Partial precipitation of these salts has been suggested as an effective procedure for carrying a number of transition elements ( I O ) ; however, this has not to our knowledge been exploited for study of trace amounts of transuranic elements. From Tables I1 and 111, it can be seen that partial precipitation induced by small amounts of sodium hydroxide is an effective and reproducible concentration step in the assay of plutonium in large volumes of sea water. Working recoveries of plutonium tracer, averaging 74% for 50 liters of sea water, are significantly better than the 3-6870 recoveries (averaging about 40%) generally experienced with the iron hydroxide and bismuth phosphate coprecipitation procedures (1-4). Moreover, americium, uranium, and polonium are also effectively carried. It should be emphasized that recoveries in Table III include losses due to ion exchange and plating techniques. Our experience suggests that these losses are-(see Table 1V)- for plutonium 16%, for americium 15%, for uranium 2570,and for polonium 33%. Thus, when the average working recoveries are corrected for such losses, the partial precipitation of the alkaline earth hydroxides and carbonates from sea water appears to carry these four elements with an efficiency of about 90%. Besides the high recoveries and ease with which the precipitate settles, the sodium hydroxide procedure has another distinct advantage over the ferric hydroxide proH. L. Healy, D. Chakravarte, and T . Koyanagi. fnviron. Sci. Techno/..1, 417 (1967).
(10) T . Joyner,
cedure: no iron carrier is added. This permits the anion exchange chromatography to be carried out in a hydrochloric acid media in which many other metals of interest to oceanographers are absorbed onto th’e Dowex-1 resin and can be selectively eluted (6). For example. the iron fraction could be analyzed for 55Fe, an important indicator of new weapons fallout and also of much biological importance. Although the sodium hydroxide procedure is a simple method for attaining high yields, no special attempt was made to achieve perfect equilibrium between the environmental nuclides and the added tracers. It is possible that refractory particulate radionuclides, collected with the precipitate, may escape detection without special treatment of the dried precipitate, for example fusion. That important study is continuing separately.
ACRNOWLEDG.MENT We gratefully acknowledge the courtesy of John Harley of the U.S. Atomic Energy Commission, Health and Safet y Laboratory of New York for the 236Pu,243Am, 232U, and 20sPo tracer solutions and also for a 239Pu standard solution (Pu-239-001-XLIV) against which the 236Pu was compared. Received for review January 7 , 1974. Accepted April 15, 1974. The work was supported by the US.Atomic Energy Commission and the Office of Naval Research of the U S . Navy.
Dichromate Reflux Chemical Oxygen Demand A Proposed Method for Chloride Correction in Highly Saline Wastes Frank J. Baumann Pomeroy. Johnston and Bailey. a Division of Jacobs Engineering Company. 660 South Fair Oaks Avenue. Pasadena. Caiif. 91 105
Any method that attempts to determine chemical oxygen demand (COD) by strong wet oxidative means in saline wastes encounters the problem of chloride interference. Chloride oxidation can be avoided by using mild oxidizing conditions but only a t the expense of inefficient oxidation of organic matter. The standard procedure ( I ) for the determination of COD by the dichromate reflux method utilizes acid concentrations and heating times which will oxidize roughly 85-95’70 of the organic matter present but will also oxidize essentially 100% of the chloride ion ( 2 ) . In the standard procedure, interference by chloride ion, a t moderate concentrations, is largely prevented through the addition of mercuric sulfate to form unionized mercuric chloride ( 3 ) . The complexing method, as currently practiced, using a weight ratio of HgS04:Cl equal to 1 O : l will yield reproducible results a t chloride concentrations up to 5,000 mg/l. Problems due to chloride interference arise in wastes of low to moderate COD with chloride concentrations approaching that of sea water. ( 1 ) American Public Health Association Standard Methods for the Examination of Water and Waste Water 13th ed 1971 pp 495 et
seq
( 2 ) W A Moore R C Kroner and C C Ruchhoft A n a / Chem 953 (1949) (3) R A Dobbsand R T Williams Ana/ Chem 35, 1064 (1963)
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21.
