The Decomposition of Alkyl Peroxides: Dipropyl Peroxide, Ethyl Hydrogen Peroxide and Propyl Hydrogen Peroxide

The Decomposition of Alkyl Peroxides: Dipropyl Peroxide, Ethyl Hydrogen Peroxide andPropyl Hydrogen PeroxideAuthor(s): E. J. HarrisSource: Proceedings of the Royal Society of London. Series A, Mathematical and PhysicalSciences, Vol. 173, No. 952 (Nov. 10, 1939), pp. 126-146Published by: The Royal SocietyStable URL: http://www.jstor.org/stable/97442 .

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The decomposition of alkyl peroxides: dipropyl peroxide, ethyl hydrogen peroxide and propyl

hydrogen peroxide

BY E. J. HARRIS

Beit Scientific Research Fellow, Department of Chemical Technology, Imperial College, London

(Communicated by A. C. Egerton, Sec. R.S.-Received 12 July 1939)

INTRODUCTION

In a previous paper (Harris and Egerton I938) the results of an investigation of the decomposition of diethyl peroxide have been published. The present communication deals with the related compounds (a) dipropyl peroxide, which closely resembles the diethyl compound as it decomposes unimolecularly, (b) ethyl hydrogen peroxide, and (c) propyl hydrogen peroxide. The two latter decompose heterogeneously at lower temperatures, while at higher temperatures they decompose luminously in absence of added oxygen.

Dipropyl peroxide does not appear to have been mentioned in the liter- ature; it was isolated from the products of a preparation of propyl hydrogen peroxide. The latter was prepared by Medvedeef and Alexeeva (I932), who found that the substance did not liberate the theoretical quantity of iodine from potassium iodide. In solution it decomposed to propaldehyde.

Ethyl hydrogen peroxide or homologues may be of importance in slow combustion reactions, and is said to have been identified in combination with formaldehyde among the products of slow oxidation of octane (Mon- dain-Monval I 932).

The difficulty in studying the peroxides lies in the poor yields obtained in the preparations. The strongly alkaline reaction mixture required to bring about the alkylation rapidly decomposes the alkyl hydrogen peroxide as it is formed, so good yields by this method of preparation are unlikely.

DIPROPYL PEROXIDE

Preparation of the propyl peroxides

The reaction mixture for the preparation is that calculated for the formation of the propyl hydrogen compound C3H70OH; changing the ratio of the

[ 126 ]



The decomposition of alkyl peroxides 127

alkyl sulphate to hydrogen peroxide used has little effect on the yields because, at best, only some 10 % of the sulphate reacts before all the hydrogen peroxide has decomposed.

A mixture of 120 c.c. 100 volume hydrogen peroxide with a solution of 56 g. potassium hydroxide in 60 c.c. water was prepared, the hydrogen peroxide having been previously cooled to -20? C. The liquid was then cooled in ice and 184 g. (or more) dipropyl sulphate run in with stirring. The stirring was continued for 8 hr., after which the liquid was separated and the lower aqueous part was made just acid to methyl red. The upper layer was used again with a fresh lot of alkaline peroxide and the process repeated for five to six runs.

The combined aqueous parts were distilled in vacuo to half bulk, and the distillate was extracted exhaustively with ether after salting out with ammonium sulphate. The ether extract was dried over sodium sulphate and the ether removed under slightly reduced pressure. The residual liquid was dried over anhydrous copper sulphate and distilled under 2 cm. pressure from a water bath. The fraction boiling at 350 C was collected separately.

The yield of propyl hydrogen peroxide was about 1 g. per run. It was stored over anhydrous copper sulphate and redistilled before use.

The dipropyl peroxide which was present with unchanged sulphate in the non-aqueous part of the reaction product was obtained by distillation of the mixture at 2 cm. pressure from an oil bath. The bath temperature was taken to 1100 C. The distillate was redistilled under reduced pressure, the liquid boiling at 51-53? C at 8 cm. being separately collected. The yield of dipropyl peroxide was about I g. per run.

Properties of propyl hydrogen peroxide, C3H700H.

Iodine liberation (methol described in text): 77 8 00 theoretical. Ele- mentary analysis: C 46-8 %0, H 112 %. (Calc. for 98-5 %? C3H7001, 1-5 % H20: C 46 8 00, H 10-8 %.) Density: 240 0 9040. Ref. index [NaD]: 1-3890 at 25 0? C. Becomes glassy at - 900 C. Boiling point: 350 at 2 cm., 38? at 3 cm.

Properties of dipropyl peroxide, C3H700C31.

Elementary analysis: C 61 3 00, H 11 5 %0. (Theory: C 60 9 %, H 11v9 9.)

Density: 20? 0-8254. Ref. index [NaD]: 1-3911 at 20 5? C, 1-3940 at 16? C.

