Thermodynamic model for reversed-phase ion-pair liquid

Laura L. M. Glavina and Frederick F. Cantwell. Analytical Chemistry 1996 68 ... Tang , John J. Stranahan , and Stanley N. Deming. Analytical Chemistry...
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Anal. Chem. 1982, 5 4 , 2251-2256

Thermodynamic Model for Reversed-Phase Ion-Pair Liquid Chromatography John J. Stranahan and Stanley N. Deming” Department of Chemistry, Universi@of Houston, Houston, Tt3XaS 77004

A quantitative, four-parameter thermodynamlc model that Is not based on any “Ion-pair” hypothesis but instead assumes simple Langmulr adsorption of an added, charged surfactant at the stationary phase-nioblle phase Interface Is derlved. The model is used to describe reversed-phase llquld chromatographic behavior that has often been attributed to “ion-palr” formation In the polar moblle phase. ‘The effects of the charged surfactant on interfacialtension and ionlc lnteractlons are discussed.

Charged surfactants have been widely used as mobile phase modifiers to improve the characteristics of reversed-phase high-performance liquid chromatography when separating charged solutes (1). The uncertainty that has existed concerning the retention mechanism in this mode of liquid chromatography (2)is reflected in the variety of names given to the technique. “ion-pair chromatography” (3), “solventgenerated (dynamic) ion-exchange chromatography” (41, “Paired-Ion Chromatography” (5), “solvophobic-ion “soap chromatography” (7), “surfactant chromatography” (6), chromatography” (B), and “hetaeric chromatography” (9). Two extreme views olf the separation process have been extensively discussed. The fiist view assumes an ion-exchange mechanism ( 4 , 6, 10, 11); in this hypothesis, the lipophilic surfactant molecules adsorb onto the bonded, nonpolar stationary phase and cause the column to behave as an ion exchanger. The second view assumes the formation of an ion pair in the polar mobile phase, prior to its adsorption onto the reversed-phase surface (5,12);solvophobic theory (13)has been used to defend thici hypothesis of ion-pair formation in the mobile phase (9). Because neitheir the simple ion-exchange hypothesis nor the ion-pair hypothesis has been entirely satisfactory in interpreting experimental results, Bidlingmeye r et al. (14) recently proposed a broader “ion-interaction” model to account for available data. Although the original ion-interaction hypothesis of Bidlingmeyer et al. (14) has been able to account for many previously unexplained phenomena in “ion-pair” chromatography (e.g., the appearance of extraneous peaks in the chromatogram), it has not been sufficiently quantitative to explain other observed effects. One phenomenon that has not been completely explained is the foldover of solute capacity factor with increasing surfactant concentration (9,15,16). Although “incipient micelle formation” has recently been suggested 8s one possible explanation for the reversal of slope (16), ithe critical micelle concentrations are generally much higher than the observed foldover concentrations (17). A second phenomenon that has not been explained quantitatively is the decrease in capacity factor of uncharged solutes with increasing surfactant concentration (14,16). Knox and Hartwick (16) have suggested that the decrease is caused by “a slight lessening of the interaction with the stationary phase as this [the stationary phase] becomes progressively more polar through adsorption of the hetaeron [surfactant].”

Finally, a third quantitatively unexplained phenomenon is the large decrease in the retention of solutes having the same type of charge as the added surfactant (14, 16). The intent of this paper is to present a thermodynamic model that provides a unifying approach to the explanations of these and other phenomena in reversed-phase “ion-pair” chromatography.

