31
Anal. Chem. 1988, 60, 31-34 (10) Blades, M. W.; Caughlln, B. L. Spectrocelm. Acta, Part6 1085, 408, 579. (11) Walker, Z. H.; Blades. M. W. Smctrochlm. Acta. Part 6 1S88. 476. 761. (12) Burton, L. L.; Blades, M. W. Appl. Spectrosc. 1088, 4 0 , 265. (13) Fuller, C. W. "Electrothermal Atomization for Atomic Absorption Spectrometry"; Analytlcel Sciences Monograph ; The Chemical Socle-
ty: London, 1977.
RECEIVEDfor review July 15,1987.Accepted September 17, l987. The authors thank the Natural ~CiM'IceSand Engineering Research Council of Canada for financial support.
Effect of Ion/Molecule Reactions on Ion Mobilities Jonathan M. Preston*' and Lena Rajadhyax Defence Research Establishment Ottawa, Department of National Defence, Ottawa, Ontario K1A 024, Canada
Ion-mobiiity spectra can exhibit significant variations of mobiilty with temperature due to ion/moiecuie reactions occurring in the drift region. I n this paper the observed mobility for reactions in equilibrium has been shown to be calculable from the mobilities of the reacting species and the enthalpy change of the reaction. Several examples are included.
fraction of that ion. The observed reduced mobility constant, then
Ko(eq),is
~ o ( e q= ) xMH+KO(~H+)
+ xMZH+KO(
MZH+)
(1)
where X and KO are the mole fractions and reduced mobility coefficients, respectively, of the ions. To relate mobility to thermodynamic constants one writes
INTRODUCTION AND THEORY The occurrence of an ion/molecule reaction in the drift tube of an ion-mobility spectrometer (IMS)can have a considerable effect on the observed ion mobilities. If the mobility coefficients are used as the prime means of identifying the ions in each mobility peak, ion/molecule reactions must be considered. Even if mass spectrometry is used to identify ions, any variation in mobility coefficient as experimental conditions are changed requires explanation. In this paper is presented an analysis of the relationship between the thermodynamic constants of ion/molecule clustering reactions and the resulting ion mobilities. This theory can be used to predict the temperature variation of ion mobilities based on thermodynamic constants of clustering reactions ( I ) and to calculate thermodynamic constants from ion-mobility spectra recorded over a range of temperatures. The former is of particular importance in the design of hand-held special-purpose IMS detectors IMS detectors ( 2 , 3 )now reaching the market, because these detectors use only mobility for ion identification and have only limited temperature control of their drift tubes. Trace components of the drift gas in an IMS can undergo clustering reactions with the drifting ion of the type
MH++Z+Y=MZH++Y Here the ion is a protonated molecule, MH+, while Z is the trace species and Y a third body. The trace component could be an impurity or a reage'nt gas added to influence the reactions in the ionization and reaction regions, particularly in field detectors in which the same gas is normally used in the drift region ( 2 , 3 ) . At atmospheric pressure, the reverse reaction is often fast enough that local equilibrium is achieved, which is the case we consider here. In the equilibrium case, the observed mobility is a weighted average of the mobilities of the MH+ and MZH+ ions (4, 5 ) . The weights are the fraction of the drift time during which the charge exists as each ion, ti/7 ( 4 ) where ti is the time spent as ion i and 7 is the drift time. The drift time of the average charge is apportioned between the ions in the same ratio as the concentration of the ions ( 4 ) ;thus for each ion t i / 7 = Xi, the mole
where K is the equilibrium constant of the clustering reaction. Combining these equations gives &(MH+)
K[Z] =
- Ko(eq)
Ko(eq)- K (Mi") 0
(3)
and the useful forms
- 7- - 1 + K[Z]
7 and --l + - 1 (4) K[ZI ~MH+ ~MZH+ In this paper, three ion/molecule reactions were studied by measuring the reduced mobility coefficient through the temperature range in which, for a particular reactant concentration, the equilibrium switches from predominantly readants to predominantly products. For these three reactions and for a fourth example from the literature the equilibrium coefficient was calculated by eq 3 at each temperature. For each reaction a van't Hoff plot was then used to calculate the least-squares best-fit AH. Since the reactant concentration, [Z], was known, AG and A S were also calculated. These calculated values of AH were then used, with eq 3 and 4, to calculate values for Kg(4)through the experimental temperature ranges. Both these values and the experimental equilibrium mobility coefficients were plotted in the figures herein. The temperature of the drift tube affects the mobility of ions through the gas density, through collision dynamics, and through reaction equilibria in some cases. Over a limited temperature range, the density effect is linear (4)and therefore mobility coefficients reduced to STP are independent of gas are used in density. Only reduced mobility coefficients, KO, this paper. Calculations using a core model ( 4 ) suggest that K Oshould be nearly independent of temperature or should decrease slightly with increasing temperature, due to changes in collision dynamics. This conclusion has been supported by experiments over almost 200 O C (6, 7). Thus substantial changes in KOwith temperature that occur in some chemical systems, such as those reported here, are due almost entirely to reaction equilibria.
