“
..
.
micrograph. Because it ir. difficult to distinguish the morphological makeup of the Ptblack in resin-impregnated thin electode cross-sections except for the loose Ptblack t h a t adheres to the replica, several were etched with aqua regia
.I..”-
of the replica, the platinum (Pt), Teflon (T), and resin (R) areas are appropriately labeled on the micrographs. This anplysis clearly shows t h a t the Pt-black is dispersed as discrete flocculates with Teflon stringers bridging across the various Pt floccu-
... .... V..”
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47,” “Fuel Cell Systems,” R. F. Gould,
ed., pp. 10615, ACS, Washington, D.C ‘ne= (4) LE search ‘Laboratbries, East Hartford, Conn., private correspondence, 1966.
RECEIVED for review February 18, 1966. Accepted March 28, 1966.
Photochemical Determination of Oxalate LtR WILLIAM M. RlGGS and CLARK E. BRICE---
Department of Chemistry, University of Kana;as, Lawrence, Kan.
b The photochemical decomposition o f ferric oxalate complexes in aqueous acidic solution has been utilized as the basis of new analytical techniques for the determination o f oxalate. Two indirect methods are described; one which i s applicable to traces of oxalate (detection limit o f 5 pg. o f Na&04), and another which i s useful in the milligram range. A number o f organic acids have been investigated for possible interference with each o f the methods. Better than 3y0 accuracy i s obtained for traces of oxalate in the absence o f interferences. Accuracy and precision are bath 2 parts per thousand or better with larger amounts o f oxalate.
T
HE PHOTOCHEMICAL ACTION of light has received only limited attention from analytical chemists. Bricker and Schonberg (8) suggested the technique of “photonometric” titration using carefully calibrated light sources. Using this technique, the amount of material titrated is proportional to the time of exposure to the light. Rricker and Schonherg determined vanadium (V) and chromium(1V) by this technique and Kuwana (6) later applied it to an oxygen determination. Beattie, Bricker and Garvin (1) determined traces of organic materials in water by photosensitized oxidation followed by mass spectrometric measurement of the COS produced. G. G. Rao and co-
workers (10, l d , 13) describe indirect methods for determination of uranium (VI) and iron(II1) using photoreduction of these ions by various organic materials followed by titration of the reduced species with sodium vanadate. b o and Aravamudan ( 1 1 ) also report the determination of oxalic acid by photosensitized oxidation with iron(II1). Part of the work reported here is a n extension of this idea. The photoactivity of iron(II1) oxalate in solution was first observed by Dobereiner (4) who detected iron(I1) in a solution of potassium trisoxalatoferrate (111) which had been exposed to sunlight. I n subsequent years the reaction has been extensively studied. PhotoVOL 38, NO. 7, JUNE 1966
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Table 1.
Method 1. Oxalate Determination
NaGOa, mg. Rel. Added Found Std. dev. error, % 0.10 13.35 0.05 13.34 0.10 26.65 0.07 26.68 0.13 0.03 40.01 40.02 0.13 53.43 0.05 53.36
Table II.
