J. Phys. Chem. 1986, 90, 6392-6400
6392
in the interaction. This portion is formed by the cycles including the OH-0 intramolecular bonds together with the rings B and C (according to the model of Figure 4 based on the X-ray crystal structure of D N R with a DNA fragment d(CpGpTpApCpG)%). The slight differences in the relative intensity decrement of some modes for DCM-DNA with respect to ADM-DNA should be explained in terms of the different positions of the two OH groups, 4 and 6 in DCM and 6 and 11 in ADM, giving rise to a different interaction. In fact by the model for the intercalation proposed in Figure 4, it can be seen that for ADM the two C-0 bonds are pretty equivalent, giving rise to a simple change of the excited-state equilibrium position, while the two OH bonds are surrounded by different groups. In this latter case the interaction must provoke also a change of the normal coordinate between the ground and excited state, causing a further intensity reduction. In fact the consequence of the inequivalent interaction of similar bonds should be the disruption of their phase-coupled local oscillators, giving rise to an excited-state vibration that mimics a mode of B, species therefore to a weaker band than the others. In DCM the OH at position 11 is replaced by a H and moved to position 4 instead of OCH3. With the assumption that for DCM the chromophore intercalates in the same geometry, both the OH bonds show a similar surrounding, while the C-O and C=O bonds become inequivalent with a further intensity reduction of the corresponding bands. In conclusion only one C=O-.HO group of DCM is involved in the interaction with DNA, giving rise to a less strong complex with respect to ADM where both C=O-HO groups are involved.
Conclusions Surface-enhanced resonance Raman spectroscopy is a powerful method to study the interaction of highly fluorescent antitumor anthracyclines with DNA. In fact the present spectral observations establish that even if the chromophore framework is slightly perturbed by adsorption onto silver particles, this does not affect its interaction with DNA. W e were able to obtain highly detailed spectra at very low concentrations without any fluorescence background for ADM, (24) Quigley, G. J.; Wang, A. H. J.; Ughetto, G.; Van der Marel, G.; Van Boom, J. H.; Rich, A. Proc. Natl. Acad. Sci. U.S.A. 1980, 77, 7204-7208.
DCM, and their model chromophores. The combined analysis of these spectra allowed us to perform a reliable and detailed vibrational assignment of the chromophoric portion, which is decisive for the interpretation of the spectral changes observed upon complexation with DNA. The intensity changes detected between the pure drugs and the complexes are due to variations of both normal coordinates and their equilibrium positions on going from the ground to the excited states. These changes on specific groups reflect the specific excited-state interactions of rings B and C together with their hydroxyl and carbonyl substitutents with the base pairs of DNA. This interpretation is fully consistent with the intercalation model found by the X-ray analysis of D N R with a DNA fragment d(CPGPTPAPCPG). The different intensity changes observed between ADM-DNA and DCM-DNA complexes are interpreted in terms of the same intercalative model, with a weaker interaction strength for DCM with respect to ADM. This confirms the observations of the UV-visible and fluorescence spectra, which predicted a complete complexation with a lower molar ratio DCM/DNA (1:20)23 compared to that of DNR/DNA (1:5).25 In conclusion the SERS method and in particular the SERRS method furnish powerful indicators to study the geometry of anthracycline chromophores and their intercalative complexes. On the basis of the spectral intensity differences observed between ADM and its complex with DNA, it is shown that the main portion of the chromophore interested in the interaction is represented by rings B and C with their hydroxyl and carbonyl substituents. As a consequence the slight intensity change observed for DCM depends on the lack of the O H at position 11 with respect to ADM, which causes a less strong interaction with DNA. Acknowledgment. This work was supported by the Italian Consiglio Nazionale delle Ricerche and Minister0 della Pubblica Istruzione. We thank Prof. M. P. Marzocchi for a helpful discussion and Dr. Valentini (Farmitalia, Milan) for the generous gift of samples of the drugs. Registry No. ADM, 23214-92-8; DCM, 81382-07-2; 1,8(0H),AQ, 117-10-2; 1,4(0H),AQ, 81-64-1. ( 2 5 ) Barthelemy-Clavey,V.; Maurizot, J.-C.; Sicard, P. J. Biochimie 1973, 55, 859-868.
Electrochemical Behavior of Dispersions of Spherlcal Ultramicroelectrodes Martin Fleischmann,t Jamal Ghoroghchian,:" Debra Rolison,! and Stanley Pons** Department of Chemistry, University of Southampton, Southampton, Hants, SO9 5NH England, Department of Chemistry, University of Utah, Salt Lake City, Utah 841 12, and Surface Chemistry Branch, Naval Research Laboratory, Washington, D.C. 20375-5000 (Received: March 7, 1986; In Final Form: June 3, 1986) It is shown that it is possible to carry out electrochemical reactions in poorly conducting and nonconducting media by means of bipolar electrolyses with dispersions. Polarization equations are predicted for highly simplified models based on the concept of the mixture potential, the surface reactions being assumed to be rate determining. Results for hydrogen evolution/oxidation and oxygen evolution/reduction show that the interpretation of polarization curves at high field strengths will have to take into account the effects of diffusion. The results also show that it should be possible to investigate, monitor, and modify heterogeneous catalyses of reactions in the liquid phase by means of the faradaic currents induced by the electric fields.
Introduction
The high rates of stationary mass transfer in the developing quasi-spherical diffusion fields allow the measurement of the The construction and behavior of microdisk,'-I9 m i c r ~ r i n g , ' ~ . ~ ~ kinetics of fast electrode reaction^^^'^^'^^^'^^^ and of fast reactions and microsphere'g921-23 have been discussed in solution coupled to electrode reactions14,15 under steady-state *Towhom correspondenceshould be addressed.
