shows a loss of CzH4to m/e 180 followed by loss of CH3. to m/e 165. The spectrum of Fraction 12 shows a metastable peak at 137.2, consistent with the loss of 56 from the 236 peak to form the 180 fragment and a metastable peak at 151.3 confirming the genesis of the 165 from the 180. A large peak at m/e 221, due to the loss of CH,. from the 236 peak, implies at least one methyl group in the molecule. Taken together, the above information indicates that the structure may be a dimethylhexahydropyrene with the 2 methyls attached to the 2 ring-carbons that are lost in forming the m/e 180 ion. The long chromatographic retention time of this material of relatively low polarity needs to be pointed out. For the present it must be explained as a gel permeation effect, namely, the occurrence of a species of very small molecular size. Summary of This and Previous Work. The classical picture of “Naphthenic Acids” as a mixture of alkanoic and l-ringnaphthenic acids must be greatly expanded for this California crude oil by the discovery of some 40 new compound classes of carboxylic acids. Table IX summarizes the results of this and the previous paper (1). The estimate of the amount of each type listed was obtained via the portion of the ion current for each (MS) and the intensities of other spectrometric absorptions combined with the amount of the subfraction in the parent fraction and of the latter in the crude oil. Because only 25 of all acids present in this oil were investigated and many similar or identical structures are expected to be present in neighboring carboxylic acid fractions of the main separation (6), the estimates are minimum values. The small quantity of each fraction available for investigation and the complexity of the acid mixtures described in this paper prohibit unequivocal identification in many cases. However, most of the structural assignments for the major
components are believed to be correct because of mutually supporting evidence from the various techniques of molecular spectrometry for most types listed in Table IX, previous finding (8) of many structures as fluoroalcohol esters and occurrence of many of the acid structures found here as hydrocarbons in petroleum (11, 13-16). In order to enable future workers in this field to clarify the unavoidable ambiguities, detailed spectral information was listed for each fraction in the Tables of this paper. The findings reported here should, it is hoped, shed light on the puzzling geochemical relationship between carboxylic acids and hydrocarbons. The next paper will deal with deuterium labeling at the site of carboxyl attachment leading to proof of the absence of contamination and identification of individual terpanoid and steranoid polycyclic naphthenic acids. ACKNOWLEDGMENT
The authors thank R. M. Bly for interpretation of fluorescence spectra. RECEIVED for review January 8, 1970. Accepted March 30, 1970. Paper presented in part at the Gordon Research Conference on Organic Geochemistry, Holderness, New Hampshire, August 1968, and in full at thejoint meeting of the American Chemical Society and the Chemical Institute of Canada, Division of Analytical Chemistry, Toronto, May 1970. (13) L. R. Snyder, ANAL.CHEM., 41,314 (1969). (14) C. J. Robinson and G. L. Cook, ibid., p 1548. (15) H. V. Drushel and A. L. Sommers, ibid., 39, 1819 (1967). (16) C. F. Brandenburg and D. R. Latham, J. Chem. Eng. Datu, 13, 391 (1968).
Potentiometric Study of Base Strengths in the Binary Solvent, Acetic Anhydride-Acetic Acid Orland W. Kolling and Wilton L. Cooper Chemistry Department, Southwestern College, Winfield, Kan. 67156 The response of the glass electrode to changes in acetate ion concentration in solutions of bases was examined in the binary solvent, acetic acid-acetic anhydride. In general the potential of the electrode is shifted toward a more positive value with increasing acetic anhydride content of the solvent. The apparent basicities of representative ionic and uncharged bases in the mixed solvent were compared to their pKB values in anhydrous acetic acid. It was found that the relative strengths of the bases are validly expressed by the half-neutralization potentials under conditions of constant concentration for the bases, with HCIO, as the reference acid. Within the separate groups of charged and uncharged bases, the order of decreasing basicity in acetic acid is unchanged by the addition of large amounts of acetic anhydride. The primary effect of acetic anhydride upon the cell potential appears to be the change of the value of Eo in mixed solvents for the mole fraction range of 1.0 to 0.28 in acetic acid. However, the quantitative statement of this effect must be expressed in terms of the mole fraction of dimeric acetic acid. 758
ANALYTICAL CHEMISTRY, VOL. 42, NO. 7, JUNE 1970
A NUMBER of analytical methods involving nonaqueous solvents have employed potentiometric determinations of weak electrolytes in media that are predominantly acetic anhydride. Representative studies have included : the titration of uncharged weak bases ( I ) , basic anions (2), quaternary ammonium salts (3),aromatic N-oxides (4), and phosphine oxides (5); the indirect determination of transition metal cations (6); and the comparative strengths of sulfonic acids (7). The polarographic behavior at the DME for several transition
J. Fritz and M. Fulda, ANAL.CHEM., 25, 1837 (1953). C. A. Streuli, ibid., 30, 997 (1958). M. Puthoff and J. Benedict, ibid., 36, 2205 (1964). C. Muth, R. S. Darlak, W. H. English, and A . T. Hamner, ibid., 34, 1163 (1962). ( 5 ) D. C . Wimer, ibid., p 873. (6) J. S. Fritz, ibid., 26, 1701 (1954). (7) D. J. Pietrzyk, ibid., 39, 1367 (1967).
