of oxidation imposed by control of the working potential. The polarographic data proved to be a valuable aid in the select,ion of electrolysis potentials. There were no significant differences in t,he currentvoltage relationships obtained with the microelectrode and with the macroelectrode. Current-Time Curves. T h e general charact.er of t,he current-time curves for the acid concentrations used is shown in Figures 3 and 4. Characteristic curves for the 1-electron oxidations occurring in 1 2 5 ac,id, illustrative of typical initial currents and electrolysis times, are shown in Figure 3. T h e current-time curves for the two-step oxidations (Figure 4) in both 1 S and 1212' acid exhibit two linear segnients with different slopes, representative of the t,wo consecut'ive oxidations. In the early part of the electrolysis the reaction expressed by Equat,ion 1 (or 3) predominates. This is indicated by the first linear portion. As the concentration of radical builds up, the react,ion expressed by Equat'ion 2 is of primary significance and is the main determinant of the slope of the second line segment. Electrode Oxidation
Mechanism.
Some observations were made regarding the nature of t h e electrode oxidat,ion mechanism. Pretreatment of the platinum macroelectrode result,ed in the open circuit oxidation of the reduced phenothiazine compounds. When the plat,inum electrode was polarized anodically for a few minutes a t normal working potentials and then
the circuit opened and the sample to be oxidized introduced, any material that came in direct contact with the electrode surface was immediately oxidized. This was indicated by a color changeviz., the reduced colorless compound was oxidized to the colored radical intermediate. This effect could be repeated only when the electrode was again subjected to brief anodic polarization. The amount of material transformed under open circuit conditions was found t o be negligible when compared to the quantities that were taken for analysis. The work of Feldberg, Enke, and Bricker ( 3 ) dealing with the formation and dissolution of platinum oxide films offers an explanation for t,his observation. These workers demonstrated anodic film formation on platinum through two single-electron steps-a very slow step, followed by a fast reversible step. The latter results in an active electrode state. This work leads us to propose that the act,ive electrode state is responsible for the oxidation of phenothiazine compounds observed under the open circuit conditions described. -1 similar situation may exist under closed circuit conditions; the active electrode state is maintained with consequent oxidation of the electroactive material. hdditional evidence supporting this oxidation mechanism was observed during the polarographic studies. Erratic oxidation waves were obtained with an untreated platinum microelectrode. However, when the electrode was first, oxidized and then reduced under controlled-potential conditions in 1.V or
12'T acid, reproducible polarographic waves were obtained. The electrode produced reproducible waves during its entire period of use after one such treatment. ACKNOWLEDGMENT
The authors thank Smith, Kline & French Laboratories, Wyeth Laboratories, Squibb Laboratories, and Sandoz Laboratories for the generous supplies of phenothiazine cornpounds made available for this work. The authors give credit' to the Research Council of Rutgers-The State University for financial assistance which enabled the constructmionand purchase of much of the instrumentation used for the coulometric work. F. H . Merkle thanks the Johnson & Johnson Laboratories for support in the form of a research fellowship for the 1963-64 academic year. LITERATURE CITED
( 1 ) Borg, D. C., Cotzias, G. C., Proc. iVatl. A c a d . Sci. 48, 623 (1962). ( 2 ) Zbid., p. 643. ( 3 ) Feldberg, S. W., Enke, C. G., Bricker, C. E., J . Electrochem. Soc. 110, 826 (1963). (4) Kabasakalian, P., McGlotten, J., ANAL.CHEM.31,431 (1959).
( 5 ) Kelly, M. T., Jones, H. C., Fisher, D. J., l b i d . , 31, 488 (1959). (6) Merkle, F. H., Discher, C. A., J. Pharm. Sci., in press. ( 7 ) Piette, L. H., Ludwig, P., Adams, R. N., ANAL.CHEM.34, 916 (1962) (8) Ryan, J. A , , J . A m . Pharm. Assoc. 48,240 (1959).
RECEIVED for review February 14, 1964. Accepted April 23, 1964.
