Sept., 1959
1457
RACEMIZATION OF A BIPHENYL HAVING ANIONIC BARRIER GROUPS
SALT EFFECTS IN THE RACEMIZATION OF A BIPHENYL HAVING ANIONIC BARRIER GROUPS1 BY J. E. LEFFLER AND B. M. GRAYBILL Contribution from the Department of Chemistry, Florida State University, Tallahassee, Florida Received February B 1 , 1969
The rates and activation parameters for the racemization of the salts of 2,2'-dimethoxy-6,6'-dicarboxydiphenyldepend on the solvent and on the concentration of added salts. Salts added to the aqueous solution decelerate the reaction by as much as a factor of 3.5,while the addition of organic solvents can cause a deceleration of a factor of ten or more. This is in contrast to the acceleration reported previously for the racemization of 0-(2-dimethylaminophenyl)-phenyltrimethylammonium ion under the same conditions. I n both reactions the added salts or solvent components lower the enthalpy and entropy of activation, but the rate effects in the reaction reported earlier are enthalpy-controlled, while those in the present reaction are entropy-controlled.
state of I. In the case of the biphenyl 11, the reSalt and Concentration Effects in Water.-The racemization of the ion I is decelerated by an in- verse was postulated; the transition state was concrease in the concentration of its disodium salt or by sidered to be de-solvated and hence able to interthe presence of added salts. The reduction in rate is act more strongly with the associated anion. We would like to suggest two reasons for the differ0ence in behavior of the two biphenyls. The OCHI h=O first is that the ion pair formed by the ground state of I with a cation is a chelate, or a t least that I it is stabilized relative to transition state ion pairs because of the lesser distance between the carboxyl-0 OCH, ate groups. The second possibility is that the I transition state for the racemization of I may be 0vibrationally less excited, and therefore less desolentropy-controlled and amounts i n some cases t.0 a vated than the transition state for the racemization factor of 3.5, while the activation enthalpies and of 11. entropies are generally lowered by the added salt, TABLE I by as much as 3.6 kcal./rnole and 11.6 cal,/moleOF CONCENTRATION ON THE RACEMIZATION OF deg. Metal ions have an effect which appears to THEEFFECT SALT OF 2,2'-DIMETHOXY-6,6'-DICARBOXYBIincrease with decreasing size or increasing charge of THE DISODIUM PHENYL IN WATER the ion. However, very large cations (detergents) Biphenyl are approximately as effective as lithium or sodium concn k X 106, sec.-*b AH* AS ions. Detergent anions do not appear to have any rnole/i: 100' 79.4O koal./mhea oal./moleldeg. marked effect. 0.0055 14.05 1.83 24.99 f 0.08 -14.22f0.22
+e
*
The behavior of the dicarboxylate ion I is quite different from that of the cationic biphenyl I1 reported earlierlb,c. Salts and organic solvent components added to aqueous solutions of I1 were found to accelerate, rather than decelerate CHS-N-CHI
a-p N +-CH3
I1 d,'CH, the racemization. The cationic biphenyl I1 also differs in that large ions were found to have a greater effect than small ones. Furthermore, the charge type of the ions added was found to be relatively less important than in the case of the present biphenyl. The single point of resemblance in behavior between the two biphenyls is that the added salts or solvent components in both cases lower the enthalpy and entropy of activation with respect to that in pure water. The difference in'direction of the effect on the rates is due to the fact, that the previously reported reaction is enthalpycontrolled rather than entropy-controlled.
The salt effects in the racemization of I can be explained in terms of a stronger association between the added cation and the ground state of I, than between the added cation and the transition (1) (a) Based on the doctoral dissertation of B. hiI. G. and presented in part at the Symposium "Solvent Effects and Reaction Mechanisms," Queen Mary College, London, July 1957. For previous papers of this series, see: (b) W. H . Graham and J. E. Leffler, THIS JOURNAL, 68, 1274 (1959) and (c) J. E. Leffler and W. H . Graham, i b i d . , 63, 87 (1959).
. 0070
13.8 1.80 25.29 f .OO -13.45f .23 .0134 1.3.3 1.77 24.72 zk .10 - 1 5 . 0 3 f .28 ,0228 13.2 1.75 24.95 f .I1 -14,433~ .32 0339 12.9 1.73 24.83 3Z .15 -14.664~ .41 ,0662 12.0 1.61 24.85 f .07 -14.90f .20 ,103 11.2 1.50 24.78 f .06 -15.223Z .22 ,005" 14.3 1.94 24.88 f .07 -15.85f .21 . 005d 14.3 1.84 25.23 f .06 -13.55k .I6 a For the significance of t,he error figures and methods of calculation see reference IC. Mean rate constants. 0 The dipotassium salt. The ditetramethylammonium salt.
