Anal. Chem. 1987, 59, 995-999
-
27688.
viding cefixime degradation standards and spectroscopic data, Carol McCoy for acquiring the DSC data, John James for providing the structure of degradation product 6, and Ving Lee for discussing degradation processes in P-lactams. Registry No. Cefixime, 79350-37-1; ampicillin, 69-53-4; penicillin V, 87-08-1.
PENICILLIN \
12=2MT 15=290.C
t
~
I
247
275
995
:99
-7
333
j?‘
.
376
Figure 9. TSPlMS spectrum of penicillin V, potassium sa#,with t2 = 230 O C and t5 = 290 O C .
regime where TSP data can be obtained tend to produce degradation ions. Under these conditions, the TSP vaporizer probe behaves as a heated flow reactor and the mass spectrometer serves to detect and analyze the product ions in situ. Useful applications of this technique besides structural elucidation of degradation products generated during accelerated stability studies include structural elucidation of high molecular weight compounds by chemical degradation, synthesis of compounds in situ a t elevated temperatures, and degradation rate studies.
LITERATURE CITED Blakley, C. R.; Vestal. M. L. Anal. Chem. 1983, 55, 750. Vestal, M. L. Mass Spectrom. Rev. 1983, 2 , 447. Vestal, M. L. Springer Ser. Chem. Phys. 1983, 25, 246. Bursey, M. M.; Parker, C. E.; Smith, R. W.; Gaskell, S. Anal. Chem. 1985, 57, 2597. Alexander, A. J.; Kebarle, P. Anal. Chem. 1986, 58, 471. Cephalosporins and Penicillins; Flynn, E. H., Ed.; Academlc: New York, 1972; Chapters 3 and 5. Indelicato, J. M.; Dorman, D. E.; Engel, G. L. J. Pharm. Pharmacol. 1981,33, 119. Lederle Laboratories, Amerlcan Cyanamid NDA 19-61 1, June 27, 1986 to FDA. James, J. C.; Ocampo, A. P.; Siegel, M. M.; Morton, G. 0. fharm. Res. 1986, 3 , 14s. Slegel, M. M.; Bauman, N.; Carter, G. T. Anal. Chim. Acta 1986, 186, 163. Vestal, M. L.; Fergusson, J. Anal. Chem. 1985, 57, 2373. Garteiz, D. A.; Vestal, M. L. LC Mag. 1985, 3 , 334. Akman, J. L.; Richheimer, S. L. Tetrahedron Lett. 1971, 4709. Meyers, A. I.; Durandetta, J. L. J. Org. Chem. 1975, 4 0 , 2021. Borders, D. B.; Carter, G. T.; Hargreaves, R. T.; Siegel, M. M. Mass Spectrom. Rev. 1985, 4 , 295. Borders, D. B.; Hargreaves, R. T. I n Blochemical Applications of Mass Spectrometry, First Supplementary Volume; Waller, G. R., Dermer, 0. C., Eds.; Wiley: New York, 1980; pp 567-610.
ACKNOWLEDGMENT The authors greatly appreciate the technical assistance of the chemists from the Fujisawa Pharmaceutical Co. for pro-
RECEIVED for review August 25,1986.
Accepted December
1, 1986.
Method for Determination of Oxygen- 18 of Hydrogen Peroxide in Rainwater Ben D. Holt* and Romesh Kumar Chemical Technology Division, Argonne National Laboratory, 9700 South Cass Avenue, Argonne, Illinois 60439 An analytical method was developed for the determlnatlon of 6”O of H,O, In rainwater to facilitate Isotopic assessment of the relative Importance of H,O, in oxidative processes in the atmosphere by which SO, Is transformed into SO:-. By this method, 20-L samples of ralnwater, acidtfied with HzSO4 and refrigerated at -5 OC untli time of analysls, were flltered; quantitatively degassed of dissolved air by a comblnatlon of evacuation, ultrasonic agitation, and helium sparglng (VUS); and treated with pulverlzed crystals of KMnO,. The O,, produced by the KMnO, oxidation of H,02 in the water, was removed from solution by the VUS treatment, converted to CO, In the heilum stream, and cryogenically separated for subsequent measurements of concentratlon and 6”O of the H,O, in the rainwater. Isotopic results obtained on the H,O,, SO:-, and HzO of rainwater samples from Illlnols, Mlchlgan, New York, and North Carolina, during the summer of 1985, indicate that 40% or more of the sulfate In precipitation was formed by H,O, oxidation of SO, in the atmosphere.
