1023
EFFECTOF NORMALALIPHATIC ALCOHOLSON ELECTRODE KINETICS
The Effect of Normal Aliphatic Alcohols on Electrode Kinetics by Katsumi Kanzaki Niki’ and Norman Hackerman Department of Chemistry, The University of Texas, Austin, Texas
78711
(Receiwed August 1 0 , 1968)
The applicability of a potential-step method for measuring the current-time variation in the electrode reaction of the V(I1)-V(II1) couple on the hanging mercury drop electrode in 1 N H2S04 was verified both in the absence and in the presence of alcohols. The effective electrode area did not change significantly with the alcohols adsorbed on the surface. This indicated mobility of the adsorbed molecules. The polarographic limiting current was attained before the desorption potential of the alcohols was reached, and it remained constant even in the presence of the adsorbate. This is explained by the formation of a uniform barrier on the electrode. The rate constants of the electrode reaction decreased with increasing alcohol concentration, but the dependence of reaction rate on coverage was not linear. Traube’s rule was found to be applicable, with a factor of one-fourth. The effect of coverage on the reaction rates by alcohols is explained in terms of the orientation of the alcohols. KMn04, using a Jones reductor in the usual analytical It is known that surface-active substances often method.12 All solutions were prepared from twicehave a marked effect on electrode processes. I n spite distilled water and were purified by activated charof numerous studies on adsorption of surface-active ~0al.13 For each experiment VOSO4 was reduced with substances on electrodes, there is still no adequate stirring on the mercury pool cathode t o V(II1) and theory of the effect of adsorbed molecules on electrode then to V(I1). The concentrations of both V(II1) processes. Early works in polarography on the effect and V(I1) were determined by polarography, and the of surface-active substances were carried out extenequilibrium potentials were measured by the hanging sively by Russian and Czech investigators.2 They mercury drop electrode as an indicator electrode. The assumed that the adsorbed layer of surface-active subvanadic-vanadous solution was maintained under nistances merely influenced the electrochemical process a t the electrode surface.2b Diffusion of reactants to trogen. Nitrogen was purified by bubbling through a concentrated vanadous sulfate solution, then through positions next to the adsorbed layer on the electrode was not taken into account, There are several papers3-’ a solution identical with that being studied, then to the electrolytic cell. This minimized vapor losses of on electrode kinetics in the presence of neutral surfacewater and alcohols from the electrolytic cell. RIeasureactive substances, and a linear relationship between ments were made a t 25.5 k 1.0’. rate constant and fractional coverage was demonstrated in most cases. A Icohols. Certified n-butyl alcohol and n-amyl alcoThis work uses a potential-step method to study the effect of neutral organic substances on electrode ki(1) Sumitomo Chemical Do., Minato-ku, Nagoya, Japan. netics. The method permits elucidation of the mode (2) For instance: (a) For a review in polarography: C. N. Reilley and W. Stumm, “Progress in Polarography,” Vol. I, P . Zuman, of diffusion of the reacting species to the electrode. Ed., Interscience, hTew York, N. Y.. 1962, p p 81-121: (b) A. N. I n addition to the potential-step procedure, polarogFrumkin, Dokl. Akad. Nauk S S S R , 8 5 , 373 (1952); Proc. Intern. Congr. Surface Activity, Rnd, London, 1957, 58 (1957): (c) J. Heyraphy, capillary electrometry, and cyclic voltammetry rovsky, F. Sorm, and J. Forjt, Collect. Czech. Chem. Commun., 12, were used. To simplify the study, the V(I1)-V(II1) 11 (1947); (d) M . Matyas, { b i d . , 16, 496 (1951). couple was selected because the reaction rate c o ~ i s t a n t ~ ~(3)~ P. Delahay and I. Trachtenberg, J. Amer. Chem. Soc., 80, 2094 (1958). at the