(1) (2) (3) (4)
metals has been reported, as well (8, 9). Generally, the solvent medium used has not been pure acetic anhydride, but a binary solvent mixture. The minor solvent component is usually acetic acid, and, if not added directly to the system, it enters in varying amounts indirectly from one or more of these sources: (a) unpurified commercial acetic anhydride used as the solvent; (b) variable small quantities of water or acetic acid added to acetic anhydride in order to dissolve compounds which are insoluble in the aprotic component; and (c) titrant acids made anhydrous by the reaction of acetic anhydride with water in the reagent acid (ex. aqueous perchloric acid or the hydrated sulfonic acids). A qualitative treatment of the role of acetic acid in acetic anhydride was reported by Pietrzyk (7) for acid-base titrations. We are presenting herein our results on the effect of solvent composition upon the apparent strengths of bases in the mixed solvent. The solvent mole fraction range considered was 0.0 to 0.936 in acetic anhydride. Since the quantitative aspects of electrolytic equilibria in anhydrous acetic acid are on a fairly firm basis, one approach to the study of the mixed solvent is to examine the effect of the addition of increasing amounts of the aprotic component upon the base-protonated base equilibrium. It is assumed that acetic acid will preferentially solvate the base in the mixed solvent over a very wide range of increasing amount of the aprotic component, and that Equation 1 applies. B .HOAc G BH+
+ OAC-;
KB =
[BH+I[OAc-I [BHOAcl
(1)
This approach was successful in the clarification of the potentiometrically measured base strengths in acetic acid-pdioxane (IO),using the glass indicator electrode. Changes in the emf for the glass-calomel electrode pair in acetic acid-acetic anhydride were determined at 25 "C for perchloric acid and the representative bases listed in Table I. Only tertiary bases of periodic group VA elements were selected for the uncharged species. Likewise, the choice of charged bases was restricted to the ionic acetates in order to avoid complicating influences of anion acetolysis reactions in the mixed solvent. EXPERIMENTAL Reagents and Solutions. Reagent grade acetic anhydride was re-distilled through a 52-cm Raschig column, and the center 80% cut, having a corrected boiling range of 139.0139.2 "C, was used in all acetic acid-acetic anhydride mixtures. Acetic acid was dehydrated with acetic anhydride, following the method of Tappmeyer and Davidson (11). The water content of these solvents was less than 0.002% by Karl Fischer titration. All solvent mixtures were prepared volumetrically at 25 "C, using burets protected with drying tubes. Stock solutions of the bases, half-neutralized bases (with HC104), and perchloric acid were made in anhydrous acetic acid at the 0.1-O.05M level. Procedures for preparation and standardization of such solutions were the same as were used in earlier work (10, 12). Apparatus and Procedures. The potential of the calomelglass electrode pair in acetic acid-acetic anhydride solutions of the several weak electrolytes was measured with a Leeds and Northrup Model 7401 pH meter. For both the reference and glass (L and N 117169) electrodes, the internal aqueous solutions were retained. As was noted before (12), aqueous leakage from the reference electrode during measure(8) W. Mather and F. Anson, Anal. Chim.Acta, 21,468 (1959). (9) J. Headridge and D. Pletcher, J . Chem. SOC.(A), 1966,757. (IO) 0. Kolling and D. Garber, ANAL. CHEM., 39,1562 (1967). (11) W. Tappmeyer and A. Davidson, Znorg. Chem., 2, 823 (1963). (12) 0.Kolling and E. Mawdsley, ibid., 9,408 (1970).
