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tolerances are not exceeded with addition of the sample to be assayed. Solvents other than Zoctanol may be used, but optimal pH values will vary with the chain length, stereoisomer, and degree of oxygenation of the organic molecule. The described procedure was designed to be used for determining changes in phosphate concentrations of reaction mixtures containing enzyme systems-e.g. adenosine triphosphatase and oxidative phosphorylation activities. Such systems contain acid-labile organic phosphates which undergo molybdenum (7) and acid-catalyzed hydrolysis. As in the procedure used by Marsh (4) citrate was added to bind excess molybdate (3) and prevent molybdenumcatalyzed hydrolysis of these labile organic phosphates. Once the free molybdate was bound by citrate, subsequent formation of phosphomolybdic acid was completely inhibited and introduction of errors due to acid hydrolysis or accidental contamination of the sample was prevented. Citrate also lowered blank values by decreasing the amount of free molybdic acid in the organic layer. Phosphomolybdic acid became adsorbed to denatured protein and was not
extracted from the aqueous phase by 2-octsnol. Therefore, i t was necessary to eliminate all protein from samples before addition to the assay mixture. Perchloric acid was used to precipitate the protein which was then removed by centrifugation. When added to the reaction mixtures perchloric acid does not interfere with the assay if the final concentration is kept within the acid tolerance of the procedure. A constant temperature must be maintained in the samples and standards during aspiration into the spectrophotometer. A change of 1' C. (in the range of 17-40') altered the rate of Z-octanol uptake sufficiently to cause a 2% change in the absorption reading. Shielding the samples from flame heat may be required if an extended aspiration time is necessary. An extensive study of interfering ions has not been attempted. Presumably, ions reported to interfere with similar extraction techniques (3, 6) will exhibit comparable effects in this system. In the presence of 0.4 pmole phosphate, equimolar concentrations of sodium silicate or disodium arsenate resulted in absorbance readings which were 6 and 4% higher, respectively, than with phosphate alone.
The absorption readings from phosphomolybdic acid in 2-octanol were quite stable with time, showing only slight decreases after standing up to 3 days. The small changes that occurred on standing had no effect on the determination of phosphate concentrations, however, because of comparable changes in the standards. The described method works well with the biological systems mentioned and its application to the analysis of inorganic phosphate in other systems is anticipated. LITERATURE CITED
(1) Berenblum, I., Chain, E., Biochem. J . 32, 295 (1938). (2) David, D. J., Analyst 86, 730 (1961). (3) Davies. D. R., Davies, W. C., Biochem. J. 26, 2946 (1932). (4) Marsh. B B., Bwchem. Biophys. Acta 32, 357 (1959). (5) Sprague, S., Slavin, W., At. Absorption Newsletter 4, 293 (1965). (6) Wadelin, C., Mellon, M. G., ANAL. CHEM.25, 1668 (1953). (7) Weil-Malherbe, H., Green, R. H., Biochem. J . 49, 286 (1951). (8) Willis, J. B., Nature 207,715 (1965).
W.
s. Z.4CGG
R.J. KNOX
Bureau of Sport Fisheries and Wildlife Western Fish Nutrition Laboratory Cook, Wash.
