J. Phys. Chem. 1984, 88. 1615-1623 case b3: PB' and PB-; Cpb(max) = (1/2)(cA E = Eo
+ (f/2)
log
cpb
- CK)
- (f/3) log CK+ cf/2) log (CA - CK - CPb) lA201
Note that the Nernst slope with respect to the concentration of PbBr2 is f/2 in all cases. A maximum in potential is observed only when the species PB' or PB- or both are assumed to determine the equilibrium at the interphase. The position of the maximum varies between 0.4 and 0.67 of the concentration of free aluminum bromide, depending on the relative amounts of PB' and PB- that determine the equilibrium. Equal amounts of both species give rise to a value of 0.5. Appendix 3 In this Appendix the results of calculations similar to those shown in Appendixes 1 and 2 are given, but for a monovalent cation. In addition to the ionic species shown in eq A l , we consider also the following different species assumed to determine the equilibrium potential a t the interphase:
-
+ 3AlzBr, 6MBr + 3Al2Br6 3MBr
-
+ [M(AlZBr7)J 2[Mz(AlBr4)]' + 2[M(A1Br4),][Mz(AlzBr7)]'
[A211 [A221
where M stands for any monovalent ion. Let us write, for brevity, CA+ and CA- for the complex ions on the rhs of eq A21 and CB+ and CB- for the ions on the rhs of eq A22. From eq A21 and A22 it can be seen that [CA'] = [CA-] = [CB'] = [CB-] = (1/3)CM [A231 and from eq A1 and A22 the concentration of free Al2Br6in cases d below is [A] = C A - C K - ( ~ / ~ ) C M
+ cf/2)
E = Eo
~4241
The listing below is arranged as in Appendix 2. case c l : CA+: no maximum
log C M - cf/6) log CK
case c2: CA-; no maximum E = E o + cr) log C M - (2f/3) log CK case c3: CA+ and CA-; no maximum E = E o (2f/3) log C M - cf/3) log CK
+
case c4: CA+ and 4CA-; no maximum E = Eo (5f/6) log CM - cf/2) log CK
+
case d l : CB+; CM(max) = (4/3)(CA - C,) E = Eo + cf/2) log C M - cf/6) log C K + cf/4)lOg (CA - CK - 0.5CM) [A291 case d2: CB-; CM(max) = CA - C K E = Eo + cr) log CM - (2f/3) log CK cr) log (CA - CK - 0.5CM) [A301
+
case d3: CB+ and CB-; CM(max) = (8/7)(CA - CK) E = Eo + (2f/3) log CM - cf/3) log CK cf/2) log (CA - CK - 0.5CM) [A311
+
case d4: CB+ and 4CB-; CM(max) = (20/19)(CA - CK) + (5f/6) log CM - cf/2) log C K (3f/4) log ( C A - CK - 0.5CM) [A321
E = Eo
+
Note that the Nernst slope with respect to a monovalent salt like CuBr varies between f/2 a n d f , depending on the relative amounts of CA' and CA- and/or CB' and CB- that determine the equilibrium. The amounts given in case c4 or d4 give rise to a value of (5/6)fas obtained experimentally for the copper salts. Registry No. Ag, 7440-22-4; AgBr, 7785-23-1; Cu, 7440-50-8; CuBr, 7787-70-4; Sn, 7440-31-5; SnBr,, 10031-24-0; Pb, 7439-92-1; PbBr,, 10031-22-8; AI2Br,, 18898-34-5; KBr, 7758-02-3; ethylbenzene, 10041-4.
Reactions of Colloidal Platinum in Aqueous Solutlons Containing Methyl Viologen, Its Cation Radical, and Hydrogen, Studied by Pulse Radiolysis Marek Brandeis, Gad S. Nahor, and Joseph Rabani* Energy Research Center and the Department of Physical Chemistry, The Hebrew University of Jerusalem, Jerusalem 91 906, Israel (Received: March 20, 1983; In Final Form: August 23, 1983)
The reactions of methyl viologen ions, MV2' and MV', and molecular hydrogen in the presence of colloidal Pt were investigated by pulse radiolysis. In the pH range 1.5-4 all the reducing species in the system were converted into molecular hydrogen. At pH 10, practically all the reducing species produced MV', while in the intermediate pH range (6-8) the predominant species were Pt particles loaded with hydrogen in the form of hydride ions. The kinetics of reactions of the Pt colloidal particles with MV' and H2 were investigated. The rate of reaction of MV' with the Pt was first order in MV' and second order in Pt. When the reaction of MV2' with loaded Pt was also important, the overall order with respect to Pt was less than 2. A loading mechanism involving particle-particle collisions is examined in relation to the effects of pH, [Pt], and repetitive pulses.
Introduction Recently, there has been a growing interest in the catalytic reactions of metal colloids.'-'3 Intensive investigations were
carried out on the catalytic production of hydrogen from methyl viologen radicals (MV') and water using Pt ~ o l l o i d s . ~The ~~-~ ( 6 ) D. S. Miller, A. J. Bard, G. McLendon, and J. Ferguson, J . Am. Chem. Soc., 103, 5336 (1981).
(1) A. Henglein, J . Phys. Chem., 83, 2209 (1979). (2) J. Kiwi and M. Gratzel, J . Am. Chem. Soc., 101, 7214 (1979). (3) D. Meisel, W. A. Mulac, and M. S. Matheson, J . Phys. Chem., 85, 179 (1981). (4) A. Henglein and J. Lilie, J . Am. Chem. Soc., 103, 1059 (1981). (5) K. Kopple, D. Meyerstein, and D. Meisel, J . Phys. Chem., 84, 870 (1980).
0022-3654/84/2088-1615$01.50/0
(7) D. S . Miller and G. McLendon, J . Am. Chem. Soc., 103,6791 (1981). (8) M. S. Matheson, P. C. Lee, D. Meisel, and E. Pelizzetti, J . Phys. Chem., 87, 394 (1983). (9) T. Nishijima, T. Nagamura, and T. Matsuo, J . Polym. Sci., Polym. Lett. Ed., 19, 65 (1981). (10) A. Moradpour, E. Amouyal, R. Keller, and H. Kagan, Nouu. J . Chim.,2, 547 (1978).
