L-AROMATIC AMINESAS HAMMETT BASES
July 5, 1964 [CONTRIBUTION
NO. 1223 FROM
THE
DEPARTMENT OF CHEMISTRY,
UNIVERSITY OF PITTSBURGH, PITTSBURGH
267 1 13, PENNA.]
Solvent Effects in Organic Chemistry. IV. The Failure of Tertiary Aromatic Amines as Hammett Bases BY EDWARD M . ARNETTA N D GEORGEW. MACH' RECEIVEDDECEMBER 26, 1963 T h e applicability of the Ha function to the protonation behavior of a series of fourteen tertiary amine indicators is tested in aqueous sulfuric acid solutions. Despite the structural similarity of the primary and tertiary amines, i t is found that t h e latter generate their own acidity function (Ha"') falling midway between t h a t for primary anilinium ions ( H o ' )and t h a t for arylcarbonium ions ( H R ' ) . Apparent average hydration numbers for the primary and tertiary anilinium ions are estimated by Taft's treatment and the difference between them is found to vary considerably with acid strength. A4rough accounting for the difference between HO' and Ho"' is given using Boyd's activity coefficients and some drastic assumptions. Attention is called to the fact t h a t the present Hafunctions may only be used with rigor for nitrated anilines in view of t h e specific influence of the nature of the basic group and other substituents (especially the nitro group) on activity coefficients of the free and protonated base. Furthermore, the fact t h a t activity coefficient behavior may be explained in terms of specific hydration of the weak base cations by no means requires t h a t this is actually the sole or even the main factor. Since the behavior of even rather similar ions and molecules cannot be predicted a t present, it is doubtful t h a t a simple solvation theory for widely differing types of BrZnsted bases and their ions will be possible. This cawed applies with even greater force against t h e interpretatioxi of transition state behavior. It is suggested, however, t h a t an exact treatment of acid-catalyzed reactions may soon be possible in terms of the activity coefficients of model ions, molecules, and activated complexes.
For over 30 years the main basis for discussion of organic processes in strongly acid solutions has been the H o acidity function, invented and developed by Hammett and his s t u d e n b 2 The Ho function provides an operational means for estimating how the degree of protonation of a Brgnsted base changes as a function of acidity in acidic media of sufficient strength to preclude the use of a pH meter. This is obviously of tremendous value for the quantitative treatment of acid catalysis in strong a ~ i d s ~and - ~ detailed but controversial proc e d u r e ~ ~have , ' been developed using correlation of reaction rates with Ho as a criterion for reaction mechanism. If all Brgnsted bases show similar dependence of the activity coefficients of their free base VB) and conjugate acid ( ~ B H + ) forms through a given region of acidity, then i t is possible to use the H o function as a thermodynamically rigorous scale for referring an equilibrium constant measured for the indicator in acid of strength Ho back to the pKa that its conjugate acid would have in pure water (the standard state). This is so because Ho is defined operationally and theoretically as
Ho = pKa - log Q = -log ( ~ H + ~ B / ~ B H +(1) ) where Q = (BH+)/(B) is the measured indicator ratio. The test of this assumption is whether the difference between log Q values for indicators of comparable basicity remains constant through a given range of acidity since it is easily seen from eq. 1 that pKal - pKa2 = log Q' - log Q' (2) for any given value of Ho. For the nitrated anilines used by Hammett and Deyrup2 this requirement was met quite well for sulfuric acid solutions up to about 65 wt. yo. At that point a tertiary aniline was employed in contrast to the primary ones which had (1) National Science Foundation Cooperative Fellow, 1963-1964. This work was supported in p a r t b y N . S. F. Grant (3-14583. (2) L. P. H a m m e t t and A. J. Deyrup, J . A m , Chem. Soc., 64, 2721 (1932). (3) L. P. H a m m e t t , "Physical Organic Chemistry," McGraw-Hill Book Co., Inc., New York, N. Y . , 1940. (4) M. A. Paul and F. A . Long, Chem. Rev., 67, 1 (1957). ( 5 ) E. M. Arnett, "Progress in Physical Organic Chemistry," Val. I , S. G . Cohen, A. Streitwieser, Jr.. and R. W. T a f t , Ed., Interscience Publishers, Inc., New York, N. Y . ,1963. (6) F. A. Long and M. A. Paul, Chem. Reu., 6 7 , 935 (1957). (7) J. F. B u n n e t t , J . A m . Chem. Soc., 83, 4956 (lQ61).
