Anal. Chem. lg82,5 4 , 2134-2136
antenna picking up vibrations from the surroundings. This pickup can be related to the sympathetic resonance seen when a vibrating tuning fork is brought close to a nonvibrating tuning fork of the same frequency. Reduction of cell height increases the resonance frequency, an effect similar to the relationship of frequency to height of organ pipes (14). The resonance frequency for the shorter cell is apparently outside the dominant room noise frequency spectrum. In summary, low-frequency acoustical noise can be dramatically reduced in liquid PAS by reducing the height of the sample container. This simple expedient can greatly enhance the signal-to-noise ratio observed in PAS. I
I
I
25
50
75
ACKNOWLEDGMENT We greatly appreciate the technical assistance of D. M. Fasano and several helpful comments from N. H. Dovichi and R. A. Keller.
TIME ( m s l
Figure 1. PAS signal generated from a 1-mJ N, laser pulse. The sample of doubly distil!ed water is contained in a 5 cm tall curvette.
LITERATURE CITED
I
25
(1) Patel, C. K. N.; Tam, A. C. Rev. Mod. Phys. 1981, 53, 517-550. (2) Tam, A. C.; Patel, C. K. N. Appl. Opt. 1979, 18, 3348-3358. (3) Fisher, M. R.; Nogar, N. S.;Schuster, S . M. And. Blochem. 1981, 113, 112-117. (4) Lahman, W.; Ludewlg, H. J.; Welling, H. Anal. Chem. 1977, 4 9 , 549-55 1. ( 5 ) Oda, S.; Sawada, T.; Kamada, H. Anal. Chem. 1978, 50,865-867 (1978). (6) Volghtman, E.; Jurgensen, A.; Winefordner, J. D. Anal. Chem. 1981, 53, 1442-1446. (7) Tam, A. C.; Patel, C. K. N. Opt. Lett. 1980, 5,27-29. (8) Fisher, M. R.; Nogar, N. C., unpublished results. (9) Fisher, M. R.; Fasano, D. M.; Nogar, N. S.Appi. Spectrosc. 1982, 36, 125-128. (10) Fasano, D. M.; Nogar, N. S.Chem. Blomed. Envlron. Instrum. 1981, 1 1 , 331-339. (11) Alkemade, C. Th. J.; Snelleman, W.; Boutlller, G. D.; Pollard, B. D.; Wlnefordner, J. D.; Chester, T. L.; Omenetto, N. Spectrochim. Acta, Part 8 1976, 338,383-399. (12) Boutiller, G. D.; Pollard, B. D.; Winefordner, J. D.; Chester, T. L.; Omenetto, N. Spectrochim. Acta, Part 8 338,401-415. (13) Alkemade, C. Th. J.; Snelleman, W.; Boutiller, G. D.; Wlnefordner, J. D. Spectrochim. Acta Part 8 ,1980, 358,261-270. (14) Sears, F. W.; Zemansky, M. W. “University Physics”; Addison-Wesley: Palo Alto, CA, 1964.
I
75
50
TIME (rns)
Same as Figure 1, except that the cell height has been reduced to 2 cm. Vertical scale is the same In both figures. A photodiode was used to monitor the laser pulse in both cases. The pulse energies used to generate these traces varied by less than 5 % . level to which that cell was filled. Relative standard deviation in blank signal was 5% for the short cell, compared to -60% for the standard cell. We rationalize these results in terms of acoustic resonances in the sample cell. Apparently the cuvette cell acts as an Figure 2.
RECEIVED for review April 19,1982. Accepted June 25, 1982. This work was supported by the U.S. Department of Energy under the auspices of Los Alamos National Laboratory.