Table I. Dissociation Constants for Weak Electrolytes in Anhydrous Acetic Acid at 25 "C
Solute PKB Perchloric acid 4.87" Lithium acetate 6.790 Sodium acetate 6.58s Potassium acetate 6.w 5 . 20b N,N-Diethylaniline Pyridine 6.10" Tribenzylamine 5.12 5.15* Tri-n-butylamine Triphenylamine 9.20c Triphen ylarsine 1O.6Oc Triphenylphosphine 8.W S. Bruckenstein and I. M. Kolthoff data (14). 0. Kolling comparative method (13). 0. Kolling and E. Mawdsley most recent values (12). Table 11. Change in Cell Potential as a Function of Dilution for NaOAc in Mixed Solvents Having Fixed Compositions Mole fraction AciO ( X ) CBrange ( M ) Slope 0.000 0.0020-0.020 0.030 0 076 0.0008-0.020 0.030 0.191 0.0008-0.020 0.030 0 030 0.288 0.0008-0.020 0.436 0.0008-0.020 0.031 0.563 0.0008-0.020 0.033 0.708 0.0008-0.020 0.034 0.870 0.0008-0.012 0.035 0.936 0.0008-0.004 0.033
ment is well below the limit of effect upon emf data in HOAc. The electrodes were pre-conditioned in anhydrous acetic acid for two weeks before a series of measurements was begun, and the electrodes were stored in the same solvent. Our experimental methods involving electrode standardization against a sodium acetate reference solution have been reported in detail (12). Emf values given in the Tables or Figuresare means derived from six determinations on solutions at 24-25 "C. The overall precision of the data on bases and half-neutralized basis is =t2 mV; and the results for perchloric acid are less reliable, having an uncertainty of * 5 mV. RESULTS AND DISCUSSION Electrode Response in Base Solutions. The potential of the glass and saturated calomel electrode pair in anhydrous (or glacial) acetic acid solutions of solvated bases is given by the form of the Nernst relationship in Equation 2, cited previously (10). EB =
+ + RFT In Ks - RT~n 2F
(Eoco E,)
-
K ~ C B
(2)
(Eoc" is the standard potential for the electrode pair, E, the liquid junction potential at the boundary with the calomel electrode, K, the autoprotolysis constant of the solvent, and KBthe dissociation constant for the base in Equation 1, having a stoichiometric concentration of CB.) The proportionality between EB and -In CB provides a simple test of regularity for the indicator electrode response in anhydrous acetic acid. Although the theoretical slope of 0.0295 volt per tenfold dilution at 25 "C is not found for all bases in pure acetic acid, those selected for the present study conform very closely to a 0.030 volt slope (12). The dilution us. emf function was determined for sodium acetate in various acetic acid-acetic anhydride mixtures of fixed composition. Representative data are listed in Table 11. ANALYTICAL CHEMISTRY, VOL. 42, NO. 7, JUNE 1970
759
1
0
0:8
0:4
Figure 1. The trend in EB for NaOAc as a function of mole fraction of acetic anhydride ( X ) in the binary solvent Fixed concentrations of NaOAc are: (1) 0.0008M; (2) 0.002M; (3) 0.004M; (4) 0.OOSM; and (5) 0.020M.
