Acid-Base Titrations in Glacial Acetic Acid Acid Potassium Phthalate as Primary Standard and Behavior of Crystal Violet Indicator WILLIAM SEAMAN AND EUGENE ALLEN American Cyanamid Co., Calco Chemical Diuision, Bound Brook, N. J .
The titration of weak bases with perchloric acid in glacial acetic acid has suffered from the lack of a primary standard which may be purchased ready for use. There has also been some confusion about the color change which would indicate the end point correctly with the commonly used crystal violet indicator. In this paper acid potassium phthalate as purchased from the National Bureau of Standards
I
N A series of pioneer investigations, Hall ( 6 ) ,Hall and Conant ( 7 ) , and Hall and Werner (8) have shown that many organic bases which are too weak to be titrated in water can be successfully titrated in glacial acetic acid, giving sharp end points similar to those obtained with strong bases in water. This method has since been applied to a variety of substances using both visual indicators and potentiometric procedures ( 1 , 2 , 4 , 5 , 9 - 1 1 , I 3 - 2 1 ) . Compounds which have acidic ionization constants of approuimately 10-6 in water are neutral in this medium and, thus, carboxyl groups do not interfere nith the titration. Substances such as perchloric, sulfuric, and hydrochloric acids react as strong acids. The titrating agent is usually a 0.1 Y solution of perchloric acid in glacial acetic acid. Perchloric acid is used in preference to either sulfuric or hydrochloric acid because it is a stronger acid in this medium. In addition, sulfuric acid introduces complications because of its diprotic character, and solutions of hydrochloric acid in glacial acetic acid are unsatisfactory because of the high rate of escape of hydrogen chloride (8). The primary standard most commonly recommended for this type of titration is sodium carbonate. This standard reacts with the medium to give acetate ion which is then titrated by the perchloric acid solution being standardized. Since sodium carbonate is not furnished as a certified standard reagent by the National Bureau of Standards, it is necessary either to prepare it in a pure form or to run tests for impurities on C.P. material before it can be accepted for use as a primary standard. Glycine ( I 7 ) , diphenylguanidine ( 1 8 ) , and 1-naphthylamine ( 2 1 ) , all of which have been recommended as primary standards, suffer from the same disadvantages, It has been found in this laboratory that acid potassium phthalate is an excellent primary standard for this purpose, This substance gives values which agree with those obtained with sodium carbonate, and it has the advantage of having a higher equivalent weight and of being available as a certified standard reagent from the National Bureau of Standards. Furthermore, it has a low hygroscopicity, and in most cases it is unnecessary to dry the material before use. The titration of dipotassium phthalate in glacial acetic acid medium has been reported in connection with a modified Kappelmeier procedure for determining phthalic anhydride in oil-modified alkyd resins ( 5 ) . Ionization of acid potassium phthalate t o potassium and biphthalate ions occurs in both aqueous and acetic acid media; however, in water, the biphthalate ion acts as a weak acid by virtue of partial ionization to hydrogen and phthalate ions, whereas in glacial acetic acid the biphthalate ion, like the acetate ion, is a base and can be titrated with strong acids.
is recommended as a standard. Crystal violet changes through a variety of colors during a titration. The correct color change at the end point depends upon the system being titrated; i t must be determined by reference to the potentiometric end point. Adoption of these recommendations should lead to greater convenience and accuracy approaching that obtainable in aqueous acid-base titrations.
Crystal violet is the indicator which is most commonly recommended in those procedures in which indicators are used. A source of difficulty in using this substance is the variety of colors through which i t changes during the course of a titration ( I S ) . The behavior of crystal violet in several typical titrations was therefore investigated in an effort to work out a valid indicator procedure. Conant and FTerner ( 3 ) have shown that the color of crystal violet in glacial acetic acid solution depends markedly on the ionic strength of the solution as well as on the pH (H4c) acid, but these conclusions have not been considered in connection with titration techniques. AClD POTASSIUM PHTHALATE A S PRIMMARY STANDARD
A series of potentiometric titrations was run on both acid potassium phthalate and sodium carbonate, using the same perchloric acid solution, and arranging conditions so that the volume of titrant used was approximately 40 ml. and the total volume a t the end point was approximately 100 ml.