THEORY The retention behavior of solutes in liquid chromatography has been treated thermodynamically by Everett (18,19) and by Locke (20). We accept the six basic assumptions that were made by Locke (20) in his work on reversed-phase liquid chromatography: either a porous microparticle or superficially porous silica support is used; surface hydroxyl groups on the silica support are all chemically bonded to an alkyl group; the organic layer is not cross-linked or polymerized; simple adsorption occurs which produces a monolayer; the eluent may contain one or more solvents; small samples are used to approximate infinite dilution. Distribution Equilibrium. Locke considers the distribution of sample and eluent components between two phases-the bulk liquid phase (mobile phase) and an “adsorbed phase” consisting of a monolayer that is adsorbed on the stationary phase (20). The basic equilibrium for this system can be described as a competition between the solvents and solutes for the available surface area of the stationary phase. For a system consisting of one solute component and one solvent component in contact with a solid adsorbent, the equilibrium is

iI1 + mMa

iIa + mM1

(1)

where I and M refer to solute and solvent molecules, respectively, subscripts 1 and a refer to the bulk liquid phase and the adsorbed phase, respectively, and i and m are coefficients of the equilibrium. The condition of equilibrium is (20) iPi1 + mPma = iPia

+ mPm1

(2) where wil and pia are the chemical potentials of component I (the solute) in the bulk liquid phase and the adsorbed phase, are the chemical potentials of respectively, and F~~ and component M (the solvent) in the adsorbed phase and the bulk liquid phase, respectively. If si and s, represent the area on the surface occupied by 1mol of I and 1mol of M, respectively, then i / m = (l/si)/(l/sm),and the equilibrium described by eq 2 can be expressed (20) Pi1 -+ si

Pma -= - + -Pia

sm

si

Pml

s,

(3)

Chemical Potentials in t h e B u l k Liquid Phase. The chemical potential of a component I in the bulk liquid phase (pill may be represented by (21)

(4) Pi1 = Pilo + RT In XilYil where pilo is the reference chemical potential of I in the pure state, R is the gas constant, T is the temperature, xilis the mole fraction of I in the bulk liquid phase, and yil is the

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ANALYTICAL CHEMISTRY, VOL. 54, NO. 13, NOVEMBER 1982

activity coefficient of I in the bulk liquid phase. The effect of pressure on the activity of I will be assumed negligible. In eq 4,yil could include effects caused by solute-solvent interactions in the bulk liquid phase (22). Chemical Potentials in the Adsorbed Phase. The chemical potential of a component I in the adsorbed phase (pia)may be represented by (19)

+ RT In x

(5) ~ - ~( u -Yuio)si ~ ~ where pk0 is the reference chemical potential of I in the pure state, xiais the mole fraction of I in the adsorbed phase, Y~~ is the activity coefficient of I in the adsorbed phase defined such that Y~~ 1 when xia 1, CJ is the interfacial tension of the system, and uio is the interfacial tension of pure I in equilibrium with the solid adsorbent. Activity Coefficient in the Adsorbed Phase. For a ternary system consisting of solvent molecules M and two kinds of solute molecules I and J in equilibrium with the adsorbed phase, an activity coefficient for solute molecule I can be defined to a first approxin the adsorbed phase (ria) imation by (23) RT In Yia = wijxj; + (wij + wmi - Wmj)XmaXja + WmiXma2 pia = piao

-

-

(6) where the w's are the energies of interaction in the adsorbed phase between the pairs of molecules indicated by subscripts and the x's are mole fractions of the indicated components in the adsorbed phase. If I and J are charged molecules, then the interactions of these solute molecules with the solvent are relatively much weaker than the strong interaction of the charged molecules with each other. Thus, eq 6 can be written

RT la

Yia

= wijxja'

+ WijXmaXja

Finally, solving eq 8 for yia and substituting into eq 12 gives -

RT Thus, two of the ways a charged surfactant J can affect the distribution coefficient of a solute I are evident in the last two terms of eq 13: by changing the interfacial tension and by ionic interaction with the solute in the adsorbed phase.

