The Autoxidation of Stannous Chloride. II. A Survey of Certain Factors

THE AUTOXIDATION O F STANNOUS CHLORIDE. I1 A SURVEYOF CERTAINFACTORS AFFECTINGTHIS REACTION’ ROBERT C . HARING

AND

JAMES H. WALTON

Downloaded by NANYANG TECHNOLOGICAL UNIV on August 26, 2015 | http://pubs.acs.org Publication Date: January 1, 1932 | doi: 10.1021/j150343a016

Department of Chemistru, Universitg of Wisconsin, Madison, Wisconsin Received September 30, 1939

The autoxidation of stannous chloride has recently been investigated by Filson and Walton (l),who studied the effect of hydrochloric acid on the reaction velocity and showed a linear relation between the velocity and the hydrogen-ion concentration. This paper is a continuation of the above work and is a general survey of certain factors influencing the rate of reaction. APPARATUS AND PROCEDURE

Twenty-five cc. of an acid solution of stannous chloride were placed in a 150-cc. Pyrex flask of the type used by Filson and Walton. The flask was placed in a thermostated shaking apparatus and connected to a waterjacketed burette containing oxygen gas a t the temperature of the thermostat, 25°C. When the system had come to constant temperature, the shaking apparatus was started and readings were taken on the volume of oxygen absorbed. The speed of shaking was selected as about 1000, since Filson and Walton had shown that higher speeds have no effect on the reaction velocity. REAGENTS AND SOLUTIONS

Several brands of stannous chloride were used and were found to give varying results. Baker’s “purified” salt gave results which were only slightly changed by three recrystallizations under nitrogen from hydrochloric acid solution, consequently this preparation was used in the experimental work. Solutions made from stannous chloride which had been dehydrated by treatment with excess acetic anhydride and washed with anhydrous ether (2) gave results which agreed with those obtained from the recrystallized hydrated salt. The solutions, made in 2-liter batches and stored under nitrogen, contained approximately 32 grams of stannous chloride per liter, and were I This research was financed by a grant from the Research Committee of the University of Wisconsin, Dean C. S. Slichter, Chairman. 133

THE JOURNAL OF PHYSICAL CHEMISTRY, VOL. XXXVII, NO.

1

134

ROBERT C. HARING AND JAMES H. WALTON

about 0.8 N in total hydrochloric acid (free acid plus that obtained by complete hydrolysis). Table 1gives the data for the solutions used in this work,


ORDER O F T H E REACTION

By neglecting the first 80 to 120 minutes of the run, during which the absorption was not always reproducible, a value of K ( K = milligrams of oxygen absorbed per minute) was obtained which corresponds to a zero order reaction, within an experimental error of 5 to 7 per cent in the course of a 4 to 5 hour run. This neglected period corresponds to about a 10 per cent absorption of the theoretical amount of oxygen required to oxidize the sample completely. Duplicate runs for solution No. 6 are given in TABLE 1 Composition of the stannous chloride solutions ~~

SOLUTION NUMBER

10 11 12

SOURCE OF

SnCh

SnClz

HCI

(TOTAL)

gramslliter

Normality

36.5 31.3 31.7 35.5 32.6

0.783 0.758 0.733 0.814 0.782

Baker’s Purified Baker’s Purified Baker’s Purified Mallinckrodt’s C.P. General Chemical Co. C.P. (dehy drated by acetic anhydride) Mallinckrodt’s C.P. (dehydrated by acetic anhydride) Baker’s Purified General Chemical Co. C.P. (recrystallized three times)

.