Moore e t a1 (2) noted that quantitative oxidation of chloride depended, among other parameters, on the acid concentration and the amount of dichromate present. Experiments have shown that the degree of oxidation of chloride is not predictable under practical conditions. This unpredictability causes the determination of small amounts of organic matter in the presence of high chloride concentrations (such as in estuarine waters) to be unreliable. It has been stated that a t the present time, COD cannot be accurately measured in samples containing more than 2000 mg/l of C1- ( 4 ) . The chlorine-chloramine cycle, as suggested by Moore ( I ) , may well be a contributing factor to this unpredictability. In light of increasingly more stringent discharge requirements being enacted by State and Federal agencies, many of which base a discharge fee, tax, or surcharge on the relative “strength” (L.e., COD) of the discharge, the need for an accurate and reproducible method of determining the oxygen demand of the waste without positive interference due to oxidation of chloride becomes evident. In experiments conducted by the author with known COD and chloride concentrations, as well as with unknown highly saline wastes, it has been shown that chloride oxidized in the dichromate reflux procedure may be ( 4 ) ’ ASTM Annual Book of Standards, Method D1-252-67, (1972).” p 219
A N A L Y T I C A L C H E M I S T R Y , V O L . 46, N O . 9 . AUGUST 1974
T a b l e I. Apparent and Corrected COD Values on S y n t h e t i c COD-Chloride Mixtures Using Various H g S O I :C1- R a t i o s [COD = 500 mg/l. ]
C 1 - Concn
None 1,000
HgSO4:Cl-
None None
1,000
10: 1
5,000 5,000 10,000 10,000 10,000 20,000 20,000 20,000
None
Apparent COD, mg 1.
C1- correction
501 722 501 1,467 486 2,749 650 632 2,450 646 622
10: 1
None 10: 1 20: 1
None 10: 1 20: 1
None 215
None 934
None 2,225 150 108 1,941 152 130
Corrected COD, mg/l.
501 507 501 533 486 524 500 524 509 494 49 2
500 ML ERLENMEYER FLASK
DISTILLING TUBE W I T H SIDE ARM
- 5 0 0 ML FLAT BOTTOMED BOILING FLASK
Figure 1. Chlorine-collection apparatus for highly saline wastes COD
quantitated and accounted for. Chlorine liberated in the oxidation of chloride is collected in acid potassium iodide solution and back-titrated us. standard sodium thiosulfate solution t o arrive a t a chloride correction.
Table 11. Apparent and Corrected COD Values in Sea Water Chemical oxygen demand, mg/l.
EXPERIMENTAL A p p a r a t u s . The all-glass apparatus, readily assembled from standard laboratory glassware, is shown in Figure 1. A controllable source of nitrogen gas is required. Procedure. To the sample and dichromate in a flat-bottom boiling flask, add mercuric sulfate. The exact amount is not critical though the mercuric ion should constitute a n excess over chloride ion. Slowly, and with cooling. add the required volume of sulfuric acid containing silver sulfate. Connect to the apparatus as in Figure 1. Immerse the tip of the delivery tube in ea. 200 ml of aqueous K I solution (approximately 2 to 3 grams of KI) acidified with 10 ml 1:2 acetic acid. Begin heating and reflux the mixture as outlined in “Standard Methods”. While refluxing. blow a gentle stream (1 to 2 hubbles per second) of nitrogen through the apparatus. Upon cessation of heating, increase the flow of nitrogen to sweep out the apparatus and prevent back-siphoning of the KI solution. Continue the flow of nitrogen until the mixture has cooled. Disconnect and rinse the delivery tube. Titrate the liberated iodine against standard thiosulfate solution, using starch indicator. Titrate reflux mixture as in normal COD determination. Calculation.
