407 1
REACTION BETWEEN BENZOYL PEROXIDE AKD RHODAMINE 6GX COLOR BASE
importance of other factors besides conjugation (such as intramolecular va8nder Waals interactions) for de\
termining rotational barriers in
C-C
/’
/ \
bonds is
borne out. The Huckel calculations lead to very similar values of AE“ in 2-furanaldehyde and 2-pyrrolealdehyde. Preliminary rate measurements on the interconversion of the rotational isomers in the latter compound have indicated the rotational barrier to be comparable to that in. 2-furanaldehyde.
Acknowledgments. The authors wish to thank Dr. E. Forslind for his kind interest in this work and Mr. Torbjorn Alm for his skillful programming of the lineshape equations. Thanks are also due to the Computing Division of the Swedish National Office for Adminstrative Rationalization and Economy for providing free access to the Swedish electronic computer BESK. The work has been financed by the Swedish Technical Research Council and the Swedish Natural Science Research Council.
Reaction between Benzoyl Peroxide and Rhodamine 6GX Color Base
by P. K. Nandi and U. S. Nandi Department of Physical Chemistry, Indian Association for the Cultivation of Science, Jadavpur, Calcutta-82, India (Received February 28, 1966)
The reaction between benzoyl peroxide and Rhodamine 6GX color base in five hydrocarbon solvents has been studied in the temperature range where the thermal decomposition of benzoyl peroxide is negligible. I n the very first step of the reaction an intermolecular complex is probably produced by electron transfer. The mechanism of the reaction has been discussed In the light of thermodynamic parameters.
Introduction The reaction of benzoyl peroxide with amines is interesting and has been studied by other workers.’ The reaction is reported to be comparatively fast at temperatures where the thermal decomposition of benzoyl peroxide is negligible. There are two mechanisms suggested for such a rapid reaction between peroxide and amines. One2is that the induced decomposition of peroxide by amine-type radicals may be an extremely rapid chain reaction, while the other314 hypothesizes a bimolecular reaction between the peroxide molecule and the amine molecule leading to the peroxide decomposition. The bimolecular reaction would most probably be a one-electron-transfer reaction. Either possibility is cogent although neither is compelling. The disparity between the proposed mechanisms of the reaction warranted the investigation for the eluci-
dation of the involved mechanism. Here we have chosen a dye base as the amine, having chromophoric >NH groups, which has the formula
A sharp change in the color of the amine dye base as the time proceeds after the addition of benzoyl per(1) A. V. Tobolsky and R. B. Mesrobian, “Organic Peroxides,” Interscience Publishers, Inc., New York, N. Y.,1966. (2) P.D.Bartlet and K. Nozaki, J. Am. Chem. Soc., 68, 1686 (1946); 69, 2299 (1947). (3) L. Homer, Angew. Chem., 61, 468 (1949). (4) M.Imoto and 5. Choe, J. Polymer rSci., 15, 486 (1965).
Volume 69, Number 12 December 1986
4072
oxide has been attributed to a possible charge-transfer phenomenon, thereby supporting the view that as an intermediate of the reaction an intermolecular complex is produced. This occurs by the donation of an electron from the amine to benzoyl peroxide, thereby decomposing the peroxide. The variation of the rate constant with temperature could be utilized for the computation of thermodynamic parameters of activation which are of immense value in elucidating the mechanism. I n the present paper both the amine dye and benzoyl peroxide were dissolved in hydrocarbon solvents, and the rates were measured at different temperatures.