32.7

0.811

33.6 33.1

0.840 0.815

table 2, this solution having an average value of K of 0.0572 mg. per minute. The 5 to 7 per cent drop in the value of K can be attributed to a decrease in the concentration of the free hydrochloric acid, since this substance is used up during the oxidation of the stannous chloride. Some hydrochloric acid will be available through hydrolysis of the stannic chloride, but since the complex equilibrium between Sn++, Sn++++,HCl, and complex chloro acids is not fully known it is impossible to define the system definitely . E F F E C T O F TEMPERATURE

Runs were made on solution No. 12 a t different temperatures. It was impossible to calculate values of K for comparison, since the apparent order of the reaction varied with the temperature. This change can be attributed to several factors. First, the solubility of oxygen is higher at bhe low temperatures, allowing a higher saturation concentration of oxygen,

135

AUTOXIDATION OF STANNOUS CHLORIDE. I1

and thus increasing the reaction rate. Then, too, the change of temperature affects the equilibrium between Sn++, SnClz and the chloro acid complexes. These factors mask the true temperature coefficient of the reaction, so that the results are of qualitative importance only. The results


TABLE 2 Data for duplicate runs at 25°C. Solution N o . 6 Stannous chloride, 31.3 grams per liter; hydrochloric acid (total) 0.758 N RUN 185 TIME

Oxygen

absorbed

minutes

milligrams

10 20 30 40 50 60 70 80 90 100 110 120 140 150 160 170 180 190 200 210 220 230 250 270

0.60 1.17 1.77 2.35 2.87 3.48 4.01 4.56 5.10 5.67 6.19 7.29 7.86 8.37 8.96 9.48 10.06 10.70 11.19 11.75 12.31 12.89 13.91 15.00

RUN

K mg./min.

0.0600 0.0587 0.0592 0.0587 0.0575 0.0581 0.0572 0.0570 0.0567 0.0567 0.0562 0.0556 0.0562 0.0559 0.0560 0.0558 0.0559 0.0562 0.0559 0.0559 0.0559 0.0561 0.0556 0.0555 Av. = 0.0568

168

Oxygen absorbed

X

milligrams

mg./min.

0.60 1.19 1.78 2.35 2.93 3.55 4.07 4.67 5.20 5.77 6.27 7.42 7.97 8.51 9.10 9.63 10.21 10.80 11.37 11.93 12.48 13.07 14.11 15.23

0.0600 0.0593 0.0594 0.0587 0.0587 0.0591 0.0581 0.0583 0.0578 0.0577 0.0570 0.0570 0.0570 0.0567 0.0568 0.0568 0.0567 0.0568 0,0568 0,0568 0.0567 0,0568 0.0566 0.0565 Av. = 0.0576

are expressed in terms of the t,ime required to oxidize one-half of the stannous chloride, starting at the beginning of the run. From the curves in figure 1, the following “half-times” were interpolated: At 45°C.. . . . . . . . . . . , . , . . . . . , . . . . . . , . , . . . . . . . , . , , . . . . . . . . , . . . 35°C.. . . . . . . . . . . . . . . . . . . . . . . . . . , . . . . . . . . . . . . . . . . . . . . . . . . . 25°C.. . . . . . . , . . , . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . , . . . . . . . . . 15°C. . . . . . . . . . . . . . . . . . , . . . , . , , , , . . . . . . . . . . . . . , . . . . . . . . . O’C.. . . . . . . . . . , . . . . . . . . . , . . . . , , , , . . . , , . . . , . , . . , . . . . . . . . .

.

144 minutes 228 minutes 320 minutes 384 minutes 330 minutes

136


The solubility of oxygen between 15" and 0°C. shows a marked increase, and this is probably the cause of the decrease in the time of half absorption for the lower temperature. EVIDENCE FOR A CHAIN REACTION


A number of tests were made to determine whether or not there is any evidence that the reaction under consideration belongs to the chain type.of

ON THE RATE OF OXYQENCONSUMPTION BY FIG.1. THEEFFECTOF TEMPERATURE

STANNOUS CHLORIDE

reaction similar to the oxidation of sodium sulfite studied by Alyea and Backstrom (3). This reaction meets the criterion of being exothermic, as was shown by Berthelot (4), whose work was done a t 500°C. A confirmatory experiment in our laboratory, a t room temperature, also indicated that the reaction is exothermic. For the purpose of confirming the chain mechanism, tests were made on: (1) the influence of light; (2) the existence of an intermediate peroxide; (3) the effect of added substances, especially inhibitors; (4) the coupled oxidation of a second molecular species present in the system.


