cohols; therefore, he did not encounter the difficulty of variable results (3). Several potential interfering side reactions were studied in depth. Results indicated that silane hydrogen (SiH), alkoxy silane (SiOR), strained siloxane, and vinyl did not interfere under any titration conditions. Acetoxy silane (SiOAc) produced an unknown reaction with the indicator which prevented the color change from being observed easily. No positive evidence was obtained that the group was actually reacting with the reagent, Potentiometric end points were measured with glass combination and platinum-graphite electrodes in an attempt to confirm the color change of the visual indicator. The curves obtained were not of excellent quality but definitely showed the indicator color change occurring within the last 20 of the curve. This was satisfactory for a visual indicator in this weak acid system. Tetrahydrofuran (THF) was used as titration solvent by previous workers (1-3). We found that initial titration r’esults for carbinol and silanol were always high with T H F as the solvent. This had not been reported previously, With two or three successive determinations, results dropped to expected values and remained there. After the first of a series of sam-
z
ples was run, no further difficulties were encountered. The problem most likely arises from production of di-n-butyl amine in the titration which changes the basic strength of the overall system, at least at first. The solution to the problem was found by employing a 1 :1 co-solvent of THF and pyridine. This largely eliminated the effects noted above. Since the standardization and sample analysis procedures employed multiple determinations, the only important difference in data obtained using the co-solvent was improved precision. RSD values were generally improved by a relative 20 to 5 0 z employing THF-pyridine as the solvent. The majority of the data given in this paper was obtained employing THF as solvent, but most of the relative standard deviation values of less than 1 were obtained with the co-solvent.
z
ACKNOWLEDGMENT
The authors thank Dow Corning Corp. for supplying the reagents, apparatus, and silanol materials for this study. RECEIVED for review June 5,1967. Accepted August 10, 1967. In partial fulfillment of requirements for the degree of Master of Science, Gene E. Kellum, June 1967. Presented at the 1967 Anachem Conference, Detroit Mich., October 1967.
Titration of Bases in Acetonitrile I. M. Kolthoff, M. K. Chantooni, Jr., and Sadhana Bhowmik School of Chemistry, Unicersity of Minnesota, Minneupolis, Minn. 55455 The neutralization curves in acetonitrile (AN) of bases of various charge and of mixtures of bases with perchloric acid can be calculated provided the dissoclation constants, K:E+, of the protonated forms of the bases are known. Knowing the dissociation constant of a large number of indicators, the best indicator for the titration of a single monoprotic base or a mixture of monoprotic bases or of a diprotic base can be selected. Bases with pKBi+in water of 4 or greater have been visually titrated accurately and precisely with p-naphtholbenzein as indicator. Bases with pK&+ in water of the order of 1 (dimethylsulfoxide, formamide, urea, anthraquinone) can be titrated spectrophotometrically using m-cresolsulfonephthalein as indicator. The pK&+ i n AN of the above bases in the order mentioned is 5.8, 6.1, 7.7, and 3.5. Mixtures of an aliphatic amine with aniline can be titrated to the first equivalence point, either potentiometrically or with neutral red as indicator. The conductometric titration of uncharged bases with perchloric acid does not give a distinct break at the equivalence point. Good titration lines are obtained by adding an excess of a weak acid, HA, which transforms the base into the homoconjugate salt, BH-AnH-, which is highly dissociated. The homoconjugate anion, itself a base, i s then titrated with perchloric acid.
NUMEROUS acid-base titrations in nonaqueous solvents, including acetonitrile (AN) have been described in the literature ( I ) . In most instances titrations have been carried out potentiometrically, usually with the glass electrode as pH indicator electrode. However, visual indicators have also been used in many instances. (1) J. Kucharskg and L. Safafik, “Titrations in Nonaqueous
Solvents,” Elsevier, Amsterdam, 1965.
