dard error of estimate for the correlation was 19.6 pg/100 ml. These studies demonstrate a close correlation between the G R A method and the reference spectrophotometric procedure. The only interferences with the G R A method resulted when the iron concentration was less than 50 pg/lOO ml and/or the serum sample was more viscous than normal; in such cases, the problem was solved by diluting the serum sample 1 : l with deionized water. From the results of the study, the G R A system In atomic absorption spectrometry is an extremely promising system providing rapid, precise, and accurate means for atomizing and measuring iron in serum with no prior sample treatment.
slope. Therefore, an aqueous standard analytical curve was used for all correlation studies. Known amounts of iron were added to two different pooled serum samples, and the results demonstrating quantitative recoveries over a wide range of iron concentration are given in Table I. A study was carried out comparing serum iron levels from 22 patients at the Shands Teaching Hospital (Gainesville, Fla. 32601) by the G R A and the automated spectrophotometric procedure. The results of the study are given in Figure 3. The equation of the regression line was y = m x ( y = spectrophotometric method, x = G R A atomic absorption spectrometry). For the G R A method, a serum iron mean of 93.3 pg/lOO ml was found and a serum iron range from 32 to 198 pg/lOO ml. For the automated spectrophotometric method, a serum iron mean of 91.8 pg/lOO ml and a serum iron range of 22 t o 215 pg/100 ml were obtained. The stan-
RECEIVED for review June 9, 1972. Accepted August 7, 1972. This work was partially supported by AF-AFOSR-70-1880-H and NIH-5RO1 GM1720-03.
Trichloride Ion Formation Constant in Acetonitrile Solutions M . C . Giordano, V . A . Macagno, a n d L . E . Sereno Departamento de Fisicoquimica, Facultad de Ciencias Quimicas, Unicersidad Nacional de Cdrdoba, Co'rdoba, Argentina
THE ELECTROCHEMICAL BEHAVIOR of X,K- (X = I, Br, C1) systems at platinum electrodes in acetonitrile solutions has been the subject of previous inevstigations (1-3). The experimental results were explained by the existence of the equilibrium Xz X- e X3-, where the relative magnitude of the equilibrium constant would determine the differences among them. The stability of the trihalide complexes in various solvents has been estimated. Voltammetry had shown that the species Xs- are produced in acetonitrile ( 4 ) ; subsequently, the following formation constants were calculated from a set of half-wave potentials ( 5 ) : triiodide, 106.6; tribromide, 10'; and trichloride, lolo. Recently, these values were reviewed ( 6 ) and it was concluded that the stability constant sequence would be analogous to that found in water:
+
Figure 1. Apparatus side view experimental method is based on the determination of the distribution ratio of a volatile solute between two miscible solvents (8).
Similar results were found in sulfolane (7) and nitromethane (6), which closely resemble acetonitrile physicochemical properties. The aim of the work described in this paper was to obtain the value of the stability constant of the C18- formation in acetonitrile under the same conditions used in the electrochemical study of the chlorine-chloride electrode (3). The
EXPERIMENTAL
(1) V. A. Macagno, M. C. Giordano, and A. J. Arvia, Electrochim. Acta, 14, 335 (1969). (2) T. Iwasita and M. C . Giordano, ibid., p 1045. (3) L. Sereno, V. A. Macagno, and M. C. Giordano, ibid., 17,
561 (1972). (4) I. M. Kolthoff and J. F. Coetzee, J. Amer. Cliem. Soc., 79, 1852 (1957). ( 5 ) I. V. Nelson and R. T. Iwamoto, J. Electroairal. Cltem., 7, 218 (1964). (6) J. C. Marchon, C . R. Acad. Sei., Ser. C., 267,1123 (1968). (7) R. L. Benoit, M. Guay, and J. Desbarres, Cau. J. CIiem., 46, 1261 (1968).
Apparatus. The device is a closed borosilicate glass apparatus containing the two separated solutions under investigation which permits by mere rotation a continuous circulation of vapor through the two solutions. The essential features (Figure 1) are two cylindrical containers of about 200ml volume each, provided with ground stoppers, and connected at the top by glass tubes. The stoppers are replaced at the end of the run by delivery tubes which are ground to fit the same openings. Reagents. Chlorine gas, LiCI, and LiC104 in acetonitrile solutions were used. Chemical and solvent purification. as well as solution preparation, have been described in a previous paper (3). (8) G. Jones and B. B. Kaplan, J. Amer. Chem. SOC.,50, 1600 (1928).