Decomposition kinetics of dipropyl peroxide

The decomposition was carried out in an all-glass apparatus similar to that shown in figure 2 of the previous paper (I938), but the pointer of the spoon



128 E. J. Harris

was made to tilt a small mirror and, with the use of this, adequate sensitivity was attained.

The initial pressures and the number of experiments are given below, with the temperatures used.

Limits of initial Temp. pressures No. of

? C cm. Hg experiments 175-4 2 -1 8 170*2 2.5-0*7 7 166.8 26-0.4 8 155*3 2 -02 8 146.5 3 -07 6

The pressure time relation followed the unimolecular law, the plot of

log10 P against time giving a straight line. The constants k=/t log10 P )

were found and are tabulated in table 1 below. The activation energy found from the slope of the line, log k against the reciprocal of the absolute temperature, was 36 5 kcal. per g. mol., which is slightly higher than the value for the diethyl compound. (31 6kcal. per mole.)

TABLE 1. VELOCITY CONSTANTS OF DIPROPYL PEROXIDE

Temp. ? C k min.-l

146-5 0-015 155.3 0036 166.8 0-117 170*2 0141 175.4 0246

The final pressure was about 2-50 times the initial pressure, and x was calculated by dividing the pressure increment by 1P5. The final pressure after the end of reaction was determined at 155? and at 175 4'. It was also checked in a packed vessel and did not deviate from 2-5 times the initial pressure.

The constant k was also calculated in a way which is independent of the value of the ratio of the pressure after complete reaction to the initial pressure and of the value of the initial pressure. The pressure increment after time intervals t2 - tl was divided by the increment over the same time intervals starting from some later time t3, and the log of the quotient was finally divided by the time difference t3 - tl. The values of k so obtained were in satisfactory agreement with those obtained graphically. For example, at 175.40 C, the following results were obtained for a particular reaction. The pressure units are, 76 = 1 cm. Hg.




Pressure of

peroxide Pressure k Time and of I-log1 P by method t2 X min. products peroxide t p_x above min. min.

0 15.0 150 -

05 205 11-34 0*244 0*250 0 1 1.0 24.7 8*5 0-248 0-242 0 1 2*0 30*15 4*9 0-243 0242 0 5 1.5 2-5 31-9 3.75 0.241 0245 1 2 3*0 33-25 2-85 0*241 0-255 1 2 4 0 35*2 1-58 0-242 0219 2 4 5.0 36*4 0*751 0*260 - - -

7-5 37-6 Ratio 2*5

The results of experiments in a packed vessel at 146 5 and 155.30 were the same as those from unpacked vessels, so it may be concluded that the reaction is homogeneous as well as unimolecular.

Explosion and ignition of dipropyl peroxide

By analogy with the diethyl compound, a luminous, explosive decomposition was expected to take place above certain limits of temperature and pressure. Trials, however, showed that the decomposition, although it could become explosive, did not give rise to visible light. Accordingly the limiting pressure at which the transition to explosive decomposition took place had to be fixed by observation of the pressure gauge. For this purpose an aneroid gauge of low inertia was employed. At 199? C the limit was between 0-69 and 0-67 cm., and at 210? C between 0-31 and 0 30 cm., the vessel diameter being 4 2 cm.

In presence of air there is emission of light when the peroxide is admitted to the vessel. The minimum amount of peroxide to give rise to a visible flash was very small as in the case of the other peroxides. The minimum pressure at which the flash was seen was 0 007 cm. with 3-5 cm. air, and 0 01 cm. with 35 cm. air at 2710 C, and at 2t1) C it was 0 058 cm. with 1 1 cm. air. The colour of the flash was blue. The pressure of peroxide was measured on a manometer containing amyl phthalate attached to the reservoir vessel (A in figure 1 of previous paper). The low-pressure end of the manometer was permanently connected to the vacuum line.

These experiments indicate that the dipropyl compound explodes at slightly lower pressures than diethyl peroxide under comparable conditions.

Vol. I73. A. 9



130 E. J. Harris

C>D Ce Ss

1'0X

-10?

-1-4-

0-7

o -18

05~~~~~~~~~~~~~~~~~~

04

03

0-2

0ii _x

0 1 2 3 4 65 time (min.)

FIGURE 1. The unimolecular decomposition of dipropyl peroxide. Inset: the line log1lok (ordinates) vs. I/T (abscissae).




Decomposition products of dipropyl peroxide The products of the unimolecular decomposition are not only those to be

expected from the reaction C3H700C3H7= C3H70H + C2H5CHO, which would lead to a pressure increase of only 100 %, the value found being 150 %.