On the basis of the acid-base properties of a given solvent it is usually possible to predict its suitability as a medium for titration of acids and bases. However, no calculation of titration curves and of the break in paH at the end point have been made in nonaqueous solvents, except in the protogenic solvent acetic acid (2). For many years acetic acid has been a popular solvent for the titration of bases; however, the break in paH at the end point is much greater in many aprotic protophobic solvents. During the last several years we have been studying acidbase equilibria in AN and have determined the dissociation constant of several acids in this solvent (3-8). Although AN is a solvent with good resolving power for acids, it is not very suitable for the successive titration of acids in their mixtures because of homoconjugation of most acids with their own anions, and heteroconjugation with other anions, homoconjugation being the cause of greatly drawn-out titration curves with an inflection at 5 0 z neutralization (5). Only in very dilute solutions (about 0.001N) homoconjugation usually becomes negligibly small and titration curves, like those in water, are obtained (5). (2) I. M. Kolthoff and S. Bruckenstein, “Treatise on Analytical Chemistry,” Part I, Vol. 1, Interscience, New York, 1959, p. 475. (3) I. M. Kolthoff, S. Bruckenstein, and M. K. Chantooni, Jr., J. Am. Chem. SOC.,83, 3927 (1961). (4) I. M. Kolthoff and M. K. Chantooni, Jr., Ibid.,85,426 (1963). (5) I. M. Kolthoff and M. K. Chantooni, Jr., Ibid.,87,4428 (1965). (6) . , I. M. Kolthoff. M. K. Chantooni., Jr.,. and S. Bhowmik, Ibid.. 88, 5430 (1966).’ (71 . , I. M. Kolthoff and M. K. Chantooni., Jr.., J . Phvs. Chem.., 66., 1675 (1962). (8) I. M. Kolthoff and M. K. Chantooni, Jr., Ibid.,70,856 (1966). VOL. 39, NO. 13, NOVEMBER 1967
e
1627
All carboxylic acids have a very small dissociation constant in A N . For example, the constant of benzoic acid is 2 x Hence their anions are relatively strong bases and yield a very large break in A N in the titration with perchloric acid. The neutralization curves and the paH changes near the equivalence point can be calculated (5). The dissociation constants K&+ of the protonated forms of a large number of uncharged bases, has been determined by Coetzee and Padmanabhan (9). The pK&. of aliphatic amines is of the order of 18, that of unsubstituted aromatic amines, like aniline, of the order of 11. Neutralization curves of such bases with perchloric acid are easily calculated, the homoconjugation B BHT B . . .HB+ usually being negligible in dilute solutions (9). It is also a simple matter to calculate the neutralization curves of mixtures of aliphatic and aromatic amines and the change in paW near the first equivalence point. For example, in the titration of an equimolar mixture of triethylamine (pK,d,+ = 18.5)and aniline(pK,d,+ = 10.6) the paH at the first end point is calculated to be 1/2(18.5 10.6) = 14.6. The change in paH near this point is calculated as it is in water. Because it is possible to calculate the neutralization curves of bases of various charge with perchloric acid, it is also possible to predict which indicators will give a sharp color change at the equivalence point. Dissociation constants of a host of indicators have been reported previously (10). Calculations of paH are more involved in the titration of a mixture of two carboxylates, A- and A'-, with perchloric acid, because of konao- and heteroconjugation:
+
+
AA'-
+ HA A . . . HA- (homoconjugation) + HA -+ A ' . HA- (heteroconjugation) ~
,
However, when the concentrations of the salts in the mixture are of the order of 0.001M, conjugation becomes negligible. Titration of bases, especially of very weak bases, has been recommended by several authors in aprotic protophobic solvents which are much weaker bases than water is. For example, nitromethane has been recommended for this purpose (11-13) in the presence of acetic anhydride, which increases the break in potential at the end point in the titration with perchloric acid. Huber (14) objects to the presence of acetic anhydride, as it can acetylate primary and secondary amines and also because it levels the difference in basic strength between strongly and weakly basic amines. Rodziewicz and Smagowski (15) titrated 0 - , nz-, and p-toluidine, o-, M - ,and p-chloroaniline, and aniline with perchloric acid in nitroalkanes, using the glass electrode as indicator electrode. Pronounced breaks in potential at the end point were observed, even with the weakest base o-chloroaniline (pK,dH- in water is 2.62). Acetonitrile as a solvent for the titration of bases with perchloric acid in dioxanes has been used by Fritz (16), who observed a good resolution in the titration of a mixture of nbutyylainine and pyridine. In addition to the glass electrode, Fritz used indicators for the detection of the end point and (9) J. F. Coetzee and 6.R. Padmanabhan, J . Am. Chem. SOC., 8'7, 5005 (1965). (IO) I. M. Kolthoff, M. K. Chantooni, Jr., and Sadhana Bhowmik, ANAL.CHEM.,39, 315 (1967). (11) J. S. Fritz and h4. 0. Fulda, Ibid.,25, 1837 (1953). (12) G . A. Streuli, Ibid., 31, 1652 (1959). (13) C. A. Streuli, [bid.,32,985 (1960). (14) W. Huber, Z . Anal. Chem., 216, 260 (1966). (15) W. Rodziewicz and H. Smagowski, Roczniki Chem., 40, 511, 727, 1369 (1966). (16) J. S. Fritz, ANAL.CHEW, 25, 407 (1953). '
28
e
ANALYTICAL CHEMISTRY
I
4.0
x
10-3
x
!
