comparison to conventional photographic detection. Sample and standard preparation (including ashing), electrode preparation, instrument preparation (setting peaks), running of ten samples and four standards, data reduction, and conversion to concentration in the whole coal required only 15 hr. This represents 84 separate determinations. This 15 hr also includes sample turn-around time and five separate 0.3 nC monitor exposures for each element in each sample and standard. The detection limits for the elements studied are on the order of 1 to 2 ppmw in the whole coal, using the previously mentioned parameters. These are not necessarily the lower limits, as such items as exposure (nC), multiplier gain, and sample dilution may increase or decrease this level. From the intensity values obtained during this investigation, absolute detection limits ranging between 0.03 and 0.1 ppm by weight with whole coal would be a realistic estimate. These lower values were not actually determined. Relative standard deviations are on the order of 6 to 15% for the mass spectrometric data and 2 to 3% for the atomic absorption values. SSMS may be used to determine those elements in coal that may be difficult to determine using AAS. The one other definite advantage to the mass spectrometric method is the lack of a dissolution step in the sample preparation.
one to make this same generalization for all elements since these were the only elements determined by both ashing techniques during this investigation. Kometani e t . al. ( 4 ) have indicated t h a t the presence of sulfates prevents the loss of many elements during dry ashing; since coal contains appreciable amounts of sulfate, this may help to explain the agreement between the two methods of ashing. Vapor pressure data also imply t h a t some of these metals could possibly be present as oxides or silicates. Since ash contents of the coals studied varied from 5 to 25% according to the geographical location, the standards for the SSMS analysis were chosen so as to coincide with the approximate area from which the actual samples were taken. Table I11 is a comparison of the AAS and SSMS data obtained on ten different coal samples from three or four different geographical areas. These values are based upon the metal concentration in the whole coal whereas the analysis was performed on the ash. Again, there is generally very good agreement between the two different methods used. These data also indicate t h a t hydrocarbon interference is not a problem a t these concentration levels. A very volatile element such as mercury could not be determined using this method as it has been postulated ( 5 ) . I t may be pumped away during heating caused by the excitation process, depending upon the operating parameters selected. Previous work (2) indicates the accuracy attainable with SSMS with electrical detection is very acceptable and t h a t considerable time savings may be realized in
ACKNOWLEDGMENT The author wishes to acknowledge the aid of E. N. Pollock and S. West for providing all atomic absorption data.
( 4 ) T. Y . Kornetani, J . L. Bove, B. Nathanson, S. Siebenberg, and M . Magyar, Environ. Sci. Techno/.. 6, 617 ( 1 9 7 2 ) . (5) W. W. Harrison and D. L. Donohue, Twentieth Annual Conference on
Received for review November 24, 1972. Accepted January 29, 1973.
Mass Spectrometry and Aliied Topics, Dallas, Texas, 1972.
Application of a Sulfate-Sensitive Electrode to Natural Waters R a y m o n d Jasinski and I s a a c Trachtenberg Texas lnstruments Incorporated, P. 0. Box 5936, Dallas, Texas 75222
In previous papers (1, 2 ) , we described the chemical and electrochemical basis of a simple and rapid potentiometric titration technique for the determination of sulfate ion in water. The pertinent concepts involved in the analysis were the following: electrodes formed from an iron-doped chalcogenide glass Fez (Ge2sSb12Seso) respond selectively to uncomplexed ferric iron (I, 2 ) ; ferric iron forms stable soluble complexes with sulfate ion ( 3 ) ;barium ion, added incrementally, breaks up the complex by precipitating barium sulfate; the electrode senses the liberated ferric ion; and the end point is signalled by the onset of a constant potential. The limit of detection of the method is approximately 150 mg of sulfate/liter. Although this value is too high to be of interest in the analysis of many fresh and drinking waters, it is more than sufficient for the analysis of many brackish waters, brines, sea waters, and waste waters. R . Jasinski and I . Trachtenberg,Ana/. Chem.. 44, 2373 (1972). C . Baker and I . Trachtenberg, d. Electrochem. Soc.. 118, 571
(1971). L. Sillen and A . Martell, "Stability Constants of Metal Ion Corn-
plexes," The Chemical Society, London, 1964.
