was caproic acid (identified after esterification with CH2N2). The splitting-out reaction was not complete, however, because a-hydroxy heptanal was also formed. Since neither the acids nor the a-hydroxy fragments were observed following roomtemperature ozonolysis in methanol, this procedure is more effective and convenient. It yields fragments which are easily identifiable without supplementary reactions.
+ CH,-JCHiJrC”o
\o-cnI
(4.6%)
CHi-(CHi),-CH-C=C-C=C-(CH~),-C= nt I OH
,
ACKNOWLEDGMENT
/o
n n n
O-CHl ICoriolatr)
The geometry of the double bonds apparently had no effect on the cleavage pattern. Splitting-off of the carbon atom adjacent to the hydroxylated atom was observed upon ozonolysis in CHzClz at dry-ice temperature (20); the fragment formed from methyl coriolate (20) W. H. Tallent, J. Harris, G. F. Spencer, and I. A. Wolff, Lipids, 3, 425 (1968).
The authors thank R. 0. Buttefield for methyl stearolate, R. G. Powell for methyl 17-octadecen-9-ynoate, W. H. Tallent for methyl coriolate and methyl dimorphecolate, and T. K. Miwa for stimulating suggestions and discussions.
RECEIVED for review June 30, 1969. Accepted September 5, 1969. Presented at the Great Lakes Regional Meeting of the American Chemical Society, DeKalb, Ill., June 5-6, 1969. The Northern Regional Research Laboratory is headquarters for the Northern Utilization Research and Development Division, Agricultural Research Service, U. S. Department of Agriculture. Mention of trade or company names is for identification only and does not imply endorsement by the Department.
Indicator Titrations in Tetramethylurea Siegmond L. Culp and Joseph A. Carusol Department of Chemistry, Uniaersity of Cincinnati, Cincinnati, Ohio 45221
RECENTLYwe reported on tetramethylurea (TMU) as a solvent for potentiometric acid-base titrations ( I ) . It was shown that TMU is indeed a useful solvent and is suitable for differentiating a wide range of acids and bases. Results of the first study indicated that perhaps other methods of detecting the end points also might be investigated. Indicators have proved to be especially helpful in the detection of end points in cases where especially steep potential curves are observed in nonaqueous titrations (2). Because the titration curves obtained in TMU tended to be rather steep, the investigation of suitable indicators seemed an appropriate study. The most commonly used indicators in nonaqueous titrimetry are thymol blue (thymolsulfophthalein), azo violet [4-(pnitropheny1azo)-resorcinol], and crystal violet (hexamethyl-prosaniline hydrochloride) (3). Crystal violet, however, is used almost exclusively in acetic acid while the first two mentioned have been found to be operative in a wide variety of solvents ( 4 ) . Fritz has recommended thymol blue as the most generally satisfactory indicator in DMF while azo violet, with a more basic transition point, was found to be preferable when very week acids were titrated (5). 1
To whom all communications should be addressed.
(1) Siegmond L. Culp and Joseph A. Caruso, ANAL.CHEM., 41, 1329 (1969). (2) I. Gyenes, “Titration in Non-Aqueous Media,” D. Van
Nostrand Co., Princeton, New Jersey, 1967, p 198. (3) Zbid.,p 204.
(4) J. T. Stock and W. C. Purdy, Chemist Analyst, 48,50 (1959). ( 5 ) J. S . Fritz, ANAL.CHEM., 24, 306 (1952).
1876
EXPERIMENTAL
All of the indicators used were obtained from Eastman and were “white label” or reagent grade. These were dissolved in a sufficient amount of purified tetramethylurea to produce a 0.3% solution. From two to five drops of indicator solution were used in titrating samples of benzoic acid, phenol, or 1,3-diphenylguanidine. All indicator titrations were monitored potentiometrically to establish the color change which best corresponded to the equivalence point. The purification of other reagents, descriptions of apparatus, and experimental procedures, have been discussed previously ( I ) . RESULTS AND DISCUSSION
The indicators evaluated for the titration of acids were thymol blue, phenolphthalein, azo violet, alizarin yellow, and curcumin. Thymol blue, phenolphthalein, and azo violet all performed satisfactorily in the titration of benzoic acid. With alizarin yellow, the color change did not occur until well after the benzoic acid equivalence point had been passed. Curcumin was evaluated only in the titration of phenol. Azo violet and curcumin were the only indicators with sufficiently basic transition points to produce a color change during the titration of phenol. However, the color transitions were not sufficiently sharp to permit satisfactory visual determination of the equivalence point. These indicators might be quite suitable, however, for the spectrophotometric titration of very weak acids in tetramethylurea.
