May, 1959
RACEMIZATION OF A BIPHENYL HAVINGA CATIONIC BARRIERGROUP
687
MEDIUM EFFECTS IN THE RACEMIZATION OF A BIPHENYL HAVING A CATIONIC BARRIER GROUP BY JOHN E. LEFFLERAND W. H. GRAHAM' Contribution from the Department of Chemistry, Florida State University, Tallahassee, Fla. Received September 10, 1968
The racemization of a biphenyl is subject to effects due to specific interactions of the biphenyl with adjacent solvent molecules. Changes in the solvent change both the activation enthalpy and the activation entropy, but the relation between these is not sim le unless the solvents belong to a closely related series. Rate ratios up to a factor of five and ranges in activation entha7py and entropy of about 3 kcal./mole and 9 cal./mole degree are reported. The solvent effects are highly specific and do not correlate with those observed for other reactions or with any of the macroscopic physical properties of the solvent.
The rate of a reaction in solution depends on two types of solute-solvent interaction. One of these, the electrostatic interaction of the solute with the non-adjacent or bulk part of the solvent, is well understood and can be described in terms of the electrical properties of the solute molecule and the macroscopic electrical properties of the solvent. However, the intimate interaction of solute molecules with adjacent solvent molecules is less well understood and is undoubtedly highly specific, correlating poorly or not a t all with solvent parameters such as the macroscopic dielectric constant of the medium. The racemization of optically active biphenyls, which is relatively free from mechanistic uncertainties, is almost ai1 ideal reaction for the study of such specific medium effects. Because the reaction involves little or no redistribution of charge, the ground and transition states will be virtually identical from the standpoint of remotely situated solvent molecules. The observed medium effects will therefore represent differences in the specific interactions of the ground and transition states with their adjacent solvent molecules. For example, the rate would not be correlated with the solubility, because the non-specific or long range solvent effects important in the latter (especially for isns) will cancel out in the transition state. The transition state for the racemization of an optically active biphenyl such as I CH3-N-CH3 I
I
CH3-k +-CH3
I
CH,
I is a set of energy levels which the two enantiomorphs have in common. I n the energy levels of the ground state, the barrier groups (occupying the 2,2'-positions) have a considerable freedom of rotation relative to their benzene rings; on the other hand, the benzene rings themselves are restricted in their relative motions to a range of orientations in which the planes of the ring are nearly perpendicular. I n the energy levels of the transition state, the relative torsion of the two benzene rings increases in amplitude to include and pass the (1) Based on the doctoral dissertation of W. H. Graham. Presented in part a t the Symposium on "Solvent Effects and Reaction Mechanisms," Queen Mary College, London, July, 1957.
coplanar orientation, but this increase in rotational freedom is achieved only a t the expense of an even greater decrease in the rotational freedom of the barrier groups relative to their benzene rings. The net effect is that the energy levels of the transition state are less densely spaced than those of the ground state, and the entropy of activation is negative. The main reason for the considerable activation energy in a biphenyl racemization is the steric repulsion of the groups during their passage through the eclipsed positions. This repulsion energy has been shown to be of the right order of mag- . nitude by the calculations of Westheimer and Mayer.2 Because of the 25 kcal. or so of vibrational activation energy the transition state molecules should be visualized as undergoing a motion which is much more violent as well as qualitatively different from that of the ground state molecules. The vibrationally excited uature of the transition state tends to make its solvation weaker and less intimate than that of the ground state. I n the present example, it is possible to explain all of the observed solvent effects a t least qualitatively in terms of the desolvation of the ground state during the activation process, ignoring the solvatioii of the transition state.3 The enthalpy and entropy increments for the desolvation of the ground state will, of course, depend not only on the forces and constraints experienced by molecules in the solvation shells but also on the forces and constraints in the macroscopic solvent, hence on the quasi-crystalline nature of the solvent. I n this paper we report the rates and activation parameters for the racemization of the biphenyl ion I a t 79.4 and 100.Oo in some seventy-seven different media. The biphenyl is one first prepared by Shaw and Turner4 and investigated by Cook and Turner16 who first showed that the racemization rates do indeed depend upon the solvent. Discussion of Results The racemization is first order or nearlys first (2) F. H. Westheimer and J. E. Mayer, J. Chem. Phys., 14, 733 (1946). (3) This simplification is not always possible, perhaps because a large part of the excess energy of the transition state in some biphenyl racemizations is potential rather than kinetic. When the transition state is relatively static and coplanar, its solvation, especially by scomplexing solvent components, cannot be neglected: J. E. Leffler and B. M. Graybill, unpublished work. (4) F. R. Shaw and E. E. Turner, J . Chem. Soc., 135 (1933). (5) D. E. Cook and E. E. Turner, ibid., 88 (1937). ( G ) I n certain solvents, notably the alcohols and aqueous-organic mixtures, there is a slight downward drift in the first-order rate con-
