expanded scale could be used to increase sensitivity and therefore only 0.02 mg would be required. In many cases, the melting point of the 3,5-dinitrobenzoate ester would be insufficient to identify it. This is readily demonstrated with the n-propyl and sec-butyl esters, whose melting points differ by only one degree. In addition, the boiling points of the parent alcohols would be of little value in identification because they differ by only two degrees. However, as shown in Table I, the second-order hydrolytic rate constants for these two compounds differ by a factor of eight, making characterization quite simple. The sensitivity of a few 3,5-dinitrobenzoate esters toward alkaline hydrolysis is shown in Figure 2 (the plots included in this representation have been normalized for comparative purposes). Through a combination of rate constant and
melting point of the 3,5-dinitrobenzoate derivative, any alcohol in Table I can be identified. Other classes of organic compounds should lend themselves to characterization through such kinetic procedures. ACKNOWLEDGMENT The author expresses his gratitude to Kenneth A. Connors for suggesting the possible utility of precise kinetic measurements in qualitative analysis and for his help throughout the study. He also thanks Jessie Armstrong for her technical assistance. RECEIVED for review March 24, 1967. Accepted May 26, 1967.
Application of Atomic Absorption Spectrometry to Extraction of Mercuric Iodide into Carbon Tetrachloride Michael D. Morris and Lee R. Whitlock’ Department of Chemistry, The Pennsylcania State Unifiersity, Unicersity Park, Pa.
INTHE STUDY of any solvent extraction system, it is necessary t o determine the concentrations of one or more of the partitioning species in one or both phases. Such analyses are needed t o determine the equilibrium constants for the system. For this purpose, many standard analytical techniques have been employed, but only spectrophotometry and radiotracer monitoring are widely used ( I ) . The present communication describes the application of atomic absorption spectrometry to extraction studies. The virtues and limitations of this technique are well known to analytical chemists, and have been recently reviewed ( 2 , 3). Of particular importance in extraction studies are the excellent (almost perfect) selectivity of atomic absorption, its applicability to the determination of metals in almost any aqueous or nonaqueous medium, its experimental simplicity, and its good sensitivity. A major disadvantage of atomic absorption spectrometry is the sensitivity of the method to matrix effects (2, 4). It is necessary that calibration curves be made from standards whose compositions resemble as closely as possible those of the solutions to be analyzed. To test the applicability of atomic absorption spectrometry to extraction studies, we have chosen the relatively simple system, extraction of mercuric iodide from 0.5M sodium perchlo-
16802
rate into carbon tetrachloride. Marcus ( 5 , 6 ) has made extensive studies of the extraction of mercuric iodide from 0.5M sodium perchlorate into benzene and has verified that the only extracted species is HgI2. [For this system, the partition coefficient is 46 f 1 (5, 6).] Because mercuric ion forms several iodo complexes, the distribution coefficient depends on the mole ratio 1:Hg. If this mole ratio is restricted to 2 :1 or less, then the formation of Hg13- and HgIIP2may be neglected. In this case the relation between the distribution coefficient, D = (Hg)o/(Hg),Q, the partition coefficient, P = (HgIz)o/(HgIz),, and the successive formation constants, KI and K2 for HgI+ and HgIz is given by Equation 1 (6). [W(l K1 - Kz D[(1 -
+ D)P - 2D(1 + P)]Z (1
+ 0 )P + D(1 + P)1
(1)
For Equation 1 to hold, the average ligand number of mercuric ion in the aqueous solution, A, must be less than or equal t o 2. For the medium of this study, 0.5M NaC1O4 a t 25”C, Marcus (6) has found Kl = 8.9 X 10l2and Kz = 7.4 X 10’0. Because these constants are so large, virtually all (99+ %) of the iodide in solution is bound to mercuric ions and A is equal to the stoichiometric ratio, I :Hg. EXPERIMENTAL
Present address, Department of Chemistry, University of Massachusetts, Amherst, Mass. (1) V. V. Fomin, “Chemistry of Extraction Processes,” National
Science Foundation (U.S.A.) and the Israel Program for Scientific Translations, Jerusalem, 1962, pp. 28-50. (2) J. W. Robinson, “Atomic Absorption Spectroscopy,” Marcel Dekker, New York, 1966. (3) R. Lockyer, “Advances in Analytical Chemistry and Instrumentation,” C. N. Reilley, Ed., Vol. 3, Wiley, New York, 1964, pp. 1-29. (4) R. Hermann and C. T. J. Alkemade, “Chemical Analysis by Flame Photometry,” Wiley, New York, 1963, pp. 278-372. 1 180
ANALYTICAL CHEMISTRY
Mercuric nitrate, standardized against EDTA, was used as the source of the mercuric ion. Sodium perchlorate was prepared by neutralizing sodium carbonate with perchloric acid. All reagents were ACS reagent grade. Distilled water was used t o prepare all solutions. All measurements were carried out a t room temperature, 3” C. 25” The initial mercury concentration in the aqueous phase was kept constant a t 2.49 X 10-4M. Various amounts of po-
*
( 5 ) Y. Marcus, Acta. Chem. Scand., ll, 329 (1957). (6) Ibid., p. 599.
