30
L. A. Cosby and G. L. Humphrey
tion will vary somewhat with different crystals and hence so will ANOz-. It appears therefore that ANOZ- is a measure of the color center concentration and that the reduction of ceric ion is another chemical method of determining these centers. The question remains, however, as to the nature of the color center. The center described by Pringsheim has its maximum a t 335 mp and cannot reasonably be attributed to either NO2 or NO in the lattice, since the absorption spectra of these molecules when observed in an inert mat r i are~ very ~ far ~ removed ~ ~ from ~ 335 mp. Pringsheim attributes the color center to an electron trapped in the vicinity of the nitrite ion. We do not think that the color center can be an electron trapped in an anion vacancy as are, for example, the color centers observed in the alkali halides. In these latter salts the color centers are easily annealed a t room temperature and below whereas in NaN03 a minimum temperature of about 160' is required. We are inclined to accept the view that the color center is an electron trapped in the vicinity of a nitrite ion; however, it must be very tightly bound and is more likely associated with a relatively stable molecular species. The results of Zeldes and LivingstonlGaclearly indicate that the nitrite ion is the precursor of the NO2 molecule observed in the esr spectra a t low temperatures and that the actual formation of nitrite ion requires some thermal activation, i.e., a t liquid nitrogen temperature the nitrite ion is nut formed but rather some intermediate which, when the irradiated sample is warmed to room temperature, produces nitrite ion. The color center, we believe, is associated with this intermediate. Since the concentration of this species a t room temperature is relatively high, -0.15 pmol/g, it cannot be NOS, N022-, or N03*- since in this concentration range
Kinetics of
cid ~
these species are easily detected by paramagnetic resonance techniques. We suggest that the center is an 0- or an 0 2 - associated with the nitrite ion/or ions formed simultaneously in some sort of a complex.
References and Notes (1) Research performed under the auspices of the U.S. Atomic Energy Commission Contract No. AT(40-1)-3660 To whom correspondence should be addressed. For a general review of nitrate radiolysis see, E. R. Johnson "Radiation Induced Decomposition of Inorganic Molecular Ions," Gordon and Breach, New York, 1970. J. Cunningham, J. Phys. Chem., 67, 1772 (1963). M. E. Shinn, ind. Eng. Chem., Anal. Ed., 13, 33 (1941). J. Cunningham, J. Phys. Chem., 70, 30 (1966). J. Cunningham, h t . J. Phys. Chem. Solids, 23, 843 (1962). J. Cunningham, J. Phys. Chem., 66, 770 (1962). S. Kaucic and A. G. Maddock, Trans. Faraday SOC., 65, 1083 (1969). W. B. Ard, J. Chem. Phys., 23, 1967 (1955). E. Bleany, W. Hayes, and P. M. Liewellyn, Nature (London), 179, 140 (1957). H. Zeldes, "Paramagnetic Resonance," Vol. ii, Academic Press, New York, N.Y., 1963, p 764. D. Shoemaker and E. Boesman, C. R. Acad. Sci., 252, 2099, 2865 (1961). R Adde and P. Petit, C. R. Acad. Sci., 256, 4682 (1963). See also ref 1, p 80. (a) H. Zeldes and R. Livingston, J. Chem. Phys., 35, 563 (1961); 37, 3017 (1962); (b) See W. G. Burns and T.F. Williams, Nafure (London), 175, 1043 (1955); J. G. Rabe, B. Rabe, and A. 0. Allen, J. Phys. Chem., 70, 1098 (1966); E. A. Roja and R. R. Hentz, ibid., 70, 2919 (1966). R. W. Holmberg, personal communication. H. Pogge, Thesis, Stevens Institute of Technology, Hoboken, N.J.; University Microfilm, Ann Arbor, Mich., No. 68-16, 222. E Ladov and E. R. Johnson, J. Amer. Chem. SOC.,91, 7601 (1969). G. Hennig, R . Lees, and M. Matheson, J. Chem. Phys., 21,664 (1953). T. H. Chen and E. R. Johnson, J. Phys. Chem., 66, 2249 (1962). P. Pringsheim, J. Chem. Phys., 23,369 (1955). E. Hutchinson and P. Pringsheim, J. Chem. Phys., 23, 1113 (1955). J. Cunningham, J. Chem. Phys., 41, 3522 (1964). G. W. Robinson. M. McCarthy, and M. C. Keelty, J. Chem. Phys., 27, 972 (1957). H. Brion, C. Moser, and M. Yamazaki, J. Chem. Phys., 30, 673 (1959).