1.2
x
I
I
10-2
2.0
x
10-2
1.6 X 10-2 Molarity of HClOa added
8.0
10-3
Figure 1. Potentiometric titration of 0.01M triethylamine with perchloric acid 1, Calculated curve; 2, curve with HC104.1.4 H 2 0 (in nitrometham); 3, curve Fith HC104 (in anhydrous acetic acid); 4, curve with HClO4 (in nitromethane, 0.84M in acetic acid). Indicator p naphtholbenzein
recommended eosin Y (pink-yellow) for the titration of aliphatic bases and methyl violet for aromatic bases. Interesting spectrophotometric titrations in acetic acid and A N of very weak bases and their mixtures with perchloric acid in acetic acid have been described by Hummelstedt and Hume (17), the weakest base titrated being o-chloroaniline. Huber (14) also recommends A N as a medium for the titration of uncharged bases and uses a standard 0.2N solution of perchloric acid in nitroethane. This is prepared by adding 70% aqueous perchloric acid and the calculated amount of acetic anhydride to the solvent and diluting with chlorobenzene to a ratio of about 1: 1 of nitroethane and chlorobenzene. This reagent was found stable for several weeks. He also used 0.5N perchloric acid in dioxane which was made water-free with acetic anhydride. The break in potential of the glass electrode was found more pronounced with the former than the latter reagent. We have found (18) that the hydrogen ion activity of solutions of acids in acetonitrile decreases on aging. 'The aging effect is particularly pronounced with perchloric acid, and is already noticeable 1 hour after preparation of the solution. The total acidity as determined by titration remains unchanged, even after a year of aging, but the hydrogen ion activity decreases more than a thousand times after long stdnding. Apparently, a polymer, P, of A N is formed which is a stronger base than AN: HaN" P -+ Hp+ A N
+
+
Solutions of perchloric acid in AN, prepared by adding a solution of the acid in glacial acetic acid to AN or in any other way, are therefore unsuitable as a titrant. Large amounts of acetic acid interfere in the detection of the (17) L,. Hummelstedt and D. N. Hume, Ibid., 32, 577 (1956). (18) M. K. Chantooni, Jr., Ph.D. Thesis, University o€ Minnesota, 1961,
first end point in the titration of a mixture of an aliphatic and an aromatic amine (Figure 1). For such a titration we have found a 0.2N solution of perchloric acid “monohydrate” in nitromethane very suitable. The comparatively weak base water does not interfere in the detection of the first end point but makes the second end point less distinct. For most purposes the reagent of Huber (14) in nitromethane is suitable, as the small concentration of acetic acid as a rule does not interfere in the titration of a mixture of pure bases. It is shown in this paper that the visual end point detection is very precise and accurate for the titration of dilute solutions of uncharged bases with a pK&+ in water of 4 or greater. For weaker uncharged bases the spectrophotometric method is recommended. The potentiometric method with the glass electrode yields a pronounced break in potential at the end point in the titration of bases with a PK;~ + in water of 4 or greater. Howeker, the observed break in potential is less than the calculated one, because the glass electrode in AN has been found to indicate too high a pH in solutions of perchloric acid ( 5 , 19). Moreover, the potential is not stable in the latter. Before the end point, water and acetic acid from the titrant do not affect the pH appreciably. For the titration of very weak bases (pK& in water of the order of 1 to 2) the glass electrode is not suitable, while the spectrophotometric method yields excellent results. Conductometric titrations in AN of bases are of little or bery limited analytical importance. For the sake of completeness, titration curkes of bases of different charge with perchloric acid are discussed briefly. In the present work we have studied the visual, spectrophotometric, potentiometric, and conductometric titration in AN of uncharged and of univalent anion bases with perchloric acid. EXPERIMENTAL
Chemicals. Acetonitrile was purified and dispensed as described before (3). Acetic acid (HAC), purified as described elsewhere (20), was found to be 5.0 X lO-3M in water. Nitromethane (Spectrograde) obtained from Eastman Organic Chemicals and N,N-dimethylformamide (DMF), Fisher Certified Reagent, were used without further purification. Triethylamine (3), pyridine (21) and dimethj lsulfoxide (22) (DMSO) were purified as described previously. Urea, Baker’s Analyzed Chemical Reagent, and aniline, Mallinckrodt Analytical Reagent, were used directly without further purification. Diphenylamine, anthraquinone of unknown origin were purified by sublimation. Tetraethylammonium bisulfate (3), 3,5-dinitrobenzoate ( 4 , o-nitrophenolate (6), salicylate (8) anhydrous lithium benzoate (8), sodium acetate (20) and p-nitrophenolate (6) were purified and dried similarly as before. Anilinium perchlorate. Four milliliters of 70 % perchloric acid (G. F. Smith Co.) dissolved in 100 ml of Du Pont glacial acetic acid were added to 2 ml of Merck Analytical Grade acetic anhydride. The solution was allowed to stand for 1 hour, then titrated with aniline to the crystal violet end point, The salt precipitated on standing; the crystals were filtered and recrystallized from ethyl acetate and benzene mixture, The salt was dried under vacuo at 50” C. Analysis by gravimetric method (23) corresponded to a purity of 99.5%. (13) J. Debarres, Bull. SOC.Chim.:1963, p. 2103. (20) I. M. Kolthoff and S. Bruckenstein, J . Am. Chem. SOC.,78, 1 (1356). (21) I. M. Kolthoff and M. K. Chantooni, Jr., Ibid.,87, 1004(1965). (22) I M. K.olthoff and T. B. Reddy, Inorn. Chem.. 1, 189 (1962). (23) H. H. Willard and 6. Smith,. IKD.E ~ GCHEM., . ANAL.ED., 11, 186 (1939).