0 1984 American Chemical Society
1616 The Journal of Physical Chemistry, Vol. 88, No. 8, 1984
potential use in photochemical reduction of water has contributed to an increasing interest in such metal colloids. Generally the catalytic formation of H2 has been attributed to the charging of the colloidal particles by reducing species (such as MV+, (CH3),COH, etc.) followed by discharging of the colloids by water or H+ ions: Although this general scheme seems simple, there are still several unanswered questions about the role of the metal colloids: What is the nature of the catalytic process, and in what form are the reducing species loaded on the Pt colloids? Is there an equilibrium between the loaded reducing species, H2, and the reductants (e.g., MV+) and, if such an equilibrium exists, how rapidly is it established? Additional complications arise from the fact that most metal colloids are metastable, and their properties depend on the preparation procedures. The results may not be reproducible, and in addition some of the reactions and species are difficult to follow directly. This report is the first in a series aimed at clarifying the situation in the MV+-Pt-H2 system. We have investigated this system using the pulse radiolysis technique mainly under conditions where the equilibration reactions of MV+, MV2+,H2, and H+ are observed. We have found conditions for reasonable reproducibility and carried out kinetic measurements as well as determined the equilibria states following the electron pulses. Our interpretation of the results is based on a homogeneous kinetic model3*,which can be used in direct relation to the phenomena that we observed in these systems. This is not meant to rule out the possible applicability of the electrochemical approach to colloid c a t a l y s i ~ . ~Moreover, ,~ we hope that the homogeneous kinetic treatment may in the future allow the calculation of the electrochemical parameters.
Brandeis et al. sults). Kinetic studies were followed at 577 nm unless otherwise stated, taking cMV+ = 1.1 X lo4 M-' cm-'. Buffers consisted of standard pH control solution^.^^ The materials used were HC104, HCl, KC1, NaOH, NaH2P04, Na2HP04,and Borax. All chemicals were of analytical grade and were used without further purification. Unless otherwise stated all solutions were deaerated by bubbling with high-purity H e (Matheson). Apparatus. UV-visible absorption spectra were recorded on a Bausch and Lomb Model Spectronic 2000 spectrophotometer. The pulse radiolysis setup consisted of a Varian 7715 linear accelerator. The pulse duration ranged from 0.1 to 1.5 ps with a 200-mA current of 5 MeV. The total concentration of MV' radicals produced per pulse was measured by the initial absorption signal of MV+ at 577 nm and was found to be in the range of 1.C-13.5 pM. Irradiation cells of lengths 0.5, 1.0, or 2.0 cm made of high-purity silica were used. A 150-W Xe-Hg lamp was used as the analytical light source and appropriate light filters were used both to avoid photochemistry and to eliminate any scattered light. The detection system included a grating monochromator and a 1P28 photomultiplier. The signal was transferred to a Nova 1200 minicomputer via either a Biomation 8100 or an analogto-digital converter where analysis of the data was carried out.
Experimental Section Materials. The Pt catalyst was prepared by the following procedure: 2.0 g of poly(viny1 alcohol) (PVA, M, 14000 purchased from BDH) was dissolved in about 300 mL of hot water. After the solution was cooled, 1.0 g of H2PtC16(BDH) with a 40% (by weight) Pt content was added. The reduction of Pt metal was carried out by slowly introducing NaBH, powder (BDH) to the continuously stirred solution until it turned black and no bubbles of hydrogen were observed (about 20 min). The black solution was then treated by dialysis in distilled water for 48 h. Finally, the resulting colloidal solution was filtered by suction through a 0.45-pm Millipore filter and diluted to a total volume of 500 mL with water to give a solution with a formal Pt concentration of 4.0 mM. This method of preparation was used rather than the more common citrate ion m e t h ~ d , ~ ,since ' , ~ it produced Pt free of any organic substances. The concentration of the Pt after dialysis was determined gravimetrically after coagulation which was achieved within several minutes following the addition of cerous ions in acidic pH, and bubbling hydrogen. The size of the Pt colloidal particles was determined to be 50 f 5 A in diameter except for batch 2605. Thus, the actual concentration of the colloids could be calculated to be 4000 times smaller than the respective formal Pt concentrations. The light scattering technique was used to determine the colloidal particle size. This was done by using 500-800-nm light, where turbidity was linear with X4. As a reference we used Ludox solution (colloidal silica, product of Du Pont, particle diameter 30 A). In addition, the diameter was cross-checked by comparison of our kinetic results in acidic pH with previously published data (Figure 4 in ref 8). The two methods resulted in the same Pt particle diameter within 10%. 2-Propanol (J. T. Baker) and methyl viologen (Sigma or Aldrich) were used as received. The MV' cation radical was generated by using the pulse radiolysis technique in solutions containing 2-propanol (see Re-
Results When an aqueous solution containing methylviologen is pulse irradiated, it reacts with eaq-,H atoms, and O H radicals, and the reaction rate constants are 7.5 X lolo, 6 X lo8, and 4 X lo8 M-' s-l, re~pective1y.l~ When 2-propanol is added to a solution containing methyl viologen, it competes for H and OH. The 2-propanol radical, CH3COHCH,, produced by abstraction of a hydrogen atom, is capable of reacting with methylviologen, MV2+, producing MV+. Thus, in the presence of a sufficiently high concentration of 2-propanol, it is possible to quickly convert practically all of the primary radical species which are produced by the irradiation to MV+. (The reaction rate constant of CH3COHCH3 with MV2+ is 3.5 X lo9 M-' s-l (r ef 3)). The following species have to be considered at the end of the electron pulse or shortly (a few microseconds) after: MV' (produced with a G value of 6), the so-called "molecular yield" of H202(G = 0.75), and H,. There are three sources of H,: (1) the molecular yield (GH2= 0.45), (2) H2 produced by the abstraction reaction of H atoms with 2-propanol, where the primary yield of H atoms is G = 0.55, and (3) H atoms produced by the reactions of e,; (G, = 2.7) with H+ ( k = 2 X 1O'O M-' s-l at zero ionic strength), in competition with the reaction of eaq- with MV2+. Therefore, the total yield of H2 will range from 1.O to 3.7 depending on pH. In addition to eaq-, H , OH, HzOz, and H2, there are two other primary products of the radiation, namely, H+ and OH-. Since in all our pulse radiolytic experiments we used either M acid or base, or a buffer, and in either case the neutralization reactions were fast in comparison with the reactions of MV+, it is justified to assume that the pH of the solutions was determined only by the acid, base, or buffers present. Moreover, as will be seen later, no net significant amounts of H+ were expected to accumulate under our conditions, even when repetitive pulses were used. Reactions in the Absence of Pt. In the absence of Pt, MV' which was produced by pulsing a solution containing 0.2 M 2propanol and 4 X 10" M MV2+ a t pH 8.1 was stable for about 1 s. A partial decay was observed later within about 1 min and amounted to about 30% of the initial absorption of the MV'. This decay took place in the pH range 1.5-11, but accurate measurements could not be carried out because of the instability of the analytical light in this time range. We attribute this process to the reaction of MV' with the molecular HzO2. Traces of oxygen, if present, could contribute to increase the magnitude of this decay. When excess HzO, was initially present ( lo-, and 5
(11) J. M. Lehn and J. P. Sauvage, N o w . J . Chim., 1, 449 (1977). (12) K. Kalyanasundaran, J. Kiwi, and M. Gratzel, Helu. Chim. Acta, 61, 2720 (1978). (13) K. I. Zamaraev and V. N. Parmon, Catal. Rev.-Sci. Eng., 22, 261 (1980), and references cited therein.
