0
for th e cel l reactio n i n (c) . (e ) ê a t 25° C fo r th e cel l Pt/H2(l atm)/HCl(0.00 1 m)/AgCl/Ag . What woul d you r answer s b e ha d th e molecula r weigh t o f 0 2 bee n define d a s 3 2 k g mol e ~1
(and m ha d bee n k g molecula r weight s pe r 100 0 k g o f solvent) ?
13-12 I t i s desire d t o separat e C d2+ fro m P b2+ b y electrodepositio n fro m a solutio n whic h i s 0.1 m i n eac h io n an d a t pH 2 . Calculat e th e sequenc e i n whic h C d an d P b metal s an d H2
are produce d an d th e concentration s o f th e variou s specie s i n solutio n whe n a ne w stag e of electrolysi s occurs . Includ e overvoltag e effect s i n considerin g H2 evolution .
13-13 R . Og g foun d ê = -0.02 9 V a t 25° C fo r th e cel l
What i s th e formul a fo r mercurou s ion ? Sho w ho w you r conclusio n follow s fro m thi s information.
The actua l dissociatio n constan t fo r H g2+ = 2Hg+ i s no t known , bu t suppos e tha t S for thi s cel l i s foun d t o b e —0.05 9 V whe n th e highe r an d lowe r mercurou s nitrat e nor -malities ar e 2 χ IO-4 and 3 χ IO- 6, respectively. Calculate KdiBB for Hg2,* from this data.
The actua l dissociatio n constan t fo r H g2+ = 2Hg+ i s no t known , bu t suppos e tha t S for thi s cel l i s foun d t o b e —0.05 9 V whe n th e highe r an d lowe r mercurou s nitrat e nor -malities ar e 2 χ IO-4 and 3 χ IO- 6, respectively. Calculate KdiBB for Hg2,* from this data.