within 30 minutes is impossible as shown in Table 11. However, this chelate is not formed at room temperature or under conditions where EDTA is added before boiling the reaction mixture. Therefore, Cr(II1) does not interfere with the determination. The mutual interference of Fe, Co, and Ni is avoided as follows : in the presence of reducing agents such as L-ascorbate, these metal ions are in a divalent state and only the color of Fe(I1)-PAR chelate can be retained; whereas, if PAR is added, after preliminary oxidation of iron(I1) to iron(II1) with hydrogen peroxide, PAR chelates of Fe(III), Co(III), and Ni(I1) are formed in the presence of dissolved oxygen and only the color of Co(II1)-PAR chelate can be retained. According to these procedures, iron and cobalt can be determined selectively as shown in Table I. Among the metal ions studied, the only interference was Pd(I1). However, this interference was easily avoided by measuring the absorbance against a blank which was prepared by adding EDTA solution first to the sample solution as described in the recommended procedure by which only palladium(I1) is able to form the PAR chelate. These reaction properties of palladium(I1) clearly indicate the possibility of developing an excellent method for Pd(I1) by PAR. This will be published in a later paper. DISCUSSION
PAR and EDTA are not selective agents. However, we have demonstrated that by combining these two agents, the reverse reaction of Equation 1 gives a highly selective color reaction for cobalt(II1) and iron(I1). For cobalt(III), Equation l seems to be irreversible even at the boiling point; however, for iron it becomes irreversible at room temperature and reversible at the boiling point, and for palladium(I1) reversible both at room temperature and boiling point. These
facts suggest that for the case of iron(I1) (at room teniperature) and cobalt(III), this selectivity is due to the kinetic nature of the masking reaction, and for the case of iron(l1) (at boiling point, as shown in Table I) and palladium(II), the selectivity may be attributed to the normal thermodynamic stability of PAR-Pd(I1) and PAR-Fe(I1) chelates against EDTA. In the case of iron, the order of the stability constants is expected from our experiment to be KT~(III)--PAR < Kre(III)--EmA and KF~(II)--FAR > KI.~(II)EDTA. This result indicates that the PAR ligand has a low valence stabilizing property such as that of the o-phenanthroline ligand. This property can be expected from the possibility that PAR forms a chelate-ring with a very similar structure to that of the o-phenanthroline complex. The above analytical methods are accurate, and are applicable in a concentration range of 0.05 ppm through to 0.8 ppm of iron and cobalt. In the case where a higher sensitivity is required than that of the recommended procedure, the extraction of the PAR chelates with tetradecyl-dimethylbenzyl ammonium chloride (TDBA+.Cl-) is recommended (15). Procedure of the extraction is as follows : 50 ml of colored solution is transferred to a 100-ml separatory funnel, to which 1.5 ml of 0.05M TDBA'.Cl- solution was added. The mixture is shaken with 10 ml of chloroform for 10 minutes. The absorbance of the chloroform layer is measured at 522 nm for iron or at 520 nm for cobalt against a blank. By this extraction, an approximately 5-fold increase in sensitivity can be obtained for both iron and cobalt. RECEIVED for review August 9, 1971. Accepted December 21, 1971. (15) T. Yotsuyanagi, R . Yamashita, and K. Aomura, Jap. Anal.
19, 981 (1970).
Loss of Mercury from Water during Storage Robert V. Coyne and James A. Collins US.Air Force Encironmental Health Laboratory, McClellan Air Force Base, Calg. 95652 THEDISCOVERY of significant levels of mercury in fish, water supplies, seed grains, and other edible materials has caused a great deal of emphasis to be placed on improvement of analytical techniques for mercury in terms of both specificity and sensitivity. Numerous modifications in analytical methods for mercury have been published relative t o techniques for urine, blood, water, sediment, geological samples, food products, and animal tissue (1-6). Some recent modifications reported sensitivities in the part per billion range and there is little doubt that such sensitivity is possible from a strict analytical standpoint (7, 8). How(1) E. Berman, At. Absorption Newslett., 6 , 57 (1967). ( 2 ) W. R. Hatch and W. L. Ott, ANAL.CHEM., 40, 2085 (1968). (3) G. W. Kalb, At. Absorption Newslett., 9, 84 (1970). (4) B. B. Mesman and B. S. Smith, ibid., p 81. ( 5 ) R. K. Munns and D. C . Holland, J. Ass. Ofic. Anal. Chem., 54, 202 (1971). (6) V. A. Thorpe, ibid., p 206. (7) R. W. April and D. N. Hume, Science, 170,849 (1970). (8) Y.K. Chau and H. Saitoh, Emir. Sci. Technol.,4,839 (1970).