standard potential is well within the range of (4) W. Lorene and W. Muller, 2. Phys. Chem., (Frankfurt a m Main) 18, 141 (1958). the application of a potential-step method. Also, vana(5) L. Gierst, “Advances in Electrochemistry and Electrochemical dium ions are said not to adsorb specifically on a Engineering,” Vol. I , P. Delahay, Ed.. Interscience. New York, mercury electrode,1° and the standard potential of N. Y., 1960, pp 55, 94. this system is close to the potential of the electro(6) W. Muller and W. Lorenz, Z . Phys. Chem. (Frankfurt a m Main) 27, 23 (1961). capillary maximum for mercury in the absence of (7) A. Aramata and P. Delahay, J. Phys. Chem., 68, 880 (1964). specific adsorption, so that the coverage by the ad(8) J. E . B. Randles, Can. J. Chem., 37, 238 (1959). sorbed molecules hardly varies around the standard (9) J. E. B. Randles, “Progress in Polarography,” Vol. I, P. Zuman and I. M. Kolthoff, Ed., Interscience, New York, N. Y., 1962, p 123. potential. Experimental Section
Solutions. The clear blue solution of VOSOl was prepared according to the method of Johns and Colvin.ll The concentration of VOSO4 was determined with
(10) K. M. Joshi, W. Mehl, and P. Parsons, Trans. Sump. EZectrode Processes, Philadelphia, P a . , 1959, 249 (1961). (11) G. Johns and J. N . Colvin, J. Amer. Chem. Soc., 66, 1563 (1944). (12) For example: G. H. Ayres, ‘$QuantitativeChemical Analysis,” Harper & Brothers, New York, N. Y., 1958, p 451. (13) G. Baker, ref 10, p 251.
Volume 73, Number 4 April 1969
1024 hol (Fisher Scientific Co.) were used without further purification, since impurities were hardly detectable by vapor phase chromatography. The n-hexyl alcohol was prepared from practical grade n-hexyl acetate; it was then purified by vapor phase chromatography. The concentrations of alcohols were chosen so that the same amount of each alcohol would be adsorbed at the mercury solution interface, according to Hansen and Craig.14 The relative concentrations (C/C,) , rather than C, are used. C, is the solubility of the alcohol in water, and the values of C , were taken from Hansen, et aZ.,16and Seidell.I6 Electrolytic Cell. The all-glass three Compartment electrolytic cell was used for all experiments. Compartments were separated from one another by fine sintered-glass membranes. One compartment contained the hanging mercury drop electrode, the mercury pool, and the dropping mercury electrode. A compartment on one side contained the anode, and one on the other was a bridge to the Hg,, HgnSOa, 1 N H2S04 reference electrode. The hanging mercury drop electrode made by Metrohm, Ltd., Switzerland, was used as the working electrode. The capillary of the hanging mercury drop electrode was silicladed. The mercury drop (maintained 10 sec in the solution) was renewed each time. Mercury was distilled twice under vacuum. Potentiostat. A Wenking 61 RS electronic potentiostat with an external reference potential eource was used. The working electrode was negatively polarized in all experiments by applying a 10-mV step over the equilibrium potential. Polarography and Cyclic Voltammetry. The linear scanning potentiostat for polarography and the cyclic triangular wave generator for cyclic voltammetry used a Philbrick Research Model 6009 operational amplifier. The scan rate for polarography was about 46 mJ7/min, and the mercury drop life was about 7 see. The sweep rate of the cyclic voltammetry was varied from 67 mV/sec to 150 V/sec. Capillary Electrometer. The capillary electrometer was constructed according to the method described by Hansen, Kelsh, and Granthan.” Measurements were made in 1 N HzSOd solution and with several concentrations of alcohols in 1 N H2S04 solution in the absence of the vanadium ions to achieve more nearly the conditions corresponding to an ideally polarized electrode. The radius of the capillary a t the null point was determined by standardization against a solution of known surface tension.18 All these measurements were also made a t 25.5 f 1.0’.
