ANALYTICAL EDITION
April, 1945 Table
IV.
Effect
of Water Content of Acetone on Recovery of Water from Methane
Water Preaent in Water Added to Dried Methane Acetone Mg. Mg./ml. 2.60 23.7 2.20 19.9 1.56 22.4 0.92 37.6 0.66 22.3 0.62 26.9 0.22 18.3 0.16 32.4 a Corrected for water present in acetone.
215
The amount of water vapor employed was determined from the difference in weight of the humidifier before and after passage of the gas. The entire known was used in each determination.
Wdter Founda
ACKNOWLEDGMENT
Mg. 19.1 10.5 13.8 25.0 12.4 27.4 17.7 31.8
V = volume of sample in cu. ft. G = gravity of sample P = barometric pressure average meter pressure - vapor pressure of water tt t , all expressed as mm. of mercury t = meter temperature, C.
+
The reason for requiring dehydrated acetone in the proposed method is unexplained at this writing. However, numerous experiments have demonstrated the fact that ordinary C.P. acetone is unsatisfactory and that reliable results are obtained if the acetone does not contain more than 0.5 mg. of water per ml. This fact is borne out by Table IV. Knowns were prepared by passing carefully dried gas through
The authors express their appreciation to B. R. Stanerson fer valuable assistance in the initial work. LITERATURE CITED
(1) Bell, R. P., J . Chem. Soc., 1932, 2903. (2) Deaton, W. M., and Frost, E. M., Bur. Mines Rept. Investigat i a s 3399 (May, 1938). (3) Evans, R. N., and Davenport, J. E., IND.ENG.CHEN.,ANAL. ED., 14, 732 (1942). (4) Fischer, K., Angew. Chem., 48, 394 (1935). (5) Hartley, H., and Raikes, H. R., J. Chem. Soc., 127, 524 (1925). (6) Henle, F., Ber., 53, 719 (1920). (7) Natural Gasoline Assoc., Natl. Petroleum News, 24, 38 (April 6, 1932). (8) Perry, c. P.,IND.ENQ.CHEM.,ANAL.ED., 10, 513 (1938). (9) Roberts, R. W., and Fraser, A,, J . SOC. Chem. Ind., 29, 197 (1910). (10) Ross, J., Ibid., 51, 121T (1932). (11) Roth, F., and Schulz, A., Brennstoff-Chem., 20, 317 (1939). (12) Smith, D. M . , and Bryant, W. M. D., J . Am. Chem. Soc., 57, 841 (1935). (13) Todd, F. C., and Gauger, A. W., Proc. A m . SOC.Testing MateriaZs, 41, 1134 (1941). (14) Vtorova, E. I.,'Sintet. Kauchuk, 4, 29 (1936).
a small bubble counter, which in some experiments was warmed,
cont,aining a small amount of water. The water vapor was carried by the gas and passed with it to the dehydrated acetone.
Assay
OF
PREBENTED before the Division of Petroleum Chemistry at the 105th lleeting of the AMERICAN C H E M I C a L SOCIETY, Detroit, Mich.
Lead and Sodium A z i d e by Cerate Oxidimetry J. W. ARNOLDI, Inspection Board of the United Kingdom and Canada
L
EAD azide is a salt of hydrazoic acid, HN, (aziomide). It is a very sensitive and powerful explosive and is used in percussion detonators and in detonators which are ignited by a flash. Various methods for determining the azide value of lead azide have been proposed and used. At least two procedures based on solution of the sample and reprecipitation of the azide as insoluble silver azide with standard silver nitrate solution have been described. Another more widely used procedure is by determination of the nitrogen, based on the reaction of ammonium hexanitratocerate and lead azide according to the following equation: Pb(N3)Z
+ 2(NHdz.Ce(NOde Pb(NOJ2
-
+ 4NH&Os + 2Ce(N03)a + 3S2
The evolved nitrogen is measured in a water-jacketed nitrometer over water (4). This report is concerned nith describing a procedure based on the reaction of ammonium hexanitratocerate and lead azide, the cerate being added to excess and the excess titrated with ferrous sulfate employing o-phenanthroline ferrous complex as the indicator. REAGENTS
Ammonium hexanitratocerate, (yH4)2Ce(N03)j,molecular weight 548.25, 0.1 N in 1 N nitric acid. Dissolve 04.8 grams of ammonium hexanitratocerate in 60 cc. of concentrated nitric acid (70% specific gravity 1.42) and 40 cc. of water. Stir well. Add 4 cc. of nitric acid plus 96 cc. of water and stir well. Dilute slowly t o 1 liter with water, stirring continuously. Protect solution from light ( 2 ) . Ferrous sulfate, 0.1 N in 6 N sulfuric acid. Store solution under hydrogen. 1
Present address, 59 Emerald Crescent, New Toronto, Ontario, Canada.