P. K. NANDIAND U. S. NANDI
0.50
0.40
B
V
3u 0"
.I
0.30
Experimental Section Rhodamine 6GX (C.I. No. 45160) was purified by using ether to precipitate the dye from a saturated solution in absolute a l ~ o h o l . ~The sample supplied was the chloride salt. The color base of the dye was extracted in hydrocarbon solvents following the procedure developed in this laboratory.6 The Rhodamine 6GX color base extract was kept over sodium hydroxide beads. The benzoyl peroxide was reprecipitated three times, using chloroform, from a solution of benzoyl peroxide in methanol.' It was kept in a desiccator containing calcium chloride in a cool, dry place. The solvents heptane, benzene, decalin (mixed isomer), and toluene were purified by the method suggested in the references.8 Cyclohexane used was B.D.H. spectroscopic grade. The characteristic yellow color of the Rhodamine 6GX color base is only obtained when it is extracted with hydrocarbon solvents. It was found that the greater the unsaturation of the solvent, the better the extraction. Any other solvent, with the exception of hydrocarbons, gives a characteristic pink color with or without fluorescence. Our experiments are limited to hydrocarbon solvents only. The concentration of the color base in hydrocarbon solvents can be determined by shaking the nonaqueous layer with a solution containing 1 M HC1 and 3 M KC1 solution. The extract in the acidic layer first turns to a pink color, which, when kept for some hours, becomes completely yellow. The concentration of this solution may be obtained from the standard graph of optical density us. concentration of Rhodamine 6GX chloride dissolved in the mixed HCl-KC1 solution. Actually, the exact concentration of Rhodamine 6GX color base is not necessary as will be seen later. The progress of the reaction of the Rhodamine 6GX color base with benzoyl peroxide was investigated by measuring the optical density a t the wave lengths of maximum absorption a t different intervals of time using The Journal of P h y s a Chemtktry
0.20 0
10
Time, min.
20
30
Figure 1. A typical plot of the decay of Rhodamine 6GX color base and consequent formation and decay of the intermediate with time in decalin: 1, measurements at 484 mp; 2, measurements at 515 mfi. Benzoyl peroxide concentration -2.5 X M; temperature 32 =k 0.1".
a Hilger Uvspeck spectrophotometer. Stoppered quartz cells of 1-em. thickness were used. The temperature in the cell compartment wais maintained to within f0.1O of the desired value by circulating water in the thermospacer set lfixed in the spectrophotometer from an external thermostat controlled by a relay. Results and Discussion The visible spectrum of the extracted color base in heptane shows absorption peaks a t 480 and 450 mp; in cyclohexane, absorption peaks are a t 480 and 452 mp; and in benzene, toluene, and decalin, they are a t 484 and 454 mp, The experiments were performed a t the wave lengths of maximum absorption, viz., a t 480 mp in heptane and cyclohexane and at 484 mp in benzene, decalin, and toluene. At these wave lengths the color base obeys Beer's law. ( 6 ) R. W. Ramette and E. B. Sandell, J . Am. Chem. SOC.,78, 4872 (1966). (6) 8. R.Palit, Chem. Ind. (London), 1631 (1960); Anal. Chem., 33, 1441 (1961). (7) c. a.swain, W. H. Stockmayer, and J. T. Clarke, J . Am. Chsm. ~ o c . ,72, 6426 (1960). ( 8 ) A. Weissberger, Ed., "Techniques in Organic Chemistry," Vol. VII, Interscience Publishers, Ino., New York, N. Y.,1966,pp. 311, 313, 317, and 319.
REACTION BETWEEN
f3ENZOYL PEROXIDE AND
4073
RHODAMINE 6GX COLOR BASE
0.60
0.70
. 0.50
5
.3
b
8
'a8
2 0.60
3
.3
u
0"
.3
i=
0"
0.40
0.50
0
0.30 5
Time, min.
10
5
10
Time, min.
Figure 2. Variation of optical density with time using 5.205 X M benzoyl peroxide at 24.9 zk 0.1' with different concentrations of Rhodamine 6GX color base in benzene.
Figure 3. Variation of optical density with time using nearly the same Concentration of Rhodamine 6GX color base with different concentrations of benzoyl peroxide in toluene at 26.1 f 0.1': 1, 5.036 X 10-4 M ; 2, 2.520 X M; 3, 1.008 X 10-4 M benzoyl peroxide.