AClD POTASSIUM PHTfiALArE
Figure 1. Color Changes of Crystal Violet Indicator in Titration of Various Substances
The titrations were conducted in a 150-ml. wide-mouthed flask provided with a rubber stopper which had openings for a glass electrode, the tip of a buret (immersed below the surface of the solution), and one leg of a salt bridge. A slit from one of the holes to the circumference of the stopper served as an air vent. (This arrangement was provided in order t o minimize absorption of 592
593
V O L U M E 23, NO. 4, A P R I L 1 9 5 1 atmospheric moisture. Although a small amount of water, such as that ordinarily found in C.P glacial acetic acid, will not interfere with the titration, the presence of too much water will result in an indefinite end point.) The salt bridge, a U-tube having a ground-glass joint, with a glass stopper to fit, a t each end, was filled a ith a supersaturated solution of lithium chloride in glacial acetic acid as recommended by Hall and Conant (7). The other end of the salt bridge, as well as a calomel reference electrode, was immersed in a saturated aqueous potassium chloride solution. (L. Lykken, in a review of this paper, stated that a nonaqueous salt brid e has not been found necessary in his laboratory, and suggeste3 that when a calomel electrode is used directly the sleeve type gives more stable potentials over a period of time than the fiber type, because it is more easily cleaned.) The glass and calomel electrodes were connected to a Beckman pH meter, laboratory model G. In the early stages of the work, a magnetic stirrer was employed in the titration vessel, but this was discontinued after erratic potential readings were occasionally encountered. I n the remainder of the work, agitation was accomplished by swirling the titration vessel by hand. Baker's C.P. glacial acetic acid was used throughout the investigation. The sodium carbonate used for the standardization was reagent grade material which had been heated for 2 hours a t 120" C. and then tested for suitability as a standard by the procedure of Kolthoff and Sandell (19). (Ten grams of the salt lost only 1 mg. when heated for 1 hour a t 300" C., showing that the diying treatment a t 120" C. was adequate.) The acid potassium phthalate was National Bureau of Standards material, used without preliminary drying. The titrating solution, 0.1 N perchloric acid in glacial acetic acid, u a s prepared as follous: 50 ml. of acetic anhydride ere placed in a flask which was submerged in an ice bath. Perchloric acid (10 ml. of 70 to 72%) was added in small portions 1% ith constant swirling, not allowing the temperature of the solution to rise above 20" C. The resulting mixture was diluted to 1 liter with glacial acetic acid. The crystal violet indicator used nas a 0.1% solution of a commercial grade in glacial acetic acid. In all titratiods employing solutions in glacial acetic acid allov-ance must be made for the high coefficient of cubical expansion of this solvent. In this investigation, it was assumed that the 0.1 -1-solution of perchloric acid in glacial acetic acid had the bailie coefficient of expansion as acetic acid itself-namely, 0.0011 per C. The temperature of the titrating solution a t the time of titiation was read on a thermometer which was attached to the buret, and observed volumes of titrating solution were corrected to an arbitrarily chosen standard temperature of 29" C. by a factor of 0.11% per O C. The values obtained potentiometrically for the normality of the perchlorir acid solution with sodium carbonate as standard nere 0.11747, 0.11736, 0.11748, 0.11749, 0.11730, and 0.11745; average, 0.11743: standard deviation, =!=0.00008. With acid potassium phthalate as standard the values were 0.11762, 0.11753, 0.11747, 0.11747, 0.11764, and 0.11740; average, 0.11752; standard deviation, *0.00009. Titration curves are shown in Figure 1. The precision of the standardization against acid potassium phthalate is substantially the same as that against sodium carbonate, and the normalities obtained by both methods do not differ significantly from each other. The indicator end point in the acid potassium phthalate titration (determined as described in later sections of this paper) was the first observable change froin violet to blue and was quite sharp. The correct end point color in the sodium carbonate titration was bluish-green, and was best recognized by comparing the color with an appropriate buffer solution. The use of acid potassium phthalate appears, therefore, to be simpler when crystal violet indicator is used. In the acid potassium phthalate titration, a precipitate of potassium perchlorate is produced upon the addition of the titrating solution; this precipitate does not interfere with the observation of the indicator end point.
BEHAVIOR OF CRYSTAL VIOLET INDICATOR
Change of Color during Titration. In order t o study the behavior of crystal violet indicator in the titration of several substance3 in glacial acetic acid, potentiometric titrationc: were run
on solutions of these substances, which also contained the indicator. The colors of the solutions were observed during the course of the titrations, and were compared with the colors of six arbitrarily prepared reference standards. To prepare the six color standards, 60 ml. of glacial acetic acid were placed in each of six 126-ml. glass-stoppered Erlenmeyer flasks. Crystal violet indicator solution (5 drops of 0.1%) was added to each flask. Then 0.1 N perchloric acid in glacial acetic acid was added to each flask as follows: h*O.
of
Standard 1 2 3 4 5
6
0.1 N HCIOd Added, Drops
Color Violet Blue Blue-green Green Green-yellow Yellow
0
2 5
7 12 1 to 2 ml.
