T H E AUTOXIDATION O F STANKOUS AND CUPROUS CHLORIDES BY AIR BY GEORGE
W.
FILSOS AND JAMES H. WALTOX
The oxidation of solutions of stannous chloride was first studied extensively by Young' who found that the reaction was somewhat dependent upon the acidity of the solution, highly sensitive to the action of certain catalysts such as copper and iron salts, and further, was inhibited by some of the alkaloids and certain other substances. The reaction was also studied by Miyamoto,? who determined the effect of acid concentration, the rate of bubbling of the oxygen, and several other factors. I t has been shown by many investigators that autoxidations are catalyzed by the presence of certain metallic ions, notably by those of copper, iron, manganese, cobalt, and nickel. This investigation was originally undertaken with the object of finding out whether or not a mixture of these ions exhibit promoter action in the case of autoxidations. Since no cases of promoter action were found, certain physical and chemical conditions affecting the autoxidation of stannous and cuprous chlorides have been studied. Reagents: In most cases no special precautions were taken to further purify the reagents inasmuch as very pure reagents were available. The stock solutions of stannous chloride were made by dissolving a weighed amount of the salt in air-free distilled water containing enough hydrochloric acid to prevent the precipitation of any of the basic chlorides of tin. These solutions were kept under an atmosphere of hydrogen. Apparatus: The apparatus used in these experiments was essentially the same as that used by W a l t ~ n .This ~ apparatus consists of a burette and leveling tube connected to a reaction flask by means of capillary glass tubing. The reaction flask was clamped in a vertical position and by means of an appropriate mechanism, it was rotated back and forth around the vertical axis. The only deviation from the original apparatus was in the shape of the reaction flask (Fig. I ) . The bottom of a 150 cc. round bottom Pyrex flask was heated in such a way that a narrow band, extending to within an inch of the end of the neck on either side and parallel to the vertical axis, was softened. This softened portion was then pressed in by the edge of an iron spatula so that a fissure about one-half inch deep was formed. A second fissure was made in a plane a t right angles to the first. When shaken, the ridge on the inside of the flask breaks the surface of the liquid and in this way drives the liquid into the gaseous phase and drives the gas into the liquid phase. Method of Procedure: The concentration of acid in the stannous chloride solution was determined by titration of a measured sample with standard
3
Young: ,J. Am. Chem. SOC., 23, 119 (1910). Miyarnoto: Bull. SOC.Chem., Japan, 2, 155 (1927) et. seq Walton: 2. physik. Chem., 47, 185 (1904).
AUTOXIDATION O F STANNOUS AND CUPROUS CHLORIDES BY AIR
741
sodium hydroxide. This gives the sum of the free hydrochloric acid and the acid that would be formed upon the complete hydrolysis of the stannous chloride or its basic salts. I n determining the rate of oxidation of the stannous chloride, a given quantity ( 2 j cc. unless otherwise noted) of the stock solution of stannous chloride was placed in the reaction flask which was then clamped into position and connected to a burette filled with pure oxygen. After the system had come to equilibrium, the shaking was started and the volume of oxygen absorbed was recorded by noting the decrease of volume in the burette. The reaction flask and capillary tubing connecting them with the burettes were filled with air. In this way a constant concentration of oxygen was maintained in the reaction flask because as the oxygen was used up in the reaction flask, it was supplied from the burette. This assumed,of course, that there is practically no diffusion of the nitrogen from the flask back through the capillary tubing, since the flow of oxygen is always toward the reaction flask and directly opposite to the direction of diffusion of nitrogen. By this method of procedure, it was found that the theoretical amount of oxygen was absorbed for complete oxidation of the stannous ions. In case a catalyst was to be added, it was placed in a capsule which could be held in the neck of the flask and dropped when desired by pulling a trigger in the side arm. The reaction was carried out at 30' unless otherwise stated. Measurement of Reactzon Velocity. At the beginning the reaction showed certain FIG.I irregularities so that uniform and reproReaction Flask. ducible results were not usually obtained until about ten per cent of the stannous chloride had been oxidized. From this period on, however, results which could be duplicated were easily obtainable. The amounts of oxygen consumed in equal time intervals were of the experiment. In view approximately equal during the first part (607~) of the fact that the solution is constantly saturated with oxygen and that the ratio of the concentration of stannous chloride to oxygen is relatively very great this would be expected. Towards the end of the run the amount of oxygen per unit time interval decreased somewhat probably due in part to decreased stannous chloride concentration and also to the using up of acid in the oxidation of the SnC12. It will be shown later that the speed of the reaction increases directly with the hydrogen ion concentration of the solu-
742
GEORGE W. FILSON AND JAMES H. WALTON
tion. For the comparison of experiments constants were calculated from the formula K = - vo - v
to
-t
in which vo is the volume as read on the burette a t the time to, while v and t are corresponding values a t later times. Thus K represents the average volume in cc. absorbed per minute. Table I gives the results of typical runs.