JOHN E.LEE'FLERAND W. H. GRAHAM
688
Vol. 63
activation enthalpy for the single-component solvents is about 3 kcal./mole and the range in entropy about 9 cal./mole degree. If attention is restricted to the alcohols (glycerol, ethylene glycol, benzyl alcohol, methanol, isopropyl alcohol and ethanol), it is found that a graph of the 26.0 enthalpy of activation against the entropy (Fig. 1) consists of a single isokinetics line. If a single isokinetic line represents a qualitatively uniform solvation mechanism, we may conclude that water and /o+C HeOH trifluoroethanol act differently from the alcohols, ,oCTHYLENE / GLYCOL since the corresponding points are not on the line.g 24.0 I n contrast t o the simple relationship between the activation parameters for the alcohol series, the / 0 QLYCERIN enthalpy-entropy relationship for the entire set of solvents, representing, a much wider variety of chemical structures, is a scatter diagram. This lack of a smooth relationship indicates the highly spe- 14 - 10 cific nature of intimate solvation and its dependAS$. ence on the chemical structure of the solvent. Fig. ~ . - - ( C H ~ ) ~ N C E H ~ +( C CHa)s ~ H ~ Ncsmphorsulfonate in A further indication of the specificity of intimate hydroxylic solvents. solvation in this reaction is the complete lack of correlation (for any wide selection of solvents) of either order in all the media studied. The rate constant the rates or the activation parameters with any in water depends on the initial concentration of the physical property of the solvent or even with the run, being higher a t higher concentrations (Table rate of any other reaction in the same solvent. I). The concentration dependence appears to van- Solvent parameters which have been compared with ish in solutions more dilute than 0.02 N . We have our data include the dielectric constant, the viscoschosen 0.02 N as a standard concentration for all of ity, Yloand 2." our runs unless otherwise specified. The anion However, in spite of the high specificity of the may be assumed to be d-camphorsulfonate unless solvation effects, the racemization rate of d-0-(2otherwise specified. dimethylaminopheny1)-phenyltrimethylammonium benzenesulfonate is nearly the same in D- and L-2TABLE I methyl-1-butanols even though the solvated species THECONCENTRATION EFFECT IN WATERAT 100.Oo are diastereoisomeric. The near equivalence of Rate Rate these rates was determined indirectly from an exInitial constant, Initial constant, ooncn. sec.-I X concn., sec.-1 X periment which compares the rates of the d- and 1moles/l. 105 moles/i. 106 biphenyls in D-2-methyl-1-butanol : a fresh solu0.008 4.92 0.033 5.18 tion of the dl-biphenyl is only slightly more levo.009 5.06 .036 5.17 rotatory than a solution which has been heated to .018 4.94 ,070 5.40 ensure equilibrium between the diastereoisomeric .020 4.93 .lo6 5.95 solvates. It also was found that the increase in .026 4.98 rotation caused by heating a solution of the 1-salt in Single-component Solvents.-In Table I1 are this solvent is only slightly greater than the decrease recorded the activation parameters for the 80-100" in rotation caused by heating a solution of the d-salt. temperature range for various pure solvents. It is Of course the D- and L-2-methyl-1-butanols must found that the range of reaction rates in single be very much alike from the point of view of a solute component solvents is about a factor of four to four interacting mainly with the hydroxyl group, and and a half depending on the temperature, trifluoro- the enantiomorphs of a more highly asymmetrical ethanol being the slowest and dimethyl sulfoxide solvent would probably show a greater difference. the fastest.' The slow reaction in trifluoro- It has been shown, for example, that the equilibrium of dl-8-nitro-N-benzenesulfonyl-N-(2-hyethanol can be attributed to hydrogen bonding. The activation parameters show a considerably droxyethy1)-1-naphthylamine is shifted in ethyl greater spread than might have been expected from (+)-tartrate as solvent.12 Two-component Solvents.-Table 111 contains the range in rate constants, the usual cancelling effect being operative to some extent. The range in the activation parameters for the racemization in
I 1
z
,
stant as the reaction progresses. This drift, which amounts to about 4% in the rate constant, is absent in water itself and in non-hydroxylic organic solvents. In order to minimize the effect of the drift o n the comparison of the rate constants, these have been compared at the same extent of reaction, corresponding to a rotation (due to the cation) about 1/2.73 of the initial value. In order to determine whether association with the optically active anion is responsible for the drift, benzenesulfonate was substituted for the d-camphorsulfonate. Since this change made no difference, we are inclined to attribute the drift to a stereospecific interaction between the positive ions. (7) The racemization without solvent in the solid state at 100" is immeasurably slow.