~ in Alkali ~ Halide~ Nlatrices1i2 1
~
~
~
s
L. A. C Q S ~and Y G. L. HMmpRley" Department of Chemistry, West Virginia University, Morgantown, West Virginia 26506 (Received June 12, 1974)
The thermal decomposition of malonic acid dispersed in alkali halide pressed pellets has been studied. The study was made by heating the pellets and observing the changes in absorption of several infrared bands of the reactant and products. The reaction was found to be first order with respect to disappearance of malonic acid with an apparent rate constant of k (seeT1) = 6.0 X 109 exp[(-23,400 f 300)/RT]. The effects of changing the matrix material are discussed.
Introduction
Although several ~ t u d i e s ~ have - ~ been carried out on the kinetics of the thermal decomposition of pure malonic acid in the molten state, none appear to have been reported for the reaction in solid solution. By employing a technique similar to that of Bent and Crawford,s we have made an infrared spectroscopic study of the thermal decomposition of malonic acid in the solid state at, or above, its melting The Journal of Physical Chemistry, Voi. 79, No. 1 . 7975
point. The reactant was incorporated in an alkali halide solid matrix and quantitative kinetic results were obtained by observing the change in infrared absorption of reactant and products as a function of time. Many investigations of the decarboxylation of malonic acid in numerous different solvents have been r e p ~ r t e d From .~ the proposed mechanisms for the reactions in solution it appeared that a study of the reaction in a highly ionic solid medium would be beneficial.
Kinetics of Malonic Acid Pyrolysis
Experimental Section Chemicals. Malonic acid was obtained from K and K Laboratories, Plainview, N.Y., and recrystallized from acetone a t low temperatures. The crystals were ground to a fine powder and placed under reduced pressure a t room temperature for several hours. Analysis showed the material to be 99.5 f 0.1% malonic acid with a melting point of 136'. The alkali halides were obtained from Brinkman Instruments, Inc., Westbury, N.V., and were manufactured by E. G. Merck, Darmstadt, Germany as Suprapur or Ultrapur. Instruments. The constant temperature bath was a Fisher unitized bath, specially insulated, and filled with UCQN fluid 50-HB-280-X obtained from Union Carbide Corp., Charleston, W. Va. Bath temperatures could be maintained by a mercury expansion controlled electronic relay system to f0.05'. A wooden rack containing holes for the reaction tubes was constructed for the bath to facilitate placing the samples in the bath. The alkali halide pellets were formed under 20,000 psi gauge pressure in a Carver hydraulic press using a 25 X 6 mm rectangular die manufactured by Beckman Instruments Co. A Mettler semimicrobalance accurate to 5 X 10-5 g was used to weigh samples. Spectra were obtained on a Beckman IR-12 spectrometer calibrated periodically by means of the polystyrene spectrum, or by scanning the atmospheric absorption in single beam mode. Absorption frequencies reported in this work are considered to be accurate to k1.0 cm-l. Absorbance values could, in general, be reproduced to better than ~k0.002absorbance unit. Pellet Preparation. Pellets were prepared both from the commercial alkali halides as received and from the commercial salts ground to a very fine powder in agate and evacuated for several days a t 250'. In the case of pellets pressed from predried salts, the salts were mixed with finely ground malonic acid in a rotating vacuum type evaporator which was simultaneously magnetically stirred. The die was filled in a dry box, closed, and sealed by applying pressure and then evacuated before pressing the pellet. Dry air was used to bring the pressure inside the die back to atmospheric after pressing the pellet. For the preparation of pellets from salts that were used as received, the die and salts were treated as above in a constant humidity room maintained a t about 40% relative humidity. Kinetics. In preparation for a kinetic run, 10 to 15 pellets were prepared as described above and weighed to the nearest 5 X 10-5 g, and the infrared spectrum of each was obtained. Absorbance values for the bands were determined by the baseline technique and the values from the initial spectra were used to establish the zero time absorbance values for each pellet. The pellets were then placed in individual sample tubes in the constant temperature bath. Each pellet was removed from the bath after elapse of the chosen specific time for the kinetic run, cooled t o room temperature either in a desiccator, or in a constant humidity room, and the spectrum ay,ain obtained. Timing for a kinetic run was made by means of a stop watch which was allowed to run continuously throughout the time period. In order to obtain good spectra after heating a pellet, it was necessary to repress the pellet. Heating of the pellets caused them to become opaque and expand slightly (possibly due to CQ2 evolution) so that the edges of the pellet
39
had to be sanded in order to replace it in the pellet die. The resulting weight loss from the repressing procedure necessitated correction of the absorbance values obtained after repressing a disk. By preparing a large number (50) of different weight pellets from a single mixture of malonic acid and alkali halide and obtaining absorbance values for each of the bands used in our study, it was determined that absorbance was a linear function of pellet weight. Furthermore, it was found that the intercept of a plot of absorbance us. pellet weight was a t the origin. These facts allowed the use of a simple weight ratio for correcting absorbance values to a value for a standard pellet weight, chosen as 0.20000 g. The average variance of the absorbance of the bands studied in this work as obtained from five different kinetic runs on the same bulk pellet mixture was 0.87 X absorbance unit.