By withdrawing from the decomposition apparatus the reacted material, it was possible to obtain a small sample of gas, which was found to be approximately 50 % carbon monoxide and the remainder hydrocarbon

(CnH2n+2) with perhaps some hydrogen having a value of n between 3 and 4. In order to collect more product, the peroxide was decomposed by re- fluxing in a tube, the upper end of which was heated to 1800 C. The analysis of the gaseous products was C02 5 %, CO 49 %, H2 6-1 %, hydrocarbon 3665 0/ (n - 3.41). The liquid part which was nearly equal in weight to the original peroxide was treated with dinitro-phenylhydrazine solution. After recrystallization, the aldehyde derivative melted at the same temperature as the propaldehyde derivative, and did not depress its melting point. This proves the presence of propaldehyde. Distillation of the filtrate from the precipitation gave a solution of some compound which could be oxidized by weakly acid dichromate, presumably propanol, but a derivative could not be obtained. After oxidation of part of the original liquid with alkaline hydrogen peroxide, it precipitated mercurous chloride from boiling mercuric chloride, indicating the presence of formaldehyde in the original.

Hence it appears that the decomposition furnishes the products of decomposition of the aldehyde besides propaldehyde and propanol, or alter- natively, the reaction characteristic of the explosion takes place to a small and constant extent because of the constancy of the pressure increment. The latter is more likely in view of the high value of n.

The explosive decomposition products were obtained by dropping during a period of about 1 min. the liquid peroxide from a tap funnel into a tube in a furnace at 2400 C. The lower end of the tube passed into a cooled trap at - 500 and any gas passed on through a water scrubber to a gas burette. The results of such an experiment are given below.

Weight used, 0.294g. Condensed, 0.199g. Gas, 10 9 c.c. non-condensed, and 36 7 c.c. condensed.

The boiling point of the condensed gas was 10 C, corresponding to that of butane. Analysis of the fractions showed that the non-condensed gas was CO 26-5 % and hydrocarbon (n = 2.1) 73-5 %, the condensed,fraction was CO 3 00, hydrocarbon (n = 4) 97 %. The yields in moles x 10-4 were HCOOH 3-5, HCHO 40, C4H1 16, CO 1-8, alcohol 4, C2H6 3-6. Most of the

9-2



132 E. J. Harris

formaldehyde was present as paraform. It is possible that a little propaldehyde was also present.

The two most important products of the explosive decomposition are clearly formaldehyde and butane, the reaction

CH3CH2CH20 CH3CH2 CH20 I= I +

CH3CH2CH2O CH3CH2 CH2O

being analogous to the one by which diethyl peroxide forms ethane and formaldehyde.

ETHYL HYDROGEN PEROXIDE AND PROPYL HYDROGEN PEROXIDE

Preparation of ethyl hydrogen peroxide The preparation of ethyl hydrogen peroxide is effected by ethylation of

hydrogen peroxide (Baeyer and Villiger I 90 I). As the procedure adopted by these authors has been modified, the method employed is briefly described.

A mixture of 120 c.c. of 100 volume hydrogen peroxide with a solution of 56g. potassium hydroxide in 60 c.c. water was prepared, the hydrogen peroxide having been previously cooled to - 200 C. The mixture was then cooled in an ice bath, and was stirred mechanically while 154g. diethyl sulphate was added in 10 c.c. lots at intervals of about 15 min. The stirring was continued for some 20 hr., at the end of which time the aqueous liquid was nearly neutral. The aqueous layer was separated off, made just acid to methyl orange and transferred to a distillation flask. About a third of the volume was distilled over at 2 cm. pressure from an oil bath. The distillate was ether extracted and the extract dried over sodium sulphate. The ether was removed at about 15 cm. pressure, and used to re-extract the aqueous solution. This was repeated until no more peroxide could be extracted.

After standing over anhydrous copper sulphate the peroxide, still containing some ether, was fractionally distilled at 5 cm. pressure from a water bath. The fraction boiling at 41-44? C was collected separately. Finally the peroxide was kept over anhydrous copper sulphate and distilled when required. The best yield obtained was IO g.

Properties.

Iodine liberation (see text): 90 00 theory. Density: 20? 0-9332. Ref. index [NaD]: 1-3800 at 20*50 C. Becomes glassy at - 100? C. Analysis: C 40-1 %, H 9-7 %. (Calc. C 38-7 %/ H 9-75. A mixture of 90 0 peroxide with 10 %/ acetaldehyde would have C 405 00, H 9-7 0g.)




Decomposition of ethyl hydrogen peroxide The preliminary experiments on the decomposition were carried out in

Pyrex tubes (volume 16 c.c.) into which about 0-06 g. of the peroxide was weighed. The tubes were cooled to - 700 C, evacuated and sealed off. They were then heated for various times either in water vapour or toluene vapour. After cooling the pip was broken off under water. The resulting solution was analysed for peroxide by the method described later.