ANALYTICAL CHEMISTRY, VOL. 45, NO, 1, JANUARY 1973
205
Table I. Analytical Concentrations at Different Stoichiometric Ionic Strengths a X 103M bx 1 0 3 ~ c x1 0 3 ~ 1.06 1.67 6.57 1.76 2.78 6.70 3.43 5.54 7.83 3.49 6.33 11.00 5.04 7.50 6.47 5.15 10.40 14.50 5.07 8.56 10.60 6.49 11.60 12.70 9.22 14.10 9.36 Without LiC104, p = 9.5 X 10-3M 5.21 8.48 15.20 6.26 10.10 15.00 14.70 22.40 17.90 18.50 27.10 17.20 With LiClO4, p = 38.8 X 10-3M 1.38 2.07 17.40 1.38 2.22 18.30 3.90 6.03 16.70 5.69 8.74 16.80 16.30 21.30 12.90 With LiC104, p = 49.8 X l(Y3M 1.83 2.56 17.70 3.65 4.89 17.10 5.22 6.74 16.60 6.12 8.26 17.60 6.46 8.82 16.00 10.50 13.60 16.00 With LiC104, p = 75.5 X 10e3M 5.34 6.38 14.90 6.87 8.00 12.40 8.24 9.96 16.20 18.10 21.50 16.00 With LiC104, p = 126.8 X 10-3M 5.03 5.34 9.73 5.39 5.68 9.73 8.73 9.13 9.92 12.40 13.30 10.50 16.90 18.00 10.50 With LiC104, p = 364.5 X 10-3M
Procedure. About 100 ml of a LiCl solution is dropped into one container (i), while the other ( j ) is half filled with a Clz solution. I n order to analyze the effect of ionic strength on the equilibrium constant under the conditions used in the electrochemical experiments (3),LiC104was added to both solutions. The stoichiometric ionic strength was changed from 9.5 X to 3.6 X 10-lM. Samples were taken at different periods of time until constant concentration is reached. The time required t o attain the equilibrium was about 6 hours. The CI2concentration was determined by adding K I to the solution. The free Iz was titrated by amperonietry with standard Na2S203. The chloride concentration was determined potentiometrically with a standard solution of AgN03 from a sample dried in vacuum to eliminate the acetonitrile and Cls, and dissolved in water. All the experiments were performed at 0 “C + 0.1 “C and the temperature was kept constant by means of a Lauda cryostat, U.K. 80 D.L.
1
J
/
50
/
100
;ci ]
a0
60
-1
x [CI
x
w 6 ma/2/-2
Figure 2. Trichloride equilibrium constants at different stoichiometric ionic strengths
Table 11. Dependence of Apparent Equilibrium Constant on p p, lo3 (molil.) K’ (l./mol) 9.5 107 i 12 38.8 53i 6 49.8 381 4 75.5 231 3 126.8 14+ 2 364.5 7 1 1
pressure of the solution in (i). Therefore, it is supposed that the C12 concentration a t equilibrium is the same in both containers. The chlorine equilibrium concentration, (a), is the value obtained in the ( j ) container, while the trichloride equilibrium concentration can be taken as the total chlorine concentration in the (i) container, (b), minus the chlorine concentration in equilibrium, (a), (9). The chloride equilibrium concentration will be the total chloride concentration, (c), minus the chloride as CIS-, ( b - a). The values obtained are compiled in Table I at different stoichiometric ionic strengths, p. From this information, the C1- $ apparent equilibrium constant of the reaction Cly CIS- can be calculated. Thus,
+
The mean value of K’ at constant ionic strength was found by fitting the CIS- equilibrium concentration against the Cla equilibrium concentration times the C1- equilibrium concentration by the least squares method, as shown in Figure 2. The slopes obtained are summarized in Table 11. It should be noted that the mean value of K’ increases as ii decreases. In order to find a functional relationship between K’ and p, the following equation must be considered: log K’
=
log K, - log
+
-t log yc1? (2)
7 ~ 1 ~ -log ~ c i -
where K, is the thermodynamic equilibrium constant and the
RESULTS AND DISCUSSION
After the equilibrium had been reached, the chlorine vapor pressure of the solution in ( j ) had to be the same as the vapor 206
(9) E. A. Guggenheim and J. E. Prue, “Physic0 Chemical Calculations,” Interscience Publishers, New York; North-Holland Publishing Company, Amsterdam, 1955.
ANALYTICAL CHEMISTRY, VOL. 45, NO. 1, JANUARY 1973
yi’s are the activity coefficients. This implies knowledge of the y’s dependence on 1.1. However, in this case it is difficult to speculate with the different functions proposed (IO), since the existence of three electrolytes, LiCI, LiC104, and LiCI3, in the solution complicates the application of the ionic interaction laws, and since the mean collision diameter of the ions does not have a clear physical picture. Furthermore, in solvents with a low dielectric constant, the majority of the electrolytes are associated (11, 12) and the “salting” effect of ions o n the chlorine molecule is very hard to evaluate (13). Although the y’s can not be precisely determined, we could still estimate K, from K‘ a t p = 0. Nevertheless, the error in the experimental method employed does not allow any
(10) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” 2nd ed., Butterworths S.P., London, 1959. (11) S . Minc and L. Werblan, Electrochim. Acta, 7, 257 (1962). (12) R. Fernandez-Prini and J. E. Prue, Trans. Faraday Soc., 62, 1257(1966). (13) H. S. Harned and B. B. Owen, “The Physical Chemistry of Electrolytic Solutions,” 3rd ed., Reinhold Publishing Corp., New York, N.Y. 1958.
decrease in the stoichiometric ionic strength without introducing noticeable uncertainty in the extrapolation.