TABLE I Data for typical duplicate runs at 30' x = Volume of oxygen absorbed in time, t. K = x/t. SnClz t
- 14.5g. per liter; HCl - .606N. X
15
0.60
25
I.I5
35
1.55
50
2.35
60 80 I35 1.50
170
240
2.80
3.75 6.I O 6.65 7.40 IO.00
K
.040 ,046 '047 ,047 .047 ,047 '045 '044 ,044 ,042
Gas vols. a t oo and 760 mm. X
0.70 I .2 0
K ,047 ,048
I . 70
,049
2.40
,048 ,047 ,047 ,044 ,045 ,044 ,042
2.80
3.75 6.00 6.70 7.40 IO.00
Factors uflecting the Reaction. I n order to find out whether the reaction measured was the rate of solution of the oxygen or the rate of reaction of the dissolved oxygen experiments were carried out in which the rate of shaking was increased from goo to 1800 shakes per minute. The data in Table I1 indicate that for these ranges the rate of oxidation is independent of the rate of shaking and therefore it is the dissolved oxygen that is reacting. This oxygen is being supplied very rapidly, the rate of diffusion through the liquid being so rapid that there is very little difference in the concentration of the oxygen at the liquid-gas interface and the interior of the liquid. Experiment 3, Table 11,shows that a non polar lining (paraffin) is without influence on the speed of the reaction. Increasing the glass surface by the addition of powdered glass (Expt. 4) is also without effect. It was of interest to compare our method of studying this reaction with that of Miyamoto, who bubbled air into 40 cc. of SnClr solution in a test tube (diameter 3 cm.) through a glass tube with a 4 mm. opening a t a rate of 7.78 1. per hour. I n our experiments air was bubbled through 1 2j cc. of SnClz solution in a tube 2.8 cm. diameter, kept a t 30'. The absorption of oxygen was followed by removing 5 cc. samples at definite times and titrating for the stannous tin. 9.5 liters of air per hour were bubbled through the solution through a jet with a 4 mm. opening. I n another set of experiments the same
AUTOXIDATION O F STANNOUS AND CUPROUS CHLORIDES BY AIR
743
volume of air per hour was bubbled through four jets having openings between 0.5 and 0.75 mm. The concentration of stannous chloride was 36.3 g. per liter in each case with total acidity of 0.85 N HC1. With the four fine jets the speed of oxidation was about twice as great as in the case of the single larger jet. This indicates that in Miyamoto’s experiments the solution was not saturated and that he was measuring two reactions-one a t liquid gas interface, the other in solution.
FIG.2 Effect of HCI Concentration,
TABLE I1 K
Expt. No.
Rate of shaking per minute
I
,040
I800
2
,041 ,040
900 900 900
3 4
’
044
Concentration of SnC12
36.3 g. per liter
Remarks
Pyrex flask 11
71
Paraffin coated flask Pyrex flask z g. powdered glass Total acidity 0.85 N HCl
+
Eflect of Acid Concentration. U’ithin certain limits, the rate’of autoxidation of stannous chloride increases with the concentration of the hydrochloric acid as shown by the data in Table I11 and also by Curve I, Fig. 2. While the solubility of oxygen decreases with increasing concentration of hydro-
GEORGE
744
w.
n. WALTON
FILSON AND JAMES
chloric acid, the change is not great as evidenced by the solubility data of Geffkm4 He gives the solubility of oxygen at 25' as 0.0308 for water and 0.026j for 2 N HCl. Below the concentration of 2 N HCl, then, the effect of the solubility of oxygen is of minor importance. There is a definite upper limit beyond which the concentration of hydrochloric acid shows little effect on the rate of reaction. If the curve is continued downward there is a minimum concentration of hydrochloric acid (0.25 N) beyond which autoxidation would not take place. These data also show that within certain limits the reaction is independent of the concentration of the stannous chloride (Table 111,A and C) but dependent upon the total acid concentration of the solution, consequently the explanation of this reaction from the standpoint of a reaction of zero order must deal primarily with the mechanism of the action of the oxygen. I t is quite possible that the oxygen functions through the formation of a perstannate which undergoes reduction leaving the tin in the tetravalent state. Sufficient data are not available, however, to test this point, which will be made the subject of further experimentation.