Application of Controlled Potential Techniques to Study of Rapid Succeeding Chemical Reaction Coupled to Electro-Oxidation of Ascorbic Acid SIR: The theory of stationary electrode polarography has been extended recently to the case of a reversible electron transfer followed by an irreversible coupled chemical reaction (8). .R - ne-
e0
k -+
Z
(1)
By applying cyclic voltammetric experiments, the value of k can be evaluated from the ratio of cathodic to anodic peak currents observed a t a particular frequency. I n addition, Schwarz and Shain (14) have developed the theory and application of a cyclic s t e p functional controlled-potential technique in which the ratio of the cathodic to anodic currents for a particular frequency reflects the kinetics of the succeeding chemical reaction. Each of these techniques was employed here to obtain kinetic data for the rapid decomposition in aqueous solution of the oxidized form of ascorbic acid. There is nearly unanimous agreement in the recent literature that the electron transfer step in the electro-oxidation of 1760
ANALYTICAL CHEMISTRY
ascorbic acid is fast and that the initial electrode product undergoes an irreversible hydration reaction (3,4,9,12, 16),as in the equations HO
0
OH
\ / c=c
i i
C CHCHOHCHzOH
0
0
involved correlations of relatively subtle changes in the polarographic half-wave potential with drop time and temperature with the kinetics of the succeeding
- 2e-
0
\ / / c-c I \ - 2H+ e C CHCHOHCHzOH // \ / 0
0
HzO
L k
HO
OH
\/
c-c
CI
0
//
CHCHOHCH20H I
/\/
0 At least t w o previous attempts have been made to evaluate the kinetics of disappearance of the initial electrode product (3, 4). One of these studies
0
reaction (4). A rate constant of 3.33 x lo3 second-' was obtained. However, Ihe possibility for large experimental error was great. Furthermore, the
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esperimental conditions for that study nere different than for the work here, and a direct comparison of kinetic data is not possible. The other study involved chronopotentiometry with current reversal (3). Those authors obtained a value of 340 second-1 for the rate constant, but there is some doubt about the validity of this value because the authors apparently did not consider the response limitations of their instrumentation and electrolysis cell. The purpose of the present study was to obtain kinetic data for the rapid hucceeding reaction in the ascorbic :wid osidation process using two different controlled potential techniques and potcntiostatic instrumentation in which response limitations were not esceeded. To the authors' knowledge, neither of the techniques employed here, cyclic voltammetry and potential-step elect,rolysis, has been applied previously with esperimental confidence to systems where the rate constant for the succeeding reaction was as large as for this system. EXPERIMENTAL
Instrumentation. -1 solid-state potentiostat utilizing operational amplifiers (Philbrick Researches, Inc., Dedham, hlass.) was used in these investigations. The time-response characterist'ics of the inst'rumentation were evaluated previously ( I O ) . The caircuit, for tmhecontrolled-pot'ent'ial experiments is Figure 4b of Schwarz and Shain ( I S ) . The measuring device was a Tektronix Model No. 536 oscilloscope with t,ypes 11) and G plug-in units for the cyclic voltammetric esperinients, and Types 1) and 1' plug-in unit,s for t8hepotentialht8ep esperiments. Oscilloscope traces were recorded with a D u l I o n t S o . 2620 Polaroid camera, using Polapan Type 47, speed 3000 film. Square-nave and triangular-wave signals were generated by a HewlettPackard Model S o . 33001 function generator with a S o . 3302.1 plug-in trigger unit,. Single cyclic sweeps or steps were used in all experiments. Cells and Electrodes. The electrode assembly used in this work consisted of a hanging mercury drop working electrode (HMDE), a reference electrode with a Luggin capillary salt' bridge which contained the samc electrolyte as the solution under investigation, and a coiled platinum wire counter electrode. The evaluation of optimum cell response by careful con.iderat,ion of the working electrode-reference electrode probe separation was carried out as previously described ( I O ) . The other details of the electrolysis cell and electrode assembly used in this work also have been discussed ( I O , 11). &ill espeiinierits were performed a t ambient temperatures b e h e e n 22' and 24' C. Kinetic data Lvere obtained for a temperature of 23' + 0.2' C. The r:iJdii of the hanging drop elec-
Figure 1.
