T H E CATALYTIC DECOMPOSITION O F HYDROGEN PEROXIDE BY SODIUM IODIDE I N MIXED SOLVENTS’
B Y VAN L. BOHNSON
1. Historical
It is well known that the velocity of a c,,emical reaction is influenced by the composition of the medium in which it takes place. The first case on record was noted by Berthelot and P6an de St. Gilles,2 who reported the velocity of esterification t o be less in ether than in benzene. The literature records a large number of experiments which illustrate an effect of the solvent, a few of the results of which are summarized in Table I. The figures represent the relative velocity constants of the several reactions in the different solvents, and are quoted from a more comprehensive table given by Patterson and Montgomerie. The reactions represented in the summary are: (I) The reaction between triethylamine and ethyl iodide. (11) The transformation of a triazole derivative into a diazoanilid ester.; (111) The inversion of menthone.B (IV) The transformation of anissynaldoxime into anisantialdoxime. Part of a thesis submitted t o the Graduate School of the University of Wisconsin in partial fulfilment of the requirements for the degree of Doctor of Philosophy. * Berthelot and Pkan de St. Gilles: Ann. Chim. Phys., ( 3 ) 66, 62 (1862). Patterson and Montgomerie: Jour. Chern. SOC.,101, 26 (1912). Menschutkin:Zeit. phys. Chern., 6,41 (189o);seealsoIbid.,I , 611 (1887). Dimroth: Liebig’s Ann., 335, I ; 338, 143; 373, 336; 377, 1 2 7 ; 399, 9 1 . Tubandt: Liebig’s Ann., 339, 41; 354, 2 5 9 ; 377, 284. i von Halban: Zeit. phys. Chem., 67, 129 (1909).
-
Van L. Bohizson
678
TABLE I Solvent
Water Methyl Alcohol Ethyl Alcohol n-Propyl Alcohol Isobutyl Alcohol
I
286.6 203.3 -
143.3
I1
0.0043 I .o 1.92
-
I11
-
.o 2.6 3.35 4.64 I
IV
I .o 1.0
1.8
-
11
8.26 2.07
1.86 1.69 0.96
With these solvents in mind, it is interesting to note that the action of sodium hydroxide on alkyl bromides is slower in methyl alcohol than in "-propyl, but fastest in ethyl alcohol, while the velocity of decomposition of iodoform1 in sunlight increases according to the solvent in the following order: methyl alcohol, ethyl alcohol, propyl alcohol. It is frequently pointed out that there is a parallelism between reaction velocity and the dissociating power (i. e., according to Nernst, the dielectric constant) of the medium. While this seems to be roughly true, for example, in reactions I and V, above, it obviously does not hold for 11, I11 and IV, as may be seen by reference to the constants in Table 11. Again, the conversion of ammonium cyanate into urea takes place with a velocity thirty times as great in 90% alcohol as in water,2 although the dissociating power of the latter is several times that of the former. The velocity of saponification of an ester with sodium or potassium hydroxide, moreover, is greatest in the medium of lowest dissociating power, and vice versa.3 Spring4 has suggested that the reason for the slower decomposition of hydrogen peroxide in ether and alcohol solution may be due to the lowering of the surface tension of the solution. Glycerine has been suggested in pharmacy as a preservative for hydrogen peroxide solutions, and it has been assumed Comanducci and Meduri: Gazz. chirn. ital., 48 I, 238 (1918). Walker and Kay: Jour. Chem. SOC.,71,489 (1897). Cajola and Cappellini: Gazz. chim. ital., j o I, 233 (1900). Spring: Zeit. anorg. Chem., IO, 161 (1895).
