Anal. Chem. IQSO, 62, 1315-1319
(25) Beldwln, J. J.; PontiCellO, 0. S.; Sugrue, M. F.; Mallorga, P. J.; Randall.
W. C.; Schwam, H.; Springer, J. P.; Smith. G. M.; Murcko, M. Book of Abstracts; The Thkd Chemical Congress of North America, Toronto, Canada, June 5-10. 1988 Amerlcan Chemical Soclety: Washington, DC, Abstract No. MEDI 0095. (26) BeMwin, J. J.; Ponticello, 0. S.; Freedman, M. 8.; Habacker. C. N.; Christy, M. E.; Sugrm, M. F.; Malbrga. P. J.; Schwam. H.; Randall. W. C.; Springer, J. P.; Smith. G. M.; Murcko. M. loth Internatlonai Medicinal Chemism Symposium, Budapest, Hungary, August 15-19. 1988. (27) Qlaham, S. L.; Shepard. K. L.; Anderson, P. S.; Baldwln, J. J.; Best, D. B.; Christy, M. E.; Freedman, M. B.; Gautheron, P.; Habecker, P. B.; Hoffman, J. M.; Lyle, P. A,; Micheison, S. R.; Robb, C. M.; Schwam, H.; Smith, A.M.; Smith, R. L.; Sondey, J. M.; Strohmaier, K. M.; Sugrue, M. F.; Varga, S. L. J . Med. Chem. 1989, 32,2548. (28) Llppa, E. A.; Hofman, H.; Feicht, B.; Bron, A.; Royer, J.; Brunner-Ferber, F.; Von Denffer, H. 3rd Congress of the European Glaucoma So-
1315
ciety, Estorii, Portugal, May 23-26, 1988. (29) Lippa, E. A.; Von Denffer, H. A.; Hoffmann, H. M.; Brunner-Ferber, F. L. Arch. Ophthelmd. (Chicago) 1988, 106, 1894. (30) Beriman, J. B. hendbodc of Fkrwesoence Spectra of AromeHc Mdecubs, 2nd ed.; Academic: New York, 1971. (31)Matuszewski, B. K.; Givens. R. S.; Srinhrasachar, K.; Carlson, R. G.; Higuchi, T. Anal. Chem. 1987, 59,1105. (32) Lee, L. L.; Zacchei, A. G. Book of Abstracts; 198th Natlonai Meeting of the Amerlcan Chemical Society, Miami Beach. Florida. Sept 10-15, 1989; America1 Chemical Society: Washington, DC, Abstract No. MEDI 37.
RECEIVED for review September 12, 1989. Accepted March 6, 1990.
Thermodynamic Properties in Micellar Liquid Chromatography Based on the Three-phase Equilibrium Model Frank P. Tomasella and L. J. Cline Love* Department of Chemistry, Seton Hall University, South Orange, New Jersey 07079
A thermodynamic model of micellar llquld chromatography based on partltlonlngtheory Is proposed. Equations are developed that provide the enthalpy of transfer of the solute from the bulk aqueous phase to elther the micelle or the statlonary phase. Retention mechanlsm Is detailed based on the comparkon of the obtained values for each of the transfers. Correlation Is estabtlshed between the thermodynamic data and the accepted three-phase equlllbrlum model of mlcellar llquld chromatography.
liquid chromatography. In this paper we describe the effect of temperature on the equilibrium constants. Thermodynamic equations are derived which account for the two principal equilibria associated with micellar liquid chromatography. The results of this study will correlate thermodynamic properties to the currently accepted equilibrium model. The retention mechanism of micellar liquid chromatography will be described utilizing a three-phase euilibrium model and thermodynamic properties.