Cyclic voltammetric curve for ascorbic acid
1 .O X 1 0 - 3 M ascorbic acid, 1 M p H 7.2 phosphate buffer. Singlesweep experiment, 200 volts/second; vertical scale, 0.25 rna./div.; horizontal scale, 0.050 volt/div. Traces correspond to three replicate experiments
trodes used in this work were the order of 0.05 c1n. Materials. -ill chemicals were reagent made used without further imrTficatiin. All solutions were prepared in water purified by distillation and passage over a mixed cation-anion ewhange resin bed (Mallinckrodt, Amberlite LIB-3). High-purity nitrogen 1\85 used to reniove oxygen froin the solutions. The nitrogen was passed through a chromic chloride Sohtion (5) to remox e traces of oxygen and through an aliquot of the buffer solution before being dispersed in the cell through a coarse-porosity sintered-glass disk. Alt1ea.t 20 minutes were allowed for initial deaeration. d w ~ r b i cacid powder (Baker Analyzed Reagent) $7 as added to the electrolybis cell after deaeration. RESULTS AND DISCUSSION
Ascorbic acid is studied most conveniently in acid or neutral solutions because of stability problems in the pH range above about 8 ( 2 ) . In this work experiments were carried out in 1.0X pH 7.2 phosphate buffer solutions. The high ionic strength of these solutions was necessary to obtain sufficient conductivity to minimize uncompensated ohmic losses ( I ) . Preliminary cyclic voltammetric experiments were carried oul rvith the ascorbic acid system varying the wan rate up to 700 volts per second. Cyclic voltammetric theory (8) for the case of a reversible charge transfer followed by an irreversible chemical reaction predicts that the peak separation for the anodic and cathodic processes, AEp, should remain constant a t about 60/n
mv. At sweep rates where the reverse peak could be observed in the studies here, AE, remained constant a t approximately 30 mv. for scan rates up to 400 volts per second. Thus, the electron transfer process can be considered reversible a t least up to that point, indicating that kinetic data can be obtained with confidence in this range using cyclic voltammetry. There are two possible explanations for an increase in A E , as the scan rate is increased further, First, the rate of the charge transfer process itself may become important a t the faster scan rates, resulting in peak separation ( 7 ) . Second, uncompensated ohmic losses become more important at the higher scan rates and can cause peak separation, even though the charge transfer step is rapid (6).
At least two methods are available for estimating the uncompensated resistance, R,. One of these involves a direct calculation from the electrolyte conductivity and reference probe placement, as described by Booman and Holbrook ( I ) . Another approach provides an estimate of the upper limit of R, by an investigation of the peak potential separation a t a scan rate where the first measurable separation greater than the reversible case of 60/n inv. occurs (IO). For the system reported here, an upper limit for R, of 1.0 to 1.5 ohms was estimated from the peak separation a t 400 volts per second. This value of R, is consistent with calculations based on the suggestions of Booman and Holbrook ( I ) , and is in agreement with values obtained under similar conditions (7, I O ) . Thus, using this value of R,, the VOL. 38, NO. 12, NOVEMBER 1966
1761
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Figure 2. Potential-step current-time curves for ascorbic acid
Figure 3. Experimental blanks for potential-step current-time curves 1M p H 7.2 phosphate buffer; vertical scale, 1.0 rna./div.; zontal scale, 0.1 rnsec./div.
1.0 X 1 0 3 4 ascorbic acid, 1M pH 7.2 phosphate buffer. Upper trace cathodic process; vertical scale, 0.5 rna./div.; horizontal scale, 0.1 msec./div. Center vertical line corresponds to the switching time, r lower trace: anodic process; vertical scale, 1 .O ma./div.; horizontal scale 0.1 msec./div.
maximum distortion of anodic and cathodic peaks could be estimated (6)' Such distortion effects are negligible for scan rates less than 400 volts per second. If the excessive separation of peak potentials was caused by charge transfer limitations in the ascorbic acid system, the acquisition of chemical kinetic data
would be limited. When the charge transfer step is not reversible, the theory developed by Nicholson and Shain (8), on which kinetic results are based, is not applicable for the case of a succeeding irreversible reaction. However, because the peak separation does not exceed the reversible value until a scan rate of 400 volts per second is reached, kinetic data acquired at lower scan rates are not subject to this limitation. Because the potential-step technique involves stepping the potential to and from the diffusion-controlled region of
Kinetic Data from Cyclic Voltammetry of 1 X in 1 .OM p H 7.2 Phosphate Buffer
Table I.
Scan rate, volts/sec.
i&a
0.45 0.55 0.66 0.73 0.77
100 200 300 400 500
I,
see. X 104" 10.0 5.0 3.3 2.5 2.0
kr' 1.283 0.775 0.475 0.336 0.271
Ascorbic Acid
k , sec-1 x 10-3 1.28 1.55 1.43 1.34 1.35
a For these experiments, is defined as the time difference between B/Iand the switching potential. From Nicholson and Shain (8).
Kinetic Data from Potential-Step Experiments with 1 Acid in 1 .OM pH 7 . 2 Phosphate Buffer
Table 11. 7,
x
see.