DISCUSS I ON Oppositely Charged Solute and Ion-Interaction Reagent. If it is first assumed that the effect of the charged surfactant (the ion-interaction reagent, or IIR) on interfacial tension is negligible, then eq 13 can be simplified to give

Ki = exp(Poi + P l i x j a )

where poi is a parameter that collects all of the assumed constant terms in eq 13 (including u and yil) and pli is a measure of the ionic interaction between adsorbed solute and the adsorbed IIR. Because the capacity factor ki is directly proportional to the distribution coefficient Ki (24),eq 14 can also be used to relate ki and the concentration of IIR in the bulk liquid phase; the proportionality constant between ki and Ki is included as part of the poi term when Ki is replaced by ki.

The mole fraction of IIR in the adsorbed phase, xja,can be related to the concentration of IIR in the bulk liquid phase by a Langmuir adsorption isotherm (25)

(7)

For a system consisting primarily of M and J, an increase in xjawill be accompanied by a proportional decrease in xma. Thus, without great error, xma= 1- xj, and eq 7 can be written RT In Y~~ = wijxja

Application to Chromatography. For chromatographic purposes, a distribution coefficient for a component I (Ki)can be described for a dilute solution as (20)

where cjl is the bulk liquid phase concentration of IIR a t adsorption equilibrium in moles per liter and a is a constant expressed in moles per liter. (a = 55.5 exp[AGo/RT], where AGO is the Gibbs free energy of adsorption at infinite dilution; the smaller the value of a, the more strongly the surfactant is adsorbed.) Although the Langmuir adsorption isotherm is strictly valid only under ideal surfactant behavior, it has been found to adequately describe the adsorption of many surfactants in real systems (25). Substituting eq 15 in eq 14 and replacing Ki with ki gives

ki = exP(Poi + Pli[cjl/(cjl + P j J l l where VI is the volume of the bulk liquid phase and nia/ Va and nil/ Vl are the molar concentrations of I in the adsorbed and bulk liquid phases, respectively. Equation 9 can be written in terms of mole fractions (20)

where Vlo = Vl/nl is the average molar volume of the bulk liquid phase, and Vao = Va/na is the average molar volume of the adsorbed phase. Solving eq 4 and 5 for xil and xis, respectively, and substituting into eq 10 gives

Ki = exp

(

pia

- piao

RT

pi1 - pilo

-~ RT

V1° Yil + In 7 + In + Va

(14)

Yia

At equilibrium, however, pia = pil; also, the reference states are chosen for pure I such that piao = pilo. Thus

(16)

where pli is related to the interaction energy of the charged surfactant J and the solute I and pjl is related to the energy of adsorption of the charged surfactant J. Figure 1 shows the graph of eq 16 fitted (26) to representative experimental data taken from a recent paper by Deelder et al. (15). The capacity factor of epinephrine, a positively charged solute under the conditions of the experiment, is plotted as a function of the concentration of negatively charged octanesulfonate in a 2-propanol-water (0.5:99.5) mobile phase, pH 3.0 in a phosphate buffer. Figure 1 demonstrates the good agreement between eq 16 and experimental data for a relatively short chained IIR opposite in charge to the sample. The estimate of pli is +2.25; the positive sign of pli indicates an attractive force between the positively charged solute and the negatively charged IIR. The estimae of pjl is 1.50; this corresponds to a AGO of adsorption of approximately -2.1 kcal/mol for the octanesulfonate. Similarly Charged Solute and Ion-Interaction Reagent. Figure 2 shows the graph of eq 16 fitted to representative experimental data taken from a recent paper by Kong et al. (27). The capacity factor of phenylethylamine, a positively charged solute under the conditions of the experiment, is plotted as a function of the concentration of positiuely charged octylammonium ion in a methanol-water (2080) mobile phase,

+

--

ANALYTICAL CHEMISTRY, VOL. 54, NO. 13, NOVEMBER 1982

2253

:

BENZYL RLCDHOL

EPINEPHRINE

+

Y

0

B OCTRNESULFDNRTE

12

(M I LL I MOL.RR

--

16 )

Y

2

6

OCTYLSULFRTE

R

I0

12

(M ILL I MOLAR )

Figure 3. Capacity factor of benzyl alcohol vs. concentration of octyl sulfate. Equation 18 fitted to data from ref 16. See text for discussion.

