CATALYTIC EFFECT OF SOLVENTS. DECOMPOSITIOK OF MALONIC ACIDS
BY ROBERT E. BURK AND WILHELMISA DAUS
While numerous instances of the influence of solvents upon the velocities of chemical reactions have been investigated, some of those which have been most carefully worked through’ have been addition reactions, e.g. the formation of substituted ammonium salts by the reaction of amines with alkyl halides, etc. In this type of reaction the possible factors which may contribute to the rate are not a minimum, e.g., there must be a third body present to carry off the heat of reaction2 in order for reaction to occur, and the possibility that the heat of reaction would be dissipated internally in the case of complex molecules is not indicated in Menschutkin’s work, since the more complex the molecules were, the slower the rates. That this factor is not the whole rate controlling one in these instances, is shown by the fact that Norrish and Smith found a discrepancy between the bimolecular collision rate of active molecules and the observed velocity, of IO-* for the reaction between nitro-benzyl chloride and trimethyl amine in benzene, whereas the factor involving dissipation of heat of reaction could not in all probability lead to any such value, although it would be in the proper direction. Norrish and Smith consider every such collision between active molecules to involve at least one solvent molecule, and if this molecule were efficient in dissipating the energy of reaction, this required dissipation would not have to be considered at all. They attribute the discrepancy in collision and reaction rates to division of activation energy with solvent molecules at the moment of collision between active molecules. Certain workers3 on oxidation reactions. in solution, however, in applying Christiansen’s theory of negative catalysis4 to such reactions, consider it not improbable that energized molecules of products may undergo of the order of IO-^ or IO-^ collisions with solvent molecules without losing their energy, and can transfer it specifically to “cold” reactant molecules at the moment of collision with them, without interference from solvent molecules of the type considered by Norrish and Smith. With addition reactions of the above type, then, the solvent could affect the probability of loss of energy of activation at the moment of collision between active molecules, could lower the energy of activation, could dissipate energy of reaction, etc. Xorrish and Smith were inclined to consider the effect of different solvents in altering the reaction rate as due to an effect Menschutkin: Z. physik. Chem., 5 , 589; 6 , 41 (1890); Norrish and Smith: J. Chem. 130, 129 (1930). 2 Born and Franck: Ann. Physik, (4) 76, 2 2 5 (192j). 3e.g. Backstrom: J. Am. Chem. Soc., 49, 1460 (1927). Christiansen: J. Phys. Chem., 28, 145 (1924).
SOC.,
‘
1462
ROBERT E. BURK AXD WILHELMINA DAGS
upon the first factor. It seemed desirable to study a simpler reaction, such as a decomposition, in order that the factors influencing reaction rates in solution could be investigated more closely. The thermal decomposition of malonic acids was chosen because they fulfill this requirement', because they are important in themselves to organic synthesis, and because t'he influence of substitution of the active hydrogens on the decomposition rate could be compared with the influence of solvents. Work on the decomposition of acetone dicarboxylic acid has appearedj since the present research was start'ed, and the decomposition of malonic acid itself has been investigated kinetically by LindnerJ6 by Hinshelwood,' and together with substituted malonic acids by Bernoulli and Wege,8 by Bernoulli and J a k u b o ~ i c zand , ~ by Jakubowicz.lo Experimental Procedure The apparatus consisted of a thermostat cont'aining a small reaction vessel of approximately 14 C.C.capacity. Temperature control to 10. j"C. was obtained. Above the reaction bulb was sealed a vessel of 41.87 C.C.capacity, which was kept immersed in ice. This bulb was followed by a manometer, then a stopcock. 811 connecting tubing to this point was capillary. The stopcock led to an evacuating system with a discharge tube t,o indicat'e \Then a sufficient vacuum had been produced in the system. The reaction rate was followed manometrically by the pressure changes caused by the evolution of carbon dioxide due to the reaction: RHC(C0OH)s = RH2CCOOH COn Due account was taken of the temperature differences of the two bulbs. When a solvent, which was somewhat volatile at, the temperature of the reaction, was used, it distilled into the cold bulb and refluxed therefrom to the small bulb, Thus only its vapor pressure at 0°C. affected the manometer, and this was always too small to measure. The sample was weighed out accurately before sealing on the reaction bulb, using the bulb as a weighing.bott1e. The amount of acid was so chosen that complete decomposition would not bring the pressure of carbon dioxide in the system above atmospheric. After the sample was weighed, (and the solvent added), the bulb was sealed onto the apparatus, and the system evacuated in order to evolve any dissolved or occluded inert gases. The materials used were Kahlbaum's benzyl malonic acid (m.pt. ~ z o . o O C . ) , ethyl malonic acid from E. Nerck, Darmstadt (m.pt. I I I . j"C.), and solvents the purest obtainable from Eastman Kodak Co. The oleic acid used was from Eastman Kodak Co. and was redistilled twice. The dimethyl aniline was also redistilled, but a redetermination with the distilled product showed no change in rate.