(8) J. E.Leffler, J . Ow. Chem., 20,1202 (1955). (9) It ia not unusual for water to differ from the alcohols in its effect on activation parameters. For example, the enthalpies and entropies of activation for the solvolysis of methyl tosylate in a series of alcohols form an isokinetic line, but the line does not pass near the point for water: J. B. Hyne and R. E. Robertson, Can. J . Chem., 84, 863 (1956). (10) E. Grunwald and 8 . Winstein, J . A m . Chem. Soc., 70, 846 (1948); A. H. Fainberg and S. Winstein, ibid., 78, 2770 (1956). (11) E. M. Kosower, ibid., 80, 3253 (1958). (12) J. Glazer, M. M. Harris and E. E. Turner, J . Chem. SOC., 1753 (1950).
RACEMIZATION OF A BIPHENYL HAVING A CATIONIC BARRIER GROUP
May, 1959
689
TABLE I1 PURESOLVENTS sec.-l X 10-
-k,
Solvent
AH $ kcal./rndlec
loO.oQ=
79.40'
AS$, cal./mole deg.c
Trifluoroethanol 0.41 3.46 26.31 f O . l l - 8.87f0.31 Deuterium oxide .56 4.74 2 6 . 3 9 f .06 8 . 0 4 f .18 .57 4.93 26.71 f .06 - 7 . 1 1 f .15 Water 1.03 7.87 2 5 . 1 7 f .04 - 1 0 . 3 0 f .lo Methanol 1.08 8.56 2 5 . 4 7 f .01 - 9 . 3 4 f .02 Acetic acid 1.10 7.18 2 3 . 0 8 f .26 - 1 6 . 0 8 f .73 Glycerol 1.18 9.14 2 5 . 2 4 4 ~ .16 - 9.831. .44 Acetonitrile 1.18 9.50 2 5 . 7 0 f .08 8 . 5 1 f .22 Ethanol 24.19 f .02 - 1 2 . 7 7 f .06 Ethylene glycol 1.20 8.51 - 9 . 3 1 f .27 Nitrobenzene 1.29 10.05 2 5 . 3 6 f .10 2 5 . 2 2 f .07 Isopropyl alcohol 1.42 10.90 - 9 . 5 2 f .20 Benzyl alcohol 1.52 11.32 2 4 . 7 9 f .06 -10.61 I . I 6 8.03 f .47 Chloroformb 1.53 12.24 25.69 f .17 Dimethylformamide 1.62 12.49 2 5 . 1 9 f .14 - 9 . 3 2 f .40 Dimethyl sulfoxide 1.82 13.39 2 4 . 6 6 f .06 - 1 0 . 6 2 1 .16 a These are mean rate constants. The individual values have been adjusted in those cases where the temperature was actually a fraction of a degree different from 79.4 or 100.Oo. It was necessary to use chloroform stabilized by a traceof The significance of the numbers indicating the precision is discussed in the experimental section. ethanol.
-
-
-
TABLE I11
Acetone-water
TWO-COMPONENT SOLVENTS Solvent, wt. % of first component
k, 8ec-1 100.Oo' X 106 79.4"
AS',b cal./mole deg.
AH* b kcsl./dole
20.9 0.95 44.2 1.19 66.5 1.37 95.2 1.54 (Solubility limit)
Methanol-water 0 3.8 16.5 20.8 44.2 70.2 94.0 100.0
0.57 .62 .72 .74 .90 1.02 1.02 1.03
4.93 5.33 6.12 6.19 7.19 7.80 7.87 7.87
-7.11f0.15 -7.39f.24 -7.35f.30 -8.10f.06 -9.27f.21 -10.34f.15 -10.19 Et .l5 -10.3OEt .10
3.8 20.8 44.1 48.7 63.5 75.9 96.0 98.0 100.0
0.64 0.93 1.10 1.16 1.22 1.25 1.23 1.19 1.18
20.7 43.9 60.9 79.7
1.02 1.26 1.37 1.41
26.24rt0.08 25.50rt.15 24.99f.06 24.91 f .07 25.07* .06 24.98f .l2 25.18f .04 25.87f.05 25.70f.08
-8,20+0.23 -9.56f.42 -10.67zt.17 -10.8Of .20 .18 -10.26f -10.45f .33 9.89 f .11 -8,01&.13 -8.51f.22
7.94 9.41 10.10 10.49
25.31f0.06 24.87 f .05 24.64f .05 24.81 f .02
-9.93*0.18 -10.76 f .14 -11.24f .13 -10.7Of .05
100.0
1.42
10.90
25.22f
- 9 . 5 2 5 .20
-
Isopropyl alcohol-water
.07
&Butyl alcohol-water 7.2 0.75 23.9 1.16 43.9 1.33 6G. 1 1.43 91.4 1.60 (Solubility limit)
