sition a t lower p H levels, with complete breakdown resulting from refluxing with mineral acid. Aliquots of the various solutions were tested for their content of free isonicotinic acid, the breakdown product, using essentially the method of Rubin and associates (4) based on the bromocyanogen reaction of Mueller and Fox ( 3 ) . With decreasing assay values for iproniazid, concomitantly increasing quantities of isonicotinic acid were found. Other aliquots of the solutions refluxed for 10 hours were subjected to paper chromatography, using butyl alcohol saturated with water as the mobile phase. The spots were developed in a chamber containing crystals of cyanogen bromide and the paper was then sprayed with a 1% alcoholic benzidine solution. Reddish spots, with an R, value of 0.16 identical with those of isonicotinic acid run in a parallel experiment, were obtained. Solution 1 showed large quantities of isonicotinic acid, solution 2 showed a distinct band of much less intensity, while solution 3 indicated only faint traces of isonicotinic acid,
In a corresponding experiment, samples of iproniazid sirup, buffered a t various pH levels, were stored for 2 months a t room temperature and a t 45” C. While no loss was found with the colorimetric method in any of the samples stored a t room temperature, the samples stored at 45” C. showed the following losses in potency: Per Cent Loss 1 month 2 months
Sample pH 4
4
5
1
Xone
6
14 2
None
If the terminal nitrogen is disubstituted, no color is produced under the conditions of the test. While the color reaction can a t this time only be conjectured as a molybdate complex formation, the mechanism remains to be elucidated. ACKNOWLEDGMENT
The authors would like to acknonledge the helpful suggestions of B. Z. Senkowski and J. A. Kapoli. Esther Critelli prepared the drawings.
Application to Related Compounds.
d series of analogs was subjected t o the colorimetric reaction a t the usual concentration of 0.2 mg. of base per ml. The results are listed in Table IV, in order of their decreasing absorbance. It would appear t h a t compounds of the following general structure will produce such a color: 0
//
R’-C-XHSHC-R where R = alkyl or aryl and R’ = pyridine.
LITERATURE CITED
(1) Bernhart, D. S.,m’reath, A. R., ANAL. CHEW27, 440 (1955). ( 2 ) Deltombe. J., J. Dharm. Belo. 8, 59 I
,
.
(1953). (3) XIueller, A,, Fox, S. H., J . Am. Phartn. Assoc., Sci. Ed. 40, 513 (1951). (4) Rubin, 8. H., Drekter, L., Scheiner, J., DeRitter, E., Dzseases of Chest 21, 439 (1952). ( 5 ) Vignoli, L., Cristau, B., Pfister, A., Chim. anal. 38, 392 (1956). (6) Wollenberg, O., Klin. Tt’ochschr. 30, 906 (1952).
RECEIVEDfor review June 12: 1957. Accepted September 14, 1957.
Coulometric Titrations with Mercury(1 and II) Determination of Cyanide EDWIN P. PRZYBYLOWlCZl and
L. B. ROGERS
Department o f Chemistry and laboratory for Nuclear Science, Massachusetts Institute of Technology, Cambridge
b The coulometric generation of mercury(l1) ion a t I O O ~ ocurrent efficiency has been successfully carried out in alkaline cyanide solutions. In solutions buffered a t pH 9 the titration was accurate to within 1% a t concentrations of cyanide of about 2 x 10-T. A comparison of generated silver ion and generated mercury(l1) ion as titrants for cyanide under identical conditions showed mercury slightly superior to silver for small amounts of cyan ide.