+
M-iig: .I. Phys. Gnem., 32, 961 !1928); 34, j96 (1930). Monatsheft, 28, 1041 (1908). J . Chem. SOC.,117, 156 ( ~ g z o ! . * Helv. Chim. Acta, 2. 511 (1919). Helv. Chim. Acta, 4, 1018 (1921). Z. anorg. allgem. Chem., 121, I 1 3 (1922).
j
6
1463
CATALYTIC EFFECTS O F SOLVENTS
Order of the Reaction The kinetic order of the decomposition of benzyl malonic acid was determined by investigating the influence of initial concentration upon the time of half life using as a solvent palmitic acid, which was found to be without appreciable catalytic influence. Table I shows the independence of the time of half life and initial concentration for benzyl malonic acid.
TABLE I Concentration (gms. in j cc. of acid) I.
0.1046
2.
0.2013
3.
0.2516
Molality
Time for half change (min.)
0.1078
35 31 30
4. 0.3004 The reaction is quite accurately monomolecular.
29
Influence of Various Solvents Table I1 shows the half life for the decomposition of benzyl malonic acid in various solvents.
TABLE I1 Decomposition of Benzyl Malonic Acid Solvent
Sone Oleic Acid Palmitic Acid ortho Sitrochlorobenzene ?2 meta I’ 7, para ” Dimethyl Aniline
Half life (minutes) 12.5”C.
75 72
135°C.
(~8.4)~ 32.1 31.0 25.0 2 2 .j
26.0
7
* This value is low, according to the temperature roefficient graph.
4.3
A remarkable independence of reaction rate and nature of solvent is evident. Determinations made with ortho, meta, and para nitrochlorobenzene indicate that the position of substitution in the benzene ring has no effect on the velocity, although accompanied by large changes in the electric moments of these molecules.ll One experiment was repeated in which about roo mm. of carbon dioxide were added a t the beginning to determine the effect of its presence on the rate of decomposition. It was found that the time of half life in this case was not appreciably different from the others. All of this is in line with other monomolecular homogeneous reactions, which in the few cases investigated have been found t o be less susceptible to catalvtic influence than bimolecular reactions.’l LLPhysik.Z., 30, 391 (1929). This was pointed out by one of us (R.E.B.) a t the Minneapolis meeting of the American Chemical Society, September 1929. Examples are the maintenance of decomposition rate of N,Os (Lueck: J. Am. Chem. Soc., 44, 757 (rgzzjj, the racemization of pinene (Smith: J. Am. Chem. Soc., 49, 43 ( 1 9 2 7 ) ) , and the decomposition of acetone (Taylor: J. Phys. Chem., 33, 1793 (1930)). For exceptions see Clusius and Hinshelwood: Z. Elektrochemie, 36, 748 (1930).
1464
ROBERT E. BURIC AND WILHELMINA DAGS
It is to be observed that solvents having a considerable electric moment such as fatty acids and nitrochlorobenzenes have no catalytic influence. The malonic acids were not soluble in completely non-polar solvents such as purified mineral oil. Dimethyl aniline was the only solvent of those investigated to show marked catalytic properties. Salt formation was indicated by a sharp rise in temperature when malonic acid was stirred with dimethyl aniline. However, reasons are given later for thinking that decomposition of the salt to free base and acid precedes the decomposition of the acid. The catalytic influence of this solvent was accompanied by a pronounced decrease in the temperature coefficient. Effect of Substituting for the Active Hydrogens in Malonic Acid This phase of the work was not investigated thoroughly because of its divergent nature, and of good correlation between such experiments as were ‘‘A summary of results is made and previous rather extensive work. given in Table 111. Other comparisons will be made in discussing temperature coefficients.
