394
J. Phys. Chem. 1983, 87, 394-399
curvature observed in Figure 2 to a single falloff curve. In that case, it may be necessary to invoke a second parallel three-body recombination involving collision along the 4A'' surface followed by a spin-forbidden curve crossing to the electronicallyexcited 2A' state. This second channel should have a lower probability, due to the spin-forbidden crossing, but it should reach its high pressure limit at much lower pressures than the recombination along the ground 2A'' surface because of the longer lifetime expected for the 2A' state. Additional experiments at higher total pressures
will be needed to decide on the necessity of this second recombination channel. Such work is in progress.20
Acknowledgment. This work was supported by the National Science Foundation (CHE78-23867 and CHE 81-20834) and by a NATO Travel Grant (1849). We thank Dr. M. J. Pilling for several useful discussions. Registry No. 02,7782-44-7; methyl radical, 2229-07-4. (20) M. J. Pilling, private communication.
ARTICLES Klnetlcs of Hydrogen Productlon from Methyl Vlologen Radlcals on Colloldal Platinum M. S. Matheson, P. C. Lee, D. Melrel; ChemlsW Dlvfsion, Argonne Netbnal Laboratory, Argonne, Illlnols 60439
and E. Pellzzettl Istltoto dl Chlmlca Analnlca, Unlversta dl Torlno, 10125 Twlno, Ita& (Recehred: June 11, 1982; In Flnal Form: September 27, 1982)
The kinetics of hydrogen evolution from colloidal Pt catalyst reacting with methyl viologen radicals was studied in detail and the results are compared with those in previous studies of the same reaction catalyzed by gold sols. The major difference between the two catalysts is in the rate of the proton discharge step. In the pt-catalyzed reaction this step follows immediately the electron transfer reaction from the reductant to the particle. The Pt particle may therefore be viewed as a storage pool of hydrogen atoms rather than electrons. The dependence of the rate on catalyst concentation in Pt, as well as in Au, is second order in the catalyst and the possibility of a particle-particle reaction is further examined.
Introduction Application of photochemical and radiolytic techniques to studies of catalytic systems has already contributed substantially to the detailed understanding of this catalysis. The observation that noble metal catalysts are useful in photochemical hydrogen production is but one of the incentives for such investigati0ns.l Henglein and coworkers have shown that colloids of noble metals will catalyze multielectron transfer reactions of free radical^.^-^ Their studies have shown that the initial step in the transformation of reducing equivalents to dihydrogen is (1) (a) Koryakin, B. V.; Dzhabiev, T. S.; Shilov, A. E. Dokl. Akad. Nauk. SSSR 1977, 233, 359. (b) Lehn, J. M.; Sauvage, J. P. Nouu. J. Chim. 1977,1,449. (c) Kirch, M.; Lehn,J. M.; Sauvage, J. P. Helu. Chim Acta 1979,62,1345. (d) Moradpour, A.; Amouyal, E.; Keller, P.; Kagan, H. N o w . J. Chim. 1978, 2, 547. (e) Kalyanaaundaram, K.; Kiwi, J.; Gritzel, M. Helu. Chim. Acta 1978,61, 2720. Angew. Chem., Znt. Ed. Engl. 1979,18,624. (0 Okura, I.; Kim-Thuan, N. J. Mol. Catal. 