should be based on dielectric constant change of the gas within the cell (IO). However, some simple calculations indicate that this detector is much more sensitive than such a detector should be. In addition the detector gives abnormally high response to Oz, Nz, and Hz when compared to C02 and CO. The use of an argon carrier gas resulted in a decrease in sensitivity for all permanent gases. Studies concerning the mechanism of response of this detector are now in progress. As mentioned in this manuscript, a number of studies are still required before this detector can be considered to compete favorably with other high sensitivity gas chromatographic detection systems (2-9). Its promise, however, as an extremely sensitive, nonselective detector has been shown.
In addition, the detector s p t e m is relatively simple in design, inexpensive to construct, reproducible in response over a considerable period of time, and free from annoying interference due to small changes in carrier gas flow rate. ACKNOWLEDGMENT
The authors thank Marvin Moss for assistance in construction of the oscillator and niixer circuits. LITERATURE CITED
(1) Clapp, J. K., Proc. I R E , p. 356, March 1948. (2) Hund, A , , “PhenV,mena In High Frequency Systems, McGraw-Hill, New York, 1936. (3) Johannson, G., ANAL. CHEM. 34, 914 (1962).
( 4 ) Karman, A., Bowman, R. L., Ann. S. Y.Acad. Sci. 72, 714 (1959). (5) Karman, A., Bowman, R. L., Gas Chromatog. Intern. Symp., 2nd, East
Lansing, Xich., 1959. (6) Lovelock, J. E., ASAL. CHEM.33, 162 (1961). (7) Lovelock, J. E., Sature 185,49 (1960). (8) Noble, F. W., Able, K., Cook, P. W., ANAL.CHEY.36, 1421 (1964). (9) Smith, Ir. N., Fidiam, J. F., 15th Annual Pittsburgh Conf. on Anal. Chem. and Appl. Spectroscopy, Pittsburgh, Pa., 1964. (10) Winefordner, J. D., Steinbrecher, D., Lear, W. E., ANAL. CHEM.33, 515 (1961). J. D. WINEFORDNER H. P. WILLIAMS Department of Chemistry University of Florida Gainesville, Fla. C. D. MILLER American Instrument Co., Inc. Silver Spring, hid.
Reduction of Disulfides with Tributylphosphine SIR: We wish to report the use of tributylphosphine as a reducing agent for both alkyl disulfides and aromatic disulfides. Aromatic disulfides are quantitatively reduced to thiols by triphenylphosphine at room temperature in aqueous methanol containing perchloric acid (2), Alkyl disulfides are reduced to only a small extent (20 to 30%) in 4 hours a t room temperature by the same procedure and are reduced to the extent of 80 to 90010 after refluxing in aqueous methanol for several hours ( 3 ) . Some alkyl disulfides (6) and benzyl disulfide (6,6) do not react with triphenylphosphine in benzene a t elevated temperatures. Tributylphosphine, as would be predicted from its more basic and nucleophilic nature ( I ) , appears to be much more effective than triphenylphosphine in reducing alkyl disulfides to thiols. Essentially quantitative reduction was found for benzyl, butyl, and propyl
Table 1.
Compound Bis(o-nitropheny1)disulfide 4,4’-Dithiodianiline Phenvl disulfide
disulfides after approximately 60 minutes a t room temperature. A number of aromatic disulfides were also quantitatively reduced in the same manner. Apparently the reaction of tributylphosphine with disulfides has not been previously reported. EXPERIMENTAL
Apparatus. T h e polarograph and amperometric titrator were described previously (2). Reagents and Solutions. Tributylphosphine was obtained from 11 & T Chemicals and was used without further purification. Benzyl disulfide, butyl disulfide, and propyl disulfide were obtained from Eastman Organic Chemicals. .Ill other compounds and solutions were the same as reported earlier (2). Procedure. h weighed amount of disulfide (0.05 mmole) a n d tributylphosphine (0.10 mmole) was dissolved in a small volume (1-2 ml.) of acetone
Reduction of Disulfides with Tributylphosphine
Time,a rnin. 20 20
5
RSSR, * mg.
Reduction, RSH, mg.