TABLE I11 Acid
Phenyl malonic Benzyl malonic Allyl malonic Malonic Malonic
Velocity Constant
x
109
I04,92 I 1,968
Diallyl malonic Methyl malonic Ethyl malonic Methyl ethyl malonic Dimethyl malonic Diethyl malonic Monochlor malonic Dichlor malonic Tartronic
IJ3”
755 477 649 598 j08 322
29;
281 18jj
585 820
Temperature (Centi- Solvent grade)
water water IOO water 100 water 100 glacial acetic acid 100 water 100 water 100 water 100 water IOO water roo nater 100 water roo water IOO water IOO
Chserver
Bernoulli and Kege
,
IOO
J )
,, ,,
,, ,,
7,
Imder Bernoulli and Kege
,,
,,
,,
11
33
1,
,!
)7
,I
9 ,
,,
7,
21
1,
Bernoulli and Jakubowicz
,,
Jakubowicz
half life (min.) hIalonic 174 125 148 125 Ethyl malonic Benzyl malonic 75 12s See Table IT for benzyl malonic acid in
7,
none Hinshelwood none Daus and Burk 1. none various solvents. ,)
1,
,I
CATALYTIC EFFECTS O F SOLVENTS
I465
Temperature Coe5cients Fig. I shows the results of plotting the reciprocal of the absolute temperature versus time of half life for the decomposition of benzyl malonic acid a t various temperatures, without a solvent, in oleic acid, and in dimethyl aniline. The temperature coefficient determined by Hinshelwood' for malonic acid and that determined by Bernoul!i and Jakubowicz9 for malonic acid in water solution are also shown.
4
x103
FIG.I A Benzyl malonic acid, no solvent. Daus and Burk. U Benzyl malonic acid, solvent-oleic acid. D a m and Burk. x Benzyl malonic acid, solvent-dimethyl aniline. Daus and Burk.
0 Malonic acid, no solvent.
Hinshelwood.
X Malonic acid, solvent-water. Bernoulli and Jakubowicz.
It is to be observed that the point a t IIS'C. ( I / T = 0.002577) for the decomposition of benzyl malonic acid without solvent is off the straight line in the sense of being too slow. This temperature was below the melting point of the acid. There was a marked induction period in the pressure change-time curve, there being a sharp acceleration when hydrocinnamic acid began to form in quantities sufficient to effect appreciable solution of the benzyl malonic acid. Thus the solid acid has a velocity constant of a lower order. A similar phenomenon was found by Hinshelwood' for the decomposition of solid malonic acid.
1466
ROBERT E. BURK AND WILHELMISA DAGS
The energies of activation calculated graphically from the slopes are : calories for the decomposition of benzyl malonic acid alone, 10,360 calories for the decomposition in oleic acid, which did not cause a marked increase in rate over the temperature range investigated, and 6,090 calories for the decomposition in dimethyl aniline between 105' and I ~ S O C . , which did have a marked catalytic effect. (The value calculated for the lower temperature range was I 1,400 calories.) The graphs leave no doubt that the marked catalytic effect of dimethyl aniline is accompanied by a definite decrease in the temperature coefficient. 11,270
~
~
-----
Q Percentage change of A Percentage change of 0 Percentage change of A Percentage change of
pressure with time at 90°C. pressure with time at 100°C. conductivity with time at 90°C. conductivity with time at 100°C.