Catalytic Determination of Vanadium in Water M. J. FISHMAN and M. W. SKOUGSTAD U. S. Geological Survey, Denver, Colo.
b A rapid, accurate, and sensitive spectrophotometric method for the quantitative determination of trace amounts of vanadium in water is based on the catalytic effect of vanadium on the rate of oxidation of gallic acid by persulfate in acid solution. Under given conditions of concentrations of reactants, temperature, and reaction time, the extent of oxidation of gallic acid is proportional to the concentration of vanadium present. Vanadium is determined by measuring the absorbance of the sample at 415 mp and comparison with standard solutions treated in an identical manner. Concentrations in the range of from 0.1 to 8.0 pg. per liter may be determined with a standard deviation of 0.2 or
less. By reducing the reaction time, the method may b e extended to cover the range from 1 to 100 pg. with a standard deviation of 0.8 or less. Several substances interfere, including chloride above 100 p.p.m., and bromide and iodide in much lower concentrations. Interference from the halides is eliminated or minimized by the addition of mercuric nitrate solution. Most other substances do not interfere at the concentration levels at which they commonly occur in natural waters.
S
and spectrophotometric techniques are the principal means currently used to determine vanadium in natural waters. Either PECTROGRAPHIC
technique requires concentration of vanadium by precipitation, evaporation, or extraction, to detect trace quantities of t'he element. Haffty ( 1 ) and Silvey and Brennan (4) described general spectrographic procedures suitable for determining vanadium and other minor elements in water. Spectrophotometric methods have been reported by Xaito and Sugawara ( 3 ) and by Sugawara, Tanaka, and S a i t o ( 5 ) . A rapid, qualit'ative test for trace amounts of vanadium was recently reported by Jarabin and Szarvas ( 2 ) . Their test is based on the fact that the oxidation of gallic acid by acid-persulfate is catalyzed by minute amounts of vanadium. In the presence of vanadium the reaction proceeds rapidly, VOL. 36, NO. 8, JULY 1964
1643
whereas the reaction is very slow when vanadium is absent. The reaction mixture develops a yellow to red color, depending upon the amount of vanadium present. As little as 0.025 pg. of vanadium can be detected. The interference effects of a number of elements were investigated by Jarabin and Szarvas ( 2 ) and found, for the most part, to offer no interference a t concentrations up to 2000 times the concentrations of vanadium. Cr(II1) up to 1600 times and Mo(V1) up to 100 times the concentration of vanadium did not interfere with the qualitative test. Quantitative procedures based on the catalytic property of the constituent being determined have been developed for the determination of trace concentrations of several substances. It seemed likely that the qualitative reaction described by Jarabin and Szarvas ( 2 ) could be made the basis of a very sensitive quantitative method for vanadium, and it was, therefore, investigated in detail. Freshly prepared ammonium persulfate solution oxidizes gallic acid only slowly, even when a relatively large amount of vanadium catalyst is present. This suggests that the catalytic oxidation depends on a reaction involving one of the hydrolysis products of the peroxydisulfate ion. Peroxymonosulfuric acid, the first hydrolysis product, is a very strong oxidizing agent, and it is this acid which is active in the catalytic oxidation. I t is proposed that the reactions involved in the catalytic oxidation of gallic acid are as follows: first, vanadate ion in acid solution forms the pervanadyl ion (V02+);second, gallic acid is oxidized by both peroxymonosulfuric acid (slow) and pervanadyl ion (fast); and finally, the resulting vanadyl ion (VO+*) is immediately reoxidized by peroxymonosulfuric acid. EXPERIMENTAL
Reagents. Ammonium metavanadate solution (1 ml. contains 0.010 pg. of vanadium) ; mercuric nitrate solution (0.035%) ; gallic acid solution (2%). Ammonium persulfate-p h o s p h o r i c acid solution is prepared as follows. Dissolve 2.5 grams of ammonium persulfate in 25 ml. of demineralized water. Bring just to boil. Remove from heat and add 25 ml. of concentrated phosphoric acid. Let stand approsimately 24 hours before using. Procedure. Prepare a blank and sufficient standards by transferring 0.0- to 8.0-ml. aliquots (0.00 to 0.08 pg, of vanadium) of the standard vanadium solution into matched 23-ml. abqorption cells (appros. 23 mm. in diameter), or other suitable absorption cell or container. Dilute each aliquot to 10 ml. with demineralized water. Pipet a volume of water sample containing less than 0.08 p g . of vanadium (10 ml. mas.) into an absorption cell 16 44
ANALYTICAL CHEMISTRY
t
'*0° 0.70 0.50
0.40 0.30 E O
n L
:: 0.20
n
4
0.I O
-
0.00
340
380
420
460
Wavelength, mu
Figure 1 .