Results By preparing pellets from mixtures of malonic acid and alkali halides in varying concentration ratios and comparing their spectra it was found that all major absorption bands in the malonic acid spectrum (3000,1700,1440,1310, 1220, 1175, 500, 760, 650 cm-l) obeyed Beer's law over the concentration range (0.05-1.0 wt %) employed for the pellets. In the initial stages of a kinetic run, the absorbance values for the pellets showed an increase. Since the working temperatures were generally above the rnelting point (136') of malonic acid, we attributed the initial increase in absorbance to malonic acid becoming more uniformly dispersed throughout the alkali halide matrix. Some workerslO have resolved the problems associated with abnormal initial changes in absorbance by subjecting pellets to an annealing process, but such a procedure was impractical in our work since temperatures as low as 80' were found to bring about thermal decomposition of malonic acid. ,Annealing a t lower temperatures had no effect on the initial absorbances of the pellets. The inability to anneal the pellets did not create major problems since the time required for the absorbance to reach maximum values was approximately 2 min a t 140°, after which time the absorbance values decreased in a regular fashion reflecting decomposition of the malonic acid. I t may be noted that the time required for the absorptions to reach a maximum was approximately equal to the time required for the pellet to reach bath temperature as measured by a thermocouple imbedded in some pellets. Since absorbances of the bands under study were found to be proportional to concentrations of malonic acid present, the order of reaction was determined from plots of the logarithm of the change in absorbance per unit of time for a particular band us. the logarithm of the absorbance for that band. The slope of the line obtained from such plots is equal to the order of the reaction. Applying this method to the 3 0 0 0 - ~ m -band ~ gave reaction orders of 1.18 a t 145", I .15 a t 161', and 1.08 a t 167'. The other bands gave values for the order of the reaction ranging from 0.59 to 1.4 with the average value for all bands being 0.57 f 0.10. Rate constants were determined by plotting the logarithm of the ratio of the initial absorbance values to the absorbance values a t time, t , each corrected to a standard weight pellet of 0.20000 g, us time, t. Five of the nine absorption bands studied for malonic acid gave good straight line plots from which rate constants could be determined, and the results are presented in Table I. The placement of The J o u r n a l of Physical Chemistry, Voi. 79, No. 1, 1975
40
L. A . Cosby and G . L. Humphrey
TABLE I: Malonic Acid Decomposition R a t e C o n s t a n t s (units = 10-3 sec-1) Temp, "C
+ 0.1'
3000
cm-1
1700
1440
cm-1
em-'
1220
cm-1
9QO
TABLE 11: Activation P a r a m e t e r s for the T h e r m a l Decomposition of Malonic Acid in Several Media
cm-1 Medium
Potassium Bromide Matrix 138.9 142.0 144.9 149.0 153.2 158.1 161.0 167.0
1.02 1.54 2.23 2.38 2.83 4.67 5.20 7.01
158.5
4.65
1.05 1.39 2.11 2.39 2.59 5.04 4.98 8.53
1.21 1.16 1.60 2.56 2.96 5.12 4.98 6.46
1.10 1.75 2.08 2.61 3.37 4.24 5.48 9.41
0.81 1.15 3.92 5.18 6.24
Rubidium Bromide Matrix 4.20
4.21
1-Butanok 1-Hexanola 2-E thylhexanol-la Diisobutylcarbitola Cyclo hexanola Acetanilide2 Molten malonic acidh Malonic acid in KBrc a
- AGS at 135O,
-AH$, kcal, mol
kcal i mol
27.2 -4.4 26.0 -7.6 24.8 -10.4 24.8 -10.7 23.0 -15.0 19.85 -22 66 35.8 111.9 2 2 . 8 i 1 - 1 7 . 4 zk 1
29.0 29.1 29.05 29.16 29.13 29.1 31.0 2 9 . 8 zk 1
Reference 13. Reference 15. This work.
