Free Energy Correlations with Solvatochromic Red Shifts for Indicators in Aprotic Solvents Sir: In a recent communication to this journal ( I ) , experimental evidence was presented for empirical correlations between the red shifts for a set of solvatochromic indicators and the Gibbs free energy values (AGf) for the equilibrium in Equation 1.
p-FC,&OH,+
:B F= p-FC&OH.
. . :B (in cc14 s o h ) (1)
It was suggested that the observed trend can be interpreted in terms of a systematic response of the solvatochromic dye to changes in the Lewis basicity for polar aprotic species (:B). Since the completion of that project, related reports have appeared elsewhere which necessitate an extension of my preceding discussion and conclusions. Kamlet and Taft (2) have identified relationships between solvatochromic shifts for 4-nitroanilines and 4-nitrophenols interacting with a wide range of hydrogen bond donor-acceptor pairs (HBA) in equilibrium 2.
+ HBA
X-C6H4-"2
X-C&-"2
. . . HBA
(2)
They found that their 0-scale for solvent basicities derived from Equation 2 is linearly related to pKf values for the same equilibrium. However, by using published data for bathochromic dyes (i.e., Phenol blue, Nile blue A oxazone, and Brooker's dye VII) in 22 solvents, it can be verified that no general correlation exists between the transition energies ( E T ) of the indicators and the 0-values of Kamlet and Taft. It should be noted that a nonlinear increasing trend does occur between ET for Brooker's dye VI1 and 0 for five hydrogen bond donor solvents (Le., the lower alcohols and dimethylformamide). A second paper which more closely relates to my previous communication is the report of Spencer, Harner, and Penturelli ( 3 )on the determination of thermodynamic differences attributable to solvation effects upon the hydrogen bonding equilibrium in Equation 3.
+
C G H ~ O H.. SA . (CH3)2SO + CeHbOH. . . OS(CH3)z
the reference base appears to have been an excellent one since the results of Arnett, Mitchell, and Murty indicate that the alkyl sulfoxides and phosphates are among the stronger Lewis bases in hydrogen bond donor-acceptor equilibria ( 4 ) ;and likewise, for a given batkochromic indicator, the lowest ET value was recorded in DMSO among all of the aprotic solvents investigated to date ( I ) . However, Spencer, Harner, and Penturelli ( 3 )found no general trends between AH, AS, or K f for Equation 3 and the dielectric constant or other macroscopic properties of the solvent (SA). Equilibrium constant values for Equation 3 are listed in Table I along with the transition energies (ET)for Phenol blue and related solvatochromic indicators in each of the six aprotic bases as SA.The graphs in Figure 1were derived from those data but expressed as ET vs. -AGf. Slopes for the lines range from as low as 1.18 for Phenol blue derivatives to 1.65 for Phenol blue and 2.72 for Brooker's dye VII. It is not fully unexpected that the empirical linear free energy (LFE) plots in Figure 1 would include within the same function the broad range of solvent types from nonpolar to polar aprotic Lewis bases as well as the weak hydrogen bond c I L
T-
5 5.-
50.-
4 5-
2
+ SA
(3)
They concluded that detectable specific interactions by the acceptor solvent (SA)occur in all instances, including cyclohexane and carbon tetrachloride. Their choice of DMSO as
3
4
-AG
Figure 1. Correlation of the transition energy (,ET) for the indicators in Table I with the AGf for the C6H50H:DMS0complex in the solvents: chloroform; 1,2-dichIoroethane; benzene; carbon tetrachloride; carbon disulfide; cyclohexane
Table I. Summary of Data Used in Correlations with Hydrogen Bonding Donor-Acceptor Equilibria Transition Energies (kcal/mol) in the solvents Bathochromic dye Phenol blue (PB)O PB derivatives:N - (4-substituted pheny1)quinone monoimimes b z"(4) -3,5-R,R-NHz ( 5 ) -NHCH3 (3) -N(CH3)2) (2) Nile blue A oxazone (NBAO)' Brooker's dye VIId K ffor C~HSOH-DMSOcomplex in
CsH6
CS2
49.72 55.2 52.6 53.5
CHCl3
C6HI2
CICHzCHzCl
50.87
50.60
48.13
51.61
48.58
.. . . .. .
56.3 53.4 53.8 52.1 55.09 48.7 215
... ... *..
56.1 53.8 54.3 52.7 55.78
55.2 52.4 52.7 50.7 53.54 45.0 33.0
,
*
51.1
54.81 46.9 106
CC14
,
..
55.58
. ..
381
...
53.24 44.2 19.4
50.0
811
a Data from J. Figueras (5) except new exptl value in CS2. b Derivatives numbered (2-5) as designated by Figueras ( 5 ) . Results obtained by M. Davis and H. Hetzer, Anal. Chem., 38,451 (1966). Values from L. Brooker et al., J . Am. Chem. SOC.,87,2443 (1965). e Formation constants (at 20 "C) reported by J. Spencer et al. ( 3 ) .
