PHILIPL. HANSTAND JACKG. CALVERT
104
Vol. 63
THE THERMAL DECOMPOSITION OF DIMETHYL PEROXIDE : THE OXYGEN-OXYGEN BOND STRENGTH OF DIALKYL PEROXIDES BY PHILIP L. HANSTAND JACK G. CALVERT Contribution from the McPherson ChemicaE Laboratory, The Ohio State University, Columbus 10, Ohio Recsiued August 11, 1968
-
Dimethyl peroxide was decomposed in a static system at various temperatures and the products measured b infrared absor tion spectroscopy. The decomposition is first order and ap roximately follows the stoichiometry 2CH84OCHa 3CH8H CO. The fact that the decomposition proceeds via metRoxy radicals was proved by the addition of nitric oxide which led t o the formation of methyl nitrite as the only major product. The presence of oxygen did not alter the decomposition seriously, although it did lead t o the roduction of some formic acid and formaldehyde. An Arrhenius plot of the data gave an activation energy (CHIO-OCHI[ond strength) of 35.3 i 2.5 kcal./mole which is in good a reement with the values obtained by other investigators for dimethyl peroxide and some of its homologs. Three sets of pubqished data on the thermal decomposition of diethyl peroxide were combined in a composite Arrhenius plot to give an activation energy of 34.1 kcal./mole. This value is in fair agreement with the activation energies of decomposition of homologous dialkyl peroxides, whereas the values obtained separately in the three investigations are anomalously low.
+
Introduction Investigations of the chemical kinetics of the decomposition of a number of alkyl peroxides are reported in the chemical literature. The activation energy is an important one of the factors which have been determined; the mechanisms indicate that it is the minimum energy necessary to cause the reaction ROOR -+ 2R0, and this energy is a qunntity one needs in the calculation of the thermodynamic properties of alkoxy free radicals as well as in the calculation of the strengths of certain other bonds. For example, this activation energy enters importantly into the calculations reported recently by Gray.' Summarized in Table I are some of the abovement,ioned results. The most notable aspect of these data is that the activation energies do not appear to be entirely consistent with each other. That is, the values for diethyl peroxide all geem to be low compared to the others. It does not seem likely to us that a change in the hydrocarbon chain length should have such a large effect on this energy. If the effect were real, however, we feel that it should show a more consistent trend. It seemed to us that there could be two possible explanations for the apparent anomaly; either the diethyl values are all too low, or the value of 36.9 for dimethyl peroxide is too high.
In view of these considerations we undertook to measure the decomposition rate of dimethyl peroxide at various temperatures and thus get another check on the activation energy. In our opinion, the anomaly in the above table is satisfactorily corrected by the results of these experiments plus some reconsideration of the data pertaining to diethyl peroxide. Experimental Technique
Thermal Decomposition of Dimethyl Peroxide.-The decompositions were conducted in a &liter Pyrex flask contained in an electrically heated oven. The roducts of thermal decomposition were analyzed periodical6 by means of infrared absorption spectroscopy using 10 cm. and 1 meter gas absorption cells on Perkin-Elmer Model 21 double beam spectro hotometer. Air was added to the cells before analysis so %at the total pressure waa one atmosphere. Concentrations were determined by comparison to reference spectra run on pure compounds under conditions of resolution, pressure broadening and per cent, absorption as near as possible to those existing during the analysis. A typical experiment was performed as follows. After the oven had come to thermal equilibrium .at a temperature indicated by a precision