Table 11. Absolute Band Intensities, Carpenteria
Band, cm-I
ALorentzian
AKswshara
- AK AK 9
860 805 740
0 .,057 0.099 0.111
0.044 0.074 0.055
30 34 100
AL
%
where the Lorentzian coefficients were obtained from the solution of Equation 5. The observed half band widths are 61 = 60 cm-I, 6 2 = a3 = 50 cm-I. [The spectral slit width for a Perkin-Elmer 337 in this region is about 3 cm-I (I4).] The results of this computation, shown in Table 11, clearly show that the use of absorbance measurements as suggested by Kawahara ( I ) as a quantitative method of differentiation of
such species could lead to erroneous results and that only a highly sophisticated computerized data reduction system which utilizes the total spectrum should be employed in this manner. ACKNOWLEDGMENT
The authors thank Linda Ruswinkle for performing the analyses, Paul A. Wilks, Jr. of Wilks Scientific for technical advice, Eugene Christianson and Sherry Weston for assistance in obtaining the seepage samples. RECEIVED for review September 22, 1969. Accepted December 5,1969. This research was supported in part by the Federal Water Pollution Control Administration, Grant NO. 16020-ELH (U. of M.) and the Western Oil and Gas Association (U.S.C.).
Differential Constant-Current Potentiometry for Analysis by Kinetics J. R. S a n d and C. 0. Huber Department of Chemistry, University of Wisconsin-Milwaukee, Milwaukee, Wis. 53201 The term differential constant-current potentiometry is applied to a technique in which the difference in potential between two voltammetric electrodes held at different constant currents is monitored. The potential at each electrode is established by electroactive species corresponding to the current density applied. The experiment is performed with stirring. Shifts in potential occur whenever concentrations of electroactive species at the electrodes change. An electroactive species undergoing depletion by homogeneous solution reaction exhibits a peak-shaped potential vs. time response. The width of the peak is a measure of the homogeneous reaction rate. The differential nature of the technique implies that many ordinary sources of error should be self-cancelling. A wide range of reaction rates are amenable to observation by suitable adjustment in current density and time base measurements. The technique has been applied to the determination of phenols by bromination and to the oxidation of hydroxylamine by ferricyanide.
MANYreactions which are too slow for application to direct titration are nevertheless potentially useful for analysis. The classical approach to such applications has been indirect analysis using a known excess of reagent. An important current area of research in this respect is use of reaction kinetics for the determination of initial solution concentrations. The experimental requirements are that the reaction rate order remain invariant during measurement, and that the progress of the reaction can be monitored without interference by the monitoring method. Recent reviews summarize progress of this area of analysis ( I , 2). Constant-current potentiometry has been used for titration involving relatively slow reactions because of its amenability (1) G. A. Rechnitz, ANAL.CHEM., 40,455R (1968). (2) H. L. Pardue in “Advances in Analytical Chemistry and
Instrumentation,” Vol. 7, C. N. Reilley and F. W. McLafferty, Eds., Interscience, New York, N. Y., 1969, p 141. 238
to higher temperatures, sealed containers, and end point detection with reproducible excess of reagent (3-5). Guilbault was the first to use the technique for direct analysis based an kinetics (6). It is the purpose here to present the technique of differential constant-current potentiometry (DCCP) and its application to analysis based on reaction kinetics. This method of rate analysis promises many attractive capabilities. The method utilizes a truly differential system to obtain the measured signal; and, presumably, small variations in controlled conditions will be self-cancelling. Anodic or cathodic observation of the appearance of products or the disappearance of reactants suggests a great deal of versatility. Most of the advantages of ordinary constant-current potentiometry are retained. Steady potentials are obtained rapidly when working with reversible or slightly irreversible systems. Nonaqueous solutions, sealed systems, exclusion of light, inert atmospheres, and many other, sometimes difficult, experimental conditions can be accommodated. If pseudozero-order in analyte is maintained, it should be possible to perform two or more repetitive runs in the same solution to increase measurement precision. The exact amount of reagent added to initiate the determination is not critical. Many reactions which are too slow for direct or indirect methods of analysis may find an application in DCCP because complete reaction is not required, and a large excess of one of the reactants is permissible. Finally, the convenient size and shape of the signal obtained from the method should permit relatively straightforward automation of time interval measurement and subsequent data treatment. ~
~~
(3) R. W. Murray and C. N. Reilly in “Treatise on Analytical Chemistry,” Part I, Vol. 4, I. M. Kolthoff and P. J. Elving, Eds., Interscience, New York, N. Y., 1963, pp 2181-2187. (4) J. R. Sand and C . 0. Huber, Tuluntu, 14, 1309 (1967). ( 5 ) C. 0. Huber and K. E. Smith, ANAL.CHEM., 40,982 (1968). (6) G . G. Guilbault, B. C. Tyson, Jr., D. N. Kramer, and P. L. Cannon, Jr., ibid.,35, 582 (1963).
