Effect of Various Alcohols upon Water Analyses on 10-MI. Samples of Vinyl Ethers
Table 111.
Vinyl
Ether Ethyl Ethyl Isopropyl Isopropyl
k%fPtp,? Isobutyl
2-Ethylhexyl 2-Ethylhexyl
Alcohol Added, Gram Ethyl (0.785) Ethyl (3,92) Isopropyl (0.785) Isopropyl (3.92) Isopropyl (3.92) Isobutyl (0,806) Isobutyl (4.03) 2-Ethylhexyl(O.811) 2-Ethylhexyl (4.06)
water, % Present Found
titrations were performed. It is possible that other solvents may prove satisfactory for this analysis. Dioxane, and 2 to 1 v./v. pyridine-ethylene glycol proved to be unsuccessful as solvents. Back-titration of excess Karl Fischer reagent with a standard solution of water in pyridine-acetic acid was not attempted. Table I shows the results obtained on 10-ml. samples of vinyl ethers to which a known amount of water had been added. INTERFERING SUBSTANCES
Alcohols constitute the most likely
0.054
0.054
0.333 0.333 0.333 0.050 0.050
0.091
0.119 0.325 0.339 0.333 0.071
0.085
0.063 0.063
0.108 0.108
Error, %
+ 68.4 +120.0 -
2.4
+ 2.0 0.0 + 42.0 ++ 70.0 71.5
+ 71.5
Titration Error, MI. +0.40 +0.83 -0.08
+0.07 0.0
$0.22 +0.37
&0.50 $0.50
interference to be encountered in this analysis. Therefore, because vinyl ethers may contain alcohols either as impurities or additives, the effects of various alcohols was investigated. Table I1 illustrates the effect of increasing amounts of methanol upon water analyses on ethyl vinyl ether. The results appear to substantiate the theory that interference is probably due to the iodoacetal reaction (1). Owing to steric, and possibly other factors, it might be expected that the alcohol interference in the higher homologs of ether-alcohol systems would be less than that experienced in the lower members. Table 111 shows the results
obtained upon adding various alcohols to their corresponding ethers. The results of Table I11 indicate that, for the alcoholether systems investigated, only isopropyl alcohol produces no appreciable interference. Unfortunately, additional higher homolog samples were not available with which to pursue the subject further. Because the amount of alcohol normally present in rectified vinyl ethers is small, the alcohol interference should cause little, if any, significant error. ACKNOWLEDGMENT
The authors express their appreciation to R. E. Feltham for performing some of the preliminary analyses, to A. H. Taylor and C. A. Wamser for helpful suggestions in preparing this manuscript, and to the Air Reduction Co., Inc., for permission to publish. LITERATURE ClTED
(1) Siggia, S., IND.ENG.CHEM.,ANAL.
ED. 19. 1025-9 (1947). ~- (2jsiggia, S., ddsbeig, R. L., ANAL. CHEM.20, 762-3 (1948). I
RECEIVED for review February 11, 1959. Accepted July 21, 1959.
Selected Triphenylmethane Dyes as Acid-Base indicators in Glacial Acetic Acid ORLAND W. KOLLING' and McCbURE 1. SMITH Deportment of Chemistry and Biology, Friends University, Wichita, Kan.
b Several p-aminophenyl derivatives of the triphenylmethane family of dyes were examined for their suitability as acid-base indicators in glacial acetic acid. The first color changes of malachite green, brilliant green, Victoria Blue, and pararosaniline can be used satisfactorily for the titration of strong and intermediate bases with perchloric acid, Rhodamine B was acceptable in photometric titrations of strong and intermediate bases in which the course of the titration was followed by the change in absorbance of the acid form of the indicator. For empirical acidity function measurements on perchloric acid solutions, the ~CHCIO,range from 2.25 to 5.10 is determinable by the overlapping indicator pairs, brilliant green and auramine, and brilliant green and Rhodamine B.
I
titration of bases with perchloric acid in glacial acetic acid as a solvent, crystal violet and malachite N THE
1876
ANALYTICAL CHEMISTRY
green are the two indicators most frequently employed. These dyes are among the stronger bases of the triphenylmethane series, and cannot be used in the titration of very weak bases. Other members of the triphenylmethane series were evaluated for their suitability as acid-base indicators in glacial acetic acid. Applications investigated were visual detection of the titration end point, empirical acidity function measurements, and photometric titrations. In each application to glacial acetic acid solutions, the neutralization of the indicator follows Equation 1, I
+ RQ e IR+Q- e IR+ + Q -
(1)
where I is the indicator base, RQ the acid, and IR+Q- the indicator salt. If the sum of the acid forms of the indicator is denoted as 11 and the base form as Ig (4, the appropriate quantitative relationship is in Equation 2.