been used (but for one secondary amine and an azo compound) for building the scale through dilute sulfuric acid solutions. The rest of the Ho scale in strong acid was mostly developed with ketones and it was noted that eq. 2 was not followed so well. This was attributed to the influence of serious medium effects on the spectra rather than to a failure of the fundamental assumption that the activity coefficient ratio fB/fBHi was independent of the base involved. Paul and Long,4 however, suggested that the discrepancy might be due to a breakdown of the activity coefficient postulate and that Brgnsted bases of widely different size or structure might follow different acidity functions even though they were of the same charge type. More recently Taft8 has focused attention on differences between the acidity function behavior of this same tertiary amine and a secondary aniline in contrast to the primary members of the Hammett indicator series. Also, Boyd'ss direct measurements of the activity coefficients of anilines of different degree and their ions in acid of varying strength leave little doubt that tertiary amines may be expected to follow a somewhat different acidity function than that for the primary anilines. where the proThere are by now numerous tonation equilibria for various kinds of Brgnsted bases do not follow the traditional Ho function4 exactly, in the sense that a plot of log Q v s . Ho does not have the unit slopez0 required by eq. 1. Most of these cases ( 8 ) R. W. T a f t , Jr., ibid., 81, 2965 (1960). (9) R. H . Boyd, ibid., 86, 1555 (1963). (IO) A. J. Kresge, G . W . Barry, K . R . Charles, and Y. Chiang. i b i d . , 84, 4343 (1962). (11) W . M. Schubert and R. H. Quacchia, ibid., 86, 1279 (1963). (12) R . I.. Hinman and J. Lang, I'elrahedron L e l l e r s , 11, 12 (1960). (13) F. A. Long and J. Schulze, J . A m . Chem. Soc., 83, 3340 (1961) (14) A. R. K a t r i t z k y , A . J. Waring, and K Yates. Tetrahedron, 19, 465 (1963). (1.5) A. R. Katritzky and A. J. Waring, J. Chem. Soc., 1540 (1962); 3760 (1963). (16) J. T. Edward and I . C. Wang, C a n . J . C h e m . . 40, 966 (1962). (17) R. H. Moodie, P. D. Wale, and T . J. Whaite, J. Chem Soc., 4273 (1963). (18) Y . Chiang and E . B. Whipple, J . A m . Chem. Soc., 8 6 , 2763 (1963). (19) E. M . Arnett and C. Y. W u , ibid., 84, 1680 (1962). (20) Since this is a log-log plot i t is not surprising t h a t straight lines are almost invariably obtained. I n fact, there are probably few areas of physical organic chemistry where there are so many fortuitous and probably meaningless linear relationships as t h a t of acidity function phenomena.
EDWARD M. ARNETTA N D GEORGE W. MACH
2672
Vol. 86
TABLE I PHYSICAL DATAA N D METHODSOF PREPARATION O F ISDICATORS USED IN THISSTUDY Compound
Reference for preparation
Name of compound
1 2 3 4 5 6
li,X-Dimethyl-4-nitroanihe h7,S-Diethyl-2,4-dinitroaniline K-(2,4-Dinitrophenyl)piperidine
i
S-Methyl-4-nitrodiphenylarnine S-Methyl-4'-bronio-4-nitrodiphenylamine
N,i'-Dimethyl-2,6-dinitro-4-nietIiylaniline r\',S-Dimethyl-2,4-diriitro-l-naplithylamine S,S-DimethyI-4-cliloro-2,6-dinitroaniline
..
8 9 10 11 12 13 14
24 25 26 27 28 29 30 31 32 33 34 31 35 36
----"l.p Lit.
163 80 92 95 88 111 68.5 113 5 163 167
, 'C
Obsd
164-165 79-80 93-94 96-97 87-88 111-112 69-70 114-115 165-166 169 5 140-141 140 162-163 210
ham ~r ",
423 393 395 445 430 452 420 417 5 402 397 385
N, S-Diethyl-%,4,6-trinitroaiiiline N-Methyl-2,4-dinitrodiplienylamine S ,S-Dimethyl-2,4,6-trinitroaniline N-Methyl-4-bromo-2',4'-dinitrodiphenylamine 110 400 N-Methyl-x ',4 '-dibromo-2,4-di1iitrodiphenylamine 382 N-Methyl-i,4,2',4 '-tetranitrodiphenylamine 210 410 a This value is that observed in the pK,, region of acidity For most compounds it scarcely differs from t h a t in water t h a t H n " ' is a thermodynamically exact acidity function
however involve molecules of quite different size or groups of quite different structure from the anilines used by Hammett. In view of the fact that the primary aniline scale has recently been subjected to a careful re-examinations1 with modern instrumentation, ? 2 we felt that it would be valuable to develop for comparison an Ho scale based entirel;; on tertiary amine indicators. Any differences between the two acidity functions should give a clear indication of how serious an error may be made by assuming that the usual H c scale may be used for molecules of slightly different structures. Furthermore, since the differences in activity coefficient behavior of primary, secondary. and tertiary amines have been attributed to different degrees of specific hydration of the corresponding ammonium ions,8 the divergence of the two acidity functions might supply values which would be useful for quantitative t.reatments of solvation in aqueous sulfuric acid solutions. We report here our results for fourteen carefully chosen tertiary amines. Because of the base-strengthening effect of steric inhibition of resonance, it was not possible to use N,N-dialkylated anilines as indicators in strongly acidic media; for example, N,N-dimethyl2,4,0-trinitroaniline is half-protonated in GGYo H2S04. A number of .N-methyldiphenylamines were therefore employed. &-e have considerable reluctance against burdening the literature with a new acidity function. Nonetheless, as will 'be -seen below, the fact is that tertiary amines generate one of their own and in order that it shall not be confused in this paper with the generally accepted Ho scale" for primary anilines we shall name i t H,"' for use in the following discussion and call that of Jorgenson and Hartter Ho'. Experimental All of the compounds except indicator 13 were prepared essentially according to published descriptions. These are listed in Table I . In several cases we moderated the reaction of polyni trochlorobenzcnes with amines by dilution with alcohol. Occasionally wc used different solvents for recrystallization from ~
(2 I! R l i1 ! l i j :i)
.