Absorption spectra of ("3)~[A), gallic acid (B), and vanadium plus reagents (C)
S20g-H3P04
and adjust the volume to 10 ml. with demineralized water. Add 1.0 ml. of mercuric nitrate solution each to the blank, standards, and samples. Place the cells in a water bath regulated to 25' + 0.5' C. and allow 30 to 45 minutes for the samples to come to the temperature of the bath. Add 1.0 ml. of ammonium persulfate-phosphoric acid solution (temperature equilibrated) , swirl to mix thoroughly, and return to the water bath. Add 1.0 ml. of gallic acid (temperature equilibrated), swirl to mis thoroughly, and return to the water bath. Add the gallic acid to successive samples a t intervals of 30 seconds or longer, to permit accurate control of the reaction time for each individual sample. Exactly 60 minutes after the addition of the gallic acid to a sample, remove it from the water bath and immediately measure its absorbance a t 415 mp, using distilled water as a reference. Colored or turbid samples must be decolorized or filtered before analysis. Construct a calibration curve by plotting the absorbance values of standards us. micrograms of vanadium. Determine the amount of vanadium in an unknown sample by reference to the corresponding absorbance on the calibration curve. A calibration curve must be prepared with each set of samples. RESULTS A N D DISCUSSION
Absorption Spectra. Absorption spectra (Figure I ) were obtained for the ammonium persulfate-phosphoric acid reagent alone (-A), the gallic acid reagent alone ( E ) , and a mivture of these two reagents after reaction in the presence of a trace of vanadium (C). A neckman Model 1113 spectrophotometer with 10-mm. cells (air
as reference) was used to obtain these curves. .Ilthough 415 mp occurs on the steep portion of the curve, this wavelength was chosen because of the maximum difference in absorbances between the unoxidized and oxidized gallic acid, and the low absorbance of the blank. Also, standards are included wit,h each set' of determinations. Stability of Reagents. The gallic acid solution used in the procedure is best prepared by adding the reagent to warm, distilled water, heating to a temperature just below boiling, and filtering the hot solution through Whatman No. 42 paper, or equivalent'. Filtration is necessary to remove traces of undissolved gallic acid which, if not removed, promote rapid precipitation of gallic acid. Even in filtered solutions, a fine precipit,ate of gallic acid appears after about 24 hours. For this reason, the reagent must, be freshly prepared for each set of samples analyzed. Samples cont'aining added known amounts of vanadium (0.0 to 10.0 pg. per liter) were treated with identical amounts of ammonium persulfatephosphoric reagent which had been allowed to age for from 1 to 72 hours. A11 other conditions were held constant': concentrat'ion of gallic acid, reaction time, t,emperat'ure, etc. The dat'a obtained show that the maximum effectiveness of this reagent is achieved after about 24 hours and that deterioration begins after about 48 hours. Therefore, it is necessary that the reagent be prepared at' least' 24 hours before its intended use, and that it should be discarded after 48 hours. Temperature. The rate of catalytic oxidation of gallic acid by a n acid persulfate solution is markedly affected by temperature. It was found t h a t temperatures a t or near 25" C. provide escellent sensitivity for trace amounts of vanadium within a reasonable reaction time. Because t'he reaction rate is so markedly affected by temperature, it is necessary to control the temperature of the reaction to within 0.5' C., and a regulated water bath is essential. Concentration of Reagents, Reaction Time. I n order to establish optimum concentrations of reagents and react'ion time, varying amounts of 27& gallic acid solution and ammonium persulfate-phosphoric acid reagent were successively added to a series of standard solutions cont,aining 0.00 to 0.10 pg. of vanadium. The reaction time was varied from 30 to 60 minutes to obtain reasonable absorbance readings for each concentration. Optimum reaction conditions were obtained when 1.0 ml. of t,he acidpersulfate reagent and 1.0 ml. of 2% gallic acid solution were added to a 10-
Table
I.