Here Lln is the half-neutralization number of the indicator, CRQthe concentration of the acid, and Yr the ion pair parameter. In a.nalytica1 uses involving perchloric acid the indicator Y r values are constant, therefore pLqn is a satisfactory measure of the base strength of the indicator (6). Likewise, the proportionality between CBQand I& . has been used by h h m and Higuchi (8) in the photometric detection of the titration end point with indicators in glacial acetic acid. An empirical acidity function L may be defined by Equation 3,
L
=
- log CRQ
5
log
IIE] - pLn 1111
(3)
and for structurally related dyes, overlapping indicator range are more probable. In selecting representative triphenyl1
Present address, Depsrtment of Chem-
ietry, Southwestern College, Winfield, Kan.
methane dyes, it was observed that the solubilities of those members of the aeries lacking an amino nitrogen and containing phenolic, sulfonic, or carboxylic acid groups, were too low for precise quantitative treatment (less than mole per liter in glacial acetic acid). However, the following paminophenyl derivatives of this series were found to be usable indicators: auramine, brilliant green, pararosaniline, and Victoria Blue. Rhodamine B was included in the study because of its resemblance to the triphenylmethane dyes in resonance structures (7).
Visual Color Changes of Indicators in Titration of Potassium Acid Phthalate with Perchloric Acid
Dye Brilliant green Malachite green Pararosaniline Victoria Blue
Color Change Blue-green to green Green to yellow-green Yellow-green to yellow Blue-green to green Green to yellow-green Yellow-green to yellow Red to violet Violet to orange Blue to blue-green Blue-green to yellow-green
Structural Formula for Triphenylmethane Dyes Ri
Etr
-NHz
-CH,
-CHI
-CsHs -C,H,-NH,
-CzHs -NH,
--C&
Y
Table 11.
- ER, Titration MV.~ Error, % 174 0.18 205 0.37 233 1.6 235 0.18 256 0.52 288 1.8 229 1.4 268 2.9 202 0.52 ( N o t reprodricibk)
Dye Auramine
Determination of Indicator Half-Neutralization Numbers
Absorption Peak, Mp Acid Base form form 435
-NH2
-CloH6NH~ -CH3 -CHa
Brilliant green
455
630
EXPERIMENTAL
All absorbance measurements and photometric titrations were made using a Bausch & Lomb Spectronic 20 spectrophotometer. Potentiometric measurements were performed with a Beckman Model H-2 pH meter, equipped with the usual glass and calomel electrodes. Reagents. Commercial glacial acetic acid (Mallinckrodt A.R. grade) was used without additional purification, as the water content of this solvent was found t o be 0.008% by Karl Fischer titration. Less than 0.4% water in the solvent is without effect upon potentiometric meaaureand reduction of the water menta content from 0.008% to less than 0.004% by treatment with acetic anhydride had no detectable influence on the absorbance values of indicator solutions. Solutions of 0.09M perchloric acid in glacial acetic acid were prepared and etsndardized as previously reported (3). These solutions were used to standardize 0.1M solutions of the bases d i u acetate, o-chloroaniline, and urea by potentiometric titration. The 0.100M potassium ac?d phthalate and sodium salicylate solutions were prepared determinately. The dyes, pararosaniline (acetate), malachite green (chloride), brilliant green (chloride), Victoria Blue (chloride), auramine (chloride), and Rhodamine B (chloride), were obtained from Fisher Scientific Co., and were used without additional purification. Glacial acetic acid stock solutions were 0.001M in the indicator. Apparatus.
(a,
Procedures.
E
a E is the measured e.m.f. for the indicator solution and E R is that for the reference solution.
Y
Dye Auramine Brilliant green PSrarOeaniline Victoria Blue
Table 1.