ANALYTICAL CHEMISTRY, VOL. 42, NO. 2, FEBRUARY 1970
A
3
LA+ i WE)
Figure 1. Idealized current-voltage plots
Time
---+
Curves 1-4, Decreasing oxidant concentration Curve 5, Residual current
Figure 2. Differential constant-current potentiometric analysis signal
The technique of differential constant-current potentiometry (DCCP) can be described using the current-voltage curves shown in Figure 1 . Curve 5 represents the residual current curve (typically hydrogen reduction in the cathodic region and oxygen evolution in the anodic region) of a system which contains some chemically reducible species. Curves I through 4 depict voltammetric waves for decreasing concentrations of some oxidizing agent-bromine in this case. The two dashed horizontal lines indicate two working electrodes polarized at different, cathodic-current densities. Before any oxidizing agent is added, both cathodes will exhibit potentials indicated by their intersection with the residual current curve; and the voltage difference between them will be quite small. To initiate the kinetic run, enough oxidizing agent is added to cause both working electrodes to move from the potential of hydrogen reduction toward the potential at which the oxidizing agent is electrochemically reduced (Curve I ) . As the oxidant is depleted in the homogeneous solution reaction, its concentration drops to a point where the working electrode with the higher current-density must revert back to the potential of hydrogen reduction in order to support the applied current. When this happens, the voltage difference between the two working electrodes begins to increase rapidly. This is the situation illustrated by Curve 2 in Figure 1 . By the time the oxidant concentration has decreased to the point represented by Curve 3, the potential separation of the two working electrodes goes through a maximum, and the low current electrode begins to shift to more negative values. The low current electrode eventually returns to the residual current-voltage curve where it started, and the voltage difference between the two working electrodes is again very small. Monitoring the potential across the two constant-current cathodes during the course of a reaction similar to the one described, results in a peak shaped signal (Figure 2). The width of this peak is used as an indication of the rate and thus the initial concentration. Determination of substance b by reaction with reactant a under conditions such that the reaction is pseudo-first-order with respect to a is described by an integrated form of the first-order rate Equation 1 for depletion of concentration C, to C,‘ over the time interval A t . In
cut --
C,
= kCbAt
Assuming that the constant currents used in the method correspond to limiting currents encountered in voltammetry, Equation 2, the basic equation for voltammetry with stirred solutions, can be used to substitute for C, and C,’, where n, F, and D have their usual designations. The Nernst diffusion i = nFADC/G
(2)
layer thickness is given by 6. Substituting C, and Cat from Equation 2 into Equation 1 and rearranging gives Equation 3. (3)
With the ratio i‘li experimentally established at a constant value and the diffusion thicknesses held constant by using a reproducible stirring rate and cell geometry, the analyte concentration is inversely proportional to the measured time interval. Equation 3 shows that slight variations in mass transfer rate and effective electrode area due to temperature changes, electrode surface effects, stirring inequities, etc., should be self-cancelling. In order to demonstrate and evaluate this new analytical technique, it was applied to two types of analysis, one representing an organic functional group analysis and the other an inorganic determination. These reaction systems were the bromination of 2,6-dichlorophenol (Equation 4) and the oxidation of hydroxylamine (Equation 5 ) with potassium ferricyanide.