Anhydrous Sodium Thiosulfate as a Primary Iodometric Standard A. A. Woolf School of Chemistty, University of Bath, Bath, BA2 7AY, United Kingdom
Whereas sodium thiosulfate pentahydrate is not recommended as a primary iodometric standard because it tends to lose water inhomogeneously under ambient conditions, Tomlinson and Ciapetta ( I ) demonstrated that the anhydrous salt, which only rehydrates slowly, is a suitable standard. This anhydrous material has not been widely accepted, probably because of the inconvenient preparation. Popular analytical texts seem to change their allegiance between editions. Thus in Vogel(2) the anhydrous sodium salt of the second edition was replaced by BaS2O3.H20 in the third edition, and the latter had disappeared from the fourth edition but has reappeared in the fourth edition of Skoog and West (3). 0003-2700/82/0354-2134$01.25/0
The barium salt seems a poor alternative since it is too insoluble to prepare 0.1 M solutions, besides requiring large amounts of material in its preparation (4). We have now found that the crystalline pentahydrate, or even concentrated aqueous solutions, can be dehydrated with hot methanol in a matter of minutes to a loose powder of the anhydrous form. This material, which can be handled in the open for hours without appreciable rehydration and is rapidly and highly soluble with only slight heat evolution, can serve as a primary standard. A considerable saving in operating time €or preparing standard solutions of thiosulfate is possible since the preparation of and titrations with standard potassium iodate 0 1982 American Chemical Society
ANALYTICAL CHEMISTRY, VOL. 54, NO. 12, OCTOBER 1982
Table I. Thermal Dehydration Rates in Na,S,03-5H,0 (25-g Specimens in 6.5 cm Diameter Crystallizing Dishes) 90 "C at 104 "C at 120 "C T/ha
a
Llmol T/h Limo1 1.00 1.29 1.00 2.26 1.58 1.57 2.00 3.84 1.96 3.00 4.72 19.00 23.5d 2.24 4.50 4.911 2.45 5.83 4.95 27.40 2.84 22.00 5.00 90.90 331.00 3.30 4.1 3 475.00 T = time. b L = water loss.
T/h 0.55 1.25 1.73 2.68 3.18 4.92
Llmol 1.33 3.83 4.80 4.96 4.97 4.99
2135
Table 11. Solvent Dehydration of an Aqueous Thiosulfate Solution at 20 "C 10 mL of a 3.26 M Solution of Na2S,0,.5H,O (H,O/Na,S,O, Mole Ratio 14.39) Shaken with 100 mL of Solvent precipitate composition I recovery I solvent mol of H,O % 0.98 77 MeOH 4.89 78 CH,OHCH,OMe 5.00 100 EtOH 100 5.00 Me, CH(OH) 4.95 100 Me,CO 4.95 100 MeCN no precipitate (CH,OH), Table 111. Solvent Dehydration of Crystalline Pentahydrate at the Boiling Point residual hydrate/ solvent solvent bp/"C mol H,O removal by
P P
Me,CO Me6H EtOH
/
P
MeCN (CHzOMe), Me,C(OMe), MeNO, CH, OHCH, OMe O(CH,CH,OMe),
i
I
1 0
20
TIME /t-'RS.
Figure 1. Vacuum dehydration of Na2S203.5H20 in tube of 3.5 cm diameter).
torr at 25 "C
or arsenious oxide solutions are bypassed. In addition it should be quicker to prepare fresh standard solutions of thiosulfate rather than frequently restandardize old solutions.
RESULTS AND DISCUSSION Thermal Dehydration of WaaS,03.EiH20(A.R.). Dehydration below 100 "C is rather slow for practical use but is possible in a few hours above 100 "C (Table I). There is some decomposition above 180 "C. Duval's figures (5) are unreliable and must represent conditions far removed from equilibrium. The water loss is parabolic after the liquid melt of Na2S2036H20 (48 "C) skins over and the exit of water is controlled by solid-state diffusion. Loss of greater than 99.9% of water at 120 "C, as recommended by Tomlinson and Ciapetta, requires at least 5 h of heating. (It i@convenient to use P.T.F.E. beakers, which can be flexed slightly after heating. This enables the cake of anhydrous salt to be easily detached from the nonstick surface.) A much slower vacuum dehydration is possible at room temperature. The curve in Figure 1shows no inflections corresponding to different hydrates as claimed by Duval(5). Indeed, the well-known superheating and supercooling phenomena i n the NazSz03-H20system (6),as well as the water gradient produced in the dehydrating crystals, would preclude any such stoichiometric inflections in an inhomogeneous system. Solvent Dehydration bf the Pentahydrate. Aliquots of an aqueous solution of the pentahydrate were shaken with solvent at room temperature. The composition and amount of precipitate are given in Table 11. There is a remarkable difference between methanol and ethanol. Ethanol forms a coarsely crystalline material, in contrast to methanol, which yields a powdered matelrial of larger volume. In Table I11 the effect of refluxing pentahydrate crysta& with different solvents is given. Methanol is unique in that the pentahydrate solid is converted to the anhydrous form without change of state. With other solvents a liquid phase of molten pentahydrate is formed above 48 "C which slowly converts to the solid.
a
56 65 78.5 81.6 83 85 101 125 162
2.29 0
3.85 a 2.52 5
0
none 0
air-drying short pumping at 0.02 torr prolonged pumping at 0.02 torr
0.73
Solid only separates on cooling.