0.4
0.8
I
X Figure 2. Change in EB for ionic acetates as a function of acetic anhydride mole fraction ( X ) The bases (at constant 0.004M) are: (1) LiOAc; (2) NaOAc; and (3) KOAc
Plots of log CB us. E B (not included herein) were definitely linear over the concentration ranges recorded, and the experimental slopes were based upon a minimum of six CB values. The experimental uncertainties in the slopes are of the order of i~0.0005V. In general the data approximate the theoretical slope up to an anhydride mole fraction of about 0.30, and at higher mole fractions the small but regular increase in slope is not sufficient to nullify the use of the glass indicator electrode for semiquantitative comparisons of acetate ion concentrations in solutions of bases in the mixed solvent. As a check upon the internal consistency of the non-concentration terms related to EB in Equation 2, measurements of emf as a function of mole fraction (HOAc) were made for each of the ionic acetates at several fixed CB values. Results obtained for sodium acetate are shown in Figure 1. The same general trend of a shift toward a more positive potential with increasing acetic anhydride content is shown by the family of parallel curves. At a fixed CB,the proportionality between EB and ~ K inB Equation 2 allows the direct comparison of basicity constants in pure acetic acid as the solvent (13). For the binary solvent, if the changes in the E, and K , terms are smaller than the changes in EGcoand KB at a constant CBas acetic anhydride is (13) 0. W. Kolling, ANAL.C H E M . ,956(1968). ~~, 760
\I
I
A
X
200J
1
I
ANALYTICAL CHEMISTRY, VOL. 42, NO. 7,JUNE 1970
Figure 3. Change in EB for uncharged bases in the binary solvent of mole fraction (X)in acetic anhydride The bases (all at 0.0040M) are: (1) triphenylarsine; (2) triphenylamine; (3) triphenylphosphine; (4) pyridine; (5) N,N-diethylaniline; and (6) tribenzylamine
added, then EB should still provide at least a qualitative parameter for comparing the relative strengths of bases. Thus, at a given fixed solvent composition, the first three terms in Equation 2 are constant, and differences in the measured acetate ion concentration for bases (having the same stoichiometric concentration) will reflect the differences of their basicity constants in that particular mixed solvent. In Figure 2 the trend in EB for the metal acetates with increasing mole fraction of acetic anhydride is again a shift toward a more positive emf, and a large change occurs only above a mole fraction of about 0.85 anhydride. The qualitative order of decreasing basicity, KOAc > NaOAc > LiOAc, found in glacial acetic acid remains the same in all solvent mixtures. However, the quantitative differences in apparent ~ K values B are slightly increased by the addition of acetic anhydride. For exampe, in HOAc [ ~ K L ~ o APKN~OAJ ~ is 0.21 and - ~ K K o Ais~ 0.43; ] the corresponding differences in solvent of mole fraction 0.87 (anhydride) are 0.84 and 0.72, based on Equation 2; and at a mole fraction of 0.94, the differences are 1.08 and 0.88 unit. For the uncharged proton acceptors, the same type of positive shift in emf is observed in Figure 3 as occurs with the metal acetates. However, there are differences between the results for the charged and uncharged bases which should be noted. When comparing a pair of bases having essentially the same pKB value in pure acetic acid (ex. KOAc and pyridine), the corresponding curves in Figures 2 and 3 cannot be exactly superimposed. The ionic base has a steeper slope, shifting to a much more positive emf at higher mole fractions of anhydride than does the uncharged base. (As an example, at X = 0.94 the E g (KOAc) is 236 mV and that for pyridine is 290 mV for 0.0040M solutions.) When comparing structurally similar uncharged bases having the same acceptor atom (ex. tribenzylamine and diethylaniline), there is little or no difference between their relative pKB values in pure acetic acid and in the mixed solvents. The same qualitative order in Figure 3 for decreasing strengths in acetic acid is valid in the presence of large amounts of acetic anhydride ; however, for pairs of bases not having a common acceptor atom, the difference in the apparent pKB values in the mixed solvent is larger than in anhydrous acetic acid. For the very weak uncharged bases, a slight maximum will be noted in the ELIus. X curves, and the positive shift in potential is visible only after the acetic anhydride content exceeds the region of 0.3 to 0.4 mole fraction.
Half-Neutralization Potentials (HNP). The values for the half-neutralization potentials in a given nonaqueous solvent for bases titrated with a reference strong acid are frequently used as a measure of base strengths because of the empirical correspondence established between the pK~(H20)and HNP in nonaqueous media. However, Bruckenstein and Kolthoff (14) have emphasized that HNP's in anhydrous acetic acid can be compared only if the titrated bases have identical stoichiometric concentrations. In pure acetic acid the essential equilibrium for the neutralization of a solvated base by perchloric acid is that in Equation 3.
B "OAC
+ HCIO4 e HOAC
+ B"C104-
+ ClO4-
(3)
For the condition of half-neutralization of the base, CB = CBc10,,and the Nernst relationship for E" has the general form of Equation 4 ( I S ) . RT + E,) + RT - In Ks + - In ( K B + K B C I OL F 2F 2F
In K B -
300.-
F!