EXPERIMENTAL SECTION INTRODUCTION With the growth of high-performance liquid chromatography (HPLC), the use of micellar liquid chromatography continues to establish unique applications. Recent applications include the use of micellar chromatography in the quantitation of hydrophobicity ( I ) , determination of heavymetal cations (2), and the determination of drug substances in biological fluids (3-5). The selectivity of micellar chromatography is due to the large variety of interactions including electrostatic and hydrophobic that are possible between the analyte and the mobile phase, the micelle, and the stationary phase. The first partitioning theory of a solute between micellar and aqueous phases in liquid chromatography is credited to Herries (6). The partitioning theory was applied to HPLC by Armstrong and Nome (7) where equations were derived which account for the chromatographicbehavior of the solutes eluted with a micellar mobile phase. Cline Love and Arunyanart expanded the model whereby the equilibrium constant can be obtained for a solute between the bulk aqueous phase and the micellar aggregate (8).Subsequently, the retention behavior in micellar liquid chromatography is described based on correlation with the equilibrium model (9,101. Refinements to the model include the influence of ionizable species (11) and a solubility limit equation (12). Dorsey and co-workers (13)have applied thermodynamics to micellar liquid chromatography. This thermodynamic approach, however, does not address the issue of multiple equilibria which has become accepted in describing micellar 0003-2700/90/0362-1315$02.50/0
LC System. A modular component liquid chromatographic system was used consisting of a Spectra Physics SP8700 extended range pump, a LDC UV monitor detector (254 nm) (Riviera Beach, FL), and a Rheodyne sample injector equipped with a 20-pL injection loop. The column was a Brownlee Labs (Santa Clara, CA) RP-18 Spheri-10 (100 X 4.6 mm). A silica precolumn was used before the injector to saturate the mobile phase with silica to minimize dissolution of the analytical column packing. Column temperature was controlled by immersing the column in a water bath where the desired temperature was maintained by a Model 73 constant-temperature circulating bath (Fisher Scientific). A Model 5000 Fisher Recordall strip chart recorder was used to record the chromatograms. Reagents. The surfactant,sodium dodecyl sulfate (SDS),was electrophoresis grade obtained from Bio-Rad, Inc., and used as received. The benzene, phenol, naphthalene, toluene, and p-xylene (Fisher Scientific Co.) and anisole and 1-naphthalenemethanol (Aldrich Chemical Co., Inc.) were used as received. Procedure, The appropriate weight of SDS was dissolved in distilled water and the solution was then filtered through a 0.45-rm Nylon-66 membrane filter (Rainin Instruments Co., Inc., MA). A flow rate of 1.0 mL/min was used throughout the studies and 20 pL of M solute solutions was injected onto the column. Retention times were measured manually by monitoring both time and volume and found to be in agreement. The void volume, V,, of the system was measured at all mobile phase concentrations by injecting pure water or NaN03 (14) and found to be 45.0 s or 0.75 mL which was used for all k’calculations. The elution volume of NaN03 was in good agreement with the elution volume of pure water.
RESULTS AND DISCUSSION The retention mechanism in micellar chromatography is based on the solubility or electrostatic nature of the solute 0 1990 American Chemical Society
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ANALYTICAL CHEMISTRY, VOL. 62, NO. 13, JULY 1, 1990