Catalytic Decovnpositiovt of Hydrogen Peroxide, Etc. 679 TABLE I1
j
Solvent
Water Methyl Alcohol Ethyl Alcohol n-Propyl Alcohol Isobutyl Alcohol Amyl Alcohol Pyridine Glycerol
Dielectric constant
81.1 at 31.2 a t 25.8 a t 2 2 . 2 at 20.0 a t 16.0 a t 12.4 a t 39.1 a t
1
Surface tension
Viscosity
rSo(a)0.00891 a t 2501 0.005564 a t z j o l zoO(b)0.01114 a t 2501 20°(b) 0.01962 at 2501 2 0 ° ( b ) 0.04034 a t z o o ( e ) z o o ( b )0.03923 a t 23.7(f) 2oo(c) 0.00885 a t 2 j 0 3 20°(d) 6.33 a t 2so4 20°(b)
70.6 a t 3oo(g) 23.02 a t aoo(g) 2 0 . 2 0 at 4oo(g) 23.82 a t 16.4(g) 22.72 a t 16O(g) 23.72 at 16.4(g) 38.4 a t 1 7 O ( g ) ,64.j 3 a t 1 8 ~ ( h )
Selected from Landolt and Bornstein. Observers: ( a ) Turner; ( b ) Abegg and Seitz; (c) Schtundt; ( d ) Drude; (e) Gartenmeister; (f)Thorpe and Rodger; (g) Ramsay and Shields; ( h ) Domke.
that this may be due to the viscosity of the s o l ~ t i o n . ~ Several investigators have adopted this point of view and brought forward evidence that the viscosity of the medium may be a factor in slowing down a reaction. Buchbock,b for example, investigated the rate of hydrolysis of carbonyl sulphide in various neutral salt solutions, and pointed out that the greater the viscosity of the solutions, the greater the retarding effect. He later repeated the work,6 revising his results according t o van't Hoff's theory (discussed below) and concluded that the velocity of decomposition of thiocarbonic acid is inversely proportional to the viscosity of the solution. According t o Raschig, monochloramin reacts in ammoniacal solutions as follows : either
+ 2NH3 Nz + 3NHdCl NHaC1 + NH3 = NHZ.NH2.HCl
( I ) 3NHzCl
or (2)
only traces of hydrazine being formed. Dunstan: Jour. Chem. SOC.,85, 817; 87, 1 2 (1905). Hartley: Thomas and Appleby: Ibid., 93, 544 (1908). 3 Jones and Schmidt: Am. Chem. Jour., 42, 37. 4 Bohm and Leyden: D. R. P. 185,597. 6 Buchbock: Zeit. phys. Chem., 23, 123 (1897). 6 Buchbock: Zeit. phys. Chem., 34, 229 (1900). 7 Raschig: Verh. Ges. Naturforscher, 2 I, IZO (1907); Eng. P. 139 (1908). 1
680
I
Van L.Bohnsox
Upon the addition of substances which increase the viscosity of the solution, such as glycerin, starch, dextrin, albumen, casein, and glue, the yield of hydrazine can be increased to 50-807~of the theoretical. Acetone, which lowers the viscosity, increases the velocity of reaction ( I ) . It has been’assumed that this effect is related to the evolution of a gas in one of the reactions. The evidence in favor of a viscosity hypothesis is balanced, however, by equally important evidence to the contrary. Arrheniusl found the velocity of inversion of cane sugar to be accelerated by the addition of salts which both raise and lower the viscosity of the solution; Euler2 pointed out the influence of neutral salts in the hydrolysis of ethyl acetate to be the reverse of that assumed by Buchbock and Reformatsky3 found the velocity of hydrolysis of methyl acetate was not influenced in the slightest by the addition of sufficient quantities of gelatine or agar agar to make the solution jell a t a temperature I O O below the working temperature. Van’t Hoff4 pointed out that the solvent exerts a twofold influence; it may act not only as a true catalyst, accelerating both the direct and the reverse reaction to the same extent, but it may have the further effect of shifting the final equilibrium. This second effect is eliminated from the velocity constant by expressing concentrations of reacting substances in terms of the solubility in the different solvents. In this manner a constant should be obtained which is the same no matter what solvent is used, provided no influences are a t work other than those considered in the development of the theory. Dimroth’s” results confirm this theory, while those of von Halban6 seem to argue against it. Bugarszky,’ studying the reaction between ethyl alcohol and bromine, Arrhenius: Zeit. phys. Chem., 2, 284 (1888); 28, 317 (1899). Euler: Zeit. phys. Chem., 36, 641 (1901). Reformatsky: Zeit. phys. Chem., 7, 34 (1891). .I Van’t Hoff: Lectures (English ed.), I, 2 2 1 . Dimroth: LOC.cit. von Halban: LOC.cit. 7 Bugarszky: Zeit. phys. Chem., 71,705 ( 1 9 1 0 ) .