6.21 8.75 9.71 10.43 11.82
26.05f0.05 2 4 , 9 6 f .06 24.52f .02 2 4 . 4 7 3 ~ .06 2 4 . 7 0 r t .13
-
8.41 rtO.13 -10.67f .18 -11.64f .05 -11.63 f .17 - 1 0 . 7 7 h .35
Ethylene glycol-water 18.0 46.7 84.5 100.0
0.69 0.86 1.09 1.20
5.2 18.7 35.6 73.4 100.0
0.63 .67 .63 .52
5.74 6.83 8.26 8.50
26.17f0.06 25.60f .05 25.02f .OS 24.19f .02
- 8.24fQ.15
-
9 . 4 4 f .13 -10.61 f .22 -12.77f .06
a-Trifluoroet hanol-water
.41
5.30 5.63 5.31 4.52 3.46
*
25.26f0.04 24.49f .09 2 3 , & 5 f .12 24.73f .12
-10.20rt0.12 -11.94f .25 -13.47f .32 -10.753Z .32
Dioxane-water
26.71f0.06 26.55f.09 26.46f.11 26.17f.02 25.62f.08 24.16f.06 25.21.f .05 25.17f .04
E thanol-wat,er 5.33 7.33 8.38 8.92 9.28 9.46 9.52 9.69 9.50
7.38 8.65 9.45 11.46
26.23 0.13 26.15 f .07 26.23 f .13 26.63 f .09 26.31 f . I 1
-
-
8 . 2 5 Et 0.35 8.34 f .20 8 . 2 5 f .35 7.52 f .25 8.87 t . 3 l
29.3 1.07 .40.9 1.27 50.9 1.37 72.2 1.63 83.8 1.87 92.6 2.11 (Solubilitylimit)
8.08 9.34 9.57 11.41 13.10 14.72
24.94Et0.06 2 4 . 6 2 3 ~ .09 2 3 . 9 2 r t .09 24.00f .05 2 3 . 9 8 5 .03 23.93f .07
Dimethyl sulfoxide-water 15.5 52.4 76.7 86.8 100.0
0.76 1.14 1.48 1.64 1.82
6.22 8.90 10.93 12.16 13.39
26.03 rt 0.05 25.31 f .02 24.69f .05 24.71 rt .02 2 4 . 6 6 r t .06
Acetic acid-water 4.0 0.40 17.2 .42 34.3 .54 51.2 .62 72.4 .78 100.0 1.08 (Solubilitylimit)
4.00 4.29 5.20 5.50 6.70 8.56
28.41 f 0.29 29.11 Et .41 28.01 Et .12 27.03* .06 26.54 f . I 4 2 5 . 4 7 f .01
-10.86rt0.16 -11.43 f .25 -13.26f .24 -12.70f .13 -12.48f .10 -12.393Z .18
-
8.46 f 0.15 .06 -10.93f .15 - 1 0 . 6 7 r t .05 - 1 0 . 6 2 3 ~ .20
- 9.69Et
- 3.00 f 0.79 - 0.93f1.14 - 3.51 Et0.35 - 6 . 0 3 f .16 - 6.96Et .38 - 9 . 3 4 f .02
Acetic acid-benzene 1.92 1.68 1.55 1.48 1.31 78.2 'O0.O 1.08 (Solubilitylimit) 2.2 19.3 37.4 54.4
14.13 12.33 11.50 11.00 10.00 8.56
24.55f0.09 2 4 . 5 7 r t .06 24.70f .07 24.70f .09 25,OOf .09 2 5 . 4 7 r t .01
-10.79f0.24 -11.02rt .18 -10.79f .18 - 1 0 . 8 9 f .23 -10.3Of .25 9 . 3 4 k .02
-
Methanol-benzene 8.3 47.4 64.3 73.0 81.8 91.8 95.7 100.0
1.62 1.27 1.19 1.15 1.07 1.07 1.03 1.03
12.80 9.75 9.06 8.85 8.67 8.47 8.00 7.87
25.52 3Z 0.06 2 5 . 2 0 f .06 25.04f .15 2 5 . 2 3 f .04 2 5 . 8 3 5 .OS 2 5 . 5 0 f .03 2 5 . 2 6 r t .03 2 5 . 1 7 3 ~ .04
- 8.41 f 0 . 1 7 9.79f .I7 - 1 0 . 3 7 2 ~ .42 9.89f .11 8.35Et .22 9 . 2 8 * .os -10,03f ,OS -10.3Of .10
-
-
These are mean rate constants. The individual values have been adjusted in those cases where the temperature was actually a fraction of a degree different from 79.4 or 100.0". The significance of the numbers indicating the precision is discussed in the Experimental section.
JOHN E. LEFFLERAND W. H. GRAHAM
(390 I
24.01
I
I - 12
I -8
- 10
As$. Fig. 2.-( CH3)2NC6H4C~H4N +( CH3)a camphorsulfonate in aqueous ethanol (wt. yo).
-
- 12
10 -8 AS$. Fig. ~ . - ( C H ~ ) Z N C ~ H ~ C+(~CH& H ~ N camphorsulfonate in aqueous t-butyl alcohol (wt. 70).
-9
-11
-7
AS:.
Fig. 4.-( C H ~ ) Z N C ~ H & ~+(HCHa)3 ~ N camphorsulfonate in aqueous ethylene glycol (wt. yo).