Triethylammonium perchlorate was prepared by neutralizing potentiometrically 5 ml of 70% perchloric acid in 100 ml of ethanol with triethylamine. This solution was evaporated to dryness under vacuo at room temperature. The salt was dissolved in a minimum amount of ethyl acetate and reprecipitated with petroleum ether. The crystals were dried in vacuo at room temperature. The only impurity was perchloric acid ( v i ) corresponding to 0.3% of the salt. Generally, 0.liM stock solutions of the salts were prepared in AN just before use. A suitable aliquot was taken by means of an ultramicroburet. Stock solutions of alkali carboxylates (0.1M)were made in glacial acetic acid. Perchloric acid “monohydrate” was a G . F. Smith eo. product. From Karl Fischer titration it was concluded that its composition was HCIQ; 1.37 K O . Perchloric acid in anhydrous acetic acid was prepared as described by Coetzee (24). Perchloric acid “monohydrate” in nitromethane was prepared by slowly adding a measured amount of the acid to nitromethane. It was standardized in AN against triethylamine both potentiometrically and visually with p-naphtholbenzein as indicator. The solution was found to be colorless and stable for at least several weeks. p-Nitrophenol (6) and resorcinol (25) were the same product as previously used. Indicators. Dibromothymolbenzein (D.Br.T.Bz.), p-naphtholbenzein, (p-N.Bz.), neutral red (N.R.) and m-cresol purple (m-Cr.P.) were the same samples of indicator as recently described (10). A stock solution of 1.0 x l O - 3 M of each indicator was made in AN. Saturated solutions of N.R. and m-Cr.P. (solubility -2.0 X 10-3M) in AN were used. Measurements. All experiments were carried out at 25 ’ = 1 C. Two Coleman tripurpose (Catalog No. 3-472) glass electrodes were used in emf measurements and calibrated in suitable picrate and nitroaniline buffers (5). For these 722 - E electrodes, the relation was found: paH = ___ in AN. 59.1 Other experimental details are similar to those described earlier (5). Spectrophotometric titrations of very weak bases (urea, DMF, DMSO, etc.) using a pair of 1.8-cm, glass-stoppered, borosilicate-glass cells and the absorbance measurements were made in a Beckman D U spectrophotometer, The analjses by spectrophotometric titration are very susceptible to the presence of traces of water. Therefore it is very essential to use freshly prepared stock solutions of the base and perchloric acid. Also, a blank in presence of the same amount of acetic acid as in the actual analysis must be performed. The experimental cells and the techniques involved for conductometric titration were the same as described previously (4). O
RESULTS
Potentiometric Measurements. K&+ of weak bases, urea, DMF, DMSO, and anthraquinone (K& - of the order of 10-1 or less in water) were determined potentiometricallj in equimolar solutions of 0.005M bases and their respective perchlorates. pK&+ of urea, DMF, DMSO, and anthraquinone were found to be 7.7, 6.1, 5.8, and 3.5, respectively, at 25’ C. No pronounced break was observed at the equivalence point in the potentiometric titration of these bases with a solution of perchloric acid in acetic acid. Using equimolar solutions of 0.02, 0.01, and 0.0018M triethylamine and its perchlorate, and also of 0.02 and 0.01M (24) J. F. Coetzee and I. M. Kolthoff,J . Am. Chem. SOC.,49,6110
(1957). ( 2 5 ) I. M. Kolthoff and M. K. Cnantooni, Jr., Zbid., 85, 2195
(1363). VOL. 39, NO. 13, NOVEMBER 1967
Q
1
aniline and anilinium perchlorate, we found their pK,Z+ to be 18.7 and 10.7, respectively. These values agree well with those (18.46 and 10.65, respectively) found by Coetzee (9). The constants were also determined by the spectrophotometric method using the indicators 2,6-di-tert-butyl-4-nitrophenol and picric acid in mixtures of the bases with their perchlorates. The results checked with those obtained potentiometrically. Potentiometric Titrations. Triethylamine was titrated with perchloric acid both potentiometrically and visually using the indicator p-N.Bz. The titration curves are shown in Figure 1. Curve 1 (Figure 1) is the calculated curve; curve 2 is the experimental curve obtained with HClOd. 1.37 WZQ (in nitromethane); curve 3 with HC104(in glacial HAC); curve 4 with €IC104 (in nitromethane, which was 0.84714in HAC). I n all cases the end point observed is sharp and precise. Curve 2 before the equivalence point is virtually identical with the calculated one. After this point the glass electrode records a considerably higher paH than corresponds to the concentration of free perchloric acid; consequently, the experimental titration