104 5.0 5.0 10.0 10.0
i& 0.36 0.26 0.22 0.17
Time ratio, (t - r ) / P 0.2 0.3 0.2 0.3
krb 0.60 0.72 1.32 1.22
X 1O-3M Ascorbic k , see.-' x 10-3 1.20 1.45 1.32 1.22
a Data refer t o three replicate determinations a t each switching time, r . t refers to the total time elapsed since beginning of experiment. From Schwarz and Shah (14). ~
~~
1762
~
ANALYTICAL CHEMISTRY
hori-
each half of the redox couple, the factors considered above limiting kinetic studies are not so important as with cyclic voltammetry. First, no charge transfer limitations are placed on the potential-step technique because data are obtained at potentials in diffusioncontrolled regions. Second, uncompensated ohmic losses are of lebs importance in the diff usion-controlled region of a voltammetric wave because small uncertainties in potential will not affect the current-time curve. Furthermore, considering the estimated value of R,, the potentiostatic rise time is the order of 10 psec., considerably less than the switching times wed here. Kinetic Studies with Cyclic VoltamKinetic data for the >ucmetry.
ceeding chemical reaction were obtained with cyclic voltammetry by comparison of the peak ratios Ivith the criteria tabulated by Sicholson and Shain (8). Data were obtained for several different scan rates, where a single trangular wave of 0.500-volt amplitude was applied to the electrolysis cell. The initial potential of the working electrode was -0.400 volt us. S.C.E., and Ex, the switching potential, was +0.100 volt, which was sufficiently anodic of E l l L to allow application of theory (8). A cyclic polarogram obtained with ascorbic acid a t a zcan rate of 200 volts per second ii s h o w in Figure 1. Cathodic peak heights were measured as suggested previously (8), and peak heightb were corrected for charging current. Experimental results are summarized in Table I. The average value o€ the first-order rate constant for the succeeding reaction ib 1.39 X lo3second+ with
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a relative standard deviation of &7.3%. The agreement of the data at 400 and 500 volts per second with the data a t sloaer scan rates indicates that the distortion of the cyclic waves is not yet significant a t 500 volts per second. This may be because of compensating dijtortions in both the anodic and cathodic peaks. Although data were obtained with scan rates up to 700 volts per second, those data mere not reported because peak separations exceeded 45 niv. and the curves were obviously distorted. Kinetic Studies with PotentialStep Electrolysis. Kinetic data were obtained from the potential-step experiments by comparison with the theory of Schwarz and Shain ( 1 4 ) . For an individual experiment a single square wave pulse of 0.350-volt amplitude was applied to the electrolysis cell. The initial potential of the working electrode was -0.250 volt us. S.C.E.. and the potential was stepped to +0.100 volt, nhich h a s well into the diffusionlimited oxidation of ascorbic acid. After a preset time, T , the potential rrverted to the initial value, where the oxidized product was reduced. Anodic
and cathodic electrolysis currents were compared to experimental blanks to obtain net values. Typical currenttime curves are shown in Figure 2, and experimental blanks are shown in Figure 3. Results are summarized in Table 11. An average value of the rate constant of 1.31 x 103 second-’ was calculated, with a relative standard deviation of 11.2%. The close agreement between the kinetic data obtained here with the two different electrochemical techniques does not appear to be fortuitous, considering that consistent results were observed with varying degreeq of time resolution for each technique. Thus, the agreement of results serves l o promote confidence in the values of the firstorder rate constant reported.
(4)Kern, D. >I. H., J . Am. Chem. Soc. 76, 1011 (1954). (5) SIeites, L., ilnal. Chim. Acta. 18, 364 (1958). (6) Yicholsoii, R. S., ASAL. CHEW 37, 667 (1965). 17) Ibid., p. 1351. (8) Sicholsori, R. S., Shain, I., Ibid., 36, 706 (1964). (9) Ono, S.,Takagi, M., M’asa, T., Bull. Chern. Soc. Japan 31, 356 (1958). (10) Perone, S. P., ANAL.CHEM.38, 1158 (1966). (11) Perone, S. P., Mueller, T. R., Ibid., 37, 2 (1965). (12) Perrin, C. L., “Progress in Physical Organic Chemistry,” S. G. Cohen, A. Streitwieser, Jr., K. W, Taft, eds., 1-01. 3, p. 165, Interscience, New York, 1965. (13) Schwarz, W. AI., Shain, I., ASAL. CHEM.35, l77p (1963). 114) Schwarz, It. II.,Shain, I., J . Phys. Chern. 69, 30 (1965). (15) Vavrin, Z., Collection Czechoslovok. Chem. Conirnuns. 14, 367 (1949).