Owing to the lag in attaining the bath temperature, the amount decomposed is too low in each case. These results were compared by the use of the unimolecular equation, for although later experiments show that the decomposition is heterogeneous, the terms not in accord with the unimolar expression are nearly equal and opposite, provided the initial pressures are not very different.

TABLE 2

Equiv. Equiv. Permanent Temp. peroxide peroxide k gas c.c.

o C Conditions used x 1O-4 after x 1O-4 min.-' (N.T.P.) 100 3 min. heating 9 2 7-55 00286 0-2

4min. heating 12-96 9 9 0-0294 0 3 ,, ll 5 min. heating 146 7.1 0.0274 0 3

6min. heating with 10-2 1*9 0120 3 0 KCI

138 1 min. heating 13.7 8.5 0.216 0.2 1 min. heating with 14-1 8-46 0.222 iP3 SiO2 gel.

1 min. heating with 15.2 1.2 1.036 2.5 KCI

I min. heating with 1141 2.7 0-879 2.5 NaCi

The amount of alkali chloride used was about 0a I g. The increase in total surface by this addition would be small and less than in the experiment in which silica gel was added. Hence it is evident that the nature of the surface plays an important part in determining the reaction rate.

Decomposition kinetics of ethyl and propyl hydrogen peroxides As the peroxides react with mercury, causing it to tail, an all-glass appa-

ratus with glass Bourdon gauge was used to study their decomposition. It was similar to that used for the dipropyl peroxide experiments. The upper pressure attainable when propyl hydrogen peroxide was the subject of investigation was rather low owing to its low vapour pressure and the con- sequent proneness to condensation. The reaction products also tended to



4034 3 //> J-S ] > It- 28

4-236 __ _ __ __ __ _0__ _ _

0 30 60 90 120 150 180 0 30 60 90 120 158 180 210 240 time (sec.) time (sec.)

FIGURE 2 a FIGU6E 2b

FIc;uRE 2 a. Pressure-time curves for ethyl hydrogen peroxide decomposition: (a) empty vessel at 166-50 C; (b) empty vessel salt washed at 1560 C; (c) packed vessel at 156? C; (d) empty vessel at 180? C; (e) empty vessel salt washed at 180? C; (f) packed vessel at 1800 C. FIGURE 2b. Pressure-timne curves at various initial pressures for propyl hydrogen peroxide decomposition: (a) empty vessel at 166.50 C; (b) packed vessel at 166.50 C; (c) empty vessel at 166.50 C; (d) empty vessel with active surface at 166.50 C.




condense if the initial pressure was higher than about 3 cm. The corresponding pressure for the ethyl hydrogen peroxide was about 5 cm.

In figures 2 a and 2 b some specimen pressure-time curves are shown for the two peroxides in empty and packed vessels, and the effect of salt washing and of a trace of impurity (most likely mercury) can be seen. The vessels were of Pyrex and the packing consisted of a number of lengths of quill tubing.

The fact that the rate is greater in vessels with increased or modified surface indicates that the reaction is heterogeneous, so an expression for the

Pt

1-055 111 22

4

PL-1O5 1.1

P~~~~~P

0 30 60 90 0 30 60 90 time (sec.) time (sec.)

FIGURE 3 a. Decomposition of ethyl hydrogen FiGURE 3 b. In empty vessel at peroxide, in empty vessel at 180450 C. 166.50 C.



136 FE. J. Harris

rate was derived as follows: if the adsorption coefficient cx determines the

rate, then applying the Langmuir isotherm, a = --1 k (p - x) , where I1?k2(P -X)?+k2x' hr

p - x is the pressure of undecomposed peroxide at time t, and p is the initial pressure. Equating this to dx/dt and integrating the equation

p lo P (k - k2)x Xpk2 lo k ogp + k + k logp-x

is obtained.

_ _ _ _ _ Pi/ , 1.5

2 1.0~~~~~~~~~~~~~~~~

i 2 + X g 40 10 12

0 S / er ; / | g f / / T0-5

0~~~~~~~~~~~~~~~~

0 30 60 90 0 30 60 90 time (sec.) time (sec.)

FIGURE 4a. Decomposition of propyl hydrogen FICiURE 4b. In packed vessel at

peroxide, in empty vessel at 166-5? C. 166.50 C.

In order to apply this equation it is convenient to choose a value of x/p (_ say n) in which case it reduces to

t = 1log 1 + (k1 -k2)nP +pk2log 1

tko1+n k -+k Ign,




which isa straight line when onlyp and t are variable. In figures 4-6 are shown the experimental values of t for different initial pressures p, and for values of n up to as much as 0O8. The value of x was calculated by multiplying the pressure increment by two, for at the end of reaction the pressure increase

4

3

b=- i2 1. 1|

4i ~ ~ ~ ~ ~ ~ 3P P

0 30 60 90 120 150 180 time (sec.)