EXPERIMENTAL SECTION 'Present address: Defence Research Establishment Pacific, FMO, Victoria, British Columbia VOS 1B0, Canada.
The IMS was similar to those used in field detectors (2,3). The sample, generated from conventional thermostated diffusion tubes,
0003-2700/88/0360-0031$01.50/0Published 1987 by the American Chemical Society
32 ANALYTICAL CHEMISTRY, VOL. 60, NO. 1, JANUARY 1, 1988
was presented to a dimethylsilicon rubber membrane, which separated the internal and external flows. Ions produced by reactions in air initiated by a 10-mCi B3Nifoil were gated into the drift region in 200-ps bursts. The drift region consisted of five gold-plated rings6.1 mm apart on ceramic spacers, plus the shutter grid and the screen grid, thus a drift length of 36.6 mm over which was applied 778 V for a field of 21.3 V/mm. The collector current was amplified in an amplifier based on an LH0042 op-amp and sampled every 31 .us with 8 bits accuracy by a custom-built sample-hold and analog-to-digital (A/D) converter board installed in a slot of an Apple IIe computer. This computer summed spectra (100 in this work), calculated KOvalues, and stored or plotted the spectra. Another slot contained a circuit, based on a 6N135 optocoupler, which generated gate pulses under computer control. The drift gas was based on air, zero gas, as supplied by Matheson, Ltd. The reactant gases were either water, present at 10 ppm in this air, or acetone, added from a permeation tube (Analytical Instruments Development, Inc.) thermostated to 25 "C, at which temperature the batch-calibrated efflux was 261 ng/min, for a concentration of 222 ppb in the 497 mL/min drift gas flow rate, at atmospheric pressure. Acetone was added to the drift gas only for the 3-(3-methoxypropoxy)propanol (DPM)acetone work. The samples were 3-(3-methoxypropoxy)propanol(dipropylene glycol monomethyl ether, DPM) as sold by Dow Chemical Co. under the tradename Dowanol DPM (maximum of 0.25% water) and pyridine and acetone (Aldrich Chemical Co.), both 99+% purity. The sample was presented to the membrane inlet at concentrations of typically 6,125, and 500 ppm for DPM, pyridine, and acetone, respectively, of which about 10% passed through the membrane, depending on the membrane temperature. A thermistor-based temperature controller maintained the desired temperature of the drift tube by powering an insulated circumferential heater. A coil carrying an antifreeze solution provided cooling. Sudden temperature changes, particularly cooling, often sharply reduced the intensities of mobility peaks. Spectra were not recorded until the peaks had been stable for 30 min or more. Peak widths, at half maximum, showed no variation with temperature, but peaks recorded at low temperatures had much reduced intensities, presumably due to larger diffusion and recombination losses. The equilibria discussed here would not have been possible without the presence of neutral reactant gas in the spectrometer. It can be shown, following Seigel (8),that the ion density in the ~ . minimum source of our spectrometer was about 2 X lo9~ m - The number of reactant molecules remains if each ion has the largest possible number of molecules clustered to it, namely two for acetone and eight or so for water. Even so, more than 99.9% of the reactant molecules remained neutral. Sample concentrations in the internal flow, behind the membrane inlet, were measured in the experiments in which acetone was used as a sample gas, by trapping the sample in a known fraction of the flow from the exhaust of the IMS in a bubbler charged with diethyl succinate and cooled in an ice bath, much as in ref 9.