1979, 5,311. (g) De Laive, P.; Sullivan, B. P.; Meyer, T. J.; Whitten, D. G . J. Am. Chem. SOC.1979,101, 4007. (2) Henglein, A. Angew. Chem., Znt. Ed. Engl. 1979,18,418. J.Phys. Chem. 1979,83,2209,2858. Ber. Bunsenges. Phys. Chem. 1980,84,253. (3) Henglein, A.; Lilie, J. J. Am. Chem. SOC.1981,103, 1059. J.Phys. Chem. 1981,85,1246. (4) Henglein, A.; Lindig, B.; Westerhausen, J. J.Phys. Chem. 1981,85, 1627. ( 5 ) Westerhausen, J.; Henglein, A.; Lilie, J. Ber. Bunsenges. Phys. Chem. 1981,85, 182. O022~3854/83/2087-0394$01.5010
the electron transfer step from the reducing radical to the metallic particle. When a metal of relatively high overpotential is used (e.g., silver), a large number of electrons could be stored on each particle and the capacitance of such microelectrodes could be measured? Recently, Miller et al.6 in an excellent report have developed an analytical expression for the electrodic behavior of the metallic particle based on the model proposed by S ~ i r o .Although ~ some details of the catalytic mechanisms are still unresolved, it seems now that a quite accurate description of such systems is given by the electrochemical approach. In a previous report we described the kinetics of the catalysis of the hydrogen evolution reaction on gold colloids using the popular methyl viologen (MV2+)radical cation as the reducing species? It was observed in this study that the electron transfer from the radical to the gold particle (reaction 1, where (M), is the metallic colloid) proceeds nMV+ + (M), + (M),"- + nMV2+ (1) at a diffusion-controlledrate, while the protonation of the charged particle (reaction 2) is very much slower than that (6)(a) Miller, D.; Bard, A.; McLendon, G.; Ferguson, J. J. Am. Chem. SOC.1981,103,5336. (b) Miller, D.; Mchndon, G. Ibid. 1981,103,6791. (7) (a) Spiro, M.; Ravno, A. B. J. Chem. SOC.1965, 78. (b) Spiro, M. J . Chem. SOC.,Faraday Trans. 1 1979, 75, 1507. (8) Meisel, D.; Mulac, W.; Matheson, M. J.Phys. Chem. 1981,85,179.
0 1983 American Chemical Society
The Journal of Physical Chemistry, Vol. 87, No. 3, 1983 395
H, Production from MV2+ on Colloidal Pt
(M),"-
+ H+ == (M)?-l)- + Hads
(2)
0.06
20
I
I
(b)
(0)
1
and very strongly pH dependent. It is clear, therefore, that in the net reaction of hydrogen evolution the protonation step or a later step such as the desorption step, be it chemical or electrochemical (reaction 3 or 4, respectively), 2Hads * H2 (3) (M),"-
+ H+ + Hads
(M)$-')-
+ H2
(4)
is the rate-determining step. In view of the slow rate of protonation of the gold (and silver3) particles it seems interesting to compare their behavior with that of Pt. In this case, due to the lower ovoerpotential and higher exchange current density on Pt, one expects a much faster protonation rate. This expectation was indeed observed in the present study. Another observation in the previous study was the strong dependence on the colloid concentration. In the case of gold, a close to second-order dependence of the rate on the gold concentration was observed. Higher than first-order dependence ofthe rate on concentration of Pt in a similar system was also r e p ~ r t e d . ~This point is further investigated in the present study.