University of Southampton. 'University of Utah. Naval Research Laboratory. 'I Present address: Chemistry Department, West Chester University, West Chester, PA 19383.
0022-3654/86/2090-6392$01.50/0
(1) Soos, Z. G.; Lingane, P. J. J . Phys. Chem. 1964, 68, 3821. (2) Ponchon, J.-L.; Cespuglio, K.; Gunon, F.; Jouvet, M.; Pujol, J.-F. Anal. Chem. 1979,51, 1483. (3) Dayton, M. A.; Ewing, A. G.; Wightman, R. M. Anal. Chem. 1980,
52, 2392.
0 1986 American Chemical Society
Electrochemical Behavior of Dispersions conditions. The spherical potential field decreases charging times so that fast transient techniques can be readily applied,'3 the low ohmic losses in solution have also allowed measurements to be made in solutions containing only low (or effectively zero) concentrations of support electrolyte as well as in glasses at low t e m p e r a t u r e ~ l and ~ . ~ ~even in the vapor phase.zs The modest scale-up of reactions at microelectrodes requires the use of spherical structures such as of embedded reticulated foamsz6 or fibemZ7 Exploitation of the special advantages of microelectrodes for synthesis (e.g., the ease of workup and the extension of the solvent range in the absence of support electrolyte'*) however requires the use of three-dimensional electrodes. Bipolar electrolyses on dispersions of spherical particles have been proposed, and the behavior of such electrodes in the presence of a single redox couple has been analyzed for a number of limiting conditions;%these systems represent an extension of the concepts used in bipolar fluidized bed electrodesz9to the area of microelectrodes. In this paper we extend the analysis to the behavior of the dispersions in the presence of two reversible or two irreversible electrode reactions using the simplest model, Figure lA,Z* and compare this behavior to that observed at the mixture potentials in systems containing relatively high concentrations of e l e c t r ~ l y t e ; ~colloidal ~ ~ ~ ' systems at the relevant mixture potentials have been extensively investigated in recent years in the context It should of the catalysis of the photodecomposition of (4) Kakihama, M.; Ikeuchi, H.; Sato, G. P.; Tokuda, K. J. Electroanal. Chem. 1980, 108, 381. ( 5 ) Wightman, R. M. Anal. Chem. 1981, 53, 1125A. (6) Oldham, K. B. J . Electroanal. Chem. 1981, 122, 1. (7) Aoki, K.; Osteryoung, J. J . Electroanal. Chem. 1981, 122, 19. (8) Heinze, J. J . Electround. Chem. 1981, 124, 73. (9) Aoki, K.; Osteryoung, J. J . Electroanal. Chem. 1981, 125, 315. (10) Scharifker, B.; Hills, G. J. J . Electroanal. Chem. 1981, 130, 81. (11) Shoup, D.; Szabo, A. J . Electroanal. Chem. 1982, 140, 237. (12) Hepel, T.; Plot, W.; Osteryoung, J. J . Phys. Chem. 1983,87, 1278. (13) Howell, J. 0.;Wightman, R. M. Anal. Chem. 1984, 56, 524. (14) Fleischmann, M.; Laserre, F.; Robinson, J.; Swan, D. J . Electroanal. Chem. 1984, 177, 97. (15) Fleischmann, M.; Lasserre, F.; Robinson, J. J . Electroanal. Chem. 1984, 177, 115. (16) Bond, A. M.; Fleischmann, M.; Robinson, J. J . Electroanal. Chem. 1984, 180, 257. (17) Bond, A. M.; Fleischmann, M.; Robinson, J. J . Electroanal. Chem. 1984, 168, 299. (18) Cassidy, J.; Khoo, S.B.; Pons,S.;Fleischmann, M. J . Phys. Chem. 1985, 89, 3933. (19) Bond, A. M.; Fleischmann, M.; Robinson, J. Extended Abstract, 165th Meeting of the Electrochemical Society, New York; Electrochemical Society: New York, 1984; p 523. Bond, A. M.; Fleischmann, M.; Khoo, S. B.; Pons, S.;Robinson, J., unpublished results. (20) Fleischmann, M.; Bandyopadhyay, S.;Pons, S . J . Phys. Chem. 1985, 89. 5537. (21) Bindra, P.; Brown, A. P.; Fleischmann, M.; Pletcher, D. J. Electroanal. Chem. 1975, 58, 31. (22) Bindra, P.; Brown, A. P.; Fleischmann, M.; Pletcher, D. J . Electroonal. Chem. 1975, 58, 39. (23) Bond, A. M.; Fleischmann, M.; Robinson, J. J. Electroanal. Chem., in press. (24) Bond, A. M., et al., personal communication. (25) Ghoroghchian, J.; Sarfarazi, F.; Dibble, T.; Cassidy, J.; Smith, J. J.; Russell, A,; Fleischmann, M.; Pons, S.,unpublished results. (26) Seeszynski, N.; Osteryoung, J.; Carter, M. Anal. Chem. 1984, 56, 130. (27) Shoup, D.; Szabo, A. J. Electroanal. Chem. 1984, 160, 19. (28) Fleischmann, M.; Ghoroghchian, J.; Pons, S.J . Phys. Chem. 1985, 89, 5530. (29) (a) Fleischmann, M.; Goodridge, F.; King, C. J. H. Br. Patent Application 16765, 1974. (b) Goodridge, F.; King, C. J. H.; Wright, A. B. Electrochim. Acta 1977, 22, 1087. (30) Spiro, M.; Freund, P. L. J. Chem. Soc., Faraday Trans. 1 1983,79,
The Journal of Physical Chemistry, Vol. 90, No. 23, 1986 6393
,
t VI
m EtIm
__
I
GyltI 4 1 e l
.VI F E R MClRm
Figure 1. (A, left) Representative particle and the coordinate system for a single redox couple for a < 0.5. (B, right) Representative particle and the distribution of cathodic and anodic reactions for two redox reactions.