Results Polarographic Studies. The V (11)-V (111) couple showed a quasi-reversible polarographic wave and gave simple reduction and oxidation waves in 1 N HzS04. Wit,h increasing alcohol concentration, the polaroThe Journal of Physical Chemistry
KATSUMIKANZAKI NIKI AND NORMANHACKERMAN graphic wave became more irreversible. However, the diffusion currents observed were more positive than the displacement potential of the alcohol, and there was no visible effect on the diffusion current on addition of n-amyl and n-hexyl alcohols but n-butyl alcohol lowered it up to 4% a t higher concentrations. The diffusion coefficients of V(I1) and V(III), determined polarographically by using the Ilkovic equacm2/sec, retion, were for 8.0 X 10-6 and 6.5 X spectively. The standard potential of the V (11)V(II1) couple was -0.953 f 0.001 V. Xone of these alcohols had any effect on the equilibrium potential. The standard electrochemical reaction rate constant, calculated from polarographic data as done by Randlesls was (1.04f 0.03)X le3cm/sec, and the transfer coefficient, a, was 0.52. A t higher alcohol concentra-
tI
I
I
I
I
-0.90 -1.00 -1.10 -1.20 Volts (V 8 . ‘1 N HzSOd, HgiSOd/Hg)
I
-1.30
Figure 1. Effect of alcohols on the charge-transfer rate constant of the V(II1) e 3 V(I1) in 1 N HZSOa obtained from polarography: * , without alcohol; -0-0-, 34.4 mM (C/C. = 3.5 X n-butyl alcohol; -0---O-, 68.8 mM (C/C. = 7.0 X 10-2) n-butyl alcohol; -0---0-,172 mM ( C / C , = 17.5 X n-butyl alcohol; -X-X-, 10 mM ( C / C , = 3.5 X 10-2) n-amyl alcohol; -X-*-X-, 20 mM (C/C. = 7.0 X IOm2)n-amyl alcohol; -X--X-, 50 mM (C/C. = 17.5 X 10-2) n-amyl alcohol; -.-e-, 2.18 m M (C/C. = 3.5 X n-hexyl alcohol; 4.36 mM (C/C. = 7 . 0 X 10-2) n-hexyl alcohol; -0- -0-,10.9 mM (C/C. = 1 7 . 5 X 10-2) n-hexyl alcohol.
.
+
-.---.-,
(14) R. S. Hansen and R . P. Craig, J . Phys. Chem., 58, 211 (1954). (15) R. 8 . Hansen, Y. Fu, and F. B. Bartell, ibtd., 53, 769 (1949). (15) A. Seidell, “Solubilities of Organic Compounds,” D. Van Nostrand Co., Princeton, N . J., 1941. (17) R. 9. Hansen, D. J. Kelsh, and D. H. Granthan, J . Phys. Chem., 6 7 , 2316 (1963). (18) C. A. Smolders and E. M . Duyvis, Rec. Tran. Chim., 80, 635 (1961).