Such aqueous solutions are, of course, considerably more complex than the simple sodium sulfate-sulfuric acid solutions used to develop the method, and, as with all methods based on barium sulfate precipitation, some interferences are to be expected from these other major ionic constituents. This note describes the evaluation of such interferences in the context of the subject method and provides some preliminary data on the application of the technique to the analysis of some natural waters high in sulfate content.
EXPERIMENTAL T h e sulfuric acid stock solution used for preparing synthetic brackish waters was standardized against a sodium hydroxide solution which in turn was standardized against potassium acid phthalate. The barium chloride titrant (generally 0.2050M) was standardized with EDTA which, in turn, was standardized against calcium carbonate. The test solutions were prepared for titration: (1) by adding the proper amount of sulfuric acid and the specific amount of the interfering ion under study; ( 2 ) by adjusting the p H to 2.1 f 0.1 to avoid subsequent precipitation of ferric hydroxide; and (3) by adding sufficient iron to produce a concentration of approximately l O - 3 M iron; the exact concentration of ferric iron need not he known. ANALYTICAL CHEMISTRY, VOL. 4 5 , NO. 7, JUNE 1973
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~~
~
Table I. Sulfate Recovery vs. Ca2+ Concentration (pH 2.1,2380 ppm SO4*- added)
Ca2+Concn. M 5 x 10-3 10-2 5 x 10-2
Recovery, % 96.5 95.0 91.5
~~
~
Ca2+ Mg2+ Na2+ so42-
c03'Fe Si02
(2380 ppm S 0 4 2 - , pH 2.1)
Potential change, Solution
+ 1 0 - 3 M Fe + 10- 3M Fe + 3 X 1 0 - 3 M Fe + 2 X 1 0 - 3 M Fe + 10-3MFe + 1 0 - 3 M Fe
Distilled H 2 0 0.05M MgCI2 0.05M MgCI2 0.05M MgCI2 0.05M CaC12 0.01M CaC12
mV
1.8 0.7 0.4 0.8 1 .o 2.1
The natural water samples evaluated in this study were prepared for titration according to steps 2 and 3 above. I t is conceivable that certain waste waters. such as acid mine drainage water, would contain sufficient ferric iron to fulfill condition 3 above, but this would be readily detectable by the initial sensor electrode potential. The occasional need for sample dilution will be discussed below. In routine use, the sensor electrode was activated prior to each working day by etching the surface with 10% KOH, washing with p H 2, 1.44 KCI. and then storing overnight in millimolar ferric nitrate ( p H 2 ) . Alternative activation procedures and the electrochemistry involved are discussed in Reference 4 . The electrodes used in the work described contained ohmic contacts on the back side of the sensor element; identical results were obtained when the electrode was constructed in the membrane configuration ( I ) . A double-junction Orion Ag/AgCl electrode was used as reference. The double junction, filled with potassium nitrate, minimized possible contamination from the reference electrode chlo.ride solution. The potentials were measured. during the titration with barium chloride, on a digital voltmeter (Hewlett-Packard Model 344OA3); impedance matching was provided by a Keithley Model 610C electrometer. The titration end point was signalled by the onset of a steady potential (*0.1 mV); approximately 3 to 5 min were required to complete one titration.
RESULTS AND DISCUSSION Calcium is known to coprecipitate with barium sulfate ( 5 ) . Shown in Table I are the results of applying the titration method to sulfate solutions containing various calcium ion concentrations. The range chosen was t o approximate that found in sea water (6). Clearly calcium ions do constitute a n interference for the titration method. Based on replicate runs, these departures from stoichiometry were reproducible within the accuracy of the method so that a correction factor based on calcium concentration becomes possible. Since routine waste water analysis generally involve simultaneous calcium analysis, this correction factor is often available. Potassium also coprecipitates with BaS04. As in gravimetric barium sulfate analysis, potentiometric analysis in KC1 solutions routinely yields less than stoichiometric re-. covery of sulfate, e.g., 94% recovery was obtained in 10-1M KC1 for 2 X 10-2 SO4. Potassium is present in sea water a t the 0.01M level (6) so that coprecipitation should be of less importance for such samples. For most accurate Jasinski and I . Trachtenberg, J . Eiectrochem. SOC.,in press. (5) W . Hillebrand, G . Lundell, M . Bright, and J. Hoffman, "Applied Inorganic Analysis," John Wiley and Sons Inc.. New York. N . Y . . 1953. ( 6 ) J. Riley and G . Skirrow "Chemical Oceanography." Academic Press, NewYork, N . Y . , 1965.