Both of these phenomena can be explained by a lowering of interfacial tension between the adsorbed phase and the bulk liquid phase caused by the adsorbed surfactant. A lowering of interfacial tension would cause a decrease in the distribution coefficient of a sample (and hence a decrease in its capacity factor) as predicted by eq 13. Thus, for systems in which the IIR is effective in reducing the interfacial tension, eq 16 would not adequately describe the retention of a solute. The Szyszkowski equation (25) describes the interfacial tension in terms of the surfactant concentration in the bulk liquid phase g = p -

-

0

1

2 3 Y OCTYLRM I NE (MILL I MOLRR )

5

6

Figure 2. Capacity factor of phenylethyiamine vs. concentration of octylamine. Equation 16 fitted to data from ref 27. See text for discussion.

pH 6.0 in 0.010 M acetate buffer. Figure 2 demonstrates the good agreement between eq 16 and experimental data for a relatively short chained IIR with the same charge as the solute, in a bulk liquid phase containing moderate amounts of organic modifiers. The estimate of pli is -3.82; the negative sign of pli indicates a repulsive force between the positively charged solute and the poaitively charged IIR. The estimate of pjl is +1.69; this corresponds to a AGO of adsorption of approximately -2.1 kcal/mol for the octylammonium ion. Interfacial Tension Effects. Experimental plots of capacity factor of a neutrcil solute vs. concentration of a surfactant often show a decrease in k as the surfactant concentration is increased (16). Similarly, experimental plots of capacity factor of a charged solute vs. concentration of an oppositely charged surfactmt often show a maximum and then an eventual lowering of capacity factor, or a folding over, as the concentration of surfactant in the eluent is increased further (9,25,16). The foldover is especially apparent when longer chained IIR's are used.

R T T ~In (1 + cjl/a)

(17)

where u is the observed interfacial tension, uo is the interfacial tension in the absence of surfactant, T~ is the maximum surface excess concentration of surfactant, cjl is the concentration of surfactant in the bulk liquid phase, and a is a constant. For a given surfactant, the constant a in the Szyszkowski equation is the same as the constant a in the Langmuir equation (28). Substituting eq 17 into eq 13 and again collecting constant terms (including ria)into the parameter poi gives where Pzi is related to the maximum surface coverage of the surfactant J and to the molar surface area of the solute component I. Figure 3 shows the graph of eq 18 fitted to data for the capacity factor of benzyl alcohol, an uncharged solute, as a function of the concentration of octyl sulfate in a methanol-water (2080) mobile phase, pH 6.00 in a phosphate buffer (16). This figure demonstrates the good agreement between the predicted effect of interfacial tension lowering and experimental data. The estimate of & is f0.195. The estimate of pjl is +0.436 which corresponds to a AGO of adsorption of approximately -2.9 kcal/mol for the octyl sulfate. Combined Interfacial Tension and Ion-Interaction Effects. Substituting both eq 15 and 17 into eq 13 and collecting constant terms gives

which accounts for both the interfacial tension and ion-interaction effects caused by a charged surfactant on a charged solute.

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ANALYTICAL CHEMISTRY, VOL. 54, NO. 13, NOVEMBER 1982

TYROSINE RHIDE ION-INTERRCTIDN EFFECT

CDHBINED EFFECT5

i

I

INTERFRCIRL TENSIDN EFFECT

I

2

3

1

4

DDDECYLSULFRTE (MILLIMDLRR)

Figure 4. Capacity factor of tyrosinamide vs. concentration of dcdecyl sulfate. Equation 19 fttted to data from ref 16. See text for discussion.