Table I. Color of Neutral Red (2 X 10-5M) and paH in Mixture 0.02M in Triethylammonium Perchlorate and 0.02M in Aniline upon Addition of Triethylamine or Anilinium Perchlorate Concn of Et3N M X 104 0 0.282 0.712 1.42 2.13
Concn of anilinium
perchlorate
Et3Nin excess over salt in 0
0.14
15.8
0.35
16.4
0.71 1.07
16.6 16.8
Anilinium perchlorate in excess, 0 0.25 0.56 1.0 2.25
0
0.50 1.13 2.0
4.5
paH 14.7
pal-I 14.7 13.4 13.1 12.8 12.5
Color red orange yellow yellow yellow
Color red red red red red
Table IT. Color of Neutral Red (2 X 10-6M)and p-Naphtholbenzein (5 X 1 K 6 M )and paH Near the Second Equivalence Point in Mixture 0.02M in TriethglammoniumPerchlorate and 0.02M in Anilinium Perchlorate Concn of Excess aniline aniline M x 104 in % paH Color of N.R. p-N.Bz. 6.2 red green 0 0 7.7 red yellow-green 0.21 0.41 8.4 red yellow 0.41 0.82 8.6 red yel!ow 0.62 1.24 8.9 red yellow 1.0 2.0 Concn
Excess
of HClOi HC104 M x lo4 in % 0 0 0.50 0.25 0.47 0.94 0.68 1.37 2.27 1.I3 Of
Color of N.R. p-N.Bz. red green green purple purple green bluish purple green ( 5 . O)a bluish purple green (4.8)a a Glass electrode indicates too high paH.
1630
e
paH
6.2 (5.2)“ (5.l)O
ANALYTICAL CHEMISTRY
curves in Figure 1 are not symmetrical around the end point, For example, assuming that perchloric acid is completely dissociated, in the presence of an excess of perchloric acid corresponding to 3.0 X 10-3M HClO4, the solution should have a paH of 2.5, while the observed paH is 4.1 in curve 2. When corrected for the concentration of hydrated protons at a concentration of 0.022M water in the system, the calculated paH is 2.8. In curves 3 and 4 lower paH values than in curve 2 were found before the equivalence point as a result of reaction of acetic acid with the free base. The end point in the titration of aniline with HClO, in HACis quite sharp and precise but not with HCIO,. 1.37 HeO in nitromethane. The water in this solvent is responsible for this behavior. Mixture of Triethylamine and Aniline. In AN (pKEtNH+PKC&H~-) = 8.0, whereas the water it is 6.1. Hence the break in potential at the first equivalence point is considerably greater in AN than in water. In order to get accurate experimental data on the change in paH at the first end point we have measured the change in paH of a mixture, 0.02M in triethylammonium perchlorate and 0.02M in aniline upon addition of small amounts of triethylamine or anilinium perchlorate. In order to get reliable results the two perchlorates must be free of traces of acid or basic impurities. The triethylammonium perchlorate appeared to contain 0.3 free perchloric acid, which was found by potentiometric and visual titration with triethylamine. The free acid in the salt solution used in the present work was neutralized with triethylamine. The anilinium perchlorate was free of acid and basic impurities. The results are presented in Table I. The calculated paH of the mixture corresponding to the first equivalence point is 14.6 in good agreement with the experimental value of 14.7. Also the experimental values of paH with a slight excess of triethylamine or anilinium perchlorate were in close agreement with the calculated values. From Table I it is evident that neutral red gives a sharp color change at the first equivalence point. Acetic acid or water in concentration less than 0.05M hardly affected the potentials or the color of the indicator. Actual titrations of mixtures of triethylamine and aniline with perchloric acid in nitromethane yielded accurate results at the first equivalence point. The sum of the concentration of the bases was titrated with perchloric acid in acetic acid. From the data in Table 11 it is clear that the measured paH changed from 7.7 at 0.2 % before the end point to 5.2 at 0.25 % after the end point. The color change of N.R. from red to purple is very sharp, p-Naphtholbenzein