o-Phenanthroline ferrous complex, indicator solution. T o 1000 ml. of ferrous sulfate, 0.025 molar solution, add 14.85 grams of o-phenanthroline monohydrate, C ~ Z H.-H20 ~ N ~(3). Sodium oxalate (primary standard grade), supply dried a t 100" C. for 2 hours. Potassium dichromate, 0.1 Y solution, prepared from accurate weight of dried reagent. STANDARDIZATION OF AMMONIUM HEXANITRATOCERATE SOLUTION
DIRECThfETHOD WITH SODIUM OXAL.4TE ( 2 ) . Weigh accurately 0.3000 gram of primary grade sodium oxalate, previously dried for 2 hours a t 100" C., and place in a 250-ml. Erlenmeyer flask. Moisten with 10 ml. of distilled water, add 75 ml. of 1.0 molar sulfuric acid, and run in 50 ml. of ammonium hexanitratocerate from a buret. Heat to 50" and allow to cool to room temperature. Add one drop of the indicator and titrate the excess ammonium hexanitratocerate with approximately 0.1 N ferrous sulfate solution prepared as above. Titrate 50 ml. of the cerate solution with the ferrous sulfate solution, following the above procedure.
Calculations.
Where a = ml. ferrous sulfate required to reduce 50 ml. of cerate solution b = ml. ferrous sulfate solution required to reduce excess cerate solution w = weight of sodium oxalate Equivalent weigiit of Xa2C20c= 67.01 w x 1000 Cerate normality = 6 7 . 0 1 X (50 ISDIRECT METHODWITH 0.1 N SOLUTIONOF POTASSIUM DICHROMATE ( 2 ) . Measure out 50 ml. of the dichromate from a calibrated buret or pipet into a 250-ml. Erlenmeyer flask, and dilute to 100 cc. with distilled water. Add 2 drops of the indicator and carry out titration with ferrous sulfate solution, 0.1 N in 6 N sulfuric acid. Having obtained the relationship between the ferrous sulfate solution and the standard dichromate solution,
q)
INDUSTRIAL AND ENGINEERING CHEMISTRY
216
the normality of the cerate solution may be determined by finding the relationship between the cerate and ferrous sulfate. PROCEDURE FOR AZIDE VALUE OF LEAD AZIDE
Accurately weigh 0.3000 gram of lead azide on a small paper scoop and place in a 250-ml. Erlenmeyer flask containing a small amount of water on the bottom to cushion the fall of the azide crystals. Add to the flask 50 ml. of distilled water and 50 ml. of standard 0.1 Ai ammonium hexanitratocerate solution from a calibrated buret. Gently agitate the flask during the addition of the cerate and wash down the sides of the flask with a small amount of distilled water a t the completion of this operation. After addition of one drop of indicator, titrate the excess cerate with ferrous sulfate solution (prepared as above) ; the end point is a sharp change from pale green to reddish yellow. At the end point the solution should be very closely (as practical) 1 molar in sulfuric acid. If it is necessary to add extra sulfuric acid for this adjustment, the previously calculated acid should be added close to the end point. A blank determination is also run. Once the titration is commenced i t should be completed as speedily as possible, because of the presence of plumbous ions in the solution. The cerate solution is more stable than the CALCULATION. ferrous sulfate, and should be considered as the standard solution in the procedure and checked periodically. For ease of calculation, the normality of the ferrous sulfate is computed daily from the blank runs using the existing normality of the cerate. Where a = ml. of ferrous sulfate to reduce 50 ml. of cerate solution of known normality b = ml. of ferrous sulfate to reduce excess cerate solution X = normality of the ferrous sulfate, computed daily from readings obtained a t a w = weight of lead azide taken Pb(N8)z = 145.63 Equivalent weight of lead azide ( a - b) X X X 14.563 Lead azide value =
-
U'
EXPERIMENTAL. The method was carefully checked against two other procedures indicated i n Table I, with one trained operator working under ideal conditions on one sample of lead azide. Table I presents the data from ten consecutive trials by each of three procedures. The standard deviation of the results is calculated by uee of the following equation: 0
=
$E
where d is individual deviation of each result from average (I). The determinations by the cerate method were completed in much less time than those by the other two procedures, owing to shorter manipulation and calculation. SUMMARY AND CONCLUSIONS. The experimental work has shown that the above procedure gives results slightly less reproducible than those obtained by refined gasometric procedure, the standard deviation and range being greater for the cerate methbd. The main advantages of this procedure over the gasometric are its simpler manipulation and shorter calculation. The outlined procedure is shown in experimental work to give more reproducible results than those obtained by the argentometric procedure and to possess the advantages mentioned above. PROCEDURE FOR AZIDE VALUE OF SODIUM AZIDE
This procedure is also adaptable to the assay of sodium azide, used in the manufacture of lead azide. Sodium azide is safe to handle. However, aqueous solutions must not be run to waste without first destroying the azide, because it reacts readily with salts of the heavy metals to give azides of the heavy metalq. many of which are sensitive detonating explosives-e.g., lcad azide. One procedure for determining the azide content is to make a neutral aqueous solution of the sample, precipitate the bulk of the azide as silver azide with standard 0.1 N silver nitrate, filter off the silver azide, and continue the titration with standard 0.1 N silver nitrate, using potassium chromate solution as the indicator.
