V O L U M E 2 4 , N O . 6, J U N E 1 9 5 2
989
Results for the determination of iron(I1) obtained by the slope-intercept method agree to within 270 of the true value. While slope and intercept have not been found reproducible for duplicate runs, the area under the curve agrees with the theoretical and may be quickly evaluated from the slope and intercept with an accuracy sufficient for many purposes.
LITERATURE CITED
(1) (2)
(3) (4) (5)
(6) ACKNOWLEDGMENT
This research was suppofied in part from funds by The the university for University Research Foundation Ohio aid in fundamental research.
Latimer, W. M., “Oxidation Potentials,” New York, Prentice Hall, 1938. Lingane, J. J., Anal. Chirn. Acta, 2, 584 (1948). Lingane, J. J., IND.ENG.CHEM.,ANAL.ED., 17, 332 (1945). Lingane, J. J., J. Am. Chem. SOC.,67, 1916 (1945). Lingane, J. J., and Small, L. A,, ANAL.CHEM.,21, 1119 (1949). MacNevin, W. M., and Martin, G. L., J. Am. Chem. SOC.,71, 204 (1949).
RECEIVEDfor review January 11, 1952. Accepted March 10, 1952. Presented before the Division of Analytical Chemistry a t the 121st Meeting of the AMERICANCHEMICAL SOCIETY, Buffalo, N. Y. From a thesis submitted to the Graduate School of The Ohio State University in partial fulfillment of the requirements for the degree of doctor of philosophy, June 1951.
Mercurimetric Determination of Chlorides and Water-Soluble ChIorohydrins WILLIARI G. DOMASK AND KENNETH A. KOBE, The University of Texas, Austin, Tex.
A relatively simple method for analyzing chloride samples was needed, which would have a clear solution at the end point rather than a solution containing a precipitate, and would give highly reproducible results. A simple procedure has been developed which makes unnecessary a blank determination or correction for complexes that would otherwise affect results. Mercuric nitrate reagent is used to obtain a sharp end point in a clear solution, thereby permitting titrations which are easier to conduct than by methods that use silver nitrate. Complexes are accounted for by the technique of “calibrating” the reagent. The proposed method is routine and easy to conduct, and it permits rapid determinations with a high degree of precision. It can be employed readily by nontechnical personnel. ARIOUS modifications of the Volhard (10, 1 6 ) and Mohr v(6, 11) methods for chloride determinations have been applied successfully to mixtures containing inorganic and certain organic chlorides. Uhrig (16) suggested a method for chlorchydrin determination based on hydrolysis in the presence of sodium hydroxide or potassium hydroxide. Trafelet (14) presented an improved method for chlorohydrin samples, which, like the Uhrig method, is based on determination of the amount of inorganic chloride formed as a product of hydrolysis. The inorganic chloride content of a sample is obtained first by the Mohr method. Another portion of the sample is heated with sodium bicarbonate to hydrolyze selectively only the chlorohydrins; the solution is then neutralized and titrated as before and the chl,orohydrin chloride is determined by difference. A third portion of the sample is heated with sodium hydroxide t o hydrolyze all water-soluble organic chlorides, and the nonchlorohydrin organic chloride is determined by another difference calculation. Iiumerous aids and improvements have been suggested for argentometric methods for chlorides. These include filtration of the silver chloride precipitate ( 8 ) , use of a layer of organic solvent to remove the silver chloride from the aqueous layer (9, i g ) , use of nitrobenzene in a similar manner ( i ) ,and regulation of concentration of the ferric ion used as the indicator (13). MERCURIMETRIC METHODS
As pointed out by Kolthoff and Sandell (6),a reaction of great practical importance, which has not received the wide application it deserves, is that between halogen ions and mercuric ions to give soluble, slightly dissociated mercuric halides: Hg++
+ 2 C1-
HgC12
(1)
Dubsky and Trtilek ( 3 )and Roberta ( 7 ) described methods for chloride determinations in which mercuric nitrate is the reagent
and diphenylcarbazide and diphenylcarbazone are indicators. The method of Roberts includes the use of bromophenol blue as an indicator for adjusting the pH of the sample. Clarke ( 8 ) reported the mercuric nitrate-diphenylcarbazone method superior from the standpoint of stability of indicator and discernibility of end point. He also presented colorimetric and titrimetric methods for determination of chlorides in water, especially in trace quantities. In the titrimetric method, bromophenol blue is used to adjust the p H of the sample to approximately 3.6, and the sample is then titrated with 0.01 N mercuric nitrate solution. The mercuric ion reacts m-ith the chloride ion to form soluble and only slightly dissoGiated mercuric chloride. Any excess of mercuric ions forms a violet-colored complex with diphenylcarbazone. DEVELOPMENT OF MERCURIMETRIC METHOD FOR CHLOROHYDRINS