The experiment was performed with a dye concentration of about 2 X M and a benzoyl peroxide concentration ranging from 5 X to M. I n the presence of benzoyl peroxide the characteristic yellow color of Rhodamine 6GX color base diminishes continuously, and a pink color develops. The pink color also disappears gradually, and a colorless solution is obtained in all solvents. At low benzoyl peroxide ill) and a t a temperature of 15' concentration the decay of the pink color is very #low (about 1% in 20 min.). The absorption maximum of the pink color, under these conditions, was found to be 515 mp. This maximum depends only slightly upon the solvents in which the experiments were performed. This absorption band corresponds with the acid color developed by the color base in acid solutions, e.g., benzoic, stearic acids, etc., in the above-mentioned solvents. The component which has the absorption band a t 515 mp has little effect on the abeiorption of the original color base as can be observed by comparing the spectra. At high temperatures the appearance of the pink color is very transitory. The fading away of the color of the intermediate compound is most rapid in heptane and slowest in toluene at any temperature. The comparison of the absorbance of a solution of benzoyl peroxide and Rhodamine 6GX color base kept
in dark, with an identical solution subjected to experimentation in the spectrophotometer, for the same period, shows very little perceptible difference, thereby eliminating the possibility of a photochemical reaction. The disappearance of the amine color, followed by the appearance of the pink color and its subsequent decay, points to the fact that the nature of the reaction is a consecutive one. However, the identification of the end product could not be made. From the comparison of spectra it seems that the intermediate compound decays giving more than one single product. Figure 1 illustrates the behavior of the reaction in decalin. Figure 2 illustrates the variation of optical density with time a t different concentration of Rhodamine 6GX color base using the same concentration of benzoyl peroxide. Figure 3 illustrates the same with a constant concentration of Rhodamine 6GX color base but different concentrations of benzoyl peroxide in benzene. Under the conditions used, the reaction obeyed zeroorder kinetics with respect to Rhodamine 6GX color base concentration and firstcorder kinetics with respect to benzoyl peroxide concentration. To establish the order of the reaction and to calculate the velocity constant K from the values of concentration a t different times, numerous methods have been sugge~ted.~The following equation was used for volume 6.9,Number 12
D E C E 1.966 ~ ~ ~ T
4074
P. K. NANDI AND U.
-
S.NANDI
in toluene. It is found that a linear plot of log K’ vs. 1/T is obtained in accord with the above equation. The slope of this plot yields a value of 13.2 kcal./mole for E*, the energy of activation in toluene. The value of the frequency factor A was obtained by computing the value of E* and K‘ a t a fixed temperature from the above equation. This gives a value for A in toluene which is 3.84 X lo8 M-l sec.-l. The values of E* and A in different solvents are listed in Table 11. Also from the data, the different thermodynamic parameters for activation were calculated using the equationslO
1.70
-1.60 d 6 I
3
-
1.50
AF+ = 2.303RT -
1.40
I
I
6
10
452O
, 16
AS* = 2,303R(log A
Time, min.
Figure 4. Variation of log (optical density) with time at different temperature in decalin.
obtaining the first-order rate constant: In C = In CO- Kt, where Co is the initial concentration, C is the concentration of the dye a t time t, and K is the firstorder velocity constant. According to this equation, the plot of log C vs. t would give a straight line for a first-order reaction, and it is actually found so, as shown in Figure 4. In Figure 4 the plots a t different temperatures in decalin are given, and the first-order behavior of the reaction is confirmed. As the concentration of benzoyl peroxide is much greater than the concentration of Rhodamine 6GX carbinol base, the real velocity constant K’ was obtained as
K’
=
- log e-
Nh
where AF* and AS* are the free energy and entropy of activation, R is the gas constant, K‘ is the rate constant at temperature T , N is Avogadro’s number, h is Planck’s constant, and e is the exponential factor. Table I: Bimolecular Rate Constant a t Different Temperatures in Toluene K‘, M -1
T,“K.
880. -1
318.4 311.4 305.2 299.1 293.1
3.57 2.32 1.46 0.91 0.58
A log [R] - A log O.D. [BIAt PIAL
where [R] is the concentration of Rhodamine 6GX carbinol base and [B] is the concentration of benzoyl peroxide. The rate constant K‘ may be assumed to actually he a bimolecular rate constant for the reaction. Dependence of K‘ on Temperature. The data given in Figure 4 show the variation of log O.D. a t different temperatures in the range 291 to 318°K. in decalin. From the slope of the plot and known initial benzoyl peroxide concentration, the value of K’ was computed and reported in Table I for the reaction in toluene. The data of Table I mere analyzed using the Arrhenius equation W*
In K’ = In A
- RT -L i
and the results are given in Figure 5 for the reaction The Journal of Physical Chemistry
Table 11: Rate Constant, Energy of Activation, and Frequency Factor of the Reaction in Different Solvents E*, kcal.