Experimental Section The Pt catalysts were prepared by the citrate method following the modifications suggested by Brugger et al.1° Typically 51.8 mg of P t C 4 was dissolved in 500 mL and brought to boiling. A 1% solution of sodium citrate (60 mL) was then added and the mixture refluxed for 4 h. After allowing the mixture to cool a mixed-bed ion-exchange resin (Amberlite MB1) was added until the conductivity of the solution dropped to l0-S s. The colloidal solution was then fiitered and 50 mg of PVA (Polyscience, 99% hydrolyzed, M, = 110000) was added to 500 mL of the solution. In some cases sodium dodecyl sulfate (SDS) was added a t this stage instead of PVA. Analysis of Pt in the colloidal system was done by evaporation to dryness, dissolution of the precipitated Pt in aqua regia, and determination of the reoxidized Pt by atomic absorption. Concentrations of the Pt particles were determined from the measured total amoupt of Pt in the solution by taking the density of the bulk metal (6 = 21.4 g/cm3) and atomic weight (M, = 195 g/equiv) and assuming the radius of the metal part of the colloidal particle to be r = 16 A.loJ3 This would yield for the aggregation number n = 47r?N6/ (3MJ = 1.14 X lo3 atoms per particle and the concentration of particles would then be [(Pt),] = [PtIbt/n. The molar concentration of surface Pt atoms, [Pt],, was calculated from the number of surface Pt atoms per particle, n,, by using the equation [Pt], = [Pt]&t,/n). For the calculation of n,, the atomic radius of Pt was taken as rR = 1.39 A and by using n, = n[? - (r - 2rpt)3/?]. In the present case this yields n, = 495 atoms/particle. It was noticed that addition of 5 X M of H2 into these colloidal sols caused coagulation of the colloid several minutes after mixing with the H2-containing solutions. No such effect was observed in the irradiated solutions, but the effect of [H2]on kinetics was not checked on externally added H2-containingsolutions. UV-visible spectra of the colloidal sols showed only an increase of the absorption in the UV region. All chemicals used were of highest purity commercially available and were used as received. Solutions, all in triply distilled water, were deaerated by Ar bubbling with the
-
~~
~
~
(9) Kiwi, J.; GrHtzel, M. J . Am. Chem. SOC.1979, 101, 7214. (IO) Brugger, P. A.; Cuendet, P.; Griitzel, M. J. Am. Chem. SOC.1981, 103, 2923.
I
'
- 00 02
1
0
01
01
1 , sec.
02
I, sec
Flgure 1. Decay of absorption signal at 605 nm (a) and conductivity signal (b) following pulse radiolytic production of MV+ in Pt containing solutions. Both are deareated and contain [MV2+] = 5 X lo4 M, [Pt] = 7.5 X 10" M, [H+] = 5 X lo4 M, 1 % 2-Propanol, and 1% acetone. Solid line is the nonlinear mean squares best fii to a firstorder decay rate law.
syringe technique. The pulse radiolysis technique with either the spectrophotometric or the conductivity detection system has been previously described.l' Pulse widths of 2-40 ns, producing 5-60 pM total concentration of radicals per pulse, were used in this study.
Results and Discussion Since all solutions in the present study contained 1% v/v of both 2-propanol and acetone, all primary radicals produced on water radiolysis are converted to the 2propanol radicals which would then produce the MV+ radical cations through the sequence of reactions 5-8.
-
H20
eaq-,H, OH, H202, H2, H+
-
eaq- + (CH3I2CO OH (H) + (CH3)2CHOH (CHJ2COH
+ MV2+
-*
H+
(CH3)&OH
(CH,),COH
(CHJ2CO
(5)
(6)
+ H2O (H2) (7)
+ MV+ + H+
(8)
Details of this sequence of reactions were previously discussed.8 A large number of reports have shown that the net reaction of the sequence 1-4, i.e., the hydrogen evolution reaction, indeed occurs in the MV+ system when colloidal Pt is used as catalyst. However, it should also be realized that a small fraction of MV+ suffers hydrogenation during the catalyzed reaction.12 This complication, which was also observed on Au colloids? is considered here as well. However, we believe that the identity of the product, H2 or a small amount of hydrogenation product, should not affect the rate law which will describe the kinetics of the system. The kinetics of MV+ disappearance (reaction 1)in the presence of colloidal Pt was studied by following the decay of its absorption a t 605 nm under a variety of initial conditions. In parallel the rate of protonation of the colloidal particle (reactions 2 and/or 4) was measured by following the decrease in the conductance of the solution. Since the main contribution to the system's conductance is from H+, the conductivity signal corresponds to H+ consumption in the system. Note that the conductivity of the solution increases initially due to production of H+ in reaction 8. In Figure 1we show the decay of MV+ absorption and the decay of H+measured in similar solutions following pulses of the same dose. As can be seen in Figure 1,although the decay rate would not strictly follow a first-order rate law, (11) (a) Gordon, S.; Schmidt, K. H.; Hart, E. J. J.Phys. Chem. 1977, 81,104. (b) Schmidt, K. H.; Gordon, S. Reu. Sci. Instrum. 1979,50,1656. (12) (a) Keller, P.; Moradpour, A.; Amouyal, E.; Kagan, H. B. N o w . J . Chim. 1980,4,377. (b) Johansen, 0.;Launikonis, A.; Loder, J.; Mau, A.; Sasse, W.; Swift, J.; Wells, D. Aust. J. Chem. 1981, 34, 981.