1814
ANALYTICAL CHEMISTRY, VOL. 48, NO. 12, OCTOBER 1976
donor, CHC13. The fundamental work on Phenol blue by Figueras (5),as well as ours (6), has demonstrated that the solvent polarity contribution to the free energy transition (ET) of the dye can be derived from the two-parameter McRae equation:, and the influence of a given hydrogen bond donor (m-cresol or CHC13) conforms to a simple additivity as a perturbation energy contribution. However, in equilibrium 3, chloroform serves as an acceptor (HBA) toward phenol and shows the regular behavior in the LFE function of the aprotic solvents. Because of the high correlation coefficients between the ET values for all of the structurally related dyes in Table I, it appears that this whole series of bathochromic dyes may share a common phenomenological response to the Lewis basicity of the solvent. Finally, it should be clear that the equilibria in Equations 1 and 3 are nonequivalent chemical processes and should have dissimilar ET vs. AG, plots. The solvent displacement reaction with DMSO assigns a measurable Lewis base strength to CCld with reference to phenols as hydrogen bond donors. The available thermodynamic data for the two equilibria are not overlapping so that direct comparisons of solvent basicity between the two processes cannot be made at the present time. When a weaker Lewis base (i.e., pyridine) is substituted for DMSO in equilibrium 3, the two hydrogen bond donor sol-
vents, chloroform and 1,2-dichloroethane, compete more effectively with phenol to interact with the HBA; and the LFE relationship fails (7).Although the LFE function for reaction 3 includes only the weaker aprotic bases, similar correlations have not been observed between thermodynamic parameters related to solvent basicity and the blue shift scales proposed in the past as measures of solvent polarity in aprotic media. LITERATURE CITED (1) 0. W. Kolling, Anal. Chem., 48, 884 (1976). (2) M. Kamlet and R. W. Tafl, J. Am. Chem. SOC.,98,377 (1976). (3) J. Spencer, R. Harner, and C. Penturelli, J. fhys. Chem., 79, 2488 (1975). (4) E. Arnett, E. Mitchell, and T. Murty, J. Am. Chem. SOC.,96, 3875 (1974). (5) J. Figueras, J. Am. Chem. SOC.,93, 3255 (1971). (6) 0.Kolling and J. Goodnight, Anal. Chem., 45, 160 (1973). (7) J. N. Spencer et al., J. fhys. Chem., 80, 81 1 (1976).
Orland W. Kolling Chemistry Department Southwestern College Winfield, Kansas 67156
RECEIVEDfor review May 17, 1976. Accepted July 14, 1976.
Application of a Vidicon Tube as a Multiwavelength Detector for Liquid Chromatography Sir: In this report, we describe preliminary data for the application of a silicon target vidicon as a multiwavelength detector for liquid chromatography (LC). Most photometers currently used as uv detectors in LC are limited to one or two wavelengths. For these systems, the selectivity of any analysis is limited to that which can be obtained via the chromatographic process, and the analytical wavelength is seldom optimized fior all components to be detected. The vidicon detector, when used with appropriate optics, is capable of monitoring many wavelength resolution elements simultaneously ( I - 3 ) , and as such should serve as a very versatile detector for liquid chromatography ( 4 , 5 ) . The vidicon spectrometer system used in this work has been described in detail previously (1,Z). The liquid chromatograph was assembled from commercially available components and is similar to that described earlier (6) except that the flow cell from a stopped-flow mixing system (Aminco-Morrow Model B30-68109; American Instrument Co., Silver Spring, Md. 20910) modified by replacing the Teflon input chamber by a Teflon exhaust chamber, was used as the observation cell. No attempt was made to optimize the performance of the chromatographic system. Spectral data were collected, processed, and dispLayed by an on-line computer. An aqueous mixture of uric acid, theophylline, and phenobarbital was selected as an illustrative example (7). Uric acid is an important biological compound and phenobarbital is often included in theophylline preparations. These compounds were eluted from an anion exchange column with an ammonia buffer at pH 10. The performance of the column had been degraded by previous operation. Spectra in the range of 225 to 450 nm were recorded every 10 s. Figure 1represents the spectra after 220,330, and 350 s for uric acid, theophylline, and phenobarbital, respectively, added to and eluted from the column separately. Figure 2 represents the absorbance measured at 10-s intervals at the absorbance maximum for each of the respective components. The spectra show that there is
no wavelength within this range at which phenobarbital is free of interference from the other two components and, similarly, that there is no wavelength at which theophylline is free of interference from uric acid. The elution peaks show that uric acid is reasonably well separated from the other two components, but that the two drugs are poorly separated. These observations will be useful in interpreting data presented below. Figures 3A and 3B represent spectra at selected times during the elution of a mixture (0.83 pg each) of the compo-
b(nm> Figure 1.
Absorption spectra of uric acid, theophylline, and pheno-
barbital. All compounds were added to and eluted from the cation exchange column separately. Eluting reagent was 0.1 M ammonia buffer at pH 10. (a) Uric acid (5fig) at 220 s, (b)Theophylline (2 fig) at 330 s,and (c) Phenobarbital (2 fig) at 350 s
ANALYTICAL CHEMISTRY, VOL. 48, NO. 12, OCTOBER 1976
1815