BHf
E H N= (Ecc'
400
RT - In (Ks 2F
+ K B C B ) (4)
(KBCIO~ is the ion pair dissociation constant for the salt, BH+Clod-.) In dilute solutions having CB N 1OPaMfor bases or larger, the term (Ks KBCB)approxiwith K B 'v mates to KBCB. Then, Equation 4 simplifies to give Equation
+
Figure 4. Comparison of EB and EHNfor a base as a function of mole fraction ( X ) of acetic anhydride in the binary solvent Base concentration is 0.0034M. Curve (1) EHNfor LiOAc; (2) EB for LiOAc; (3) EHNfor tri-n-butylamine; and (4) EBfor tri-nbutylamine
Figure 4 persists at all solute concentrations in the range, 10+ to and it is the only base showing this phenomenon for the B-BHC104 mixture. A quantitative statement of the relationship of the HNP to EB for a given base having the same concentration in each solution is obtained by subtracting Equation 2 from Equation 5 .
5.
E H N= (Eoc'
2F
+ E,) + RT In Ks +
RT RT In (KB + K B C I OJ - In K B - - In KBCB ( 5 ) 2F 2F
Upon comparison of Equation 5 to Equation 2, the mathematical similarities suggest that as a rough approximation the function, E" ES. X,for the binary solvent should exhibit a general trend analogous to the EB us. X curve for a given base. and EB are shown for identical In Figure 4 plots of both concentrations of the moderately strong bases, tri-n-butylamine and lithium acetate. Graphs of the same kind were prepared for sodium acetate and N,N-diethylaniline; however, KOAc, pyridine, and triphenylphosphine could not be used because their perchlorate salts precipitate in the half-neutralized mixtures. On dilution of a given base-salt system, a family of parallel curves like those in Figure 1 are obtained for E" us. X . The slight maximum in the HNP curve for LiOAc in (14) S. Bruckenstein and I. M. Kolthoff, J . Amer. Chem. Soc., 78, 2974 (1956).
Table 111. Difference in Potential As (E"
For pure acetic acid solutions, known ion pair dissociation constants for the base and its perchlorate salt can be used to calculate (E" - EB). Using reliable published constants (13) for two representative bases, one computes this difference to be 44 mV for LiOAc and 33 mV for NaOAc. Although Equation 6 is of little importance when applied to anhydrous acetic acid solutions, it has significance for the mixed solvent, acetic acid-acetic anhydride. If the sum of the RT - In Ks), changes in an three quantities, (EGc' E3 F identical manner as a function of X for both the base and halfneutralized base (at the same concentration), then the difference, (EH,v - EB), will be essentially constant over a wide range of solvent composition. The experimental results in Table 111confirm this prediction for the stronger bases. The analytical consequence of this extension of Equation 6 to the binary solvent is that the half-neutralization potentials for a series of bases determined with the same reference strong acid (and at identical C B )are valid measures of the basicity constants for the bases in the mixed solvent. The empirical correlation of pK,(H20) to HNP(Ac20) made by Streuli (2) gave two separate linear plots for the
+
+
- E B )at a Fixed CB(0.00133M) in Acetic Acid-Acetic Anhydride
AmV (Emf) in solvent mixtures having mole fractions of AGO: 0.08 0.19 0.29 0.44 0.56 0.70 0.87
Half-neutralized base 0.00 0.02 0.94 Mean 44 46 48 49 LiOAc 47 45 48 41 44 57. 47 i 2 27 28 24 21 19 18 16 14 15 18 17 i 2 a NaOAc 14 14 15 16 N,N-Diethylaniline 15 13 13 13 15 ... 14f. I 7 8 6 6 Tri-n-butylamine 8 8 9 10 9 9 8&2 a First three values are excluded from the computed mean because of slight increase in the absolute value of EHNin the presence of small amounts of acetic anhydride.
ANALYTICAL CHEMISTRY, VOL. 42, NO. 7, JUNE 1970
761
trode pair in the mixed solvent without an added solute (curve 2, Figure 5). Clearly, the minimum is associated with a change in the indicator electrode response to the activity of the protonic species upon the addition of small amounts of Ac20 to acetic acid, since this feature is not present in the corresponding curves for the bases and half-neutralized bases. In extending the implications of Equation 7 to the binary solvent it is assumed that the same protonated species predominates in the mixed solvent as in pure acetic acid, namely, [CH~COOHZ+] and its perchlorate ion pair. Obviously, this is not correct, since acetic anhydride can compete as a proton acceptor in solutions containing strong acids. CH3COOHzf I
0
r
I
X
I
Figure 5. Curve 1-Change in emf for the electrode pair in 0.004M perchloric acid with mole fraction ( X ) of acetic anhydride. Curve 2-Change in emf for the electrode pair in the binary solvent alone
ionic and uncharged bases. The fact that the change in slope for EB us. X(Figure 2) is more rapid for charged bases than for uncharged species (Figure 3) is consistent with the conclusion of Streuli. The normalized HNP data of Pietrzyk (7)confirms the trend shown in Figure 4, namely, that the HNP decreases in magnitude as the acetic anhydride content of the solvent increases. Electrode Response in Perchloric Acid Solution. The potential of the glass-calomel electrode pair in solutions of H C I O ~in anhydrous acetic acid is given in Equation 7 (14).