Models indicate that a chelate structure for the ion-pair made from the ground state conformation of I is sterically quite reasonable. The ring formed by the chelation is a large one, but the ready formation of large rings by aromatic compounds is not unusual, since rigid molecules lose comparatively less entropy by ring closure than do flexible ones. -O--C=O
- I- =
A+ I \
C-0
//
-
0
The chelate ion-pair
The role of the metal ion in the proposed chelate
1458
J. E. LEFFLERAND B. M. GRAYBILL
Vol. 68
TABLE I1 THI EFFECT OF ADDEDSALTS ON THE RACEMIZATION O F THE DISODIUM SALT OF 2,2'-DIMETHOXY-6,6'-DICARBOXYBIPHENYL I N WATER Added substance and ita concn., moieji."
None NaOH
14.05 14.05 13.5 12.3 12.1 11.5 10.9 11.3 8.93
0.0148 ,0302
,0500 ,0811
NaCI
,1800 ,3134 .2 .5 .75
8.48
1.0 LiCl
7.71 11.4 8.31 5.50 4.30 7.37 6.27 5.16 4.01 6.97 8.26 11.45
0.075 .20 .74
BaClp
MgClz
k X 108,seo.-lh 1000 79.4'
1.00 0.01 .02
.05 .10 -01
1.83 1.83 1.77 1.69 1.65 1.56 1.49 1.59 1.27 1.16 1.04 1.40 1.10 0.75
0.65 1.19 1.00 0.906 0.596 1.17 1.18 1.53 1.53 1.50 1.91 1.85 1.89 1.80 1.69 1.19
*,
A H * , kcal./molec
AS
24.99f .08 24.99f 0.08 25.11f .06 34.46f .10 24.57i .10 24.45f .ll 24.78f .08 24.24i -05 24.04f .13 24.56f .12 21.80f .IO 25.18i .18 24.96f .ll 24.62f .09 23.25f .08 22.50 f .12 22.65f .10 21.40f .10 23.73f .I5 22.01f .14 24.0if .10 24.88f .12 24.56f .IO 24.35f .05 24.88f .10 24.91 f .I1 25.19 f .08
-14.22 f . 2 ' ! -14.22 f 0.23 -13.95 f .16 -15.90 f .27 -15.62 f .26 -16.05 f .31 -15.22 f .22 -16.643~ .12 -17.61 f .37 -16.36 f .33 -15.91 f .27 -14.12 i .33 -15.32 f .32 -17.05 f .24 -21.23 f .23 -22.17 f .32 -22.08 f .27 -25.83 f .28 -20.03 f .43 -23.58 f .36 -17.72 f .27 -14.91 f .35 -15.81 f .26 -16.44f .I4 -14.46f .27 -14.46 f .30 -13.61 f .22
cal.jmole-deg.
1 .o NaBr 0.2 NaOAc .2 11.2 Sodium tosylate .2 10.8 (CHs)dN+C11.0 14.3 CeHbN +( CH3)sCI- 1.0 13.85 Polydiallyldimethylammonium bromide, 20 g./l. 14.4 Sodium lauryl sulfate, 0.11M 13.05 .......... ............ Dodecyltrimethylammonium chloride, 0.2Af 11.75 23.99 f .10 -17.23 f .28 Cetyltrimethylamnionium chloride, 0.2&f 8.53 24.32f .15 -17.00f .43 a The coticentration of the disodium salt of 2,2'-dimethoxy-6 6'-dicarboxybiphenyl is approximately 0.005mole/l. Mean rate constants. For the methods of calculation and the sidificance of the error figures, see reference IC.