be noted that the size range of the particles which will be most cm) frequently used in dispersion electrolyses ( l w cm C a C is intermediate to that of the colloidal systems (typically 10-6 cm) and fluidized bed electrodes (typically cm). The behavior in the presence of two redox systems is compared to that observed for a single redox reaction and to experimental data for two electrode reactions on pt dispersions: hydrogen evolution/oxidation in conductivity water and oxygen evolution/reduction in dilute potassium hydroxide solutions. The behavior of other systems and applications to synthesis4zand the use of supported metal catalyst particles43 will be discussed elsewhere. Mathematical Description of Dispersion Electrolysis Control by a Single Redox Couple. Dispersion electrolyses will usually be carried out using low or zero concentration of support electrolyte. For the generation of a cation from an uncharged substrate O++e+R
(A)
or the converse case of the generation of an anion O+e==RR-
(B)
the charge density generated in the reactions will be compensated by counterions either from deliberately added low concentrations of support electrolyte or from impurities or ions generated by autoionization of the solvent. For the likely range of particle radii, a, cm C a C cm, the steady-state mass-transfer coefficient, k,,
k,
N
D/a
(1)
is so high that the surface reaction is likely to be rate controlling at low and intermediate overpotentials for the majority of reactions. In the simplest possible model we assume2*that the reactions on the anodic and cathodic parts of a representative spherical particle do not significantly perturb the equipotentials (i.e., the surface of a particle experiences the mean field). The overpotential v(V) at the position (a,O) is then given by
where V,,,, (V) is the applied voltage, L (cm) is the distance between the feeder electrodes, and X defines the position at which t) = 0 (see Figure 1A for the coordinate system). For a reaction following the simple polarization equation
1649.
(31) Freund, P. L.; Spiro, M. J . Phys. Chem. 1985, 89, 1074. (32) Kiwi, J.; Gratzel, M. J . Am. Chem. Soc. 1979, IOZ,7214. (33) Gratzel, M. Faraday Discuss. Chem. SOC.1980, 70, 359. (34) Keller, P.; Mooadpour, A. J . Am. Chem. SOC.1980, 102, 7193. (35) Meisel, D.; Mulac, W. A.; Matheson, M. S . J . Phys. Chem. 1981,85, 179. (36) Johansen, 0.;Launikouis, A.; Loder, J. W.; Mau, A. S.-H.; Sasse, W. H. F.; Swift, J. D.; Wells, D. Aust. J . Chem. 1981, 34, 981. (37) Miller, D. S.;Bard, A. J.; McLendon, G.; Ferguson, J. J . Am. Chem. Soc. 1981, 103, 5336. (38) Miller, D. S.; McLendon, G. J. Am. Chem. Soc. 1981, 103, 6791.
~
~~
~~
(39) Albery, W. J.; BartGtt, P. N.; McMahon, A. J. In Photogeneration of Hydrogen; Harriman, A., West, M. A., Eds.; Academic: London, 1982. (40) Matheson, M. S.; Lee, P. C.;Meisel, D.; Pelizzctti, E. J. Phys. Chem. 1983, 87, 394. (41) Albery, W. J.; Bartlett, P. N.; McMahon, A. J. J. Electrwnal. Chem. 198.5, 182, 7.(42) Ghoroghchian, J.; Pons, S.; Fleischmann, M. unpublished results. 143) Rolison. D. R.: Nowak, R. J.; Pons, S.; Fleischmann, M.; Ghoroghchian, J. unpublished results.
The Journal of Physical Chemistry, Vol. 90, No. 23, 1986
6394
10
ID
J 1
im ID
in
0
.i
Fleischmann et al.
. I
in im 110
.I. im
r
-10
2
! I D
m m 4
m m m o
I M
11
u
u
r
u
U
4d
u
+
,
M
,
,
,
u
11
,
,
,
u
1 M
I
Y
Figure 2. (A, left) Effect of Q on the predicted normalized polarization curve for the dispersion for a single redox couple. (B, middle) Effect of the difference in redox potential between two redox couples on the predicted normalized polarization plots for dispersion electrolysis and for iol = iO2and a1 = a2 = 0.5 and for the values AEF/4RT indicated. (C, right) Effect of the difference in redox potential and of the applied field between two redox couples (expressed as y) on the predicted net normalized current density IN/3ioVforone of the couplex and on the ratio IN/Ipfor io, = iO2and al = a2 = 0.5 and for the values AEF/4RT indicated.
we obtain the total current into a single particle
-IP- 3i0u
1 a(.
- 1)y
[(a- 1) exp(ay cos X) exp(ay) -
a exp((a
-
2 r i a TLa FVa,,P(a - 1)
[(a- 1) exp(ay cos A) exp(ay) a exp((a
- 1)y cos A) exp((a - 1)y) + 11 (9)
For the special case a = 0.5 and h = x / 2 , we obtain
-
’[
- l)y cos A) exp((a - 1)y) + 11
3i0v = y cosh
(4) where
(z)
-11
while the special case a = 1 gives exp(y cos X) =
and the division between the net cathodic and anodic areas is defined by exp(y cos A) =
sinh ((a - 1)y) a-1 sinh (ay)
(5)
(6)
In the dispersion each bipolar particle presents a shunt resistance to the overall solution resistance and we obtain the overall polarization curve
r r
- a exp((a - 1)y cos
A) exp((a
- 1)y) + 11 (7)
where IT (A cm-*) is the cross-sectional current density through the dispersion, N is the number of particles per unit volume, and p s ( Q cm) is the solution resistivity. The nonfaradaic bypass can be determined from the background current current Vappl/psL in the absence of the dispersion. We can therefore cast (7)in the dimensionless form
Ipy 3 iou = (IT - $)L 3i0V =a(.