TABLE I11 The Effect of Acid Concentration on the Reaction A B Sormality HC1
7.04 2.48 I .41 0.87 0.61 0.47
Normality
K 0.40
HCl
K
Norrnali ty HCI
I .oj
0.096
3.23
C K 0.54
.31
.io
,054
1.63
'15
'52
1.22
. I29
,088
.43
'035 .027
1.09
,118
.20
,045
.032
A-Contains 14.j g. SnClp per liter. B-Same as A but with another sample of acid. C-Contains 29.3 g. SnClz per liter. The relation of the speed of oxidation to the actual hydrogen ion concentration has been calculated and graphed (Curve 11) from data (Table IV) supplied by Dr. Art,hur Weber. The hydrogen ion concentrations were measured by means of the glass electrode.
TABLE IV Hydrogen Ion Concentration of Stannous Chloride Solutions Concentrat,ion of SnCL gm/l
Acidity as HCl (N)
29.3 29.3
.98 .96 2.80 4.08 1.95 3.91 I
4
Geffken: 2. physik. Chem., 49, 2j7 (1504).
Mols of Hydrogen Ion per liter
PH
+o.
20
0.631
-0. I 2
1.32
-0.25 -0.31
1.78 2.04
-0.Oj
1.12
-0.25
1.78
AUTOXIDATION O F STASNOUS AND CUPROUS CHLORIDES BY AIR
745
It, is plain that the oxidation of stannous chloride is directly proportional to the hydrogen ion concentration between p H of +o.z and - 0 . 2 5 . Metallic Ions as Catalysts. Of the several salts whose effect on the reaction was tried (Table 1') the outstanding example of positive catalysis was cupric chloride. With this salt the rate of reaction became a function of the rate of shaking. An increase from 900 to 1800 shakes per minute increased the constant from 1.28to 2.49. The action of the cupric chloride can be explained by the assumption of a reduction of the cupric to the cuprous ion by stannous chloride with a subsequent rapid oxidation of the cuprous to the cupric ion, this latter being the principal reaction. The effect of increased shaking could then be due to a catalysis of this reaction by the glass, or the rate of oxidation of the cuprous chloride might be so great that there was an appreciable drop in the oxygen concentration between the liquid surface and the interior of the liquid. TO test the action of the surface of the glass, two sets of experiments were performed, one using paraffin-lined flasks and the other adding an excess of powdered pyrex glass (Table VI). Although there seems to be quite an appreciable drop in the reaction velocity in the case of the paraffined flasks there is not the expected increase in reaction velocity in the case of the powdered glass. In neither case was the reaction velocity changed enough to explain the great difference occurring with the given change in the rate of shaking. I t appears as if the wax and paraffin act as inhibitors rather than that the glass acts to any great extent as a positive catalyst. I t seems probable therefore that the area of the liquid gas interface exposed per minute is the controlling factor. TABLE T Catalytic Effect of Certain Chlorides (40 millimolw per liter)
+
+
I
K
Fait added
K
,+dt added
c'uC12 FcCh CuC12 FeCI3 lInC12 31nC12 CuC12 Sone
CaC12 CuClz ThC1, XC12 Sone
.69
0.45 1.22
0.66 I .30
0.79 I . 2: 0
049
0.07j
0.040
0.048
3 0 g. SnCL per liter HC1 Concn. 2 . 1 r\'
Volume of solution-1
j
36.3 g. SnC12per liter HC1 concn. 0.8 j N cc.
The reaction may be expressed by the following equations:
+ +
+
(a) 2 CuClz SnCI2 = z CuCl SnClr (in solution) (b) 4 CuCl 4 HC1 02 = 4 CuC12 2 H20 (liquid gas interface). By adding cupric chloride to the solution of stannous chloride, then, the reaction becomes one of the autoxidation of the cuprous chloride instead of the stannous chloride, which autoxidation, at the rates of shaking used, becomes n reaction at the liquid gas interface.