Oxygen isotopy can be used to elucidate the role of HzOz in the oxidation of SOz to sulfates in the atmosphere. In previous experiments (1) we found that the 6l80 [deviation ratio 0 of the sample in parts per thousand (760) of the 180/16
from the same ratio in the reference material, standard mean ocean water] of sulfate varied with the 6l80 of the solvent water when the 6l80of the HzOz was held constant. The results of those studies suggested the formation of an interwhich decomposes mediate five-oxygen species, Hz0z-S032-, to sulfate of unchanged oxygen isotope ratio. This structure of the intermediate species is consistent with the observed dependence of the 6l80of the sulfate on the 6l80 of the water, which suggests that three oxygen atoms of the intermediate molecule are isotopically controlled by the solvent water and the remaining two oxygens by the HzOz. Although the 6l80of the reagent-grade H202that was used in the earlier experiments was not known, the P O of the sulfates formed by oxidation of SOzwas generally lower than the 6l80of sulfates formed by catalyzed aqueous air oxidations of SOz in laboratory preparations and much lower than the ~ 3 of ~ sulfates ~0 found in rainwater. These results suggested the need for isotopic analysis of HzOzin dilute solutions and for a methodology whereby the 6l80 values of HzOz,HzO, and SO?- in rainwater could be compared, for assessment of the importance of H2O2as an oxidant of atmospheric SOz. Chemical properties of HzOzrelevant to isotopic studies are the following: HzOzdoes not appreciably exchange oxygen atoms with solvent water (2); the oxygen atom of the solvent water molecule does not participate in decomposition reactions
0003-2700/87/0359-0995$01.50/0@ 1987 American Chemical Soclety
996
ANALYTICAL CHEMISTRY, VOL. 59, NO. 7, APRIL 1, 1987 ANALYTICAL TRAIN
h FUNNEL COLD WATER CONDENSER
7
Figure 1. Flow
diagram of the method.
GAS
VENT
VICUUV
20-LITER FLASK ,ULTRASONIC B A T H
G45 SAMP-E BULB
7 MOLECULAR 'W COCO TRAP I !-780C'
Figure 2.
. . . MCNOMETER
1
Analytical train. Figure 3.
of H202,and no exchange occurs between the water and any molecular oxygen that may be generated (3);the 0-0 bond is not broken in the decomposition of H202by oxidation to molecular oxygen, and no oxygen fractionation occurs ( 4 ) ; when H202is oxidized by KMn04, the oxygen of the Mn0,does not isotopically interfere with the O2 product (5,6);H202 is susceptible to autodecomposition to O2and H20,catalyzed by impurities in the solvent water and accompanied by isotopic fractionation (3). This paper describes the development of a method for the conversion of the oxygen of H202in water (C25 wM) to CO, for measurement and for isotopic analysis. As shown schematically in Figure 1, the method consists of four steps: The O2 of air, dissolved in the very dilute aqueous solution of H202, is removed from the solvent water; the H202is oxidized to 0,; the newly formed 0, is recovered from the water in a carrier gas stream; the H202-derived O2 is quantitatively converted to CO, and cryogenically collected for subsequent measurement and mass spectrometric analysis.
EXPERIMENTAL SECTION Diagrams of the apparatus are shown in Figures 2 and 3. The all-glass analytical train (Figure 2) consists of the following: a bed of molecular sieve (-196 "C) for removal of traces of O2 from the helium carrier gas; a gas pipet for the injection of known amounts of 0, or C 0 2 into the carrier gas stream during standardization; the reaction chamber (detailed in Figure 3); a 20mm-0.d. cold trap (-78 "C) for the removal of residual water vapor from the gas stream; a cold trap (-196 "C) for the removal of CO, that is scrubbed from the water sample by the helium stream; a bed of activated charcoal (3 g, 8-10 mesh, coconut grade, in a vertically mounted quartz tube, lined with platinum gauze, 600 "C) for the conversion of O2to COP in the carrier gas stream; another cold trap (-196 "C) for the collection of the newly formed CO, from the gas stream; a capillary open-well mercury manometer for measurement of the CO,; a gas sample bulb, attached t o the train for transfer of the COz to a mass spectrometer for isotopic analysis; and another cold trap (-196 "C) in the vacuum manifold (not shown in the diagram) to protect the analytical train from contamination by oil vapors from the mechanical vacuum pump. The reaction chamber, Figure 3, consists of a 20-L round-bottom flask, a side-arm tube for the addition of crystalline KMn04, a cold-water condenser to limit the amount of water vapor swept
Reaction chamber.