EFFECTOF NORMAL ALIPHATICALCOHOLSON ELECTRODE KINETICS
1025
tions, the V(I1)-V(II1) system may be considered t o be totally irreversible. Thus the reaction rate constant cmi be calculated in the usual manner.lg The results are shown in Figure 1. Potential- Step Method: the Electrode Kinetic Parameter of the V ( I I ) - V ( I I I ) Couple. The current-time relationship for the V (11)-V (111) oxidation-reduction couple for mass transfer controlled by semi-infinite linear diffusion given by
It
=
A F[kfC*v(III)
Q
- kbC*V(II)]
= [kf/D1”V(III)
+
exp(Q2t)erfc(Qt1/2) (1) kb/D’12V(II)]
where I , is the current a t the time, t, elapsed from electrolysis beginning; F, the Faraday; A , the electrode area; kr and k b , the reaction rate constants for the cathodic and the anodic reactions, respectively; C*V(III) and C*V(II),the bulk concentrations of V(II1) and V ( I I ) , respectively; and DV(III)and Dv(II),the diffusion coefficients. When Q P 0. For the V(II)-T7(III) couple in the presence of alcohol, the apparent rate constant, f ( r )kr, is small. Thus, the polarographic wave becomes more irreversible with increasing alcohol concentration. As shown by polarographic measurements the limiting current is not affected by the alcohols, probably because the effective electrode area is unchanged by coverage. The slight change of limiting current a t higher concentrations of n-butyl alcohol is likely due to change of bulk properties of the solution by the alcohol, since a much more concentrated solution of n-butyl alcohol is required to attain the same relative concentrations used for the other alcohols. If k” is very large, as in the reduction of Ag+, even though f(r) is very small, f ( r )k rmay still be large enough to be polarographically reversible, as was observed by 31atyas.2d When organic molecules are strongly adsorbed on the electrode surface, the adsorbed molecule may become immobile; accordingly, the polargraphic wave becomes more irreversible and the limiting current becomes smaller. Effect on Reaction Kinetics of the Time Interval between the Formation of a Mercury Drop and Initiation of the Step Function of the Potential. The reaction rate of the system V (111-Y (111) drops markedly a t higher concentrations of the alcohols.z6 This is true also even a t very low alcohol conceritrations when the mercury drop is maintained a t the equilibrium potential for a long time (5-20 min) before applying added voltage. Electrocapillary measurements in the system with the same concentration of alcohol, but without V(I1) and V ( I I I ) , show surface coverage of lO-ZO% and an almost instantaneous attainment of adsorption equilibrium. Since the inhibition effect is accelerated by increasing the concentration of V(I1) and also by stirring the solution, the inhibition may be associated with V(I1) ions, such as by preferential adsorption of the ions with the alcohol molecules or a reaction inter(22) G. Torsi and P. Delahay, “Double Layer and Electrode Kinetics,” P. Delahay, Ed., Interscience, New York, N. Y., 1965, p 233. (23) T. Biegler and H. A. Laitinen. J. Electrochem. Sac., 113, 852 (1966).
(24) J. H . de Boer, “Physical Chemistry of Surfaces,” A. W. Adamson, E d . , Interscience, New York, N. Y . , 1960, p 461. (25) According to E . Blomgren. J. O’M. Bockris, and C. Jesch, J. P h y s . Chem., 65, 2000 (1961), the heat of displacement of BuOH is 3.7 kcal/mol a t 8 = 0.25. (26) At a higher ratio of CV(II)/CV(III)(especially greater than 1. When the ratio of CV(II)/CV(III,is less than 1, this type of inhibition effect is quite small.) Volume 76,Number 4
April 1969
1028
KATSUMIKANZAKI NIKI AND NORMANI~ACKERMAN ~
Table 11: Surface Excess of the Alcohols at Mercury-Solution Interfaces --Butyl
Concn, m M
alcohol-
-------?+Amyl alcoholConcn, m M 10-10 mol/cmn
10-10 mol/cmz
6.9 17.9 34.4 69
0.5 1.3 2.3 3.9
2.0 5.0 10 20
mediate on the electrode. The stirring effect suggests a diffusion-controlled step. Although the nature of these effects is still not clear, no further study has been made of this. Cyclic Voltanametry. The shift of cathodic potential in the cathodic direction with increasing sweep rate in the presence of alcohol reflects the presence of adsorbed molecules on the electrode making the electrode reaction irreversible. On the other hand, adsorbed moleculcs may desorb completely from the electrode a t the potential where anodic peak appeared. Hence, the anodic reaction rates are large enough to be reversible, and the anodic peak potential would not shift with increasing sweep rate. Traube's Rule for the Adsorption of Alcohols and Reaction Kinetics. As othersz7observed, Traube's rule is applicable to adsorption of normal fatty acids (acetic to caproic) on mercury from 1 N Na2S04. The present electrocapillary measurements also showed that adsorption of n-butyl, n-amyl, and n-hexyl alcohol from 1 N HzS04 on mercury follows Traube's rule. The coefficient was about one-fourth as shown in Table 11. The solubilities of these alcohols in decreased by one-quarter for each additional CH2group. Therefore, the concept of the relative concentration of the homologous series of normal fatty alcohols proposed by Hansen, Fu, and B~irtell'~ as an extended Traube's rule may be equivalent t o the original Traube's rule. The model of t ~ 7 ocondensers in parallel is assumed as the double layer structure and the lateral interaction between adsorbed molecules and other particles on the
Table 111: Apparent Fractional Coverages" of the Electrode in the Presence of Alcohols
S.