Extent of Reaction with Time of Irradiation
Irradition time (min.) 5.0 7.0 9.0 10.0 15.0 20.0
Amount reacted, 7c 87.3 95.8 98.0 99.3 100.2 99.9
(26.68 mg. sodium oxalate present)
chemical action has been observed with light of wavelengths from 579 to 254 mp (6). The quantum yield is low at long wavelengths and has a value if 1.00 or more [based upon iron(I1) xoduction] only a t wavelengths shorter than 436 nip, Recently Parker and Hatchard (6, 7') pointed out the uselulness of the reaction as a chemical actinometer. They also suggested a reaction mechanism ( 8 ) and later, after a more detailed study using flash photolysis (9),they modified their suggestions somewhat. Some details of the mechanism are still in doubt. Iron(II1) forms mono-, di-, and trioxalato compleyes in solution, all of which are photochemically active. The overall reaction may be represented by Equation 1 with the monoxalate species as the example. 2Fe(C204)+ 5 2Fe+*
+ 2C02 + Cz04-2 (1)
The analytical methods described in this report are based on this reaction. The first of these (Method I) depends upon titrating the photochemically produced iron(I1) with dichromate. It is applicable to quantities of sodium oxalate in the 10- to 100-mg. range. Method I1 rests upon the spectrophotometric determination of the iron(I1) by formation of the 1,lO-phenanthroline complex and measurement of the intensity of the red color formed. This method works well with samples of 5 to 150 pg. of xa2C204. EXPERIMENTAL
Apparatus. T h e light source used throughout was a 100-watt medium pressure mercury vapor lamp from Ultra-Violet Products Inc. T h e radiation from this lamp mas filtered 898
ANALYTICAL CHEMISTRY
through a soft glass plate which effectively eliminated wavelengths shorter than 305 mp. Spectrophotometric measurements were made with a Bausch and Lomb Spectronic 20 except for an absorption spectrum of the above-mentioned glass plate which was obtained on a Bausch and Lomb Spectronic 505. For Method I the sample vessel was a 125-ml. Erlenmeyer flask equipped with a side arm a t the bottom so that the solution could be de-aerated with nitrogen. After irradiation, titrations were performed directly in this vessel. Sample containers for Method I1 were 50-ml. beakers also fitted with side arms near the bottom for de-aeration with nitrogen. Solutions were stirred with a Teflon covered magnetic stirring bar. Reagents. Standard dichromate solutions were prepared from Fisher Certified Reagent potassium dichromate, dried overnight a t 110" C. and stored in a desiccator over calcium chloride. Standard oxalate solutions were prepared from Mallinckrodt analytical reagent grade sodium oxalate, dried and stored in the same manner. Standard iron(II1) sulfate solutions were prepared from reagent grade material, and were standardized by the stannous chloride reduction-dichromate titration method ( 3 ) . Sodium diphenylaminesulfonate indicator solution was prepared by dissolving 0.3 gram of barium diphenylaminesulfonate in 100 ml. of water, adding excess 0.1JI sodium sulfate and decanting the solution after the precipitate settled. Orthophenanthroline solution (0.1%) was prepared by dissolving the solid in hot water. Standard iron(I1) solutions were prepared by weighing G. F. Smith primary standard ferrous ethylenediammonium sulfate tetrahydrate. The compounds used in the interference studies were of various grades depending upon availability. This caused no trouble except in the case of formic acid which appeared to have a small amount of oxalate impurity. All other reagents used were analytical reagent grade. Solutions were de-aerated when necessary with tank nitrogen. Rocedures. A . .Analyses of sodium oxalate by Method I n e r e carried out as follows: T h e sodium oxalate was prepared as a 0.01005f solution, n-hich was also 1 9 HzS04. Amounts varying from 10.00 to 40.00 ml. were pipetted into the reaction flask. d 3- to 5-fold excess of ferric iron was added from a buret as 1-2 ml. of 0.lJ.I solution of iron(II1) in 1.V sulfuric acid. The solution was swept for 5 minutes with nitrogen and then irradiated for 20 minutes with constant stirring and continued nitrogen sweeping. The mercury vapor lamp was approximately 6 inches from the surface of the solution. After irradiation, 2 ml. of 85% phosphoric acid and 5 drops of sodium diphenylaminesulfonate indicator were added to the flask, and the iron(I1) present was titrated with 0.0100N dichromate. Compounds studied as possible interferences were added to the mixture before irradiation