(4) R.
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~
Table Ill. A Roswell Test Facility Well WaterOSW Analytical Results item Concn. ppm
CI yc03-
Table I I . Potential Change at the End Point vs. Solution Composition
~~
Total dissolved solids Total hardness (as CaC03)
51 7 163 4,650 1,490 7,340 206 0
0.22 15 14,400 1,980
pH 7.5
results, however, potassium, as well as calcium ion, would have to be removed as, for example, via passage through a cation ion exchange resin. Magnesium and calcium also form soluble complex ions with stability constants which approximate those of the ferric iron-sulfate complexes ( 3 ) . A series of titrations were made in a medium of 0.05M MgC12; the magnesium content of sea water is a nominal 0.05M (6). Complete recovery was achieved in all cases, indicating little or no coprecipitation of magnesium sulfate a t this concentration level. However, the break in the titration curve a t the stoichiometric addition of barium ion was not so large as that observed with distilled water and sodium ion solutions. This is shown in Table 11, which presents the potential difference between the end point and a volume of titrant 0.3 ml before the end point. [Although these potential changes are small they are well within the stability range of the electrode and the electronics ( I ) . ] Also shown in Table I1 are data for the titration of calcium chloride solutions of comparable concentrations. A somewhat less sensitive end point is also obtained in titrating the 0.05M CaC12 solution, besides the lower recovery discussed above. With natural sea water and other waters high in sulfate concentration, this effect can be minimized by dilution of the sample prior to titration; salt concentrations at and below 10-2M do not cause problems in determining the end point. Finally, the data in Table I1 also indicate that there is little to be gained from manipulating the ferric ion concentration. The method was then applied to a waste water, a brackish water, and sea water. The waste water (-950 ppm SO42-) was analyzed by the standard sulfate-turbidimetric method as well as the subject method. Both methods agreed within experimental error; the titrimetric end point was similar to that found in distilled water. A sea water sample (taken near Galveston, Texas) was run following a 10:1 dilution. Recovery was 2950 ppm of sulfate, which is well within the sulfate range expected for sea water (e.g., 2800 ppm given in Reference 6). Brackish water from the "Roswell Test Facility" (Office of Saline Water) was titrated without dilution: 1360 ppm was found, compared to the 1490 ppm reported (Table 111). However, this water also contained 0.05M calcium, which can be expected t o yield a lower sulfate recovery (Table I). Correcting this sulfate result by the factor shown in Table I yields a value of 1490 ppm, which compared with that reported in Table 111. The end point, although readily detectable was not so sharp as in distilled water samples. A 5/1 dilution of the sample before titration improved the end point without altering the accuracy of the analysis. This is as predicted from the data in Table I and 11.
CONCLUSIONS Among the more ubiquitous ions to be found with sulfate ion are the alkalis and the alkaline earths. The resulting difficulties thereby introduced into the subject method are essentially the same as in all barium sulfate methods. Sodium ion and chloride ion a t a hundredfold excess cause no problems ( I ) . Potassium ion interferes a t this concentration level via coprecipitation; 94% recovery was obtained for 2 x 10-2M S042- in 10-1M KC1. Magnesium and calcium distort the end point because of simultaneous complexation with sulfate ion; this does not
cause a major problem with the titration technique. Although calcium also coprecipitates, this effect is sufficiently reproducible to allow for correction if the Ca content is known. At high sulfate concentrations the sample can be diluted to minimize interference.
Received for review November 17, 1972. Accepted January 22, 1973. Financial support of this work by the Office of Saline Water, U.S. Department of the Interior, is gratefully acknowledged.
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