Vol. 63
a few of the sixty-one two-component solvent mixtures. The behavior of alcohol-water mixtures is about what might have been expected in view of the fact that the point for water does not fall on the isokinetic line for the series of pure alcohols. The activation parameter for the mixtures of water with various alcohols shows a linear enthalpy versus entropy relationship from zero up to about fifty or more weight per cent. of the alcohol. Beyond this point the line breaks sharply, in most cases doubling back nearly upon itself. Figures 2 and 3 for aqueous ethanol and aqueous t-butyl alcohol are given as examples.la I n the series of monohydric alcohols the values of the enthalpy and entropy a t which the line breaks are in the order methanol > ethanol > isopropyl alcohol > t-butyl alcohol, Le., in the order of increasing molecular volume rather than the order in which the points fall on the isokinetic line for the pure alcohols. The somewhat different behavior of aqueous ethylene glycol is shown in Fig. 4. Although the enthalpy-entropy diagram for the monohydric alcohol-water mixture shows sharp breaks, we prefer to interpret the results in terms of a gradual change in the solvation mechanism rather than in terms of an abrupt change. One reason for this preference is that the graphs of enthalpy and entropy against weight per cent. of the alcohol are smooth, the minimum found in these curves in most cases being very broad. The relationship is shown in Fig. 5 for the isopropyl alcohol-water system. The minima in the activation parameters can be rationalized by simply assuming that the activation process involves considerable desolvation and release of solvent molecules to the bulk solvent structure. For methanol and ethanol it is known that the excess enthalpies and entropies of mixing with water pass through a minimum.14 If this is general for the monohydric alcohols of our series it affords a sufficient explanation for the minima in the activation enthalpy and entropy. The enthalpy-entropy relationship for the 6trifluoroethanol-water mixture is much more complicated than that for the unsubstituted alcohols. It has both a maximum and a minimum in the enthalpy and entropy, and a single maximum in the rate. On the other hand, the enthalpy-entropy relationship for the acetone-water mixtures is very much like those of the alcohol-water mixtures except that the single minimum occurs at a considerably lower value of the enthalpy and a more negative value of the entropy. The activation parameters for the dioxane-water and dimethyl sulfoxide-water mixtures resemble each other. I n both cases the enthalpy of activation decreases as a linear function of the entropy up to about 70% (by weight) of the organic component. Above this composition, however, there is a sharp break. The enthalpy of activation remains almost constant, but the entropy increases, causing the rate (13) For the even more complicated pattern of the solvolysis reaction (9.1-type) in mixed solvents the reader should refer to the excellent paper by S. Winstein and A. H. Fainberg, J. Am. Chem. Soc., 1 9 , 5937 (1957). (14) A. G.Mitchell and W. F. K. Wyone-Jones, Dzsc. Faraday Soc., 18, 161 (1953).
May, 1959
RACEMIZATION OF A BIPHENYL HAVING A CATIONIC BARRIER GROUP
to continue to rise. The rates in 92.6% dioxanewater mixtures were the fastest observed for any medium, being 5.15 times as fast as for trifluoroethanol at 79.4' and 4.26 times as fast a t 100.0'. I n aqueous acetic acid the dimethylamino groups are protonated to a considerable extent, as can be seen from the low value of the specific rotation. I n aqueous hydrochloric acid the specific rotation is also low and the rate of racemization is too slow to measure. The latter result presumably means that the repulsion of the two positive charges, which should accelerate the racemization, is of lesser importance than the idcreased solvation of the protonated dimethylamino group and the change in hybridiza tion. Although the low specific rotation reduces the accuracy of the results for aqueous acetic acid mixtures, there is no doubt about the direction of the effects. The addition of very small amounts of acetic acid to water drastically reduces the rate, but slightly larger additions of acetic acid increase the rate. The enthalpy-entropy relationship (Fig. 6) is V-shaped, but with a maximum rather than the minimum encountered in the aqueous alcohol mixtures. Since our experiment with hydrochloric acid shows that the protonated form does not racemize at all, the activation parameters in aqueous acetic acid include contributions from the deprotonation as well as the desolvation of the biphenyl. The addition of benzene to acetic acid results in a steady increase in rate. The enthalpy and entropy decrease as benzene is added and maintain their linear relationship with each other until a proportion of about soy0 benzene is reached. At about that composition, there is a sharp break in the linear relationship and there appear to be a t least two minima and one maximum in the curve before the solubility limit is reached near 98y0 benzene. The addition of benzene to methanol also causes a steady increase in rate and the plot of enthalpy against entropy of activation is N-shaped, with breaks a t 50 and SOY0 methanol. The initial increase in enthalpy and entropy of activation on adding benzene to methanol cannot be due to increased solvation of the biphenyl, since the rate actually increases. It could, on the other hand, be attributed to a change in the nature of the solvent structure to which the molecules of the solvation shell are returned during the desolvation which accompanies the activation process. Relative to pure methanol, the desolvation in a solvent containing a little benzene is more endothermic and also involves a greater decrease in the constraint of the solvent molecules. This seems to be in the wrong direction if the effect is due merely t o the presence of benzene instead of methanol in the solvation shell, but could be due to a change in the structure of the bulk solvent. The required change is that solvent molecules are less tightly associated with each other in the solvent mixture than they are in pure methanol. Hence the differences in enthalpy and entropy between solvent and solvation shell are greater for the mixture than for the pure alcohol. Miller and Fuoss have suggested that benzene added to methanol breaks down the
691
12
26
11
25
9 8
20 40 60 80 100 Wt. % ' isopropyl alcohol. +( CH3)3 camphorsulfonate in Fig. 5.-( CHS)ZNC~"C~,H~N aqueous isopropyl alcohol. The filled circles represent A H $ , the open circles represent AS$. 0
28.0
26.0 '
/
0H O A ~ -8
-4 AS$* Fig. 6.-( CH&NC6H4C6H4N+( CH3)3 camphorsulfonate in aqueous acetic acid (wt. yo).