RESULTS AND DISCUSSION DPM-Acetone. Typical spectra of DPM with acetone as the reactant gas appear in Figure 1, and reduced mobility coefficients over the temperature range studied are plotted in Figure 2. DPM is of interest in IMS primarily because it mimics the ion-mobility spectra of a particular chemical warfare agent, and for this reason it has been studied on an IMS/MS (10). The mobility spectra of acetone have also been mass-analyzed (10)because it is used as a reactant gas in a portable chemical-agent detector (2). The agreement between the reduced mobility coefficients obtained in the IMS/MS studies and in the present work allow the ion identifications shown in the figures to be made with confidence. The only mobility peak which exhibits a signifcant variation of KOwith temperature is that due to the equilibrium of the clustering reaction shown. and Values of the limiting mobility coefficients, KO(MZH+), are required in order to calculate the equilibrium
t
t
i
P & E
9
10
11
I2
13
14
15
DRIFT TIHE Ins1
Figure 1. Ion mobility spectra of DPM in air containing 0.2 ppm acetone, at the temperatures, pressures, and vertical magnifications indicated. Numbers atop peaks are K O values. Ion identities are known through IMS/MS work (70) and are given in Figure 2.
4 14t
TEMPERATURE ('C)
Flgute 2. Reduced mobility coefficients of ions of DPM and acetone in air containing 0.2 ppm of acetone (Ac). The solid llne is KOfor the
reaction equilibrium shown, calculated by the method presented here with AH " derived from the experimental points. constants. For this reaction KO(MH+) was found easily by operating the IMS without acetone; KO(MH*) = 1.70 cm2 V-' s-' throughout the temperature range (Figure 2). The mobility of the cluster ion can be found only by operating in conditions in which the charge is on the cluster ion throughout, or for almost all of, the drift time. The key is to select the reactant-gas concentration such that no further decrease in Kg(eq) is seen below some temperature, as exhibited in Figure 2. The limiting value was taken to be = 1.43 cm2 V-' s-'. The thermodynamic constants were then derived by calculating the equilibrium constant at the temperature of each experimental point from the limiting KOvalues and the KO value at that temperature, by eq 3. The least-squares best fit of the van't Hoff plot (correlation coefficient 0.97) was In (K,) = -8.774 + 8070/T, leading to the constants listed in Table I. This best-fit equation was then used, with eq 3 and 4, to calculate values for the reduced mobility coefficient of the equilibrium ion cluster throughout the temperature range, and these values were plotted as the solid line in Figure 2 for comparison with the experimental values. Acetonewater. Clustering with water can be avoided only by using drift gases with parts per million or less water concentrations. IMS/MS results (10) show that the predominant charged acetone species in a drift tube near room temperature are the hydrated monomer, (Ac)(H,O)H+, and the dimer, (AC)~H+, with reduced mobility coefficients near 2.06 and 1.82
ANALYTICAL CHEMISTRY, VOL. 60, NO. 1, JANUARY 1, 1988
33
Table I. Thermodynamic Constants of Gas-Phase Clustering Reactions MH+ + Z = MZH+ as Measured by IMS (This Work) and by High-pressure MS (11, 12)' IMS M
-AH"
2
DPM
Ac
Ac
H2O PYr H20
PYr Pyr
67 99 139 44
high-pressure MS
-AGO
f7 f 10
45 32 38 29
f 25 f9
f f f f
-AHo
-AS"
6 4 9 6
71 f 40 226 f 50 260 f 80 50 f 50
-AGO
68 f 3 32 f 2
103 f 2 63 1
*
-ASo
T
118 f 6 104 f 4
25 25 127 25
[ :
;r;
'Ac acetone, Pyr = pyridine. Energies in kJ/mol, entropies in J mol-' K-l. AGO is at temperature ("C) given in last column. Mass confirmation of ion identities is discussed in the text and in the caDtions. Errors in IMS data were estimated from van't Hoff dots. 2.2 21
I
f
br
2
u"
20
I
10
1
1
20
30
I
I
I
I
I
40
50
80
70
80
TEMPERATURE ('C)
F w e 3. Reduced mobility coefficients of Ions known by IMSIMS (70) to contain a single acetone molecule, in air containing 10 ppm of water. The solid line was calculated as in Figure 2.