in the form of 0.1 gram/ml. solutions in 1.Y sulfuric acid. B. The general procedure for Method I1 was as follows: A standard curve was obtained by dissolving 0.0684 gram of ferrous ethylenediammonium sulfate tetrahydrate in water and diluting to 1 liter with 1 S H2S04. This solution contained 0.010 mg./ml. of iron(I1). Further dilutions were made by pipetting amounts varying from 1.00 to 50.00 ml. into 100-ml. volumetric flasks. Ten milliliters of 0.1% orthophenanthroline was added, the p H was adjusted to about 5 with sodium acetate, and the solution was diluted to volume. The absorbance was read a t 508 mp. Samples for analysis were prepared and analyzed by pipetting amounts varying from 0.10 ml. to 10.00 ml. of 1.00 x 10-4M sodium oxalate solution into the irradiation vessel, adding 5.00 ml. of 1 X l O - 3 J f iron(II1) solution and enough 1 S H2S04to bring the volume to 20.00 ml. The solutions were stirred and suept with nitrogen for 5 minutes in the dark and then for 20 minutes under irradiation with the lamp positioned about 6 inches from the surface of the solution. At the end of this time a 5.00-ml. aliquot of the irradiated solution was pipetted into a 10.00-ml. volumetric flask; 2 ml. of orthophenanthroline were added, and the pH was adjusted to about 3.5 with sodium formate solution. The solution was then diluted to volume and the absorbance determined 21s. the appropriate blank. For interference studies, appropriate solutions were added before the irradiation. RESULTS A N D DISCUSSION
Method I. The results of the analysis of several different solutions of sodium oxalate are presented in Table I. The precision and accuracy are good and are of the magnitude of ordinary volumetric errors. This supports the hypothesis t h a t the photochemical reaction is quantitative and stoichiometric under these conditions. Experiments carried out without nitrogen de-aeration gave low and erratic results due to air oxidation of iron(I1). The solution acidity is not critical as long as it is high enough to prevent precipitation of ferrous oxalate or ferric hydroxide. The recommended concentration of acid is 0.1 to 1.0-Y. Sulfuric rather than hydrochloric acid should be used since HC1 prevents the reaction from going to completion in a reasonable period of time. This is due to the chloro- complexes of iron(II1) which form in HC1 solution and absorb a large proportion of the incident radiation. Results are independent of the iron(II1) concentration provided a t least 2 moles of iron (111) are present for each mole of oxalate. The time of irradiation required is somewhat dependent upon the amount of oxalate in the sample as well as upon the lamp intensity and geom-
etry. With the lamp and geometry used here, 15 minutes was found to be a sufficiently long irradiation time over the oxalate concentration range studied. For a sample containing 26.68 mg. of sodium oxalate, the percentage found as a function of irradiation time is shown in Table 11. Attempts were made to analyze sodium oxalate in the presence of a variety of other compounds. Table 111gives a list of these compounds and a tabulation of the concentration levels at which interference was found. The slight interference found with formate is believed to be due to traces of oxalate impurity in the formic acid used. All of the hydroxy acid anions studied form photochemically actix e complexes and interfere by producing iron(I1) in a competing reaction. Method 11. This method is essentially a n extension of Method I, with spectrophotometric determination of iron(I1) replacing the titrimetric method. -4s with Method I, t h e amount of iron(I1) produced is found to be directly proportional t o t h e amount of oxalate initially present. A plot of absorbance of t h e ferrous 0-phenanthroline solution us. amount of sodium oxalate in the sample is linear. It is in excellent, agreement with a similar plot prepared using iron(I1) solutions prepared from primary standard ferrous ethylenediammonium sulfate with the assumption t h a t two iron(I1) ions \\ill be produced per oxalate ion. Thus, entirely satisfactory calibration curves may be prepared using this convenient standard. The same restrictions regarding the absence of atmospheric oxygen and the proper solution acidity apply in this case as in Method I. Twenty minutes was found to be sufficient irradiation time for complete reaction in all cases. Fif teen-minute irradiations sometimes gave lo\\ results. Compounds which interfere M ith AMhod I also interfere in this case for the same reason. Compounds which do not interfere at the levels investigated with Xethod I may interfere at the higher concmtrations which might be encountered when attempting to analyze traces of oxalate in other materials This possibility was investigated in the case of formate, acetate, propionate, and malonate. I n the presence of formate, high, consistent results were obtained which were, as with larger amounts of oxalate, probably due to o\alnte impurities in the formate. Assuming that the high results obtained in the presence of formate are due to oxalate impurity, the per cent of oxalate impurity calculated to be in the formate Laried between 0.06 and 0.13%. Results were less than 3‘3 off in the
Table 111.