and is more complex than reversephase chromatography. For neutral moderately water soluble solutes, the three-phase model as described by Armstrong (7) and Cline Love (8) is applicable. For ionizable species, the electrostatic model (11) must be utilized. For neutral water insoluble solutes, a direct transfer of the solute from the micellar pseudophase to the surfactant-modified stationary phase is proposed (12). Chromatographic retention in nonmicellar HPLC has been described in terms of thermodynamic properties (15-18). The thermodynamic theory has satisfactorily accounted for the interactions of the solute with the mobile phase and the stationary phase. The use of thermodynamic properties in describing micellar liquid chromatography has been limited. Due to the multiple retention mechanism that exists for micellar chromatography as cited above, a thermodynamic treatment must consider which mechanism is in effect. Hinze and co-workers conclude that the retention mechanism is analogous to reverse-phase liquid chromatography for water-insoluble solutes (12). Chromatographic retention of such solutes would be described by the same thermodynamic equations that describe a nonmicellar reverse-phase system. For moderately water soluble solutes that adhere to the three-phase equilibrium model, the thermodynamic equations derived for a conventional reverse-phase chromatography would not be appropriate. These investigations are undertaken to develop a thermodynamic model of micellar liquid chromatography based on partitioning theory. Initially, the effect of temperature on the equilibrium constants will be described. The retention mechanism will be detailed in terms of thermodynamic properties which correlate to the three-phase equilibrium model of micellar liquid chromatography. Equilibrium. The three-phase equilibrium model derived by Arunyanart and Cline Love is utilized in describing the interactions of the solute with the micelle and the stationary phase (8). The two principal equilibria are a reversible equilibrium of solute in the bulk solvent mobile phase with the stationary phase sites, K,, and a reversible equilibrium of the solute in the bulk solvent mobile phase with the surfactant in the micelle in the mobile phase, K2. The third reversible equilibrium involving the direct transfer of solute in the micelle to the stationary phase is neglected, K,. The model relates capacity factor to micellar mobile phase concentration expressed as
Table 1. Equilibrium Constants Measured from Equation 1
where [M,] is the concentration of surfactant in the micelle in the mobile phase, [L,] is the stationary phase sites, and 4 is the chromatographicphase ratio, i.e., the ratio of the volume of the stationary phase to the volume of the mobile phase in the column (8). This study limits itself to neutral molecules thereby eliminating the complexity of the interactions caused by ionization (11). Additionally, the solutes under investigation all have sufficient water solubility to conform to the three-phase model where K , is neglected. Hinze and co-workers (12) recently concluded that the direct transfer of the solute from the micelle to the stationary phase predominates when a solute is water insoluble. K1 and K2 are negligible and can be neglected under these conditions. Effect of Temperature on the Equilibrium Constants. Equation 1 is used to examine the data. In accord with eq 1, linearity is obtained from the plots of llk'vs [M,], resulting with regression coefficients as good as 0.99. The value of K2 is obtained from the slope/intercept ratio, where the intercept is (d[L,]K1)-' and the slope is K2(4[L,]Kl)-'. Berthod and co-workers observed that the absorbed surfactant on the bonded phase remains constant with micellar mobile phases (19). Therefore, a change in the slope results in a change in
Figure 1. Inverse capacity factor versus concentration of surfactant in the miceMe in the mobile phase for phenol (open symbols) and anisole (closed symbols) at column temperature of .25 OC (+), 35 OC (A), and 45 OC (0).The concentration of surfactant in the micelle is [Mm]=
intercept
(io3)
K2
temp, "C
slope
phenol
25 35 45
0.840 0.868 0.914
88.7 104 120
benzene
25 35 45
0.335 0.335 0.346
16.5 18.0 19.0
20.3 18.6 18.2
anisole
25 35 45
0.382 0.415 0.447
13.1 15.2 17.5
29.1 27.3 25.5
toluene
25 35 45
0.281 0.285 0.293
4.98 5.31 5.58
56.5 53.7 52.5
1-naphthalenemethanol
25 35 45
0.610 0.659 0.704
8.13 9.46 11.0
75.0 69.7 64.0
p-xylene
25 35 45
0.230 0.235 0.240
1.28 1.33 1.22
180 177 197
naphthalene
25 35 45
0.304 0.320 0.331
0.862 0.884 1.33
353 362 248
compound
9.47 8.35 7.62
0.24
0.18
0.12
I
0.00 0.04
0.06
0.08
0.10
0.12
0.14
(SDS MICELLES). M
[surfactant] - cmc, where [surfactant] is the total concentration of surfactant in solution and cmc is the critical micelle concentratlon. The cmc for SDS is 8.1 X M (ref 20).