T
HE most widely used volumetric method for the determination of cyanide is that of Liebig ( I s ) ,which depends on formation of the dicyanatoargentate(1) complex. This titration was applied to the determination of small amounts of cyanide by Laitinen, Jennings, and Parks (11), who titrated
Present address, Research Laboratories, Eastman Kodak Co., Rochester 4,
x. Y.
cyanide in 0.18’sodium hydroxide using an amperometric end point. At concentrations between 0.1 and 0.001F cyanide the method was accurate to within 0.27,; a t lower concentrations the accuracy decreased, but a 4 X lO-5F solution could be titrated within 2.3y0 accuracy. An end point was found for 8 x 10-68’ cyanide. The present study using niercury(I1) seemed advisable because comparison of the dissociation constants of the complexes Ag(CN)?- and Hg(CN)2, whose values (12, 16) are 1.38 X 10-19 and 1.26 x 10-a4, respectively, indicates that a titration using mercury(I1) should be somewhat more sensitive for small amounts of cyanide than one employing silver. Several others (1, 4, 19, 20) have studied mercury(I1) cyanide complexes and reported their formation constants. The most recent work was done by Sewman (16), who reported formation constants for the 2-, 3-, and 4-cyanide complexes with mercury(I1) as 1031.9,1038.1, and 1040.6, respectively,
39, Mass.
a t an ionic strength of 2.0. Because the values of these three constants are bunched closely together, in a titration of cyanide with mercury(I1) these complexes will coexist to varying degrees. Experimentally a potentiometric break i; observed a t a 2 to 1 cyanide-mercury ratio, with no evidence of potential breaks a t ratios of 4 to 1 , 3 t o 1, or 1 to 1. The volumetric titration of cyanide with mercury(I1) chloride in alkaline solution was first reported by Hannay (8),who detected the end point by the formation of a precipitate of the oxide or a basic ammono salt. Thompson (21) and Wick (25) later proved that Hannay’s results were 3% high, becmse of the methods employed for indicating the equivalence point. These authors and others ( I O , 22) have shown that for 0.18’ solutions of cyanide, the potentiometric method gives results accurate to better than 0.2%. Kolthoff and Verzijl (IO) found that the change in potential a t the equivalence point was greater when mercuric perchlorate or VOL. 30,
NO. 1, JANUARY 1958
65
mercuric nitrate was used. These reagents, however, required the presence of acid t o minimize the hydrolysis of mercury(I1). As the acid led to losses of hydrogen cyanide, they preferred mercury(I1) chloride as a titrant. Votocek and Kotrba (22) used acidified rnercury(I1) nitrate as the reagent, but used back-titration to minimize the loss of hydrogen cyanide. Muller and Aarflot (15) attempted to titrate cyanide directly with mercury(1) perchlorate; a rapid quantitative decomposition of mercury(1) cyanide took place with resulting formation of free mercury and niercury(I1) cyanide. Therefore, it is evident that in the presence of cyanide ion essentially all the mercury electrolytically dissolved goes into solution as mercury(I1) ion. Experimentally, the fact that mercury(I1) rather than mercury(1) was formed in the titration was borne out by the absence of a visible precipitate during a cyanide titration. I n the absence of cyanide, either mercury(1) or mercury(I1) can be generated, depending upon pH. In an acid solution mercury(I) is generated; in alkaline solution mercury(I1) is generated, with the resulting formation of hydrated mercury(11) oxide. This is consistent with the findings of Kolthoff and Miller (9), who studied polarographically the anodic wave of cyanide and reported that the reaction with mercury was reversible in 0.1F sodium hydroxide. The present study was undertaken with the idea of extending the use of mercury(I1) titrants to alkaline solution, and providing a more sensitive procedure for small amounts of cyanide than the evisting procedure with silver. Because &hemercury indicator electrode is sensitive to changes of pH and hydrogen cyanide has a pK of 9.39 ( 7 ) , the potentiometric titration curve of cyanide with mercury(I1) should be pH dependent. For that reason, it was necessary to determine optimum conditions under which the titration should be made as well as to examine the usual operating variables. EXPERIMENTAL