137


1. E f e c t of light

Visible light had no effect on the reaction rates. Two flasks were painted black, and the runs made in these showed no difference from those in unpainted flasks, In addition, some experiments were made in which a 250-watt tungsten light was placed adjacent to the reaction flask, and other experiments in which the direct beam of the strong green and yellow light from a high intensity capillary mercury vapor lamp was used. I n both cases no acceleration of the reaction was noted. It was discovered, however, that ultra-violet light of wave length below 3070A. was absorbed completely by the solution and that it speeded up the reaction very considerably. The details of these experiments will be discussed in a subsequent paper.

2. Detection of peroxide

A sample of the partially oxidized stannous chloride was found to give a peroxide test with titanium sulfate. This test could not be obtained after the solution had been allowed to stand for 2 or 3 hours, showing that the peroxide had been used up by reaction with more stannous chloride. No attempts were made to isolate the peroxide, because of the small amount present and its relative instability. 3. Catalytic egects of certain substances

One of the most important of the criteria of photochemical chain reactions, which can also be applied to thermal chain reactions, is the great effect of negative catalysts. The reaction of stannous chloride with molecular oxygen is very sensitive to the action of added substances, as shown both by our work and by that of Young (5). In attempting to duplicate Young’s work on inhibition by alkaloids, it was found that cocaine, morphine, and brucine had little effect in 0.001 M solution. In 0.01 M solution a reduction to about 90 per cent of the normal rate was shown, and in a saturated solution, a reduction of 55-65 per cent of the normal values. Because of the effect of the alkaloids it was of interest to investigate the action of other nitrogen compounds; accordingly a series of experiments was carried out in which many varieties of these compounds were used. The results are shown in table 3, in which the value of K is given, together with the percentage value of K in relation to the uncatalyzed value of K for the stock solution used in each experiment. It will be noticed that nitrogen compounds, in general, were inhibitors, but that the inhibiting power is affected by the position of the nitrogen and the nature of the other atoms bound to it. Amino compounds had little effect, some of them slightly inhibiting the reaction and some accelerating it slightly. Nitro compounds were the most effective of the inhibitors studied, and the mole-

138


TABLE 3 The efect of certain nitrogen compounds Normal K values for the solutions: No. 5, K = 0.0626; No. 6, K = 0.0572; No. 7, K = 0.0583 mg. per minute PER CENT OB

COMPOUND

NORMAL


moles jliter

5 5 5 5 5 5 6

-

0.001 0.001 0.1 0.01 0.1 0.01 0.01

0.01 0.001 0.01 0.001 0.01 0.01 0.001 0.001 0.01 0.001 0.01 0.00001 0.0001 0.0005 0.001 0.01

-

..... 91.7 55.1 90.3 65.0

..... 96.5

p-Nitrotoluene Nitronaphthalene Nitromethane o-Nitrophenol p-Nitrophenol m-Nitroaniline p-Nitrobenzoic acid Nitrophenylhydrazine m-Dinitrobenzene sym-Trinitrobenzene Trinitrotoluene Picric acid Picric acid Picric acid Picric acid Picric acid

0.0112 0.0147 0.0114 0.0106 0.0156 0,0012 to 0.0043* 0.0120 0.0119 0.0029 0.0081 0.0058 0.0330 t o 0.0444 0.0170 0.0154 0.0122 to 0.0179 0.0145

19.6 25.7 19.6 18.6 26.8

About 5 21.0 20.8 5.1 14.2 10.1 About 60 29.7 26.9 About 25 25.3

Amino and imino compounds

5 5 5 6 6 6 7 7 7

0.0620 0.0574 0.0345 0.0565 0.0407 0.0622 0.0552

Nitro compounds

6 6 7 6 7 6 6 6 6 6 6 6 6 6 6 6

mg.jmin.

Cocaine hydrochloride Brucine Brucine Morphine hydrochloride Morphine hydrochloride Nicotine Verona1

0.01 0.01 0.01 0.01 0.01 0.01

0.002 0.0045 0.01

Glycine p-Toluidine hydrochloride Phenylhydrazine hydrochloride Carbazole Hydroxylamine hydrochloride Urea Thiourea Thiourea Thiourea

0.0631 0,0564 0.0382 0.0535 0.0582 0.0595 0.0656 0.1120 About 1 . 2

..... 90.1 61.0 93.5

..... ..... 112.5 192.0 2060,O

*The two values of K in some experiments denote a rise in the value of K , due to the destruction of the inhibitor.