In the analysis of chlorohydrin samples it is generally desirable to use reagents that are approximately 0.1 AT. I n this concentration range, it was found that the diphenylcarbazone-bromcphenol blue mixed indicator recommended by Clarke ( 2 ) lacked the sharpness that was desirable for a high degree of precision. Hickman and Linstead (4)had found that xylene cyanole FF, when mixed with methyl orange, served to sharpen the methyl orange color change for base-acid titrations. Experiments showed that xylene cyanole FF, when added to the diphenylcarbazone indicator in optimum amount, gave excellent results in sharpening the end point without shifting the mercuric chloride equivalence point. When the mixed indicator containing diphenylcarbazone, bromophenol blue, and xylene cyanole FF is prepared as indicated below, it may be used both for pH adjustment and for obtaining a sharp mercurimetric end point. I n order to control the influence of pH, it is desirable to begin all titrations a t approximately the same pH. Bromophenol blue
ANALYTICAL CHEMISTRY
990
undergoes a color change from blue to yellow around pH 3.0 to 3.6. The blue color of xylene cyanole FF causes this color change to be from blue to green at about pH 3.5. After the pH of the sample has been adjusted, titration is conducted with mercuric nitrate reagent. During most of the titration, the sample retains a bluish-green color; a t a point which is usually 0.5 to 0.2 ml. before the end point, the solution acquires a grayish tinge, and as titration continues this becomes a faint slate color; a drop or two before the end point, the solution is essentially waterwhite. The end point is taken as the f i s t permanent tinge of lilac, as determined against a white background and with the aid of a titration illuminator. If more xylene cyanole FF is used than the amount specified, the blue color is intensified; if much less is used. the effectiveness is lost. I
I
I
I
I
1
0.1030 0.1020
0.1010
0
5 IO I5 20 25 30 35 40 45 5 0 MILLILITERS OF MERCURIC NITRATE REAGENT REQUIRED
Figure 1.
Calibration Curve for Mercuric Nitrate Reagent
If a sample is known to be either strongly basic or acid, it is advisable to neutralize it approximately against phenolphthalein indicator before adding the mixed indicator. The presence of phenolphthalein does not affect the mercurimetric end point. The mixed indicator may be safely used in the pH range of 1 to 11. In the region of pH greater than about 8, the mixed indicator gives a violet color, whereas it is blue between 3.5 and 8 and green below a pH of 3.5. It is necessary to stabilize the mercuric nitrate reagent with nitric acid; accordingly, when using concentrations in the range of 0.1 A‘, an appreciable quantity of nitric acid is introduced into the sample solution with the mercuric nitrate reagent. The amount of nitric acid present a t the end point has a definite influence on the amount of reagent required. Roberts ( 7 ) suggested that a correction blank be titrated which should contain not only the specified amount of indicator and nitric acid present a t the beginning of the titration, but also an additional amount of nitric acid equivalent to that introduced during titration. However, such a correction is incomplete when working in this concentration range. As pointed out by Kolthoff and Sandell (6), the amount of mercuric nitrate required will depend in part upon the concentration of mercuric chloride a t the end point, because a small portion of the mercuric ions added reacts with mercuric chloride, forming complex ions in accordance with the following equilibrium reaction: HgC1,
+ Hg
++
2HgC1+
(2)
Furthermore, experiments have shown that the influence of nitric acid is dependent to a great extent upon the amount of mercuric chloride present; accordingly, a correction blank for nitric acid would be inadequate, and the possible practice of neutralizing the excess nitric acid just before the end point js reached would not account for the complex ion formed according to Equation 2. The problem can be handled in a simple manner by “calibrating” the mercuric nitrate reagent, using a common starting point and following a routine procedure in all titrations. The method presented here involves the use of reagent grade water of reasonably uniform quality. The total volume of solution a t the end point should be between 100 and 125 m]. It is suggested that the indicator be added by means of a medicine dropper, to ensure reasonable uniformity in the amount added.