indicates even a sharper end point. Spectrophotometric Titrations. Urea, DMF, DMSO, and diphenylamine were titrated spectrophotometrically using m-Cr.P. as indicator and HClO, (in glacial HAC) as the titrant. All solutions and the reagent were freshly prepared and blanks were run at the same time of the titration. As example, Figure 2 shows the plot of [HI+]/[I] us. CHCIOa for the blank and the titration of 5.0 X 10-3MDMSO. The results with the various bases are given in Table 111. Visual Titration. Indicators with pK& in AN of the order of 6 indicate very sharply the end points in the titration of salts of weak acids and uncharged bases with pK&+ of the order of 10 or greater. The pK& values of p-naphtholbenzein, dibromothymolbenzein and neutral red have been found to be 6.8, 5.9, and 6.0, respectively (IO). The color change of D.Br.T.Bz. from yellow to red is very pronounced. This indicator is not commercially available. The color change of p-N.Bz. from yellow to green is equally sharp and this indicator is highly recommended. Neutral red is also quite useful
Table 111. Spectrophotometric Titration of Very Weak Bases with HCIOd (in Glacial Acetic Acid) Using rn-Cr.P. (1.6 X 10-6M) as Indicator Concn Concn taken found Base M X 108 M X l o 3 Error, % 4.80 4.80 0 Urea 4.75 4.70 +1.1 Diphenylamine 9.90 10.0 $1.0 N,N-Dimethylformamide 5.1 5.1 0 Dimethylsulfoxide 26.0 26.4 +1.5
L
0.4
L 0.0 X 10-8
6.0 X 10-8
4.0 X 10-3
1.0 X 10-2
1.2 X
8.0 X
Molarity of HClOa added
Figure 2. Spectrophotometric titration of 5.1 X 10-3M DMSO with perchloric acid (glacial acetic acid) m-Cresol purple as indicator. Wavelength 533 mp. Plot 1 is blank.
but its sensitivity for hydrogen ions becomes considerably less in the presence of much acetic acid as is seen from data in Table IV. On the other hand, even in a mixture of 1 :1 ANHAC, the color changes of p-N.Bz. and D.Br.T.Bz. are very sharp. However, the red color of the latter fades on standing; while the green color of the former is stable. As seen in Table IV the sensitivity for hydrogen ions of N.R. in the 4 : l ANHACmixture is some 20 times less than in AN. Table V presents the results of the same titrations of bases with perchloric acid. Neutral red is not recommended for the titration of alkali carboxylates because of the interference of acetic acid used in making a solution of the salt. I n all instancesp-N.Bz. appears to be a n excellent indicator. The sensitivities in 4 :1 , l :1 AN-HAC and in pure HACare difficult to determine as the color changes of N.R. in Table IV are indistinct in these media.
Table V.
Indicator p-N.Bz. p-N .Bz. N.R. p-N.Bz.
NaOAc NaOAc NaOAc LiBz
D.Br.T.Bz.
LiBz
+ + +
+ +
For the sake of completeness it may be added that in the visual titration of picrates with perchloric acid in HAC, no sharp end point is obtained, the picrate ion serving as its own indicator. On the other hand, solutions less concentrated
Visual Titration of Bases with HClOd (in HAC) Concn of indicator, 2.0 X 10-6M
Base&
Medium 4AN lHAc HAC 4AN lHAc 4AN 1HAc 1AN lHAc 4AN 1HAc 1AN 1HAc 4AN lHAc AN AN AN AN AN AN AN AN AN AN AN AN
+ + + + + + +
N.R.
p,-N.Bz. D.Br.T.Bz. N.R. p-N.Bz. D.Br.T.Bz. N.R. p-N.Bz. D.Br.T.Bz. N.R. p-N.Bz. D.Br.T.Bz. N.R. a OAc = acetate, Bz = benzoate, Sal * Color change not sharp,
Table IV. Sensitivity for [Hf] of (2.0 X lW5iM) g-N.Bz., D.Br.T.Bz., and N.R. in AN and in Mixtures of AN with Acetic Acid (HAC) [H+]at distinct color change, Indicator Color change Medium M AN 6 x 10-6 yellow-green 4AN lHAc 8 X 10-6 p-N.Bz. yellow-green 1AN lHAc 1.6 X yellow-green yellow-green HAC 4 x 10-5 AN 8 X 10-6 y ellow-red 4AN IMAc 1.5 X 10-6 D.Br.T.Bz. yellow-red I A N + lHAc 2 X 10-6 yellow-red HAC 6 x 10-6 yellow-red red-bluish purple AN 3 x 10-6 N.R. red-bluish purple 4AN 1HAc 6 X red-bluish purple 1AN lHAc -3 X 10-3 red-bluish purple HAC -3 x 10-3
=
Concn taken iM x 103 4.74 4.74 4.74 1.72 4.30 2.90 2.90 4.30 3.92 3.92 3.92 0.550 0,550 0,550 1.04 1.04 5.20 8.50 8.50 8.50
Concn found (corrected for blank) M x 103 4.12 4.75 4.80 1.75 4.38 2.92 2.91 4.40 3.96 3.96 3.96 0.540 0.540 0.540 1.04 1.03 5.19 8.55 8.55 8.55
Error, % -0.4 $0.2 i-1. l b
+1.7 +1.9 $0.7 Jr0.3
1-2. 3b +1.0 +1 .o +1.0b
+1.8 +1.8 + l , 8b 0
-1.0
-0.2 +0.6 1-0.6
+-0.6
salicylate.