by the helium stream into the -78 "C cold trap, three stopcocks for manipulation of the helium stream, and a funnel for the addition of solutions of acid and oxidants (used only during the development of the method). The round-bottom flask is supported by a stainless steel rack in an 83-L stainless-steel tank (interior 40.6 X 50.8 X 40.6 cm, Model ATH 1620-24) of an ultrasonic cleaning system (Model EMa 50-24) manufactured by Branson Cleaning Equipment Co., Shelton, CT. The water in the ultrasonic tank is maintained at 10 "C by circulation through a refrigeration unit. The components of the glass analytical train are connected by No. 18 ball joints, sealed with solidified black wax. The hollow plug of stopcock 22 was modified to provide adequate volume in the capillary manometer ( 7 ) .
RESULTS AND DISCUSSION Conversion of O2to COP. The first part of the method to be developed was the conversion of O2 to COSin the carrier gas stream (4th box in the flow sheet, Figure 1). The procedure was a modification of a vacuum technique by which air is circulated at low pressure over graphite and platinum at 600 O C , and the resulting COz is cryogenically removed from the closed system for mass spectrometric analysis (8, 9). The yield of CO, was found to be affected both by the mean bed temperature and by the axial temperature gradient in the charcoal bed. With a furnace suitably long to provide uniform temperature throughout the charcoal bed, the maximum recovery of O2 (as C O z ) ,determined as a function of bed temperature, was -97% at -600 "C. At lower temperatures, the yield of CO, was lower due to the incomplete reaction 02 + c co2 (1) at higher temperatures, the yield of C 0 2 was lower because of the competing reaction co2 + c 2co (2) The effect of reaction 2 was further demonstrated by injecting C 0 2 into the carrier gas instead of 02.The results, confirmed that about 3% of the COS decomposed a t 600 "C. The 97% recovery was adequate for isotopic studies, and, a t 600 "C, oxygen isotope fractionation, within the equilibrating system of COS,CO, unconverted 02,and fixed oxygen
-
ANALYTICAL CHEMISTRY, VOL. 59, NO. 7, APRIL 1, 1987
on the charcoal, was negligible. The demonstrated reliability of the 97% recovery at 600 "C also allowed application of the technique to the quantitative determination of O2 in inert gas streams. Removal of Dissolved O z from Water. Next, the technique for the quantitative removal of dissolved oxygen from multiliter quantities of water was developed (first and third boxes in Figure 1). After various combinations of sparging with helium, vacuum pumping, and ultrasonic agitation were tested, the best results were obtained by a combination of all three (IO). A low-pressure sparge by helium, combined with ultrasonic agitation, yielded removal of -99% of the dissolved oxygen in 1 h for 3-L water samples and in 5 h for 20-L samples. Oxidation of H z 0 2t o O2 in Water. Bromine water was at first tested as an oxidant of H z 0 2to Oz in the degassed water sample. Although it was effective in converting known amounts of H202to 02,its use was undesirable because of its , , contamination of the analytical train. Potassium per- , manganate was used in subsequent experiments. The average recovery of H202,added as 460-pmol spikes to seven samples of degassed water, was 96 f 3%. This recovery was adequate for the determination of oxygen isotope ratios and for the measurement of micromolar concentrations of H 2 0 2in rainwater. Procedure blanks were 3-5 pmol of O2 in 20 L of water. Isotopic Interference. As a test for isotopic interference by oxygen exchange between the HzO and either the HzOzor the Oz, before, during, or after the oxidation reactions, the oxidation was carried out in the presence of three different water supplies of various 6l80. The results showed that the 6l80 of the product C 0 2 was unaffected by the of the water solvent (11). Autodecomposition of H z 0 2 . Results from two sets of experiments, performed to assess the isotopic effects of autodecomposition of H202(added to 20-L samples of rainwater) during storage periods of up to 11 days, showed that when the rainwater was stored at room temperature, without being acidified, the H z 0 2 concentration declined rapidly ( -30% depletion in 2 days; -98% in 11 days) and that the 6l80 of the undecomposed HzOzunderwent a corresponding increase (-5%0 during the first 5 days). However, when the rainwater was first acidified (1mL of concentrated H2S04/Lof water) and kept cold (1-9 "C) during storage, the changes in the concentration and in the 6 l 8 0 of the HzOz were negligible. Consequently, the adopted procedure for storage of rainwater prior to analysis was to fill a 25-L polyethylene bottle with rainwater, add 25 mL of HZSO,, and place the bottle in a refrigerated room at -5 "C. For rainwater samples obtained at sites other than Argonne, the 25-L samples were first cooled by dry ice to just above freezing, shipped by air express to Argonne in insulated containers, and then stored in the cold room until analysis. KMnO, Decomposition. Decomposition of excess KMnO, by organic matter in rainwater was experimentally shown to produce only COz (cryogenically removed in the analytical train prior to conversion of the H20z-derivedO2to COz). Since the reaction between KMn0, and organic matter did not produce 02,it was not a source of interference in the method. Application of t h e Method. Relationship of 6180s04~to 6 1 8 0 ~ zand ~ 2 d ' 8 0 ~ , ~ .Four stock solutions of H202 (-200 pequiv/mL), each of different 6 1 8 0 H 2 0 2 , were prepared from water supplies of correspondingly different 6180H,o by a high-voltage discharge technique (12). From these four stock solutions of hydrogen peroxide, each of different 6l80, sulfate solutions of correspondingly different 6l80 were prepared by oxidation of SOz [which was in isotopic equilibrium with its solvent water of constant P O = 7.9%0 ( I ) ] . When the 6l80 of each resulting sulfate was plotted vs.
997
A
0
Argonne
0 Dearborn
v
A Whiteface Mountain V
Research Trianale Park
lo]
-10O
++
)a*
jsb
@G
5e
1985
Figure 4. Isotopic data for SO,'-, sampling sites, summer 1985.
H,O,
and H,O in rainwater at four
the 6l80 of the HzOz,the equation of the best-fit regression curve was
+
6180s04~-= 0.436180~20, 3.5%0
(3)
The regression curve of the previously determined relationship between 6180sotand P O H in aqueous-phase oxidation of SOz by Hz02of constant S1%0 was 6180s04~-= 0.576180~,o - 2.4%~(I)
(4)
Since all significant effects of the 6l80 of the SOz on the P O of the are lost by rapid isotopic exchange between the SOz and the large excess of solvent water, prior to appreciable oxidation ( I ) , 6180H2qand 6180H20zremain as the only comin the complementary variables in the equation for 6180s~4~prehensive regression curve, 6180s04~= 0.576180~,o+ 0.436180~2~, +C
(5)
The average value of C for five preparations of sulfate was found to be 8.4 f 0.6%. The resulting relationship 6180s042-
= 0.576180H,0
+ 0.436180H20,+ 8.4 f O.6%0
(6)
can now be used to calculate 6l8OSo4z-from measured 6180H202 and 6l8OHZo of rainwater for comparison with corresponding measured values of 6180s04z-. This comparison is uniquely useful in the assessment of the importance of H202 in the oxidation of SO2 to in the atmosphere. 6180Hz02, 8180S04z-, and 6 1 8 0 H 2 in ~ Rainwater. Our newly developed analytical method was applied to rainwater samples, collected during the summer months of 1985 a t four sites: Argonne, IL; Dearborn, MI; Whiteface Mountain, NY; and Research Triangle Park, NC. The isotopic data are presented in Figure 4. These results show that a unique characteristic of atmospheric peroxides is the very high P O (45-60%0), relative to the 6l80's of air oxygen (-23.5%0, not shown), sulfates (13-17%0),and precipitation water (-6% to O%O). Since the relatively high 6l8O of atmospheric peroxides is probably related to the chemical mechanism(s) by which they are formed, further isotopic studies may be uniquely applicable to investigations of the origin(s) of peroxides in the atmosphere. 6l80 of Atmospheric Sulfates: Measured us. Calculated. The measured 6l80 values of sulfates, Figure 4, are coplotted in Figure 5 with values calculated for H202oxidation (using eq 6), for catalyzed aqueous oxidation in air (1) and for pri-
998
ANALYTICAL CHEMISTRY, VOL. 59, NO. 7, APRIL 1, 1987
40
Z Argonne 0 Dearborn
1
A Whiteface Mountain
V R e s e a r c h Triangle Park
60 L
20
z W
-1
d
0
50-
L1L W L
40-
'@I
~
0-6".Vko
/._.