2NHzOH
+ 2[Fe(CN)6J-a + 20H- -+ Nz
+ 2[Fe(CN)sl-4 + 4 H z 0
(5)
Bromination has been used by many authors for the determination of phenols (3, and several direct and indirect methods exist for the analysis of hydroxylamine (8, 9).
(7) C. 0. Huber and J. M. Gilbert, Anal. Chem., 34,247 (1962). (8) B. R. Sant, 2.Anal. Chem., 158, 116 (1957). (9) A. J. Clear and M. Roth in “Treatise on Analytical Chemistry,” I. M. Kolthoff and P. J. Elving, Eds., Pt. 11, Vol. 5 , Interscience Publishers, New York, 1961, pp 288-290.
ANALYTICAL CHEMISTRY, VOL. 42, NO. 2, FEBRUARY 1970
239
IA'!
-
< t Figure 3. Differential constant-current potentiometry circuit
vs. Phenol C o n c e
.533
.2w
/
P N
A = ta'! B =
-
at I'
'I
F, 1:l voltage follower for impedance matching
'I
Rec, Potentiometric recorder
600 m ~ .o 5 0 0 rnV. 400 mV. 0
pH, pH meter for monitoring individual electrode
potentials EXPERIMENTAL
- V
Apparatus. The two working electrodes were constructed
by sealing 1.0 cm lengths of 20 gauge platinum wire in glass tubing. A 16.0-cm length of 20 gauge platinum wire, sealed in glass and spot welded in the form of a loop, was used as the counter electrode. The saturated calomel reference electrode (SCE) was connected to the reaction cell by means of a salt bridge filled with 0.10M nitric acid. The reaction cell itself consisted of a water-jacketed 250-ml beaker fitted with a ground glass top into which were sealed: both working electrodes, the counter electrode, a nitrogen purge tube, a standard-taper joint for the salt bridge, and a solution inlet port. Solutions were stirred magnetically with a synchronous motor. A diagram of the constant-current and potential measuring circuit is shown in Figure 3. The constant current to the working electrodes is supplied by two 90-volt batteries in series with appropriate resistors. A one-to-one voltage follower based on a single operational amplifier in a conventional feedback circuit serves to ensure impedance matching to the recorder used for readout. The recorder used was a 1250-mV, 1 sec full-scale model. Also indicated on the diagram is a pH meter which can be used to monitor the potential of either working electrode individually relative to the SCE reference electrode. The pH meter is not essential to the analysis technique, but is convenient when studying untried systems. The temperature of the reaction cell was kept at 23.0 L 0.1 "C. Reagents. The 2,6-dichlorophenol (mp 67 "C) was recrystallized twice from high-boiling petroleum ether. Reagent grade hydroxylamine hydrochloride was washed with anhydrous ether and dried in a desiccator before being used. All other chemicals were reagent grade and used as supplied. The pH 8.1 borax buffer used for the hydroxylamine determination was prepared by dissolving 6.0 grams of boric acid and 9.0 grams of borax in a liter of solution. Prepurified nitrogen was used as supplied for purging the solutions. 2,6-Dichlorophenol Determination. The working electrodes were pretreated before determinations by anodizing at 3.0 V for one minute in reaction solvent. The anodization pretreatment was terminated by draining the solution from the cell. Both electrodes were switched to open circuit while new solution was being added and purged with nitrogen. 240
I
20
so
4 0
Phenol
Conce
I
-
no
IO 0
IO3
Figure 4. Reciprocal time DS. 2,Cdichlorophenol concenlration A , A t measured at 600 mV B, 500 mV C, 400 mV D,At half-peak potential x , Sample results
After briefly passing a 30-pA current through both electrodes they were switched to the respective constant-current values of 8 and 18 pA. Then 15 pl of approximately 0.1M bromine moles) solution in 97% v/v (approximately 1.5 X acetic acid-water was added from a pipet fitted with a Teflon (Du Pont) tube. Oxidant was added at a relatively hlgh concentration and corresponding small volume in order to minimize dissolved oxygen addition. The 2,6-dichloroto 1.2 X phenol amounts ranged from 0.12 X moles in 97 % vjv acetic acid-water solvent which was 0.20M in NaC104. The solution volume during analysis was 120.0 ml. Hydroxylamine Determination. The working electrodes were pretreated by a procedure similar to the one described for the phenol determination before every