398
The Journal of Physical Chemistry, Vol. 87,No. 3, 7983
the decay curve does fit such kinetics after a small fraction of the MV+ radical has decayed. In fact at doses higher than those shown in Figure 1 a larger fraction of MV+ would decay by first-order kinetics. At high enough doses the decay rate could be described by a pure first-order process. We therefore will analyze the kinetics in terms of the first-order rate constant although it is realized that at the early stages the kinetics are more complex. Several authors have already emphasized that the rate of the catalytic reaction should in principle follow the ButlerVolmer This would generally lead to fractional power dependence of the catalytic rate on the concentration of the electroactive species. However, under mass transfer limiting conditions, the catalytic rate will become first order in the concentration of the rate-limiting diffusing reactant. The conditions favoring such dependence in the present case will be high exchange current density and a relatively low concentration of one of the reactants,% vide infra. The first point to note in Figure 1 is the close correspondence between the decay rate of MV+ absorption and the consumption of H+ in the system as implied by the decay of the conductivity signal. This correspondence was found to hold over the whole pH range studied (pH 5-1) and under a variety of conditions of [Pt] and [Mv2+l0and on changing the stabilizer from PVA to SDS. We conclude therefore that reaction 2 follows immediately reaction 1 and only a small overpotential is necessary before protonation of the charged Pt particle occurs. This is in clear distinction from the observation when either Aua or Ag2-5 is used as a redox catalyst, where large amounts of electrons could be "stored" on the particle for long periods of time before protonation occurs. Perhaps not surprisingly, Pt, due to its high exchange current density in the hydrogen evolution reaction, can be viewed as a storage reservoir for adsorbed hydrogen atoms rather than electrons. Clearly in the case of Pt, the proton discharge step cannot be the rate-determining step, and when MV+ is used as the reducing radical, neither could the heterogeneous electron charging step be the rate-determining step. I t has already been shown previously that the rate of electron transfer from MV+ to Au colloids is diffusion controlled and there is no reason to expect that this reaction will be slower on Pt. In fact, if we take the conditions of lowest ([MV+], 2 X lo* M) doses and highest [Pt] (4.5 X M) used in this study, the second-order rate constant for reaction 1 is 2.2 X 1O'O M-' s-l (114 cm s-l). The above-measured value is probably slowed down by some coverage of the particle surface by H atoms; nevertheless, it is clear that the initial step of electron transfer is diffusion controlled. The rate of MV+ disappearance under a variety of conditions was checked. In general, behavior similar to that previously observed when Au was used as a catalyst8 was observed in the present case. It should, however, be noted that in the Au case the proton discharge step was well separated in time from the heterogeneous electron transfer step. The effect of accumulation of the products of the net catalytic reaction 9 (and probably of the hy2MV+ + 2H+ 2MV2++ Hz (9)
-
-
drogenation products) is shown in Figure 2. Not only is the rate of MV+ decay slowed when H2 accumulates but the equilibrium in eq 9 is established as can be deduced from the small but increasing amounts of MV+ which are left over at the end of the reaction. The rate of decay of MV+ depends inversely on the amount of [H,] accumulated as can be seen in Figure 3. (Results in Figure 3 were normalized to unity [PtI2because of the square depen-
Matheson et al.