AUTOXIDATION O F STANNOUS CHLORIDE. I1

139

TABLE 3-Concluded PBR CENT OF NORMAL

COMPOUND

Amino and imino compounds-Concluded


moleallifei

0.01 0.01 0.01 0.01 0.01

mg.lniin.

0.0465 0.0584 0.0571

Ammonium thiocyanate Sulfanilic acid Acetanilide Semicarbazide hydrochloride p-Aminophenol

0.0501

0.1028

79.8

..... ..... 85.9 176.3

Other nitrogen compounds 6 7 7 7 7 7 7 7 7

7 7

0.001 0.01

0.01 0.01 0.01 0.01 0.001 0.01 0.01 0.01 0.01

Cellulose nitrate Potassium nitrate Potassium cyanide Butyl cyanide Acetoxime Azobenaene Aminoazobenzene Aminoazobenzene Butyl nitrite Azoxybenzene Guanidine nitrate

0.0470 0.0564 0.0517 0.0536 0.0567 0.0208 to 0.0298 0.0587 0.0642 0.0576 0.0523 0.0421

..... 96.7 88.7 91.9 97.3 About 45

..... 110.0

..... 89.7 72.2

cules with two or three nitro groups were more effective than those with one. In contrast with this, organic and inorganic nitrate salts showed little or no effect. The azo and hydrazine groups were fair inhibitors, while the cyanide, oxime, azoxy, amino, and nitrite groups had very slight inhibitory power, or none at all. In general no attempt was made to discover whether or not the catalyst, either positive or negative, had been destroyed during the course of the reaction. In the case of picric acid the strong yellow color disappears during t,he course of the oxidation, and at the same time, a rise takes place in the rate of oxidation. Both of these facts indicate that the picric acid is destroyed during the reaction. Thiourea is a marked exception to the other nitrogen compounds since it showed a strong acceleration of the oxidation in concentrations as low as 0.002 M . As the concentration was increased the amount of acceleration of the reaction increased, until at 0.01 M the initial rate was about 1750 per cent of the uncatalyzed value. However, at this concentration the rate of oxidation drops off rapidly, indicating that the thiourea is being destroyed. In studying the action of thiourea, two special runs were made, one using 0.01 M thiourea in 0.8 N hydrochloric acid, and the second similar, except for the addition of a small amount of stannous chloride


140

ROBERT C. HARING AND JAMES €I WALTON .

(0.007 M ) . In the first case, no absorption was noticed, indicating that under these conditions thiourea alone does not absorb oxygen. In the second run oxygen was absorbed until the stannous chloride was used up, when the absorption stopped. The accelerating action of thiourea is apparently due to the sulfur atom, since amino compounds have been shown to have little effect, and yet this action is very specific since some of the other sulfur compounds used had no effect (benzyl sulfide, sulfanilic acid, benzenesulfinic acid, and powdered sulfur). Some sulfhydryl compounds were used, such as cysteine, thiocresol, thiobarbituric acid, and thiosalicylic acid. These were all inhibitors, reducing the rate to 4 0 4 5 per cent of its normal value. Thioglycollic acid however had practically no effect on the rate. This is in agreement with Schoberl’s work ( 6 ) in which he found a negative catalytic effect with cysteine and glutathione on the autoxidation of leuco methylene blue, but is opposite to the accelerating effect of sulfhydryl compounds in biological phenomena (7). The oxygen compounds used had little effect. The aliphatic alcohols increased the absorption rate 5 to 10 per cent and phenol had no effect, but hydroquinone, pyrogallol, and mannite decreased the value of K to about 90 per cent, and quinhydrone to 67 per cent of its uncatalyzed value. Other oxygen compounds, of widely different types-acids, ketones, thymol, ethyl ether, benzoyl peroxide, and paraformaldehyde-showed no appreciable effect on the rate of absorption. Tetraethyllead gave a value for K of 455 per cent of the uncatalyzed value. Ammonium thiocyanate was a good inhibitor; potassium cyanide, potassium nitrate, benzene, and anthracene decreased the value of K , 5 to 15 per cent. Certain other compounds had no effect; this group includes arsanilic acid, magnesium chloride, magnesium pyrophosphate, dibasic sodium phosphate, sodium arsenate, sodium pyrophosphate, and mineral oil. Silica gel, ground to 180 mesh, had no effect on the rate, showing that the reaction does not take place on the surface of the flask, but is a true solution reaction. Catalysis by charcoal The catalytic acceleration of many autoxidations by charcoal suggested its use in this study. Powdered willow charcoal accelerated the reaction but not in proportion to the weight used. For example, 0.12 mg. per CC. increased the rate to 167 per cent of its normal value; 0.6 mg. per cc. increased it to 228 per cent; and 1.2 mg. per cc. increased it to 253 per cent. This differs from the results of Lamb and Elder (8) who found for autoxidation of ferrous sulfate that the increase in oxidation rate was proportional to the amount of charcoal used. As in other autoxidations in which charcoal has been used it was found that the “active centers” on the charcoal can be promoted or poisoned.