M -1
Solvent
M-1 sec.-l at 300OK.
mole-1
sec.-l
%-Heptane Cyclohexane Decalin Benzene Toluene
4.04 4.6 2.97 1.42 1.01
6.3 6.6 9.8 10.4 13.2
1.64 X 2.94 X 4.12 X 7.76 X 3.84 x
K’,
A,
lo6 los 10’ 10’ 109
The thermodynamic quantities are entered in Table 111. It is to be noted that the intermediate com(9) A. A. Frost and R. G. Pearson, “Kinetics and Mechanism,” John Wiley and Sons, Ino., New York, N. Y., 1963, p. 40. (10) S. Glasstone, E(. J. Laidler, and H. Eyring, “The Theory of Rate Processes,” McGraw-Hill Book Go., Ino., New York, N. Y.,
1941.
REACTION BETWEEN
13,ENZOYL PEROXIDE AND
4075
RHODAMINE 6GX COLOR BASE
Table I11 : Different Thermodynamic Parameters of Activation for the Reaction oal. deg. -1 mole-1
AS*, oal.
-36.7 -35.5 -25.7 -25.5 -16.3
-24.7 -23.7 -13.9 -13.7 -4.5
A#+,
A H * , koal. mole-1
Solvent
n-Heptane Cyclohexane Decalin Benzene Toluene
5.7 6.0 9.2 9.8 12.6
AF*, kcal.
mole-1
AF*, koal. mole-1
16.7 16.6 16.9 17.3 17.5
20.8 20.7 21.0 21.4 21.6
deg. -1 mole -1
&
9.66 1.76 2.43 2.68 2.71
x x x x x
10-9 10-8 10-6 10-6 10-4
base and oxygen atoms of benzoyl peroxide. The N atom in >NH donates an electron to the 0 atom, resulting in a charge-transfer complex having a characteristic pink color; the reaction may be termed as the acid-base14 reaction, benzoyl peroxide behaving as a Lewis acid and Rhodamine 6GX color base as a Lewis base. The reaction is catalyzed by benzoyl peroxide. The presence of the appreciable dipole moment of >C=O, which probably arises as
0.50
60.25 I3
TJ- :o:oI .. @
-C-
0
:O
7.75 3.1
Figure 5.
3.2
3.3 1/T X 103.
3.4
Linear variation of log K' with 1/T in toluene.
pound (maxima at 515 mp) is more stable in the solvent having double bonds, which implies the possibility of its formation by ionization.ll The differences in the values of the entropy of activation12 in different solvents may be attributed to differences in the degree of orientation of the solvent molecules under the influence of a polar transition state and to the differences in the degree of orientation of the reactants and products. The comparison of the stability of the intermediate compound mdth the values of the entropy of activation listed in Table I11 suggests that the aromatic solvent molecules are capable of further stabilizing the charge separation owing to their high capacity for polarizability of the T electron in the aromatic ring. The activation energy in such ionogenic reactions should be antibatically related to the solvating power of the rnedi&.'O The value of the rate constants, as well as the activa,tion energy (Table 11), leads to the fact that specific effects13 other than the dielectric constant may be more important. The possibility of charge formation is favored by the presence of >NH groups in the Rhodamine 6GX color
implies that the coordination of the electron takes place between central peroxy oxygen atoms and N atoms of the NHCzHsgroups. The charge-transfer interaction most probably takes place between the highest occupied orbital of the donor and the vacant antibonding orbital of the 2pu, around the 0-0 bond in the peroxide. The lowering of the energy level of the antibonding orbital of the 0-0 bond and the elevation of that of the lone pair of N atom facilitate the charge transfer.15 The product of the reaction, which has a maxima at 515 mp, is most probably of inner complex type as has been classified by NIulliken.16 It has been found that such complexes are always unstable. The formation of a charge-transfer complex lowers the activation barrier. The complex formation is the precursor of the chemical reaction, an idea which was earlier set forth by Barckmann.l' A similar case of charge-transfer complex occurs when pyromellitic dianhydride (11) D. P. Aimes and J. E. Willard, J . Am. Chem. SOC., 73, 164 (1961). (12) W. F. K. Wynne-Jones and H. Eyring, J . Chem. Phys., 3, 492 (1935). (13) J. H. Beard and P. H. Plesch, J . Chem. Soc., 3682 (1964). (14) G. N. Lewis, J. Franklin Inst., 226, 293 (1938). (15) K. Tolcumaru and 0. Simamura, Bull. Chem. SOC.Japan, 36, 333 (1963). (16) R. 9. Mulliken, J . Am. Chem. SOC.,74, 818 (1952). (17) W. Barckmann, Rec. trau. chim., 6 8 , 147 (1949).