1
- 1)
[(a- 1) exp(ay cos A) exp(ay) a exp((a
- 1)y cos X)
exp((a - 1)y) +
11 (8)
where I p (A cm-*) is the cross-sectional faradaic current density through the particles and v (cm3 ~ m - ~is )the volume of the particles of the dispersion added per unit cell volume. Equation 8 compares dispersion electrolysis most directly with the conventional form, ( 3 ) , of the polarization equation. Direct comparisons of experimental data with the predicted forms of the polarization equations may be carried out by writing (8) (and subsequent equations) as
7 sinh y
and IP exp Y = -3i0u sinh y
(y ) 1+7
-
ln
(2) (12) sinh y
It should be noted that in view of the bipolarity of the particles the polarization curve for a given value of a is identical with that for (1 - a). In consequence, the dependence of the polarization curves on a is diminished compared to that observed for planar electrodes, Figure 2A. Effect of Particle Rotation. The formation and removal of surface layers (e.g., Pt-H or Pt-0) due to particle rotation will contribute further faradaic components to the current. The rotational frequency will be2*
where N (number mol-’) is Avogadro’s number and p (g cmW3) is the density. The frequency will lie in the range 10-2-104 revolutions s-! for the likely range of particle radii. If a monolayer charge nFQ is continuously formed and stripped by the particle rotation, then at such a limiting condition the current flow will be IR
zz 4ra2nFQv
(14)
where Q (mol cm-2) is the saturation monolayer coverage and n charges are transferred for each adsorbed species. The current flow in a slice of thickness 2a will be ~ ~ N =I 8?ra3nFQvN R This current can be compared to Z p by taking
(15)
Electrochemical Behavior of Dispersions
zm
odor Electrodes
2D
-
26l
-
240
-
2102m
-
i
1.80 1.4 im
Flow In
la tll Wl
1.00
P
P
-
030
am a m
1m.m
0.00
Y
hlnrwrI.n
Effects of Changes in the Polarization Equation Simple changes in the polarization equation such as the replacement of (3) by
[ (-y;F) -
- exp( ( ' - : y ) ]
i = io exp
/Vcm-'
(17)
of any externally applied electric field the particles will adopt a potential such that the net anodic current density of the couple having the more negative reversible potential is equal to the net cathodic current density of the couple with the more positive reversible potential (in this simple model we assume that there is no interaction between the two reactions as would be caused, for example, by competitive adsorption of reactants, intermediates, or products of the two component reactions or the reaction of adsorbed reactants or intermediates of the two redox couples)
)
io:[ exp( (1 - RT 4 7 1 m F -exp(
y)
lead to equations of the form 3ioU
4m.m
Figure 4. Polarization curves for hydrogen evolutionloxidation in conductivity water at low values of the applied field: 6.3 X lo8 Pt particles ~ m - diameter ~, 2.5 fim: (0) background current, + 10%H2 + 90% N2; ( 0 ) 100% H2 total pressure 1 atm.
Figure 3. (A, top) Cylindrical cell. The inner cylinder diameter was 2.2 cm, and the outer cylinder diameter was 3.2 cm. The platinum electrode cylinder height was 5.5 cm. (B, bottom) Plane Parallel capillary gap cell and the electrolysis system. The electrodes are 2.5 cm square, and the gap thickness is 0.5 mm.
-ZP- -
sw.m
1
- 1)y
O!(a
CY
io.[ exP( -"z'tzmF -exp(
[(a- 1) exp((any cos A) exp(any)
exp((a
- 1)ny cos A)
exp(a - 1)ny)
+ 11 (18)
sinh ((a- 1)nr) sinh (an?)
(CY: 1)
=
)]
(1 - RT a2)72mF
(20)
-
with X being defined by exp(ny cos A) =
=)I
where 7,"'- 72mare the overpotentials at the mixture potentials. Equation 20 can be solved with 72m
(19)
It follows that if the parameter ny is replaced by a parameter $, then the shapes of the derived polarization curves will be the same as those for a one-electron transfer reaction; if the currents are plotted against 7 , then the shapes of the predicted curves will be changed. However, in practice the effects on the polarization equation of a sequence of reaction steps will usually be more complicated than is indicated by simple changes such as from (3) to (17). Control by Two Redox Reactions. We discuss this case as the simplest possible example of a catalytic reaction and assume that homogeneous electron transfer is slow compared to the rates of the electrode reaction^.^' We develop the analysis by analogy to the concept of mixture potentials" which has been extensively applied to a range of electrochemical problems.45 In the absence (44) Wagner, C.; Traud, W. Z . Elektrochem. 1938,44, 391. (45) Spiro, M. In The Physical Chemistry ofSolutions; Fenby, D. V., Watson, I. D., Eds.; Massey University: New Zealand, 1983.
- 71m = +2,r
- &,r) = AE
(21)
to give the mixture potential; AE will be a negative quantity. The rate of the two redox reactions can then be derived from (20). It has been shown that a large number of paired redox reactions in conventional electrolyte solutions are catalyzed by noble metals which adopt this mixture potential.46 An extensive series of investigations has shown that the reactions may be controlled by slow electrode reaction^^'*^^ or else by diffusion with a Nernst equilibrium at the See also references cited in ref 30,31, and 47-49. An exact and general analysis of the problem in the presenc of an externally applied electric field requires the solution of the differential equations governing diffusion and migration to the particles, the reactions on the surface being boundary conditions for these solutions.28 To simplify the analyses we will assume here (46) (47) (48) 481. (49) 491.