Solutions. Deciformal potassium cyanide solution was prepared by weighing out 6.5 grams of reagent grade salt and dissolving it in 1 liter of distilled water. The solution was standardized by titration with standard silver nitrate. This stock solution was diluted with distilled water that had been deaerated for 15 minutes with prepurified nitrogen in order to minimize oxidation of cyanide to cyanate (12). A standard solution of silver nitrate was prepared by weighing out a known amount of the reagent grade salt, dried a t 110” C. for 2.0 hours. This solution was employed to standardize the solution of cyanide, using both a visual end point and a potentio-
66
0
ANALYTICAL CHEMISTRY
metric end point obtained with a silvertrode recorded. The resulting potentiocalomel system. Under three different metric curve was plotted and the inflectitration conditions-in the presence of tion point determined by a graphical (1) 3.0 ml. of ammonia and 0.1 gram of method (14) commonly used in polaropotassium iodide, ( 2 ) 3.0 ml. of amgraphic analysis for determination of a monia, and (3) 0.53’ tribasic phosphatehalf-wave potential. the results had an over-all coefficient of PRETITRATION METHOD. To the variance of 0.05 with no significant diftitration cell was added 90 ml. of 0.5F ference between visual and potentiodibasic sodium phosphate solution, folmetric end points. lowing which the solution ~ 7 a deaerated s A buffer solution prepared by diswith prepurified nitrogen for 10 minutes, solving 71 grams of dibasic sodium phosA 0.001F potassium cyanide solution phate in distilled water and diluting to was added dropn-ise until the potential liter (0.5F) had a pH of 9.05. A similar of the mercury indicator electrode was 0.53‘ solution prepared by dissolving -0.05 volt us. S.C.E. In general, one 190 grams of sodium borate decahydrate drop (about 0.05 ml.) was sufficient. in distilled water and diluting to 1 If slightly too much cyanide was added, liter had a pH of 9.21. Addition of the potential was readjusted by gener15 ml. of 1M sodium hydroxide to 85 ating a small amount of mercury(I1). ml. of either buffered solution resulted Once the electrode potential had been in solutions of pH 11.0. adjusted t o the proper 1-oltage and reII’ith unbuffered solutions, 0.53’ somained stable, a sample to be analyzed dium perchlorate was employed. It was added by either a calibrated 1.000was prepared by dissolving 17 grams of ml. hypodermic syringe or a 10.00-ml. sodium perchlorate monohydrate in disvolumetric pipet and the solution was tilled water and diluting to 1 liter. titrated back to the same potential, Apparatus. The coulometric apNitrogen was passed over the solution paratus and the titration’ cell have during the titration. been described (18). I n addition t o In the absence of added cyanide, the the usual stationary indicator elecpotential observed for the phosphate trode used in potentiometric titrabuffer was always more noble than tions, a rotating electrode was used for -0.0500 volt us. S.C.E. Thus, the amperometric studies. This consisted solutions had what might be called a of a rotating mercury-coated gold wire “negative blank.” Unless the buffer (0.15 X 0.4 em.) sealed into a 3-inch was pretreated d t h cyanide before polystyrene rod perpendicular to the addition of a sample, the results of axis of rotation. cyanide determinations were lo^ by 3 Procedures. DIRECT TITRATION. to 5%. For studies in which the potassium cyanide solutions were directly tiRESULTS trated, 10 ml. of a 0.1F sodium hydroxide solution was added to 90 ml. Pretitration Method. Aqueous of the sample. This solution was then samples containing between 0.032 deaerated with pre-purified nitrogen and 2.6 mg. of cyanide were analyzed for 10 minutes without significant loss of using generating currents from 0.50 hydrogen cyanide. Both indicator and t o 50 ma. By titrating these solugenerator electrodes xere precathodized tions to the end point potential, the in this solution a t 30 ma. for 60 seconds typical results summarized in Table I each. Dibasic sodium phosphate (7.1 were obtained. At the lowest current, grams) was added to lower the pH to the slope of the potentiometric curve a t 9.2 and generation of mercury(I1) begun. During the titration, nitrogen the inflection point was 0.9 mv. per was passed over but not through the second; a t a generating current of 50.00 solution. because direct deaeration of ma., a black substance [apparently 0.5F dibasic sodium phosphate solumercury(I1) oxide] formed on the tion containing potassium cyanide had generator electrode. When the generabeen found to cause a 2 to 3% loss of tion of reagent was discontinued before cyanide in 10 minutes. the equivalence point. the black maThe generation of mercury(I1) reaterial slowly dissolved during 2 to 3 gent was interrupted from time to time minutes, following which the titration and the potential of the indicator elec-
Table I. Titrations of Cyanide in 100 MI. of Pretitrated 0.5F Dibasic Sodium Phosphate
(Mercury-coated gold electrode as genwator electrode) Std. Dev., Av. Amount Added,a Current, Elapsed No. of Amount Mg. Ma. Time, Sec. Trials Found, M g . Mg. Error, yo 2.550 1.275 0.6375 0.6375 0.0638 0.0638 0.0319
50.95 20.00 20.00 10.01 1.084 0.500 0.500
185.7 236.5 118.5 236.4 217.6 459.0 232.4
5 5
5 5 5 3 5
2.550 1.276 0.6388 0.6380 0.0635 0.0624 0.0313
0.004 0.0007 0.0011 0.0004 0.0005 0.0006 0.0004
+0.03 f0.08 $0.20 i-0.09 -0.48 -2.20 -1.90
A proximately 0.0013 mg. of cyanide in 100 ml. of eolution ( 10-6nl) waa present pnor to adition of sample.
-307
will then be accurate to within 3%. On the other hand, in the pretitration procedure the sample is diluted tenfold, but the titration can be carried out with an error of only about 2% of the amount added. Thus the pretitration method is to be preferred, not only because of its somewhat better accuracy but also because the direct method is more timeconsuming, as it requires a plot of the complete potentiometric curve.
lh
-20
,
[ OO
Figure
7 1
1.0 2.0 C N - ( F x I09 in Titrating Medium
1.
End
point
-06 -
error
4 0
6 PH 8
4 10
B. C.
1 L4
100% titrated 150y0 titrated
D. Calculated curve for HgO
was completed in the usual way. At generating currents less than 50 ma. this phenomenon was not observed. The method appeared t o be more accurate in very dilute solutions than the argentimetric procedure reported by Laitinen, Jennings, and Parks (11), as concentrations of 1 X 10-V' cyanide could be determined with a 1.9% error using mercury, nrhereas a 2.3% error was reported for a 4 x 10-517 solution using silver. Direct Method. The consistently low results obtained by direct titration appeared to be due in part to the fact that the inflection point does not correspond t o the equivalence point in the cyanide titration with mercury(I1). The inflection point, however, is the only point on a potentiometric curve which can be related
40
easily and accurately t o the concentration of cyanide present in the sample especially when the curves are somewhat flat. If the inflection point were used as the end point in a direct titration of cyanide, the resulting error depended on the amount of cyanide present, as shown in Figure 1. An empirical calibration curve could be prepared which would relate the inflection point to the amount of cyanide present, but this was not entirely satisfactory. Results obtained by the direct titration (sample not diluted before analysis) and the pretitration methods (sample diluted) can be compared by referring to Figure 1. Thus, if a 10-4F solution is to be analyzed by both methods, it can be seen from Figure 1 that a 10% correction must be applied to the data from a direct titration. The corrected value
60
80
100
Time ( Seconds1
Figure 3.
I2
2 X 1O-V' potassium cyanide A . 50% titrated
using inflection point as end point
20
2I
Figure 2. Potential-pH diagram for mercuric cyanide system
As a function of the amount of cyanide titrated in direct method
0
DISCUSSION
Cathodic polarization curves
Constant 100-pu. current and 1-sq.cm. mercury-coated electrode A . After solution had been deaerated with ureuurified nitroeen . ., for 10 minutes B. After solution had been purged with prepurified nitrogen for 15 minutes and mercury electrode cathodized for 5 minutes at 1 ma.