the ratio of K 2 / K 1since 4 and [L,] are constant. Equilibrium constants for the seven solutes used are given in Table I. It should be noted that as the column temperature is increased, the slope increases slightly. The intercept increases with a temperature increase and is more pronounced. This means that K, decreases more rapidly than does K2 as the temperature increases, within the temperature range investigated. Hence, the solute favors the bulk aqueous phase with an increase in column temperature. To a lesser extent, the solute favors the micelle as opposed to the stationary phase. Thus, k'is observed to decrease as the temperature increases. These results are graphically illustrated in Figure 1 for phenol and anisole. In addition to the effect caused by temperature, the hydrophobicity of the solutes can be assessed by examining the intercept as its value approaches zero. The most hydrophobic
ANALYTICAL CHEMISTRY, VOL. 62, NO. 13, JULY 1,
solute in the study is naphthalene, whose intercept at 25 "C is O.OOO9. Phenol, the most hydrophilic solute, has an intercept of 0.0887. The hydrophobicity decreases in the order naphthalene > toluene > anisole benzene > phenol. These results are in agreement with the more vigorous quantitation of hydrophobicity with micellar chromatography as reported by Khaledi and Breyer (1) with the solutes that are common to both studies. Greater error is evident in the equilibrium constants associated with hydrophobic solutes such as naphthalene and p-xylene as the intercept approaches zero. The data in Table I reflect these errors. Hinze and co-workers have proposed the solubility limit theory for near-"zero" intercepts where the retention mechanism shifts to a single equilibrium defined as the direct transfer of the solute from the micelle to the stationary phase (12). The drawback to the solubility limit theory is that an exact limit is not defined for a broad range of solutes. The error that is apparent in the case of p-xylene and naphthalene indicates a transition to a single equilibrium retention mechanism is probable for solutes of greater hydrophobicity. Thermodynamics of Micellar Chromatography, The effect of temperature on retention provides thermodynamic data describing the chromatographic process. Melander, Campbell, and Horvath applied the Van't Hoff equation to reverse-phase chromatography (15) In k ' = -AHO/RT AS"/R In 4 (2)
8.00
-
+
+
where the chromatographic retention is measured by the capacity factor, k', as related to the equilibrium constant given as k' = r$K. The resultant Van't Hoff plots give enthalpies and entropies of transfer for the solutes from the mobile phase to the stationary phase. Dorsey and co-workers applied eq 2 directly to micellar chromatography (13)and obtained enthalpies of transfer from the mobile phase to the stationary phase for several solutes. These results do not consider the multiple equilibria which are possible in micellar chromatography as described by the accepted three-phase model. A more vigorous approach accounts for the two principal equilibria in micelle chromatography and provides the thermodynamics associated with each transfer. The reciprocal of eq 1 relates the capacity factor to the thermodynamic equilibrium constant expressed as k ' = (4[LsIK1)/(1 + K,[MmI) (3) Solving for K1, one obtains Kl = &'(I + K,[M,I)J/4[L,I (4) For a given transformation, two expressions for the standard Gibbs free energy are AGO = -RT In K (5) AGO = AH" - TAS" (6) Combining eqs 4,5, and 6, the following Van't Hoff expression is obtained which describes the transfer of the solute from the bulk aqueous phase to the stationary phase (K,) In k'(1 K2[Mm])= -AHl0/RT ASlo/R In $[La] (7) where k', K,, [M,], 4, and [LJ are as defied for eq 1. If AHl" and ASl" are independent of temperature over the temperature interval of interest, the resulting Van't Hoff plots will be linear. The slope of the line gives standard enthalpies of transfer of the solute from the bulk aqueous phase to the stationary phase. Figure 2 illustrates plots of eq 7. Straight lines were obtained for each solute. The change in enthalpy for the transfer of the solute from the bulk aqueous phase to the stationary phase becomes enthalpically favored with more negative values of AH,".Table I1 gives the change in enthalpy obtained in a micellar system
+
+
+
1317
1990
4.80E==q
rN f
t
"60
I
1
0.00
3.10
3.22
3.16
3.28
Temp. ( 1O3/T(K-'
3.34
3.40
))
Flgure 2. Van't Hoff plots for In (k' + k'K, (M,)) versus inverse temperature for six solutes: phenol (+), benzene (A),toluene (0), anisole (a),1-naphthalenemethanol(A),and naphthalene (e);mobile phase is 0.10 M SDS.