Spiro, M.; Ravno, A. B. J. Chem. SOC.1965, 78. Spiro, M. J. Chem. SOC.,Faraday Trans. 1 1979, 75, 1507. Freund, P.L.;Spiro, M. J . Chem. Soc., Faraday Trans. I 1983, 79, Freund, P.L.; Spiro, M. J. Chem. SOC.,Faraday Trans. 1 1983, 79,
6396 The Journal of Physical Chemistry, Vol. 90, No. 23, 1986 that io, and iO2 are sufficiently small so that the potential of the surface follows the applied external field, Figure l A , and so that the effects of mass transfer can be neglected. We now define vIm and vzm as being the overpotentials at which there is no local net faradaic current so that we can write
Fleischmann et al. Equation 28 can be solved with (20) and ( 2 1 ) to give X as a function of the experimental conditions, and the polarization curve can then be derived from ( 2 5 ) . It has been shown above that the polarization curves for a single redox couple do not depend markedly on a in the middle range of this parameter. We therefore restrict attention here to the special case a1 = a2 = 0.5 and iol = iO2giving lr
12
X =2
=
Vappla
n2m
-7 (cos 0 + cos A)
and
~
=qlm+AE--
( V*ppla (cos 0 L
+ cos A)
(23)
The total external cathodic (or anodic) current is obtained by taking into account the contributions from both redox couples (see Figure 1B for the general case (2Vapp1a)/L> A E )
Z, = -2aa2Lr-'i1 sin 0 d0 + 2aa2Lr-'i2 sin 0 d0 ( 2 4 ) With some rearrangement we obtain
'[ (
5 = 3i0v y
cosh
4RT
for all y
+ I2) + cosh
(29)
(- -) AEF - Y 4RT 2
2 cosh
-
(z)]
(30)
We are also interested in the total net current for each redox couple e.g.,
ZI = 2 a a 2 L r i l sin 0 d0 =
I, =
exp((al - 1)y cos A) sinh ( ( a l - 1)y)
1
(31)
For the special case a 1= a2 = 0.5 and iol = iO2,we obtain
_ "3i0v y
The angle X is defined by the equality of the total net cathodic and anodic currents:
+ Lr-'i2sin 0 d0 =
-Lr-'ilsin 0 d0
S,:i il sin 0 d0 - LI+ sin 0 d0 (26) Le. L"ilsin 0 d0 =
s,"
i2 sin 0 d0
giving exp(a,y cos X) X
cos A) sinh ( ( a l - 1 ) y )
1
=
exp((a2 - fly cos A) sinh (a2- 1) - ( a z- 1 )
X
1
exp(a2y cos A) sinh ( a 2 y )
(28)
sinh
(-)
AEF 4R T
sinh
(i)
(32)
where ZN ( A cm-2) is the net cross-sectional current density for each of the redox couples. Zp/(3iou) is plotted against y in Figure 2 B (eq 30). It can be seen that Zp increases rapidly with AEFI4RT this increase is so marked that it is unlikely that the simple model illustrated in Figure 1 B will be adequate for large values of AE. The shapes of the polarization curves however are nearly independent of AEF/4RT so that it is only possible to distinguish the case of control by two redox processes from that by one redox reaction, Figure 2A, by virtue of the observation of an anomalously low value of io. The plots of ZN/3i$ against y, Figure 2C, show that it is possible to increase the net rate of each redox process by applying an electric field to the particles, and this could prove to be useful in certain examples of catalysis. It can be seen that the ratios I N / I p approach a limit with increasing values of AEF/4RT even for low values of y: the ratios are already independent of AE for AE as low as -0.25 V. Since Z, is measured experimentally it follows that the form of the I N - y relationship can be derived indirectly. It should therefore be possible to investigate and monitor catalytic reactions involving redox processes by measuring the polarization curves for the dispersed catalyst particles. It should be noted that such measurements can be made in solvents having effectively zero ionic conductivity (see further below) and, equally, on dispersions made of materials showing low electronic conductivity (e.g., semiconductor^).^^ Control by Two Irreversible Reactions. For sufficiently slow redox reactions or, more usually, in the case of the intervention of chemical reactions leading to the formation of electroinactive species, one or both of the component reactions must be regarded as being irreversible. There will be many special cases; here we restrict attention to two of these. (A) The cathodic reaction of redox couple 1 is balanced by the anodic reaction of redox couple 2. (B)The anodic reaction of redox couple 1 is balanced by the cathodic reaction of redox couple 2.
The Journal of Physical Chemistry, Vol. 90, No. 23, 1986 6397
Electrochemical Behavior of Dispersions
s)
Special Case A . At the mixture potential we have iol exp(
)
= iO2exp( (1 - RT az)7zmF
(33)
which can be solved together with (21). In the presence of an applied field, the overpotentials are given by (22) and (23) and the total cathodic and anodic currents on each particle are obtained by integrating over the range 0-T. We obtain I, =
I7 mF 7) exp(aly cos A) sinh ( q y )
4aa2iOl
1
--cy
-exp(
(34)
and determine X from (21) with
!?! exp(
7) exp(aly cos X) sinh (sly) = -a1 71mF
+
io2
az)(71m
RT exp((a2 - 1)y cos A) sinh ((1 - a&) (35) The special case al = a2 = 0.5 and iol = iO2gives X = ~ / and 2
I’ = 3i0u
y
exp(
AEF E )sinh
(i)
(36)
Special Case B. At the mixture potential we have
which is solved in conjunction with (21). In this case we obtain
I, =
(1 exp((a, - 1)y cos X) sinh ((1 - a J y ) (38)
with X being determined from (21) with
exp((al - 1)y cos. X) sinh ((1 - a l ) y ) = a2
exp(a2y cos X) sinh ( a 2 y ) (39)
RT
The special case al = a2 = 0.5 and iol = iO2gives X = a / 2 and
I’ = 3i0v
y
exp(
E ) (i) -AEF
sinh
(40)
Comparison of (36) with (40) shows that the rates are much higher for case B than for case A.
