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A. Keenan and I. J. Ferrer-Vinent
3, 1979
(32) T. Osa and M. Fujihira, Nature (London), 264, 349 (1976). (33) M. S. Wrighton, R. G. Austin, G. B. Bocarsly, J. M. Bolts, 0. Haas, K . D. Legg, L. Nadjo, and M. C. Palazzotto, J. Am. Chem. Soc., 100, 1602 (1978). (34) D. F. Untereker, J. C. Lennox, L, M. Wier, P. R. Moses, and R. W. Murray, J. Elecfroanal. Chem., 81, 309 (1977). (35) P. R. Moses, L. M. Wier, J. C. Lennox, H. 0. Finklea, J. R. Lenhard, and R. W. Murray, Anal. Chem., 50, 576 (1978). (36) S. R. Morrison, "The Chemical Physics of Surfaces", Plenum Press, New York, 1977, p 320. (37) H. Gerischer, Adv. Electrochem. Necfrochem. Eng., 1, 139 (1966). (38) A small bump was also apparent in the cyclic voltammograms at
(39) (40) (41) (42)
approximately the same potential as the peaks in Rand C(Figure 3). These peaks dd not appear consistently, but were more prevalent at high pH's. Their shape and dimunition at high frequencies suggest that they may be caused by charging and discharging of a surface state. The crystal face was roughened by the polishing procedure and probably a variety of crystallographic planes were exposed. This would lead to a mean dielectric constant between" 173 and 89. P. T. Boddy, J. Electrochem. SOC.,115, 199 (1968). M. A. Butler and D. S. Ginley, J. Ekctrochem. Soc.,125, 228 (1978). J. R. Lenhard, R. Rocklin, H.Abruna, K . Willman, K . Kuo, R. Nowak, and R. W. Murray, J . Am. Chem. SOC.,100, 5213 (1978).
Transition MetaVMetal Oxide/Carbonate/Carbon Dioxide/Oxygen Electrodes at 350 in Fused Potassium Nitrate'
O C
A. G. Keenan" and Ignaclo J. Ferrer-Vinent Department of Chemistry, University of Miami, Coral Gables, Florida 33 124 (Received August 18, 1978) Publication costs assisted by the University of Miami
--
Previous work on Pt/PtO and Cu/CuO electrodes has been extended to a series of ten transition metals. Co, Ni, and Pd, like Pt, follow the simple reaction M 4 C032- MO + COz 2e-. Fe, Zn, Ag, and Ta, like Cu, follow a reaction which may be generalized as MO + CO2- M + COz + O2 + 2e- although a more complex oxide may be involved in some cases. Standard potentials for the electrodes are reported. Evidence is presented to support the hypothesis that a carbonato complex of the metal ion is the critical species in the electrode reaction mechanism. The reason the metals fall into two distinct classes is the large differences in the crystal field stabilization energies in the carbonato complexes.
Introduction In previous work2 it has been shown that a P t / P t O electrode in fused nitrate at 350 "C responds to the electroactive species C032- and C 0 2 according to the equation Pt C032- PtQ COz 2e(1)
+
-
+
+
Zambonin et ala3have confirmed this result and shown that Au behaves similarly. In earlier work4 it was shown that a Cu/CuO electrode responds to the three species C03*-, COz, and O2 according to the equation CuO + CQ2Cu C 0 2 + O2 2e(2)
-
+
+
In the present work these investigations have been generalized to a series of ten available transition metals. With the exception of three metals which are not usable because they corrode and dissolve, the remainder fall neatly into one or the other of the two classes characterized by eq 1 and 2. Experimental Section The apparatus and procedures were identical with those already described in detai1.2,4 The metals used were of 99.0-99.95% purity. They were polished with fine emery cloth and oxidized in a nitrate melt as previously described. The reference