Table 11. Change in Enthalpy (kcal/mol). Transfer of the Solute from the Bulk Aqueous Phase to the Stationary Phase Obtained at Differing Micelle Concentrations solute phenol slope
0.06 M 0.08 M 0.10 M 0.12 M 0.14 M SDS SDS SDS SDS SDS 1.44 -2.85
ml0
1.50 -2.97
1.37 -2.71
1.42 -2.81
1.46 -2.89
average AHlo = -2.85 & 0.10 kcal/mol benzene slope
0.721 0.648 0.628 0.646 0.703 -1.43 -1.28 -1.24 -1.28 -1.39
W0
average Wl0 = -1.32 f 0.08 kcal/mol anisole slope
1.29 -2.55
mlo
1.44 -2.85
1.33 -2.63
1.38 -2.73
1.35 -2.67
average AHlo = -2.69 i 0.11 kcal/mol toluene slope
0.513 0.455 0.437 0.455 0.513 -1.02 -0.90 -0.87 -0.90 -1.02
Ul0
average AHl0 = -0.94 f 0.08 kcal/mol l-naphthalenemethanol slope M1°
1.32 -2.61
1.36 -2.69
1.32 -2.61
1.34 -2.65
1.30 -2.57
average AH1' = -2.63 f 0.05 kcal/mol p-xylene slope
W0
-0.492 -0.293 -0.506 -0.506 +0.85 +0.58 +1.00 +1.00 average MIo = +0.86
naphthalene slope M1°
2.05 -4.06
-0.448 +0.89
0.17 kcal/mol
2.14 -4.24
2.07 -4.10
2.14 -4.24
2.07 -4.10
average MIo= -4.15 f 0.11 kcal/mol
for the transfer of the solute from the bulk aqueous phase to the stationary phase for aromatic solutes with five micelle concentrations ranging from 0.06 to 0.14 M SDS. At all five concentrations of SDS, the change in enthalpy is the same for each solute within experimental error. The micelle con-
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ANALYTICAL CHEMISTRY, VOL. 62, NO. 13, JULY 1, 1990 8.001
Table 111. Change in Enthalpy and Entropy. Transfer of the Solute from the Bulk Aqueous Phase to the Micelle
I
i I
AHz", kcal/mol
6.40e
phenol anisole benzene toluene 1-naphthalenemethanol p-xylene naphthalene
-2.22 -1.20 -0.99 -0.53 -1.22 f1.41 -3.41
f 0.19 f 0.10 f 0.08 f 0.16
f 0.16 f 0.16 f 0.33
AS2O, cal/mol -2.99 +2.57 +2.61 +6.20 +4.49 +14.90 +0.34
f 0.19 f 0.10 f 0.08 f 0.16 f 0.16 f 0.16 f 0.33
Table IV. Comparison of the Change in Enthalpy for K 1vs
K2 MIo,kcal/mol 0.00 1 3.10
I 3.76
3.22
3.20
Temp. ( ?03 /T(K
-'
3.34
3.40
))
Flgwe 3. Van? Hoff pbts for In K2 versus inverse temperature for six solutes: phenol (+), benzene (A),anisole (O), toluene (0),1naphthalenemethanol(A),and naphthalene (e);mobile phase is 0.10 M SDS.
centration has no effect on the transfer of the solute from the bulk aqueous phase to the stationary phase. The average value of the change in enthalpy for each of the solutes with its standard deviation is provided. Knox and Vasvari reported values of -1.76 to -5.85 kcal/mol as the enthalpy of transfer for substituted benzenes from the mobile phase to the stationary phase in a reversephase system (16). Typical values for a reverse-phase system for a variety of compounds are accepted as -1 to -10 kcal/mol (17). The values in Table I1 range from +0.86 to -4.15 kcal/mol for the K1 transfer. The less negative values obtained with micellar liquid chromatography as compared to nonmicellar reversephase chromatography indicate that the effect of temperature on k is less pronounced in micellar chromatography. The general trend is a more negative AH with increasing k ', but instances where larger k'are associated with larger AH have been reported for conventional reverse-phase liquid chromatography (18). Similar results are evident with micelles as in the case where phenol and anisole have more negative values than benzene and toluene. The +0.86 kcal/mol obtained in the transfer of p-xylene from the bulk aqueous phase to the stationary phase conflicts with the generalization that hydrophobic solutes are associated with negative enthalpy transfers. The errors associated with K2values as the intercept approaches zero may account for the discrepency with solutes of high hydrophobicity. These observations are in agreement with the causes of error reported in the equilibrium data and may serve as additional support of a single equilibrium retention mechanism for solutes which have no water solubility as proposed by Hinze (12). In addition to obtaining the change in enthalpy for K,, one can readily obtain the thermodynamic values for K2 which describe the transfer of the solute from the bulk aqueous phase to the micelle phase as given by In K2 = - M , " / R T -k AS2"/R (8) Equation 8 will apply to a micellar chromatographic system since Kzis obtained via eq 1 which has related the capacity factor to the equilibrium constants. The resultant value of Kzis obtained from the plot of l/k'vs SDS concentration from 0.06 to 0.14 M. The values of K z are obtained for five temperature intervals according to the Van't Hoff equation given above. Graphic treatment of eq 8 for the solutes is shown in Figure 3 where In K2 is plotted vs 1/T. The plots are indeed linear and give the changes in enthalpy and entropy, which are independent of temperature over the temperature interval