ever, as early as 1941, reference was made to adsorption of mercuric ions from dilute solutions onto glass surfaces, but no data were presented to support the statement (9). More recently, Shimomura et al. (10, 11) reported their results using isotopic mercury ; the loss of mercury from stored solutions was observed to be proportional to pH, greater losses occurring with higher p H values. In another study using radioactive isotopes, losses up to 82 of added mercury were encountered in unpreserved samples of natural water (8). Greenwood and Clarkson reported loss of mercury in dilute solutions, again using a radioactive mercury tracer; these authors studied a number of container materials under total artificial conditions and, relative to polyethylene, observed mercury adsorption of about 38x in 46 days (12). Other references have been (9) A. E. Ballard and C. D. W. Thornton, Ind. Eng. Chem., 13, 893 (1941). (10) S. Shimomura, Y. Nishihara, and Y. Tanase, Jup. Anal., 17, 1148 (1968). (11) Zbid., 18, 1072 (1969). (12) M. R. Greenwood and T. W. Clarkson, J. Amer. Znd. Hyg. Ass., 31, 251 (1970). ANALYTICAL CHEMISTRY, VOL. 44, NO. 6, MAY 1972
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Table I. Loss of Mercury from Creek Water Stored in Polyethylene Containers Time after Mercury, mg/l. addition of Preserved 0.05 mg/l. mercury Unpreserved (HAC:HCHO) Minutes 0 0.018 0.040 0.010 0 040 15 25 0.010 0.030 40 0.006 0.030 55 0.002 0.030 145 0.002 0.030 Days 3 0.003 0.022 5 0.001 0.021 1 NM” 0.019 9 0.001 0.012 12 0.001 0.010 14 NM 0.009 16 NM 0.007 19 NM 0.005 22 NM 0.002 24 NM 0.001 a NM = None Measurable; generally equal to or less than 0.001 mg/l. I
Table 11. Loss of Mercury from Distilled Water Stored in Polyethylene Containers Time after Mercury, mg/l. addition of Preserved 0.05 mg/l. mercury Unpreserved (HAC:HCHO) Minutes 0 0.050 0.054 1 0.050 0.050 13 0.050 0.049 40 0.040 0.040 41 0.040 0.044 57 0.041 0.035 75 0.040 0.032 108 0.029 0.034 138 0.032 0,032 147 0.032 0.025 200 0.028 0.030 230 0.028 0.028 Days 2 0.005 0.004 7 NM“ NM a NM = None Measurable; generally equal to or less than 0.001 mg/l.
made to the problem of mercury adsorption t o glass surfaces, but quantitative data t o support the observations were not provided (5,6). From a survey of the current literature, it would seem that loss of mercury from dilute solutions might represent a serious problem relative t o analyses performed as part of environmental quality surveys of natural waters, but quantitative information as to the magnitude and rate of this loss is difficult t o obtain. The purpose of the present study is t o investigate the stability of dilute mercury solutions stored in polyethylene containers and t o document the rate of loss under routine conditions of sample collection, storage, and measurement by atomic absorption spectrophotometry. EXPERIMENTAL Apparatus. Perkin-Elmer Model 303 Atomic Absorption Spectrophotometer, with a Perkin-Elmer Model 165 strip chart recorder was used. 1094
ANALYTICAL CHEMISTRY, VOL. 44, NO. 6, MAY 1972
Absorption Cell. The cell was constructed from borosilicate glass tubing, 20-mm 0.d. x 20 cm. The ends were ground perpendicular t o the longitudinal axis, and quartz windows (25-mm diameter x 2 mm thickness) were cemented in place. G a s inlet and outlet ports were attached approximately 2 cm from each end. For further details, refer to Hatch and Ott ( 2 ) . Reagents. STOCKSTANDARD MERCURY SOLUTION. To 50 ml of concentrated nitric acid was added 1.3603 grams of 99.5 % mercuric chloride. The resultant solution was then diluted t o 1 liter with distilled water. One milliliter equals 1 mg of mercury (1000 ppm). INTERMEDIATE STANDARD. One milliliter of stock mercury was pipetted into a 1000-ml volumetric flask and made to volume with distilled water. One milliliter equals 1 pg of mercury (1 ppm). WORKINGSTANDARDS. The 0.01, 0.04, and 0.10 ppm standards were prepared by successively pipetting 2.5, 10, and 25 ml of intermediate standard into 250-ml volumetric flasks and diluting to volume with distilled water. Working solutions were prepared fresh daily. STANNOUS CHLORIDE.A 20% stannous chloride solution was prepared using 6Nhydrochloric acid as the diluent. SAMPLE CONTAINERS. One-gallon polyethylene jugs (manufactured by PII, distributed by Van Waters and Rogers, San Francisco, Calif.) were used for all experiments. SAMPLEPREPARATION. Creek water samples were filtered through Whatman No. 40 filter paper prior to use. “Distilled water” was prepared by glass distillation subsequent to preliminary distillation and deionization. Unless specified otherwise, spiked water samples were prepared by first adding the appropriate preservative to a 1- or 3-liter portion of water and then adding the necessary volume of mercuric chloride standard to achieve the desired final mercury concentration. RESULTS AND DISCUSSION The first phase of the study was designed t o investigate the loss of mercury from a series of creek water samples t o which had been added mercuric chloride solution to a final concentration of 0.5 and 0.05 mg/l.; the preservative (acetic acidformaldehyde. 