electrode was a silver wire in a 0.1 m AgN03/KN03 solution. Following previously described method^,^^^ the electrode response to the various electroactive species was studied by varying the concentrations, one at a time, while holding the other(s) constant a t various arbitrary values. Linear regression computer analysis of the data gave the n values of the Nernst slope 2.3RT/nF. Runs were then made in which the concentrations of all electroactive species were varied simultaneously. These data were used to calculate 0022-3654/79/2083-0358$0 1.OO/O
+
TABLE I: N e r n s t Slope n Values a n d Standard Potentials Relative to 0.1 m Ag+/Ag
std n value and MDM
potential, 424.1 430.8 404.5
class
metal
I
co
1.93 i: 0.03 2.02 i: 0.09 1.98 i: 0.17
Fe
1.98 0.06 2.08 i: 0.11 1.98 i: 0.10 2.02 r 0.10
Ni Pd
I1
Zn Ag Ta
*
mV
416.0 476.9 474.6 783.5
the standard electrode potentials using the applicable Nernst equations. Carbonate concentrations were varied from 2 X to 4 X m. C 0 2 and O2 partial pressures ranged from 0.04 to 0.90 atm. Results All the metal electrodes were found to be anodic to the silver reference electrode. Table I gives the averages of the experimental values of n and the mean deviations from the average as determined from the least-squares analysis. The value of n = 2.00 is seen to be confirmed well within experimental error for all cases. As stated, the metals fell into two distinct classes. Class I metals, as listed in Table I, showed response to COZ- and C 0 2 only and were insensitive to O2 within experimental error. Their Nernst equation would be of the form2 (3)
This puts them in the same class as Pt and Au and their electrode reactions would be analogous to eq 1.
0 1979 American Chemical Society
Transition Metal Electrodes
Class I1 metals form electrodes which respond to CO?-, C02, and O2 and would conform to a Nernst equation of the form4
The Journal of Physical Chemistry, Vol. 83, No. 3, 1979
359
TABLE 11: Elec,trode Reactions, Oxidation States, and dn Electron Configurations of Metal Ions in the Oxides Shown oxdn electrode reaction state dn
-.-
They are in the same class as Cu and their electrode reactions would be analogous to eq 2. Cr, Mo, and W showed extensive corrosion, flaking, and dissolution and could not be used as electrodes. Also, as with Pt and Cu, all class I and I1 metals showed erratic behavior in pure COz. Table I also contains values of the standard potential E"' obtained from runs in which all electroactive species were varied. Since the Nernst equations are written in terms of partial pressures, Eo' includes the applicable Henry's law constants for dissolution of the gases in the melt. Graphs of emf against log concentration when the electroactive species were varied one a t a time gave good straight lines with no more scatter than shown in the figures in the earlier paper^.^^^ Except for silver, the points all lay within less than three standard deviations of the theoretical Nernst line for n = 2.00. The scatter for Ag was somewhat greater. This may have been due to the lower stability of Ag,O, resulting in a less reproducible electrode surface, or due to an unfavorable kinetic factor. All metals of class I and I1 were also run with the melt in equilibrium with the ambient atmosphere while varying the carbonate concentration. As with Pt and Cu good, reproducible Nernst lines with n = 2.00 were obtained. The electrodes can thus be used to measure carbonate in nitrate melts in air without the elaboration of preparing a synthetic atmosphere.