When compared to the EBresult in Equation 2, the only differences in the mathematical forms for the two equations are that Equation 7 lacks the constant term for the solvent autoprotolysis constant and the slope of EHCIO,, us. In CHCIO, is positive rather than negative. If the electrode pair is to be applied to acetic acid-acetic anhydride mixtures containing perchloric acid, one would would change in much the same way as anticipate that EHCIO, EB with the exception that the direction of the shift would be opposite to that of EB. The experimental results for E H C I O ~ us. X (anhydride) shown in Figure 5 conform to this expectation only at a solvent mole fraction greater than 0.20 to 0.25. The minimum in the curve for perchloric acid occurs at a mole fraction of 0.075 acetic anhydride and is found for the elec-
XAmO
0.000
0.024 0.076 0.131
0.191 0.288 0.436 0.563 0.708 0.875 0.936
762
+ (CHpC0)zO e (CH4C0)20Hf
+ CH3COOH
Alternate acidic species have been discussed by Pietrzyk (7). The concentration term for perchloric acid in Equation 7 was tested by making emf us. log C plots for mixed solvents having mole fractions of 0.022, 0.075, and 0.103 anhydride. Although the linear graphs were not highly precise, the slopes were much less than the theoretical value at 25°C and ranged from 0.020 to 0.024 volt. Role of the Solvent. The simplest model for assigning a role to the aprotic solvent is the Born expression. The Born relationship for the electrostatic contribution in the free energy change to E" has been reviewed by Amis (15) and by Popovych and Dill (16). An essential consequence of the model is that E" is proportional to the reciprocal of the macroscopic dielectric constant (D)of the binary solvent. Because of the many variables operating in the present acetic acid-acetic anhydride system, no attempt was made to evaluate Eoco. However, if electrostatic effects are dominant, the measured potentials should show decisive trends paralleling changes in the dielectric constant of the medium. For example, HNP values for phenols correlate qualitatively with the dielectric constant for a wide range of solvent types (17). Representative data are given in Table IV for uncharged and ionic bases. Dielectric constant values at 25 "C were published previously (18). From a survey of Table IV, it is clear that the magnitude of the cell potential decreases by only a few millivolts while the dielectric constant increases from (15) E. S. Amis, J. Electround. Chem., 8,413 (1964). (16) 0.Popovych and A. Dill, ANAL. CHEM., 41, 456 (1969). (17) A. Fischer, G. Leary, R. Topsom, and J. Vaughan, J . Chem. SOC.( B ) , 1967, 846. (18) 0. Kolling, K . O'Hara, and T. Stevens, Trans. Kuns. Acad. Sci., 66, 435 (1963).
Table IV. EMF Data for Bases as a Function of Dielectric Constant of Mixed Solvent Cell potential (in mV) for 0.00133M solutions Tri-n-Butylamine Diethylaniline LiOAc D EB EHN EB EHN EB 430 405 392 6.2 394 40 1 429 404 391 7.3 393 398 428 403 388 8.8 390 396 426 402 385 388 10.1 394 426 398 382 392 385 11.3 422 393 378 380 13.6 388 41 1 380 368 368 15.3 376 395 368 355 355 18.2 364 3 76 351 332 342 339 20.2 3 52 ..* 318 ... 304 20.9 304 ... 256 21.1 266
ANALYTICAL CHEMISTRY, VOL. 42, NO. 7, J U N E 1970
EHN 474 475 476 477 475 469 456 443 423 40 1 357
Table V.