polymeric structure of the solvent and increases the amount of monomeric methanol. l6 Experimental qnd Mathematical Procedures 2,2 '-Bisdimethylaminobiphenyl.-2,2 '-Dinitrobiphenyl was made by means of a modified Ullmann reaction16 using pretreated copper bronse.17 The reduction to the diamine with iron filings proved to be erratic and most of the reduction was carried out by a catalytic procedure similar to that of Blood and Noller.l* The methylation was done by the procedure of Shaw and Turner . 4 dl-o-( 2-Dimethylaminophenyl )-phenyltrimethylammonium Iodide.-The following procedure is easier than the sealed tube method of Shaw and Turner. One equivalent of 2,2'bisdimethylaminobiphenyl is refluxed for several hours in anhydrous acetone containing several equivalents of methyl iodide. The white solid product is filtered off, more methyl iodide is added, and the refluxing is continued, producing a second crop of solid product The product needs no recrystallization, melts at 184-185", and is obtained in quantitative yield. It is essential that the acetone be dry or the product will be contaminated with its hydriodide salt. d-o-( 2-Dimethyl aminopheny1)-phenyltrimethylammonium d-Camphorsu1fonate.-This substance was pr:pared by the method of Shaw and Turner,4 m.p. 184-185 , [ C Y ] % 48.6 in water a t 1.32 g./100 cc. (15) R . C. Miller and R. M. Fuoss, J . A m . Chem. Soc., 7 6 , 3076 (1953). (16) N. Kornblom and D. L. Kendall, ibid., 1 4 , 5782 (1952). (17) E. C. Kleiderer and R. Adams, ibid., 66, 4225 (1933). (18) A. E.Blood and C. R. Noller, J . O r g . Chem., 22, 711 (1957).
JOHNE. LEFFLERAND W. H. GRAHAM
692
Vol. 63
heating at the usual racemization temperature are identical, indicating that the reverse reaction does not take place. Demethylation by wat,er would produce hydrogen ion; none is liberated, as shown by titration to a phenolphthalein endpoint, when the d-camphorsulfonate is heated in water for several hours. Disproportionation to a doubly quaternary salt and 2,2’-bisdimethylaminobiphenylcan be ruled out as a side mechanism in view of the fact that a large excess of a much more nucleophilic substance, dimethylaniline, produces only a 10% increase in racemization rate in ethanol and a 4% increase in chloroform, probably medium effects. An intramolecular methylation of the dimethylamino group by the trimethylammonium group would not lead to racemization in any case. Kinetics.-Samples in nitrogen-filled ampules were placed in thermostats of conventional design with a temperature control of f0.02”. Absolute values of the temperature, known to within 0.1’, were obtained from N.B.S. calibrated thermometers. The ampules were quenched by immersion in ice-water, opened, and the contents placed in the polarimeter tube. The Bellingham and Stanley polarimeter, equipped with a sodium vapor light source, was read :i a darkened room whose temperature was controlled to f 2 For most points a total of sixteen readings, eight 6AS+ each with the tube turned in opposite directions, were taken. Fig. 7. Generally four to six points, in addition to an initial and dl-o-( 2-Dimethylaminophenyl )-pheny ltri methylammonium final point, were taken for each run. These points were Benzenesu1fonate.-The racemic benzenesulfonate is pre- clustered in that part of the curve which gives maximum acpared by warming a solution of silver benzenesulfonate in curacy for the rate constant. A standard concentration of 0.02 N was used in most of the acetone-ethanol with an equivalent amount of the dl-o-(2dimethylaminopheny1)-phenyltrimethylammonium iodide. experiments. This concentration was chosen for maximum Silver iodide is removed by centrifugation and the solution conservation of material consistent with adequate rotation. The concentration of the d-o-(2-dimethylaminophenyl)is diluted with heptane which causes the racemic salt to phenyltrimethylammonium d-camphorsulfonate had little crystallize. d-o-( 2-Dimethylaminopheny1)-phenyltrimethylammonium effect on its specific rotation Iodide.