10: 1) was that used routinely by this and other laboratories for the preservation of water samples for heavy metals. I n most instances, the first analysis was performed about two hours after the mercury solution was added t o the water sample: this figure was used for the calculation of initial recovery. The loss of mercury from these samples was so exceedingly rapid that a second set of samples was prepared and analyzed as a check on the overall procedure of sample preparation, addition of mercury solution, analytical technic, standardization, etc. In view of this high rate of loss, spiked water samples were subsequently analyzed on an hourly basis. The composite results of these experiments are presented in Table I. In a n unpreserved sample, 60% of the added mercury was lost before the first analysis could be made and the loss continued with time; within three t o five days virtually all the added mercury was lost. The use of acetic acidformaldehyde as a preservative resulted in the immediate loss of about 20% of the added mercury; the loss continued over the next several days so that by the third day only about 50% of the added mercury remained. By twenty-four days almost n o mercury was detectable in the preserved sample. Samples of glass-distilled water were prepared and analyzed in the same fashion as were the creek water samples. The results of these analyses are presented in Table 11. The same general pattern was observed with distilled water as with creek water-ix., a 40 % loss of mercury over a period of four hours, the losses being about equal for preserved and for unpreserved samples ; by seven days, no mercury was detectable.
Table 111. Loss of Mercury from Creek Water and Distilled Water Stored in Polyethylene Containers-Effect of Common Preservatives Mercury concentration in mg/l. after addition of 0.05 mg/l. mercury Days 0 3 5 7 9 10 12 Sample description Natural water, plus HNOJ HSOI
0.040
HC1 HC1 HzPOa Distilled water, plus HzSOa
HC1 HC1 NM
0.041 0.043 0.040 0.051 0.051 0.052 0.058
“03
a
0.040
=
0.037 0.034 0.040
0.036 0.021 0.020
0.034
0.015
0,040 0,010
0.049 0.011 0.002 0,001
0.038
0.016 0.013
0.027
None measurable--i.e., generally equal to or less than
0.025
0,010
0.040
0.020
0.040 0.013 0.013 0.006
0.008 N M“
NM NM
19
NM NM
NM
NM
0.0(?1mg,’l.
Additional samples were prepared and analyzed t o investigate other commonly used acid preservatives. Both creek and distilled water samples preserved with either hydrochloric or sulfuric acid t o a pH of 1 showed mercury loss in the 20-40Z range during the first few days; by the ninth or tenth day, these losses were 100% in distilled water and 80% in the creek water sample (Table 111). Phosphoric acid t o a final p H of 1 was also a n ineffective preservative, 7 0 z of the mercury being lost within five days. Of the acids studied, only nitric acid t o a final p H of 1 was effective t o any appreciable extent in preserving the mercury concentration in a creek water or a distilled water sample (Table 111). The nitric acid experiments were extended to examine the effect of the sequence of addition of the constituents t o the sample container. If the water and mercury solution were added t o the container before the preservative, 20% of the mercury was lost immediately, 90% \VIIS lost by the fifth day, and the loss was 100% by the ninth day. If the acid was present in the bottle before the water and mercury were added, a much higher percentage of the mercury was immediately recoverable and the loss rate was greatly reduced. The loss of mercury from solution appeared t o be related to the initial concentration of mercury as well as to t h e presence or absence of a preservative. A preliminary study indicated a loss of 90% of the mercury within three days from a n unpreserved sample of creek water spiked with 0.5 mg/l. of mercury; a sample similarly prepared but preserved with acetic acid-formaldehyde maintained an appreciable amount of mercury through 38 days of testing. Such a high level of mercury is of little practical value, however, since the significant level in environmental monitoring programs is less than about 0.05 to 0.1 mgil. For this reason, our investigations were confined entirely t o the 0.05 mg/l. concentration of mercury. Disappearance of metal ions from solution during a storage period or during a n analytical process is not a new observation. As a check on the validity of their methodology, Gage and Warren found recoveries ranging from 75 to 100% of the mercury added t o rat urines (13). Vickers and Merrick reported recoveries of mercury added to urine in the range of 80-120%, mostly in the 80-90z range (14). Ling performed recovery experiments on mercury added to acetone and re(13) J. C. Gage and J. M. Warren, A/ui. Occup. Hyg., 13, 115 ( 1970). (14) T. J Vickers and S. P. Merrick, Tulunta, 15, 873 (1968).