Effect of Added Ions on Determination of Vanadium
A S o interference occurs for the following elements at the concentrations listed:
Element added
Added as
Sodium Na2HP04 Potassium KNO3 Calcium CaS04 Maenesium hleSO1 Sulfite (NH4);so4 Kitrate K"O8 Aluminum AlCl3 Manganese( 11) MnS04 Zinc ZnClp Arsenic(II1) As203 Selenium( IV) H2Se03
Concn. of element, mg./l. 1000 600 500 1000 1000 1000 5.0 5.0 5 0 5 0 5 0
B Interference ocrurs for the following rations when concentrations listed are exceeded: Concn. of Element element, added Added as mg./l. Silver Uranium( VI) CobaltiII) Iiickel(I1 j Copper( 11) Chromium(l'1) Molybdenum (VI) Iron( 11) Iron( 111)
Ag?jOs UOz( C2HLh)z CoCL NiCl2CuC12 K2Ci-207
2.0 3.0 1 .o 3.0 0.05 1.0
MOO3
Fe(NHaLFe( NHa)2-
0.1 0.3
(SO4)Z FeSH4(S04)2
0.5
C Interference occurs for the following anions when concentrat,ions listed are exceeded: Concn. of Element element, added Added as mg./l. Chloride Bromide Iodide
SHaCl KBr KI
100 0 10 0 001
ml. aliquot of sample, and the reaction allowed to continue for 60 minutes a t 25" i. 0.5' C. The blank value was reasonably low, indicating slow oxidation of gallic acid in the absence of catalyst, and the data provided a calibration curve approaching a straight line. Because the absorbance readings generally fall off rapidly above 0.08 or 0.10 pg. of vanadium, calibration curves were limited to the range of from 0.00 to 0.08 p!. of vanadium. The concentration range can be extended tenfold by reducing t,he reaction time to 10 minutes while all other conditions are kept constant. Although the calibration curve is not linear over this range of vanadium concentrations, it does provide a satisfactory determination of vanadium, providing the calibrat,ion curve is established each time a set of samples is analyzed. Interferences. Table I summarizes the results of a n evaluation of inter-
ferences in which the substances noted were added to a standard solution containing 4.0 pg. of vanadium per liter. If the measured value of vanadium differed from the calculated value by more than 0.4 pg. of vanadium per liter, the added substance was considered to contribute intolerable interference. No interference was noted for the substances listed in Part A of Table I , a t the concentrations specified. Concentrations greater than those specified were not tested because these substances do not commonly occur in fresh water a t concentrations exceeding these values. The substances in Part B of Table I will interfere in the quantitative determination of vanadium if the specified concentrations are esceeded. This is not a serious problem for Ag, U(VI), Co(II), Ni(II), Cr(VI), and Mo(VI), since the tolerable concentration in each case is greater than that commonly encountered in fresh water. On the other hand, the tolerable concentration of Cu(II), and Fe(I1) or Fe(II1) may be expected to be exceeded in some samples. Because of the high sensitivity of the method, interfering substances present in concentrations only slightly above t,olerable limits, can be rendered harmless by simple dilution of the original sample. Traces of iodide and bromide (Part C of Table I) interfere seriously, and simple dilution of the sample will not always reduce the concentration of these ions to values below tolerance limits. Mercuric ion may be added to complex these halides and minimize their interference; however, Hg(I1) ion itself interferes if an excess is present. The addition of 1 ml. of 0.03570 mercuric nitrate solution to each sample'does not interfere, and permits the determination of vanadium in the presence of up to 100 p.p.m. of chloride ion, and up to 0.25 p.p.m. each of bromide and iodide. If bromide and iodide are absent, up to 200 p.p.m. of chloride may be tolerated. Samples containing more than 0.25 p.p.m. of bromide or iodide, or more than 100 p.1i.m. of chloride, must be diluted to reduce the concentrations of these ions below the above values and mercuric nitrate must be added. Accuracy and Precision. Twentyone natural-water samples were analyzed for vanadium by the proposed method. T h e results obtained are shown in Table 11. An indication of the accuracy of the method was determined by re-analyzing these samples after a known amount of vanadium had been added to each. T h e amounts of vanadium found were in good agreement with the amounts calculated to be present. For samples containing less than 8.0 pg. per liter, the difference between the cal-
Table 11. Analysis of Samples of W a t e r Containing Known Added Vanadium
Vanadium, pg. per liter Water Origisample nally No. present Added Totala Found 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16
0.4 9.5 0.8 2.3 82 0.4 13 15 0.0 3 8 0 2 0 5 0 5 0 4 0 4 1.5
1.0 5.0 1.0 1.0 50.0 1.0 5.0 5.0 1.0 1 0 1 0 1 0 1 0 1 0 1 0 1.0
1.4 14 1.8 3.3 130 1.4 18 20 1.0 4 8 1 2 1 5 1 5 1 4 1 4 2.5
1.4 15 1.8 2.8 140 1.4 18 20 0.9 4 6 1 2 1 7 1 6 1 8 1 8 2.3
Two significant figures only.