,
-
..- ->- - - - 0 -.----=-? .7 -
v .. &.
CALCULATED Kl-( a q u e o u s air oxidation)
-
30-
"-.
0-
0 Argonne
A Whiteface Mountain
3 Dearborn
5 Research Triangle Pork
;1/ 10
1985
Figure 5. 6I8Oof S O : -
in rainwater: measured vs. calculated values.
mary sulfates (13). The measured values are clearly higher than those calculated for catalyzed aqueous oxidation in air and clearly lower than those calculated for H202oxidation or for primary sulfates. In 1982, we proposed that since measured values for 6180~~,s in precipitation water were consistently lower than the calculated values for primary sulfates and consistently higher than those of secondary sulfates produced by air oxidations, the rainwater samples contained mixtures of primary and secondary sulfates, and, from the isotopic data, the fraction of primary sulfates a t a given site could be estimated (14). At that time, however, the isotopic qualities of atmospheric peroxides, and of the sulfates which they might produce, were not known. Since then, we have performed isotopic field experiments near a strong source of primary sulfates (15). Results indicate that precipitation scavenging was -300 times more efficient for sulfates than for SOz. We would expect, therefore, that during precipitation, scavenging of primary sulfates would be essentially complete within a relatively short distance (- 10 km) from of the source and that beyond that distance, the sulfates in precipitation would be largely secondary sulfates, formed within the storm system by one or more mechanisms of SO2 oxidation, along with a relatively small fraction of secondary sulfates scavenged from the air. Fraction of Atmospheric Sulfates Formed by Peroxide Oxidation. Assuming that primary sulfates are effectively scavenged by rain within a short distance from the source, the fraction of sulfates in rainwater (not near a strong source) which can be attributed to HzOzoxidation can be estimated from the relative deviations of the measured 6180so,2from the respective values calculated for peroxide oxidation and for catalyzed aqueous oxidation by air (see Figure 5 ) . The calculated percentages of sulfates formed by peroxide oxidation, a t our four sampling sites during the summer of 1985, are plotted vs. time in Figure 6. The peroxide oxidation is shown to represent an average of -40% of the sulfates in precipitation. The average of -40% for all of the values at all of the sites is probably a lower limit for peroxide oxidation. The 6l80 of the peroxides that caused the oxidation of SO2 in the cloud droplets could have been somewhat lower than that which was finally measured in the ground-based rain samples. That is, autodecomposition of the peroxides, between the time of
Figure 6. Fraction of rainwater sulfate formed by peroxide oxidation, assuming negligible amounts of primary sulfates. oxidation of sulfur(1V) in cloud droplets and the time of collection in the rain sample, could have been accompanied by a corresponding enrichment of P O in the residual peroxides. In that case, the calculated for sulfates formed before autodecomposition would be correspondingly lower and the percentage of sulfate formed by peroxide oxidation correspondingly higher. The technological significance of our findings, that 40 70 or more of acid sulfate in summer rains in the eastern U S . originates from peroxide oxidation of SOz, is that further investigation of the origin(s) and of possible methods of control of atmospheric peroxides is needed. Although the method was developed primarily for isotopic analysis of H202in dilute solutions, it also yields concentration data. The usefulness of the method for simultaneous measurement of concentration should be further investigated by comparison to other methods currently in use for the analysis of rainwater (16-22).