Preliminary Studies. The usefulness of pH-potential diagrams for predicting the nature of titration curves has been pointed out (3, 6, 6, 17). The characteristics of the cyanide system, shown in Figure 2, were constructed by combining published data for equilibrium constants with experimental data. To establish optimum conditions with relatively few experiments, data were obtained on three solutions which represented a 2 x l0-V' solution of potassium cyanide which was 50, 100, and l50%, respectively titrated with mercury(I1) nitrate. The pH of each solution was increased by adding appropriate amounts of 0.1F or 1.OF sodium hydroxide. The difference bet\$-een curves A and C represents the approximate potential break a t the equivalence point in the titration of lO+F cyanide with mercury(I1). The optimum pH appears to be close to 9.0, because a t lower pH values hydrogen cyanide losses were substantial, whereas a t higher pH the formation of mercury(1I:l oxide decreased the break. -4s one mould expect, addition of a complexing agent such as ammonia or acetone to the solution to prevent the formation of the oxide decreased the limiting potential even more. From the stability constants for the Hg(CS)2 complex and the dissociation constants for hydrogen cyanide, one can calculate that the equivalence point potential should change 10 mv. for every tenfold change in the initial concentration of cyanide. This factor did not introduce a significant error into determinations carried out by the pretitration method, though it could in the direct method. Effect of Oxygen. The presence of oxygen in the solution could cause the consumption of cyanide by formation of cyanate or by reaction with metallic mercury t o form mercury(I1) oxide, which, in turn, could react with cyanide to form Hg(CN)2 plus hydroxide. By using the coulometric technique developed by Campbell and Thomas (2) to study fi!ms on metals, it was possible to show that the negative blank found in direct titration was caused by dissolved oxygen. TT71~ena mercury electrode was placed in an alkaline solution (pH 9 or greater) and the solution was VOL. 30, NO. 1, JANUARY 1958
67
thoroughly deaerated with prepurified nitrogen for 10 minutes, cathodic polarization of the electrode indicated the presence of a small amount of oxide on the surface (Figure 3). Saturating the solution with air produced a much larger blank. On the other hand, when the electrode was cathodized for 5 minutes a t 1 ma. while the solution was being deaerated, the cathodic polarization curve of the electrode showed an almost negligible amount of oxide film on the electrode. I n addition to decreasing the blank due t o oxide film, the combination of deaeration and cathodization left only extremely small amounts of oxygen in solution. Though the negative blank could be decreased by a combination of deaeration and cathodization of the electrode, it could not entirely be eliminated. This indicated that the technique did not differentiate between the reduction of dissolved oxygen and the reduction of mercury(I1) oxide. Comparison of Amperometric and Potentiometric End Points. The potentiometric and amperometric end points were compared by using lO-5F solutions of cyanide. The amperometric procedure used a rotating mercury-coated gold electrode on which a potential a t 0.0 volt us. S.C.E. was applied, to follow the anodic diffusion current of cyanide. I n addition to considerable curvature in the vicinity of the equivalence point (Figure 4), the amperometric data had a negative bias of 4 to 7%, which may have been due in part to consumption of cyanide ion a t the indicator electrode. (Error from such a source could be minimized by connecting the indicator circuit only during the few seconds required to obtain a current measurement ) Thus, in spite of the fact that the potentiometric titration curve had a very shallonbreak. the potentiometric pretitration method appeared to give more accurate results than direct amperometric