FIGURE 5. Decomposition of propyl hydrogen peroxide in empty vessel at 166.50 C, having active surface.

was 45 % in the case of the ethyl compound (which was 90 % pure) and was nearly 50 % in the case of the propyl compound. Thus a value of

pressure at time t_ 1 10 initial pressure

is equivalent to a value of n = 0-2 or 20 % decomposed. The constants k, etc., are calculable from the slopes and intercepts of the

lines, but as the factor np varies at nearly the same rate as p log1o 1 -



138 E. J. Harris

not possible to obtain accurate solutions of the simultaneous equations for k1 and k2. In table 3 some values of the constants are given for the various experimental conditions. It appears that the two peroxides behave very similarly. It is noteworthy that the acceleration in the vessels with active surface is due more to a lessening of the protective effect of the peroxide and its decomposition products as shown by low k1 and k2, than to an increase in k. The value of k is the product of the rate of the surface reaction and the adsorption factor, so its comparative constancy in this case may be due to decreased adsorption coupled with increased reaction rate.

TABLE 3

Units: Pressures in cm. logs to base 10. Time in sec. Temp.

Substance 0 C Conditions k ki k2

C2H500H 156 Packed vessel 0-04 1-6 4 166.5 Packed vessel 0.055 4-5 7 180 Packed vessel 0067 2-8 3.8 166-5 Empty vessel, active surface 00092 0 45 0 180 Empty vessel, clean surface 0 011 1.5 2.6

C3H700H 166.5 Packed vessel 0O05 3.8 6.4 166.5 Empty vessel, active surface 0.0113 0.52 0.22 166.5 Empty vessel, clean surface 0-008 3-6 4-8

The agreement of the experimental data with the equation used and the acceleration of the decomposition by surface show that the ethyl and propyl hydrogen peroxides decompose heterogeneously.

Luminous decomposition of ethyl and propyl peroxides

In the course of the preliminary experiments on ethyl hydrogen peroxide, it was found that if the tubes containing a little peroxide were heated rapidly to about 150? C, an explosive reaction took place with emission of light, and some of the tubes were shattered. On breaking off the pip of those which had remained intact under water, it was possible to collect a sample of gas. The high pressure of gas in such cases (2 atm. corresponding to 30 c.c. at 1 atm. when the tubes were at room temperature) is in marked contrast to the small volume formed as a result of slow decomposition (vide table 2). By in- jecting approx. 5 mg. of the liquid peroxide through a fine capillary at 160?, explosion occurred before it reached the attached reaction vessel.

Luminous decomposition of the vapour could be observed, using the apparatus shown in figure 1 of the previous paper (I938). The low pressures of the peroxide requisite to give rise to a visible glow were measured, as in the case of dipropyl peroxide, by means of an amyl phthalate manometer. On




account of the feeble luminosity it was necessary to work in the dark and to rest the eye before making an observation. The results are only quoted to show that exceedingly small pressures can suffice, and may be determined by the limiting sensitivity of the eye. This applies particularly to those in table 5, and are not necessarily the lowest at which the ignition will take place. The vessel diameter was 4-2 cm.

TABLE 4

Explosion pressures of ethyl hydrogen peroxide (mm. Hg)

Temp. Flash not ? C Flash seen at seen at

195 8.9 8.6 201 2*3 2.1 230.5 0*92 0*89 241*5 0*80 0.80 252*5 0*56 0*50

Explosion pressure of propyl hydrogen peroxide 233 1*5 1*4

The light emitted was of a sky-blue colour, and seemed to persist longer than the luminescence associated with the explosion of diethyl peroxide.

The pressures at which ignition could be observed in the same vessel were as follows:

TABLE 5

Ignition pressures of ethyl hydrogen peroxide (mm. Hg). Pressure of air added 11 cm. Temp. Flash not

O C Flash seen at seen at 201 0-18 0-13 230*5 0-18 0-13 241-5 0.19 0*06 268 0*10 0-04

Ignition pressure of propyl hydrogen peroxide 233 0*175

The luminosity was not observed in a I1 - cm. diameter vessel at 237? C at pressures up to 9 5 mm. of propyl hydrogen peroxide.

The pressures for ignition at a given temperature are always lower than those for critical explosion; both, however, for propyl hydrogen peroxide appear to be slightly higher than for ethyl hydrogen peroxide, so far as these experiments indicate, which is contrary to the result for the dialkyl compounds. Both the explosion pressure and ignition pressure of the dialkyl are lower than those of the alkyl hydrogen peroxides, but the explosion of the



140 E. J. Harris

latter may not be the same type of process, for it may be -an ignition due to oxygen produced in the decomposition as is shown by the analyses (see table 6).