It is seen that there is a change in slope in the sense of the reaction being too slow at the lower temperatures, beginning about 105' for benzyl malonic acid in dimethyl aniline. I t was suspected that t,his was due to a dependence upon dissociation of the salt to free base and acid, which would lead to lower concentration of free acid and consequent diminution in rate. I n order to investigate this point conductivity measurements were made, using a small cell with platinum electrodes, and a Wheatstone bridge. Both high frequency alternating current with phones and elect'ron tube amplification in the detecting arm, and direct current with a galvanometer in the detecting arm were used. Fig. z shows the percent change in conductivit,y plotted against time at 90' and 100°C. The percent pressure change is also plotted against time for
1467
CATALYTIC EFFECTS OF SOLVENTS
benzyl malonic acid in dimethyl aniline a t the same temperatures. The fact that the conductivity would approach a value so small as to be immeasurable with our equipment shows that the end product, hydrocinnamic acid, has a negligible conductivity under these conditions. That the conductivity should drop to one percent of its original value a t 90 minutes and IOO’C., while the pressure change curve indicated that sixteen percent of the original benzyl malonic acid remained, showed that this acid also has a very small conductivity in dimethyl aniline under these conditions. These facts justify the conclusion that the observed conductivity is due almost entirely to the salt. The assumption of direct proportionality between conductivity and concentration of salt should also be justifiable as a first approximation.
TABLE IV Temperature (“C.)
80 90
95 IO0
105
Half life in minutes Salt Acid decompn. decompn.
65.3 19.0 (10.24) 7.5 (3,236)
-
73.5 38.5 26.0 17.2
Temperature (“C.) I10 I1 j
125
I3 5 I45
Half life in minutes Salt Acid decompn. decompn.
(1.85) (1.084)
10.2
6.8 3.9 2.6
Table IT’ shows a comparison of half lives of decomposition of the salt to free base and acid and of the acid to hydrocinnamic acid and carbon dioxide. The values in brackets are calculated from the temperature coefficient as determined from the measured points, which lay on an excellent straight line when log t1I3was plotted against I ~ T . Fig. 3 shows the values of the time of half change in conductivity plotted against the reciprocal c +the absolute temperature with the corresponding half pressure change values for the decomposition of benzyl malonic acid in dimethyl aniline as solvent and without a solvent. It is clear from the table and graphs that the rate of salt decomposition becomes very rapid compared with the acid decomposition at a point corresponding to that a t which the temperature coefficient of the acid decomposition settles out to new and lower, but constant value. Furthermore the temperature coefficient for this reaction at low temperatures tends to assume the corresponding value for the salt decomposition. With concurrent reactions, that reaction having the higher temperature coefficient may predominate a t higher temperatures in case it is homogeneous, and that with the lower temperature coefficient is heterogeneous.13 This phenomenon depends upon the reaction with the lower temperature coefficient being heterogeneous, in which case but a small proportion of the total reactants is undergoing heterogeneous reaction a t a given instant. 13 This type of composite reaction is discussed by Hinshelwood: “Kinetics of Chemical Change in Gaseous Svstems,” 2nd edition, p. 5 . Examples are the combination of hydrogen and sulfur (NorrLh and Rideal: J. Chem.%oc., 123, 696 (1923)) and the decomposition of hydrogen iodide in pyrex (H. A. Taylor: J. Phys. Chem., 28, 984 (1924).)
I 468
ROBERT E. BURK AND WILHELYINA DAUS
’/T
*103
FIG.3 A Benzyl malonic acid, no solvent.
0 Benzyl malonic
acid in dimetbyl aniline X Conductivity due to salt.
This behavior would not be found with two homogeneous reactions of the same type, for if their velocity expressions were:
where
I.
kl = pl e-A1 RT
2.
kz
= pz e
-A2 IRT
or
kl = plnle-A1’RT k2
=
p2n2e-A2/RT
p is a probability factor k is the velocity constant n is the number of collisions per second A is the energy of activation.