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
-201
AZO VIOLET INDICATOR
T H Y M O L B L U E INDICATOR
-
Orange- yellow
-301
- 30'
E
-> -
LL
W
-401 c
->
Red - oronge
4-
-401
E
LL
I
P
-501
W
-50'
1.-
Green
Violet
-60( -60
-70' I
1
I
1
2
3
4 5 m.l T B A H
6
7
8
1
I
I
1
2
3
I I 4 5 ml T B A H
6
8
Figure 2. Titration of benzoic acid in the presence of azo violet
Figure 1. Titration of benzoic acid in the presence of thymol blue
B
BROMOPHENOL B L U E INDICATOR
The titration curves and corresponding color transitions obtained with thymol blue and azo violet, are shown in Figures 1 and 2. Phenolphthalein produced its usual colorless to violet transition at the benzoic acid equivalence point and is not illustrated. Visual detection of the equivalence point was easiest with thymol blue, because of the green transition immediately preceding the end point. Conversely, the azo violet end point, with its red to violet color change, was the most difficult to determine visually. Recoveries also were calculated for indicator titrations by comparing with potentiometric titrations made without any indicator present. Both titrations were performed on the same day. In this manner, any blank due solely to the presence of indicator could be differentiated from the solvent blank (should the latter be appreciable) and determined under the conditions normally prevailing at the end point. On this basis, the recoveries obtained in titrations of benzoic acid with thymol blue, phenolphthalein, and azo violet indicators were 99.80 (0.33), 100.29 (0.37), and 99.34 (0.04), respectively. (The numbers in parentheses represent the average deviation of triplicate determinations). Thus it appears that there is no appreciable indicator blank. For the titration of bases, bromophenol blue and methyl orange indicators were evaluated in the titration of 1,3-diphenylguanidine. Bromophenol blue gave excellent end points and a titration curve showing the color transition for this indicator is presented in Figure 3. A recovery of 99.372 was obtained for the titration of 1,3-diphenylguanidine with bromophenol blue indicator (individual titration). Methyl orange did not change color until after the equivalence point had been passed. The potentiometric curves plotted in the presence of thymol blue and azo violet were obtained using a different glass elec-
7
i
Yellow
I
I
I
1
2
3
I
I
4 5 rnl HC104+
I
I
I
6
7
8
Figure 3. Titration of 1,3-diphenylguanidine in the presence of bromophenol blue
trode then that used in the previous titration study ( I ) . The addition of indicator just prior to titration did produce a small shift in potential (-25 mV for thymol blue and - 3 mV for azo violet). Even without indicator, the initial potentials for the
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
1877
benzoic acid titration curves were about 145 mV more positive with the glass electrode used in the indicator studies even though the same make and model of electrode was used. The differences in the potentials spanned in the corresponding benzoic acid curves (Figure 1 and 2 of this study as compared to Figure 1 of reference I ) , are therefore due primarily to differences in electrode efficiency rather than the presence of indicators. The potentials at the equivalence point were more consistent than those preceding it and varied from - 561 to -596 mV in its presence of indicators. Values obtained in the absence of indicators and with a different glass electrode
also fell within this range. The effect of indicators on the shape of the titration curves obtained with benzoic acid was evidenced by a slight steepening of the plateaus at the beginning of the titration. It may be concluded that thymol blue, phenolphthalein, and azo violet can be used as visual indicators in titration of acids in TMU. Thymol blue is perhaps the most generally satisfactory, while azo violet is preferred for very weak acids. Bromophenol blue is satisfactory for the titration of bases.
RECEIVED for review July 7, 1969. Accepted August 20, 1969.