KCI
is somewhat like that of a hydrogen atom in a hydrogen bond.3 Partial neutralization to the monocarboxylate ion (Table 111) does in fact have an effect on the rates similar to that of added salts, although the deceleration might also be expected as a result of the loss of the intraionic repulsion of the two carboxylate groups.
in the activation process and if both of the ions thus formed are considerably more solvated than the ion-pair. Solvation of the dicarboxylate transition state will he favored if the activation energy for this racemization is largely potential rather than kinetic. An indication that this might indeed be the case is the fact that trinitrobenzene accelerates the racemization of the related dimethyl TABLE I11 ester. The formation of a-complexes is favored THE RACEMIZATION O F THE MONOSODIL~M S A L T OF 2,2'- by coplanarity and conjugation in the substrate DIMETHOXY-6,G '-DICARBOXYBIPHENYL IN WATER and would seem to be difficult in a molecule having ksea. X -106, la AH* As *,tal./ free internal r ~ t a t i o n . ~Evidence suggesting conjuConditions 1000 79.40 kcal./mAleb niole-deg. gation between the two rings in the transition state The acid plus NaOH, is the acceleration produced by the substitution both 0.014 AI 2.73 0.433 22.71 f 0.11 -23.57 f 0.32 X = CH30- ill the biphenyl 111.5 The disodium salt, 0.005 U .plus NHdCI, 1.0hf
3.77 0,739 20.01
&
significmce Mean rlttjc of error coiistatits. figures, *see Forref. Inetlioda lc.
,IO
-30.1F & .27
of c:dcul;ttjioti nntl
The decreased entropies of activatioii for the ion-pairs can be explained if the ion pair dissociates (2) Tri-o-thymotide is an example. W. Bake:, B. Gilbert and W. D. Ollis, J . Chem. Sac. 1443 (1952): A. C. D.Newman and E. M. Powell, ibid., 3747 (1952) (3) Spectroscopic evidence indicates the existence of intermolecular lithium bonds. A. Rodionov, D. Sliiaorin, T. Talelaeva and K . ICochesl:oe, Izuestia Akad. Nauk S.S.S.R., Oldel. Chirn., 120 (1957).
+-& NO2
&OH
OCH,
I11
(4) Hindered biphenyls (in the grouid state) do not form complexes with TNB. C. E. Castro, L. J. Andrews and R . M. Keefer, J . Am. Chcm. Sac., 80,2322 (1958). (5) W. E . Hanford and R. Adams, i b i d . , 57, 1592 (1935). Sec also M. Calvin, J . Ow. Chem., 4, 256 (1939) and F. W.Cagle and H. Eyring, J . Am. Chcm. Soc., 73, 5628 (1951).
The Enthalpy-Entropy Relationships in the Salt Solutions.6-As the concentration of added salt is increased the reaction rate continues to decrease, but the enthalpy and entropy of activation pass through a minimum. However, the net effect is always a lowering of the enthalpy and entropy with respect to that of water. Figure I shows the behavior of the activation parameters with varying concentrations of sodium chloride. The Effect of Organic Solvent Components.Addition of alcohols or acetone to aqueous solutions of the dicarboxylate salt decelerates the reaction and decreases the enthalpy and entropy of activation (Table IV and Fig. 2 ) . Factors that might be expected to cause the deceleration are an increase in ion association and a medium effect on the intramolecular repulsion of the two carboxylate ion groups. The interpretation is complicated, TABLE IV THE EVFECTOF ORGANICSOLVENT COMPONENTS ON THE RACEMIZATION OF THE DICARBOXYLATE SALT I N AQUEOUS MIXTURES Solvent, wt. %,of organlc component CHaOHC 0.0 25.4 53.4 73.1 82.0 93.7 100 EtOHd 0.0 26.6 52.2 EtOH' Acetone/
100 0 0 12 9 24.9 62.7
1459
RACEMIZATION OF A BIPHENYL HAVING ANIONICBARRIER GROUPS
Sept., 1959
k X 1 0 6 , sec.-kQ
looo
13.8 12.95 12.2 9.16 5.66 3.15 1.43 14.3 13.6 12.0
AH* 79.4 koal./mheb 1 . 8 0 25.29 f 0.09 1.80 24.41 I: .lO 1.79 23.68 2Z .07 1.62 21.35 2~ . n 1.15 19.51 f .10 0.714 18.15 Zt .16 0.289 19.67 f .06 1.84 1.80 1.76
- .15
14.05 12.12 9.57 5.42
AS* tal./
moll-deg. -12.45 f 0.23 -15.91 f .28 -18.00 .20 -24.80 f .3n -30.71 f .26 -35.5 i .40 -33.00f .16
*
25.23 f .06 -13.55 f .16 25.01 f .10 -14.22 f . 2 8 23.70 f .07 - 1 7 . 9 7 2 ~ .21
..