0 C
o n L v) 0
n
a
0.0 0.0
1.0
2 .o
t , sec
Figure 2. Decay of MV+ absorbance following a string of four pulses, 40 ns each at 0.5s intervals. Same conditions as in Figure 1, except [Pt] = 4.5 X M, pH 4.9 (1 mM acetate buffer).
'0
[MV+I,,ILM
Figure 3. Observed decay rate constant, normalbed to the same [Pt]' at different total doses deposited in the solution. Obtained from experiments similar to those in FI ure 2. All conditions are as in Figure 2 except (O)/Pt] = 2.2 X 10- M, (0[Pt] = 3.1 X M, ( 0 )[Pt] = 4.5 X 10- M .
f
dence on Pt concentrations, see below.) Very similar to the situation on colloidal Au, the rate of the MV+ decay depends on [PtI2rather than being first order in Pt concentration. Higher than first-order dependence on [Pt] has been previously reported by Kiwi and G r a t ~ e l .We ~ have also found that this second-order dependence on [Pt] holds when SDS is used as a stabilizer rather than PVA. On the other hand, the dependence on pH is much less pronounced than in the Au case. Very little dependence on [H+] was observed in the pH range 1-6 when all other reaction conditions were kept constant. Hardly any effect of buffer concentration (10-4-10-3 M acetate, pH 4.8, buffer) on the rate of the reaction could be measured. The weak dependence on pH in this range is entirely predicted as was emphasized by Miller et al. Under the present experimental conditions ( [MV2+]= 5 x lo4 M; and assuming the standard rate constant for the heterogeneous electron transfer to be k: = 3.8 X lo4 cm/s, &?m+ = -0.44V; transfer coefficient a = 0.5; mass transfer coefficient mR = cm/s; all taken from ref 6), the pH at which the rate would drop to half its limiting value was calculated to be PH,,~6.9. From comparison with the experimental pHll2 data in ref 6b, we also conclude that
H, Production from MV2+ on Colloidal R
PH'/~,which is dependent on [MV+]/[MV2+],is always greater than 6.0 for all our experiments. Under similar conditions but on colloidal Au, pH1/2 -5.5 was observed and thus the strong dependence on pH with Au catalyst. Most of the experimental observations could be well accomodated with the electrochemical m0de1.w~ It is difficult, however, to explain the second-order dependence on the concentration of the catalyst, i.e., on the surface area of the electrode material, solely on electrode kinetics grounds. In our earlier studies on colloidal gold, we suggested that this dependence on colloidal concentration could result from a rate-controlling step involving collisions between two colloidal particles. In this sequence of reactions, the invoked interparticle reaction is suggested to enhance hydrogen desorption from the surface of the particle. In the following we further examine this hypothesis. We set the rate at which sites on the platinum particles are cleared as equal to twice the rate of particle-particle collision times the hydrogen atoms adsorbed per particle. Thus, one H atom covers one site. Further, this assumes that each collision clears all sites on both partners in the collision. This is reasonable since one or more of the -6000 A length PVA polymer molecules is attached to the small platinum particle (32 A diameter): thus the polymer coils extend an appreciable distance outward from the metallic particle. When the coils collide they will brush the attached section of the polymer across the platinum surface, releasing the hydrogen in the process. The rate of uncovering sites from the surface, vd,, in such a process will then be proportional to the frequency of particle-particle collisions and to the number of sites per particle already covered by hydrogen atoms:
In eq 10 [(Pt),)] is the molar concentration of particles calculated as outlined in the Experimental Section; [2H2], is the molar concentration of adsorbed hydrogen atoms. It should be noted that eq 10 predicts a first-order dependence of vd, on [(Pt)J at low concentrations of hydrogen which will change to second order in [(Pt)c] in excess H,. The rate at which sites on the platinum particles are occupied, V,&, is taken as twice the rate of encounters of H2with the particles times the fraction of unoccupied sites: Veda = ~ , [ ~ H ~ I I [ ( P ~-)6) CI(~ (11) where the fraction of unoccupied sites is
The molar concentration of sites is taken to equal the molar concentration of Pt surface atoms, [Pt],, which is calculated as shown in the Experimental Section. [ 2H2I1 in eq 11 is twice the molar concentration of H2 in the solution and the total concentration of H2, [2H2],,is then [2H2I1 [2H2],. At equilibrium Vad, = vd,, which yields I P t 1 8 - [2H218 [2H218 = -kd (13) [Ptl, k a [2HAt - [2H21s from which [2H2],can be calculated once the rate constants are known. The rate constant for the particleparticle collisions can be evaluated in two ways. Weigner and Tuorila14found the rate constant for collision between gold particles to be
+
(13) Aika, K.; Ban,L. L.; Okura, I.; Namba, S.;Turkevich, J. J.Res. Inst. Catal. Hokkaido Unio. 1976, 24, 54. (14) Wiegner, G.; Tuorila, P. Kolloid-2. 1926, 38,3.