ported values of 90-100% recovery (25). Mesman et al. presented excellent recovery data relative t o mercury added t o urine, with values ranging from 95-99% (26). Such reports are useful and necessary not only to prove the analytical methodology, but also to indicate that the recovery of added mercury from a solution will not always be as high as is desirable or anticipated. Furthermore, our studies would suggest that the failure t o recover 100% of the added mercury may not necessarily be the fault of the analytical procedure, but rather may be due to actual loss of some fraction of the mercury which was added t o the test solution. Both adsorption t o the container wall and volatilization have been suggested as mechanisms for the loss of mercury from dilute solutions (5, 6, 8-23, 17). Numerous container materials have been studied relative to their tendency to adsorb mercury; based on these reports, polyethylene definitely seems not t o be the container material of choice for maintaining low concentrations of mercury ( 5 , 6 , 8 , 9 , 1 2 ) . Nonetheless, it is common knowledge that polyethylene (or one of its close relatives), because of the low cost, light weight, and durability in transit, is used routinely for collection of water samples for environmental surveys. In view of this knowledge, it becomes vitally important t o determine what rate of loss is observed for mercury solutions stored in polyethylene, and what preservation technique might be of benefit in preventing or at least decreasing that loss. From the standpoint of the environmental testing laboratory, one which analyzes water samples collected at distant sites and mailed to the laboratory for measurement of mercury, the possibility of severe loss of mercury with time presents a very real problem. The reports published to date all seem t o indicate that some loss of mercury from dilute solution can be expected as storage time increases. Unfortunately, none of these reports were based on studies performed under “field conditions”--i.e., on natural water samples stored in typical polyethylene containers and analyzed for mercury by the flameless atomic absorption technique. The present report would seem t o be unique in attempting to document the actual rate of loss of mercury from “typical” natural water samples analyzed by commonly accepted methods. The results indicate that, at the level of 0.05 mg!l. which is a n environmentally (151 C. Ling. ANAL.CHEM., 40,1876 (1968). (16) B. B. Mesman, B. S. Smith, and J. 0. Pierce, J. Amer. I d . Hyg. Ass., 31, 701 (1970). (17) T. Y . Toribara, C. P. Shields, and L. Koval, Talarita, 17, 1025 (1970). ANALYTICAL CHEMISTRY, VOL. 44, NO. 6, MAY 1972
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significant level, the losses of mercury can be severe, possibly leading to serious misinterpretation of the actual conditions within a given water source. Since the error in this case is always on the low side, it would seem appropriate to recommend extreme caution in the interpretation of results from a water sample over about seven days old unless the conditions of collection and preservation are known accurately. T o the degree that these conditions are questionable, any result on a sample over about 24 hours old probably would be suspect. In terms of preservation, the present investigation indicates that of the preservatives studied, nitric acid to a final p H of 1 (about 10 ml per liter of sample) seems to be the only one that is effective to any significant degree. The effectiveness, however, is realized only if the acid is present in the container before the water sample is introduced so that the sample is acidified immediately upon entering the container. Water samples collected in polyethylene containers and subsequently acidified, perhaps even at some time long after collection,
probably would not yield valid results for mercury. The use of acid potassium permanganate has been suggested recently by several authors as an effective preservative for natural waters and standard mercury solutions (17-19). However, since acid potassium permanganate is not a preservation technique which is commonly used or recommended by environmental testing laboratories, it was not included as a portion of the present report.
RECEIVED for review August 2, 1971. Accepted December 27, 1971. The views expressed herein are those of the authors and do not necessarily reflect the views of the US.Air Force or the Department of Defense. (18) S. H. Omang, A ~ a l Chim. . Acta., 53, 415 (1971). (19) T. J. Vickers, Department of Chemistry, Florida State Univer-
sity, Tallahassee, Fla. 32306, personal communication, 1971.