Visual color changes
of the indicators were observed in the titration of 20-ml. portions of 0.025M
Pararosaniline
420
550
Rhodamine B
470
555
potassium acid phthalate (0.20 ml. of indicator stock solution was added) with perchloric acid. At each visual end point, the potential difference of the glass-calomel electrode pair was recorded, and- the e.m.f. for a 0.1M sodium salicylate reference solution was measured before and after each indicator solution. Four determinations were made on each indicator end point, with an over-all precision of f 1.5 mv. The absorption spectra for the acid and base forms of each indicator were obtained from solutions that were 0.025M in perchloric acid and potassium acid phthalate, respectively, and a p proximately 10-bM in indicator. Beer's law plots were made for the acid and base forms in solution concentrations ranging from 1 O - W to 104M. T o determine the indicator constants, 5.00 ml. of 0.05M potassium acid phthalate was placed in a 50-ml. volumetric flask and 0.200 ml. of indicator stock solution was added; then the
Mean
PCHCIOC 0.398 0.0755 -0.0325 -0.845 -0.0595 -0.556 -0.793 -0.913 -0.0306 -0.342 -0.535 -0.726 0.343 0.0899 -0.0223 -0.542
3.80 3.48 3.38 2.58 4.46 3.98 3.76 3.64 4.16 3.90 3.61 3.48 3.61 3.38 3.23 2.75
PLLIL -3.40 (f0.002) -4.53 ( f O ,011
-4.20 (f0.03) -3.27 (*0.01)
contents were titrated with acid to a color change, and diluted to volume.
The absorbance values of the indicator acid and base forms were measured on their respective peaks, with the exception of auramine whose acid peak is not in the visible region. A minimum of 16 determinations were made for each indicator.
In the photometric titrations 5.00 ml. of 0.1M base and 25.0 ml. of solvent were placed in a 25 X 150 mm. test tube in the spectrophotometer. The instrument waa adjusted to zero absorbance a t the absorption maximum for the acid form of the dye, and 0.20 ml. of indicator stock solution was added. The standard perchloric acid titrant waa added in 1-ml. portions until the end point was approached, then in 0.1ml. incrementa for the remainder of the titration. After being corrected for dilution, the measured absorbance value and that for the pure acid form were used to calculate the ratio of IJIs from Equation 4: VOL 31, NO. 11, NOVEMBER 1959
1877
where A . is the absorbance of the acid form and A , that of the pure acid form. The graphic titration end point was determined from the intersection of the extrapolated linear portions for the function I d , us. milliliters of HClO,. Four titrations of each base were performed with each indicator. Titrations were made rapidly to minimize the heating of the cell contents by the instrument. RESULTS AND DISCUSSION
Visual Color Changes. The potential differences for the glastwxlomel electrode pair that correspond t o each indicator color change are listed in Table I. These e.m.f. values are expressed as differences between the measured potential and that of a standard reference solution, in order t o circumvent any variability among glass electrodes. Malachite green is included in this indicator group, as Comparable data on this dye have not been reported previously. However, auramine and Rhodamine B are excluded, because of their unabrupt visual color changes. The suitability of a given indicator can be determined not only by the experimental titration error as recorded in Table I, but also by comparison of the quantity ( E - EB) to the equivalence point potentials of representative bases titrated with perchloric acid. This potential difference for sodium acetate (a strong base) is 178 my., and that for urea (a very weak bas4 is 227 mv., when measured against the same reference solution as used above. For the titration of bases having strengths between potassium acetate and urea, only the first color change for any of the indicators in Table I should be used. Malachite green and brilliant green are the better indicators of this group, although Victoria Blue and pararosaniline are acceptable when their higher indicator blank corrections are applied. Indicator Constants. The conventional method for maintaining fixed ionic strength by the addition of inert salts t o aqueous solutions is severely limited when applied t o nonaqueous solvents of low dielectric constant. The extensive solvolysis of salts into the conjugate acid and base of differing strength, and the tendency toward ion pair (and ion multiple) association in preference to ion separation in solutions of the salt, leaves little real meaning for the term, inert salt, when applied to glacial acetic acid solutions. Likewise, the calculatcd ionic strength is devoid of meaning for systems involving two or more salts undergoing simultaneous ion separation equilibria, as is the case in a solution 1878
0
ANALYTICAL CHEMISTRY
Table 111. Photometric Titrations of Bases with Perchloric Acid Relative Error, P.P.T. 0-
Dye Brilliantgreen Pararosaniline Rhodamine B VictoriaBlue
Sodium Chloroacetate aniline Urea 3.1 -1.7 -790 3.8 0.1 -110 3.0
-0.8
3.4
0.01
- 53 - 81