Table I.
Vol. 17, No. 4
Comparison of Lead A z i d e Values Obtained b y Three Methods
Cerate Gasometric Argentometric Trial No. Nethod .\fethodn Jfethodb 1 97 0 97.1 95 7 2 97 4 97.4 96 1 3 96 9 97 3 95 5 4 96 9 97 3 96 0 5 97.4 97.4 95 5 6 96 9 97.2 96 1 97.1 7 97.3 95 9 97.1 97.4 95 4 8 9 97.1 97.6 95 6 10 96.8 97.4 95 7 Arithmetical average, 97.1 97.3 95.8 Standard deviation, o 0.20 0.13 0.25 High value 97.4 97.6 96.1 Low value 96.8 97.1 __ 95.4 Range 0.6 0.5 0.7 a By measurement of evolved nitrogen from reaction of ammonium hexanitratocerate with lead azide. b Involving precipitation of azide as silver azide. ?A
-
-
REAGENTS, as dcscrihed for lead azide. PROCEDURE. Accurately weigh 0.3000 gram of sodium azide previously dried to 100" C. and cooled, and place i n a 250-ml. Erlenmeyer flask. Dissolve the sodium azide in 100 ml. of distilled water and add 50 ml. of the standard ammonium nitratocerate by buret. Gently agitate the flask during the addition of the cerate and wash down the sides of the flask with a small amount of distilled water a t the completion of this operation. Add one drop of o-phenanthroline indicator and titrate the excess cerate with ferrous sulfate solution; the end point is a sharp change from pale green to reddish yellow. A t the end point the solution should be very closely 1 molar in sulfuric acid. The previously calculated acid may be added just prior to the final titration or, as in lead azide, just prior to the end point. Titrate 50 ml. of the cerate solution with the ferrous sulfate solution following the above procedure. ~.%I.CUIATI~N.
a = ml. of ferrous sulfate to reduce 50 ml. of
the standard cerate solution b = ml. of ferrous sulfate to reduce excess cerate X = normality of ferrous sulfate, computed daily from readings obtained in a w = weight of sodium azide sample Equivalent xeight of sodium azide (XaK3) = 64.997 ( a - b) X X X 6.4997 Sodium azide valuc = W
Table 11.
Comparison; of Sodium A z i d e Values b y Two Methods
Cerate Method Titrationa 99.3 99.0 9 99.6 99.2 3 99.3 99.3 4 99.3 99.7 5 99.4 99.8 6 99.4 99.4 7 99.3 99.7 8 99.4 99.4 9 99.6 99.4 10 99.3 99.1 Arithmetical average, ?i 99.4 99.4 Standard deviation, u 0.11 0.25 High value 99.6 99.8 Low value 99.3 99.0 Range 0.3 0.8 Titration with standard 0.1 S silver nitrate, bulk of silver azide being filtered off prior t o end point and titration continued, using potassium chrom a t e solution as t h e indicator. Trial No 1
EXPERIMENTAL. The cerate titrimetric procedure was carefully checked against one other procedure as indicated in Table 11, with one trained operator working under ideal conditions on one sample of sodium azide. Table I1 presents the data from ten consecutive trials by each of two procedures. P v m f A R I ' A X D COXCLLXIOXS. The method as applied to sodium azide has none of the disadvantages of existing argentometric procedures. The end point is sharp and clear, and no filtration is required. Experimental work shows that results are
April, 1945
ANALYTICAL EDITION
more reproducible by the cerate procedure and are obtainable with less manipulation and in a shorter time. ACKNOWLEDGMENT
The writer gratefully acknowledges the collaboration of members of the Inspection Board Staff. This paper is published with the permission of the Inspection Board of the United Kingdom and Canada.
217 LITERATURE CITED
(1) Moran, R. F., IND.ENG).CHEM.,ANAL.ED.,15, 361 (1943). (2) Smith, G . F., “Cerate Oxidimetry”, Columbus, Ohio, G. Frederick Smith Chemical Co., 1942. (3) Smith, G . F., and Richter, Frederic, “Phenanthroline and Sub.
stituted Phenanthroline Indicators, Their Preparation, Properties, and Applications to Analysis”,Columbus,Ohio, G. Frederick Smith Chemical Co.. 1944. (4) Sommer and Pincas, Ber., 48, 1963 (1915).