Apparent COD =
(a
- b) c
8000
ml sample
Corrected COD = Apparent COD
chloride correction in
-
8000 ( d ) ( e ) ml sample
Where Apparent COD = apparent chemical oxygen demand from dichromate, mg,’l.; a = Fe(NH4)2(S04)2 used for blank, ml; b = F e ( N H 4 ) 2 ( S 0 4 ) 2used for sample, ml; c = Normality of ferrous ammonium sulfate; d = Thiosulfate used for titrating the iodine, ml; e = Normality of the thiosulfate. Corrected COD = Apparent COD - Chloride correction, mg/l. (K2Cr207 = 0.25N).
RESULTS AND DISCUSSION Table I shows some typical results obtained with known COD and chloride mixtures using the recovery method described. Known COD’S were obtained from standardized potassium acid phthalate solutions; sodium chloride was used to obtain chloride concentrations. No significant interference, due to chloride, is encountered a t chloride concentrations of 5000 mg/l. or less in the presence of excess mercury. Conversely. a t chloride concentrations of 10,000 mg/l. or more, even a 20:l ratio of HgS04:Cl (twice the nor-
Method
HgSOd: C1 ratio
Std.
10: 1
Std. Std.
10: 1 20: 1 20: 1 10: 1 10: 1 20: 1 20: 1
Std. Std. Std. Std.
+ recovery + recovery + recovery Std. + recovery
Apparent
Corrected
149 144 163 176 152 147 168 165
58 48 66 54
Table 111. Apparent and Corrected COD Values on T u n a - C a n n i n g Wastes [HgSO:: C1- R a t i o 1O:l) Chemical oxygen demand, mg, I. Run No.
1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24
Chloride, m g i l
12,900 12,500 17,700 15,700 18,500
Unknownu Unknown,’ 18,300 18,700 14,500 16,800 18,600 16,000 18,300 14,750 18,200 15,660 17,300 18,700 18,750 11,280 18,400 18,700 14,700
Apparent
Chloride correction
567 1,710 1,220 541 460 3,000 612 474 480 1,115 554 69 1 762 830 474 672 812 895 250 445 320 142 147 2,896
199 814 698 141 98 778 430 218 307 85 180 131 176 209 102 195 95 45 186 365 61 89 73 1,520
Corrected
368 896 522 400 362 2,222 182 256 173 1,030 374 560 586 621 372 477 717 850 64 80’ 259 53 74 1,376
Sample size, ml.
25 25 23 23 25 25 25 25 25 25 25 25 25 25 25 25 25 25 50 50 50 50 50 10
a Chloride concentration unknown. Sample assumed to be sea water, HgS04 added accordingly.
ANALYTICAL C H E M I S T R Y , VOL. 46, NO. 9, A U G U S T 1974
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mal ratio and a 4.8:l theoretical Hg:Cl ratio) did not prevent the oxidation of some chloride, thereby yielding erroneously high COD results. These results, however, can be corrected by collecting, and accounting for, the liberated chlorine. Table I data would indicate the feasibility of determining dichromate reflux COD in salt water without chloride sequestration by mercuric sulfate. While this possibility is intriguing, the care and extra trouble required in the manipulations would render this variation unsuitable for routine COD determination. Table I1 illustrates the serious error to which the standard method is prone. COD was determined in samples of sea water obtained from Marineland of the Pacific, Palos Verdes Peninsula, Calif. Having established the repeatability of the recovery method on synthetic COD and chloride mixtures as well as on sea water, we used the method to determine COD on waste water from a tuna-canning operation. Table I11 shows some apparent and corrected COD values covering chloride concentrations between approxi-
mately 12,000 and 19,000 mg/l. Table III also shows that a 10:1 ratio of HgS04:C1 will not entirely prevent some chloride from being oxidized. The fraction of chloride oxidized does not appear to be dependent on chloride concentration, sample us. dichromate and acid volumes, nor the amount of mercuric salt added; rather, it appears to vary with the substrate. True values for oxygen consumed from dichromate can be determined in highly saline wastes and sea water by the use of the chlorine-recovery method outlined herein without the need for prior chloride determination or dependence upon a large excess of mercury which, as demonstrated, will not achieve complete sequestration of chloride.
ACKNOWLEDGMENT For inspiration and guidance in the conduct of this study, the author is indebted t o Richard D. Pomeroy. Received for review January 7, 1974. Accepted March 25, 1974.