Dehydration is hot a purely thermal effect, since it is unrelated to the boiling poiht of the solvent, except possibly for the demethoxyethyl ether of highest boiling point, (dimethoxypropane dehydrates by a hydrolytic reaction, unlike the other solvents). , Preparation of Anhydrous Sodium Thiosulfate with Methanol. The time required for dehydration was examined by heating Na2S2O3.5Hz0(21 g) under reflux with MeOH (100 mL) for periods ranging from 3 h to 10 min. In all experiments, a 94-95% yield of anhydrous salt was recovered after filtration and rinsing with 10-20 mL of fresh MeOH a t room temperature. Residual solvent was evaporated either in vacuo (0.02 torr) or by holding at 70 "C for half an hour; a mass balance was obtained by evaporating the filtrate. It is advisable to keep the mixture stirred during the dehydration to avoid bumping. Dehydration is unaffected by crystal size. Obviously, there is no need to dry the methanol. In the above dehydrations the final mole ratio MeOH/H20 was 6.13/1. This ratio should not be allowed to drop below 3/1, in order to keep the solubility of the thidsulfate below 1% w/w. This solubility is linearly related to the H20/MeOH ratio (0.262 g/lOO g solution at ratio 0 (methanol); 0.521 a t 1/12; 0.972 at 1/6; 1.51 a t 1/3; 3.71 at l/l). It should be noted that it is still possible to obtain the anhydrous salt when the methanol contains a considerable water content, because of the slow equilibration in the Na2S203-H20-MeOH system, of which the binary section NazS2O3-H2Ohas been examined (6). The composition of the solid equilibrated for 3 days at 25 "C with saturated MeOH-H20 solution of NazS2O3-3.86 H,O at 2/1 MeOH-H20, 1.85 at 3/1, and 0.09 at 6/1-shows that the anhydrous form is only stable above a ratio of 6/1, yet the anhydrous salt can be filtered off after refluxing solutions of final composition as low as 3/1. I t should be noted that the anhydrous salt crystallizes even from aqueous solutions above 70 "C (6). Yields can be marginally improved (ca. 2%) either by cooling the solution below room temperature or by maintaining the MeOH/H20 ratio whilst refluxing. This is
2136
Anal. Chem. 1982, 54, 2136-2139
Table IV. Heats of Solution of Thiosulphates AH"/kcal mol-' by calofrom solubility NBS 500 rimetry data a ( 9 ) value NazSzO,
t 2.03
Na,S20,.5H2O -11.35
2.47 (over range
t 2.0
50-80 "C)
10-25 "C) a Both have positive temperature coefficients of solubility. ~~
E 3
10.44 (over range
~~~
achieved by using a Soxhlet column filled with molecular sieve (3A) below the condenser. Thus NazSz03.5HZ0(32 g) in MeOH (100 mL) after 3 h left an amount of soluble thiosulfate corresponding to a 5.0/1 ratio rather than the 3.911 ratio calculated from reactants. The improvement is rather better than indicated, because the vapor phase is impoverished in water content (7).A mole ratio of 3.9/1 in the liquid phase is in equilibrium with a 9.011 ratio in the vapor being dried. Properties of Anhydrous NazSz03of Analytical Interest. Rate of Solution. The fine powder produced from methanol dissolves rapidly enough for solutions to be made up in situ in standard flasks, in marked contrast to potassium iodate solutions. Heats of Solution. These were determined at 25 "C by injecting gram samples into 200 mL of water and comparing the temperature changes with those caused by injecting potassium chloride samples. Magnitudes of heats of solutions were also calculated from slopes of the logarithm of mole fraction solubility against 1/T (K) curves (Table IV). The maximum heat rise in making up a molar solution of NaZSZO3 would be 0.47 "C in comparison with a 2.53 "C cooling for a molar pentahydrate solution. Rate of Water Uptake. This is the crucial property for analytical use. Previous workers by not defining their parameters have given nonunique figures which could fit a variety of conditions. In a series of experiments, samples were exposed to different relative humidities (R.H.) over sulfuric acid solutions at constant temperature. Samples were weighed to the nearest 0.1 mg immediately after removal from the constant humidity chamber. Air in the balance case was ionized with an 241Ama-source to overcome electrostatic charging effects. Moisture uptake was linear for short times, but decreased with time, especially at high R.H. Uptakes were averaged over 21-24 h periods, assuming a linear uptake. In a given cylindrical tube the uptake was independent of weight-upper layers act as desiccant for lower layers-but increased proportionally to the cross-sectionalarea of the tube. It was also strongly dependent on R.H. above 12% R.H., below which the salt remained anhydrous (Figure 2). A t a given humidity the moisture uptake can be minimized by using the smallest practical cross section. Thus a cylindrical
-
60t /
0 5 M OISTURE JPTAKE/ rng/hr/cm2
I
.o
Flgure 2. Effect of humidity on rate of hydration of Na2S20,.