thermometer, dimethyl peroxide was admitted simultaneously to the reaction flask and to the previously evacuated infrared cell. The timing of the run was then started; the initial pressure was read on a mercury manometer a t room temperature, and the spectrum of the reactant was recorded. After that, the infrared cell was evacuated, and when the desired time had elapsed, the hot gases were allowed to expand out of the reaction vessel and into the infrared cell where the concentration of unreacted dimethyl peroxide and/or reaction products was measured. TABLE I This sampling and analysis process was repeated to get the SUMMARY OF PREVIOUS RESULTSON DECOMPOSITION OF successive points on the curves. It was necessary to correct for the small pressure dro which occurred on expansion by DIALKYL PEROXIDES multiplying each measure3 pressure by the ratio of the presPre-exponential Activation sure at the start of the experiment to the pressure after factor, A energy sampling. The pressure dro was measured for each of the Compd. dec. (sea. -1) (kcal./mole) Ref. infrared cells in separate calil%ation runs using tetramethylDimethyl peroxide 4 . 1 X 1016 36.9 2 ethylene, which did not decompose at the temperatures used, Diethyl peroxide 5 . 1 X 1014 31.5 3 and was found to be about for the 10 em. cell and about Diethyl peroxide 2 . 1 X 1018 31.7 4 23% for the 1 meter cell. I n order to apply the correction successively, it was necessary to assume the first-order rate Diethyl peroxide 1 . 1 X 1013 29.9 5 law. The 1 meter cell was used only in a few experiments Di-n-propyl peroxide 2 . 5 X 10'8 36.5 6 where a single measurement was made. The tem erature Di-t-butyl peroxide 4 x 1014 36 7 usually fluctuated over a range of about A0.2' &ring a Di-t-butyl peroxide 3 . 2 X 10'8 39.1 8 run, and an average was estimated from periodic readings. Preparation of Dimethyl Peroxide.-The dimethyl perI (1) P. Gray, Trans. Faraday SOC..62, 344 (1956). oxide was prepared as follows. Ten cc. (about 0.1 mole) of (2) Y.Takeraki and C. Takeuchi, J . Chem. Phys.. 2 2 , 1527 (1954). dimethyl sulfate and 15 cc. (about 0.14 mole) of 30% aqueous r (3) E. J. Harris and A. C. Egerton, PTOC.ROY.SOC.(London),Alee,
1 (1938). (4) R. E. Rebbert and X.J. Laidler, J . Chcm. Phys., 20, 574 (1952). (6) I(. Moriya, Xav. Phys. Chem. (Japan), Horiba Vol., 143 (1940). , (6) E.J. Harris, PTOC.Roy. SOC.(London), 8118,126 (1939).
1
(7) J. Murawski, J. S. Roberts and M. Szwarc, J . Chem. Phys., 19, 698 (1951).
M(8) J. H. Raley, F. F. Rust and W. E. Vaughan, J . Am. Chcm. Boc., 7 0 , 88 (1948).
OXYGEN-OXYGEN BONDSTRENGTH OF DIALKYL PEROXIDES
Jan., 1959
hydrogen peroxide were placed in a 250-cc. 3-necked flask fitted with a stirrer, a 125-cc. dropping funnel and an outlet tube. From the dropping funnel 20 cc. (about 0.15 mole) of 40% aqueous potassium hydroxide was added slowly with stirring and cooling in ice. The dropping funnel was then replaced with a tube leading to a nitrogen tank, a very weak flow of nitrogen was started, and the outlet tube was attached to a cold tra immersed in liquid nitrogen. Heat then was applied gent$ to the flask, until at about 60" the contents began to bubble smoothly, and the reaction then maintained itself. At times, some moderating with the ice-bath was necessary. The dimethyl peroxide vapors were carried in the nitrogen stream into the cold trap. When bubbling in the reaction flask had ceased, the trap containing the product was removed to the vacuum line where the volatile dimethyl peroxide was distilled into a storage bulb. This distillation served to se mate the peroxide from any HzO, Hz02, etc., which may fave been carried over during the preparation. The infrared spectrum of the final product showed no detectable amounts of possible contaminants. The peroxide was stored as a vapor in a clear bulb a t room temperature and showed no appreciable decomposition over periods as long as one week.