Decomposition products of ethyl hydrogen peroxide

It has already been mentioned that the slow decomposition did not give rise to important quantities of gas, whereas the explosion formed much gas. In order to obtain the slow decomposition products, some of the peroxide was distilled through a tube either containing glass wool (at lower temperatures) or empty (higher temperatures). The liquid products were condensed in a cooled trap, and the gas was drawn off by means of a Toepler pump arrangement. The presence of unreacted peroxide and the formation of paraldehyde rendered analysis difficult, and the small quantity of material available rendered identification of the various products uncertain.

The most important products were acetaldehyde (mostly polymerized), formaldehyde and alcohol. The latter was not methanol, so it may reasonably be assumed to be ethanol. The isolation of oxygen at the lower temperatures shows that such reactions as

2C2H500H--2C2H50H +02

and/or 2C2H500H--2CH4 +?2 + 2CH20

may be taken as primary processes, followed by oxidation of the acetaldehyde formed by the reaction C2H500H--CH3CH0 + H20. It seems probable that the hydrogen arises from the decomposition of the intermediately formed C2H500CH20H, because a trial with this compound showed a greatly increased hydrogen yield.

TABLE 6

Decomposition products of ethyl hydrogen peroxide Units, moles x 10-4 from 10 moles x 104 used

Temp. Unde- 0 C composed C02 CO 02 CnH2n+2 n H2 HCHO Ald. Alc. 170 P ? 0.062 0-086 0.633 0-298 1.18 - 1*13 (1-82) (4.86) 185 P 0-42 0 50 0 49 0-31 0-50 1.35 0-16 1.01 (0-89) 2.93 200 P 0.77 0-26 0.83 0.25 1-05 1.5 1-88 1-76 4582 320 E 0-20 0 33 3.64 - 4-11 1-32 0-30 1.15 1.95 1-25

Decomposition products of C2HrOOCH20H 305 E 0.52 0-09 1.56 - 1-75 1-67 2-08 488 3-43 4-83

In confirmation of the predominantly surface character of the decomposition, it was found that passage of the peroxide in a nitrogen stream through




an unpacked tube at 180? C with a contact time of 4 sec., did not lead to more than 5 % decomposition. This also shows that the peroxide is protected from decomposition under such conditions and exists at comparatively high temperatures.

At 200? C the yields from a tube containing salt-coated glass wool were identical with those from the untreated wool. The decomposition in the empty tubes was no doubt of the explosive nature: there was considerable formation of fog. The essential difference is the formation of more carbon monoxide and hydrocarbon at the expense of aldehyde and alcohol. The figures in the table in parentheses are low in the case of aldehyde owing to the method of estimation causing polymerization to resins instead of oxidation to acid when peroxide was present; the alcohol figure is high owing to conversion of some peroxide to alcohol during analysis.

The absence of unsaturateds is notable; they were only formed if the peroxide exploded in bulk. Thus the gas from one of the tubes in which explosion had taken place at 150? had the following analysis: C02, 30 %; CO, 37-9 %; H2, 19-5 %; C2H4, 70 %; CH4, 29.1 %.

Decomposition products of propyl hydrogen peroxide

The decomposition products were obtained in the same way as those of the ethyl compound, viz. by distillation through a hot tube. In the first three experiments the tube was packed with glass wool. These are marked P; the other runs were carried out in an unpacked tube and are marked E. The run at 247? was in the presence of carbon dioxide, which was used as carrier; the reaction is, therefore, likely to be preponderantly homogeneous owing to the buffering effect of the gas. The last experiment at 300? was carried out by dropping the liquid peroxide from a tap funnel into a hot tube. The products were sucked out of the other end of the tube.

TABLE 7

Decomposition products of propyl hydrogen peroxide Yields in moles per 10 moles decomposed

Unde- Temp. C composed CO2 CO H2 C2H. C3H8 HCHO Ald. Alc. Unsat.

180 P 3-61 0-06 0.21 - 0.35 - 2.79 (2.82) (7.25) -

200 P 0-14 0 57 0-62 0-60 1-08 - 2-80 2.70 5.39 250 P 1-42 0 70 1-06 1.06 1.39 0-14 3-51 (1.48) 4 40 - 247 E 0 40 ? 0*20 0 40 - 0-81 3.55 4.2 4 03 0.2 300 E 0.2 0.36 3 10 1-43 - 3.26 3-62 2.0 2-68 0.2 300 E - 0-35 3-01 4.93 _. 3.35 ? 2-40 0.2 300 E 0-48 0.29 2*24 0 75 1.88 0 53 3.75 2*30 3.53 0-4



142 E. J. Harris

The parentheses have the same meaning as those in the preceding table. The dimedone derivatives of the aldehydes could be separated into the

high-melting formaldehyde compound (m.p. 1880) and a small amount of Low-melting material which could not be purified. The 3, 5, dinitrobenzoic ester of the alcohol was made, and it melted at 730, corresponding to propanol. The peroxide itself was found to form an ester by the same method (Schotten-Baumann); it melted at 620.