Without alcohol 13.5 mM n-butyl alcohol 3.5 mM n-amyl alcohol 0.85 mM n-hexyl alcohol
x
100 56 47 38
(100 = S'),
%
S', %
0 44 50 56
100 56
50 44
-4ssumed that the apparent fractional coverage by n-butyl alcohol is 44% and the alcohol molecule lies flat on the electrode. lpparent fractional coverage (100 - B')%
- projected area of normal aliphatic alcohol x 44% projected area of n-butyl alcohol
The Journal of Phuaical Chemiatry
-?%-Hexyl Concn, m M
0.7
alcohol10-10 mol/cm*
0.424 1.07 2.13 4.3
1.6 2.8 4.3
0.6 1.5 2.6 4.0
80
20
20
Surface Coverage e
Figure 4. Influence of coverage on rate constant: 0-0, n-butyl alcohol; X-X, n-amyl alcohol; 0-0, n-hexyl alcohol.
electrode is neglected in order to simplify the discussion. When the adsorbed molecules lie flat on the electrode a t lower concentrations of alcohols, the eff ecLive area of the electrode decreases regularly with increasing number of CHZ group if the surface excess of the alcohols stays constant. This effect i s observed experimentally as shown in Figure 4 and Table 111. Here the ratio of the available electrode area, X, is assumed to be given by alcohol) & f = -k" (with____le" (without alcohol)
x
100%
(5)
Hence the fractional coverage was (100 - 8 )o/;;. In medium alcohol concentrations, the reaction rate does not change regularly with number of CHZ groups a t a given relative concentration of the alcohol, i.e., a t a given surface excess. With increasing alcohol concentration, the reaction rate constants of the V (11)V(II1) couple with different alcohols tend to be congruent with each other. The effect of the alcohols a t higher concentrations may be explained in terms of the orientation of the alcohols. Oriented parallel to the surface, n-butyl alcohol exhibits the smallest area, and n-hexyl alcohol the highest. The inhibition effects on the rate constant should follow this order for a (27) R . I. Kaganovich, V. AT. Gerovich, A k a d . Nauk SSSR. 155, 893 (1964).
and T. G.
Osotova, Dokl.
givcii amount of adsorbed alcohols on the mercury electrode. In more nearly perpendicular orientation, the projected area of normal aliphatic alcohols on the mercury electrode is about the same. Thus, the inhibition effect in this orientation is proportional to the amount of adsorbed alcohols on the mercury and not on the alcohol itself. The deviation of the inhibition effect from regularity in the medium concentration range is explained by some perpendicular orientation and some parallel. With increasing alcohol concentration, it is reasonable that perpendicular orientation predominates; therefore, the rate constants in the presence of the different alcohols become super-
imposable on each other a t a given relative concentration of the alcohol. dcknowledgment.
We deeply appreciate support of
this work by The Robert A. Welch Foundation of Houston, Texas, and by the Office of Naval Research, Contract NONR375( 15). The authors also express their thanks to Dr. Allen J. Bard and Mr. David Jones for their extensive discussion and reading, and t o many others for their helpful discussion. One of the authors (K. K. N.) also thanks Sumitomo Chemical Company of Japan for the leave of absence which permitted him to undertake this work.