containing both a simple salt and an indicator salt. To overcome these difficulties, constant activity coefficients for the indicator components of the equilibrium in Equation 1 were maintained bv saturating the solution with potassium perchlorate. The molar solubility of potassium perchlorate is 1.95 x 10-4 at 25' C. (9). Measurements used in the calculation of indicator half-ncutrslization numbers were made in the presence of an excess of this salt. The concomitant ion separation equilibria for the two perchlorate salts results in a t least an approximately invariant ionic environment in these solutions. Also, the close structural similarity of the Ayes used in this study made it appear likely that all of the indicator perchlorate salts have ion separation equilibrium constants of the same order of magnitude. In Table I1 representative data used in the evaluation of indicator half-neutralisation numbers are summarized. A plot of log Is/IA US. pCacio4 gives a straight line having a slope 1.00 It 0.01 for each indicator, thereby confirming the conclusion of Kolthoff and Bruckenstein (6) that the dissociation constant for the indicator perchlorate is equal to the over-all dissociation constant for perchloric acid. The simplified treatment in which each indicator base reacts with only a single equivalent of acid, according to Equation 1, is experimentally justified for low concentrations of perchloric acid. For Rhodamine B, a b sorbance measurements were made a t 470 rather than on the acid peak a t 495 mp, owing to the alight overlap of the absorption band of the base form. The five indicators listed in Table I1 were applied to the empirical acidity scale measurement defined by Equation 3. In perchloric acid solutions in glacial acetic acid, the usable range for log Is/II was 0.5 to -1.0, and outside of this range the experimental values showed considerable scattering from the linearity predicted by the defining equation, regardless of the indicator selected. All five dyes have sufficiently close values for p L n to permit their use as a series of overlapping indicators, covering the range of L values from 2.25 to 5.10. It is possible to measure L values in this
range by using only two indicators. TWOpairs suitable for this purpose are brilliant green and auramine, and brilliant green and Rhodamine B. A comparison of the values of pLln for structurally comparable triphenylmethane dyes in Table I1 shows that the introduction of the alkyl group into the aromatic amine increases the basicity of the indicator, while the substitution of a second aromatic nucleus on the amino nitrogen decreases basicity. Photometric Titrations. Bases having representative strengths, including sodium acetate (strong), ochloroaniline (intermediate), and urea ( v e q weak), were titrated photometrically with each of the indicators listed in Table 111. The absorption maximum for the acid form of each indicator was used to follow the course of the titration, as these indicators exhibit a longer range of ideal response to acidity in the acid form then in the base form. For auramine, the strong overlap of the absorption bands of the acid and basc forras made it impossible t o obtain quantitative titrations for any of the bases. Several methods for the determination of the equivalence point in photometric titrations have been proposed ( I , 9,8). The most general titration graphs derived by Goddu and Hume ( I ) are analogous to curves obtained in conductometric and amperometric titrations, where the end point is the intersection of the extrapolated linear portions of the curve. TI& method was utilized in the titration results in Table 111. The major applications of photometric titrations to glacial acetic acid solutions involve the titration of bases of weak or intermediate strengths with a strong acid. The plot of IdIB us. volume of titrant yields a hyperbolic function corresponding t o the curve of two weak bases, in which K , for the colorless baae is 100 to lo00 times greater than Rb of the indicator base. Only in the titration of bases as strong as sodium acetate is there no detectable curvature in the immediate vicinity of the end point. The alternate graphical method of Higuchi, Rehm, and Barnstein @),in which the curve I B / I ~08. 1 per d. of titrant is extrapolated to the z-intercept, waa capable of high precision and acceptable accuracy in the titration of sodium acetate, using malachite green indicator. However, the application of this method to the titration of urea and o-chloroaniline gave precise photometric results that were much lower than potentiometric titrations. In the original work of Higuchi, Rehm, and Barnstein (8) there is no mention of a comparison of their values from photometric titratione with results obtained by an independent method. The mean relative errors in the titration of 0.5-mmole samples of the baeea
are listed in Table 111. All four dyes are suitable indicators for the photometric titration of strong and intermediate bases, as the relative error values fall within the probable error range of k4.0 p.p.t. for these determinations. Urea behaves as a base that is slightly weaker than the indicators used in the titration. Because the dyes included in this study extend from the strongest to the weakest bases of the triphenylmethane dyes, it is doubtful that any dye Of fsmily malachite green) is suitable for indicat-
ing the exact equivalence point in the titration of urea with perchloric acid. ACKNOWLEDGMENT
m e authors express their indebtedne88 to Research Corporation for financia1 support of this investigation. LITERATURE CITED
(1) @'ddU, R., Hume, D., A N A L -
26, 1681 (1954). (2) Higuchi, T., Rehm, C., Barnskin, C., M . ,28, 1509 (1956).