From table 7 it appears that the decomposition at surfaces forms considerable amounts of aldehyde and alcohol; the formaldehyde may be the result of the secondary oxidation of the higher aldehyde, the oxygen being provided by the reaction 2C3H700H-*2C3H70H+ 02. The high temperature decomposition (which is attended by the emission of light) forms a hydrocarbon with the apparent n value of 3; but as the boiling-point has not been determined owing to the small quantity of gas available, the possibility of the gas being a mixture of equal parts of ethane and butane remains.

Energy changes in the decomposition of ethyl and propyl hydrogen peroxides

On the basis of the determination of the heats of combustion by Dr Stathis in this laboratory, it is possible to evaluate the heats of some alternative reactions of decomposition.

The heat of combustion of the ethyl hydrogen peroxide giving an iodine liberation of 90 % theoretical was 5-4 kcal. per g. In view of the carbon per- centage being slightly over the theoretical, no correction has been applied to this figure. The corresponding value per g. mole is 335 kcal. and the heat of formation would then be 64 kcal. per mole. As the following heats of reaction refer to all compounds in the gaseous state, it has been necessary to assume a value for the latent heat, the figure used was 10 kcal. per mole.

TABLE 8. HEATS OF REACTION

Reaction R R' - zH kcal.

ROOH ->R'CHO + H20 C2H5 CH3 64 C3H7 C2H5 55

ROOH -+ROH + fO2 C2H5 - 8 C3H7 5

ROOH -+R'H + CH20 + 2O2 C2H5 CH3 - 0 C3H7 C2H5 - 6

ROOH -RH + 02 C2H5 - 23 C3H7 -30

The heat of combustion of propyl hydrogen peroxide was 6X3 kcal. per g. Allowing in this case for a purity of 98-5 % (this being based on the elemen-




tary analysis is no proof that the compound is in fact 98-5 % peroxide) the heat offormation (liquid) comes to 704 kcal. per mole,the heat of combustion being 486 kcal. per mole. The latent heat has been taken as 10 kcal. per mole.

These values will not hold for surface reactions and the effective value will depend on the heats of adsorption. The products of the slow decomposition indicate that the first and second reactions are involved, and the third may be the source of the lower hydrocarbon. At high temperatures when a homogeneous decomposition is more likely, the fourth process seems to take place.

DISCUSSION

The decomposition of dipropyl peroxide proves to be of precisely the same nature as that of the diethyl compound. The phenomenon of transfer to explosive reaction takes place at similar pressures to those for the latter, and is luminous in the presence of oxygen. Hence it appears safe to conclude that all the dialkyl peroxides have these properties. The importance of this in combustion chemistry is that the property of possessing a critical explosion pressure has been suggested by several workers as an explanation of the formation of cool flames during the course of slow-combustion reactions and was referred to in the previous paper (Harris and Egerton I938).

The fact that ethyl and propyl hydrogen peroxides decompose heterogeneously at once suggests a connexion between these compounds and the initiation of slow combustions. The factors which accelerate the decomposition of the peroxides (increased surface and salt washing) tend to inhibit the oxidation, and reduce the value of the exponent 0 in the expression for the reaction rate, at time t, w = AeOt. If the oxidation proceeds via an alkyl hydrogen peroxide as the precursor of the formation of radicals, then the catalysis of the decomposition to inert products will slow down the rate of formation of radicals which is likely to be a homogeneous process, for, if not, the wall would act as a third body in permitting their recombination.

The very faint general luminescence (not cool flame), which can be observed when slow combustion is taking place at 270-350?, may well be due to the continuous decomposition of an alkyl hydrogen peroxide. As is shown by this investigation, very small amounts would be required to account for the degree of luminosity observed.

Another reaction which is accelerated by salt washing is the oxidation of acetaldehyde (Pease I933). Whether a compound similar to ethyl hydrogen peroxide is involved in the reaction is not yet known, but the fact that the salt coating reduces the yield of peracetic acid to zero in the final oxidation



144 E. J. Harris

products makes it appear that the intermediate product (CH3CHO1. 02) is decomposed on a salt surface to other products such as carbon dioxide, instead of becoming stabilized as peracetic acid. The only alkyl peroxides that could reasonably be formed from acetaldehyde are the mono- and di-methyl compounds. Whether the cool flames which have been observed in the oxidation of acetaldehyde are to be attributed to one of them is uncertain, but methanol is a product of the reaction, as would be the case if the reactions CH300CH3 = CH30H + HCHO, or 2CH300H = CH30H +02 were taking place. Preliminary experiments which have been made on the decomposition of peracetic acid show that when it is decomposed at low pressure at 5000, methanol is one of the products, which can be explained if the equilibrium CH3COOOHtiCH3CHO. 02 is shifted to the right at high temperatures, and is followed by a breakdown of the aldehyde peroxide to methanol either directly or via the peroxides as above.