VOL. 39, NO. 13, NOVEMBER 1967
e
1631
/
8
9 7
8
6
7
5
z X -I4 6
3 5
2 4
/
1
'
3
0
I
I
II
0.001
0.OOP
1
I
I
0.003
0.004
0.005
Eq. Pts. I
I
0
9
I
0.001 Molarilv HClOi added
0.002
Figure 4. Conductometric titration of tetraethylammonium salts in AN with 0.489M perchloric acid (in acetic acid) 1
1, 1.94 X 10-aM3,5-dinitrobenzoate; 2,1.19 X 10-SMinpresence of 0.32M resorcinol; 3, 1.00 X 10-aM bisulfate. Upper abscissa to 1; lower to 2,3. No viscosity corrections I
0
I 0.004
0.002
I 0.006
Eq. Pt.
Molarity HCiOa added
Figure 3. Conductometric titration of 0.0030M triethylamine with 0.489M perchloric acid (in acetic acid) 1, in absence of HDNB; 2 , in presence of 0.015.V; 3, of 0.350M HDNB
of the titration curve in very dilute solution is easily calculated assuming both BHClOd and BH+HA2- are completely dissociated. Before the equivalence point the following reactions occur: B 2HA + BHHAaHA*- H+ + 2HA
+
+
+
The electroneutrality condition is than 10-3M can be titrated spectrophotometrically at wavelengths of the order of 380 mp. Conductometric Titrations. AMINES. The mobility of the solvated proton in AN is 80 (23),and not much different from that of ions of similar size. In the conductometric titration of uncharged bases there is therefore no break in the titration curve at the equivalence point (see Figure 3). However, by transforming the base into a highly dissociated homoconjugate salt of a relatively weak acid, a pronounced break is observed at the equivalence point. The dissociation constant of such salts is of the same order of magnitude as that of tetraethylammonium salts ( 210-2) (4,and good titration curves of an aliphatic amine with perchloric acid are obtained in presence of a large excess of a weak-acid hydrogen-bond-donor such as 3,5-dinitrobenzoic acid. In fact, K E ~ ~&N (DN HB ) ~ has previously been found equal to 3.0 X IO-* ( 4 ) . The shape 1632
e
ANALYTICAL CHEMISTRY
[BH+] = [ClO,-]
+ [HA*-],
(1)
[H+] and [A-] being negligible. Substituting the expression for the equivalent conductance, L,
I03L = ABE+HA~[BH'] into Equation 1, Equation 2 results. 103L = [BH+] - ~ B H + H A ~-k - [C~OI-]{boa-
- h~.42-]
(2)
In Equation 2 [BH+]is equal to the analytical concentration of amine transformed into BH+HA2-. From Equation 2 the intercept and slope of the titration line before the equivalence point is 1 0 - 3 [BH+] ABH+HA?-and 10-3{hc10p- - AHA%-], respectively. At the equivalence point the conductance is due only to BHClOd, hence 1031, = [BH+]ABHCIO~. In Figure 3 is shown conductometric titration curves of 3.0 x 10-aM triethylamine with perchloric acid in absence and in
Table VI. pK and paH at First Equivalence Point in AN PH
at 1st
18.46 1,2-Diaminoethane 1,3-Diaminopropane 19.70 20.12 1,4-Diaminobutane 18.29 Triethylenediamine
13.01 15.0 15.34 10.16
5.5 4.7 4.8 8.1
3.0 2.0
1.5 4.0
15.7 17.4 17.7 9.1
presence of 0.015 and 0.35M 3,5-dinitrobenzoic acid. Corrections for viscosity were applied in the latter system (v = 0.412, that of pure solvent 0.345 cp). From the intercept of the titration line before the equivalence point in presence of 0.35M 3,5-dinitrobenzoic acid AEts" tH(DNB)2- was found equal to 117. In a previous publication ( 4 ) it was reported that X O E ~ ~ N H L = 88 and XOH(DNB)?- was 46. Taking into account the effect of ionic strength, p, X E ~ ~ S H = + 78 and XEI(DNB)~- = 38 at p = 0.003. From the conductance at the equivalence ~ O 167 ~ while a value of 172 is calculated from point A E ~ ~ N H C = XEtsNH- and Xclo,- = 94 (18) at p = 0.003. From the good agreement it is apparent that both triethylammonium perchlorate and 33-dinitrobenzoate homoconjugates are essentially completely dissociated in 3 X lO-ZM solution. The observed and calculated slopes of the titration line in 0.35M HDNB (Figure 3) up to the equivalence point are 5.0 X 10-2 and 5.8 x lo-*, respectively, while beyond the equivalence point the values