Discussion From the data of Fig. 1 it is seen that dimethyl peroxide decomposition is first order in the temperature range used. Takezaki and Takeuchi also found this t o be the case.2 The only major products were methanol and carbon monoxide. Table I1 shows that although the measurements were somewhat erratic, the products appeared to be in the mole ratio of approximately 3 to 1 in agreement with the over-all reaction 2CHsOOCHs +3CHsOH
CH300CHs
CHIO CHI0 CHaO
+2CH30
+ CH3O +CH30H + HzCO + H2CO +CHsOH + HCO + HCO +CH3OH + CO
(1) (2) (3) (4)
The possibility of chain decomposition of the dimethyl peroxide through the repeated occurrence of the cycle (5) and (6) CHIO
0.8
. 0.6
2 a
0.4
0.2
0.0 0
120 180 Time, min. Fig. 1.-Plots of integrated first-order rate law for the thermal decomposition of dimethyl peroxide.
+ CHIOOCH +CH30H + CHzOOCHs CHzOOCH, +HzCO + CH30
(5)
(6)
can be eliminated. This follows from the facts that formaldehyde does not build up as a final product of the reaction and the stoichiometry follows closely the relation 2CHaOOCHs +CO + 3CHSOH It appears that methoxy radicals react with the product formaldehyde in the reaction 3 to the practical exclusion of (5). I n other words, the possible chain decomposition of dimethyl peroxide is selfinhibited by the formaldehyde product. To test the mechanism 1 t o 4,dimethyl peroxide was decomposed in the presence of nitric oxide with the results that methanol and CO formation was suppressed, and the only major product was
60
-3.2
-3.6
+ CO
The lack of reproducibility of the measured mole ratios is believed to be due to the fact that the CO band is a very weak one and the spectrometer background noise was somewhat high. Although it was difficult to measure the amount of decomposed dimethyl peroxide accurately because of interference from the methanol, it appeared that the amount used up could be approximately accounted for in the observed products. The products are assumed to form in the reactions
105
4
,-4.0 0
c
-4.4
-4.8
2.4
2.5 2.6 1/T x 103. Fig, 2.--Arrhenius plot for rate constants of the thermal decomposition of dimethyl peroxide.
methyl nitrite. When it was assumed that the measured amount of methyl nitrite formed was twice the dimethyl peroxide decomposed, a rate constant was obtained which was in good agreement with the rate constants calculated directly from dimethyl peroxide disappearance and from methanol formation. The open circle on the activation energy curve in Fig. 2 was obtained from this experiment with nitric oxide. The decomposition also was conducted in the presence of oxygen, although no rate constant could be measured in this case. I n the presence of the oxygen, the methanol and carbon monoxide still formed to a large extent, but in addition there appeared a small amount of formic acid and apparently a fairly good yield of formaldehyde as indicated by its infrared spectrum and by the formation of a white solid deposit on expansion of the hot gases into the infrared cell. Because of this deposit, no quantitative measurements of products were attempted. It appears likely that the large
PHILIP L. HANSTAND JACKG. CALVERT
106
Vol. 63
TABLE I1 PRODUCTS OF THXRMAL DECOMPOSITION OF DIMETHYL PEROXIDE Starting CHsOOCHa (mm. H d
Temp. ("C.)
Elapsed time (min.)
Pressure of CHaOH (mm.)
Pressure of co bm.)
PcHaoH
130 61 22 15 45 20
1.85 12.6 4.6 1.99 3.14 11.4 20 t o 40
0.54 5.0 1.8 0.6 1.3 3.9 10
3.4 2.5 2.6 3.3 2.4 2.9 2 to 4
119.7 127.2 135.2
7.25 30.0 14.3 3.16 same run 3.16 162 11.5 167 (Takezaki and Takeuchi2) 23
}
0.0 1 \
I
I
'I
I
...