(14) "Lange's Handbook of Chemistry", J. A. Dean, Ed., McGraw-Hill, New York, 1979, Tables 5-23, pp 5-77-5-78. (15) S. Solar, W. Solar, N. Getoff, J. Holcman, and K. Setested, J . Chem. SOC.,Faraday Trans. 1, 78, 2467 (1982).
The Journal of Physical Chemistry, Vol. 88, No. 8, 1984
Reactions of Colloidal Pt in Aqueous Solutions
1617
TABLE 1: Effect of pH on the Reducing Speciesa PH
[MV+1 efdb
5.8 6.1 6.8 7.3 7.65 7.95 8.7 9.0 9.55 9.95 10.9 11.0
0.18 0.38 0.28 1.03 1.23 2.40 3.56 1.28 9.67 9.36 12.26 12.16
tl,,
fd
2.1 2.3 2.8 2.1 ~1.7
Z(PtC)efdd 13.07 12.87 12.97 12.22 12.02 10.85 9.67 5.97 3.58 3.89 0.99 1.09
E0(PtHnn-)efde [MVclmbuf -82.7 -92.4 -64.8 -83.2 -77.5 -85.7 -74.1 -83.5 -74.9 -62.4 -41.8 -29.9
[H2 lmbug
Z(Ptc)mbuh
Pt batch no.
0.895
15.70 15.29 17.34 16.58 16.42 15.25 14.02 9.58 6.12 6.29 2.78 2.99
3012 2408 3012 2408/3012 3012 3012 2408 3012 3012 2408/3012 3012 3012
1.00 0.02 0.03 0
0 0 0 0 0 0 0
3.65 8.09 11.55 11.38 14.89 14.68
[MV'] at the end of the fast decay, a Concentrations normalized t o [MV+], = 1.33 X M and [H,], = 2.21 X M at pH >4.1. First half-life of in wM. Corrected for the small residual absorption (up to 5% correction, see Discussion of the Results in Acidic pH). Reducing equivalents loaded on the Pt at the end of the fast decay calculated as the difference [MV'], the fast decay in milliseconds. [MVtjefd, in MM. e The standard redox potential of loaded Ptc vs. NHE (see eq 6) in millivolts. Note that this redox potential may [MV*] at the end of the buildup reaction (wM) given whenever the buildup process was observed. It depend on the degree of Pt loading. was assumed that, at this stage, equilibration with both H, and loaded Pt was established. The effect of H,O,, which usually reacted later, was neglected. g [H,](WM) at the buildup process (see also footnote f). Equivalents loaded on the Pt (wM)at the end of the buildup process (see also footnote f ) .
IO4 M), an enhanced decay of the MV' was observed (Figure l ) , the rate of which depended on the pulse intensity and was proportional to the [H202]added. The order of reaction with respect to MV+ was not 1, as might be expected if a simple reaction between MV+ and Hz02took place, but rather close to zero order. This indicates a complex mechanism which is beyond the scope of this manuscript. However, we can conclude the following: (a) The lifetime of the decaying MV' in the absence of initially added H202was consistent with the assumption that this partial decay is due to the molecular H20z. (b) The rate constant was found to be almost linearly dependent on [H202], when excess H,O, was initially present. Under these conditions the MV' absorption decayed to zero. (c) The reaction of MV' with H202did not proceed (at least under our conditions) via O H radicals, since, if this had been the case, O H would have reacted with 2-propanol by H abstraction, and the CH3COHCH3radical thus formed would have produced MV'. Since these reactions are extremely fast in comparison with the reaction between MV' and H20z,the net result would have been the destruction of H,02 by a chain reaction involving MV' and CH,COHCH,, without any net decay of MV+. (d) Dimerization of MV+ was ruled out under our conditions since the concentration of MV' was too 10w.l~ Results in the Presence of Pt. When a deaerated solution containing MVZ+(4 X lo-" M) and CH3CHOHCH3(0.2 M) was pulse irradiated in the presence of colloidal Pt, the MV' absorption changed with time until equilibrium was obtained. The time profile depended on the pH, [Pt], [MV"], and pulse intensity and on the number of pulses given to a solution prior to the test pulse. In the following we will discuss these effects separately. Effect of p H at Near Neutral and Alkaline Solutions. Typical results are shown in Figure 2 for the pH range 6.1-12. At pH 6.1 most of the MV' decayed away in the millisecond time scale, by a nearly first-order rate law. At higher pHs, the decaying fraction became gradually smaller and its lifetime longer, and there were greater deviations from a first-order rate law. This process (from now on termed "the fast decay") was followed by additional changes in the optical absorption. These changes were observed when the absorbance that remained at the end of the fast decay (abbreviated Defd)was relatively large. At pHs above 9, the fast decay was followed by an increase in the optical density which took place during about 0.5 s (in the following, "the buildup"). This process is shown in Figure 2. It amounts to an increase in the absorbance by up to 33%, and its rate seems to be pH independent. Finally, the absorbance partially decayed away in the time range of several seconds (in the following, "the slow decay"). The rate of the slow decay seems to depend on the pH and possibily also on the pulse intensity, and on the [Pt]. However, we enX
(16) E. M. Kossower and J. L. Cotter, J . Am. Chem. Soc., 86, 5524 (1964).
b
I
I
I
I
I
-
-
I
Figure 1. Reaction of MV' with H202: 4 X M MVZ+, 0.2 M 2-propano1, no Pt present; 1.5-ps pulse, 1.5-cm light path; (a) no H202 initially added, (b) 1 X M H2O2initially added, (c) 5 X M H 2 0 2initially added.
countered some difficulties in the investigation of this process because of the instability of our analyzing light over such long time scales. The decaying fraction of this process was always over 30%. We would like to stress that, in different series of experiments, the same general features were observed as described above. However, experiments carried out on different dates or using different Pt stock solutions often showed certain quantitative differences. This will be discussed later. In Table I we show the effect of pH on the equilibrium concentrations of MV', H2, and the reducing equivalents loaded on the Pt colloidal particles. This was calculated as the difference between the total concentration of reducing equivalents produced
1618
The Journal of Physical Chemistry, Vol. 88, No. 8,1984
Brandeis et al.