Conductometric Analysis of Oleum SIR: During the course of our investigations with fuming sulfuric acid for the synthesis of organic compounds, it became necessary to develop a rapid, simple method for the determination of the composition of oleum. Many methods for the analysis of oleum have been described in the literature: alkalimetric ( I ) , water ( 2 , 3), and potentiometric titration (4-6), specific gravity ( 7 ) , conductivity measurements (8-12), and others (13-15). Because all of these methods possessed one or more features that did not match our requirements, we undertook t o develop a more appropriate procedure. The method we chose t o investigate was the conductometric titration of oleum with H20. Mixtures of H2S04/SOsform conducting solutions (16), and the composition at minimum conductance corresponds closely t o 100% H2SO4 (17, 18). Orgell (16) has observed (1) F. J. Welcher, "Standard Methods of Chemical Analysis," Vol. I1 (A), 6th ed., Van Nostrand, Princeton, N.J., 1963, p 552. (2) I. Dragusin and D. Trifu, Rec. Chim.,15(3), 164 (1964). (3) J. C . D. Brand, J . Chem. SOC.(London), 1946, 585. (4) L. Giuffre, E. Losio, and A. Castoldi, Chim. Ind. (Milarr),
50(6), 628 (1968). (5) H. B. Van Der Heijde, Chem. Weekbl., 51(46), 823 (1955). (6) W. P. Banks, S. L. Holt, and R. L. Every, U.S. Patent 3,294,652 (C1240-l), Dec. 27, 1966; C.A., 66, PC 8719e. (7) R. Takeno, Ryusau, 17(2), 29 (1964). (8) Ibid., 17(9), 187 (1964). (9) Y . Lacroix, Chim.Anal. (Paris), 46(9), 442 (1964). (10) Ibid., 46(10), 495 (1964). (11) G. M. Fialko and V. M. Bronnikov, Tr. Ural. Nauch.-Issled. Khim. Inst., 1958, 261. (12) A. J. Arvia, J. A. Bolzan, and J. S. W. Carrozza, A I .Soc Cient. Argent., 191(3-4), 91 (1971). (13) R. Takeno, Ryrrsar?,16, 114, 127, 237 (1963). (14) Ibid., 18(10), 211 (1965). (15) Zbid., 18(11), 231 (1965). (16) C. W. Orgell, Diss. Abstr., 24, 96 (1963). (17) R. J. Gillespie and S. Wasif, J. Chem. SOC.(London), 1953, 204. __ .
(18) R. J. Gillespie, J. V. Oubridge, and C. Solomons, ibid., 1957, 1804. 1096
ANALYTICAL CHEMISTRY, VOL. 44, NO. 6, MAY 1972
that the conductance of H2S04/S03 mixture increases with increasing SO3 concentration until a concentration of 12 SOa is reached. At this concentration, the conductance passes through a maximum and decreases with increasing SOs concentration. Similarly, Darling (19) and Record (20) have observed that the conductance of H 2 S 0 4 / H 2 0mixtures increases with increasing HzO concentration until a maximum at about 70% H20 is reached. Thus, although the literature clearly shows that conductometric titration of H 2 S 0 4 / S 0 , with H?O t o a definite end point is theoretically possible, the method has not been described. Gillespie, Oubridge, and Solomons (18) have investigated the electrical conductivity of the mixture resulting from the combination of water and sulfur trioxide with special reference to the composition of minimum conductivity. They show that the point of minimum conductivity occurs slightly on the aqueous side of the composition H2S04; at 9.66 " C , the composition of minimum conductance is 0.0023 i 0.001 mole H?O kgsoln-l,while at 40.00 "C, the value is 0.0015 i 0.001 mole HzO kgSoln-l. These data indicate that over the temperature range 9.66-40.00 "C, the cornposition of minimum conductance changes from 99.996 to 99.997 H2SO4. A schematic drawing of the 3-neck, 1.25-in. diameter X 2-in. high conductance cell used in this work is shown in Figure 1. A 10-cc sample conveniently covers the Teflon (Du Pont) stirring bar and the two platinized 1-cm2 Pt electrodes positioned 5 mm apart. The cell constant was not determined. The three necks of the cell contain a drying tube, a 10-cc buret equipped with a drying tube, and a thermometer. The cell is held in an ice bath and maintained at a temperature of about 10 "C. Conductance readings are obtained with an Industrial Instruments, Inc. conductivity bridge, Model RC 16B2. (19) H. E. Darling, J . Chem. Efig. Data, 9 , 421 (1964). (20) R. G. H. Record, Instrum. Eiig., 4(7), 131 (1967).