J
~
.
-
~
Ji,igrnson and I )
~
R
Hal ttel-, J
Am
Chrii? Soc
,
85, 878
( 2 2 ) H a m m e t t and 1)ey;up made theiv sixctral measurements with a ci-udc COIOI imeti ic piocedul-e t h e best availnhle in the ea]-ly 1 9 ~ 4 0 ' s . T h e f a c t t h a t f u i the>-studies with modern ecluipment have confirmed mo*t of theii- measui-ements is a hiQh tribute t o theii- caie a n d technique
pKah
+0 +0 -0 -1 -2 -3 -3 -4 -5 -6 -6 -6 -8 -10
66 21
38 66 59
12 4% el
71 19 55 93 1; 56
Assuming
those described in the original articles. Aside from these mincr changes the methods were quite straightforward. N-Methyl-r',4'-dibromo-2,4-dinitrodiphenylamine (Indicator 13).--To a stirred suspension of 10.9g . (0.040mole) of S-tnethyl2,4-dinitrodiphenylaminein 100 ml. of glacial acetic acid was added 6.4 g. (0.040mole) of bromine. The solid dissolved to give a clear red solution from which material soon precipitated and another 6.4g. of bromine was added. The mixture was brought to reflux over the period of an hour and was refluxed for another hour. Upon cooling, 11 .O g. of mixed yellow and reddish orange crystals was obtained. After six recrystallizations from acetic acid and t w o from ethanol, 3.0g . , a 175~; yield, of yellow prisms melting :it 162-163" was obtained. Anui. Calcd. for C13HYS3BrcOd:C , 36.22; H, 2.10;Br, 37.08. Found: C, 36.25, 36.45;H , 2.16,2.04;Br, 36.96,37.28. Spectrophotometric Measurements.-Standardized sulfuric acid solutions were prepared and handled according to methods described by hrnett and n'uZ3and indicator solutions were also prepared in the way described there with the exceptions of the following small details. Firstly, indicator stock solutions were prepared in anhydrous methanol or glacial acetic acid and introduced into the aqueous sulfuric acid solution with a 0.25-1x11. Hamilton microsyringe ( S o . 725) equipped with a Chaney adapter to deliver 0.200nil. .in independent experiment showed that an optical density of 0.671 f 0.001could be obtained with this instrument for three replica measurements. As is usual in this field, insolubility of some of the indicators was a problem. T o deal with low solubilities, 10-cm.cells were used routinely in the Cary-14 spectrophotometer, the long path length and large molar absorptivity indices of most of the conipounds allowing the use of highly dilute solutions ( 10-5-10-6 ,U). The long path-length cells, however, required preparation of 50-ml. samples with considerab1.e consumption of standard acid. Often it was necessary t o warm the acid after the indicator had been added t o it and agitate vigorously in order t o effect complete solution as shown by a Tyndall beam. In these cases there was no indication of reprecipitation after cooling to rooni temperature (26 =t2 " ) a t which all measurements were made. Several other indicators not mentioned here were rejected because of insolubility. Instability was not a problem as was shown by repeated scans of the acid solutions of t h e indicators. ( 2 8 ) E hf. Ai-nett and C. 1'. Wu. J A m , Chrm. Soc , 82, 5660 (ICJ60). (21) V Me)-tz and W Weith. Be? , 10, 746 (1877) (2.5) P. van Romburgh. Rec lrnv. chrm , 2 , :