-This compound, prepared by the method pre[ a l p 3 D in HpO, viously reported,4 was recrystallized from chloroformNormality degree hexane: m.p. 181-182’ dec. 0.01025 48.8 d-o-( 2-Dimethylaminopheny1)-phenyltrimethylammonium .02726 48.6 benzenesulfonate6 was crystallized from acetone-hexane, m.p. 124-125’, [ c Y ] ~ ~43.5 D in HzO a t 0.84 g./100 cc. ,05126 48.6 Lo-( 2-Dimethylaminophenyl)-phenyltrimethylammonium Benzenesu1fonate.-The l-iodide is obtained by adding It was also observed that variations in room temperature sodium iodide to the mother liquor from the crystallization had no detectable effect on the specific rotations in the orof the d-o-(2-dimethylamino heny1)-trimethylammonium dinary solvents. However, in experiments using D-2d-camphorsulfonate, and cry stahzing the product by adding methyl-1-butanol as solvent, it was necessary to control the hexane; [ a I z 6-40.0, ~ m.p. 124-125’. temperature to f 0 . 2 ’ because the specific rotation of the Solvents.-Unless otherwise noted, all solvents were solvent changed appreciably with temperature, about 0.01 carefully dried and purified by distillation and/or fractional per degree. crystallization. Dimethyl sulfoxide (99.9yo) was used as Asymmetrical Environment Experiments.-The optical supplied by the Stephan Chemical Go., Chicago. “Spectro rotations of solutions of d and 1 and dl-o-(2-dimethylaminograde” chloroform, containing a trace of ethanol, was used phenyl)-phenyltrimethylammoniumbenzenesulfonate in Dwithout further purification since it was found that pure 2-methyl-1-butanol were measured before and after heating chloroform tended to produce phosgene during the runs. in sealed ampules. The concentration of the dl-salt was Absolute cthanol and methanol were prepared by the methods described by Fieser.19 Reagent grade benzene and Angle, Angle, before after analytical reagent grade glycerol were used without purifiheating Conditions heating Aa cation. Dioxane was purified by refluxing over sodium 16 hr. at 100’ 167.151’ 0.884 d-salt 168.035’ hydroxide for several days and then refluxing and distilling over sodium. “Spectro grade” acetone was used directly. 168.268’ 11 days a t 80’ 167.170’ .912 2-salt Since all acetone solutions used were aqueous mixtures, 16 hr. a t 100’ 167.370’ ,029 dl-salt 167.341’ traces of water were unimportant. D-2-Methyl-1-butanol, [ c Y ] ~ ~ . -5.79, ~D was supplied by the Dow Chemical Com- different from that of the others, but the concentrations of pany * d- and 1-salts were nearly equal. The error expected in Investigation of Side Reactions.-Equimolecular amounts the the ACYdue to the fluctuation of the rotation of the solvent of methyl iodide and 2,2’-bisdimethylaminobiphenylheated with the temperature a t which it is read, is about 0.004. in chloroform for eight hours gave the quaternary iodide in Racemization in the Solid State.-A sample of d-o-(2only 2.18% yield as determined by Volhard titration. When dimethylaminopheny1)-phenyltrimethglammonium d-camd l - o-(2-dimethylaminophenyl)-phenyltriniethylammonium phorsulfonate was sealed in an ampule under nitrogen iodide was heated in chloroform, however, the back reaction left in a steam-bath for 55 days. The spyific rotationand in to 2,2’-bisdimethylaminobiphenyl was almost complete aqueous solution before heating was 47.9 , after heating, after a fcw hours. The racemization of the quaternary 47.8”. ammonium biphenyl can therefore not be measured in the The Calculation of Rate Constants. Activation Parampresence of iodide ion. The corresponding reactions in- eters and their Errors.-The rate constant is given by the volving methyl camphor sulfonate do not take place. The equation infrared spectra of o-( 2-dimethglamino heny1)-phenyl trimethylammonium d-camphorsulfonate gefore and after mutarotation in chloroform are idcnt,ical. Furthermore, the spectra of a chloroform solution of methyl d-camphorsulfonate and bis-dimethylaminobiphenyl before and after where 1 is the elapsed time, eothe initial rotation of the solution, and CY the rotation a t time t . The rotations are corrected for that due to the camphorsulfonate. The method (19) L. F. Fieser, “Experiments in Organic Chemistry,” 3rd ed., D. C. Heath and Co., Boston, Mass., p. 281. of least squares was used to calculate the best values of the
1
AH*
-
.