Volume 60,Number 18'
December 1965
4076
P. K. NANDIAND U. S. NANDI
(PlMDA), which acts as a T acceptor or T acid, reacts with aniline producing a red complex, but the color disappears as amine reacts with PMDA.18 Complex formation is reduced with increasing temperature. The reaction between N,N-dimethylaniline and chloranil produces crystal violet salt by a charge-transfer phenomenon. Diamagnetic donor-acceptor complexes and paramagnatic semiquinones are two observed intermediates.lg From the values of AF* and AS* the terms AF* and AS* can be calculated by means of the equations
- RT In 1000 - R f R In 1000
AF* = AF* AS" = AS*
The terms AF* and AS* are the excess free energy and the excess entropy of activation, respectively. The quantity AF* is the free energy of formation of the intermediate state for reacting particles in excess of what it would be if the state were simply the usual transient collision complex of two neutral nonreacting particles. AS* is the corresponding entropy term.20 The entropy of activation, as well as other parameters, has been calculated assuming the transmission coefficient to be equal to unity. I n the case of an electron-transfer reaction, the best compromise configurations are those giving frequent electronic transitions without too high a free energy of activation. Thus, any measurable rate for an electron-transfer reaction involves a transmission coefficient of less than unity since it is arrived at using an idealized transition configuration.21 The electronic transmission coefficient K , is related to the apparent entropy of activation by the formula AS* = R In K,. The values of K, in the solvents have been listed in Table 111. Isolcinetic Temperature. Figure 6 shows a plot of the activation parameter for the reaction between benzoyl peroxide and the amine. The slope gives a temperature of 335.4"K. (62.3"C.) for the reaction which indicates that the rate of the reaction a t this temperature is independent of the The Validity of the AH*-AS* Plot. Peterson, et aLlZ3have pointed out that, for a reaction series involving a narrow range of AH" values and considerable error in the rate constant, little validity can be assumed in any observed AH *-AS relationship. The larger the range of the AH values and the smaller the experimental error, the greater is the reliability of any observed relationship. If- the range of AH values be designated by dAH and the maximum fractional error in AH* by the symbol 6, Peterson, et ab., have shown that the ratio of dAH*/26 must be equal to or greater than unity if
*
*
The Journal of Physical Chemktry
*
*
6.0
- 40
-30 AS
*,
-20
- 10
0.u.
Figure 6. Least-square plot of enthalpy-entropy of activation for the reaction.
any validity can be assumed in an observed A"AS* plot. This ratio must be much greater than unity if any details of the relationship are to be inferred. The maximum fractional error in AH* can be calculated using the equation 6 = 2RT'(r/(Tf - T ) , where a is the maximum fractional error in the rate constant, R is the gas constant, and T' and T are the upper and lower temperature limits, respectively. The value of a may be assumed to be equal to the reproducibility of the rate constant K'. The range of the AH* value is about 6.8 kcal./mole, and 6 turns out to be about 0.205. The ratio of dAH*/26 therefore turns out to be about 7. This value is much greater than 1, and the validity of the AH*-AS+ plot may be assumed correct.
Acknowledgment. The authors are indebted to Professor Santi R. Palit for constant encouragement during the course of the work. Thanks are due to the Council of Scientific and Industrial Research. for financial assistance to P. K. N. (18) L. I. Ferstandmg, W. G. Toland, and 6.D. Eeaton, J . Am. Chem. Soc., 83, 1151 (1961). (19) J. W. Eastman, G. Englesma, and M. Eelvin, J . Am. Chem. Soc., 84, 1939 (1962). (20) R.A. Marcus, J. Chem. Phys., 26, 868 (1957). (21) R. J. Marcus, B. J. Zwolinski, and H. Eyring, J. Phys. Chem., 58, 432 (1954). (22) S, L.Fries, E. 8. Lewis, and A. Weissberger, Ed., "Techniques of Organic Chemistry," Vol. VIII, Part I, 2nd Ed., Interscience Publishers, Inc., New York, N. Y.,1961,p. 207. (23) R. C.Peterson, J. H. Markgraf, and 8. D. Ross, J . Am. Chena. Soc., 83, 3819 (1961).