titration. Effects of pH and Buffers. Titrations were carried out a t p H 9 and 11 with both borate and phosphate buffers. There was no apparent difference between titration curves in borate and phosphate buffers a t a given p H value. However (Figure 2), a titration a t pH 11 gave a smaller break a t the end point than one a t p H 9. It was thought that by titrating cyanide in a n unbuffered medium it might be possible to obtain a larger jump in potential at the equivalence point because of the sensitivity of the mercury electrode to hydrogen ion as well as to mercury(I1). This was found true a t levels of lO+F or higher; however, the advantage decreased in somewhat more dilute solutions, as shown by the fact that a 5 X 10-6F solution of cyanide
68 *
ANALYTICAL CHEMISTRY
I
a
l -
7 I
t
\
+0.1-
1, Point
W
0.0-
-
--I>n -O.I,7 2
-
L
-
z-0.2-
I
20
0
I 40
I I 100 120 Time (seconds)
I
I
60
80
I 140
1 160
1 180
Io 200
Figure 4. Amperometric and potentiometric end points 0.0288 mg. of cyanide (1.15 X 10-5F) in pH 9.0 buffer solution using a mercury-coated gold electrode and generating current of 1.096 ma. 0 Amperometric 0 Potentiometric
I
‘“‘-1
I
cp
-03t L_I__L
0
20
40
60
A
I
I
’
80
lo0
I20 Time (seconds)
140
I60
180
200
Figure 5. Titrations of 0.288 1 rng. of cyanide (1.1 5 X 1 O-4F) In pH buffer solution generating current of 10.40 ma. 0Silver electrodes 0 Mercury-coated gold electrodes
gave about the same size of break in buffered and unbuffered solutions. Comparison of Mercury and Silver Titrations. I n solutions buffered a t p H 9 the magnitudes of t h e potential breaks a t the equivalence points are about the same for silver and mercury (Figure 5 ) , probably as a result of the greater solubility of silver oxide. The change in slope is somen-hat steeper for the mercury curve, which leads to somewhat better sensitivity and accuracy in analyzing very dilute solutions. ACKNOWLEDGMENT
One of the authors (E.P.P.) wishes to thank the Eastman Kodak Co. for a fellowship; the other, the U. S. Atomic Energy Commission, for partial support. LITERATURE CITED
(1) Brigando, J., Job, P., Compt. rend.
222,1297 (1946).
(2) Campbell, 11‘. E., Thomas, U. B.,
Trans. Electrochem. SOC.76, 303 (1939). (3) Charlot, G., “ThBorie et MBthode Nouvelle d’Analyse des Reactions en Solution,” 3rd ed., Masson, Paris, 1949. Charlot, G., Gaugin, R., “M6thodes d’Analyse des Reactions en Solution,” Masson et Cie, Paris, 1951. Clark, W. M., L‘Determinationof Hydrogen Ions,” 3rd ed., p. 287, Williams & Williams, Baltimore, 1928. Delahay, P., Pourbaix, M., van Rysselberghe, P., J. Chem. Educ. 27, 683 (1950).
Gregg, E. C., Tyler, W.P., J . Am. Chenz. SOC.72, 4661 (1950). Hannay, J. B., J . Chem. SOC.,33,’ 245 (1878). Kolthoff, I. M., Miller, C. S., J . Anz. Chem. SOC.63, 1405 (1941). Kolthoff, I. M., Veraijl, E. J. A . H., Rec. trav. chim. 42, 1055 (1923). Laitinen, H. A., Jennings, W. P., Parks, T. D., IND.ENG. CHEM., ANAL.ED. 18, 574 (1946). Latimer, W. SI., “Oxidation Potentials,” 2nd ed., Prentice-Hall, New York, 1952. Liebig, J., Ann. 77, 102 (1851). Meites, L., “Polarographic Tech-
niques,” p. 60, Interscience, New York, 1955. Muller, E., Aarflot, H., Rec. trav. chim. 43, 874 (1924). h’ewman, L., Ph.D. thesis, Massachusetts Institute of Technology, March 1956. Pourbaix, M., “Thermodynamics of Aqueous Dilute Solutions,”English tr. by J. K.Bgar, Arnold, London, 1949.
(18) Przybylowicz, E. P., Rogers, L. B., ANAL.CHEM.28, 799 (1956). (19) Sherrill, M. S., 2. phys. Chem. 43, 705 (1903). ( 2 0 ) Ibid., 47,103 (1904). (21) Thompson, M. R., Bur. Standards J. Research 6, 1051 (1931). (2’2) Votocek, E., Kotrba, J., Collection Czechoslcv. Chem. Communs. 1, 165 (1929).
(23) Wick, R. AI., Bur. Standards J Research 7, 913 (1931). RECEIVEDfor review May 22, 1957. Accepted August 27, 1957. Abstracted from a thefiis presented by Edwin P. Przybylowicz in partial fulfillment of the requirements for the degree of doctor of philosophy at Massachusetts Institute of Technology, September 1956.