If traces of any of the alkyl peroxides be added to a mixture of propane and oxygen and then admitted to a quartz reaction vessel at 280-350?, the time before a measurable pressure change takes place is greatly reduced, but it is not always possible to reduce it to zero even by adding comparatively large amounts of peroxide. Provided the reaction is not very fast, the slope of the pressure-time curve is not affected by the peroxide. The condition of the surface of the vessel is of great importance in determining the efficiency of the peroxide and the slope. It has been observed that enough diethyl peroxide to promote an immediate flash as the propane-oxygen mixture is admitted to the vessel could be admixed without starting the propane oxidation in a vessel which had become inactive (probably owing to a trace of mercury or an alkaline substance). In the same vessel after activation with hydrofluoric acid, the reaction would proceed after a lapse of 18 min. even in absence of added peroxide. These facts indicate that the initiator of the hydrocarbon oxidation is not the peroxide itself, but most probably a decomposition product of one.

The decomposition products (both surface and homogeneous) of the peroxides are not such as could be distinguished from the oxidation products of hydrocarbons, and as in any case the amounts would need to be only very small to account for the luminous phenomena, it is unlikely that evidence will be found analytically for their intermediate formation. Nevertheless, the behaviour of the peroxides which have been investigated is providing very valuable information towards the elucidation of the nature of the combustion and ignition of hydrocarbons.




NOTES ON ANALYTICAL METHODS

The iodometric analysis of the ethyl and propyl hydrogen peroxides was carried out by weighing about 007 cm. of the substance into a weighing bottle containing 1 g. potassium iodide which had been dissolved in a little water, and 3-4 c.c. glacial acetic acid. After standing 15 min. the liquid was diluted and titrated with decinormal thiosulphate. A blank determination was carried out at the same time. The results when propyl hydrogen peroxide was used were apparently too low, as the maximum strength found was 78 %, whereas elementary analysis indicated that the composition was nearly that calculated for the pure compound.

The elementary analysis was carried out in a stream of nitrogen, as recommended by Rieche (I 93 I ).

The methods used for analysis of the aldehydes and alcohols have already been described (Harris and Egerton I937).

The author wishes to express his thanks to Professor Egerton for having suggested and taken a close interest in this work, and to the Trustees of the Beit Research Fellowships for a Fellowship which rendered it possible.

SUMMARY

The decomposition of dipropyl peroxide like that of diethyl peroxide (Harris and Egerton 1938) is a homogeneous, unimolecular reaction below a critical pressure: above this pressure explosive decomposition takes place. The limiting pressure varies with the temperature according to the law for thermal explosions. The products of slow decomposition are complex, and include propaldehyde and an alcohol: the explosive decomposition leads to formation of butane and formaldehyde.

The decomposition of ethyl and propyl hydrogen peroxides is heterogeneous and accelerated by increasing the surface or by coating the surface with salt. The vapours of these peroxides luminesce feebly when admitted to a hot vessel, probably due to combustion of part of the peroxide in oxygen liberated from the rest. The reaction products include aldehydes and alcohols, and at low temperatures oxygen is found, while at high temperatures the hydrocarbon having the same number of carbon atoms as the peroxide is formed.

All the peroxides ignite in air at very low pressures in the temperature range 200-300?, and the connexion between these results and the luminous phenomena of slow combustion is discussed.



146 E. J. Harris

REFERENCES

Allen, A. 0. and Rice, 0. K. I935 J. Amer. Chem. Soc. 57, 310. Baeyer and Villiger I9OI Ber. dtsch. chem. Ges. 34, 738. Harris, E. J. and Egerton, A. C. G. 1937 Chem. Rev. 21, 287.

- - I938 Proc. Roy. Soc. A, 168, 1-18. Medvedeef and Alexeeva 1932 Ber. dtsch. chem. Ges. B, 65, 133. Mondain-Monval 1932 Chim. et Ind. 27, 770. Pease, R. N. 1933 J. Amer. Chem. Soc. 55, 2753. Rieche 193I Alkyl Peroxyde und Ozonide. Dresden: Steinkopff.



Documents

The Decomposition of Alkyl Peroxides: Dipropyl Peroxide, Ethyl Hydrogen Peroxide and Propyl Hydrogen Peroxide