Experimental Section Two suitable cell designs which have been used in the present investigation are illustrated in Figure 3. The first had concentric platinum foil feeder electrodes with a 0.5-cm interelectrode gap. The microparticles were kept in suspension by sparging gas through the fine sintered disk at the base of the cell. The cell was used for measurements of the evolution/oxidation of hydrogen; hydrogen-nitrogen gas mixtures of the appropriate composition were therefore sparged into the solution, the total gas pressure in the cell being kept at 1 atm. The second cell, Figure 3B, had a plane parallel capillary gap configuration with an interelectrode gap of 0.5 mm. The feeder electrodes were again made of platinum foil (0.5 mm). Particles were kept in susp&sion in a small reservoir by bubbling gas, and the dispersion was pumped through the flow circuit by using a Monostate Ministaltic pump. Typical flow rates were 15 cm s-l
and were measured by a Gilmont Instruments flow meter. Controlled voltages were applied to the cells by using a Kepco regulated power supply Model H B d M . Current-potential curves were recorded point-by-point, a HI-TEK digital integrator/DVM being used to smooth out current fluctuations. Prior to each experiment the background current due tc) the solvent (water) or electrolyte (potassium hydroxide solution) in the absence of the dispersion was measured. An appropriate amount of platinum powder (Johnson-Matthey) was then added to give a particle concentration in the range 108-109 cm-3 (based on the tap density quoted by the manufacturer). The required amount was of the order 3 mg of Pt cm”; the average particle diameter was 2.5 pm; the geometry was assumed to be spherical. The polarization curves were then redetermined for the chosen conditions. Conductivity water was wepared by triple distillation in a glass still from “in-house” distilled water. The first stage contained 0.1 mM potassium permanganate and potassium hydroxide to raise the pH to -8.0; the second stage contained phosphoric acid. Potassium hydroxide solutions were made up using J.T. Baker Chemical Co. KOH. Hydrogen and nitrogen, or the gas mixtures used were U H P grade from IGP and were used as supplied. Glassware was cleaned by using a mixture of H N 0 3 / H 2 S 0 4and was then rinsed thoroughly by using triply distilled water.
Results and Discussion Hydrogen EvolutionlOxidation. Figure 4 illustrates the polarization curves obtained for hydrogen evolution/oxidation in conductivity water at low values of the applied field. The background current in the absence of the platinum particles shows that the water maintained a specific resistivity in excess of lo7 Q cm in the cells and ci, culation systems used in these experiments. Electrolyte impurities therefore remained at very low levels (the specific resistivity of pure water has been given as 1.69 X lo7 Q cm,50and the theoretical limit is 1.82 X lo7 0 cm). Addition of the platinum particles gave the expected increase in the faradaic current; this current was found to be proportional to the amount of particles added (expressed here as the volume of the dispersed phase per unit volume of liquid) and to the partial pressure of hydrogen sparged into the cell, Figure 4. Figure 5 parts A and B show that at low applied fields the experimental data fit closely to the special caselo (for a = 0.5) of the predicted equation (8) based on the simple polarization equations (3) or (1 7) (with n = 2). The data do not permit a clear distinction to be made between the kinetic forms although it appears that (17) is to be preferred as (3) consistently gave plots lying above the maximum permitted rate at higher values of the applied field. The simple first-order dependence of the polarization curves on the partial pressure of hydrogen as well as the fit of the curves to (10) shows that the behavior of the bipolar system is controlled by a single overall hydrogen evolution/oxidation reaction such as (A) and (B). Control of hydrogen evolution by 2H+ 2Hz0
+ 2e
+ 2e
H2
(A)
H2 + 2 0 H -
(B)
=i
(B) and oxidation by (A) would require a fit of the experimental data to (30) or to the special case (A), (36). In Figure 5 parts A and B the experimental data have been fitted to the predicted plots by using the apparent exchange current density, io, as the adjustable parameter. The required io is high, even allowing for the large real surface area of the particles (the ratio of the real to the apparent area lies in the range 5-10). This large exchange current density shows that the particles maintained a very high level of catalytic activity and that as in other investigations of the hydrogen evolution/ionization reaction we must assume that the surface reaction is close to reversibility since i, approaches the value predicted from the-mass-transfer coefficient, (50) Haller, W.; Duecker, H.C.J . Res. N a f . Bur. Stand., Sect. A 1960, 6th, 521.
Fleischmann et al.