-5
1
I
I
I
I
l
l
1
'
-
20-
19-
9
1716 15
cm2 V-l s-l, respectively. As the temperature is increased, the peak due to the hydrated monomer drifts more quickly, even after correction to STP, almost certainly due to the shifting equilibrium of the hydration reaction. However we know of no IMS/MS data to confirm the assignment of this faster ion. The experimental points shown in Figure 3 were analyzed as before, and the van't Hoff plot (correlation coefficient 0.986) leads to the values in Table I. As before, the experimental best-fit van't Hoff equation was then used to calculate values for the reduced mobility coefficient throughout the temperature range, for comparison with the experimental points. As with the DPM-acetone studies, there are no literature values for thermodynamic constants of this reaction in the gas phase. Comparisons could be made with similar reactions (13). Pyridine-Water. The reduced mobility coefficients of a pyridine-containing ion, as measured in this work from 11to 83 "C, are plotted in Figure 4 along with mobility data measured by Lubman (7) and by Lubman and Kronick (14) between 85 and 220 "C. The two sets of data agree precisely at 83-85 "C. Also, recent work by Kolaitis and Lubman (15) at 210 OC have identified the main ion at that temperature (KO= 2.17 cm2 V-' s-l ) as (Pyr)H+. As discussed above, the two limiting values of KO,corresponding to the mobilities of the two ions in the reaction, can be obtained (if the reactant gas cannot be eliminated from the IMS) only by selecting the reactant gas concentration and the temperature range such that there is no further noticeable variation in KOat the temperature extremes. At these extremes the equilibrium is almost all either reactants or products, so further changes have little effect on the equilibrium reduced mobility coefficient. With appropriate selections, the plot of Kg(Bq)vs T resembles a flattened "s", horizontal at both asymptotes, as in the two examples described above. In this case, the plot is suggestive of two such shapes, from 2.16 to 1.69, and from 1.69 to 1.54 cm2 V-l s-l. If these are analyzed as two separate reactions, the van't Hoff plots (correlation coefficients 0.9998 and 0.967) lead to the data in Table I. As before, the values of KO(4 based on these enthalpies appear as the solid line in Figure 4 and fit the observations well. The identity of the high-temperature reaction is known through IMS/MS data published recently (15)which showed that the two major ions at 105 "C were (Pyr)H+ and ( P ~ T ) ~ H + . The reaction is thus the dimerization of pyridine. Thermodynamic constants for this reaction are available from high-
I
1 4
'
1
1
1
'
1
1
1
I ' I I ' I ' I I
Anal. Chem. 1980, 60,34-37
34
experimental values of KO(* agree very well, for the most part, with theoretical values based on thermodynamic constants derived by the method presented. Agreement with literature values for the constants is an essential requirement for the establishment of the method. While it is clear that thermodynamic constants cannot be calculated from ion mobility spectra with accuracies comparable to the established MS techniques, the reasonable agreements listed in Table I lend strong support to the method presented here.
ACKNOWLEDGMENT L. J. Hart designed and implemented the computer interfaces and programs without which this work would have been much more difficult.
LITERATURE CITED Castleman, A. W., Jr.: Keesee, R. G. Chem. Rev. 1986, 86, 589-618. Blyth, D. A. I n Proceedings of the International Symposlum on Protection Against Chemical Warfare Agents, Stockholm, 1983; pp 65-69. Carrico, J. P.; Davis, A. W.; Campbell, D. N.; Roehl, J. E.;.Slma, G. R.; Spangler, G. E.; Vora, K. A,; White, R. J. Am. Lab. (Fairfield, Conn.) 1986, 18, 152-163.
(4) Mason. E. A. I n Pklsma ChromtOaretW: Carr, T. W.. Ed.; Plenum: New York, 1984; Chapter 2. (5) Carroll, D. I.; DzMIc, I.; Stlllwell, R. N.; Hornlng, E. C. Anal. Chem.
im
47 1956-1959 .__ ._.