culated values and observed values was found to range from 0.00 to 0.5 pg. per liter. For samples containing more than 8.0 pg. per liter, excluding sample No. 5, the greatest difference between the calculated and the observed values was only 1.0 pg. per liter. Sample No. 5, after adding 50.0 p g . of vanadium per liter, was Falculated to contain 130 pg. per liter. The concentration of vanadium found was 140 pg. per liter. Replicate determinations were made on four natural-water samples (Table 111). The standard deviation using the
Table 111.
Analysis of Natural-Water Samples
Water sample Yandium, No. Found
pg.
per liter Average
Std. dev.
la
0.6 0 9 0.8 10 0 8
0 0 0 0 0
9 6 9 5 7
0 8
fO 2
2O
6.5 6.8 6.5 6.6 7.0
6.8 7.0 6.9 7.1 6.8
6.8
f0.2
40
91 93
92 (32
92
10 8
Calibration: 0.00 t o 0.08 H g . of vnnadium. Calibration: 0.0 to 1.Opg. of vanadium.
VOL. 36, NO. 8, JULY 1964
1645
0.00 to 0.08 pg. vanadium calibration curve was 1 0 . 2 for samples 1 and 2 which contained, respectively, 0.8 and 6.8 pg. of vanadium per liter. Using a 0.0 to 1.0 pg. vanadium calibration curve, the standard deviations for samples 3 (18 pg. per liter) and 4 (92 pg. per liter) were 1 0 . 6 and 1 0 . 8 , respectively.
LITERATURE CITED
(1) Haffty, J., U . S.Geol. Survey WaterSupply Paper 1540-A,1 (1960). (2) Jarabin, Z . , Szarvas, P., Acta b'niv. Debrecen 7, 131 (1961); ( ' . A . 57, 9192e (1962). ( 3 ) Naito, H., Sugawara, K., Bull. Chem. Sac. Japan 30, 799 (1957); C . A . 52, 5706i (1958).
(4) Silvey, FV. D., Brennan, R., ANAL. CHEM.34,784 (1962). ( 5 ) Sugawara, K., Tanaka, M., Saito, H., Bull. Chem. Sac. Japan 26, 417 (1953); C.,4.49. (3842b(1958). RECEIVEDfor review Rlarch 9, 1964. Accepted April 27, 1964. Division of Water and Waste Chemistry, 147th Meeting, ACS, Philadelphia, Pa., April 1964. Publication approved by the Director, L.S. Geological Survey.
Fluorometric Method for the Determination of Urea in Blood JOSEF E. McCLESKEY Clinical Chemistry Branch, U . S. Naval Medical School, National Naval Medical Center, Bethesda, Md.
b It has been observed that the compound produced by the reaction of urea with diacetyl monoxime exhibits fluorescence. A study of this fluorescent property has resulted in the development of a quantitative procedure for the determination of urea. The variables studied include the method of deproteinization of blood and serum, the time of heating, the concentration of diacetyl monoxime and the effect of pH. Comparison studies with the AutoAnalyzer and recovery studies show the method to b e valid.