Figure 4 shows the graph of eq 19 fitted to representative experimental data taken from a recent paper by Knox and Hartwick (16). The capacity factor of tyrosinamide, a positively charged solute under the conditions of the experiment, is plotted as a function of the concentration of negatively charged dodecyl sulfate in a methanol-water (2080) mobile phase, pH 6.0 in a 0.020 M phosphate buffer. The experimental points show a maximum in capacity factor a t approximately 1.6 mM dodecyl sulfate followed by a decrease or folding over as the dodecyl sulfate concentration is further increased. Figure 4 demonstrates the good agreement between eq 19 and experimental data for a relatively long chain IIR in a bulk liquid phase containing moderate amounts of organic modifier. The estimate of &i is +3.83 and indicates an attractive force between the positively charged solute and the negatively charged IIR. The estimate of Ojl is +0.0582; this corresponds to a AGO of adsorption of approximately -4.1 kcal/mol for the dodecyl sulfate. The estimate of & is +0.132; this indicates that the interfacial tension (and thus the capacity factor) decreases with increasing surfactant concentration. Figure 5 shows the relative contributions of the interfacial tension and ion-interaction effects as they contribute to the retention of tyrosinamide shown in Figure 4. The effect of interfacial tension is to decrease the capacity factor of the solute as the surfactant concentration is increased. The effect of ion interaction is to increase the retention of the solute. Because the ion interaction is the stronger effect, the net effect is increased retention of the tyrosinamide. However, at higher surfactant concentrations, the change produced by the interfacial tension effect is greater than the change produced by the ion-interaction effect, with result that the net effect now is to decrease the capacity factor, hence the “foldover” of k a t higher concentrations of surfactant shown in Figure 4.

Figure 6 shows graphs of eq 19 fitted to additional data taken from Knox and Hartwick (16). The capacity factor of tyrosinamide is plotted as a function of octyl sulfate, decyl sulfate, and dodecyl sulfate concentrations with other eluent conditions the same as in Figure 4. The experimental points for dodecyl sulfate are the same in Figures 4 and 6. The data shown in Figure 6 were used to fit simultaneously the three separate models (eq 19 for each surfactant) with the constraint that poi(essentially ki in the absence of added surfactant) be common to the three models.

2 3 Li DODECYLWLFRTE (MILLIMDLAR)

Figure 5. Contributions of interfacial tension and ion-interaction effects to the retention of tyrosinamide shown in Figure 4: (upper curve) 3.83[cjl/(cl,+ 0.0582)]; (lower curve) -0.132 In (1 cl,/0.0582); (middle curve) 3.83[cl,/(c,, 0.0582)] - 0.132 In (1 + cll/0.0582).

+

+

R

T TYRD5lNE RHIDE

I 0

Li

E

12

16

2E3

5URfRCTflNT (MILL I MIlLflR )

Flgure 6. Capacity factor of tyrosinamide vs. concentrations of dodecyl

sulfate (upper curve), decyl sulfate (middle curve), and octyl sulfate (lower curve). Equation 19 fitted to three sets of data from ref 16. See text for discussion. The estimate of for octyl sulfate is +2.50; for decyl sulfate it is +3.04; and for dodecyl sulfate it is +3.52. The sign of this parameter indicates an attractive force between the negatively charged alkyl sulfates and the positively charged tyrosinamide. The magnitude of this parameter appears to indicate an increase in attractive force between the surfactant modifier and the solute molecule with increasing chain length. This might be due to nonionic interactions between the surfactant modifier and the solute molecule, which would tend to increase with increasing chain length (23),although alternative explanations are possible. The estimate of pjlfor octyl sulfate is +1.57; for decyl sulfate it is +0.232; and for dodecyl sulfate it is +0.060. These values correspond to AGO values of adsorption of -2.1, -3.2, and -4.0 kcal/mol, respectively. The magnitudes of these parameters indicate an expected change in the free energy of adsorption of the surfactant as the chain length is increased, and corre-