O
May, 1959
,
KINETICS OF THE. STEAM-CARBON
activation parameters for the experimenlal temperature range, fitting the equa.tion k a AS+ AH* 1 log~=logjE+2.303R 2.303R T The calculations were performed on a desk calculator a d checked with an I.B.M. model 650 digital computer, using a program for which we are indebted to Prof. H. C. Griffith of the F.S.U. mathematics department. The expected maximum errors in A H $ and ASS are about 10.14 kcal./mole in the enthalpy of activation and about f 0 . 4 0 cal./mole degree in the entropy of activation, based on a conservative estimate of f0.002° as the maximum error in the average of the sixteen readings of CY for each point. The calculation of the least-squares activation parameters gives us a measure of their precision as a by-product. These numbers indicated as precision measures in the tables have the form of probable errore, 0.6745 times the standard deviations. I t should be remarked that these numbers are not, in fact, probable errors, because the statistical sample of points for each solvent is too small, some of the reported error quantities no doubt being fortuitously small. On the other hand, the mean of the error figures for a large number of related solvents very likely is a measure of the probable error of the activation parameter for that series. In the graphs of activation enthalpy versus entropy, the experimental points are represented by small circles. These circles are not probable error contours. The actual probable error contour is a narrow ellipse20whose major axis
693
REACTION
has a slope numerically equal to the mean experimental temperature in degrees Kelvin. The reason for the elliptical shape of the probable error contour is that the precision with which log k , or the free en rgy of activation, is known is such that if A H $ has the InTgest value permitted by its probable error, then A S S must also have nearly its largest permitted value. That is to say, the largest value for the enthalpy within its probable er or and the lowest value for the entropy within its probabL error, are a combination which would correspond to an kxtremely improbable value for the rate constant. The effect of this restriction the probable error contour is shown in Fig. 7. The practical consequence of the sloping elliptical .error contour is that points having elative positions like those for 44.1 and 75.9% alcohol iAFig. 2 are more likely to differ significantly than points having relative positions like those for 44.1 and 48.7% alcoho
01
. Iinvestigation was aided
Acknowledgments.-This by the Office of Ordnance Research, U. S. Army, through a research contrakt and by a National Science Predoctoral Fellowship (held by W. H. G., 1957-1958). We also wish to thank Professors E. Grunwald and F. H. vestheimer for helpful discussions. (20) J. Mandel and F. J. Linnig, &aZ.
Chem., 29, 743 (1957).
ICINETICS OF THE STEAM-CARBON REACTION BY GEORGEBLYHOLDER' AND HENRY EYRING Department of Chemistry, University of Utah, Xalt Lake City, Utah Received September 1 1 , 1968
I
The dat,a for the steam-graphite reaction from 900 to 1300' and in the 5 to 100 p pressure range are presented as a function of both pressure and temperature. The rate-determining processes are the adsorption of ater vapor and the desorption of molecular hydrogen from the surface. Absolute rate theory calculations lead to the c$clusion that the adsorbed species have a limited mobility upon the surface.
natural graphites which cbntain impurities and carbonized filaments which are not 100% graphitized would react differently/from pure graphite. With these conditions in mind an apparatus was designed in this Laboratory to study the reactions of graphite with oxidizing gases between 900 and C h ) H2O(,) +CO(,, H2k) 1300' and in the 5 to 100 p pressure range. The apBeyond the agreement that CO and Hz are the pri- paratus is a flow system in which water is admitted mary products of the reaction a t one atmosphere or through a capillary before tye furnace and the flow less of HzOpressure, there is little agreement on the is maintained by a mercury diffusion pump after activation energy, order of the reaction and the the furnace. The furnace is such that a four inch magnitude of rate constants. A number of experi- long zone is maintained at1 12'. The incoming mental conditions which have not been met by pre- vapor is heated on alumina chips in the central reviously reported data are deemed necessary if the gion of the furnace. The g aphite samples, which inch.' iameter spectroscopic resultant data are to be suitable for kinetic inter- were cut from a pretation. The amount of HzO vapor decomposed electrode (National Carbon Company), are susshould be small to avoid the retarding effect of Hz pended by means of a platin m wire in the constant on the reaction. This condition also makes inter- temperature zone of the fur ace. The rate of the pretation of the order of the reaction unambiguous. reaction was followed by me suring the rate of presThe HzO vapor and the carbon surface should be at sure build up of CO and 2 after the HzO was the same temperature. If this is not done the tem- trapped out. Results with t is apparatus have apperature of the activated complex in the reaction is peared in a n earlier publication. Unfortunately, not clearly defined. The nature of the carbon it has recently been discovered that, due to the high should be adequately defined. I n this study we flow rate of vapor (over 1000 cm./sec.) through the are interested in the reaction of graphite with water furnace and an unfortunate choice of location of the vapor so the carbon samples are from a spectro- thermocouple gage used to determine pressure, the scopic graphite electrode. It is to be expected that pressures of water vapor in the furnace at the samIntroduction There have appeared a large number of papers on the steam-carbon reaction. It is generally agreed that the reaction may be represented by the equation
+
+
1 1
i
(1) Presently a Research Associate at the Johns Hopkins University. Baltimore 18, Maryland.
(2) J. 8. Binford, Jr., and H. Eyring, THrs JOURNAL^ 60, 486
(195G).