14 1

Promotion was shown by copper salts. The initial absorption rate for charcoal alone was 1.72 mg. per 10 minutes, for cupric chloride alone, 2.01 mg. per 10 minutes, while cupric chloride and charcoal together gave a value of 10.80 mg. per 10 minutes. These values are given in grams, since a value of K could not be calculated for the promoted catalyst. The poisoning effect was shown by picric acid and by sodium arsenate. The latter compound had no effect on the reaction rate when used alone; charcoal alone increased the rate to 228 per cent of normal, but when the two were used together the rate was reduced to exactly the uncatalyzed value. This is similar to the results of Lamb and Elder (S), in which they found that phenylurea, amyl alcohol, and acetanilide when used alone showed no effect on the autoxidation of ferrous sulfate, but that they destroyed the accelerating influence of charcoal on the reaction. Picric acid showed a strong inhibition for the stannous chloride reaction, and when it was used in conjunction with charcoal, absorption rates were obtained which lay between the values for the two substances singly. According to the Taylor theory (9) these compounds are preferentially absorbed a t the most active points and prevent the usual accelerating action of the charcoal.

4. Coupled oxidation It was found that while an aqueous solution of allyl alcohol does not absorb oxygen, there is coupled oxidation of the allyl alcohol in the presence of stannous chloride. In order to determine the amount of excess absorption, the apparatus was arranged so that solid stannous chloride could be held above the solution in a glass capsule until the gas had come to constant temperature. The capsule was dropped into the acid solution with the first shake of the machine. I n this way, it was possible to measure the total amount of oxygen absorbed by the system. By using pure oxygen as the atmosphere, the experiment was completed in a short time, Table 4 gives the results of this work, showing the effect of varying amounts of stannous chloride, allyl alcohol, and acid. The effects of temperature and of added substances are also given. The per cent excess absorption increased with increasing amounts of allyl alcohol, and decreased with increasing amounts of stannous chloride and hydrochloric acid. These effects may be explained by the assumption that the allyl alcohol is oxidized by the primary peroxide. With increased concentration of allyl alcohol more of the chains will involve alcohol molecules, and thus a larger amount will be oxidized. Increased amounts of stannous chloride or hydrochloric acid cause an increase in the rate of consumption of oxygen, and a t the same time cause a decrease in the amount of induced oxidation. Increased temperature, and homogeneous catalysts such as cupric chloride and thiourea, also have

142


TABLE 4 Induced oxidation of allyl alcohol (a) Effect of varying amounts of stannous chloride 0.854 gram allyl alcohol in 25 cc. of 0.440 N HC1

I


SnCh

EXCESS 0 2 ABSORBED

grana8

per cent

0.127 0.381 0.635

45.1 34.8 28.1

(b) Effect of varying amounts of allyl alcohol

I

ALLYL ALCOROL

EXCESS 0 2 ABSORBED

(1) 0.381 gram SnCl:! in 25 cc. of 0.440 N HCl grains

per cent

0.171 0.854 1.708

19.4 34.8 36.2

(2) 0.635 gram SnClr in 25 cc. of 0.440 N HCI 0.171 0.854 1,708

14.8 28.1 31.3 (c) Effect of temperature

0.381 gram SnClz in 25 cc. of 0.440 N HCI EXCESS

0%ABSORBED

ALLYL ALCOEOL

15°C.

grams

per cent

per cent

per cent

0.085 0.171 0.854 1.708

16.8 20.1 37.2 40.1

19.4 34.8 36.2

16.4 27.6 36.3

....