If p1 and p2, and nl and n2 are not greatly different, then if A1 < A P , reaction ( I ) will always exceed reaction ( 2 ) in speed in spite of a relative increase in the speed of ( 2 ) . With consecutive reactions, as in our case, the reaction with the lower temperature coefficient not only increases in speed relative to the other, but likewise predominates as the temperature is increased. It is seen in Fig. I that in the results of Bernoulli and Jakubowicz, there is likewise a falling off in the rate of decomposition of malonic acid in water at the lower temperatures. Here salt formation is out of the question. However, since these temperatures are far below the melting point of malonic
CATALYTIC EFFECTS O F SOLVENTS
I469
acid, association of these molecules may occur, leading to a diminished reaction rate for the complexes, as was found for the solid malonic acid by Hinshelwood, and for benzyl malonic acid by us. Any residual effect due to incomplete salt decomposition a t the higher temperatures could only make the temperature coefficient appear higher than its true value, so that the conclusion that the temperature coefficient is lowered by the catalyst cannot be jeopardized. Such demonstrations of temperature coefficient lowering with catalysis, though suspected, are not frequent in the 1iterat~re.l~ Interpretation of the Lowering of the Temperature Coefficient A great deal of evidence has been accumulated to indicate that a major action of catalysts is a lowering of the energy of activation, the dominant factor governing reaction rates.I5 An interpretation which has been given to this is that it may be brought about under proper conditions by the stretching of bonds by a particular kind of association between catalyst and reactant involving more than one point of contact, a point of view which has been shown to be in excellent agreement with those phases of catalysis previously most difficult to explain, namely specificity, and promoter action.16 The preponderance of evidence was in favor of this type of distortion rather than the atomic distortion mechanism’’ for the influence of a catalyst upon a bond. Recent work on band spectra,18 and particularly upon Raman Spectra,lg tends to minimize the possibilities of the atomic distortion mechanism. Some new suggestions arise from Quantum Mechanics, according to which the principal component of non-polar bonds is the resonance or exchange force arising from a coupling of the spin moments of unpaired electrons in the respective atoms making up the bond, such a force being demanded by one of the solutions of Schrodinger’s wave equation, resulting from an exchange of the two electrons which potentially can form a bond.m According to Londonz1the following notions regarding chemical activation result. If we have a pair of atoms forming a non-polar bond and a third l4 Some new cases, e.g., the decomposition of ethers catalyzed by halogens and organic halogen compounds have recently been reported. (Clusius and Hinshelwood: 2. Elektrochemie, 36, 748 (1930) Summarizing paper.). l5 See e.g., Hinshelwood: “Kinetics of Chemical Changes in Gaseous Systems,” 2nd Edition; Burk: Sixth Report of the Committee on Contact Catalysis, J. Phys. Chem., 32, 1601 (1928). l6 See Burk: J. Phys. Chem., 30, 1134 (1926); 32, 1601 (1928); Balandin: 2. physik. Chem., B2,289 (1929). li See, e.g., Langmuir: Trans. Faraday Soc., 17, 610 (1922). See Burk: Sixth Report (Loc. Cit.) for a partial summary. lo See e.g., Smekal: 2. Elektrochemie, 36, 618 (1930); Andrews: Phys. Rev. (21, 36, 546 (1930). * O Heitler and London: 2. Physik, 44, 455 (1927); Rouark m d Urey: “Atoms, Molecules and Quanta,” p. 684 (1930); Heitler: Physik. Z., 31, 186 (1930); Mulliken: Chem. Rev., 6, 536 (1929). This relatively simple concept seems, however, to be growing more complex, because of a “degeneration” of orbital momentum of electrons as well as of the spin moments leading to binding forces of the same order of strength. (Heitler: Naturwissenschaften, 17, 546 (1929)). *lZ.Elektrochemie, 35, 5 5 2 (1929).
ROBERT E. BURK A S D WILHELMINA DAUS
1470
atom approaches which is capable of “resonance” combination with either of the two atoms forming the bond, the bond will be weakened, and if the potential reaction is exothermic the new combination will be effected, and the second atom which was a part of the original bond ejected. However, the third atom experiences a repulsion before approaching sufficiently closely to cause the phenomena just described. The energy of activation would thus be that necessary to overcome this “potential wall” in the case of exothermic “atom” reactions, and has been found by Polanyi and co-workers?* to be vanishingly small for reactions of the type:
AB
+ (‘ = AC‘ + 3 + energy
London then says, p. 5 5 3 , “If the distance between the atoms of the original bond is increased, as by increase in vibration or through the influence of a nearby catalyst, the resonance repulsion against the approach of the third atom will be weakened, and the height of the potential wall lowered.’’ This, in common with previous attributes the activation process to increasing the distance between the atoms of the bond concerned. However, in contradistinction, the energy of activation as seen by London is external, Le., the energy necessary to overcome the potential wall rather than the remaining energy necessary to break the bond (for exothermic reactions). In case the reaction is endothermic, the heat of react’ion per rnolecule must be added to that necessary t o overcome the potential wall in order to have a successful elementary act of react,ion. The small activation heats of atomic reactions are attributed by London to the existence of “sensitive spots” on the diatomic molecule easily reached by the approaching atom. For a reaction of the type:
XY
+ ZIT’ = YZ + sw
both X and Y cannot reach the sensitive spot of Zn‘, etc., so activation heats of higher order result, which London says should be close to the energy requirements of: (I) ST + Z I T = Y-z s IT’ or (2) X T zn- = Y z.