Figure 2. Time profiles of the optical absorption of MV+ at various pHs: 1.5-ps pulse, 2-cm light path; 4.5 X no Pt, the same profile was observed in the entire pH range investigated; (b)-(g) 4 X M Pt, batch 2408.
+
in the system (G = G, + GOH + 3GH 2GH2= 8.0) and the sum [MV+] 2[H,] at equilibrium. [MV+] was calculated from its absorption, and [H,] calculated from the [MV’], by using eq 1.”
+
M MVZ+,0.2 M 2-propanol; (a)
a
EH+/H2= 0.0919 -0.059 log ([H2]’/2/[H+]) = EMVZ+/MV+= -0.44 - 0.059 log ([MV+]/[MV2+]) (1)
H2 was generated by the pulse with a G value of 1.O under the conditions of Table I; namely, initial [H,] was 2.2 X 10” M. From Table I it is clear that under equilibrium conditions at pH 6.1 very little or no H Zwas present. Only about 4% of the initially produced MV+ remained at pH 6.1 when equilibrium was established. The concentration of H2 (equivalents) expected to be present in equilibrium with MV+ was only 18% of the total reducing equivalents. Table I shows that, in the pH range 6.1-8.7, most of the reducing equivalents are loaded at equilibrium on the Pt. The nature of the loaded species will be discussed later. At pH 9, a sharp decrease of the number of reducing equivalents loaded on the Pt was observed and, upon further increasing of the pH, a growing fraction of the reducing species remained in the form of MV+. As the pH decreased below pH 5, H2 was expected to become the predominant species in the system. More details about our results in acid solutions will be presented later. Effect of Pulse Intensity. Solutions containing 4 X M MV2+ and 0.2 M 2-propanol were irradiated with different pulse intensities. It was foufid that decreasing the pulse intensity, namely, decreasing [MV’], resulted in increasing the fraction of absorbance decaying away in the fast decay, and speeding up the rate of this process. Typical results are presented in Figure 3. The rate law of the fast decay became closer to first order when the pulse intensity decreased. Effect of MP’. Increasing [MV2+]resulted in decreasing of the fast decaying fraction. At sufficiently high [MV2+],the fast decaying fraction could be completely eliminated at moderately high pHs. The rate of the decay was not much affected. In Figure (17) [H,] is in M and the value 0.0919 is the correction factor between pressure and concentration units.
-
..
20ms
I
I
I
I
I
1
I
I
J
0
b k
slopes160 WC-I
^I
1
rc
1
Figure 3. Effect of pulse intensity on the fast decay: 4 X lo4 M MV2+, 0.2 M 2-propano1, 4 X M Pt, batch 1805, pH 9; 1.5-cm light path; (a) 0.75-ps pulse, (b) 0.2-ps pulse. Right: the first-order plots. Note the deviations at the higher pulse intensity (discussed in the text). 4 we present typical examples for this effect.
Effect of Pt Concentration. In alkaline solutions, the fast decaying fraction increased with [Pt], Typical results are shown in Figure 5, where increasing the [Pt] from 1.6 X to 1 X M resulted in an increase of the fast decaying fraction from about 0.6 to nearly 1. Additional results on the effect of [Pt] on the fast decaying fraction at various pulse intensities at pH 8.1 are presented in Figure 6. The ratio [MV+]/[Pt] has a strong effect on the magnitude of the fast decaying fraction. Thus, with relatively high [Pt] and relatively low initial [MV’], the fast decaying fraction was near to 1, while at low [Pt] and high initial [MV’] it was considerably lower.
Reactions of Colloidal Pt in Aqueous Solutions
The Journal of Physical Chemistry, Vol. 88, No. 8, 1984 1619
4ms
I
I
I
I
I
0
I
I
I
I
04
0.8
12
1.6
20
[Ptl,m M 4ms CI I
I
1
I
Figure 4. Effect of MV2' on the fast decay: 0.2 M 2-propanol,4 X lo4 M Pt (batch 607), pH 8.3; 1.5-ps pulse, 1.5-cm light path; (a) 4 X M MV2'. M MV2', (b) 4 X
r i
rt2
b
L -4
Figure 5. Effect of [Pt] on the fast decay: 4
X M MV2', 0.2 M 2-propano1, pH 8.1, Pt batch 1805; 1.5-ps pulse, 1.5-cm light path; (a) M Pt, (b) 1.6 X lo4 M Pt. 1.0 X
Figure 6. Fast decaying fractions as a function of the [Pt]: 0.2 M 2-propanol,4 X M MV2*,pH 8.1, Pt batch 1805; (+) 0.2-ps pulse (producing on the average 4.6 X lod M MV'), (m) 0.75-ps pulse (producing on the average 8.2 X 10" M MV'), ( 0 ) 1.5-ps pulse (producing M MV'). on the average 1.4 X
than those extrapolated from results in previous work? indicating that our Pt particles (in the following, Ptc), were about 50% larger in diameter, resulting in approximately 3.5 times lower [Ptc] at a given [Pt]. Effect of Repetitive Pulses. In the acidic pH range the absorption produced by the pulse decayed away to nearly zero, even when repetitive pulses were given. The decay rate became slower with every additional pulse. After 100 pulses of 1 . 5 - ~duration s the rate of decay was about 100 times slower than in the first pulse. The rate of the decay which was proportional to [PtI2 in the first pulses gradually became proportional to [Pt] after 100 pulses. When shorter pulses were applied, the same change required more pulses. Typical results are shown in Figure 9. In the pH range 7.0-9.0 the effect of repetitive pulses resembled the effect of increasing pH; Le., the fast decaying fraction decreased in magnitude with every additional pulse and the decay kinetics became gradually slower until after several pulses no decay was observed. Instead, a buildup process of MV' took place in the same time range as it occurred in the first pulse at higher pHs.