ANALYTICAL CHEMISTRY, VOL. 54, NO. 13, NOVEMBER 1982

2255

Table I. Summary of Data Sources and Parameter Estimates AGO,

fig ref 4 16 6 16 6 16 6 16 1 15 2 27 3 16 8 9

IIR C,,SO,C,,SO,ClOSO,C,SO,C,SO,C,NH,+ C,SO,C,SO,-

solute tyrosinamide tyrosinamide tyrosinamide tyrosinamide epinephrine phenylethylamine benzyl alcohol epinephrine

PH 6.0 6.0 6.0 6.0 3.0 6.0 6.0 2.55

modifier 20%methanol 20% methanol 20% methanol 20%methanol 0.5% propanol 20% methanol 20% methanol

0o i

Pli

P zi

t0.0261 t0.272a t 0.272a t0.272a -0.342 t 0.967 t 1.73 -0.913

t3.83 t3.52 t 3.04b t2.50b t 2.25b -3.82b

+0.132 +0.119 t 0.023O tO.OOOc

+3.16

+0.195 t0.438

mM

kcall mol

t0.0582 +0.0601 +0.233 +1.57 +1.50 t1.69 t0.436 +7.40

-4.05 -4.03 -3.22 -2.10 -2.13 -2.06 -2.86 -1.19

Pjl,

a Constrained to be shared by three surfactant models. Probably includes interfacial tension effects also. low because ion-interaction term includes interfacial tension effects also.

CONCENTRRTiDN

Flgure 7. Relative contributions of interfacial tension and ion-interaction effects: (upper curve) [cjl/(c,, 20)]. (lower curve) In (1 cj1/20).

+

+

spond very closely to changes in AGO of 500-700 cal per -CH2group reported by others (25). The estimate of pzi fior octyl sulfate is +O.OOO 001 07; for decyl sulfate it is +0.0224; and for dodecyl sulfate it is +0.118. The sign of this parameter indicates a lovvering of interfacial tension by each of the surfactants. The magnitude of this parameter appears to indicate that the slhorter chained surfactants are less effective in reducing the interfacial tension, but this runs counter to known surfactant behavior (25). It is probable that the decreasing magnitude of Pli and the decreasing magnitude of BZias chain length is decreased are related and are an artifact produced by a combination of the model being fitted and the span of data in each case. At relatively low concentrations of surfactant, the Szyszkowski and Langmuir terms produce plots that are very similar in shape (see Figure 7). From a mathematical point of view, a model fit to data obtained at these low concentrations might account for the observed behavior by making use of the Langmuir term, the Szyszkowski term, or both terms. Apparently in the case of decyl sulfate, and even more so in the case of octyl sulfate, the fitting process resulted in much of the interfacial tension effect appearing in the ion-interaction term as a smaller than expected Pli estimated parameter value. This utilization of the Langmuir term to express interfacial tension effects also appears to have occurred in Figure 1where the Szyszkowski term was intentionally omitted; compare the value of Pli = +2.25 for Figure 1with values higher in magnitude in the other figures. This confounding again appears to have occurred in Figcure 2 where the Szyszkowski term was

0

z0 LIB 60 HEXYLWLFATE < M 1 LL I MOLRR )

Probably

E8

Figure 8. Capacity factor of epinephrine vs. concentration of hexyl sulfate. Equation 19 fitted to data from ref 9. See text for discussion.

also intentionally omitted; compare the value of PI1 = -3.82 with values lower in magnitude in the other figures. Apparently the Szyszkowski term can only become separately estimated with confidence if data have been obtained at higher concentrations where the plots of the Szyszkowski and Langmuir terms are very different in shape (see Figures 5 and 7). This is seen in Figure 8 (9) where data have been obtained over a broad range of surfactant concentrations and the estimated p2i value is more in accord with accepted values (25). Finally, Table I summarizes the results of fitting eq 19 or its abbreviated forms (eq 16 and 18) to data taken from four widely different literature references. The highly consistent values of pli, &i, and AGO lend support to the validity of the four-parameter thermodynamic model presented in this work.