Interference Studies in the Determination of Oxalate
by
Method
I
NazCnOl mg. found Oxalate: interference ratio
Na2C204, mg. added 1:l 1:lO 1:100 26.67 Formate 26 68 26 76 27 69 26.67 Acetat e 26 62 26 64 26.67 Propionate 26 62 26 68 26.67 Malonate 26 68 26 74 26.67 Succinate 26 66 26 66 26 66 26 67 26,67 a Maleate 26 58 b 26.67 Acrylate 38 41 26.67 Citrate a 26,67 ... Tartrate 26.67 31.45 ... Lactate 26.67 32.50 ... Mannitol The amount present of the compound being studied as a possible interference is expr.essed as the ratio: moles oxalate/moles interference. a S o end point color change occurs in the dichromate titration. b Acrylate polymerizes under these conditions upon exposure to light.
presence of a 250-fold excess of acetate, a 125-fold excess of propionate, and a 5000-fold excess of malonate. Accuracy of & l o % could be obtained in the presence of a 1250-fold excess of acetate, a 2500-fold excess of propionate, and a 10,000-fold excess of malonate. I n these experiments the interference was in the direction of production of less iron(I1) than that corresponding to the amount of oxalate added. However, irradiation of solutions with interfering ion but no oxalate present produces small yields of iron(II), the amount of which increases with time over a wide range. It is thus apparent that at least two types of interference are operative; one which inhibits the iron (111) oxalate reaction and another, less important one which produces iron(I1) in a side reaction. The exact nature of these processes is far from clear and deserves further study. Conclusions. Using techniques described in this paper, it is now possible to determine rapidly, milligram t o microgram quantities of oxalate with good accuracy and precision. These indirect techniques have certain definite advantages over the direct titrimetric methods for oxalate which are presently in common use. For example, the standard dichromate solutions used in the back-titration of iron (11) are stable and can be prepared directly by weighing. This is in contrast to permanganate which is not stable indefinitely in solution and which should be standardized before use. Ceric nitrate and perchlorate solutions are also unstable on standing, although sulfate solutions are stable. The iron(I1)dichromate titration can be carried out at room temperature without catalysts, and the indicator color change is sharp. On the other hand, permanganate and ceric titrations of oxalate must be carried out at elevated temperatures or with a catalyst present, or both. The
indicators used with ceric titrations must be added just before the end point is reached, or if the yellow color of the ceric ion is used as the end point indication, a blank must be used. Furthermore, most organic substances interfere with permanganate and ceric titrations because of the strong oxidizing nature of the titrants. However, the presence of many of these organic compounds can be tolerated in large excess with the photochemical method. The spectrophotometric technique described makes it possible to determine oxalate in microgram quantities with good accuracy. N o other method is presently available which is applicable to such small amounts of oxalate. This trace determination may be successfully carried out in the presence of large excesses of some compounds. LITERATURE CITED
(1) Beattie, J., Bricker, C. E., Garvin, D., ASAL. CHEM.33, 1890 (1961).
(2) Bricker, C. E., Schonberg, S. S., Ibid., 30, 922 (1958). (3)1‘Day, R. A., Underyood, A. L., Quantitative Analysis,” pp. 116, 121, Prentice-Hall, Inc., Englewood Cliffs, N. J.. 1958. (4) Dobereiner, J., Schweigg. Journal 62, 19 (1831). (5) Kuwana, T., rZx.1~.CHEM.35, 1398 (1963). (6) Hatchard, C. G., Parker, C. A., Proc. Roy. SOC. (London) A235, 518 (1R;i61.
( i ’ P a r k e r , C. A., Ibid., AZ20, 104 (1953). (8) Parker, C. A., Trans. Faraday SOC. 50, 1213 (1954). (9) Parker, C. A., Hatchard, C. G., J . Phvs. Chem. 63. 22 11959). (10) kao, G. G.,’AraVamiidan, G., Anal. Chim. Acta 13, 328 (1955). (11) Ihid n 41.6 (12) E R., Ibid., 15, 97 (1956). (13) Rao, G. G., Rao, Sr. P., Venkatamma, N. C., 2. Anal. Chem. 150, 178 (1956).
RECEIVED for review February 2, 1966. Accepted April 7 , 1966.
VOL. 38, NO. 7, J U N E 1966
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