Discussion Near Neutral and Alkaline Solutions. Two schemes have been proposed to explain the activity of a colloidal catalyst in redox reactions. Henglein' in his outstanding work interpreted his results in terms of a diffusion-controlled electron transfer from the reducing species to colloidal Ag particles which act as electron sinks. On the other hand, Matheson, Lee, Meisel, and Pelizzetti? in their excellent work on hydrogen production from MV+ on colloidal Pt, concluded that the Pt was loaded with adsorbed hydrogen atoms, but did not rule out the electron-charging model for metals other than Pt, such as Ag and Au. Differences between the electron loading and H atom loading may be due to differences in the nature of the metals. The low overpotential of Pt, as compared to Ag, may favor a higher degree of protonation in the case of the Pt8J8 as compared to Ag. We have tested the two extreme cases: electron loading with no protonation, on one hand, and full protonation on the other, in the light of our results. Firstly we considered the electron-charging model. On the basis of this model, the fast decay of MV+ is attributed to the charging of the Ptc, until a charge Z is obtained and equilibrium is established (eq 2, where Pt stands for a colloidal particle). The
The Pt concentration also had an effect on the kinetics of the fast decay, as may be observed in Figure 5 . When [Pt] was increased, while the pH, [MVz+], and the pulse intensity were kept constant, the decay kinetics became faster. The effect of [Pt] on the pseudo-first-order rate constant of the fast decay at pH 8.1 is shown in Figure 7. When deviations from a first-order rate law were considerable, particularly with the higher pulse intensities, only the initial decay rates were taken. The log-log plot was used to derive the order with respect to [Pt]. From Figure 7 we concluded that, at the lower pulse intensities, the pseudofirst-order rate constant is proportional to [Pt]' while at the highest pulse intensity it is proportional to [Pt]' O. These results differ from our observation a t acidic pHs (Figure 8) where the PtZ- MV+ ~t Pt(Z+l)- MV2+ (2) fast decaying fraction was near to 1, and the rate law of the decay of absorbance was near to first order. In the first pulse the reaction rate was proportional to [FYI2,in agreement with earlier s t u d i e ~ . * ~ ~ ~ ~(18) A. Henglein, B. Lindig, and J. Westerhauser, J . Phys. Chem., 85, 1627 (1981). Usually our rate constants were found to be about 12-fold lower
+
+
1620 The Journal of Physical Chemistry, Vol. 88, No. 8, 1984
Brandeis et al.
a
J
1.5
--
I
1
1
J
-
I
I
I
parts of 'the
?
100
-1 0 -45
I
-40
-35
101
Number of pulses
I
-30
-25
102
-20
Iog CPtl 01 -4.5
I
- 4.0
I
-3.5
I
-3.0
Figure 9. log kobsd, measured in acidic solutions for the 100th pulse, plotted vs. log [Pt]. Conditions as in Figure 8. Insert: kobrdvs. the number of pulses plotted on a log-log scale.
I
-2.5
log CPtl Figure 8. log kOM,measured in acidic solutions, plotted vs. log [Pt]: Pt batch 2605 (50%larger in diameter), pH 1.5,0.2 M 2-propanol, 4 X lo4 M MV''; (e) 0.2-ps pulse, ( 0 ) 1.5-ps pulse. reactivities of the Ptz-species toward MV+ and MVZ+are expected to depend on 2. However, it is difficult to explain the large changes in the magnitude of the fast decaying fraction above pH 6 from nearly 1 to practically 0, on the basis of this model, because the redox potential of neither Ptz-nor MVZ+is expected to change much with pH. Note that the charge of the Pt colloids prior to the irradiation was negative (see Experimental Section), in the entire pH range investigated (1,5-12), and no other changes in the colloid properties (e.g., light scattering) were observbd in this pH range, so that we assume that the solutions contained the same colloidal particles in the entire pH range. Let us now consider the H adsorption model where the transfer of electrons is believed to be followed by fast protonation, according to eq 3. Equilibrium is achieved when the redox potential of the PtH,
+ MV+ + H+ e PtHZ+l + MV2+
(3)
loaded Ptc species equals that of the MV+ (eq 4). The value of E(PtHZ/PtHZ+,) = Eo(PtHZ/PtHz+l) ( R T / F ) In (l/[H+]) = E(MV2+/MV+) = Eo(MV2+/MV+) - ( R T / F ) In ([MV+]/[MV2+]) (4)
Z (independent of the mechanism) can be calculated from the amount of [MV'] which disappeared in the fast decay. From comparison of our kinetic measurements with previous ones8 and assuming that the rate of the reaction of MV+ with Pt is controlled by the number of Pt colloidal particles, we confirmed that our average diameter of the colloidal particles was about 50 A. A simple calculation shows that 2 was about 150, under our typical conditions ([MV'] = 1.5 X M, [Pt] = 4.0 X M). Assuming a normal distribution around the average Z , for large Z values, [PtH,] = [PtHt,,], so that the redox potential of the loaded Pt at a given Z value appears to depend only on pH (variation of 2 is expected to affect Eo(Pt)). Henceforth, E(Pt) and Eo(Pt) will represent the redox potential and the standard redox potential, respectively, of the loaded Pt, irrespective of the nature of the loading. In Figure 10 we present a plot of log ( [MV+],fd/[MVZ+],fd)vs. pH. Linear dependency is expected in the pH range 6-9. In this pH range (see Table I), most of the reducing equivalents were loaded on the Pt at the end of the fast decay. Above pH 9, Z sharply decreased, resulting in changes in Eo(Pt). This introduced nonlinearity into Figure 10. Since 2 depended also on the size of the colloidal particles, different batches may have somewhat different pH limits for the linear part of Figure 10. According to the H adsorption model, on the basis of eq 3 and 4, the slope of the line in Figure 10 is expected to be 1. In reality, the measured slope was 0.5. This means that the H adsorption model, although in qualitative agreement with the
The Journal of Physical Chemistry, Vol. 88, No. 8, 1984 1621
Reactions of Colloidal Pt in Aqueous Solutions
- 4' 4
I
I
I
I
6
I
8
I
I
10
I
12
PH
Figure 10. pH dependence of [MV'],,: 0.2 M 2-propano1, 4 X lo4 M Pt, 4 X lo4 M MVZ+, 1.5-ps pulse. The straight line corresponds to the loading of Pt with hydride ions. Pt batches 2408 and 3012.