The Journal of Physical Chemistry, Vol. 87, No. 3, 1983
397
4.1 x 109 M-1 s-1 (* dependent of particle size in the range of 29-970 A in diameter). We have used this value for '/&d (both collision partners are uncovered in a collision) in the present study. The value of k d can also be estimated by using the Smoluchowski equation and the measured diffusion coefficient (D= 2.05 X lo-' cm2/s) and colloidal radius (r = 110 A; Table I in ref 9 indeed verifies the independence of Dr on r and thus the constant value of kd at different radii) to yield '/&d = 3.4 X logM-' s-', close to the value taken above from Wiegner and Tu0ri1a.l~ The rate constant for collisions of H2with Pt particles k, can be estimated in a similar way by assuming a diffusion-controlled rate. For this rate constant we take a collision radius of 18 A (16 A for the Pt core and 2 8, for H2. Note that for adsorption to occur H2 and MV+ must collide with the small Pt core and not the envelope of the polymer chain). The diffusion coefficient is taken to be that of H2, Le., 3.6 X 10" cm2/s to yield k, = 4.9 X 1O'O M-' s-'. The desorption-adsorption equilibrium constant is then obtained to be kd/k, = 0.168. It should, however, be emphasized that this is a desorption-adsorption equilibrium which is determined by the particle-particle collision and is substantially bigger than the one under collision-free conditions, i.e., on the bulk metal. As mentioned above, MV+ reacts with the Pt (and Au) particles at a diffusion-controlled rate. The rate of its disappearance is therefore given by -d[MV+]/dt = k1[MV+][(Pt),](l - 0)
(14)
where kl is the rate constant of collisions of MV+ with the particles. In analogy with the gold particles we set this rate constant to be 7.0 X 1O'O M-I s-', a somewhat high value which reflects the negative charge on the particle. To apply the above mechanism to the experimental results in a given pulse, we calculate [2H2],at the point where half the MV+ in that pulse has already reacted. This (2H2), includes all the radiolytic H2 and all H2 produced from MV+ up to that point, including the H2 produced by previous pulses, if any, on the same solution. We assume that this [2H2], has attained the absorption equilibrium on the Pt particles and then calculate [2H2],. This value of [2H2], is then put in the expression
to calculate an average rate constant, k', for comparison with the experimentally measured rate constant for MV+ disappearance. The calculated and measured k's are compared in Figure 5 for those experiments where at midreaction [2H2],> [Pt],. For variations of measured k' from 3.1 to 1605 s-' the calculated k's are in reasonable agreement considering the approximations made and the experimental scatter sometimes found in the measured k's. A best fit of these results would yield kl = 8.3 X 1O'O M-' s-' as compared to the assumed value of 7.0 X 1O'O M-l s-l. We next treated the results at pH 1in a similar manner. A reasonable fit between the calculated and measured rate constants was obtained for this pH (and for all other pH's studied) as well. However, the best fit at pH 1 was obtained when kl = 3.5 X lo9 M-' & is assumed. The 20-fold decrease in Itl at pH 1 as compared to the one at pH 4.9 may be a result of the small positive charge on the particle at the former pH which is below the isoelectric point of this colloid. Furthermore, in Figure 6, the plot of k vs. [PtI2at pH 1has an intercept suggesting some contribution from a first-order term in the desorption reaction. The dependence of k1 on pH over the entire pH range studied is shown in Figure 7. The decrease below pH 5
396
Matheson et al.