ACKNOWLEDGMENT We thank W. E. Wentworth and M. Tang for helpful discussions. LITERATURE CITED (1) Bidiingmeyer, Brian A. J. Chromatogr. Sci. 1980, 18, 525-539. (2) Kissinger, P. T. Anal. Chem. 1977, 49, 883. (3) Fransson, B.; Wahiund, K.-G.; Johansson, I. M.; Schiii, G. J. Chromatogr. 1978, 125, 327-344. (4) Kraak, J. C.; Jonker, K. M.; Huber, J. F. K. J . Chromatogr. 1977, 142, 67 1-608. ( 5 ) "Paired-Ion Chromatography, an Alternative to Ion Exchange"; Waters Associates: Milford, MA, 1975. (6) Hoffman, N. E.; Liao, J. C. Anal. Chem. 1977, 49, 2231-2234. (7) Knox, John H.; Laird, George R. J . Chromatogr. 1976, 122, 17-34. (8) Tomiinson, E.; Jefferies, T. M.; Riley, C. M. J . Chromatogr. 1978, 159, 315-350.

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(9) Horvath, Csaba; Melander, Wayne: Molnar, Imre; Molnar, Petra Anal. Chem. 1977. 49. 2295-2305. (10) Konljnendijk, A. P.; van de Venne, J. L. M. I n "Advances in Chromatography 1979"; Zlatkls, A., Ed.; Chromatography Symposium: Houston, TX 1979; pp 451-462. (11) van de Venne, J. L. M.; Hendrlkx, J. L. H. M.: Deelder, R . S.J. Chromatogr. 1978, 767, 1-16. (12) Wittmer, D. P.; Nuessle, N. 0.; Haney, W. G., Jr. Anal. Chem. 1975, 4 7 , 1422-1423. (13) Horvath, Csaba; Melander, Wayne; Molnar, Imre J. Chromatogr. 1976, 725,129-256. (14) Bidlingmeyer, B. A.; Deming, S.N.; Price, W. P., Jr.; Sachok, B.; Petrusek, M. J. Chromatogr. I97g, 786,419-434. (15) Deelder, R. S.;Linssen, H. A. J.; Kronijnendijk, A. P.; van de Venne, J. L. M. J. Chromatogr. 1979, 785,241-257. (16) Knox, John H.; Hartwick, Richard A. J. Chromatogr. 1981,204,3-21. (17) MukerJee, P.; Mysels, K. J. Natl. Stand. Ref. Data Ser. ( U S . , Natl. Bur. Stand.) 1971,No. 36. (18) Everett, D. H. Trans. Faraday SOC. 1964,6 0 , 1803-1813. (19) Everett, D. H. Trans. Faraday SOC. 1965,67, 2478-2495. (20) Locke, D. C. J. Chromatogr. Sci. 1974, 72,433-437.

(21) Lewis, G. N.; Randall, M. "Thermodynamics"; McGraw-Hili: New York, 1961; Chapter 20. (22) Elkoshi, 2 . ; Grushka, Ell. J. Phys. Chem. 1981,85,2980-2986. (23) Lucassen-Reynders, E. H. I n "Progress in Surface and Membrane Science"; Cadenhead, D. A., Danielli, J. F., Eds.; Academic Press: New York, 1976; Vol. IO, pp 253-3130, (24) Johnson, E. L.; Stevenson, R. "Basic Liquid Chromatography"; Varian: Palo Alto, CA, 1978; p 37. (25) Rosen, M. J. "Surfactants and Interfacial Phenomena"; Wiley: New York, 1978. (26) O'Neill, R. Appl. Statist. 1971,20,338-345. (27) Kong, R. C.; Sachok, 8.; Demlng, S. N. J. Chromatogr. 1960, 799, 307-316. (28) Davies, J. T.; Rideal, E. K. "Interfacial Phenomena"; Academic Press: New York, 1963; Chapter 4.

RECEIVED for review November 30, 1981. Resubmitted June 11, 1982. Accepted July 30, 1982. This work was supported in part by a grant from Chevron Research Co.