6398 The Journal of Physical Chemistry, Vol. 90, No. 23, 1986 15
6.0
1
1.3
/
5.0
1.2 1.1
i.0
.-n>
4.0
03 b’
Od
ii
0.7
L L
\
z
-
\
P
LO
0.6 0’5
20
0.4 0.3
0.2
0.1 0.0 0.0
0.2
0.4
0.6
0.1
1.2
1.0
1.4
1.6
Y 1W lSo.0
11,m
140.0 10,m
150.0
9,w
120.0 110.0
-
I
Lmo
1m.o
>
...n
Boll
\
m.0
A L
70.0
.c
n
8.0
50.0 40.0 50.0 20.0 10.0 0.0 0.0
20
4.0
6.0
LO
x,
0.0
20
4.0
to
LO
v
10.0
120
14.0
16.0
11.0
Figure 5. Comparison of experimental polarization plots for hydrogen evolution/oxidation with the predicted forms of these plots (full lines for a =
0.5, a 0.25, or 0.75 and 0 or 1): 6.3 X lo8 Pt particles cm-), diameter 2.5 j” (A, top left) Low values of y; apparent exchange current density 120 mA cm-*. (B, top right) Low values of $; apparent exchange current density 55 mA cm-2. (C, bottom left) high values of y; apparent exchange current density 50 mA cm-*. (D, bottom right) high values of $; apparent exchange current density 22 mA
(1). The behavior of dispersion electrodes under conditions which include the effects of diffusionz8 will be discussed further elsewhere.51 It should be noted that in view of the high rate of the reaction, the data show no evidence for the deposition/removal of adsorbed hydrogen, the expected value of ( 16) being 10-2-10-1 at y = 1 with the same value of io as that used in fitting the experimental data to the predicted equations and a value s2 = lo-’ mol cm-* allowing for the high real area of the particles (see also section on oxygen evolution/reduction). ( 5 1) Ghoroghchian, J.; Pons, S.;Fleischmann, M., unpublished results.
At higher values of the applied field, Figure 5 parts C and D indicate at first sight that behavior according to the kinetic law (3) is to be preferred to that according to (17) but the fit obtained in Figure 5C must be regarded as being fortuitous. At these high rates the reaction on the particles will be diffusion limited by the anodic process as is indicated by Figure 5D.Indeed the continued fit to (10) implied by Figure 5C would require a kinetically controlled current density in excess of 500 mA cm-* on the apparent surface area of the particles if the current within each layer of thickness =2a is regarded as becoming wholly faradaic, i.e., if each layer is effectively screened by the adjacent layers of the
Electrochemical Behavior of Dispersions
The Journal of Physical Chemistry, Vol. 90, No. 23, 1986 6399
tm
0
D
7.m
I
I
am '
I
I'
0
d
I
'f,
I
D
I'
rm
-
I
im
I D
=*a zoo
I
1m*
om moo ma a m a m i m m ~ ~ m i ~ ~ m m ~ m m
M Q a0
2
4
8
10
12
V m-'
Figure 7. Polarization curves for oxygen evolution/reduction in 10" M KOH. Diameter of Pt particles 2.5 pm: 1.8 X lo* and 3.6 X lo8
particles cm-).
dispersion. Figure 5D requires that the diffusion-limited current should reach this value using the same assumption of the screening of adjacent layers of the dispersion. However, the diffusion-limited current will be maximally
FD ilim= 2 x 3 . 2 2 a ( C H J h l k = 50 mA cm-2
(41)
the factor 3.22 being due to generation of hydrogen on the cathodic areas adjacent to the anodic areas of the bipolar particles.28 The discrepancy between the two values points to the particles being effectively screened a t a distance 5-6a a t the limiting current density rather than at the value 2a which was assumed previously.2B It should be noted that the proportionality of the observed current to the number concentration of added particles does indeed show that the particles are screened from the feeder electrodes by adjacent layers of particles. Oxygen EuolutionlReduction. Figure 6 shows that as for the case of hydrogen evolutionfoxidation the rate of the bipolar oxygen evolution/reduction is proportional to the amount of the added dispersion for low values of the applied polarization. (In view of the complexity of the phenomena observed in the oxygen system, these are presented as a plot of Zpy/3v vs. y rather than as plots of the dimensionless parameter Z y/3i$.) At higher values of y the plots diverge: the decrease o/ipy/3u with increasing u shows that the solution is being depleted of oxygen during passage through the cell even though this is being regenerpfed in the anodic reaction on the bipolar particles. It should be doted that in these experiments the solution was presaturated with oxygen outside the cell in distinction to the experiments with hydrogen where this gas (or the hydrogen-nitrogen gas mixture) was directly sparged into the cell. The form of the polarization curves, Figure 6 , is clearly more complex than for the case of hydrogen evolutionfoxidation, Figures 4 and 5 . Three separate steps can be distinguished on the plots a t low hydroxide concentrations, Figure 7. The magnitude of the first of these can be attributed to the formation and reduction of an oxide layer due to the rotation of the particles, (15) and (16), provided Q in these equations is scaled by the true surface area of the particles and the screening distance =2a is replaced by =6a. This process can be observed in the case of the oxygen system because of the irreversibility of the oxygen evolution/reduction:
for the case of hydrogen evolutionfoxidation the corresponding process is masked by the high rate of the main faradaic reaction. Oxide formationfreduction can also be seen on the plots in Figure 6 which show that the reaction is detected a t surprisingly low . evolutionfreduction is also seen at low values of y ( ~ 2 ) Oxygen values of y, Figures 6 and 7. This early onset of the reactions can probably be attributed to aggregation of the particles in the alkaline media so that each aggregate experiences a higher potential difference than is indicated by the dimension of the individual particles. The height of the polarization curves increased only slowly with hydroxide concentration in the accessible range 10"-104 M indicating that the behavior of the dispersion electrolysis was controlled by the reduction of oxygen rather than oxygen evolution. Figure 7 shows that two limiting currents could be detected at low hydroxide concentrations in addition to that due to oxide formation/reduction. It is likely that the two waves are due respectively to
O2 + H 2 0 + 2e
-
H02- + OH-
and
O2 + 2 H 2 0 + 4e
-
-
40H-
(C) (D)
together with intervening decomposition of H02-: 2H02-
0 2
+ 20H-
(E)
Detailed interpretation of the kinetics must therefore be based on the control of the dispersion electrolysis by two irreversible electrode reactions. at least for the case of the first wave.5'
Conclusions The data presented for the hydrogen evolutionfoxidation and oxygen evolutionfreduction reactions show that it is possible to study electrode processes in poorly conducting or nonconducting media by means of bipolar electrolyses of dispersions. The results obtained for the first system fit predictions derived from a simple model based on the concepts of the mixture potential;" it is clear, however, that both for this system and for oxygen evolution/reduction the effects of mass transfer2* must be taken into account. The conditions used for the experiments reported in this paper are close to those of heterogeneous liquid-phase catalyses; it is
6400
J . Phys. Chem. 1986,90, 6400-6404
clear therefore that catalytic processes could be investigated (and monitored) provided electron transfer is involved. Catalytic reactions could also be investigated if no electron transfer is involved provided one couple undergoes a redox reaction. As measurements can be made in nonconducting media it is evident that such investigations could be extended to semiconductor (and possibly insulator) catalysts. Results for catalysts on oxide and zeolite (where the amount of precious metal per unit volume of solution is greatly reduced) as well as applications to syntheses42 will be reported elsewhere. Investigations of catalytic systems will require extensions of the modeling which will have to include the effects of reactions and concentration changes in solution. The results obtained also show that it should be possible to accelerate (or modify) catalyses by means of externally applied fields and the consequent faradaic current through the dispersion.