effect of pH, fails to account quantitatively for it. We invoke partial protonation of the charged Pt. A slope of 0.5 (Figure 10) means that only one-half of the loaded electrons were neutralized by protonation, according to eq 5 . The proMV+ MV+
+ PtH,"-
+ PtH,("+')-
+ PtH,("+')MV2+ + PtH(,+l)("+')-
G MV2+ H+
(5a) (5b)
tonation step is assumed to be fast as compared with the transfer of electrons from MV+ to the Pt. The average number of loaded equivalents, Z , equals 2n. In Figure 11 we present the [MV'] either at the end of the fast decay or at the end of the buildup process, whenever observed ([MV+],,,), as a function of pH. The circles are the experimental observations (taken from Figure 2 and Table I , together with results obtained with other Pt batches). The curves demonstrate the differences between the H-adsorbed and partial-protonation models. The solid curve on the left side represents the expected results if "true" catalysts took place, namely, if no significant concentrations of loaded Pt were present. In view of the constant ratio of 2 between Z and n over a relatively wide range of pH, we are tempted to suggest that this ratio of 2 represents hydrogen adsorbed in the form of a hydride ion (see eq 6). The standard redox potential of the Pt colloidal PtH,"-
+ H+ + 2e- .e PtH(,+l)("+l)-
(6)
particles loaded with H, Eo(Pt), as defined by eq 7 can be calE(Pt) = Eo(Pt) - ( R T / 2 F ) In ([PtH(,+l)(fl++l)-]/[PtH,"] [H+]) = Eo(Pt) - ( R T / 2 F ) In ( l / [ H + ] )
(7)
culated from Figure 10, extrapolating the straight line to = 1. This yields Eo(Pt) = -0.082 f 0.008 [MV+lerd/[MV2+lefd V, based on Eo(MV2+/MV+) = -0.446 V (both values are vs. "E). In the calculation of Eo(Pt) we again assumed that, for large n, [Pt,"] = [Pt(,+l)("+')-]. Eo(Pt) was constant at pH 6-9 where there were only moderate changes in Z. Results below pH 6 are inaccurate because the fast decaying fraction was too close to 1, and a small residual absorption introduced a large error. Above pH 9, Z decreased quickly with pH and Eo(Pt) is not expected to remain constant. The storage of the reducing equivalents in the form of hydride ions can be considered in the context of the more general hypothesis, according to which adsorbed e-, H, and H- are always formed together and the relative abundance of each species de-
Figure 11. Equilibrium concentrations as a function of pH. Experimental conditions as in Figure 10. Circles: experimental results for MV+. Solid lines: (right) best empirical fit with the experimental results; (left) M V + calculated from the redox potentials of methylviologen and hydrogen, assuming that the Pt was a "true" catalyset (no loading). Dots: the concentration of hydride ions loaded on the Pt (n[Ptc]). Dashed lines: initial and equilibrium [HJ.
pends on the nature of the metals.19 According to this view Heven plays a role at colloidal silver particles.20 The presence of both adsorbed e- and adsorbed H has already been proposed for gold.21 The hydride model is also in agreement with the buildup reactions observed in the alkaline solutions (Figure 2 ) . The buildup is due to the loading of Pt (in the form of H-) by H2, followed by reduction of MV2+to MV+. As already mentioned above, the G value of H2 produced during the pulse is around 1.O under these conditions. Our results show (Table I) that, in the pH range 6-8, almost all the reducing species are in the form of PtH,". At pH 10 and above, MV+ is the predominant species, with only traces of PtH," and H2. The pH abundance curves of the other reducing species are also shown in Figure 1 1. The relative concentrations of PtH,", H,, and MV+ vary with pH because the redox potentials of these systems are differently affected by pH: the redox potential of PtH,fl-/PtH~,+l)("+l)-becomes more negative by 0.03 V per pH unit, that of H+/H2by 0.06 V per pH unit, and that of MV2+/MV+is pH independent. The standard redox potentials are such that at stronger acid media (under our conditions, at pH 4) all reducing species are practically converted to H2, in the intermediate pH range (6-8) they are in the form of PtH,", and in the stronger alkaline pH range as MV'. In conclusion, we observed two equilibrium reactions. The faster one, in the millisecond time range, under our conditions (typical [Pt] = 4 X M), is the reaction of Pt with the reduced methylviologen species. The slower process, the buildup in the time range of 100 ms, was due to equilibration of the system with H2. This is observed in the stronger alkaline solutions. The nature of the pH-dependent slow decay (taking place within minutes in the stronger alkaline solutions down to seconds at the (19) A. Henglein, private communication. (20) A. Henglein, Ber. Bunsenges. Phys. Chem., 84, 253 (1980). (21) J. Westerhauser, A. Henglein, and J. Lilie, Ber. Bunsenges. Phys. Chem., 85, 182 (1982).