CONCLUSIONS Results show that the higher value of the constant a t the lower experimental ionic strength is about lo2. This agrees with the trihalide stability series already postulated, 1 3 - > Br3- >> Cls- (6). Considering the ionic strength used in the electrochemical experiments ( 3 ) , it is perfectly justified that the chloride oxidation o r chlorine reduction had shown only one wave t o complete the reaction 2 C1- S Clz
+ 2e
(3) The existence of a low stability constant for C13-, as compared with Br3- and 13-, makes the CI3- concentration too low to be involved in the electrode process (14).
RECEIVED for review May 1, 1972. Accepted July 17, 1972. This work has been partially financed by the Consejo Nacional de Investigaciones Cientificas y TCcnicas of Argentina. (14) R. Guidelli and G. Piccardi, Electrochim. Acta, 12, 1085
(1967).
Determination of 17-Hydroxycorticosteroids with p - HydrazinobenzenesuIfonic Acid-P hosphoric Acid Ajit Sanghvi, Luigi Taddeini,’ and Carl Wight Depurtment of Putlzologj., Dicision of Clinical Chemistry, Uni6ersit.v of Pittsburgh, 201 DeSoto Street, Pittsburgh, Pa. 15213
THEPORTER-SILBER REACTION ( I ) has been widely used for the detection and determination of 17-hydroxycorticosteroids in biological fluids to provide an assessment of adrenal cortical activity. In this reaction, steroids which possess a 17,21-dihydroxy-20-keto group a t C-17 react with phenylhydrazinesulfuric acid-absolute ethanol to form 17-deoxy-20-keto steroid 21-phenylhydrazones which show characteristic absorption a t 410 nm. Although numerous modifications (2-#) of the procedure have appeared since it was first proposed by Porter and Silber in 1950 (9,reaction with phenylhydrazinesulfuric acid has remained central to these efforts. In a brief communication in 1956, Hadd and Perloff (6) indicated that substitution of p-hydrazinobenzenesulfonic acid (HBS) and phosphoric acid for the conventional Porter-Silber reagent decreased technical problems with high blanks and turbidity. Present address, Department of Medicine. University of Minnesota Hospitals, Minneapolis, Minn. 55414. (1) R. H. Silber and C. C. Porter, J . Bid. Chem., 210,923 (1954). (2) D. H. Nelson and L. T. Samuels, J . Cliu. Etidocriuol. Metah., 12, 519 (1951). ( 3 ) W. J. Reddy, D. Jenkins. and G. W. Thorn, Metabolism, 1, 511 (1952). (4) E. M. Glenn and D. H. Nelson, J. C l h Endocri/io/. Metab., 13. 91 I (1953). (5) C . C. Porter and R. H. Silber, J . Biol. Chem., 185, 201 (1950). (6) H. E. Hadd and W. H. Perloff, Abstracts, 126th Meeting of the American Chemical Society, Division of Biol. Chem., 1956, No. 77.
The use of HBS was also reported to increase the sensitivity of the procedure. To our knowledge, however, a detailed study evaluating and optimizing the various conditions of the reaction using HBS has never been published. In this communication, we present the optimum conditions and supporting data for the use of HBS-phosphoric acid reagent system in the analysis of 17-hydroxycorticosteroids.
EXPERIMENTAL Apparatus. Spectra were obtained with a Beckman Acta V double-beam recording spectrophotometer with a wavelength reproducibility within =t5 x 10-2 nm and a wavelength accuracy of f1 X 10-1 nm. The spectra were recorded at ambient temperature (-25 “C) using 1.O-cm optical pathlength matched fused silica cells (Precision Cells, Inc., 221 Park Avenue, Hicksville, N.Y. 11801). Melting points were determined with a Thomas-Hoover capillary melting point apparatus (A. H. Thomas and Co., Philadelphia, Pa.). Reagents. Analytical grade reagents were used throughout. Steroids used in this work were obtained either from Sigma or from Steraloids, Inc. Ab-Pregnenolone was a gift of R. N. Dexter of Indiana University, School of Medicine. p-Hydrazinobenzenesulfonic acid was purchased from Eastman Organic Chemicals. Celite was from Johns-Manville. All other reagents were from Fisher Scientific Co. Procedure. HBS was recrystallized to a constant mp 283283.5 “C [lit. 286 “C (7)] prior to use by the following pro(7) “Handbook of Chemistry and Physics,” R. C. Weast. Ed., 52nd ed., The Chemical Rubber Co., Cleveland, Ohio, 1972.
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