Resolution of Acid Mixtures in Nonaqueous Solvents Potentiometric Titration of Dibasic Acids with Quaternary Ammonium Titrants GERALD A. HARLOW and GARRARD E. A. WYLD Shell Development Co., Emeryville, Calif.
b
The influence of the solvent on the resolution of acids by potentiometric titration has been studied. Curves a r e presented for the titration of sulfuric, oxalic, and succinic acids in mixtures of water and isopropyl alcohol, and for sulfuric acid in isopropyl alcohol, ethyl alcohol, methanol, ethyl ether, and pyridine. The manner in which the properties of the solvent may b e utilized in the analysis of acid mixtures is discussed. Mixtures containing acids of different charge types which cannot b e resolved in water can b e readily resolved in nonaqueous solvents.
A
problem in analytical chemistry is the resolution of mixtures containing dibasic acids. -4s such determinations are more readily carried out when the dibasic acid gives tn-o sharp inflections, the factors that influence the resolution of dibasic acids were investigated to determine how they can be varied for the best practical results. This investigation vas greatly aided by the development of methods for the preparation of anhydrous quaternary ammonium hydroxide titrants (6, 8). Many dibasic acids which give only a single inflection in water give two sharp inflections in nonaqueous solvents. Little information is available, however, on the behavior of these acids in iiiixtures of water and another solvent. An attempt has been made here to provide data relating the degree of resolution of some common dibasic acids to the composition of mixtures of water and isopropyl alcohol. These two solvents were chosen for this study because they are miscible in all proportions and the mixtures are good solvents for acids, a wide range of FREQUENT
dielectric constant from 80 to 18 is provided, and the tetra-n-butylammonium titrant can be prepared in anhydrous isopropyl alcohol as well as in water. Ester formation was minimized by carrying out the titrations immediately after dissolving the samples. The two major considerations in the choice of a solvent for the resolution of acid mixtures are its acidity and basicity and its dielectric constant. Acidity is very important because it determines to a large extent whether or not a weak acid can be titrated. Phenol, for example, cannot be titrated as an acid in aqueous solution because water is too acid and present in too high a concentration to permit the phenolate ion to be formed stoichiometrically by titration with a base. In less acid solvents such as ethylenediamine this titration can be carried out readily (6). In solvents more acidic than water, such as glacial acetic acid, some acids which appear very strong in waterphosphoric acid,. for example-cannot be titrated. The solvent must not be strongly basic if resolution of the strong and moderately strong acids is to be achieved, because of the “leveling effect” of a basic solvent on the stronger acids. Thus in a solvent such as ethylenediamine all acids stronger than acetic acid appear to be equally strong and are not differentiated (6). An ideal solvent for the titration of an acid mixture should be sufficiently weak in acidity to permit titration of the most weakly acid component and sufficiently weak in basicity to permit resolution of the strongest components. The influence of dielectric constant on the relative strength of acids has been discussed by Bronsted (4) and other workers (3, 7, 9, 17). Wolff (1416) has outlined the analytical possibil-
ities of this effect. Most of the common acids can be classified as uncharged, positively charged, or negatively charged. The members of any one class may vary in relative strength to a certain extent as the dielectric constant of the solvent is changed, but in general they behave in a similar manner. Acids of different charge type change greatly in relative strength as the solvent is changed. Acetic acid, benzoic acid, and phenol are uncharged acids. If acetic acid, for example, is chosen as reference, the other members of this group show little change in relative acidity with changes in the dielectric constant of the solvent. There may be small changes due to differences in structure but.
,
I
0
.l”
I
I
1
2
,
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lScpT.p).l
I 4 coho 1 4 “me a, T l i t r a n l rnl
,
5
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Figure 1. Titration of sulfuric acid in water-isopropyl alcohol mixtures VOL. 30, NO. 1, JANUARY 1958
69