run. Constant currents of 9 and 6 pA were used in this system. to The hydroxylamine amounts ranged from 1.2 X moles. Solutions were made up in the reaction 6.0 X vessel by adding a measured volume of nitrogen-purged 0.50M hydroxylamine hydrochioride stock solution, previously neutralized with sodium hydroxide, to a measured volume of purged borax buffer. The buffer solution was also 0.10N in NaC104 and 0.10N in NaCN. The addition of the hydroxylamine was accomplished by means of a 25-ml buret fitted with 0.7 mm diameter Teflon tubing. The total volume in the reaction cell was always 120.0 ml. The kinetic run was initiated by 10 pl addition of l.OMK,Fe(CN)o (1.0 X moles). RESULTS AND DISCUSSION
Figure 4 shows plots of reciprocal time in minutes us. 2,6dichlorophenol concentration. The lines are least-square fits
ANALYTICAL CHEMISTRY, VOL. 42, NO. 2, FEBRUARY 1970
~~
~
Table I. DCCP Analysis of 2,6-Dichlorophenol Concentrations Taken Found Std. dev. AE (M X loa) ( M X IOa) (M X lo3) Detns. 2.65 3t0.06 8 600 mV 2.67 8.10 2~0.07 6 600 mV 7.97 Ed2 Ell12
2.67 7.97
2.73 8.06
3t0.07 3tO.08
8 6
of the data. Table I summarizes accuracy and precision for samples analyzed. Figure 5 contains similar reciprocal time US. concentration plots obtained for the hydroxylarninepotassium ferricyanide system. No attempt was made to obtain the best least-square, straight line because of the remarkable linearity obtained. The data shown are for the second DCCP run in each solution; results from the first run were spurious. Apparently, the first run acts in some way to condition the system for subsequent runs. The accuracy and precision of this method are indicated by six determinations on solutions with the same hydroxylamine concentration. These showed a relative standard deviation of approximdte!y 3x. These data indicate that it is indeed possible to get analytically useful measurements from the technique and that the accuracy and precision obtained are comparable to that obtained Irom other kinetic methods (10). The major limitaticn on the precision obtainable seems to stem from the non-inert character of the indicator electrodes used ( I l - 13). The inability to reproduce electrode surfaces restricts the ultimate precision to approximately one per c a t of the substance being measured. When gold plated platinum electrodes were Lsed with the bromine-phenol system, no significant improvement in peak shape or reproducibility was obtained. A slight decrease in the height as well as the width of the DCCP peaks was noted when solutions with faster reaction rates were analyzed. This was attributed to sluggishness in the movement of the high current electrode as it approached hydrogen reduction potentials. A more marked efect is seen in the faster reacting solutions because the disappearance of oxidant causes the low current electrode to start its movemeni toward more negative potentials before the high current electrode has moved out of this sluggish region. In an attempt to compensate somewhat for this phenomenon, time interval measurements were made ai a point which was related to the peak height, specificaliy, at the haif-peak potential. The results of these measurements are shown as two straight-line plots on Figures 4 and 5. Table I indicates that the accuracy and precision obtained by this procedure are not significantly better than that obtained by measuring time intervals at fixed potentials. The potentials chosen for time interval measurement should correspond to points on the DCCP peaks a where the rate of potentiai change is most rapid-Le., potential range where the sides of the signal peak are most vertical. It is in this range of potentials that surface effects are minimal. Some of the practical considerations which must be weighed in choosing the constant current levels to be used are the reversibility of the electrode couple being observed, the reaction rate, electrode surface effects, and any theoretical advantages associated with a certain ratio. If relatively ir(IO) G. E. James and W.1,. Pardue, ANAL.CHEM., 41,1618 (1969). (11) S. D. James,J.Elecr;ochem.Soc., 114, 1113 (1967). (12) M. W. Breiter, ibid.,8,230 (1964). (13) I. M.Kolthoff and N. Tanaka, ANAL.CHEY.,26,632 (1954).