Titrations were run on the following solutions: 0.2522 gram of sodium carbonate dissolved in 60 ml. of glacial acetic acid 0.9727 gram of acid potassium phthalate dissolved in 60 ml. of glacial acetic acid 0.3574 gram of glycine dissolved in 60 ml. of glacial acetic acid with the aid of gentle heating, - to room _. followed by cooling temperature 0.6531 gram of methyl nicotinate dissolved in 2 ml. of benzene followed bv 60 ml. of nlacial acetic acid A b1ank"of 60 ml. o r glacial acetic acid
lr --
4:
ACID POTASSIUM PHTHALATE
L -
.4 -
Q -
i-
4 -
7:-
GLYCINE
SODIUM CARBONATE
METHYL NICOTINATE
--
2
t- I
I I I 'I, 1 I I I I I I I I I I I I I 1 I M L . 0. I N PERCHLORIC
ACID
Figure 2. Variation of Slope of Titration Curve at Inflection Point for Various Substances
To each of these solutions were added 5 drops of 0.1% crystal violet indicator solution. The titrations were run potentiometrically using the equipment described. The strength of the perchloric acid solution was 0.1134 X. After the addition of each increment of perchloric acid, the color of the solution was compared visually with the color standards, and judged as being either identical with one of the standards or between two adjacent standards. The results of these titrations are shown graphically in Figures 1 and 2. Figure 1 gives the entire titration curves and shows the color changes. Figure 2 shows curves for the titration of the same substances in which the portion of the curve a t the immediate vicinity of the inflection point is plotted on an expanded scale to demonstrate more clearly the differences in slope. The curves in both diagrams are arranged in the order of decreasing slope a t the inflection point from left to right. The potential a t the inflection point has been placed on the same horizontal line for all the plots, because the absolute potential readings a t the inflection point varied approximately 30 to 50 mv. for successive titrations of the same substance. The color of the solution a t
ANALYTICAL CHEMISTRY
594 the inflection point appears to vary with the slope of the curve a t this point; those substances having curves with greater slopes are on the blue side a t the inflection point and those with lesser slopes are on the yellow side. The slope of the curve a t the inflection point is influenced by the strength of the base being titrated and, possibly, by the concentration and nature of the ions in solution a t this point-for example, in the titration of acid potassium phthalate, potassium ion is precipitated as potassium perchlorate, whereas with sodium carbonate, the corresponding sodium ion is not precipitated. Because of the differences in color a t the end point, a preconceived color change cannot be used for all titrations, but the correct color change for each individual titration must be determined by observing the color of the indicator a t the true end point as determined potentiometrically. If this is done, it is believed that titration in glacial acetic acid medium can be performed with a precision and accuracy approaching titration in aqueous medium (as indicated by data on the titration of methyl nicotinate). The blank titration curve on glacial acetic acid alone, shown in Figure 1, is interesting because it indicates that a blank determination on the indicator as was used by some workers (17) is erroneous. The shape of the curve shows clearly that the consumption of perchloric acid required to produce the color changes is caused by an effect of the medium rather than by the presenre of basic impurities. Use of Color Reference Solution. Titrations of substances which do not give very sharp breaks a t the end point are somewhat confusing, because the color changes are not abrupt and occur over a considerable range of addition of perchloric acid I n such cases, it has been found convenient to use a reference solution having the correct crystal violet color for the end point of that particular titration. Such a solution could obviously be prepared by bringing a solution of the substance being titrated to the potentiometric end point and then adding crystal violet indicator. However, such a system would not be well buffered, so it was decided to use instead a buffer containing urea and perchloric acid adjusted to the proper crystal violet color. Such a solution was prepared as follows for titrating a benzene solution of methyl nicotinate: A 0.6gram portion of niethyl nicotinate was dissolved in 2 ml. of benzene. To the solution were added 60 ml. of glacial acetic acid and 5 drops of 0.1% crystal violet indicator solution. In a second flask, a solution of 0.5 gram of urea and 20 ml. of 0.1 N perchloric acid solution in 40 ml. of glacial acetic acid containing 5 drops of crystal violet indicator solution was prepared. The flask containing the methyl nicotinate solution was set up for potentiometric titration with the buret tip above the surface of the solution. The flask containing the urea solution was stoppered tightly and was placed beside the first flask so that they might be viewed simultaneously under a strong light. The 0.1 iY perchloric acid solution was added to the titration flask until the methyl nicotinate solution had turned the same shade of green as the urea solution. The 0.1 11-perchloric acid solution was now added to the methyl nicotinate solution in 2-drop increments. Before the addition of each increment, the galvanometer needle of the Beckman p H meter was brought to zero by means of the large p H knob. After the addition of the increment, the deflection of the galvanometer needle to the nearest tenth of a division was recorded. (This was found to be a more sensitive method of obtaining the slope of the curve than recording millivolt diffprences.) After the addition of each increment of perchloric acid solution to the methyl nicotinate solution, the color of the urea buffer solution was adjusted to match that of the methyl nicotinate solution by the addition of 0.1 N perchloric acid. When the maximum galvanometer deflection was reached, addition of the perchloric acid t o the urea solution was stopped, and the titration of the methyl nicotinate solution was continued somewhat to make sure that the true peak in the curve had been attained. The volume of the buffersolution was continuously adjusted during the titration, by adding glacial acetic acid, in order to keep it approximately equal to the volume of the methyl nicotinate solution. When the buffer solution had been brought to the correct color, it was immediately transferred to a flask of the same size and shape as the one in which the methyl nicotinate titrations were to be carried out, and the flask was stoppered securely.