+
+ +
xn-+ +
This works out excellently with the homogeneous bimolecular reaction:
+
I 2 Energy of activation, 44,000 calories2‘ H P = zH - 98,100 i 2,300 caloriesz6 1 2 = z I - 3 5 , 6 0 0 =k 70 calories?j
zHI = Hp
HI = H HI = H
+ I - 66,000 calories (spectroscopic data)zfi + I - 6 9 , 3 0 0 calories (thermal determination)?fi
22 Beutler and Polanyi: 2. physik. Chern.. B1, 3; Bogdandy and Polanyi: and Schoy: 30; Outuka and Schoy: 62 (1928~. 23 Burk: J. Phys. Chem.. 30, 1 1 3 4 (1926). 24 Lewis: J. Chem. SOC.,113, 471 (1918). *5 Franck: Z. Elektrochemie, 36, j 8 7 (1930). z6 Sponer: Z. Elektrochemie, 34, 488 ( 1 9 2 8 ) .
21;
Polanyi
1471
ChTdLYTIC EFFECTS O F SOLVENTS
+
~ € 1 1= Hz 2 1 - 34,900 calories if the spectroscopically determined heat of dissociation of HI is used, and 40, joo calories if the thermally determined value is used z H I = Iz zH gives the wrong order of values for the heat, of activation. Polanyi?’ in applying these principles to activation at surfaces, considers again that activation involves stretching of bonds, and in still closer analogy to pre-quantum mechanical concepts, considers that activating adsorption of a molecule AB involves two surface atoms; surface complexes are formed in which the bond AB is either stretched or broken, and the bonds catalyst-A and catalyst-B must of course be of such strength that they lead to increased reaction rates. The correlation with London’s “sensitire spot” concept is not so lucid. If the catalyst atoms are spaced somewhat farther apart than the atoms of AB, and we have two point adsorption, then clearly the resonance weakening of the bond would be assisted by mechanical stretching.?l The most recent contribution of quantum mechanical theory to the concepts of factors participating in catalytic activation is due to Born and Franck,?g and has been but very superficially revealed. Referring to the association of energy of activation with that required to overcome a potential wall, they say that if on a catalyst the reacting molecules can be held a t a distance from each other comparable with the distance reached in a gaseous collision, and for a time long compared with the duration of a gaseous collision, there is a certain probability of reaction although the energy necessary to overcome the potential wall is not present. This probability depends upon the height and breadth of the potential mall, and upon the mass of the atoms, and is exceedingly sensitive to distance apart a t which the atoms find themselves as a result of adsorption and to the time during which they are held in this position. This point of view emphasizes once more the correlation between the spacing of catalyst atoms and their specific action, which various chemists have sought to bring out.30 However, the energy of activation which is found for catalytic reactions from this viewpoint presumably would be due to the energy requirement for the adsorption of the reactants in the sensitire configuration, in case the reaction probability is itself independent of temperature, although the previously described processes could go on simultaneously and involve temperature coefficients. The energy of activation may thus be: ( I ) that which is required for the heat of reaction in the case of endothermic reactions (including in this class the first step of exchange reactions, e.g., XY Z W = XZ Y IT),( 2 ) that required to overcome a potential wall, (3) that required t o produce a sensitive configuration on a surface. In catalysis, then, either ( I ) or ( 2 ) (principally (I)?) is lowered as compared with the homogeneous values, or (3) is substituted.