Polarographic Determination OF Vitamin in Fruits and Vegetables W. S. GILLAM,
Purdue University, Lafayette, Ind.
A polarographic method for determining ascorbic acid in fruits and vegetables affords a reliable means of determining vitamin C in quantities ranging from 4 to 85 micrograms per ml. of solution with an accuracy of 3.3 to 4.39& depending upon the supporting electrolyte used. Variable results were obtained on some extractse.g., dehydrated beets. The half-wave potential and the diffusion current, respectively, were found to b e independent of, and proportional to, the ascorbic acid concentration in the four supporting electrolytes studied. Diff usion current constants, half-wave potentials, and diffusion coefficients were determined in four supporting electrolytes. Results of vitamin C analyses on fruits and vegetables b y four different methods-polarographic, visual titration, photometric, and Roe’s-were reported and discussed. The polarographic method showed good agreement with the visual titration and photometric methods on certain fruit extracts. By making use of the diffusion current constant polarographic determination of vitamin C in certain m.%ials i s possible without the necessity of calibrating the instrument.
M
C
OST of the methods that have been proposed for the estimation of vitamin C have been based upon the oxidation of ascorbic acid to dehydroascorbic acid by a specific oxidant such as 2,6-dichlorophenolindophenol. Roe (11) has developed a method which is based on an entirely different reaction of ascorbic acid-namely, its reaction with 2,+dinitrophenylhydrazine. The use of a dye as a specific oxidant has certain limitations. The indophenol dye is not ideal because of its relatively high oxidation potential; materials other than ascorbic acid may be osidized. Its reaction with ascorbic acid is not instantaneous and therefore the time of reading the end point and the rate of addition of the dye are factors that may introduce errors. In the absence of the vitamin the dye will fade, and cause a drift in the galvanometer if used in connection with a photoelectric colorinieter. For highly colored extracts the visual titration method cannot be used, and some modification of the method must be resorted to, such as the one suggested by Bessey (2). Potentiometric methods have been published for estimating vitamin C (1, 6) but a continuous drift in potential throughout the whole course of the titration with the dye makes it very difficult to locate the end point with any degree of accuracy. Harris et al. (6) using a platinum-mercury electrode greatly improved the potentiometric method. Their electrode, however, was found to be slightly sluggish a t the end point, and when working with materials that were high in content of reducing substances they used the less desirable bright platinum electrode. Lewis (S), in this laboratory, found the platinum mercury electrode acceptable on extracts from fresh materials but unsatisfactory for dehydrated products. The ’polarographic method, described in this paper, differs
from most methods in that a specific compound is not used as the oxidant. Instead, ascorbic acid is oxidized at the dropping mercury electrode. Essentially, it is based on current voltage curves and since ascorbic acid possesses striking reducing properties this method should be applicable to vitamin C determination. The determination is specific, is sensitive to small concentrations of ascorbic acid, and colored extracts do not necessarily interfere. Furthermore, an inspection of the anodic wave will definitely indicate whether or not interfering materials are Dresent in the solution being analyzed. Kodicek and Wenig I ’ ‘ 1 ( ? ) stated that dehydroascorbic acid was not reducible a t the dropping mercury cathode, but a characteristic wave was obtained when the dropping mercury electrode was polarized &s an anode. They suggested using a 0.066 N phosphate buffer (pH 7.0) in all polarographic analyses for vitamin C. Their method of analysis was I I ,I , suitable for orange juice +0.3 +0.2 t O . 1 0 but they were not sucVOLTS cessful in applying it to (vs. s.c. E.) extracts of plant or Figure 1. Shift of Half-Wave animal tissues. Other Potential with Change of pH constituents apparently 1, 28 micrograms of ascorbic acid per ml. at hindered the electrode PH3.41. 2, 28 micrograms of ascorbic acid per ml. at p H 2.17. 3, residual current, reaction. supporting electrolyte II (phosphate bufferCozzi (3) reported the 1.5% HPOa). Broken line, residual current after aeration use of the polarograph in the analysis of fruit juices and concluded that it gave high results. The present paper presents a polarographic method for determining vitamin C in plant materials vhich eliminates the necessity of frequently standardizing solutions, is not affected by colored extracts, and is sensitive to very low concentrations of the vitamin. Diffusion current constants and the half-wave potentials of ascorbic acid in four different supporting electrolytes are given. The results of analyses run in this laboratory by four different methods-polarographic, visual titration, photoelectric, and Roe’s dinitropnenylhydrazine method (f2)-are also reported.
-0.1
PRELIMINARY INVESTIGATION
In the oxidation of organic compounds hydrogen ions are usualiy involved in the electrode reaction-.g., RH,,==