1622 The Journal of Physical Chemistry, Vol. 88, No. 8, 1984
lowest alkaline pHs) is not clear. It seems to involve the molecular H202,which slowly oxidizes MV' even in the absence of Pt. Pt may catalyze this reaction by decomposing H 2 0 2to water and 02,a well-known reaction. Thus, both O2and H202may oxidize MV' either directly or through partial unloading of the Pt, enabling additional loading by MV'. If this is so, the effect of pH can be qualitatively explained if the unloading of the Pt is faster than both the direct oxidation of MV+ by H202and the decomposition of H,02 to water and 02.The faster reaction of PtH," with Hz02at the lower pHs, where n is relatively large, may take place, at least in part, simultaneously with the H2 reaction. The effects of pH, [Pt], [MV2+],and pusle intensity on the fast decaying process suggest the following: (a) Process 5 is an equilibrium process, where the rate of equilibration depends on the rates of both forward and backward processes. The observed rates and reaction orders will depend on the relative contributions of the forward and backward reactions. (b) The rate of loading is expected to become gradually slower when n approaches its upper values. A quantitative model for these effects is beyond the scope of this manuscript. The effect of accumulated pulses is undoubtedly connected with the accumulation of H, and the building up of P t H R with higher n values. At slightly alkaline solutions, the effect of repetitive pulses was similar to the effect of increasing pH (Figure 2 ) ; namely, the fast decaying fraction became gradually smaller and its rate of decay slowed down. The increase of n made the Eo(Pt) more negative until no additional loading takes place. This is M Hz was initially supported by an experiment where 7 X present at pH 8. The fast decaying fraction, which usually occurred following the first pulse, was not observed. This can be explained if the prior addition of H2 loaded the Pt with a sufficiently high n which prevented additional loading by the MV+. The effect of [Pt] on increasing the fast decaying fraction (Figure 5) is in agreement with the hydride, as well as the H , loading models. Under the conditions of our experiments the concentration effects of MVZ+and MV+ can also be explained in terms of our model: The concentrations of both Pt and MV' can be considered quite constant while that of MV+ changed from the initial value to the equilibrium value. The loading reaction by MV+ (eq 4) is expected to be pseudo first order in [MV'], while the back-reaction (unloading by MV2+) is expected to be of zero order as neither [MVz'] nor [PtH(,+I)("+')-]changed much during the reaction under our conditions. However, the loading and unloading reactions are of a particular nature; namely, although [Pt] did not change with time, n did. The reactivity of PtH,"toward MV+ and MVZ+may, of course, depend on n: with increasing n, the reactivity toward MV+ may decrease, and the reactivity toward MV2+ may increase. Therefore, no simple dependence of the rate of the fast decay on either [MV'] or [MV2+] is expected. The order of the fast decay with respect to [Pt] was 1.00 f 0.15 and 1.35 f 0.1 for initial [MV+] = 1.5 X and (5-7.5) X lo6 M, respectively. At the lowest [MV+], the fraction of MV+ decaying in the fast process approached 1 at the highest Pt concentrations, and the decay kinetics obeyed a first-order rate law with respect to [MV']. This differs from the high [MV'], where deviations from first-order rate law were considerable. In the latter case only initial rates were analyzed. This effect of [Pt] will be further discussed together with the results in acidic pHs. Discussion of the Results in Acidic p H . In general, our results agreed with the earlier rep~rts.~,'JThe rate of the MV' decay was slowed down considerably (Figures 8 and 9) with increasing of the number of pulses given prior to the test pulse. However, the fast decaying fraction was nearly 1 at all pHs below 6 and was not affected by the number of pulses given. This is not surprising as, at pH 6 and below, the predominant reducing species are loaded Pt and H2. An unexpected "residual absorbance", of up to 5% of the initial MV+ absorbance, was observed. We attribute this to changes in light scattering by the Pt colloidal particles. Indeed, in acidic solutions, occasionally we could observe, by eye, precipitation of the Pt after 200 pulses. This precipitation was presumably caused by the catalytic reduction of the poly(viny1
Brandeis et al. alcohol) stabilizer, as hydrogen accumulated in the system. This hypothesis was supported by an experiment where an aqueous solution containing excess of poly(viny1 alcohol) was pretreated with hydrogen in the presence of Pt colloids. The Pt and the hydrogen were removed and the remaining polymer was tested and found uneffective in the stabilization of a freshly prepared Pt colloid. In the initial stages of the precipitation, the amount of light which reached the photomultiplier decreased due to aggregation and was observed as the residual absorbance mentioned above. We found (Figures 8 and 9) that the rate of MV' decay was second order in [Pt], using first pulses, changing to first order dependency on [Pt] a t the 100th 1.5-ps pulse, or to an order of 1.3 at the 100th pulse of 0.2 p s . The second order in [Pt] is in agreement with earlier observation^.^^^-^ Matheson et a1.,8 using their H-adsorbed model, attributed the [PtI2 dependency to a collision between two H-loaded Pt colloidal particles as the rate-determining step in the reaction of Pt with MV+. According to this model, Pt quickly becomes loaded with adsorbed H, so that no additional sites are available for further reaction, unless unloading takes place by collision of two loaded Pt particles. We suggest that, in the stronger acidic solutions, the platinum hydride reacted quickly with H+ and adsorbed H or H, was produced. It is not possible, on the basis of the presently available results, to distinguish between adsorbed H and adsorbed H,. Thus, the H (or H,) adsorption model is possibly valid for the stronger acidic solutions. This is supported by conductivity measurements showing that the conductivity decayed back to its prepulse value8 while the MV' decayed away. Had the hydride been stable in the acidic solutions, a net residual conductivity amounting to half of the conductivity change by the pulse would have remained, in contrast to the observations.* No conductivity data are available for neutral and alkaline pHs. The suggestion that the R unloading takes place by collisions8seems to be less certain. It is clear from our results (see Figure 2 ) that 4 X IO4 M Pt can be loaded with up to at M reducing species. This can be concluded from least 1.5 X the fast decaying fraction approaching 1 at initial [MV+] = 1.5 X M, at pHs such as 6.1 and 7.3, where no free H2 formation is expected (see Table I). Therefore, as long as loading alone occurs, its rate cannot be controlled by an unloading process. These conditions exist in the pH range above 6 and at the initial stages of the MV' decays at all pulse intensities even under the strong acidic conditions. The assumption that the Pt unloading takes place by collision would require that at sufficiently low [ MV'], before unloading becomes the rate-determining step, the decay of MV' will be faster. Such a fast process was not observed. Moreover, some of our results at pH 8.1 also indicate a reaction order of 1 in Pt. An order greater than 1 in [Pt] at this pH, where the end product is loaded Pt, cannot be attributed to an unloading process. The decaying order of Pt in acid solutions became first order in [Pt] after 100 pulses. This is also difficult to understand if the Pt becomes completely unloaded after every pulse. On the other hand, we cannot advance an explanation for a second-order dependency on [Pt], unless a collision between two colloidal Pt particles takes place in a way that affects the rate-determining step of the loading process. It is conceivable that the colloid stabilizer (in our case, poly(viny1 alcohol)) closely surrounded the Pt particles and inhibited their reaction with MV'. A prior collision between two particles, or between the stabilizing polymer molecules, may have produced a "naked" Pt particle, so that part of the surface would become available for an efficient reaction with the MV' ion. Consequently, we observed a second-order dependency of the MV+ decay on [Pt] in acid and near neutral solutions. At the higher pHs, where the back-reaction of loaded Pt with MV2+ was also important, the overall order of the equilibration rate was less than 2. If this interpretation is correct, the results at the higher pHs may mean that the back-reaction of the loaded Pt particles with MV2+ was not second order in Pt. A possible reason for this is the relatively high concentration of MVZ+. In the bulk of the solutions, the [MV2+] was usually 4 X M, as compared with 1.5 X M MV'. Moreover, the colloidal particles which were negatively charged, even before loading took place, probably had some of the doubly positive MV2'