+
+
GEORGE: w. FILSON AND JAMES
746
n. WALTON
TABLE VI Factors influencing the Cu++ catalysis of the autoxidation of SnC12’ 30 millimoles of CuClz per liter Shakes/min
Glass capsule, glass flasks Wax capsuleI2glass flask >I
JJ
JI
,>
” ” paraffined ” Glass capsule and flask 2 gms. of powdered glass Glass capsule and flask I gm. of powdered glass Glass capsule and flask
+ +
1 f 1 f
K
1800 1800 900 900 900
2.49 I .83 ,98
900
1.53
900
1.3
‘73 I
.60
Conc’n of SnCL 36.3 g/l. Total Rcidity 0.85 N HCI. 1 5 cc. sample. The capsules were used to add CuC12 to the SnClz solution. Those of wax were made of equal parts of beeswax and paraffin.
Temperature Coeficient. By changing the reaction temperatures from 30’ to 40”, the velocity constant for the autoxidation of SnClp changed from 0.041 to 0.085. This is again in disagreement with the findings of Miyamoto
who reported only a very slight change in the rate of reaction with a change of IO’. This difference may be explained by the fact that Miyamoto was not working with a reaction in solution. The change in reaction temperature of ten degrees would not cause a great change in the rate a t which oxygen diffuses across the liquid gas interface. While the temperature coefficient is affected by the solubility of oxygen, this effect would be a decrease of about 1 5 per cent only, consequently the temperature coefficient would be still higher if the concentration of oxygen was the same for the two temperatures.
The Autoxidation of Cuprous Chloride The catalysis of the autoxidation of SnClz shown by CuC12 led to a study of the autoxidation of CuCl solutions. Below a rate of 1800shakes per minute, the rate of shaking determines the rate of autoxidation of the cuprous ions. (Table VII.) I t was impractical to attempt a rate greater than 1800shakes per minute, because even at this rate the reaction flasks were often broken. The data in Table VI1 indicate that, under the conditions of these experiments, saturation was not being maintained and that the reaction was taking place mainly at the liquid gas interface. As a consequence these experiments are largely empirical, necessitating a constant rate of agitation. I t was found that reproducible results could be obtained with 900 shakes per minute. Experiment 3, Table VII, in which powdered glass was added and 4, in which a paraffin-lined flask was used, show that the glass surface of the flask is practically without effect in this reaction. Experiments j and 6 show the effect of the degree of dispersion of air bubbles on the autoxidation of CuCl by bubbling air through the solution.
AUTOXIDATION O F STANNOUS AND CUPROUS CHLORIDES BY AIR
747
TABLE VI1 Physical Factors affecting the Rate of Autoxidation of Cuprous Chloride (1.1 N HCl) CuC1- I S g. per liter No. I
2
3 4
*j *6
Volume of Solution
Rate of Shaking
K
cc. 2 0 cc. 2 s cc. 25 cc.
1800
3.2s
95 cc. 95 cc.
1j
900
1.24
900
I .32
900
0.94
-
0.087 0.052
Remarks
I gram of powdered glass Paraffin lined flask
Three jets, 0.5 to 0.75 mm. bore One jet, 4 mm. bore
* I O liters of air er hour bubbled through 9j cc. of solution, 1 7 cm. deep. The reaction was followed hy adsing asample of cuprous chloride to an excess of potassium dichromate and titrating the excess with standard titanous sulfate.