The Journal of Physical Chemistry, Vol. 87, No. 3, 1983
A *0°
I 5 O l - - - - - 7
ti I
I
I
1
50
IO0
[Pt]'x IO",
M2
Figure 6. The dependence of MV+ disappearance on [Pi]' at 0.1 M HCiO,. Conditions as in Figure 4 except a second 24s pulse following a 40-ns pulse was used.
""
1
I
I
5
10
15
[Pt]'X IO'',
t
M2
Figure 4. The dependence of MV+ decay rate on [Pi]'. Experimental conditions were as follows: [MV2+]o = 5 X l o 4 M; pH 4.9; 40-ns pulse producing 4.1 X M MV'.
c L
, I
I J
IO0
10
I
I
I
I
I
I
000
k m e o s , sec-'
Flgure 5. Comparison of the calculated rate constants with the measured ones for the decay of MV'. The error bars represent experiments in which [2H,], < [Pt], where [2H,], is difficult to estimate. Experimental conditions were as follows: [MV2+Io= 5 X lo4 M; pH 4.9; [Pt] = (1.5-45) X M; (0) 2 ns; (0)4 ns; (0)10 ns; (0)20 ns; (0)40 ns pulses.
was already attributed to the decrease in the negative charge on the particle. The decrease at pH > 5 is probably due to the decrease in the free energy of the net reaction at the higher pH's. Comparison with Other Studies The effect of the free energy of the reaction on the rate brings back the question of the effect of the redox potential of the two reacting couples as well as that of the colloid. As has already been shown3v6the rates of the charging reaction 1 and the proton discharge reaction 2 should follow Butler-Volmer kinetics. In a rigorous treatment, therefore, the effect of the change in redox potential as the reaction proceeds should be included in the rate expression of eq 14. A complete treatment of the system was given by Miller et al. However, since the desorption rate is the slowest step in the system, and since our experiments were
(15) Gratzel, C.
K.;Gratzel, M.J . Am. Chem. SOC.1979, 101, 7741.
J. Phys. Chem. 1983, 87, 399-407
core is -20 A and furthermore, at the higher [Pt] the reaction observed in their time-resolved experiments is primarily the charging reaction 1 with the bare metal surface and before an appreciable amount of H2 is adsorbed. Miller and McLendonebprepared Pt-PVA colloids by the same method of Kiwi and Gratzel and with moderate centrifugation have obtained hydrodynamic radius of 360 A (asmeasured by the latter authorsg) and Pt core radius of 300 A,measured by electron microscopy. Such a large radius for the Pt core would then yield extremely low particle concentration M) and preclude any particle-particle collision during the reaction. On the other hand, for this method of preparation, Kiwi and Gratzel found 20% of the initial amount of Pt colloid before centrifugation had an hydrodynamic radius of 110 A and our calculations above indicate a core radius of -20 A. With the same preparation method Miller and McLendon should have more than 20% such small particles in their partially centrifuged samples. For 40% by weight the concentration of 20-A particles would be (for 6 pM Pt) 1.1 X lo4 M [(Pt),]. With 1/2kd= 3.4 X lo9 M-I s-l calculated above for these particles, this concentration gives a rate of particle-particle encounters of 3.8 s-l, and an even higher rate for encounters of 1.1X M 20-A particles with 0.5 X 10-l2M 3 0 0 4 particles. Miller and McLendon give a first-order rate constant for MV+ disappearance in 6 pM Pt of 0.078 s-l.