‘I I
50W
/
/4
Figure 5. Reciprocal time us. hydroxylamine concentration A , At measured at 350 mV B, 250 mV C, At half-peak potential
reversible systems are used, the separation between the high and low current must be increased. A small current separation will permit both working electrodes to shift potential together, and virtually no DCCP peak will be observed. Systems with quite fast reaction rates, in which the disappearance of reactants is being observed, necessitate the use of smaller currents. They permit the initial shift of both electrodes to the proper poteiltial before reaction rate-order conditions are violated. Conversely, observing the formation of products from a fast reaction will necessitate larger currents to permit adequate mixing of the reagents before the DCCP peak is observed. Higher currents with a large current separation should be avoided because this condition tends to create dissimilar surface conditions on the two electrodes and destroys the advantages obtained by the differential aspect of the technique. It is possible to calculate a current ratio which should minimize the random errors in reading the time interval. Starting with Equation 3 in which Cois related to the reciprocal of the time interval 2nd using a procedure similar to the method of calculating photometric absorbances that should give a minimum relative reading error in concentration determinations ( I d ) , one can determine this current ratio to be 0.368. Usually, however, reversibility and reaction rate considerations, along with electrode surface effects, are more important considerations. It was nored that extensive cathodization of the working electrodes before initiation of a kinetic run had to be avoided. Prolonged evolution of hydrogen seemed to slow response of the e!ectrode to the presence of excess oxidant. Consequently the reduction of the electrodes prior to use was minimized. (14) G . H. Brown and E. M. Snllae, “Quantitative Chemistry,” Prentice-Hall, Inc., Ziiglewood Cliffs,N. J., 1963, pp 399--W.
ANALYTICAL CHEMISTRY, VOL. 42, NO. 2, FEBRUARY 1970
*
241
Purging the reaction solutions with nitrogen prevented any kinetic complications arising from independent reactions with dissolved oxygen. Hydroxylamine in neutral solution is oxidized by dissolved oxygen to a variety of products which are reduced at platinum electrodes (15). Removing dissolved oxygen was also necessary from the standpoint of preventing irreversible oxygen reduction at the working cathodes, The potential at which this process occurred was found to be less reproducible than hydrogen reduction. From the two systems explored, some of the basic capabilities of the method of DCCP can be stated. Care must be exercised in choosing the systems to which the technique is to be applied. It is essential that the products of the desired reaction are not themselves capable of reacting at the working electrodes. If the analyte is electroactive, the potential at which it reacts at the working electrode must be significantly different than the potential at which the species added to start the kinetic run interacts with the working electrode.
(15) P. C. Noews, Jr. and L. F. Audrieth, J . lnorg. Nucl. Chem., 11, 242 (1959).