In this work, 5 drops of indicator were used, but one drop more or less did not affect the results. Essentially the same conditions prevail when standardizing the mercuric nitrate reagent as when titrating a sample; accordingly, the ratio of mercuric nitrate to chloride will vary in the same manner in the standardization procedure as when chloride samples are titrated. In Figure 1 is shown a calibration curve which presents the prevailing normality of a mercuric nitrate reagent as a function of the amount required for titration. Khen this method is used, no correction blank is needed, and the normality to be used in calculating results is that which prevails for the amount of mercuric nitrate reagent used. The curve in Figure 1 was obtained by using four different standard potassium chloride solutions. More points were obtained for presenting this curve than are required for routine work. -4s seen from Figure 1, it is desirable to have the sample of such size that 5 ml. or more of reagent are required for easily read values of N p . The curve rapidly approaches a straight line when more than 15 ml. of reagent are used. Essentially chloride-free sodium bicarbonate and sodium hydroxide are used in the hydrolysis steps, and these require no correction blanks. Sodium, potassium, calcium, and nitrate ions do not interfere. Tests indicate that cupric ions can be tolerated to a concentration of only about 25 p.p.m., whereas the ferric ion concentration should not exceed about 10 p.p.m. Preparations of mixed indicator have remained unaffected in clear glass bottles for periods up to 6 months, and preparations of mercuric nitrate reagent have maintained standardization in brown glass bottles, without precipitation of basic salts, for more than 4 months. REAGENTS
Mixed Indicator is prepared by dissolving 0.5 gram of C.P. crystalline diphenylcarbazone, 0.05 gram of C.P. crystalline bromophenol blue, and 0.12 gram of crystalline xylene cyanole FF in 100 ml. of 95% ethyl alcohol. Mercuric Nitrate Reagent is prepared by dissolving 53 grams of C.P. mercuric nitrate, Hg( N03)2.HzO, in 100 to 150 ml. of reagent grade water plus 6 ml. of C.P. concentrated nitric acid (70%) and diluting to 3 liters in a brown glass reagent bottle. (Note. If less than 2 ml. of concentrated nitric acid per liter are used for 0.1 N mercuric nitrate reagent, there is a tendency for a sediment to form in the solution. If much more than this amount is used, the pH of the reagent is lowered unnecessarily.) Other Reagents include: solutions of 0.1 N C.P. sodium chloride or potassium chloride, to serve as standards for calibration of the mercuric nitrate reagent; C.P. solid sodium bicarbonate; C.P. sodium hydroxide pellets; approximately 2 K nitric acid for coarse neutralization of strongly basic samples; approximately 0.1 N nitric acid for neutralization of slightly basic samples: and phenolphthalein indicator. PROCEDURE
Calibration of Mercuric Nitrate Reagent. The reagent need be calibrated only over the working range. An appropriate amount of standard chloride solution is pipetted into a 250-ml. Erlenmeyer flask, and an amount of reagent grade water is added such that the final volume of liquid a t the end point will be about 100 ml. Five drops of mixed indicator are then added. As this solution is approximately neutral, the color will be blue. A few drops of 0.1 S nitric acid are added to adlust the pH to about 3.5; this is indicated by a color change from blue to bluieh green. The solution is then titrated with mercuric nitrate reagent to the first permanent tinge of lilac. The calculated value of the prevailing normality, ATp, is plotted as a function of the milliliters of reagent used, as shown in Figure 1, and this operation is repeated until several points are obtained to determine the curve over the desired range. Analysis of Samples. The samples are prepared by follon Ing the general procedure outlined by Trafelet (14), and titrations are conducted as described above for the calibration procedure. For each titration, the prevailing normality, N p , of mercuric nitrate is obtained from the calibration curve. To report the per cent chloride as C1-, the calculation is made as follows:
(MI. of mercuric nitrate reagent) ( N p )(3.5462 (Actual weight of sample titrated)
-
%Cl-
V O L U M E 24, NO. 6, J U N E 1 9 5 2
991
The chloride from inorganic compounds, chlorohydrin, and other organic chlorides is determined by difference calculations. I n the hydrolysis of organic chlorides, approximately 0.5 gram of sodium hydroxide pellets may be used instead of a sodium hydroside solution.
However, reasonable precautions, such as are exercised with silver nitrate to prevent staining of the skin, are adequate in the handling of this reagent. ACKNOWLEDGMENT
DISCUSSION
The same trend in data as is shon-n in Figure 1 was obtained when xylene cyanole FF was omitted from the indicator; therefore, the use of xylene cyanole FF does not affect the titration r w q ) t to make the end point sharper and easier to detect. Inasmuch as titrations in the calibration procedure and in the analysis of sam.ples have a comm.on initial pH, there is no concern about controlling the pH during titration or at the end point; changes in p t l resulting from titration are accounted for by using the appropriate value for .Vp. -1s seen in Figure 1, there is a reasonably large region over the 5at portion of the curve for which an average normality would give good accuracy. I n many cases of routine industrial analyses, such an average value will be ent,irely satisfactory. The practice of establishing a common initial point and a controlled final volume for all titrations is by no means a disadvantage of the method. Such procedure is desirable for all volumetric methods and is necessary for most methods to ensure reliable results. However, it was found that a variation of +25 ml. from the specified 100 ml. did not noticeably affect results. This mercurimetric method yields excellent precision and reproducibility of results. Because the sample solution remains clear, titrations can be made rapidly. The series of color changes which occurs just before the titration is completed serves as a signal to help prevent overrunning the end point. .Sttention is called to the toxic naturr of mercuric nitrate.