(6) Parent, D: C.; Bowers, M. T. Chem. Phys. 1981, 6 0 , 257-275. (7) Lubman, David M. Anal. Chem. 1984, 56, 1298-1302. (8) Slegel, M. W. I n Plasma Chromatcgraphy;Carr, T. W., Ed.: Plenum; New York, 1984; Chapter 3. (9) Casselman, A. A.; Gibson, N. C. C.; Bannard, R. A. B. J. Chromatogr. 1973, 78, 317-322. (IO) Harden, C. S., Chemlcal Research Development and Engineering Center, Maryland, unpublished work, 1985. (11) Davidson, W. R.; Sunner, J.; Kebarle, P. J. Am. Chem. SOC. 1979, 101, 1675-1680. (12) Moet-Ner (Mautner), Michael; Sieck, L. Wayne J. Am. Chem. SOC. 1983, 705, 2956-2961. (13) Lau, Y. K.; Saluja, P. P. S.;Kebarle, P. J. Am. Chem. SOC. 1980, 102, 7429-7433. (14) Lubman, David M.; Kronick, Me1 N. Anal. Chem. 1983, 5 5 , 1486-1492. (15) Kolaitis. Leonidas; Lubman, David M. Anal. Chem. 1986, 5 8 , 1993-2001.
RECEIVED for review February 5, 1987. Accepted August 31, 1987* This work was presented in part at the Canadian Chemical Conference, Saskatoon, SK, 1986.
Isotopic Analysis of Lithium as Thermal Dilithium Fluoride Ions L. W. Green,* J. J. Leppinen, and N. L. Elliot
General Chemistry Branch, Chalk River Nuclear Laboratories, Chalk River, Ontario, Canada KOJ 1JO
A U Isotopic analysrS method based on measurement of Up+ lons by thermal lonlzation mass spectrometry was developed. The Li,F+ ions were volatlzed from sample fllaments by radiant heat from the center filament; the optimal mole ratio for fllamenl loading was 211 LI/F. The M o p e fractionation rate with this molecular specles was low and was negligible over the normal analysis perlod. For appllcatlon of the method to samples wlth complex mafrlxes, a two-column lonsxchange process was developed for separatlon of LI from the sample matrix and conversion to LlOH for mlxlng wlth HF. The preclslon of the entlre procedure, InclUcAng the chromatography, was 0.2 % relatlve standard deviation, and the accuracy was wkhln the same range provided m a s and column blas factors were corrected for.
The isotopes of lithium have many practical uses in the biomedical ( I ) , geological (2), and nuclear industries (3). For example, in the nuclear industry, 6Li is used as a tritium breeder in fusion reactor blankets because of its high cross section for the n,a reaction. The natural abundance of the mass six isotope is only -7.5%, thus for tritium generation enrichment of X i is desirable. Single-stage enrichment fadors of up to 4 % have been obtained by chemical exchange techniques ( 4 ) and currently several of these are being studied in our laboratories. For all of the above uses of Li, precise measurements of the 617 ratio are required. A number of laboratories (5-8) have reported isotopic analysis techniques that achieved reproducibilities between 0.1% and 0.7% RSD; in some cases mass spectrometers specifically designed for Li isotopes were used ( 5 , 6 ) . Recommended filament loading forms were the iodide,
chloride, and nitrate salts of Li, and atomic ions (Li') were generated by interaction of the sample vapors with an ionizing filament, except in the procedure of Brown et al. (6) in which electron impact ionization was used. For the technique used with large magnetic sector instruments designed for heavier elements, very rigid control of sample deposition and analysis parameters was required, especially for the ionizing filament temperature (8). Control of isotope fractionation is the most important factor in Li isotopic analysis, because of the low atomic mass and high relative mass difference between the isotopes. To obtain reproducible results, the volatilization rate of each ionic and molecular species must be constant or controllable during the analysis and reproducible from one loading to the next (9). However, determination of the volatilization rates of all of the important species is not possible with thermal ionization, because the neutral species are invisible. The practical approach has been establishment of reproducible volatilization rates for the atomic positive ions followed by further optimization of procedures until reproducible results were obtained (5-8). An alternative approach is to use a molecular species that greatly reduces the fraction of lithium vaporized as atomic ions or neutrals. Theoretical and experimental data have shown that the ratio of molecular to atomic species volatilized is a major factor in the fractionation rate, expecially for triple filament sources in which the evaporation and ionization processes are spatially separated (9-1 1). Furthermore, estimated decreases in fractionation rate caused by increases in proportion of molecular species volatilized were largest for lithium, which was the lightest of the elements modeled by Kanno (9). Of the compounds of lithium that would be of practical use for thermal ionization techniques, the fluoride compounds are
0003-2700/S8/0360-0034$01.50/0 @ 1987 American Chernlcal Society