phenol anisole benzene
toluene 1-naphthalenemethanol naDhthalene
-2.85 -2.69 -1.32 -0.94 -2.63 -4.15
f 0.10 f 0.11
f 0.08 f 0.08 f 0.05 f 0.11
A&O,
kcal/mol
*
-2.22 0.19 -1.20 f 0.10 -0.99 f 0.08 -0.53 i 0.16 -1.22 f 0.16 -3.41 f 0.33
of interest. The slope of the line gives the standard enthalpies of transfer of the solute from the bulk aqueous phase to the micelle phase. The intercept provides the standard entropy for the same transfer. Table III gives the changea in enthalpy and entropy for each of the solutes plotted in Figure 3. The change in enthalpy in the case of p-xylene stands out as an exception to the other solutes presented as having a positive value. The explanation lies in the fact that p-xylene has the greatest error associated with the Kzterm which is inherent in the enthalpy calculation. As the value for the change in enthalpy increases for the various solutes, the change in entropy also increases. Similar observations have been reported for nonmicellar reversephase chromatography (15, 18). The change in entropy becomes greater with increasing hydrophobicity as is evident in the data presented in Table 111. As a result, the energy associated with the entropy term becomes a dominant factor when calculating the change in free energy as per eq 6. Such calculations indicate a more negative AG" associated with the hydrophobic solutes. The free energy increases with increasing hydrophilicity as follows: p-xylene < naphthalene < 1-naphthalene methanol < toluene < anisole < benzene < phenol. The trend is indicative that a more hydrophobic solute will favor the stationary or micelle phase as opposed to the aqueous phase, which is in accordance with retention. Comparisons of the change in enthalpy for each of the two principal equilibria are presented in Table IV for six solutes of increasing hydrophobicity. The changes in enthalpy for the transfer to the stationary phase or the micelle are very similar. Therefore, the transfer of a solute from the bulk aqueous phase has an almost equal affinity for either the stationary phase or the micelle with a slight preference for the stationary phase based solely on enthalpy. It is also evident that the change in enthalpy associated with the K , process renders more experimental error than that of the K2process. This is due to the fact that Kzis a factor in determining the change in enthalpy for the K , process, as evident in eq 7.
CONCLUSION The results presented in this study have correlated equilibrium and thermodynamic properties in terms of the three-phase equilibrium model. Thus, equilibrium and thermodynamic data can be calculated for the transfer of the solute to either the stationary phase or the micelle from the bulk aqueous phase. The retention mechanism described from
Anal. Chem. 1990, 62, 1319-1324
equilibrium constants and thermodynamic properties are in good agreement. As the hydrophobicity of a solute becomes very large, the results obtained in this study support the proposal that the retention mechanism is shifting toward a direct transfer of the solute from micellar pseudophase to the surfactant-modified stationary phase. Future studies involving the effect of organic modifier and its role on efficiency can be evaluated in terms of equilibrium and thermodynamics. ACKNOWLEDGMENT The authors are grateful to Paul Ander and Joseph Fett for many helpful discussions. LITERATURE CITED (1) Khaledl, M. 0.; Breyer, E. D. Anal. Chem. 1989. 67, 1040-1047. (2) Okada, T. Anal. Chem. 1988, 60, 2116-2119. (3) Palmisano, F.; Guenieri, A.; Zambonln. P. G.; Cataldl, T. R. I . Anal. Chem. 1989. 61, 946-950. (4) Posluszny, J. V.; Weinberger, R. Anal. Chem. 1988, 60, 1953-1958. (5) DeLuccia, F. J.; Arunyanart, M.; Cline Love, L. J. Anal. Chem. 1985, 57, 1564-1568. (6) Herries, D. G.; Bishop, W.; Richards, F. M. J . phys. Chem. 1964, 68,
1842- 1852.