13.6 13.9 9.5 8.9 10.4 9.7 18.2 18.2 13.3 13.5 11.3 this period of time to react
13.2 13.6 9.0 8.9 10.4 9.5 17.1
7%
97 98 95 100 100 98 94
30 5 p-Tolyl disulfide 15 Benzyl disulfide 30 60 Butyl disulfide 30 60 Propyl disulfide 60 Elemental sulfur added after with remaining phosphine and stop reaction. b Sufficient tributylphosphine added so ratio of phosphine to disulfide was approximately 2 .
164
ANALYTICAL CHEMISTRY
and diluted to approximately 100 ml. with 10% aqueous methanol. Alternatively, aliquots of stock solutions of the disulfide and phosphine in methanol were taken. Solutions were allowed to age for periods of up to 1 hour. Before titration 1-2 ml. of a 0.10M sulfur solution in benzene, equivalent to the phosphine introduced, were added to remove excess phosphine. Then 1 ml. of 70% HCIOa was added and the thiol titrated amperometrically with 0.01M AgNO3. RESULTS A N D DISCUSSION
Essentially quantitative reduction of the disulfides investigated was found within 5 to 60 minutes after reacting these compounds with tributylphosphine. This is shown in Table I. The thiols were detected by their polarographic oxidation wave, which for the aromatic compounds occurs a t about -0.05 volt vs. S.C.E. and for the alkyl thiols is found in the region of +0.05 volt us. S.C.E. in 0.1M HC10,. Most of these substances also exhibit characteristic odors which can be recognized. The fate of the tributylphosphine has not been established but the reaction is presumed to proceed as shown in Equation 1, analogous to the triphenylphosphine reaction. RSSR
+ (CdHg)3P + H20 = 2RSH
+ (C39)d’O
(1)
hpproximately 100% excess of the theoretical amount of tributylphosphine was used in almost all cases to compensate for loss of this substance caused by oxidation by dissolved oxygen. Contrary to results found with triphenylphosphine (Z), very little reduction occurred with tributylphosphine
in solutions 0.1JI in HClO,. The reduction with tributylphosphine qeerns to be efficient in solutions containing no electrolyte r r with acetic acid and ammonium acetate or sodium hydroside present. Refluxing did not appear to aid the redurtion. Elemental sulfur was added prior to titration to remove any excess phosphine which would cauze some difficulty in the location of the amperornetric end point (3). hlso, the end point was sharper if the solution was made 0.1Jf in HClOl prior to titration and after the reduction had occurred. The presence of elemental sulfur with the thiol causes an apparently complex
reaction to occur during addition of silver nitrate, but quantitative results are still possible (4). The solutions turn yellow or orange during titration and a dark precipitate, presumably silver sulfide, is produced. ACKNOWLEDGMENT
The authors thank 11 & T Chemicals for supplying the tributylphosphine. LITERATURE CITED
(1) Henderson, W. A., Jr., Buckler, S. A., J . Am. Chem. Soc. 82, 5794 (1960).
( 2 ) Humohrev. It. E.. Hawkins. .T. %I --, ANAL.'CHE;;. 36, 1612 (1964). ( 3 ) Humphrey, It. E., Webb, R. M., I
Sam Hbusthn State College, unpublished work, 1964. ( 4 ) Karchmer, J. H., ANAL. CHEM.29,
425 (1957). ( 5 ) Moore, C. G., Trego, B. R., Tetrahedron 18, 205 (1962). ( 6 ) Schonberg, A , , Barakat, M. Z., J . Chem. SOC.1949, p. 892.
R A Y E. HUMPHREY JAMESL. POTTER
Department of Chemistry Sam Houston State College Huntsville, Texas SUPPORT of the Robert A. Welch Foundation of Houston, Texas, is gratefully acknowledged.