(3) Kolling, 0. W., J . Am. Chem. Sor. 79, 2717 (1955). (4) Kolling, 0. W., J . Chem. Educ. 35,452 (19%): (5) Kolling, 0. W., Trans. Kansas h a d . Sa'.59,422 (1956). (6) Kolthoff, I., Bmckenskirl, s., J . A m . Chem. Soc. 7 8 , 7 (1956). (7) Rsmeth, R., sandell, E., ibid,, 78, 4872 (1956). (8) Rehm, C., Higuchi, T., ANAL.CHEM. 29,367 (1957). (9) Seward, R., Hamblet, C., J . A m . Chent. Soc. 54,557 (1932).
RECEIVEDfor review January 12, 1959. Accepted .4ugust 5, 1959.
Titrimetric Determination of 2-Mercaptoacetic (or Thioglycolic) Acid by Copper(l1) SUSEELA B. SANT' Carr Chemical laboratories, Mount Holyoke College, South Hadley, Mass. BHARAT R. SANTl Codes Chemical laboratories, Louisiana State University, Baton Rouge, la.
b A mettrod for the determination of 2-mercaptoacetic (or thioglycolic) acid in aqueous solution is based on the titration of a standard cupric salt solution with thioglycolic acid. At the end point a permanent yellow precipitate is formed. The method is accurate within 0.3%.
D
of mercaptans by reaction with cupric alkyl phthalate is rapid and accurate although somewhat less accurate than the iodine method (2, 7, 9, 10). This procedure is effective in the presence of many substances which would otherwise interfere in the iodine method. The reaction between copper(I1) and mercaptan takes place as follows. 2Cu++ 4RSH + 2CuSR RSSR 4H+ (1) ETERMINATION
+
+ +
The method is of general applicability. Titrations are carried out in a nonaqueous m e d i p , usuaily Pentawl or a hydrocarbon solvent. The appearance of the green color of cupric ion marks the end point. However, thioglycolic acid cannot be determined by this method, because the deep blue cupriccuprous mercaptide produced obscures the end point (2). In the c o m of these investigations, thioglycolic acid was found to undergo quantitative reaction with cupric ion in aqueous solutions. Experimental con'Present addresa, Department of Chemiatry, University of Toronto, Toronto, Ontario, Canada.
ditions for a direct titration method using a standard cupric salt solution for the estimation of thioglycolic acid are described. REAGENTS
Copper sulfate (J. T. Baker analyzed) was used to prepare a 0.1N solution in water. An aqueous solution of thiowas glycolic acid, HS-CHxCOOH, prepared by dissolving a 98 to 99% pure sample. Its mercaptan content was determined by oxidation with a known e x c w of iodine solution, followed by titration of the excess with standard sodium thiosulfate (6):
+
~HS-CHICOOH Is -C (S--CHzCOOH)r
+ 2HI
PROCEDURE
A measured portion of the copper sulfate solution was transferred to a 150ml. Erlenmeyer flask or, preferably, a porcelain dish. The white background of the latter is convenient for discerning the correct end point. The thioglycolic acid solution was added from a buret until the deep violet precipitate first formed changed to a permanent yellow. Titrations near the end are performed slowly with vigorous shaking or stirring of the solution. RESULTS AND DISCUSSION
The reactions taking place during the titration are: complexstion of copper with thioglycolic acid, reduction of copper(I1) to copper(1) with simultaneous oxidation of thioglycolic to dithioglycolic acid and, at the end point,
Table 1. Titrimetric Determination of Thioglycolic Acid by Copper(l1) Thioglycolic Acid, Mg. Iodine Copper(I1) method method Difference 15.91 0.02 15.89 0.04 31.78 31.82 45.82 0.15 45.67 0.22 * 91.87 91.65 137.7 137.5 0.2 183.2 183.3 0.1 228.8 229.1 0.3
formation of yellow cuprous mercaptide. The over-all process represented by the following equation
+
~ C U + + ~HSCHICOOH --c 2CuSCH&OOH (S--CHzCOOH)* 4H'
+ +
(2)
is the .same as that of Equation 1. I t follows from Equation 2 that 1 ml. of a 1M cupric sulfate solution is equivalent to 92.12 mg. of thioglycolic acid. Experimental results are given in Table I. Cupric nitrate, chloride, or acetate may be used instead of cupric sulfate. A pH of 4.0 to 4.4, which is the case with aqueous solutions of cupric salts except the acetate, is best suited for the titration. At higher pH valuea the reaction is faster, but the end point is not stable because of the oxidation of cuprous mercaptide. At too low a pH the reaction is slow. When cupric acetste is used, the desired pH m y be obtained readily by the addition of a reqVOL 31, NO. 11, NOVEMBER 1959
0
1879