-0.5
\
-1.0
0
-1.5 4
-
M
0
-2.0
Pco
as 10% from this stoichiometry would not affect the calculated rate constants significantly. The logarithms of the rate constants obtained from these curves (slope = IC) are plotted against 1/T in Fig. 2. The slope of this curve, calculated by the least squares method, leads to an activation energy of 35.3 =t 2.5 kcal./mole, where the uncertainty is twice the statistical standard deviation and represents 9501, confidence limits. This value is in satisfactory agreement with the 36.9 =k 1.1 kcal./mole obtained by Takezaki and Takeuchi2 The pre-exponential factor calculated from the intercept by least squares is 1.6 X 1015 set.-'. The open circle obtained in the experiment with added nitric oxide was not used in these calculations. This agreement with the results of Takezaki and Takeuchi emphasized the anomaly of the low values for diethyl peroxide shown in Table I. I n view of this, we have checked the results of the three different investigations on diethyl peroxide against each other by making the combined Arrhenius plot shown in Fig. 3. This treatment has the advantage of a wider temperature range than a plot from one of the sets of data alone. The results of the different studies appear to be consistent. The least squares method of calculation applied to this curve gives an activation energy of 34.1 kcal./mole and a frequency factor of 1.6 X 1014sec.'l. I n summary, we suggest the following values as the most reliable estimates of the pre-exponential factors (sec.-l) and activation energies (kcal./mole), respectively, for the four peroxides discussed : dimethyl peroxide, 2.4 x 1016,36.1 (average of our values and those of Takezaki and TakeuchP); diethyl peroxide, 1.6 X lo1*,34.1 (values taken from the combined Arrhenius plot, Fig. 33,435); di-n-propyl peroxide, 2.5 X 1016, 36.5 (Harriss); di-t-butyl peroxide, 4 X 10'5, 37.5 (average of values from Murawski, et aE.,' and Raley, el aZ.* Acknowledgment.-The authors gratefully acknowledge the financial support received from the United States Public Health Service, National Institutes of Health, Bethesda, Md., and the helpful communication with Dr. M. H. J. Wijnen.
-3'0 -2.5
-3.5
2.3
2
1.9
2.1 I / T x 103. Fig. 3.-Composite Arrhenius plot for rate constants of the thermal decomposition of diethyl peroxide.
amount of formaldehyde in the experiments with added oxygen has its origin in reaction 7 CHsO
+
0 2
+HOz
+ HzCO
(7)
Figure 1 shows the integrated first-order rate law plots for the decomposition at the various temperatures. The measurements at 139.9' (square points) were made on the disappearance of dimethyl peroxide using its infrared band a t 8.7 p ; all others were made on the rate of formation of methanol using the OH band a t 2.7 p and assuming the stoichiometry 2CH300CH3 + 3CH30H CO. A sample calculation showed that a deviation of as much
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Jan., 1959
CHANGE POTENTIAL ON SILVER ELECTRODE IN ALXALINE SOLUTION
107
THE EFFECT OF CONTINUOUSLY CHANGING POTENTIAL ON THE SILVER ELECTRODE rN ALKALINE SOLUTIONS BY THEDFORD P. DIRKSEAND DALEB. DE VRIES Department of Chemistry, Calvin College, Grand Rapids, Mich.igan Recebed Auoust IS, 1968
When a silver wire is oxidized by continuously increasing its otential in alkaline solutions, four reactions take lace. others are due to the formation of Ag20, of Ag8 and of oxygen. The first and second reactions are indistinguishable in solutions having a high p H . The results show that the formation of Ag20 is not due to the action of oxygen on silver, and the formation of Ago does not involve the HOz- ion.