(d) Effect of additional substances with allyl alcohol 0.381 gram SnClz 0.171 gram allyl alcohol in 25 cc. 0.440 N HC1

+,

ADDED SUBST.4NCE

I

EXCESS 0 2 ABSORBED

per cent

None +O. 1 gram thiourea +O. 1 gram charcoal +O. 1 gram CuCli

19.4 7.2 23.7 8.7

143

dUTOXIDATION O F STASNOUS CHLORIDE. I1

EXCESS

ACID

01

ABSORBED

per cent

None (25 cc. HsO). . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 10 CC. 0.440 N 15 C C . HzO. . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 25 CC. 0.440 1Y . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . 25 cc. 4 . 0 N . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Downloaded by NANYANG TECHNOLOGICAL UNIV on August 26, 2015 | http://pubs.acs.org Publication Date: January 1, 1932 | doi: 10.1021/j150343a016

+

28.1 44.4 34.8 7.7

TABLE 5 Induced oxidnlion i n non-aqimous solvents TOTAL ORA118 OF SOLVENT

SPEED OF REACTION

OXIYOEN IXED B Y 0.320 GRA?J O F ANHTDROUS STANNOUS CHLORIDE

EXCESS 3XYQEIN USED

per cent

Ethyl alcohol. . . . . . . . . . . . . . . . . . . . . . . Isopropyl alcohol. . . . . . . . . . . . . . . . . . . Amyl alcohol. . . . . . . . . . . . . . . . . . . . . . . n-Heptyl alcohol. . . . . . . . . . . . . . . . . . . . C yclohexanol . . . . . . . . . . . . . . . . . . . . . . . Benzyl alcohol, ..................... Phenylethyl alcohol.. . . . . . . . . . . . . . . . Cresol . . . . . . . . . . . . . . . . . . . . . . . . . . . . . .

Rapid Rapid Rapid Rapid Rapid Slow

Acetic acid. . . . . . . . . . . . . . . Propionic acid. . . . . . . . . . . . . . . . . . . . . . . Ethyl acetate.. ................. Butyl butyrate.. ..................... Scetone. .............................

Rapid Rapid Rapid Rapid Rapid

0.04194 0.04450 0.04777 0.04912

About 50 55.3 64.3 76.9 81.9

SlOlV

SlO\V 0.03585 0.04088 0.03169 0.03454

32.3 51.4 17.4 28.0 About 30

SlO\V

Aniline . . . . . . . . . . . . . . . . . . . . . . . . . . . . . . Glycerine . . . . . . . . . . . . . . . . . . . . . . . . . . . . Benzene

SlO\V SlO\V Slow Slow Very slo\v

the dual effect of speeding up the rate of consumpt'ion and of decreasing the amount of induced oxidation. Charcoal, on the other hand, while it is also a positive catalyst, increased the amount of induced oxidation. This may be explained by the higher concentrations of allyl alcohol produced a t the surface of the charcoal by adsorption. Several other substances were used in place of allyl alcohol in investigating the phenomenon of induced oxidation. It was found that there was no


144


excess oxygen used when acetone, phosphorus acid, maleic acid, or paminophenol were used, but that butyraldehyde used 14.8 per cent excess. Coupled oxidation was also found when stannous chloride was oxidized in non-aqueous solvents. The solvents used are shown in table 5 together with the total grams of oxygen used and the per cent excess over the theoretical amount necessary for oxidation of the stannous chloride. There seem to be two classes or types of solvents among those tried, one class in which the oxidation is fast and the second in which it is slow. In the former group the oxidation of the sample was complete in 3 to 10 minutes, in spite of the fact that no excess hydrochloric acid was present as in the case of oxidation in aqueous solution. I n each solvent of this group an excess of oxygen was used. In the rest of the solvents the reaction was comparatively slow, so that it was not followed to completion. There is no relation between the phenomenon of molecular compound formation and this division into types, for substances which form molecular compounds with stannous chloride are found in both classes. Neither is there a relation between ease of oxidation and the separation into types, since compounds which are easy to oxidize and others which are hard to oxidize are found in both classes. DISCUSSION