. ....
,039 ... 1 . 8 3 24.99 f .08 1 . 7 0 24.32 f .09 1.35 2 4 . 2 8 4 ~ .10 0.812 23.47 2Z . l l
- 18
- 16
- 14
AS$.
Fig. 1.-Effect of added NaCl on the activation parameters for the racemization of the disodium salt in water.
partly racemized disodium salt tended to precipitate. This particular difficulty was not encountered with the disodium salt in absolute ethanol. However a troublesome yellow discoloration eventually developed in that solvent, and the approximate rates obtained were so extremely slow (about 1/70) compared to the reaction is water that we suspect the formation of aggregates of some kind even though no precipitation occurred.8 Experimental Synthesis of 2,2 '-Dimethoxy-6,6'-dicarboxybiphenyl.The synthetic scheme outlined below proved to be more convenient and gave better over-all yields than that previously reported.8
OH
............
-14.22 f -16.30f -1f3.88f -20.16 f
.22 .23 .26 .29
v
IV
a Mean rate constants. b For the method of calculation The and the significance of the error figures, see ref. l o . disodium mlt at 0.008 $1. The bis-tetramethylammonium The sodium salt a t 0.009 M . f The disalt a t 0.007 M . sodium salt at 0.005 211.
however, by the probability that the dicarboxylate ion partly solvolyzes to monocarboxylate ion in the more organic media. For example, in the case of phthalic acid there is a change in second ionization constant of three pK units as the solvent is changed from water to 80% ethanol.' Whatever the mechanism, the solvent effect is quite a large one. In 93.7% methanol the enthalpy of activation is reduced by 7 kcal./mole and the entropy of activation by 23 cal./mole-degree. The rate is reduced by more than a factor of four. I n pure methanol the enthalpy and entropy have risen somewhat over the values for 93.7% methanol, but the rate a t 100" is only about one-tenth that in water. It should be noted (Table 111) that the inonosodium salt racemizes a t about one-fourth thc rate of the disodium salt, both in water. For the rate studies in aqueous ethanol it was necessary to use the bis-tetramethylammonium salt rather than the sodium salt and to limit the solvents t o the more aqueous mixtures, because the (6) For a general discussion of enthalpy-entropy relationships see
J. E. Leffler, J. Org. Chem., 2 0 , 1202 (1955). (7) N. Mizutani, Z. physik. Chem., 118, 318 (1925).
VI
VI1
\*
b VI11
IX OCH, I
COOH I
9-p OOH OCH3
x
( 8 ) I n spite of several repetitions of the experiment we were unable to reproduce the much higher rate reported in the literature. The solutions were made by dissolving the acid in a slight excess of a freshly prepared solution of sodium in ethanol. C j . W. Stanley, E. McMahan and R. A d a m , J. Am. Chem. Soc., 66, 706 (1033).
1460 25
J. E. LEFBLER AND B. M. GRAYBILL
The 2-amino-3-methoxybenzoic acid (5 g., 0.03 mole) is ground in a mortar with 15 cc. ofoH20 and 10 cc. of concd. HCl, the suspension cooled to 0-5 ,and to it is added a solution of 2.9 g. (0.04 mole) of NaNOz in 35 cc. of H20during 15 to 20 minutes. The resulting diazonium salt solution is filtered if necessary and kept below 5' until used. The reducing agent is repared from 15 g. (0.06 mole) of CuS04.5H20 in 50 cc. of E20 and 25 cc. of 28% ammonium hydroxide. This solution is cooled to 10" and treated stirring) with a fresh, cold solution consisting of 4.25 g. [O.OB mole) of H2NOH.HCl in 15 cc. of H20plus 9 . c ~ .of. 6 N NaOH. A gas is evolved and the color changes immediately from dark to light blue. The diazonium salt solution is introduced, with stirring, below the surface of the reducing solution during about 20 minutes and the mixture stirred in its ice-bath for an additional 10-15 minutes. The solution is then boiled while 30
-
wAr70 0
G
'20 -
.