It is suggested that the first is due to the formation of AgOH. !he
Introduction When silver is electrolytically oxidized in alkaline solutions, AgzO is formed, then, a t a higher voltage level Ago is produced. As the voltage rises still higher oxygen is evolved. There is doubt whether Agz03is formed in such a treatment. Most of the work that has been carried out in an attempt to understand the mechanism of these processes has been done by the use of a constant current technique. l--4 The work reported here was carried out in an effort t o shed more light on these reactions. This was done by applying a continuously increasing potential to the silver electrode. So far as could be determined, no such work on these reactions has been reported in the literature. Such results cannot be subjected to strict quantitative interpretations because of a variety of variables present, e.g., convection and changing electrode surface area. However, significant qualitative conclusions can be drawn from such data. Experimental
measurements. The evolution of oxygen was observed visually. For each reaction, the rise in current indicates an increase in the rate of the reaction and a decrease in the current represents a decrease in the rate of the electrode reaction. For reactions a and b the decrease in rate of reaction is rather gradual and this may correspond to the gradual shutting off of the reaction due to the formation of a film of the reaction product. With reaction c the decrease in rate of reaction was very sudden. This was accompanied also by the formation of a visible film of gas (oxygen) on the surface of the electrode. As soon as this film was broken and bubbles of gas were evolved from the electrode surface, the current rose again, peak d. I n some of the runs a vigorous stream of air was passed through the solution during passage of current. However, in general, the nature of the curves was unaffected by such agitation. The effect of dissolved oxygen was observed by carrying out a run after purified nitrogen had been passed through the electrolyte for some time. Then a similar run was made after air had been bubbled through the electrolyte. I t appeared that none of these phenomena was affected by the presence or absence of dissolved oxygen, see Table I. There is a little variation in the potential at which oxygen is evolved but these potentials are a t best hardly reproducible in a given system. Thus it appears that none of the reactions noted involves dissolved oxygen. This indicates, e.g., that in the formation of Ag20 it is not the action of dissolved or liberated oxygen on the silver that is responsible.
The circuit employed in this work was similar to that in polarography. A Sargent Model 111 polarograph was used as the voltage control. It was equipped with a 0.5 r.p.m. motor to increase the voltage continuously and uniformly. The current, instead of passing through the galvanometer, was fed into a precision resistor and the voltage drop across this resistor was measured by a Brown recording potentiometer. Only relative current values were necessary in this work. An H type cell contained the electrodes. In one branch there was a saturated calomel electrode with a plug of agar and KNOa solution in the cross piece. The other electrode was a silver wire about 2 cm. long fitted in a polystyrene holder. This electrode was the anode in all theoruns. Measuraments were made a t room temperature and 2 . I n addition to temperature, the other variables studied were: agitation, potassium hydroxide concentration and dissolved TABLE I oxygen. Since rather high concentrations of potassium hyEFFECT OF DISSOLVED OXYGEN ON POTENTIALS FOR OXIDAdroxide were used, corrections were made for junction poTION OF SILVER IN KOH SOLUTIONS AT ROOM TEMPERATURE tentials.5
Results On Fig. 1 are shown results that are typical of all those that were obtained. One of the first things noted was the presence of another oxidative stage in addition to that of the formation of Ag20, Ago and oxygen. This is indicated by peak a on Fig. 1. The formation of AgtO takes place a t peak b, of Ago at c, of oxygen at d. The presence of Ag20 and Ago at the peaks indicated was confirmed by e.m.f. (1) R. Luther and F. Pokorny, 2. anarg. allgem. Chem., 51, 290 ( 1908). (2) I. A. Denison, Trans. Electrochem. Soc., 90, 387 (1946). (3) Hiokling and D . Taylor, Disc. Faraday Sac., No. 1, 277
A.
(1947). (4) P. Jones, H. Thirsk and W. F. K. WynneJones, Trans. Faraday Soe., 62, 1003 (19513). (5) T.P. Dirkse, 2. p h y s i k . Chem., N . F . , 6 , 1 (1955).
mKOH
Soln. satd. with
1.0 1.0 1.6 1.6 4.9 4.9 13.2 13.2
Air Nitrogen Air Nitrogen Air Nitrogen Air Nitrogen
-Minimum potential for---FormaFormaFormaFormation of tion of tion of tion of AgOH Agio Ago OP
0.08 .08 .12 .12 .09 .08
...
...
0.16 .16 .18 .18 .15 .15 .16 .15
0.43 -43 .46 .46 .41 .41 .42 .42
0.68 .66 .69 .68 .61 .65 .57 .60
The effect of KOH concentration is shown on Fig. 2. The point of interest here is that the lines representing points a’, b’ and c’ on Fig. 1 are parallel. Except at the highest concentration these lines are also approximately parallel to that €or oxy-