Determination of Nitrated Polycyclic Aromatic Hydrocarbons by Fused Silica Capillary Gas Chromatography/Negative Ion Chemical Ionization Mass Spectrometry Thomas Ramdahl" and Kjell Urdal Central Institute for Industrial Research, P.O. Box 350, Blindern, Oslo 3, Norway

Gas chromatography/electroncapture negatlve Ion chemical lonlzatlon mass spectrometry is shown to be a very sensltlve and Selective determlnatlon method for nltrated polycyclic aromatic hydrocarbons (PAH). The detectlon llmlt was 1 pg Injected to the GC column uslng slngle Ion monltorlng. The mass spectra are characterized by an intense molecular Ion, the base peak In all spectra. The second most abundant ions were (M 16)- and (M 30)- for mono- and dlnltro-PAH, respectlvely. The NO,- Ion ( m / z 46) Is observed in most spectra. There are few dlfferences In the mass spectra of various Isomeric compounds. The pressure and the temperature In the ion source had llttle Influence on the mass spectra and sensitivity within the llmlts tested. The method is demonstrated on a synthetic mixture containlng nltro-PAH, and on a toluene extract of a formerly commercial carbon black. The appllcatlon to urban alr partlculate extracts Is discussed.

-

-

Nitro aromatic compounds can be readily formed by reaction between nitrogen oxides (NOx) and polycyclic aromatic hydrocarbons (PAH). Exposure of PAH adsorbed to different carriers to NOx has been shown to produce nitro-substituted reaction products which are directly mutagenic in the Ames Salmonella test (1, 2). Nitro-PAH have been identified in several environmental samples like air particulate matter (3, 4 ) , diesel exhaust particles (5),and carbon black (6). These findings suggest that nitro-PAH like PAH may be ubiquitous in the environment, but in comparatively lower concentration. Environmental samples are often extremely complex. The method of capillary gas chromatography (GC) has proved to be an extremely useful tool in analyzing multicomponent mixtures (7). Capillary GC is characterized by high sensitivity, excellent resolution power, and good reproducibility. Because

of the often very low level of nitroarenes in environmental samples, such a high resolution method is needed, but no satisfactory detection method has yet been devised. Capillary GC/MS and high-resolution mass spectrometry have recently been employed to determine nitro-PAH in diesel exhaust particles (5). Another recent approach by capillary GC is by using a nitrogen-selective detector (8). Here the identification was made by comparing the retention times with those of reference compounds. As the level of nitro-PAH in environmental samples is very low, a more sensitive and selective determination method is wanted. It is known that many nitro compounds exhibit strong response to the electron capture (EC) detector. The theoretical detection limit for an EC detector is estimated to 330 am01 (9). The ionization mechanisms of the EC detector and in methane (electron capture) negative ion chemical ionization mass spectrometry (ECNICIMS) are similar. Hunt and Crow have stated a 10- to 100-fold increase in sensitivity by ECNICIMS compared to the GC/ECD (10). Recent work on polyhalogenated hydrocarbons and mycotoxins shows the great versatility of this negative ion technique (11,12). These results indicate that ECNICIMS may be a very sensitive and selective method for the determination of nitro-PAH. The purpose of this study is to investigate the application of methane negative ion chemical ionization mass spectrometry to the determination of these compounds.

EXPERIMENTAL SECTION GC/MS System. All spectra were recorded on a Finnigan Model 4021 quadrupole mass spectrometer equipped with a standard EI/CI source, Primary ionization of the CI reagent gas was accomplished with a 70-eV beam of electrons generated from a heated rhenium filament with an emission current of 0.25 mA. Methane was used as reagent gas, the ion source pressure was maintained at 0.15 torr, and source temperature was 250 "C, unless otherwise stated. The electron mulitplier voltage was 1700 V.

0003-2700/82/0354-2256$01.25/00 1982 American Chemical Society