Effect of Auxiliary Complexing Agents on the Rate of Extraction of Zinc(l1) and Nickel(l1) with Diphenylthiocarbazone P. R. Subbaraman,’ Sr. M. Cordes,2 and Henry Freiser Department of Chemistry, University of Arizona, Tucson, Ariz. 85721
IT HAS
BEEN SHOWN PREVIOUSLY ( I , 2) that the rate of extraction of various metal chelates was limited by the rates of formation of the 1 :1 chelates from the hydrated metal and the ligand anion. Because the rates of these homogeneous chemical reactions were shown to be closely related to the rate of loss of coordinated water from the hydrated metal ion ( I ) , it was decided to investigate the effect of replacing at least one water molecule by another (auxiliary) ligand on the rate of chelate formation. Both mono- and bidentate ligands were selected for the study which focused on the zinc(I1) and nickel(I1) dithizone systems.
EXPERIMENTAL
Apparatus. The apparatus is essentially the same as that previously reported ( I ) . The samples were agitated at the high speed setting of the Eberbach shaker. A Cary Model 14 spectrophotometer was used for the absorbance measurements of dithizone. Estimations of zinc(I1) and nickel(I1) were made with a Model 303 Perkin-Elmer atomic absorption spectrophotometer using resonance lines at 214 and 232 nm, respectively. Materials. Solutions of dithizone in CHCI3 were prepared weekly from purified (1) reagent grade diphenylthiocarbazone (dithizone) [Fisher Scientific Co.] and confirming assays obtained at frequent intervals by spectrophotometry using A,, = 606 mp and E = 4.06 x 104 (1). Deionized water and reagent grade chemicals were used throughout the study. The zinc(I1) and nickel(I1) solutions were prepared from stock solutions (1.00 x 10-3M) of their perchlorates. Solutions of mercaptoacetates were stored under a nitrogen atmosphere. Buffer components used to cover the pH range from about 4 to 7 included formate, phthalate, and phosphate at sufficiently low concentrations so that metal complexation with these ions was absent. Kinetic Measurements. The procedure is similar to that previously described ( I ) . A batch of an aqueous phase was prepared so as to contain measured concentrations of the metal cation, auxiliary complexing agent, an appropriate 1 2
National Research Laboratory, Poona, India. Creighton University, Omaha, Neb.
buffer, and enough NaC104 to give an ionic strength of 0.25. Ten-milliliter aliquots were carefully placed over 10-ml portions of CHCl, solutions of dithizone in a series of about twenty vials which were then sealed and positioned in the shaker, which was thermostated at 25 f 0.1 OC, to permit the removal of individual vials at given time intervals without disturbing the others. The extraction time was taken as the interval between starting the shaker and the time the particular vial was removed. Once removed, the vial was allowed to stand a few minutes to complete phase separation and then a 5-ml aliquot of the aqueous layer was transferred for metal analysis uia atomic absorption. The pH of the aqueous phase was checked both before and after extraction. Measurement of such a series of extractions resulted in a single value of an apparent rate constant which was obtained from a plot of logarithmic values of the metal concentration us. time. These were in all cases found to be linear, indicating first order dependence on total metal ion concentration throughout. In any series of kinetic measurements with a given metal and auxiliary ligand pair, the dithizone concentration in the CHCI3 was generally kept constant (except in those cases when it was desired to test the validity of the assumption of first order dependence on dithizone) and the pH was varied whenever expedient to bring the apparent rates within the optimum range of measurement. The auxiliary ligand ion concentration at any given pH was computed using the appropriate published p K , values (3). The validity of the previously employed (1) rate expression
was here reconfirmed for several cases in which an auxiliary ligand (acetate or thiocyanate) was present, in the following modification: d[W/dt =
- kppC.dDz-1
(2)
where C,Mrepresents the total metal ion concentration in the aqueous phase. Regardless of the nature or concentration of auxiliary ligand used, the extracted species was the expected simple metal dithizonate in every case.
(3) L. G. Sillen and A. E. Martell, Ed., “Stability Constants of
(1) B. E. McClellan and H. Freiser, ANAL.CHEM., 36, 2262 (1964). (2) J. S.Oh and H. Freiser, ibid., 39, 295 (1967). 1878
Metal Ion Complexes,” Special Publication No. 17, Chemical Society, London (1964).
ANALYTICAL CHEMISTRY, VOL. 41, NO. 13, NOVEMBER 1969