Far-Ultraviolet Spectra of Hydrogen and Hydroxyl Radicals from Pulse Radiolysis of Aqueous Solutions. Direct Measurement of the Rate of H
+ H1
by P. Paysberg, H. Christensen,28J. Rabani,2b G. Ni1sson,2aJ. Fenger, and S. 0. Nielsen Danish Alomic Energy Commission Research Establishment. RisB, Denmark, and AktieboEagel Atornenergi, Sludsuik, Nykuping. Sweden (Received August 8 2 , 1 9 6 8 )
Pulse radiolytic absorption transients have been observed in aqueous solutions between 200 and 300 nm using an 1 1 - M e ~Linac and an optical detection system that allowed accurate measurements (a) down to 200 nm M HCIOh 0.027 M HZ (p(H2) = 35 atm) transients and (b) 0.2 psec after the electron pulse With with second-order decay were observed which had amplitudes that decayed monotonically in the region from 200 to 240 nm. Assigning these transients to free H atoms, the molar decadic absorptivities E at 200, 210, and 240 nm of I3 were found to be 900, 560, and 0 cm-', respectively, and %H+H = (1.55 =!= 0.10) x 1OO ' M-1 sec-l from measurements at 200 and 210 nm. The transients could be completely quenched by addition of O2 resulting in a species with the absorption spectrum of HOz. Furthermore, the transient at 210 nm was M ) was added innot affected when IfClOr was left out of the Hz-saturated solution and E20 ( > 2 X stead. The apparent OH transient in 2 X lW3 M NzO (no Hz) decayed according to second-order kinetics with a calculated rate constant that after correction for the reaction of H withOH was found to be (1.04 rt 0.10) X 1O'O Jl-l sec-' independent of the wavelength used. The calculated e for OH showed, after correction for the absorbance of € 1 2 0 2 , H., and OH-, one broad absorption maximum near 230 nm with E 530 M-l cm-l and one below 200 nm. The measurements at 200 nm had to be corrected for a substantial contribution from OH- to the observed optical absorption. The calculated values of EH and EOH account quantitatively at all M HClOd (no 132) if it is assumed that waveiengths used for the initial absorption of the transients in H30, if formed, decomposes to yield H f H20 after no longer than 0.2 psec. The light absorption of aqueous solutions of H and to some extent of OH at 200 nm is attributed to a red shift of the water absorption continuum beginning at 186 nm, caused by a partial electron transfer from the first excited singlet state of water to a neighboring H or OH free radical in analogy with the optical transition associated with the @ bands in alkali halide crystals.
+
Introduction Free hydrogen atoms in the gas phase do not absorb light in the ground state a t longer wavelengths than 121.57 nm,3 and it has, in general, been assumed that solutions of free hydrogen atoms in water do not absorb 1igh.t in the far-ultraviolet spectral region (>200 nm). It is the purpose of this report to demonstrate that solutions of free hydrogen atoms in water, produced by pulsed radiolysis in solutions saturated with pressur-
ized Hz, do in fact absorb light appreciably a t 200 nm and to suggest the type of optical transition involved, (1) Presented at the Symposium on Photochemistry and Radiation Chemistry at Natick, Mass., Aprii 22-24, 1968. A preliminary communication has appeared: 9. 0. Nielson, P. Pagsberg, J. Rabani. H. Christensen, and G . Nilsson, Chem. Commun., 1523 (1968). (2) (a) Aktiebolaget Atomenergi, Studsvik, Nykoping, Sweden: (b) Department of Physical Chemistry, The Hebrew University, Jerusalem, Israel. (3) G . Herzberg, "Atomic Spectra and Atomic Structure," 2nd revised ed, Dover Publications, New York. N. Y., 1944, p 25. Volume YS, Number 4 April I960