T h e Eflect o j Salts o n the Rate of Autoxidation of Cuprous Chloride. I n every case the addition of a salt diminished the rate of reaction, even in the case of added cupric chloride. Using 2 s cc. of solutions of cuprous chloride of acid concentration 2.1 N, the following values of K were obtained with the following salts; 3 0 millimoles per liter of each salt being used: Co++ 1.18; Ni++ 1.10; V++++ 1.03; U+++ 0.95; Th++++ 1.1j; Mn++ 0.97; Fe+++ 1 . 2 6 ; CuCl alone 1.28. This does not mean that some of these salts would not act as positive catalysts provided the reaction observed was one in solution rather than one taking place at the liquid gas interface. I n a reaction of the type a t hand, an increase in the rate of reaction when the oxygen was being used up a t a greater rate than the rate at which it was being supplied through the liquid gas interface would be improbable. No doubt the decrease in the rate of reaction in the presence of these salts was due t o the formation of complex ions which are not oxidized by the dissolved oxygen or, at least, are oxidized at a much slower rate than the cuprous chloride. Similar results were obtained by the addition of sodium chloride to the cuprous chloride solution in the following amounts: 0.5 g, I g, z g, which gave values of K as follows: 1.25, 1.10, 0.93, while I g of CuC12 gave a value of 1.10. K for cuprous chloride alone was 1.38. A change from 30' t o 40' had very little effect on the rate of reaction, which is characteristic of a surface reaction of this type. The acid concentration had very little effect upon the rate of reaction, the value of K for normality 3.45 being 0.87 while for 5.73 N HCl the value was 0.84. Whatever effect was observed indicates that the reaction tends to decrease in speed with increase in acid concentration, which may also be due to the complex ion formed in solution. Induced Reactions. The oxidation of cuprous chloride in the presence of the following compounds: arsenious acid, phosphorous acid, hydrazine sulfate, citric acid, and tartaric acid was tried. I t was found that the oxidation of
7 48
GEORGE W. FILSON AND JAMES H. WALTON
cuprous chloride caused a slight oxidation of all the materials tried with the exception of hydrazine sulfate which was not oxidized and phosphorous acid which autoxidizes in the presence of sufficient hydrochloric acid and cupric salt. The effect of the concentration of the salts used in these experiments upon the solubility of oxygen has not been discussed. From consideration of the work of MacArthurj it seems, however, that as the concentration of the salt increases the solubility of the oxygen decreases. I t follows then that the change in concentration of the stannous chloride in Table IV would cause some change in concentration of the dissolved oxygen, but this change in concentration was negligible. The autoxidation of stannous chloride is believed to be unusual in that the rate of autoxidation increases with the increase of acid concentration while the general tendency is for compounds to become more stable towards oxygen with increasing acid concentration.6 The increase in reactivity with higher concentrations of acid may be due to catalysis by hydrogen ions, in much the same way as many other reactions are catalyzed by hydrogen ions. It has been shown that the rate of reaction is in direct proportion to the hydrogen ion concentration. I t is also possible that at the higher concentrations of acid the formation of salt solvent complexes may play a part. There is also the possibility that the oxidation of the Sn++ ion occurs at a different rate than the SnOn-- ion, which would also account in some measure for the effect of the hydrogen ion concentration of the solution. The whole problem is complicated by the extent of hydrolysis of the tin salts, a subject which is now under investigation. These experiments with stannous and cuprous chlorides illustrate two extreme possibilities of autoxidation. The one is the case in which the oxygen is being supplied at a rate greater than the rate of autoxidation in which case the dissolved oxygen comes to equilibrium with the gaseous oxygen at the partial pressure in the gaseous phase. I n the other case, the autoxidation is more rapid than the supply of oxygen and the reaction takes place a t the liquid gas interface while the interior of the liquid may be almost void of dissolved oxygen.
Summary The autoxidation of stannous chloride occurs in solution as is evidenced by the high temperature coefficient and the fact that an increase in the liquid gas interface does not cause an increase in the rate of autoxidation. 2. Within certain limits the rate of autoxidation of stannous chloride increases directly with the increase in the acid concentration, and is directly proportional to the hydrogen ion concentration. 3 . Within certain limits the rate of autoxidation of cuprous chloride is independent of the acid concentration. I.
@
MacArthur. ,J Phys Chem, 20, 4 9 j (1916). See Lamb and Elder. J. Am. Chem. SOC.,53, 147 (1931,
AUTOXIDATION OF STANNOUS AND CUPROUS CHLORIDES BY AIR
749
4. The catalytic effect of a number of salts on the rate of autoxidation of stannous chloride has been investigated. Cupric chloride was outstanding in its catalytic activity] which is explained by a cycle of oxidation and reduction involving the cupric and stannous chlorides and dissolved oxygen. 5 . Of the investigated salts, none was found to increase the rate of autoxidation of cuprous chloride. 6 . The autoxidation of the cuprous chloride induced the oxidation of citric and tartaric acids as well as of arsenious acid. 7. Certain published results on the autoxidation of stannous chloride which differ from the data obtained by the authors are apparently due to experiments with unsaturated solutions of oxygen. M a d i s o n , Wisconszn