Catalytic Deconzposition of Hydrogel2 Peroxide, Etc. 68 I obtains by this method nearly identical constants for the solvents carbon tetrachloride and carbon disulphide, but a greatly increased value for benzyl bromide, and concludes, therefore, that the latter takes part in the reaction. Comparatively little exhaustive work has been done on velocity of reaction in mixed solvents. Hemptinne and Bekaertl measured the velocity of the reaction between triethylamine and ethyl bromide in half-and-half mixtures of various solvents. Euler2 points out that they obtained a maximum deviation from the mixture line in those cases which show the greatest deviation of the dielectric constant. Cohen3 studied the velocity of inversion of cane sugar in several water-alcohol mixtures, and Patterson and Forsythi the reaction between iodic and sulphurous acids in aqueous mixtures of methyl alcohol and of acetone. Millar” and Braune6 studied the decomposition of diazoacetic ester in alcohol-water mixtures, finding a minimum velocity a t about I 1% water. The products of this reaction are different with alcohol than with water, bath substances taking part in the reaction. This paper presents the results of a study of the velocity of decomposition of hydrogen peroxide as influenced by the medium in which it takes place; the reaction differs from most of the reactions heretofore studied in that one of the products is a gas. The solvents used were mixtures of water with varying concentrations of one of the following : methyl, ethyl, 12-propyl, isobutyl and amyl alcohols, glycerine, and pyridine. Experimental For determining the velocity of decomposition of hydrogen peroxide, the gasometric method was used. It has been shown by Walton7 to be fully as accurate as the titrimetric method, Hemptinne and Bckaert: Zeit. phys. Ch-m., 28, 2 2 5 (1599). Euler: LOC.cit. Cohen: Zeit. phys. Chem., 28, 1 4 j (rSgg). Patterson and Forsyth: Jour. Chem. SOC.,101, 40 (1912). Mi!lar: Zeit. phys. Chem., 85, 129 (1913). Braune: Ibid., 85, 1 7 0 (1913). Walton: Ibid., 47, I S j (1904).
682
Vaw L.Bohlzson
and can be used in cases where catalyst or solvent would be oxidized by the permanganate used in the latter method. Apparat.us.--The apparatus was essentially that described by Walton, constructed in quadruplicate to permit four determinations a t one time. Reaction flasks of special design were supported by a shaking device in a thermostat, and connected by means of capillary tubing with the tops of jacketed burettes in which the evolved oxygen was collected. The catalyst, held in a glass capsule, was supported in the neck of the reaction flask in such a way that it could be dropped into the reaction mixture a t any desired instant. Thorough stirring was insured by placing glass beads in the bottom of the flask. The evolved gas was collected and measured over water, and a leveling tube connected with the bottom of each burette made possible the rapid adjustment of the water levels. Water from the thermostat was pumped through the burette jackets, thus obviating a variable temperature correction. Ail1 experiments were performed at 25 C. Method-A measured volume (25 cc) of hydrogen peroxide solution in the solvent mixture to be investigated was pipetted into the reaction flask, which previously had been thoroughly cleaned, steamed, and dried. Exactly I cc of the catalyst solution was then placed in the glass capsule, which,was supported in the neck of the reaction flask. The temperature in thermostat and burette jackets being constant, and the burettes and leveling tubes having been filled with water, the flask was placed in the thermostat and connected with a gas burette. An initial reading of volume was made, the catalyst dropped into the reaction mixture, and the shaker started. After waiting until from 5 to 8 cc of gas had been evolved, in order to be sure of a regular evolution of gas, volume readings were taken at measured intervals until about three-fourths or seven-eighths of the hydrogen peroxide had been decomposed. A final reading was taken when there was no appreciable change in volume after a half-hour interval. Table I11 shows a typical data and result sheet. Complete data for the duplicates are omitted.