4.28
the lines from which the slopes (rate constants) were determined was made by a least-squares computer program ~ ~ ~ The avwhich was a modification of the L S K I N program. erage variance of absorbances for the absorption bands of malonic acid used in determining the rate constants, which is based on five kinetic runs made from the same bulk pelabsorbance unit. This results in let mixture, was 1.0 X an uncertainty of 10% in the rate constant values. The first-order rate constants for the decomposition of malonic acid in a rubidium bromide matrix gave results comparable to those obtained in the potassium bromide matrix. The spectra for malonic acid obtained in other matrix materials (NaC1, NaBr, NaI, KCl, KI, CsC1, CsBr, CsI, RbC1, RbI) were indicative of reaction of malonic acid with the matrix material.12 The 900-cm-l absorption band in some instances (144.9, 149.0, and 167.0' in KBr and 158.5' in RbBr) gave such random points that no sensible least-squares fit for the data could be found. From the experimentally determined rate constants and their corresponding temperatures, the thermodynamic parameters shown in Table I1 were calculated on the basis of the absolute reaction rate theory equation, k = (constant)(T) exp(AS#/R) exp(-AH*/RT) and AGf = A@ T AS#. A comparison of values obtained by other investigators in other media is also presented in Table 11. A plot of the negative logarithm of T/k us. 1/T based on the absolute reaction rate theory equation, from which AS2 (intercept) and -AH7 (slope) were determined for the five absorption bands of malonic acid used in this work, is presented in Figure 1. The activation energies and frequency factors for the bands studied are summarized in Table 111.
Discussion The thermal decomposition of malonic acid in potassium bromide and in rubidium bromide, as in some other media,l*J5 appears to be a straightforward decarboxylation to yield carbon dioxide and acetic acid. The infrared spectrum of carbon dioxide trapped in the pellets was evident in the early stages of heating and the infrared spectra of the pellets in the final stages of heating showed only acetic acid to be present. Fraenkel, et a1 , 1 6 have proposed a mechanism for the thermal decarboxylation of acids which appears to have been generally accepted. The mechanism involves the formation of an activated complex between an electrophilic carbonyl carbon atom and a nucleophilic center in the solvent medium, illustrated as follows The Journai of Physicai Chemistry, Voi. 79, No. 7 , 7975
TABLE 111: Frequency Factor a n d Activation Energy for Malonic Acid Decomposition in Potassium Bromide Ir band frequency, em - l
Activation energy, kcal, mol
3000 1700 1440 1220
Frequency factor, see-'
23.0 23.2 23.6 23.9
2.0 1.2 3.9 6.2
x x x x
109 1010
109 109 V
o
g
-J
o
n
0
4.E
4 -
I
1
,
I
?.2j
2.3:
2,;;
1
.zy
- -~
I 2.39
i.?7
I ;.-I
Flgure 1. Combined enthalpy of activation plot for the thermal decarboxylation of malonic acid in KBr at temperatures from 138.9 to
167.0'.
R\-/
I
H
R'
R\ /R'
1
H
H-e\ H-0
/"
F=% K'
R
+
CO?
A --t
\
H-c-c
//O
'/E
'e-H
Support for the above mechanism is found in the fact that AH# decreases with solvent basicity, and there is a linear dependence of the rate of decarboxylation upon concentration of a strong amine base, such as pyridine,16 the reaction rate dependence disappearing a t a 1:1 mole ratio of base to malonic acid, indicating a bimolecular reaction. Kenyon and Ross17 observed that unsymmetrically disubstitued malonic acid did not produce optically active di-
41
Kinetics of Malonic A c i d Pyrolysis
substituted acetic acid, which is further support for the above mechanism. The production of a racemic mixture of disubstituted acetic acid would require formation of a carbon-carbon double bond a t some point in the reaction mechanism. Blades and Wallbridge18 observed that deuteriomalonic acid exhibits no kinetic isotope effect. The absence of a kinetic isotope effect suggests that the rate-determining step in the decarboxylation is the breaking of a carbon-carbon bond rather than the formation, or dissociation, of a hydrogen bond. Loudon, e t al., also conclude from l3C/lzC isotopic effects that both carbonyl groups must be involved in the transition state, and that bond breaking is almost complete between the leaving carbon and the central carbon atom. These workers also investigated an oxygen isotope effect by measuring the variation with per cent reaction in the '80 content of the COz released. Since no variation in I S 0 content was found, bond breaking and bond making in the transition state is again suggested. Clarklg has shown that when AH? is plotted against AS; for a given series of reactions in various solvents, the