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(7) Armstrong, D. W.; Nome, F. Anal. Chem. 1981. 53, 1662-1666. (8) Arunyanart, M.; Cline love, L. J. Anal. Chem. 1964, 56, 1557-1561. (9) Berthd, A.; Girard. I.; Gonnet, C. Anal. Chem. 1986, 58, 1359-1362. (10) Khaledl. M. G.; Peuler, E.; Ngeh-Ngwainbi, J. Anal. Chem. 1987, 59. 2738-2747. (11) Arunyanart, M.; Cline Love, L. J. Anal. Chem. 1985. 57, 2837-2843. ( 12) Borwding, M. F.; Quina, F. H.; Hinze, W. L.; Bowermaster,J.; McNair, H. M. Anal. Chem. 1988, 60,2520-2527. (13) Dorsey, J. G.; DeEchegaray, M. T.; Landy, J. S. Anal. Chem. 1983, 55. 924-928. (14) Wells, M. J.; Clark, C. R. Anal. Chem. 1981, 53. 1341-1345. (15) Melander, W.; Campbell, D. E.; Horvath, C. J . Chromatogr. 1978, 158, 215-225. (16) Knox, J. H.;Vasvari, G. J . Chromtogr. 1973, 83, 181-194. (17) Melander, W. R.; Horvath, C. High Performance Liquhi Chromatogrephy: Advances and Perspectives; Horvath, C., Ed.; Academic Press: New York, 1980; Voi. 2, p 198. (18) Colin. H.;Diez-Masa. J. C.; Guiochon, G.; Czajkowska, T.; Miedziak, I. J . Chrometogr. 1978. 167, 41-65. (19) Berthod. A.; Girard. I.; Gonnet, C. Anal. Chem. 1986, 5 3 , 1362-1367. (20) Mukerjee, P. J . Phys. Chem. 1982, 66, 1733-1735.
RECEIVED for review November 27, 1989. Accepted March 2, 1990.
Liquid Chromatography/Time-of-flight Mass Spectrometry with High-speed Integrated Transient Recording W. Bart Emary, Ihor Lys, and Robert J. Cotter*
Department of Pharmacology and Molecular Sciences, The Johns Hopkins University, Baltimore, Maryland 21205 Richard Simpson
Department of Chemistry, University of Maryland Baltimore County, Baltimore, Maryland 21228 Andrew Hoffman
Kratos Analytical, Manchester M31 2LD, England
High-performance iiquld chromatography (HPLC) has been interfaced to a tkneof-flight mass spectrometer. The interface Is a continuous flow probe and ions are desorbed from the iiquld matrix by energetlc ion bombardment. Pulsed and delayed ion extraction from the source permits the use of broad ionization times, results in the production of analog signals in each tkne-of-fllght cycle, and provides both energy and spatial focusing. A high-speed Integrated transient recording system has been developed and isabm reported. TMS instrument Is the prototype for devebpment of a high-speed, highinass range LC detector with high duty cycle. Its performance is demonstratedfor the separatlon of several mixtures of small peptides.
INTRODUCTION A number of techniques have been developed for interfacing bonded phase high-performance liquid chromatography (LC) to mass spectrometers. The most common (and commercially available) instrumental configurations employ moving belt interfaces ( 1-3), thermospray ionization (4),and continuous flow fast atom bombardment (5-8)with quadrupole and sector (double focusing) mass analyzers. Because the ionization sources in sector instruments are generally at a very high 0003-2700/90/0362-1319$02.50/0
electrical potential (4-8 kV),the high (>1mL/min) flow rates used with the thermospray technique have made this method unquestionably easier to accomplish on quadrupole-based mass analyzers. Quadrupole instruments also have the advantage that they can be scanned more rapidly than sector instruments but are limited in mass range to (typically)around 3000 amu. Alternatively, lower flow rates (1-10 pL/min) and microbore columns may be used with the continuous flow fast atom bombardment (FAB) technique and can therefore more easily take advantage of the high resolution and high mass range (10000 amu) of high-performance sector instruments. Recently electrospray ionization has been introduced as a means for interfacing liquid chromatography and mass spectrometry (9). This technique has been used primarily with quadrupole analyzers and, since it produces ions with very high charge states, has extended the mass range of such instruments so as to permit the recording of molecular ions of proteins ( 1 0 , l l ) . In addition, both continuous flow FAB and electrospray have been used as interfaces for capillary zone electrophoresis (CZE), while thermospray has been used for high-performance anion exchange (HPAE) chromatography (12). The Time-of-Flight Mass Spectrometer as a n LC Detector. Time-of-flight (TOF) mass analyzers have not been employed as detectors for high-performance liquid chromatography. They have, however, the ability to record mass 0 1990 American Chemical Society