Use of a Potentiometric Sulfur Dioxide Electrode for Lamp Sulfur Determinations John A. Krueger Orion Research lncorporated, I 1 Blackstone Street, Cambridge, Mass. 02139
Sulfur in petroleum is commonly determined by the Lamp Combustion Method ( I ) , in which the sample is burned under controlled conditions and the gaseous sulfur combustion products are collected and analyzed. In the past, the instability of sulfur dioxide solutions and the unavailability of interference-free methods for sulfur dioxide measurement have made it more convenient to use a hydrogen peroxide absorbing solution, and to measure the sulfate formed by oxidation of sulfur oxides by acidimetric titration or gravimetric determination of precipitated barium sulfate. The recent introduction of a gas-sensing electrode which is selective for sulfur dioxide makes possible the measurement of sulfur combustion products directly as sulfur dioxide in the absorbent. The theory of gas sensing electrodes has been discussed in another publication (2). Briefly, the sulfur dioxide electrode is a gas detecting device which senses the level of dissolved sulfur dioxide in aqueous solutions. A gas permeable membrane separates the sample solution from an internal filling solution in contact with a pH-sensing electrode. As the partial pressure on each side of the membrane equilibrates, the pH of the filling solution varies. The internal filling solution contains a high level of bisulfite, so that the variation in pH is reversible and reproducible. The electrode has a slope of 58 mV per decade over a concentration range of 10-6 to 2 x 10-2M. (1)
"Sulfur in Petroleum Products by Lamp Combustion,'' in "Standard Methods of Chemical Analysis," F. J. Welcher. Ed , (Abstracted from former ASTM Methods D90-55T) p p 2031-33, Vol. 2. part B.
1963 ( 2 ) J W . Ross, J r . , J H Riseman, and J. A . Krueger. "Potentiometric Gas Sensing Electrodes," J Pure Appi. Chem., 36, 473-387 (1973) (paper presented at the IUPAC Conference in Wales. 1973).
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This paper describes a method of collection and analysis of oxidized sulfur products which is considerably faster than earlier methods. The sulfur is measured directly as sulfur dioxide in the absorbing solution-particulate and ionic by-products of the lamp combustion do not interfere.
EXPERIMENTAL Apparatus. A VWR Scientific Lamp Sulfur Apparatus, Cat. No. 51934-000, was used for sample combustion, with a single absorber or, in later experiments. two absorbers in series. Air was drawn through the system a t about 300 ml/minute using a vacuum pump; the air flow was controlled by a 5-turn stainless steel valve. Sulfur dioxide measurements were made using a n Orion Model 95-64 sulfur dioxide electrode and Orion Model 801 digital pH/mV meter, reading to 0.1 mV. A liquid membrane chloride electrode, Orion Model 92-17, p H electrode, Orion Model 90-01, and double junction reference electrode, Orion Model 91-02, with 10% K N 0 3 in the outer chamber, were also used. Reagents. Sulfur dioxide standards were made by appropriate dilution of Orion sulfur dioxide molarity standard, Cat. No. 9564-06. Osmolality and p H of samples were adjusted using Orion "SO2 buffer," made by adding 190 grams of anhydrous N a ~ S 0 4to about 800 ml of water in a 1-liter volumetric flask, mixing thoroughly, then adding 53 ml of concentrated (96-9770) and diluting to the mark with distilled water. The p H of the buffer was about p H 1.2 after 1 + 10 dilution. All other chemicals were Fisher Certified Reagents. The sodium hydroxide used was tested and found to be free of sulfur dioxide. Procedure. Electrode Operation and Storage. Correct operation of the electrode was confirmed by preparing a calibration curve (3).Between measurements, the electrode was kept in the recomniended storage solution to reduce changes in osmolality inside the electrode. The storage solution was prepared by adding about (3) Sulfur Dioxide Instruction Manual, Orion Research Incorporated, 1973
A N A L Y T I C A L C H E M I S T R Y , VOL. 46, N O . 9, A U G U S T 1974