T
fop a simple and accurate method for quantitatively determining urea nitrogen in blood has provided impetus for the reevaluation of several urea nitrogen procedures. A variety of methods have been developed utilizing enzyme reactions to form ammonia with subsequent titration (3). Several objectionable features to these methods are evident, via., loss of ammonia and time consumed. Several methods have also been published which utilize the reaction of urea with diacetyl ( 4 , 6), diacetyl monoxime ( 5 , 7-10)] or a-isonitrosopropiophenone (1) in acid media to form a colored compound. Color intensities are then measured to yield urea concentration. The methods using diacetyl or diacetyl derivatives usually suffer from the fact that the color produced does not conform to Beer's Law and is photosensitive. Experimentation oriented toward develoliment of a blood urea nitrogen method with more desirable characteristics led to the investigation, in our laboratory, of the method of Richter and Lapointe (9). During the course of this work a fluorescent property of the ureaalpha diketone compound was disHE NEED
1646
ANALYTICAL CHEMISTRY
millilit,ers of water. Dilut'e to 100 ml. This reagent is stable for several months at' room temperature. DI~CETY ~ ~LO N O X I M E .Dissolve 5.0 grams of diacetyl monoxime (Eastman Organic Chemicals) in 500 ml. of distilled water. Add 150 grams of sodium chloride, 100 ml. of distilled water and shake well until dissolved. Dilut'e to 1000 ml. a n 3 filter. If stored in an EXPERIMENTAL amber bottle a t room temperature this The fluorescent properties of the comreagent is stable for about six weeks. pound formed by the reaction of urea Procedure. The recommended proand diacetyl monoxime were studied cedure for the fluorometric determinawith a n Aminco-Bowman spectrophototion of blood urea nitrogen is as fluorometer. The fluorescent peak was follows: observed a t a wavelength of 415 mp Add 0.2 ml. of whole blood or serum when the compound was activated a t an to 4.0 ml. of distilled water. Mix well optimal wavelength of 380 millimicrons. and allow to stand for 10 to 15 minutes. Reagents and Solutions. VREA Add 5.0 ml. of 30Y0 trichloroacetic acid NITROGENS TOCK STANDARD. 1 MG. to precipkate the prot,ein. Mix well PER ML. Weigh out 2.14 grams of and centrifuge at 2500 r.p.m. for 15 urea (c.P., A.c.s.) and transfer to a minutes or, alt,ernat'ively, filter t'hrough 100-ml. volumetric flask. Add 50 ml. Whatman #40 filter paper. Pipet 2.0 of 0.Ol.V sulfuric acid and swirl to ml. of the protein-free filtrate into a dissolve. Then dilute to 100 ml. screw cap test t'ube (16 X 125 mm.). with 0.01,V sulfuric acid. .4dd 2.0 ml. of diacetyl monoxime and UREA NITROGENWORKINGSTAXD- 0.30 ml. of concentrated sulfuric acid. ARD, 40 MG. PER100 ML. Dilute 4 ml. Screw the cap down tightly and mix of the stock standard to 100 ml. with well. Heat the tube in a boiling water 0.01N sulfuric acid solution. bath for 15 minutes. After heat'ing, TRICHLOROACETIC ACID REAGENT, release the pressure by loosening t'he 30%. Dissolve 30 grams of reagent cap, mix well, and allow to cool to grade trichloroacetic acid in a few room temperature. Run a standard (40 mg. per 100 ml.) in the same manner by using 0.2 ml. of the working standard in place of blood or serum. Measure Table I. Typical Calibration Data for the fluorescence in the spectrophotoDetermination of Urea in Blood and fluorometer a t a wavelength of 415 mp Blood Serum with an activation wavelength of 380 Urea nitrogen Fluorometer mp, Use the 40 mg. per 100 ml. concn. mg./100 ml. reading standard to adjust the instrument to a predetermined fluorescence. An 8070 0 0 0.0 adjustment is used in t8his laboratory. 9 0 10.0 18 0 15.0 Read t,he results from a previously 20.0 26 0 prerared calibration curve. 37 0 25.0 51 0 30.0 RESULTS A N D DISCUSSION 64 5 35.0 40.0 80 0 Table I shows a typical calibration Each sample contained 2.0 ml. of diadata using the procedure described cetyl monoxime and 0.30 ml. of concenabove, but substituting standard urea trated sulfuric acid. nitrogen solutions for blood and serum covered. Subsequent work indicated that this fluorescence could be adapted to a quantitative procedure for urea. A more complete study of reagents and reagent concentrations provided an accurate, simple, and reproducible procedure for urea nitrogen in blood.