3001 0.5
0
Experimental Slopes for Linear Functions in Figure 6 Bases (0.0040 M) Slope, mV Lithium acetate 109 Sodium acetate 129 Potassium acetate 121 N,N-Diethylaniline 111 Pyridine 121 Tribenzylamine 118 Tri-n-butylamine 120 Triphenylphosphine 112 Half-Neutralized Bases (not included in Figure 6) Lithium acetate 111 Sodium acetate 118 N,N-Diethylaniline 107 Tri-n-butylamine 105 Mean 116 f 7 (std dev)
-LOGX Figure 6. Cell potential as (E& related to the logarithm of the mole fraction ( X ) of acetic acid in the mixed solvent. The bases are in descending order: triphenylphosphine, LiOAc, NaOAc, KOAc, pyridine, N,N-dieth ylaniline, and tribenzylamine
roughly 6 to 11 ; and the region of largest change in emf is in the D-range from 18 to 21. A plot of either EB or ERN us. 1/D (not included herein) is highly curved. Thus, like the acetic acid-dioxane system, the influence of the aprotic component on measured potentials of basic solutes in acetic acid-acetic anhydride is not primarily an electrostatic one (IO). Although the total polarization us. mole fraction curve for HOAc-Ac20 does not have the unusual features of p-dioxane mixed solvents (19), important solvent-solvent interactions are operative. Quantitative NMR results for the equilibrium in Equation 8 under anhydrous conditions indicate that the formation constant of the hydrogen bonded complex (HOAcAS) HOAC AS e HOAC-AS (8)
+
is considerably larger when the aprotic solvent (AS) is p dioxane than it is for acetic anhydride (20). However, solvent-solvent interactions of this kind are more important factors in competitive equilibria involving perchloric acid than with basic solutes. An alternate description of the function of the added acetic anhydride on the cell potential for base solutions is that proposed by Feakins and French (21). Although used originally for aqueous-organic solvent mixtures, their thermodynamic arguments are general and are based on the assumption that the dominant effect exerted by the addition of varying amounts of the minor solvent component is the alteration of the E" value of the reference electrode. When applied to the acetic acid-acetic anhydride system, the Feakins-French relationship becomes Equation 9,
where the subscripts, s and HOAc, refer to the HOAc-AcsO mixture and to pure acetic acid, respectively. For a dilute solution of a base in the binary solvent, Equation 2 can be generalized to give Equation 10. (19) 0. Kolling and C . VanArsdale, Trans. Kans. Acad. Sci.,68,65 (1965). (20) M. Muller and P. Rose, J. Phys. Chem., 69,2564 (1965). (21) D. Feakins and C . French, J . Chem. SOC.,1957,2581.
+ + RTF In K, +
( E B =~ (EOC')HOA~ Ej
-
If this result is correct, ( E B )should ~ be proportional to the logarithm of the mole fraction (HOAc) and have a slope of 0.059 at 25 "C. The graph of experimental data is shown in Figure 6 . Linear plots were obtained for the mole fraction range 1.0 to 0.28 in HOAc; however, at higher acetic anhydride mole fractions the points scatter to a considerable degree. In a few instances the data at the mole fractions from 1.O to 0.9 deviate slightly from the curve. The empirical slopes are listed in Table V. It will be noted that the mean experimental slope is essentially twice the theoretical value of 59 mV; and it follows that (EB), == RT/F In X H O A ~This ~ . result is compatible with the monomer-dimer equilibrium for acetic acid in acetic anhydride (20),as given in Equation 11,
XD
2 HOACe (HOAC)~;KO = - = 103 X M 2
(11)
in which M and D refer to the monomer and dimer, respectively. In effect the solvent parameter for acetic acid in this system must be calculated as the dimer because of the size of KD. RT (EB)~ = (EOC')HOA, Ej -In K, F
+ +
+
(XHOA~' refers to acetic acid expressed as the dimer.) The same requirement applies to (EHN)#, as well. Then, the mean experimental slope becomes 0.058 V for the mole fraction range of 1.OOto 0.40 HOAc dimer. It is certain that Equation 12 is approximate relationship, since KB, K,, and Ej will be subject to changing electrostatic and solvation influences with changing solvent composition. The fact that there are significant variations in the slope of the mole fraction term among the several bases studied attests to such complications. On the other hand, the acetic acid dimer composition is the predominant variable influencing EB and EHNover a very broad range for an increasing amount of acetic anhydride in the binary solvent.
RECEIVED for review February 16, 1970. Accepted April 13, 1970. This project was supported financially by the National Science Foundation Grant, GP-10015. ANALYTICAL CHEMISTRY, VOL. 42, NO. 7, J U N E 1970
763