the gas constant, J mol-I temperature, deg volume of particles added per unit volume of solvent or solution, cm3 cm-3 voltage, V
R
T V
V
Greek Symbols
transfer coefficient polarization parameter overpotential, V angular position, rad angular position, rad rotational frequency, SKI resistivity, R cm potential, V polarization parameter angular velocity, rad s-' monolayer coverage, mol
a Y rl
0
x
Y
P
#J $ w
R
Acknowledgment. The support of the Office of Naval Research and the provision of precious metals by Johnson-Matthey PLC is gratefully acknowledged.
Subscripts
Glossary of Symbols radius of a particle, cm a concentration of a given species, mol cm-) c diffusion coefficient of a given species, cm2 s-I D E electrode potential, V F Faraday's constant, coulombs mol-' i current density, A cm-2 I current, A km mass-transfer coefficient, cm s d distance between the feeder electrodes, cm L number of electrons transferred per molecule n number of particles per unit volume, number cm-3 N N Avogadro's constant, number mol-' oxidized species of a redox couple 0 reduced species of a redox couple R
aPPl
0 192 C
lim N r
R S
T
P
denoting the exchange current pertaining to species 1 and 2 pertaining to the applied voltage cathodic pertaining to the diffusion-controlledcurrent pertaining to the net cross-sectional current density through the dispersion for one redox couple pertaining to the reversible potential pertaining to the current due to rotation of the particles pertaining to the solution pertaining to the cross-sectional current density pertaining to the cross-sectional faradaic current density through the dispersion
Superscript m pertaining to the mixture potential Registry No. H2, 1333-74-0;02,7782-44-7; Pt, 7440-06-4.
An ESR Study of Free-Radical Protonatlon Equiilbrla In Strongly Add Media H. F. Davis, H. J. McManus, and Richard W. Fessenden* Radiation Laboratory and Department of Chemistry, University of Notre Dame, Notre Dame, Indiana 46556 (Received: February 3, 1986)
Protonation of ascorbic acid radical, HP03', and SO,' in strong perchloric acid solution has been studied by means of changes in ESR parameters. The behavior of the ascorbic acid radical shows that the use of published acidity functions does not well describe the observations in that changes in ESR parameters occur over a much wider range of acidity function than expected. Treatment in terms of an equilibrium involving H+(H20), works much better. The pK so determined is -0.86 f 0.1 in contrast to an earlier value of -0.45. A continuous increase in hyperfine constant for HP03'- was observed from 650 G at pH 2.4 to 772 G at 61.2% HClO,. This change could not be explained by a single protonation to H2P03', but if a second step giving H3P03'+ was also included a reasonable fit to the data could be obtained. Approximate pK values of 0.0 and -1.7 were found for the first and second steps, respectively. No systematic change in the g factor for phosphite radical outside of the error limits f0.00008 was found over the whole range. The radical SO3'- was produced over the range from pH 11.3 to 61.2% HClO, and no change in the g factor larger than &0.00001 was found. It is possible that this radical does protonate to HSO,' but that no change in g factor accompanies this transformation.
Introduction
The present study is an outgrowth of an attempt to determine the pK corresponding to protonation of the radical SO3'- which is of some interest with respect to the chemistry of sulfite and SOZ.' In fact, this radical shows a single ESR line at the same g factor over a very wide range of conditions from normal aqueous solution to very strongly acid conditions (see below). We wish to know whether this observation can be used to show that SO3'- remains (1) Neta, P.;Huie, R. E.EHP, Enuiron. Health Perspect., 1985.64, 209.
unprotonated in strong acid. Ideally, one would measure the 33S hyperfine constant over the same range as a much more sensitive probe of the state of the radical. However, 33Sis sufficiently expensive and difficult to obtain that we chose to study the isoelectronic radical HP03'-. As a further outgrowth, the behavior of the ascorbic acid radical was studied under similar conditions in order to better define the behavior of radicals in strong acid. Various studies have established ESR as a very convenient method for investigating acid dissociation equilibria for radicals particularly in both strong acid2 and strong base.3 If the dynamics of the protonation equilibrium is sufficiently fast, then the ESR
0022-3654/86/2090-6400%01.5010 0 1986 American Chemical Society