Catalytic Decomposition of Hydrogeiz Peroxide, Etc. 683
TABLE I11 Sodium Iodide Catalysis of Hydrogen Peroxide in Water I . NaI = 0.02053 mols per liter Bar. = 738 mm Temp. = 25' C a = 36.28 . 1
t (Minutes) 0
5 9 I7 27 38
60 79 Final reading 3 hrs later I cc 0 2 ovei water a t 732 mm 5nd 2 5 ' C = 1.225; mg Ks= 0.4343 X X 104 =128.;
Observed volume (cc)
Milligrams of dry 0 2
4.00 8.05 15.75 20.40 24.00 28.65 30.83
4.90 9.87 13.36 19.30 25.00 29.42 35.12 37 * 79
33.60
41.18
10.90
Y
4.97 8.46 14.40
a-x
31.31 27.82
.20. IO
21.88 16. I8
24.52 30.22 32.89
11.76 6.06 3.39
36.28
-
o.4343K
0.01279 0.01281 0.01292 0.01298 0.01287 0.01295 0.01302
-
0 . 0 1290 Average Duplicates, 0.01285 -0.01286 Mean, 0.01287
CaZcUZations.-Waltonl has shown that the sodium iodide catalysis of hydrogen peroxide in aqueous solution is a monomolecular reaction. The velocity constant can, therefore, be calculated from the well known expression I a 0.4343 K =
t'
1%
where x represents the amount of substance decomposed (i. e., the amount of oxygen evolved) in the time t, and a represents the concentration of hydrogen peroxide at t = 0 ; a may, therefore, be represented by the total amount of oxygen evolved after the time selected as the starting point ( t = 0 ) . As the table shows, the weight of dry oxygen in milligrams was calculated for each reading, in order that the velocity constants obtained may be independent of the barometer pressure. Incase the barometer vaned before the ob1
Walton: LOC.cit.
.
.
Vaiz L. Bohnsoi?
684
taining of the final reading, appropriate correction was made. In order to conserve space, the remaining results of this research have been condensed to the form shown in Table IV, where K, represents 0.4343 K X 104.
TABLE IV Sodium Iodide Catalysis of Hydrogen Peroxide in Water 2 . NaI = 0.1026 mols per liter t
a
3 5 7 9 13
T
=
46.85 16.70 24.29
638 635 638 643 622
30. I I
34.51 40.51 .
635 Duplicate, 605 Mean, 620
Materials.-In all cases the purest materials obtainable were used. The water was distilled from a large copper still, over permanganate, and the middle portion of the distillate was used. The hydrogen peroxide was Merck’s “Perhydrol,” sometimes from glass bottles, sometimes from paraffin containers. Samples from two such sources gave constants which checked within the limits of error, and checked as well against a constant obtained from a sample of the middle portion of a lot distilled in vacuo. A measured amount of perhydrol (generally about 1.5 cc) was diluted to IQO cc with the solvent t o be used. 2 5 cc of such a solution usually gave a total oxygen evolution of about 40 to 45 cc or approximately 50 to 55 mg. The initial concentration was, therefore, approximately 0 .I 2 to 0.15 mols per liter. The sodium iodide had been recrystallized several times and completely dried in an air bath. Standard catalyst solutions were made up by weighing out the dry salt and making up the solution to a known volume. The concentration of catalyst in the reaction flask is expressed in mols per liter.
Catalytic Decomposition of Hydrogen Peroxide, Etc. 685 The alcohols used as solvents were distilled, dried with anhydrous copper sulphate, and redistilled, and in some cases further dried with sodium and again redistilled. The glycerine