Discussion An extensive search of the literature was made to discover parameters which might explain the grouping of co, Ni, Pd, Pt, and Au into class I response and Fe, Cu, Zn, Ag, and T a into class 11. Obviously periodic table family relationships give no clear correlation except that the three corroding metals, Cr, Mo, and W, constitute family 6B. Among the properties investigated were electron work functions, band gaps, ionic radii, lattice energies, ionization potentials, bond energies, crystal structures, cohesive energies, semiconductor properties, general physical and thermodynamic properties, and activity of the metals as hydrogenation catalysts. Although the present experimental data were taken a t 350 "C in a nitrate melt, properties in aqueous solution a t room temperature, such as standard electrode potentials and hydrogen and oxygen overvoltages, were also investigated. No correlation with any of these properties was found. Discussion and references are given in the Ph.D. dissertation of Ferrer-Vinent.5 Interestingly enough, a very persuasive correlation was found with crystal field stabilization energies of metal ion carbonato complexes. Metal surfaces in fused nitrate a t 350 " C contain an oxide layer. In the case of transition metals, the oxides contain mixed valence states of the metal ions and probably have an indefinite composition. An extensive discussion with literature references is given in ref 5. Since the oxide concentration cannot be varied, it is of course not possible to determine, by electrometric measurements alone, the particular oxide species which is electroactive. However, electrode reactions which conform to the applicable Nernst equation may be written for all of the metals using known oxides. These of course
Class I Co -t C0,'- -+COO t CO, t 2eNj t C0,'- -+ NiO t CO, + 2ePd + C0,2- + PdO t CO, t 2ePt t CQz- PtO t CO, t 2eAuO t CO, + 2eAu t C0,'Class I1 Fe,O, t COS2--+ 2 F e 0 t CO, t 0, t 2eCU,O t CO," -+ 2Cu t CO, t 0, t 2eZnO t COS2--+ Zn t CO, t 0 , t 2eAg,O t C0,'- -f 2Ag t CO, t 0 , t 2eTa,O, t CO,'--+ 2Ta0, + CC), t 0, t 2e-f
-ic
Co(I1) Ni(I1) Pd(I1) Pt(I1) Au(I1)
d7 ds d* d8 d9
Fe(II1)
db
Cu(1)
d10
Zn(I1)
dIo
Ag(1)
d"
Ta(V)
do
TABLE 111: Crystal Field Stabilization Energies in Dq Units for Weak Field Carbonato Complexes and Various Metal Ion d" Electron Configurations d" do db d7 d8 d9 d'O
-
octahedral 0 0 8.0
square planar
square pyramid
0 0
0
-
0
12.0 6.0
10.3 14.6 12.3
10.0 9.1
0
0
0
9.1
need not be the most stable ones a t room temperature or in contact with aqueous solutions. These are shown in Table I1 together with the oxidation state of the metal ion and its d" electron configuration. It is interesting that class I reactions all involve oxidation of the metal by CO2whereas class I1 involve reduction of the metal species. Overall, these reactions are all anodic oxidation reactions. It is seen from Table I1 that the metal ions in class I are d7, d8, or d9 whereas those in class I1 are do, d5, or dlO. It is well known in inorganic chemistry that members of the latter class show similar properties. The equations in Table I1 represent overall or stoichiometric electrode reactions. The elementary mechanistic steps could only be discovered by the methods of electrode kinetics. It is nevertheless reasonable to propose that a critical step in the mechanism is the formation of carbonato complexes with metal ions on the surface. Such complexes are known from spectroscopic studies to be of I11 gives values for the crystal the weak field t ~ p e . ~Table J field stabilization energies of weak field carbonato complexes of various geometries and for various metal ion electron configurations.8 These values are calculated theoretically from d-orbital energy levels and are independent of the medium. It is immediately apparent that the stabilization energies for class I metal ions are all uniformly zero whereas class I1 metal ion complexes all have quite substantial crystal field stabilizations. This direct correlation between the crystal field stabilization energies of the carbonato metal ion complexes and the classification of the metals into two distinct types of electrode response suggests strongly that such complexes are a critical speciies in the mechanisms of the electrode reactions. The measurements reported are of course under equilibrium conditions but variations in intermediates or in activated complexes could change the reaction mechanism to yield different stoichiometries. This possibility
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The Journal of Physical Chemistry, Vol. 83, No. 3, 1979
M. Almgren and R. Rydholm
suggests an interesting field of study of these reactions by the methods of electrochemical kinetics.
(3)
E. Desimoni, L. Sabbatini, and P. G. Zambonin, d. Elechoanal. Chem.,
71, 73 (1976). (4) A. G. Keenan and C. G. Fernandez, J. Phys. Chern., 78, 2670 (1974). (5) I. J. Ferrer-Vinent, Ph.D. Dissertation, University of Miami, May, 1978. (6) L. E. Orgel, "An Introduction to Transition-Metal Chemistry: Ligand Field Theorv". Wilev. New York. 1970. (7) D. D. Perrin, kev. pure Appl. Chern., 9, 257 (1959). (8) F. Basolo and R. G. Pearson, "Mechanisms of Inorganic Reactions: A Study of Metal Complexes in Solution", Wiley, New York, 1967.
References and Notes (1) This work comprises part of the Ph.D. Dissertationof I.J.F.V., University of Miami, May, 1978. (2) A. G. Keenan and Thomas R. Williamson, J . Phys. Chem., 82, 46 (1978).