0.10
-l -
i
l
l
l
l
l
l
l
l
l
l
i
l
l
l
l
I
1
0.08 -
z
0.06
0 C
e 2
s
0.04
0.02
-
r
-
-
-
.-,
2.8 2.4
2.0
L
-
-
* P i
-
I
l
l
l
l
l
l
l
l
l
l
l
l
I
I
I
1.6
1.2
0,
1
3 0.8 0.4
1.0
2.0
3.0
4.0
5.0
6.0
7.0
8.0
9.0
([-)in, Mx104
tassium iodide were used to obtain the desired mole ratios. The aqueous phase in each extraction experiment contained 0.5M sodium perchlorate to maintain constant ionic strength. Equilibration of 6 ml each of aqueous mercuric iodide solution and carbon tetrachloride was effected by shaking the two liquids in a stoppered 15-ml centrifuge tube for several minutes. One or two hours' settling time was allowed before concentration measurements were made. In a few instances, settling times of 15-17 hours were employed, with no noticeable effect on the extraction. The quantity of mercury in each phase wa? determined by atomic absorption spectrometry (Hg 2536-A line) using a Perkin-Elmer Model 303 atomic absorption spectrophotometer equipped with a 10-cm slot burner and a mercury vapor lamp. It was necessary to use a very lean flame when aspirating carbon tetrachloride in order to minimize the amount of unburned carbon in the flame. Calibration curves were prepared for the aqueous phase by measuring the absorbance of known solutions of mercuric nitrate in 0.5M sodium perchlorate saturated with carbon tetrachloride. Organic phase calibration curves were prepared for solutions of known amounts of mercuric iodide dissolved in carbon tetrachloride which had been previously equilibrated with 0.5M sodium perchlorate. In all cases, the sum of the aqueous phase and carbon tetrachloride phase mercury concentrations, measured after equilibration of equal volumes, was within 1 relative to the known concentration of mercury in the aqueous phase prior to extraction.
Figure 2. Distribution of 2.49 X lO-,M Hg between OSM NaC104 and CC14 as a function of initial I- concentration Curve 1. [HglBq Curve 2. [HgI0
Table I. Distribution of HgIz between 0.5M NaC104 and CCl, at 25" i. 3" Ca Initial (I-)Bq n D P 1.57 x 10-4 1.97 2.36 2.86 3.15 3.94 Av.: P = 16.4 f 0.9. a
0.389 0.587 0.85 1.08 1.78 3.17
14.8 15.9 16.4 18.3 16.0 17.1
The initial aqueous phase mercury concentration is 2.49 x
10-4~.
RESULTS
The presence of iodide in the solution does not affect the absorbance of a given concentration of mercury (aqueous phase) until the 1:Hg ratio approaches 3 : l as shown in Figure 1. The decrease in absorbance at high iodide concentration reflects formation of a difficultly dissociable mercury-iodide complex, which persists in the flame. However, in the region of interest for this study, 1:Hg 5 2:1, the mercury absorbance does not depend on the iodide concentration. The results of the extraction experiments are summarized in Figure 2. At very high initial iodide concentrations (not shown), the extraction becomes less efficient, because much of the mercury exists as Hg13- and Hg14-*, which do not extract. Table I shows the calculated values of the partition coefficient, P,calculated from Equation 1, using the data of Figure 2 and Marcus' values (6) of Kl and K z .
0.63 0.79 0.95 1.06 1.26 1.58
.
The average value of P,16.4 =t 0.9 is somewhat higher than the value of D obtained at ii = 2, where D = 11.4 + 0.3. The value of P obtained from Equation 1 is more reliable. At ii = 2, the aqueous phase mercury concentration is small, about 2 X lO-SM, and the calibration curve is less reliable than at higher concentrations. Equation 1 is employed on solutions where rf is between 0.6 and 1.6 and where more accurate measurements of mercury concentration can be made. The relative error in the determination of P, + 5 . 5 % , is greater than that expected from the accumulation of errors in the atomic absorption measurements. The empirical distribution coefficient, D,is given by the ratio of two mercury concentrations, each determined by atomic absorption spectrometry. Because each of these measurements has a relative error of + 1 %, D is known with a relative error of =t2 %. The manipulation of the data to obtain the partition coefficient, P, increases the relative error to about +4%. That the observed relative error (f5.5 %) is greater than this limit is probably a result of the crude temperature control used in this experiment. RECEIVED for review February 16, 1967. Accepted May 19, 1967. Work supported in part by Public Health Service Grant 1 R 0 1 D H 00101-01.
VOL. 39, NO. 10, AUGUST 1967
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