J . Phys. Chem. 1984,88. 1623-1627 ions in their ionic atmosphere. Therefore, the concentration of MVZ+near a colloidal particle was even higher. The reaction of MV2+ with loaded Pt may be first order in Pt, if the reaction proceeds sufficiently fast without the enhancement by the particles' collisions. The effect of repetitive pulses described above is in agreement with this interpretation. Destabilization of the Pt colloidal particles resulted in an increase in the size of the particles, and weaker interaction of the Pt particles with the polymer, resulting in more naked areas on Pt particles. Hence, the reactivity of the Pt toward the MV+ decreases and the rate of reaction gradually changes from second- to first-order dependency on Pt, with increasing the number of repetitive pulses. Hydrogenation of the methylviologen species by Hz in the presence of Pt colloids has been r e p ~ r t e d . This ~ ~ ~process ~ ~ was reported to be relatively slow and is not expected to affect our results when single pulses are applied. Upon application of repetitive pulses, hydrogenation gradually becomes important, and indeed, after about 200 pulses (1.5 p s ) , no MV+ formation could be observed following electron pulses. This is attributed to full hydrogenation of MV2+. Steady-State Experiments. When hydrogen is bubbled through solutions containing MV2+ (4 X M) no blue color, which might be attributed to MV+, was observed (pH 5-1 1). However, when Pt (5 X 10"-4 X 104M) was added to a solution containing both MV2+ and HZ,the typical blue color of MV+ immediately developed (pH 8.2). The absorption spectrum of the solution was exactly the same as the previously reported spectrum of MV+,s,16 and the one observed in a solution containing 0.2 M 2-propanol M MV2+ which was irradiated with a y source. and 4.0 X The intensity of the blue color increased with pH and seems to have decreased at higher Pt concentrations (no quantitative measurements were carried out). The blue color of MV' disappeared quickly by bubbling with ultrapure H e and reappeared upon rebubbling HZ. This cycle could be carried out many times. These experiments, although qualitative in nature, demonstrate that all the reactions in this system are reversible, and that the ~
~
~~
~~
(22) P. Keller and A. Moradpour, J . Am. Chem. SOC.,102, 7193 (1980), and references cited therein. (23) A. Launikonis, J. W. Loder, A. W.-H. Mau, W. H. F. Sasse, and D. Wells, Isr. J. Chem., 22, 158 (1982), and references cited therein.
1623
Pt colloidal particles can be loaded, both by MVf and by H,, and can also be unloaded by MVZf. When acidic solutions were saturated with H2, precipitation of the Pt occurred after about 10 min. All these results are in full agreement with the pulse radiolytic data and their interpretation. An interesting result was obtained in a solution containing Pt colloidal particles (4 X M Pt) in the absence of methylviologen. Helium gas was bubbled to eliminate air and the pH adjusted to 10.1. Then, H 2 was bubbled, resulting in a decrease of the pH to 9.0, within several minutes. This experiment is in qualitative agreement with the hydride model, because the formation of H+ must be connected with the formation of a net electrical charge on the Pt. On the other hand, loading of the Pt with H, or H is not expected to affect the pH. Control experiments showed that neither the Pt present alone (bubbled with He instead of H,) nor H2in the absence of Pt caused a decrease in pH. Further work is in progress on the equilibria involving colloidal Pt, using steady-state and flow techniques. Relation with the NHE. A question may arise whether the hypothesis that Pt colloidal particles become loaded with hydride ions, when in contact with H,, is not in contrast with the wellknown properties of the hydrogen electrode. In the hydrogen half-cell, the redox potential changes by -0.0592 V with every pH unit, which corresponds to the reduction of H+ to H2 and not to H-. The answer to this question is that we are dealing with different systems. The effective number of sites available for Hon the surface and inside a Pt electrode is by several orders of magnitude smaller than in the colloidal systems. Therefore, no significant amounts of H- may be present at the electrode, and the predominant reducing species are H 2 molecules (which are at equilibrium with Hz or H dissolved in the Pt). Therefore, irrespective of the mechanism, the redox potential is that of H+/H2. On the other hand, the Pt colloid, when present at sufficiently high concentrations, may act as a storage of hydride ions (at the appropriate pHs). Acknowledgment. We are indebted to Professor V. Barboy for most useful discussions. This research was supported by the National Council for Research and Development (Israel), and by the Schreiber Foundation. Registry No. Pt, 7440-06-4; H2, 1333-74-0;H202,7722-84-1; MV2+, 4685-14-7;MV', 25239-55-8; 2-propano1, 67-63-0; water, 7732-18-5.
Ion-Exchange Kinetics on Ion Exchanger Using the Pressure-Jump Technique. 1. Protonation-Deprotonation and CI- Adsorption-Desorption on Diethylaminoethyl Group of Sephadex A-25(CI) Kazuaki Hachiya, Minoru Sasaki, Yoshiharu Nabeshima, Naoki Mikami, and Tatsuya Yasunaga* Department of Chemistry, Faculty of Science, Hiroshima University, Hiroshima 730, Japan (Received: March 23, 1983; In Final Form: August 30, 1983) Single relaxation was found in an aqueous suspension of Sephadex A-25(C1), which has a diethylaminoethyl group as an anion-exchange site (at pH 4.2), by using the pressure-jump technique with conductivity detection. In the acidified suspension of Sephadex A-25(CI) in the pH range of 2.5-3.2, another slow relaxation was also found. Fast relaxation time decreases with Cl- concentration, while slow relaxation time increases slightly with Hf concentration. Taking into account the electrostatic potentials created by the CI- and H+adsorbed, the fast and slow relaxations were attributed to the C1- adsorption-desorption on a protonated diethylaminoethyl group and the protonation-deprotonation of the diethylaminoethyl group, respectively. The intrinsic values of the protonation and deprotonation rate constants were determined to be (1.7 f 0.7) X 1O4 mol-' dm3 s-' and (1.8 i 0.8) X lo-' s-l, and those of the C1- adsorption and desorption rate constants (9 f 2) X lo3 mol-' dm3 s-' and 8 zt 5 s-', respectively, at 25 OC. Introduction Anion and cation exchangers derived from Sephadex and cellulose are useful materials for the analysis, separation, and preparation of various biochemical materials such as amino acids, 0022-3654/84/2088-1623$01.50/0
proteins, and nucleic acids.'-3 The dynamic properties of the ion-exchangeable site existing on the surface Of these ion ex(1) Peterson, E. A.; Sober, H. A. J . Am. Chem. Sor. 1956, 78,751, 156.
0 1984 American Chemical Society