-
ACKNOWLEDGMENT We are grateful to Mark Dubois, Whiteface Mountain Field Station, Atmospheric Sciences Research Center, to Cary Eaton, Research Triangle Institute, and to William Pierson and Wanda Brachaczek, Ford Research Engineering Staff, for collection and shipment of rainwater samples to Argonne for analysis. We are also grateful to Marcia C. Dodge, U.S. Environmental Protection Agency, for assistance, review, and interpretation of results. Registry No. HzOz,7722-84-1; H20,7732-18-5; SO2,7446-09-5; "0, 14797-71-8. LITERATURE CITED (1) Holt, 8.D.; Kumar, R.: Cunningham, P. T. Atmos. Environ. 1981, 15,
557-566. (2) Anbar, M.; Guttman, S.J . Am. Chem. SOC. 1961, 83, 2035-2037. (3) Schumb, W. C.;Satterfleld, C. N.; Wentworth, R. L. Hydrogen Peroxide; Reinhold: New York, 1955; Chapter 7. (4) Cahill, A. E.: Taube, H. J . Am. Chem. SOC. 1952, 7 4 , 2312-2318. (5) Cahill, A. E.: Taube, H. J . Am. Chem. Soc. 1952, 7 4 , 2312-2318. (6) Dole, M.; Rudd, D. P.; Muchow, G. R.; Comte, C. J . Chem. fhys. 1952, 20, 961-968. (7) Hok, B. D. Anal. Chem. 1955, 1500-1501. (8) Horibe, Y.; Shigehara, K.; Takakuwa, Y. J . Geophys. Res. 1973, 78, 2625-2629. (9) Holt, B. D. Anal. Chem. 1977, 4 9 , 1664-1667. (10) Holt, B. D.; Kumar, R. Report No. EPA/600/3-86/035, July 1986; Unlted States Environmental Protection Agency: Research Triangle Park, NC: pp 1-34.
Anal. Chem. 1987, 59, 999-1002
(11) Holt, B. D.; Kumar, R. I n Fossli Fuels Utlllration Environmental Concerns; Markuszewski, R., Blaustein, B. D., Eds.; ACS Symposium Series 319; American Chemical Society: Washington, DC, 1986; pp 277-283. (12) Voihov, I. I.; Tsentsiper, A. B.; Chamova. V. N.; Latysheva, E. I.; Kuznetsova, 2 . I. Russ. J . Phys. Chem. (Engl. Trans/.) 1964, 38, 645-648. (13) Wok, B. D.; Kumar, R. Atmos. Environ. 1984, 18, 2089-2094. (14) Holt, B. D.; Kumar, R.; Cunningham, P. T. Science (Washington, D . C . ) 1982, 2 1 7 , 51-53. (15) Holt, 6. D.; Nielsen, E.; Kurnar, R. In Precipitation Scavenging, Dry Deposition, and Resuspension; Pruppacher, H. R., Semonin, R. G., Siinn, W. G. N., Eds.; Elsevier: New York, 1983; Vol. I , pp 357-368. (16) Zika, R.; Saltzman, E.; Chameides, W. L.; Davis, D. D. J . Geophys. Res. 1982, 87, 5015-5017.
999
(17) Yoshizumi, K.; Aoki, K.; Nouchi, I.; Okita, T.; Kobayashi, T.; Kamakura, S.; Tajima, M. Atmos. Envlron. 1984, 18, 395-401. (18) Schone, E. 2. Anal. Chem. 1984, 33, 127. (19) Lazrus, A. L.; Kok, 0. L.; Gltiln, S. N.; Lind, J. A.; McLiaren, S. E. Anal. Chem. 1985, 57, 917-922. (20) Klockow, D.;Jacob, P. Chemistry of Multiphase Atmospheric Systems; NATO AS1 Series, Vol G6; Springer-Veriag: Berlin, 1986. (21) Kok, G. L.; Thompson, K.; Lazrus, A. L. Anal. Chem. 1986, 58, 1192-1194. (22) Lazrus, A. L.; Kok, G. L.; Lind, J. A,; Gitlin, S. N.; Heikes, B. G.; Shetter, R. E. Anal. Chem. 1986, 58, 594-597.
RECEIVED for review May 30,1986. Accepted November 7, 1986.
Liquid Secondary Ion Time-of-Flight Mass Spectrometry James K. Olthoff, Jeffrey P. Honovich, and Robert J. Cotter* Department of Pharmacology and Molecular Sciences, Johns Hopkins University, Baltimore, Maryland 21205
A pulsed ion beam gun is interfaced to a pulsed drawout (Wiiey/McLaren) type thesf-flight analyzer. The field-free ion source is well-adapted to the use of a liquid matrix. Large primary ion pulses are used and the resulting secondary ions are recorded by us8 of fast analog to dlgltai techniques. Time delays between ion formation and extraction from the source region improve focusing and allow observation of fragment ions from metastable decompositions.