Two separate color reference solutions were prepared in this manner, and were labeled reference solutions 1and 2, respectively. Methyl nicotinate solutions were prepared by, dissolving accurately weighed O.6gram portions of the compound in 2-ml. portions of benzene, which was follo~vedby adding 60 ml. of glacial acetic acid containing 5 drops of 0.1% crystal violet indicator solution to each solution. The flasks containing these solutions were stoppered until immediately before the titration was started. Titrations were carried to the point where the color of the methyl nicotinate solution exactly matched that of the reference solution. The methyl nicotinate used in t,hese experiments as a synthetic material which had been carefully purified by fractional distillat,ion. Duplicate assay values (by saponification and backtitration of excess alkali) were 100.2 and 100.3%. Moisture (by Karl Fischer titration) was 0.027& -1test for nicotinic acid by direct titration in aqueous sblution to the phenolphthalein end point was negative. The values obtained using reference solution 1 were 100.3, 100.2, 100.1, 100.0, 100.2, and 100.5; average, 100.2; standard deviation, k0.2. Using reference solution 2, the values obtained were 99.9, 100.2, and 100.3; average, 100.1; standard deviation, 10.2. These figures show that the two reference solutione did not give significantly different values. ACKNOWLEDGMENT
The authors are indebted to Herman Cherlow, who purified the methyl nicotinate used in this work. LITERATURE CITED
Blunirich, K., and Bandel, G., Angew. Chem., 54, 374 (1941). Conant, J. B., Chow, B. F., and Dietz, E. M., J . A m . Chcm. Soc., 56, 2185 (19343.
Conant, J. B., and Werner, T. H., Ihid.. 52, 4436 (1930). Cropper, F. R., J . Oil &. Colour Chemists’ Assoc., 28, 207 (1946). Goldberg, A. I., IND.ENG.CHEM.,ANAL.ED.,16, 198 (1944). Hall, N. F., J . Am. Chem. Soc., 52, 5115 (1930). Hall, N. F., and Conant, J. B., Ibid., 49, 3047 (1927). Hall, N. F., and Werner, T. H.. Ibid., 50, 2367 (1928). Harris, L. J., Biochem. J., 29, 2820 (1935). Harris, L. J., J . Biol. Chem., 84,296 (1929). Haslam, J., and Hearn, P. F., Anaz2/st, 69, 141 (1944). Kolthoff, I. AT., and Sandell, E. B., “Textbook of Quantitative Inorganic Analysis,” p. 518, New York, Macmillan Co., 1937. Nadeau, G. F., and Branchen, L. E., J . Am. Chem. Soc., 57, 1363 (1935).
Oppenheim, J. C., Soc. Chem. f n d . ‘C’ictoria, Proc., 45, 647 (1945).
Russel, J., and Cameron, A. E., J . Am. Chem. SOC.,58, T74 IlR3AI. \ - - - - ,
Spengler, H., and Kaelin, A . , Hundert Jahre Schweiz. Apoth.Ver., 1843-1943, 542 (1943).
Toennies, G., and Callan, T. P., J . Biol. Chem., 125, 259 (1938). Toennies. G.. and Kolb, J. J., Ibid.. 144, 219 (1942). Tomifiek, O., Collection Czmhoslou. Cheni. Communs., 13, 116 (1948).
Wagner, C. D., Brown, R . H . , and Peters, E. D., J . A m . Chem. Soc., 69, 2609 (1947).
Wilson, H. N., J . Soc. Chem. I n d . (London), 67, 237 (1948). RECEIVEDJune 30, 1950. Presented before the Division of Analytical CHEMICAL SOCIETY, AtChemistry a t the 116th Aleeting of the AMERICAX lantic City, N. J.
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