weighing bottle, 7 cm by 2.2 cm diameter, filled with thiosulfate dehydrated by methanol, would only take up 0.01 % of ita weight of water per hour when exposed to an atmosphere of 50% R.H. at 25 "C, or alternatively it could be left open 10 h before changing weight by f0.1%. This thiosulfate packs by light tamping to 43% of the maximum X-ray density of 2.334 g (8). Standardization against KIO,. Specimens of anhydrous thiosulfate prepared from methanol were checked against iodine solutions (0.1 M in I) prepared from KIO, (AR) using both starch and amperometric end points (for the latter, the current produced between fixed Pt foil and Ni wire electrodes in a constantly stirred solution was observed). All volumetric ware was calibrated by weight. The Na2S20, and KI03 were equivalent within 0.07%, which is within the limits set by the stated KI03 purity (a buoyancy correction was ignored, since it was less than 0.05% and in any case would further reduce the discrepancy). Note Added in Proof. A patent covering methanol dehydration has just come to our attention (10). LITERATURE CITED (1) Tomlinson, H. M.; Ciapetta, T. C. Ind. Eng. Chem., Anal. Ed. 1941, 13, 539. (2) Vogel, A. J. "A Textbook of Quantitatlve Inorganic Analysis", 4th ed.; Longman: London, 1978. (3) Skoog, D. A.; West, D. M. "Fundamentals of Analytical Chemistry", 4th ed.; Saunders: New York, 1982. (4) MacNevln, W. M.; Kriege, 0. H. Anal. Chem. 1953, 25, 767. (5) Duval, C. "ThermogravlmetricAnalysis", 2nd ed.; Elsevier: New York, 1963. (6) Young, S. W.; Burke, W. E. J . Am. Chem. SOC. 1906, 28,315. (7) Cornell. L. W.; Montana, R. E. Ind. Eng. Chem. 1933, 25, 1331. (8) SBndor, E.; Csordis. L. Acta Crystallogr. 1961, 14, 237. (9) Linke, W. F. "Solubilltles of Inorganic and MetaCOrganlc Compounds", 4th ed.; Amerlcan Chemical Society: Washington, DC, 1965; Vol. 11. (10) Dlmsdale, W. H.; Kendall, J. D.; Axford, A. J. British Patent 737295, 1955.
RECEIVED for review April 14, 1982. Accepted July 14, 1982.
Extraction of Various Forms of Sulfur from Coal and Shale for Stable Sulfur Isotope Analysis Linda M. Westgate" and Thomas F. Anderson Department of Geology, Universiv of Illinois, 130 1 West Green Street, Natural History Building, Urbana, Illinois 6 180 1
Studies of sulfur isotopes may provide unique information about the timing and mechanism of sulfur incorporation in organic-rich sediments. The sulfur isotope geochemistry of sediments is complicated by the fact that sulfur can exist in 0003-2700/82/0354-2136$01.25/0
several oxidation states and chemical forms. In coal and shale, for example, sulfur can exist as pyrite (FeSZ),other metallic sulfides (e.g.,ZnS, PbS, FeS), sulfates (e.g., FeS04,Fez(S04)3, or CaSO,), organic sulfur, which is not well defined in coal 0 1982 American Chernlcal Society