The use of primary ion beams, with energies in the kiloelectronvolt range, for the desorption of intact molecular ion species from nonvolatile organic compounds began with the introduction in 1970 of “molecular” or “static SIMS” (secondary ion mass spectrometry) by Benninghoven ( I , 2 ) . Primary ion current densities on the order of 1nA/cm2 used in this technique result in less sample damage and avoid the charging effects of “dynamic SIMS” which employs current densities of 1 pA/cm2 or more (3, 4 ) . Because of the low secondary ion currents which are a consequence of this approach, the use of time-of-flight analyzers, which have high ion transmission and the ability to record ions of all masses simultaneously, is particularly advantageous. In the SIMSTOF instrument first reported by Chait and Standing ( 5 ) ,a 1-nA beam was deflected onto the sample surface for approximately 10 ns, producing secondary ions which could be recorded by single ion counting techniques. The fast atom bombardment (FAB) technique employs primary neutral atom currents that are of the order of those used in dynamic SIMS (6),but the sample damage is reduced through the use of a liquid sample matrix (7,8). The use of neutral atoms is not essential (9),but convenient for highvoltage ion sources of double-focusing sector instruments on which the technique has been most successfully employed. Thus FAB has also been termed ”liquid SIMS” (IO). Recently we introduced the possibility of combining the FAB, or liquid SIMS, technique used on scanning instruments with the time-of-flight analyzer (11). Primary ion currents of 1pA were directed at a liquid sample in a field-free source for periods of 5 ps. This liquid SIMS-TOF instrument is also referred to as a “high flux TOF”, since the large primary ion pulse produces conditions which are similar to that of dynamic
SIMS, with secondary ion currents which are best recorded by analog (rather than ion counting) techniques. Spatial and energy focusing of secondary ions emitted from the less-defined surface are achieved by time-lag focusing as described by Wiley and McLaren (12). In this paper, the addition of fast (100 MHz), repetitive (50 spectra/s) transient recording techniques, postacceleration detection, optimization of primary ion pulse characteristics, and extension of the drawout pulse width, which leads to mass ranges above 13000 amu are reported. In addition, extension of the time-lag focusing technique provides a means for monitoring metastable fragmentation.
EXPERIMENTAL SECTION The time of flight mass spectrometer is a standard CVC (Rochester, NY) Model 2000 electron impact mass spectrometer with a 1-m fight tube, which has been modified as described below as to carry out high-flux SIMS measurements (Figure 1). A Kratos (Ramsey, NJ) Minibeam I ion gun has been fitted to the source chamber through a standard 2.75411. conflat port. The ion gun consists of an electron impact source and lens elements for accelerating, focusing, and rastering of the ion beam. The beam is focused onto a direct insertion copper probe tip located 2 in. from the front of the ion gun. A Keithley 410B electrometer connected to the probe is used to measure the average ion current reaching the probe tip. An additional 4-in. diffusion pumping system has been added to the source chamber for differential pumping of the ion gun. Argon or xenon gas pressures of up to 1 mtorr can be maintained in the ion source of the gun, while pressures of lo4 torr or less are maintained in both the source and analyzer regions of the mass spectrometer itself. In the Minibeam Control Unit, the emission regulator circuit has been replaced with a pulse amplifier which floats at the accelerating voltage. The 30-V electron grid pulse from the CVC timing circuitry is transmitted through a capacitor (680 pF, 6 kV) and amplified to 100 V. The pulse width can be varied from 1 to 10 ps and is used to control the grid to filament voltage of the ion gun source to turn the ion beam on and off. Following the primary ion pulse, the secondary ions are extracted by two negative voltage pulses (-150 and -300 V) applied, respectively, to the drawout and fiist acceleration grids. The rise time of the drawout pulse is 40 ns. The “time lag focusing” circuit on the CVC instrument has been modified to permit time delays of up to 20 ps between the primary ion beam pulse and the drawout pulse. In addition, the drawout and accelerating pulses have been lengthened to 8 ps in order that heavy ions leave the source with the full extraction field. The ions are accelerated t o their full
0003-2700/87/0359-0999$01.50/00 1987 American Chemical Society