Investigation of 1,3-Diaminopropanol-2 Interactions with Copper and Nickel Ions SIR: During an investigation of enzyme-chelate interactions, we were concerned with the chelating properties of 1,3-diaminopropanol-2 (AOH) with Cu+2 and Ni+2. Isolation of a 2-1 AOH-CU+~complex has been reported (3) and Gonick ( 7 ) found that the C u f Z chelate reached a maximum li value of 1.5 on the basis of the following equilibrium : Cu+2
+ ;iOH
~
Cu.40Hf2
Later, Duertsch, Fernelius, and Block ( I , 2) showed that if the ethanolic hydrogen is released during titration with Cu +2, analysis of the interaction yielded steady values for the stability constant and a readily explainable maximum i~ value of 1.0. The stability constant for the interaction Cu+2
+ AOH e C u h O + + H+
was reported as
K,
=
(CuXO+) (H+)
/ (Cut*) (XOH) (1)
because the ionization constant of the ethanolic hydrogen cannot be easily ascertained. The experimental method used by both Gonick and Bertsch consisted of titrating the metal ion in the presence of mineral acid with a solution of the amine. We found that the nature of the interaction is readily apparent if solutions of various amine to metal ratios are titrated with base as described by Chaberek and Martell ( 4 ) . When the differences in the two methods are taken into consideration, our results are in close agreement with the results obtained by Bertsch, Fernelius, and Block ( 1 , 2 ) . In the N i i 2 systems, there is evidence that the ethanolic group participates in chelation but that the
ethanolic hydrogen acts as a very weak acid. To clarify the interactions further, spectrophotometric methods were used to determine chelate ligand to metal ratios and the k , of the ethanolic group. Although the latter experiments were not entirely successful, we were able to obtain an approximate k , value which agreed well with a theoretical calculation based on inductive effects in organic compounds according to the method of Hine ( 8 ) and Dewar (6). EXPERIMENTAL
Analytical reagent grade hTi(NO& and Cu(NO& were used and the solutions standardized according to the chelatometric titrations described by Wilson and Wilson (11). T h e dihydrochloride salt of AOH was prepared by dissolving the amine in excess concentrated HC1, analytical reagent grade; flash-evaporating to dryness; and recrystallizing several times from methanol-ethanol-water mixtures, m.p. 163-4' C. All titrations were made a t 30 i 0.1' C. and at an ionic strength of 0.162 with NaCl as the supporting electrolyte to correspond with our enzyme studies. Ligand concentration was kept constant a t 0.02JP in all studies and the metal concentration was varied appropriately. I n all titrations, 0.1933M NaOH was the base employed. The experimental setup and procedure were as previously described (4). A Beckman Zeromatic pH meter was used to follow the potentiometric titrations of solutions containing various ligand to metal ratios. When necessary, a sufficient interim of time was allowed to elapse between additions of base until steady pH values were attained. Beckman Models DK-2 and DU spectrophotometers were used to analyze solutions of various ligand to
metal ratios according to the method of Job (9). The DK-2 wa5 also used in the experiments to determine the k , of the ethanolic group. This latter experiment consisted of preparing a series of AOH-SaOH solutions of known OH- concentration prepared with Fischer reagent ION NaOH. T h e ionic strength was kept constant with NaCI. The p H of these solutions was calculated using known activity values for XaOH solutions of known molarity. The solutions were scanned from 200 to 350 mp. T h e peaks for the dissociated and undissociated forms were not separable. The poor resolution observed may be attributed to the high OH- concentrations used in these studies. RESULTS A N D DISCUSSION
Titrations of Cu +2-AOH(HCI)2 mixtures and Ni +2-;\OH(HC1)2 mixtures, as well as free amine di HCI are shown in Figures 1 and 2. Ligand-Cu+2 ratios of 1-1 and greater all yield exactly three equivalents of H + per metal ion. This can only be explained on the basis of formation of a 1-1 terdentate chelate with simultaneous release of the ethanolic proton. That a 2-1 chelate does not form is evident from a comparison of curves .I ! B , and C in Figure 1 ; curve C is a summation of curves .I and B . Since a 2-1 chelate does not form, the possibility that the third proton comes from Cu+2-bound water is excluded. This system was also analyzed according to the method of .Job, utilizing the peak at 623 mp exhibited by the C u f 2 chelate. Unfortunately, mixtures containing a mole fraction of amine lethan 0.5 precipitated a t the pH of the experiment. Therefore no absorption VOL. 37, NO. 1 , JANUARY 1965
165