The lower determination limit for the electrodes used in these applications is about 10-BM. Because of the pseudofirst-order requirement on the reaction (pseudo-zero-order on analyte concentration), the smallest analyte concentration determinable by the method would be about 10-4M. For cases where initial increasing concentrations of product rather than decreasing concentrations of reagent are monitored, this limit could be lower. Increasing electrode area would increase sensitivity, but the associated increase in the contribution of undesirable electrode surface reactions to the overall signal is often intolerable. Maximum concentrations determinable depend on rates at which electrode surface effects permit electrode response to solution concentrations. This same criterion limits the selection of applicable systems to those where rate of reaction allows adequate time for the electrode response. Systems with slow reaction rates are prohibitive only when competing solution reactions are significant. RECEIVED for review October 20, 1969. Accepted December 1, 1969. Presented at the 3rd Great Lakes Regional ACS Meeting, DeKalb, Illinois, June 1969. Part of a thesis for the M.S. degree by J. R. S.
Direct Current and Alternating Current Polarographic Response of Some Pharmaceuticals in an Aprotic Organic Solvent System Albert L. Woodson Food and Drug Administration, Chicago District Laboratory, Room 1222, Post Ofice Building, Chicago, III. 60607
Donald E. Smith Department of Chemistry, Northwestern Uniaersity, Ecanston, 111. 60201 Extensive investigations of electrode reaction mechanisms in aprotic organic solvents have revealed that considerable potentialities and advantages should attend application of such solvent systems to voltammetric analysis of organic compounds in general. Because they have received little exploitation, the present work was undertaken primarily to demonstrate these potentialities in a more explicit manner than provided by earlier mechanistically-oriented investigations. Interest was focused on organic compounds of significance in the pharmaceutical field. t h e dc, fundamental harmonic ac, and second harmonic ac polarographic responses of 24 pharmaceutically-important compounds in acetonitrile-tetrabutylammonium perchlorate media were investigated. Barbiturates, salicylates, corticosteroids, alkaloids, sulfa drugs, and estrogens are among the compound classifications included in this work. The majority of the compounds yielded ideal responses for analytical purposes with one or more of the techniques employed. Many gave ideal, one-electron reversible (diffusion-controlled) waves, whereas the corresponding aqueous solution response is irreversible or nonexistent. ONEof the most significant advances in modern electrochemistry has involved the study of electrode processes in aprotic organic solvents. The past decade has witnessed this field develop from infancy to a state of incipient maturity as a result of rapid evolution of knowledge in three principal areas: instrumental techniques, usable solvents and appropriate 242
schemes for their purification, and electrode reaction mechanisms. The measurement problems attending the high resistance characterizing most aprotic solvent-supporting electrolyte systems (1, 2) have been overcome to a large extent by the development of potentiostats and galvanostats which permit automatic potential and current control with threeelectrode cell configurations (3-10). This development began in the late 1950’s with the efforts of Booman (3),DeFord (3, and Kelley and coworkers (6, 7). With the latest refinements
(1) J. Heyrovsky and J. Kuta, “Principles of Polarography,” Academic Press, New York, N. Y., 1966, pp 61-63. (2) P. Delahay, “New Instrumental Methods in Electrochemistry,” Interscience Publishers, New York, N. Y., 1954, pp 132-135, pp 166-168. (3) G. L. Booman, ANAL.CHEM., 29,213 (1957). (4) G. L. Booman and W. B. Holbrook, ibid., 37,795 (1965). (5) D. D. DeFord, Division of Analytical Chemistry, 133rd Meeting, ACS, San Francisco, Calif., April 1958. (6) M. T. Kelley, D. J. Fisher, and H. C. Jones, ANAL.CHEM., 31, 1475 (1959). (7) M. T. Kelley, D. J. Fisher, and H. C. Jones, ibid., 32, 1262 (1 960). (8) D. E. Smith in “Electroanalytical Chemistry,” A. J. Bard, Ed., Vol. 1, M. Dekker, Inc., New York, N. Y.,1966, Chapter 1. (9) E. R. Brown, D. E. Smith, and G. L. Booman, ANAL.CHEM., 40,1411 (1968). (10) E. R. Brown, H. L. Hung, T. G. McCord, D. E. Smith, and G. L. Booman, ibid., 40,1424 (1968).
ANALYTICAL CHEMISTRY, VOL. 42, NO. 2, FEBRUARY 1970