The authors \?--ish to thank Granville TV. Burtt, Research Laboratories, Jefferson Chemical Co., for his helpful suggestions and trial of the method. LITERATURE CITED (1) Caldwell, J. R., and Moyer, H. V., IND.ENG.CHEM.,ANAL.ED.. 7, 38-9 (1935). (2) Clarke, F. E., ANAL.CHEX.,22, 553-5, 1458 ( 1 9 3 ) . (3) Dubsky, J. V., and Trtilek, J., Mikrochemie, 12, 315-20 (1933). (4) Hickman, K. C. D., and Linstead, R. P., J . Chem. Soc., 121, 2502-6 (1922). (5) Kolthoff, I. hl., and Sandell. E. B., “Textbook of Quantitative Inorganic Analysis,” 2nd ed., p. 576, New York, Macmillan Co., 1946. (6) Mohr, F., Ann., 97, 335-8 (1856). ED., 8 , 365-7 (1936). (7) Roberts, Irving, IND.ENG.CHEM.,AN.AL. (8) Rosanoff, M A, and Hill, A. E., J . Am. Chem. SOC.,29, 269-75 (1907). (9) Rothmund, V., and Burgstaller, A., Z. anorg. Chem., 63, 330-6 (1909). (10) Scott, W. D.. “Standard Methods of Chemical Analysis,” Val. I, 5th ed., p. 271, New York, D. Van Nostrand Co., 1939. (11) Ibid., p. 272. (12) Stschigol, M. B., 2. anal. Chem., 91, 182-5 (1932). (13) Swift, E. H., Arcand, G . M., Lutwack, R., and Meier, D. J., ANAL.CHEM.,22, 306-8 (1950). (14) Trafelet, Lella, Ibid., 20, 68-9 (1948). (15) Uhrig, Karl, IND. ENQ.CHEM.,ANAL,ED., 18, 469 (1946). (16) Volhard, J., J . p m k l . Chem., 117. 217-24 (1874). RECEIVED for review October 20,1951. Accepted -4pril7, 1952
Small Amounts of Copper in Dyes and Rubber Chemicals Colorimetric Determination with Zinc Dihenzyldithiocarbamate R. I. MARTENS
AND
R. E. GITHENS, SR.
Jackson Laboratory, E . I . du Pont de N e m o u r s &
T
co.,Inc., Wilmington, Del.
HE method in use in this laboratory for the determination of
This work was undertaken to provide a more specific method for the routine determination of small amounts of copper in rubber chemicals, dyes, and dyed fabrics. A dithiocarbamate-type reagent -a 0 . 0 1 7 ~ solution of‘ zinc dibenzyldithiocarbamate in carbon tetrachloride-will selectively extract copper from acid solutions containing relatively large amounts of most other metals. The method .is accurate and reproducible to ahoiit 0.5 microgram in the range of 0.5 to 40 microgranis of copper. No precipitations or filtrations are required. The reagent is commercially available, the reagent solution and the color of the extracted copper dibenzyldithiocarbamate are stable, and the acid solution remaining after extraction may be used to determine other metals. Because copper dibenzyldithiocarbamate is selectively extracted from solutions containing relatively large amounts of many other metallic ions, the reagent may be used to determine small amounts of copper in metals and their salts.
small amounts of copper (together Rith a method for manganese on the same sample) in dyes and other organic materials was described in 1940 by Palfrey, Hobert, Benning, and Dobratz ( 3 ) . Sulfuric and nitric acids and hydrogen peroxide are employed for destruction of organic matter, and copper is determined by measurement of the intensity of color developed with sodium diethyldithiocarbamate in an ammoniacal medium. Although provision is made for preliminary separation from a number (but not all) of the interfering metals, the separations involving filtrations are troublesome and there has often been a question whether copper is completely recovered. Even small quantities of nickel, cobalt, and bismuth interfere, providing they exceed somen.hat the amount of copper. A better method was needed. The advantages of extraction of copper diethyldithiocarbamate into an immiscible organic solvent have been summarized by Sandell (4). I t seemed likely that many interferences might be eliminated if extraction could be made from acid rather than from ammoniacal solution as is usually done. Because sodium diethyldithiocarbamate appeared to be too unstable to be uscd with acid solution, a search was made for other dithiocarbamates which might be sufficiently stable, sensitive, and selective. Selection. A number of dithiocarbamates were tried as reagents in n-ater, chloroform, or carbon tetrachloride solutions of