399
Our conclusion is that in our experiments (well below pHlI2)using the preparation methods of Brugger et al.1° the particle-particle interaction can reasonably explain the [PtI2dependence. Further, if the Miller and McLendon preparation is similar to that of Kiwi and Gratzel, then a number of particle-particle collisions occur during the reaction of MV+ in their experiments and our hypothesis may be applicable. Using our theory, we calculate a first-order rate constant of 0.29 vs. 0.078 measured. However, as this experiment appears to be near pH1 2, we have probably used too high a value for k l (lo1' M-( s-l), and our simple kinetic theory needs modification. If particle-particle collisions occur frequently during MV+ oxidation, then the proposed thin-cell a n a l o d b loses significance in expIaining the [Pt] dependence. However, this thin-cell reaction may be totally dominated by pairs of particles which are essentially in collision so that at any moment the number of such pairs is proportional to the collision rate, i.e., proportional to [PtI2.
Acknowledgment. We are much obliged to D. Ficht and G. Cox for their dedicated operation of the linac. Support of this research by a NATO Research Grant No. 1780 to E.P. and D.M. is gratefully acknowledged. Work at ANL is performed under the auspices of the Office of Basic Energy Sciences, Division of Chemical Science, US.DOE, under Contract No. W-31-109-ENG-38. Registry No. H2, 1333-74-0; Pt, 7440-06-4; MV+, 25239-55-8.
One-Electron Transfer Equilibria and Kinetics of N-Methylphenothiazine in Micellar Systems Claudlo Mlnero,+ Edmondo Pramauro, Ezlo Pellzzetti, Istituto di Chimica Anaiitica, Universiti di Torino, 10 125 Torino, Italy
and Dan Melsel" Chemistry Division, Argonne National Laboratory, Argonne, Illinois 60439 (Received September 20, 1982)
The kinetics and equilibria of electron transfer between N-methylphenothiazine and aquoiron(II1)or octacyanomolybdate(V) in the presence of various micellar-forming surfactants (CTAN,Triton, SDS) were investigated by using several spectrophotometric techniques. The presence of micelles strongly influences the equilibria and the kinetics of the reactions. Binding constants are determined and electrostatic and hydrophobic effects in the various micellar systems are analyzed.
Introduction Electron transfer reactions involving phenothiazine derivatives are currently under intensive investigations. The interest in this class of compounds stems from their exceptional photoredox properties' as well as from their physiological activity.2 Micellar effects on the electron transfer reactions of these compounds are of relevance to both these areas of interest. On one hand, micelles are often used as model systems to mimic cellular membranes while the interactions of the phenothiazine drugs with such membranes may affect their therapeutic properties. On the other hand, several phenothiazine derivatives have been used in micellar or vesicular designs for photochemPresent address: CISE, Segrate (Milano), Italy.
ical energy storing systems. In view of this widespread interest we report here on the effects of micelle-forming surfactants on the equilibrium and kinetics of the electron transfer reactions between N-methylphenothiazine (MPTZ) and aquoiron(II1) or octacyanomolybdate(V). Although micellar effects on electron transfer rates have been extensively studied, reports on their effects on oneelectron transfer equilibria are rare. Kinetic studies have already provided a wealth of information on the mechanistic details of the micellar effects on electron transfer reaction^.^!^ When kinetic effects are combined with redox (1) Infelta, P. P.; Gratzel, M.; Fendler, J. H. J. Am. Chem. SOC.1980, 102, 1479. Moroi, Y.; Infelta, P. P.; Gratzel, M. Ibid. 1979, 101, 573. (2) Forrest, I. S.;Carr, C. J.; Usdin, E. Adu. Biochem.Psychopharmcol. 1974, 9. Gasco, M. R.; Carlotti, M. E. Pharm. Acta Helu. 1977,52, 296.
0022-3654/83/2087-0399$01.50/00 1983 American Chemical Society