Influence of Counterion Binding on Micellar Reaction Rates. Reaction between p-Nitropnenyl Acetate and Hydroxide Ion in Aqueous Cetyltrimethylammonium Bromide Mats Almgren" and Robert Rydholm" Department of Physical Chemistry, Chalmers University of Technology and University of Gothenburg, S-4 12 96 Gothenburg, Sweden (Received July 6, 1978) Publication costs assisted by Chalmers University of Technology
The reaction between p-nitrophenyl acetate (PNPA) and hydroxide ion has been studied in aqueous solutions of cetyltrimethylammonium bromide (CTAB), in the absence of buffer substances. To account for the results, a model is presented. The fundamental factors affecting the rate are (i) the solubilization equilbrium of PNPA and (ii) the counterion binding equilibrium. The latter is treated as a Langmuir adsorption, with competition between OH- and Br- for the available sites. The experimental results show that the total degree of counterion binding might not be constant. The individual counterion binding constants are dependent on the surface potential of the micelle. Due to the preferential binding of Br-, not only the ionic strength but also the composition of the ionic solution determines this potential. The change in surface potential is accounted for in two ways: (i) changes in surface charge and (ii) changes in other surface properties. The results are compatible with the model if Br- is bound stronger than OH- by a factor of 40 f 10 and if the second-order rate constant is about 6.5 M-' s-l, somewhat smaller than in the aqueous solution (10.9 M-l s-'). The catalysis is therefore in this case a mere concentration effect.
Introduction The reaction between p-nitrophenyl esters and hydroxide ion has earlier been extensively used to study the effects of micelles on reaction rates.'-7 Anionic micelles are expected to inhibit the reaction, whereas cationic micelles should catalyze it.8a The results for catalyzed proce~sesl~"~ often show a maximum in the first-order rate constant when the surfactant concentration is varied and the initial reactant concentrations are held constant. This has hitherto not been found for this system, but at higher hydroxide ion concentration and without buffer we found a rate maximum (Figure 1) similar to those obtained for other reactions. The appearance of a rate maximum is the consequence of two counteracting effects: (i) the rate enhancement, due to the concentration of the reagents to the micelle surface and possibly a specific catalytic effect there, and (ii) the rate decrease, due to the replacement of the reactive counterion by the normal micelle counterion with increasing surfactant concentration. Berezin et al.9 and Romsted5 have presented general theories for micellar catalysis. Berezin et al. do not treat systems where specific ionic effects are important. Romsted deals with such systems and has convincingly shown that the two effects mentioned above are sufficient to explain the dependence of the reaction rate on the surfactant concentration. However, a major shortcoming of his theory is that he makes the somewhat arbitrary assumption that the total number of counterions bound 0022-365417912083-0360$0 1.OO/O
to a micelle is constant and, furthermore, he has to assign a specific value to this degree of counterion binding in order to obtain a numerical value for the rate constant. We will here present a theory similar to Romsted's and test it especially regarding the constancy of the degree of counterion binding.
Experimental Section Materials. CTAB (Merck, >99%), PNPA (Merck, -99%), NaOH (EKA, 99.0%), and KBr (Merck, >99.5%) were used without further purification, and fresh stock solutions were prepared for each series of experiments. Apparatus and Kinetic Methods. The kinetic measurements were carried out in an Aminco-Morrow stopped-flow apparatus at 25.1 f 0.2 "C and the transmittance was measured a t 400 nm, where the product p-nitrophenolate ion (PNP-) has its absorption maximum. In series I an oscilloscope trace was photographed, and in series I1 and I11 the signal from the photomultiplier was recorded by a Biomation 805 waveform recorder. The greatest ratio between the initial concentrations of ester and hydroxide was 31800, and thus the hydroxide ion concentration can safely be regarded as constant. Contrary to earlier investigations, no buffer substances were used, thus restricting the possible counterions to Br- and OH-. The first-order rate constants were determined with the Guggenheim intercept method as the slopes of plotted straight lines. If appropriate, the slope value was determined by the method of least squares. The rate con@ 1979 American Chemical Society