are 1.6 X 10-l and 1.5 X respectively. When B can form BzH+(26) the dissociation constant of the homoconjugate B2H-HA2- is somewhat greater than that of BH+HA2-. On the other hand < AB=+ and BzH+ formation has little effect on the shape of the conductometric titration curve. This is especially true in the titration of very dilute solutions when BzH+ formation can be neglected for present purposes. ANIONBASES. When the salt is highly dissociated, such as tetraalkylammonium salts, the conductivity usually increases slightly up to the equivalence point as K R & O ~> K&A when hclo,- 2 Xa-. After the equivalence point there is a pronounced increase in conductivity. Even a 0.001M solution can be titrated with good accuracy and precision. In Figure 4 are presented conductometric titration curves of 0.0194M tetraethylammonium 3,5-dinitrobenzoate and of 0.0010M tetraethylammonium bisulfate with perchloric acid. The error in this type of titration was found to be & O S %. The conductometric titration of 6.84 X 10-4Msodium p nitrophenolate with perchloric acid is illustrated in Figure 5. I n the absence of any added hydrogen bond donor a fairly sharp minimum was found at about 10 neutralization. p-Nitrophenol and acetic acid have about the same dissociation constant ( 6 ) and the minimum is accounted for by acidbase interaction of the salt anion with acetic acid, with formation of the slightly dissociated and slightly soluble sodium acetate. Upon further addition of perchloric acid (in acetic acid) the conductance increases because of the formation of the highly dissociated homoconjugate salt. Without addition of a hydrogen bond donor the end point is hardly detectable (curve 1). Curve 2 illustrates the beneficial effect of a hydrogen bond donor, p-nitrophenol, which even when added in a small concentration of 0.03M to a solution of 7 X 10-4M sodium p nitrophenolate gives rise to straight titration lines and an easily detected end point. (26) J. F. Coetzee, G. R. Padmanabhan, and G. Cunningham, Talanta, 11,93 (1964).
0.2
0.4
0.6
0.8 1.0 1.2 1.4 Molarity HClOi X I O 3
1.6
1.8
Figure 5. Conductometric titration of 6.84 X 10-4M sodium p-nitrophenolate with 0.489M perchloric acid (in acetic acid) 1, in absence of p-nitrophenol; 2, in presence of 0.029M p-nitrophenol. No viscosity corrections made DISCUSSION
From the introduction and experimental part it is evident that neutralization curves with perchloric acid of dilute solutions of uncharged bases in AN can be calculated in a similar way as in water. Similar calculations for monovalent anion bases can be made in dilute solutions ( ~ 0 . 0 0 1 M ) . In more concentrated solutions conjugation may have to be considered. For diprotic uncharged bases the calculations again are similar to those in water, taking activity coefficients into account. Coetzee and Padmanabhan ( 9 ) have determined PKBH + and pKB& of some diprotic bases in AN. It is evident from Table VI that pKBIld- - pKB&+zdenoted as ApK is much greater in AN than in water. The first three bases in Table VI cannot be titrated as monoprotic bases in water, but should give a detectable break in potential in AN. pKB&Taof the first three bases in water are 6.90, 8.49, and 9.20; hence in the titration as diprotic bases, much sharper end points are found in AN than in water. Triethylenediamine in water has a pK&,+? of 4.2 and cannot be titrated as a diprotic base. On the other hand, in AN the end point should be as sharp as observed with aniline. It is also seen from Table VI that, in its titration as a monoprotic base, the end point is much sharper in AN than in water. RECEIVED for review June 26, 1967. Accepted August 21, 1967. This work was supported by the Directorate of Chemical Sciences, Air Force Oflice of Scientific Research under Grant AF-AFOSR-1223-67, and by the National Science Foundation. VOL. 39, NO. 13, NOVEMBER 1967
e
1633