result is a straight line of slope T , the isokinetic temperature. Other workers20 have pointed out limitation in the use of isokinetic temperatures for drawing valid conclusions, and have established requirements which must be met before valid conclusions can be drawn. The data obtained by Clarkz1 appear to satisfy the requirements, and for the solvents studied thus far for the decarboxylation of malonic acid, AH$ and AS# values, when plotted, form a straight line with a slope of 422 K (149') for the isokinetic temperature. The AH$ and AS* values for the decarboxylation of molten malonic acid also fall on the same straight line as those for various solvents including the solid alkali halides used in this work. Apparently the same mechanism of reaction prevails in each case, and the fact that the AH# value for the decarboxylation of malonic acid is significantly lower in potassium bromide than for molten malonic acid could be the result of the activation effect of the bromide ion in the alkali halide medium. A comparison of the reaction rates for the decarboxylation of malonic acid in KBr to those in RbBr reveals no significant differences. This might be expected if the mechanism does involve the nucleophilic halide ion. Investigation of the effect of the anion upon the reaction rate would be most meaningful in this case, but is impractical due to the interfering side reactions taking place between malonic acid and other alkali halides.12 An interesting observation made in this research was that solid malonic acid undergoes decarboxylation at mea-
surable rates well below its melting point. We found that malonic acid decomposes slowly in a KBr matrix at 80' and the reaction half-life is approximately 6 weeks. This temperature is well below the reported solid state phase transiand therefore it appears that tion temperaure (130°),22~23 malonic acid can decompose from the (3 phase as well as the LY phase. The multipellet technique developed for the kinetic studies in this work has enabled us to collect data for the decsrboxylation of malonic acid at temperatures considerably higher than were previously possible, and allows the study of reactions with half-lives as low as 4 min or less without significant problems.
Acknowledgment. We are indebted to the National Science Foundation for funds to purchase the Beckman IR12 spectrophotometer, and to Dr. G. A. Hall for many helpful discussions. References and Notes Abstracted in part from the Ph.D. Dissertation of L. A. Cosby. West Virginia University, 1969. This work was supported by the National Aeronautics and Space Agency, Grant No. NLG-4900-001, formerly NsG-533. C. N. Hinshelwood, J. Chem. SOC.,117, 156 (1920). J. Laskin, Trans. Sib. Acad. Agr. for., 6, 7 (1926). J. Gyore and M. Ecet, J. Therm. Anal., 2, 397 (1970). A. S. Wipf. Diss. Abstr., Int. T, 31, 4623 (1971). A. G. Loudon, A. Maccoii, and D. Smith, J. Chem. SOC.,faraday Trans. 7, 69, 894 (1973). H. A. Bent and B. Crawford, J. Amer. Chem. Soc., 79, 1973 (1957). L. W. Clark, "Decarboxylation Reactions," in "The Chemistry of Carboxylic Acids Esters," S. Patai, Ed., Interscience, New York. N.Y.. 1969, pp 589-622, Chapter XII. K. 0. Hartman and I. C. Hisatsune, J. Phys. Chem., 69, 583 (1965). C. E. Detar and D. F. Detar, Program 80, Quantum Chemistry Program Exchange, Indiana University, 1967. L. A. Cosby and G. L. Humphrey, to be submitted for publication. L. W. Clark, J. Phys. Chem., 66, 125 (1962). L. W. Clark, J. Phys. Chem., 71, 302 (1967). L. W. Clark, J. Phys. Chem., 67, 138 (1963). G. Fraenkei, R. L. Belford, and P. E. Yankwich, J. Amer. Chem. Soc., 76, 15 (1954). J. Kenyon and W. A Ross, J. Chem. Soc., 3407 (1951). A. T. Blades and M. G. H. Wallbridge, J. Chem. SOC., 792 (1965). L. W. Clark, J. Phys. Chem., 68, 3048 (1964). R. C. Petersen, J. H. Markgraf, and S. D. Ross, J. Amer. Chem. Soc., 83, 3819 (1961). L. W. Clark, J. Phys. Chem., 71, 2597 (1967). W. W. Wendlandt and J. A. Hoiberg, Anal. Chem. Acta, 28, 506 (1963). The temperature of the phase change was determined by differential thermal analysis2' and is probably about 25-30' too high judging from our infrared spectral studies where we detected the phase change by means of reversible changes in the absorbtion spectrum in going through the transition temperature. From a comparison of the temperature of fusion for malonic acid as determined by the standard melting point technique (136') and that determined by differential thermal analysis (160°),22 we also conclude that the reported temperature of the phase change is too high. Our spectral studies of malonic acid in alkali halides show the /3 to a phase change occurring at about 90'.
The Journalof Physical Chemistry, Vol. 79, No. 1, 1975