Vol. 63
NETHANOL
I
I
2,J-Dimethylanisole (V).-The phenol (89 g., 0.73 mole) and NaOH (30 g., 0.73 mole) are stirred in 200 cc. of water while dimethyl sulfate (46 g., 0.37 mole) is added during about one-half hour. After two hours of heating on the steam-bath, an additional 46 g. of dimethyl sulfate is added slowly and the mixture allowed to reflux for 15 hr. Alkali is added periodically in sufficient amount to keep the reaction basic. After cooling, the upper layer is extracted with ether, washed with alkali, dried over CaC12and vacuum distilled; b.p. 90' at 20 mm.; yield 70 g. 3-Methoxyphthalic Acid (VI).--& la King.9 3-Methoxyphthalic Anhydride (VII) .-The mixture of KCl and 3-methoxyphthalic acid obtained in the oxidation step is refluxed for 5 minutes in 45 cc. of acetic anhydride, Altered to remove KCl, refluxed an additional 2 hours and allowed to cool. The anhydride is then removed by filtration and washed with acetic acid. Additional product is obtained by adding ice-water to the filtrate;' yield 13 g. from 20 g. of 2,3-dimethylanisole, m.p. 160-161 3-Methyoxyphthalimide (VIII).-The anhydride (50 9.) is fused with 60 g. of ammonium carbonate a t about 220'. The heating and occasional stirring are continued until no more vmor is evolved and the melt is auiescent. The product,. crystallized from methanol, melt's a t 219-221"; yield 45 g. (90%). 2-Amino-2-methoxybenzoicAcid (IX) .-A 1.O N solution of NaOCl is prepared by dissolving 35 g. of Cln in one 1. of cold 10% aqueous NaOH. To 50 g. of the 3-methoxyphthalimide (0.28 mole) dissolved in 200 cc. of 10% NaOH is added 800 cc. of the NaOCl solution. The solution is then heated to about 70" for 45 minutes, the color changing from yellow to red-brown. The solution is then cooled, precisely neutralized with moderately concentrated HzS04 and cooled in ice. The first crop of crude brown crystals is removed and the filtrate extracted with ethyl acetate. The precipitated and the extracted fractions are combined and crystallized from aqueous ethanol; yield 28 g. (60%), light brown needles, m.p. 169-171", not depressed by a sample made by reduction of 3-methoxy-2-nitrobenzoic acid. 2,2'-Dimethoxy-6,6'-dicarboxybiphenyl.-The following method was found to be quite superior to the Ullman reaction.
.
(9) H. King, J . C h e n . Soc., 1157 (1939).
-The procedure of Adamss and T;rnerlo gives the le& sohble quinine salt, m.p. 176-177' (from acetone), [ ( u ] ~ ~ D $119.5' and the more soluble salt, m.p. 90-100" [ ( U ] ~ ~ D -56.7'. A mixture of the less soluble salt (3 g.) and 40 cc. of 5% NaOH is ground in a mortar, shaken vigorously and extracted with four 15-cc. portions of cold CHCl8 to remove the quinine. Acidification precipitates the l-acid a8 a white, spongy mass which after drying melts at 293-295', [ ( u ] . ~ ~ D 109". It may be clarified if necessary by treatment with charcoal in acetone. The d-acid, similarly prepared, melts at295-297', [ C Y ] ~ ~f104'. D Solutions for the Kinetic Experiments.-The disodium salt solutions of the 2,2'-dimethoxy-6,6'-diphenic acid are prepared by adding a 5% excess of standard NaOH solution to the acid and diluting to a diphenate ion concentration of 0.005 M . For the kinetic procedure, see reference IC. Error Analysis.-Based on a conservative estimate of f0.003' as the maximum error in the average of the twelve readings of the optical rotation a t each point, the expected erroP in an individual rate constant is about f 2 % . The tabulated rate constants are mean values computed from three or more individual rate constants. The expected errors in AH and A S + for a single pair of rate constants are f 0 . 3 kcal./mole and f 0 . 9 cal./mole degree. The tabulated activation parameters should be considerably better than that since they are the result of least squares calculations using all of the rate constants (six or more points). The tabulated measures of recision have the form of probable errors and are a by-profuct of the least squares calculation. Although the number of points involved is too small for the individual error quantities to have much meaning, their average value for any given series of related solvents is probably a fair indication of the precision. The values for aqueous solutions seem to be about f O . l kcal./mole and f 0 . 3 cal./moledeg. For a discussion of the shape of the error contour in the AH*, AS*plane, seereference I C .
-
*
Acknowledgments.-This investigation was supported in part by the Office of Ordnance Research, U. S. Army and by a Tennessee Eastman Company fellowship held by B. M. G. (10) E. E. Turner, ibid., 2348 (1928).