The results of the foregoing experiments are best explained by the wellknown peroxide theory of autoxidation, in which it is held that oxidations of this type are attended by the intermediate formation of metastable or dative peroxides which are characterized by high energy content and great instability. Such peroxides are very active and account for many of the phenomena of autoxidation (10). Mention has been made of the fact that ' in the autoxidation of stannous chloride enough peroxide is formed to be detected qualitatively. Further evidence of the existence of a peroxide is afforded by the induced oxidation of many compounds, especially alcohols. That this reaction is a chain reaction is supported by the great influence of added substances, particularly inhibitors, on the course of the oxidation, and also by the fact that the reaction is exothermic. There is also further support from photochemical data. I n the autoxidation of stannous chloride there is reason to believe that a complex chloro acid is the form in which the Sn++ is oxidized. Prytz (11) has shown that the complex ions SnC13- and SnClr-- exist in aqueous solutions of comparatively low hydrochloric acid concentration. This conclusion is supported by the fact that an excess of hydrochloric acid has a great accelerating action on the autoxidation of stannous chloride. Complexes are also known between stannous chloride and potassium chloride or sodium chloride, and these substances act in a manner similar to that of hydrochloric acid in accelerating the autoxidation of stannous chloride, but to

AUTOXIDATION O F STANNOUS CHLORIDE, I1

145

a less degree than does hydrochloric acid. I n another paper further evidence will be given to show that in the photochemical autoxidation the substance which actually absorbs the light quanta and becomes activated is a complex ion of this type.


SUMMARY

1. Increased temperature increases the rate of autoxidation of stannous chloride, but the true temperature coefficient is masked by the change in the solubility of oxygen and changes in the complex equilibria existing in the solution. 2. Visible light has no effect on the reaction but ultra-violet light below 30701. is absorbed and accelerates the reaction. 3. A peroxide was detected in the fresh, partially oxidized sample of stannous chloride. This peroxide disappears by reaction with further stannous chloride. 4. Many organic substances? particularly nitro compounds, were found to be inhibitors of the reaction. Picric acid was the strongest inhibitor found, and this inhibitor was shown to be destroyed during the oxidation. 5. Thiourea was found to be a strong positive catalyst for the autoxidation. 6. Poisoning and promoter action were found for the catalysis of the reaction by powdered willow charcoal. 7. Coupled oxidation of a second molecular species was found in aqueous solution in the cases of allyl alcohol and butyraldehyde, and in many nonaqueous solvents such as alcohols, acids, esters, and ketones. REFERENCES (1) FILSON,G. W.,AND WALTON,J. H. : J. Phys. Chem. 36,740 (1932). (2) STEPHEN, H.:J. Chem. Soc. 1930,2786. (3) ALYEA,H. N.,AND BACKSTROM, H. L. J. : J. Am. Chem. SOC.61,90(1929). (4) MELLOR,J. W.: A Comprehensive Treatise on Inorganic and Theoretical Chemistry, Vol. 7,p. 427. Longmans, Green and Company, London (1927). (5) YOUNG, S. W.: J. Am. Chem. SOC.23,119 (1901). (6) SCHOBERL, A. : Ber. 64,546(1931). (7) MILAS,N.A. : Chem. Rev. 10,295(1932),especially p. 312. (8) LAMB,A.B., AND ELDER,L. W., JR.: J. Am. Chem. SOC.63,137 (1931). (9) TAYLOR, H.S.: J. Franklin Inst. 194,l (1922). (10) MILAS,N.A.: Chem. Rev. 10,295(1932),for an excellent review of autoxidation. (11) PRYTZ, M. : Z. anorg. Chem. 172,147(1928).

The Autoxidation of Stannous Chloride. II. A Survey of Certain Factors

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