+
+
2i
+ +
Z. Elektrochemie, 35, 561 (1929j.
? 8 C f .Burk: J. Phys. Chem. 30. 1134 (1926). Z. Elektrochemie, 36, 588 (1930). e.g. Emil Fischer in his 1ock.and key concept of the specificity of enzymes; Adkins and coworkers: J. Am. Chem. Soc., 45, 809 (1923); 46, 130, 2291 (1924); Burk: 1,oc. cit.; Balandin: Lo?. cit. 2g
30
I472
ROBERT E , BURK AND WILHELMINA DAUS
Just what the activating mechanism is, for dimethyl aniline on benzyl malonic acid, is a matter of speculation. The principles developed for heterogeneous catalysis may be used in the main for homogeneous catalysis, and a large organic molecule as a catalyst, gives plenty of opportunity for spacing effects, etc. The catalytic action is probably dependent, however, upon the “free valencies” of the nitrogen in the present case. Furthermore, since the salt has been found to be less reactive than the free acid, the carboxyl hydrogen atom is probably involved in the first elementary act. The activation may consist in a mild excitation of the electron taken up from the hydrogen atom by the carboxyl group leading to dissociation of the carboxyl into neutral hydrogen and free RCHCOOH COO radicals;31 this would be a still different type of activation. This may be catalyzed by dimethyl aniline if the nitrogen of the aniline gives up an electron t o the hydrogen ion, and then takes one again from the carboxyl ion, in case these two energy s k p s are appreciably smaller. There are no evident unpaired electrons in dimethyl aniline available for quantum mechanical coupling which might loosen the OH bond of a nonionic form of the free acid through resonance action, although a collision with an acid molecule might render them available for coupling so that the dimethyl aniline could fulfill its catalytic duty of bond loosening in the quantum mechanical sense. That there is a type of quantum mechanical perturbation which does not depend on unpaired electrons is indicated in London’s new interpretation of van der Waals’ forces.32 That secondary valence forces have much to do with the present case is, however, not indicated since most solvents, including some which are quite polar, have little influence. With respect to the effect of replacing the hydrogen atoms alpha t o the carboxyl groups by other groups, there is no certain evidence of a lowering of the heat of activation, though a slight lowering is indicated when one benzyl group is introduced, and the rate is certainly increased. h plot of Bernoulli and Jokubowicz’ values showed that only their work on malonic acid gave a good straight line when I / T was plotted against log k, the data for the other acids being erratic. I t has been said that on account of results from investigations of band spectra and Raman spectra, effects corresponding to a distortion carried through the chain should be minimized. This is confirmed by the observation of Bernoulli and Jakubowiczg that when both alpha hydrogens are replaced by like substituents the decomposition rate is less than for malonic acid alone, in all cases. Whereas when only one is substituted, the decomposition rate is greater, for phenyl, benzyl, allyl, chlorine and hydroxyl. This would scarcely follow if the effect of substitution on decomposition rate were due to distortion carried through the chain. Furthermore, the halogen substituted acids did not decompose particularly rapidly. This is contrary to 31 See Franck: LOC. cit., for such processes of dissociation induced photochemically in polar unions. 32Z. Physik, 63,245 (1930).
CATALYTIC EFFECTS OF SOLVENTS
I473
the effect of halogen substitution on the strength of acids. This, however, may be a space effect rather than a chain effect.33 Since in solution there would always be enough activating collisions to maintain the Maxwellian distribution of energy between acid molecules, the reaction rate would not fall off with diminished concentration and the number of participating internal degrees of freedom cannot be calculated. The heat of activation, reaction rate, and temperature coefficient do, however, fall into line fairly well with other monomolecular reaction^.^'
Summary The decomposition rate of malonic acids, especially benzyl malonic acid, have been determined both for the acid alone and in various solvents. Dimethyl aniline was the only solvent, of those investigated, to show marked catalytic action. This was accompanied by a definite lowering of the temperature coefficient. The effect was shown not to be due to salt formation. Possible mechanisms for activation by the catalyst have been discussed. Morley Chemical Laboratory, 1Vestern Reserae Uniuersity, Clewland, Ohio. Decemher, Pi?, 1930.
33This is suggested by the work of McInnes: J. Am. Chem. SOL,50, 2j87 (1928). 34 See Hinshelwood: "Kinetics of Chemical Changes in Gaseous Systems, 2nd. Ed., p. 159. The reaction rate of XzOSin Hinshelwood's table would be too high for the purposes of the table if a chain mechanism applied to this case.