8. P. PERONE
W.J. KRETLOW
LITERATURE CITED
(1) Booman, G. L., Holbrook, R. B., A s . 1 ~ CHEW . 35, 1793 (1963). (2) Brezina, AI., Zuman, P., “Polarog-
raphy in Medicine, Biochemistry. and Pharmacl-,” rev. English ed.. Chtp. XXIT, Interscience, New York, 1938. (3) Jaenicke, W., Hoffman, H., Z.Electrochernze 66, 814 (1962).
Department, of Chemistry Purdue University Lafayette, Ind. 47007 WORK supported in part by Purdue Research Foundation. Division of Analytical Chemistry, 15211d AIeeting, ACS, Xew York, September 1966.
Determination of Refractory Oxide Elements by Atomic Absorption Spectrometry with the Plasma Jet SIR: In atomic absorption spectrometry and flame photometry, determination of those elements that form thermally stable molecular species has been limited by the lon concentration of free atoms, the formation of molecular ovides being particularly troublesome in this regard. T o minimize this difficulty, various workers in atomic absori7tion spectrometry have attempted to increase the atom concentration by raising flame temperature, reducing available oxygen concentration in the flame, and facilitating side reactions that rvill release the metal from the oxide ( 1 , 5,4, 6). -1 inore attractive alternative would be to use a sj-stem which provides an inert atmosphere. Such systems have been studied ( 6 ) , but are not readily amenable to samples in solution form. Recently, Wendt and Fassel, ( 7 ) have reported the use of the induction coupled plasma as a n atomizer. We have been norking u i t h a plasma jet, and our results agree with their conclusions that the formation of refractory oxides is virtually eliminated when a plasmabased system is utilized. EXPERIMENTAL
Apparatus. Atomizer: A plasma jet solution analyzer (Spex Industries, Inc.) powered by a P-26 source control unit (Baird-Atomic, Inc.) was
used for sample atomization. The tangential gas n-as helium a t 80 c.f.h. and the aspirating pas was argon a t 14 p s i . A wide-bore aspirator (Beckman Instruments, Inc.) n-as used to introduce solutions a t about 1 ml.,’ min. Plasma current was 17 amp., and the gap between cathode and upper control ring was 0.7 em. The effective path length was about 0.5 em. Optics. The spectrometer was a Perkin-Elmer Corp. Model 214 atomic absorption spectrometer, modified for single beam operation by removal of the beam splitter. Slit width was 25 microns and an R C h 1P 28 photomultiplier tube was the detector. Electronics. The 60 c.p.3. signal from the half-wave modulated hollow cathode was amplified by the sample
Table 1.
channel of the Perkin-Elmer autoratiometer, further amplified and tuned by an SP656 operational amplifier (Philbrick Researches) with a 60-C.P.S.twin-T filter in the negative feedback loop, and rectified prior to recording on a Fisher Recordall.
RESULTS AND DISCUSSION
The sensitivity found for Ak1(?;0a)3 in 95%) ethanol was 2.7 p g , of Xl/ml. for 1% absorption using the 3092.7-1. line. The sensitivity found for vanadyl acetylacetonate in methanol was 18 p g . of V/nil. using the 3185.4-A. line. Relative standard deviations were about 6%. Work is under progress for the deter-
Electronic Energy Distribution for Lanthanum and Aluminum
Energy level, c m . 3
3000’ K.
S i / S Lanthanum
0.00 1053.2 2668.2 3010.0 3494.6 4121.6 7011.9 0.00 112.04
4 0 0 0 5000
6000
7000
8000
0.29 0.26 0.08 0 10 0.11 0.10 0.02
0.22 0.23 0.09 0.11 0.13 0.13 0.03
0.16 0.18 0.08 0.12 0.14 0.15 0.04
0.14 0.17 0.08 0.11 0 14 0.15 0.05
0.13 0.16 0.08 0.11 0.14 0.16 0.06
0.34 0.66
0.34 0.66
0.34 0.66
0.34 0.66
0.18 0.20 0.09 0.11 0.14 0.14 0.04
S i /SAluminum 0.34 0.66
0.34 0.66
VOL. 38, NO. 12, NOVEMBER 1966
1763