25
50 Fig.
75
I
used was Merck’s, sp. gr. 1.26, British Pharmacopoeia. Samples of pyridine from two sources were used, one which had been prepared by the writer for a previous research; and one prepared by Professor Koenig in this laboratory. Both samples had been dried repeatedly over potassium hydroxide and re-
Van L. Bohnson
686
distilled, and gave velocity constants that checked each other well within the limit of error. All the solvents used were made up by mixing known volumes of the components; the composition of the mixture in the reaction flask was then calculated in percentage of alcohol by weight, in which all the results are expressed. 3. Results Velocity of Reaction in Water-Alcohol Mixtures.-Tables V, VI, VI1 and VI11 give the results of experiments using as solvents binary mixtures of water with methyl alcohol, ethyl alcohol, and n-propyl alcohol. The results are graphically shown in Fig. I . TABLE V Catalysis in Mixtures of h4ethyl Alcohol with Water NaI = 0.02053 mols per liter C = yoAlcohol by weight KS
a = 50.02
15 23 34 47 64 87
10.78 15.54 2 1 . I4 26.59 32.23 37.75
70.I 70.3 70.2 70. I 70.2 70.2
70.2
Duplicate, 71.6 Mean. 70.9
a= 47.46 9 23
45 64 94 =I7
2.97 7.24
13.10
17.64 23.65 27.47
23 36 53 73
4.91 9.53 13.79 18.48 22.60
a = 49.40 IO
31.3 31.2 31.5 31.9
22
2.83 5.88
41 60 91 114
14.24 19.92 23.53
32. I
31.5
Duplicate, 3 1 . 2 Mean, 31.3
55.2
11
55.2
23 39 59 83
55.1 55.3 54.7
-
55.1 Duplicate, 56. I Mean, 55.6
31.2
10.32
I
X
KS
-
a= 51.I O
a = 37.62 I1
t
-
25.6 25 .o 24.8 26.3 24.6 24.6
-
25. I Duplicate, 24.4 Mean , 24.7
107 128
5.44 10.67 16.83 23 *
12
29.28 34.08 37.50
44.4 44.2 44.5 44.3 44.5 44.6 44.9
44.5 Duplicate, 4 4 . 1 Mean, 4 4 . 3
Catalytic Decowposition o f Hydrogen Psroxide, Etc'. 687
TABLE VI Catalysis in Mixtures of Ethyl .41cohol with Water NaT = 0.02053 mols per liter C = yo alcohol by weight
Ks
a= 49.80 6 4.97 14 10.88 34 22.4 47 28.12 65 34.22 81 38.31
76.I 76.5 76.3 76.8 77.6 78.6
76.9 78.5 77.7
a = 47.22
7.51 44.2 I7 27 11.36 44.3 43 16.83 44.5 60 21.63 44.3 92 28.87 44.6 126 34.35 44.8
t
a= 16 27 43 60 ' 92 126
j
r
Ks
-
48.16 8.75 54.3 13.87 54.6 20.21 54.9 25.73 55.2 33.45 55.9 38.72 56.I 55.2 52.8 54.0
a= 39.95 12 4.72 45.5 28 1 0 . 1 0 45.1 50 16.22 45.4 74 21 * 53 45.5
a= 7 15 26 36 53 74
52.90
50.2 50.8 .13.61 49.7 18.01 50.2 24.26 50.3 30.56 50.6 4.11
8.51
6. c= 93.4% a= 48.69 6 2.91 =7 7.79 29 12.48 5 1 19.77 SI 27.45
1
44.6 44.5 44.4 44.4 44.5
45.4 44.5
44.5 Duplicate, 45.3 Mean, 44.9
Duplicates, 44.2
46.5 45.6
Mean, 45.2
Van L. Bohnson
6 88
TABLE VI1 Catalysis in Mixtures of n-Propyl Alcohol with Water N a I = 0.0~053mols per liter C = yo alcohol by weight
l x I
Ks 2.
a = 48.76
6
94.2 93.4 93 2 93.1 93 .'4 93.4 94.
5.95
11
10.27
18 25 34 43 56
15.63 20.23 25.29 29.42 34.26
'
__
93.5
81.4 80. j
28 37
48 63 71
C=19.8% a= 48.22 4 3.37 12 9.4 20 14.6 30 20.15 40 24.76 58 31.34
80.3 80.4 80.2 80.3 80.7
19.38 23.80 28.29 33.16 35.30
80.6 Ihplicate, 80.2 Mean, 1 8 0 . 4
Duplicate, Mean, 5 . C=61.0% a= 47.18 4 3.64 12 10.18 20 15.61 30 21.41 40 25.98 58 32.63 Duplicate, Mean,,
78.7 78.5 78.3 78.3 78.2 78.6
78.4 77.8 78. I 87.2 88.0 87.2 87.5 86.9 88. I 87.5 88.0 87.7
Duplicate, Mean, e
142.5 142. 141.6 141.1
140.9 140.6 '40.7 '47.3 142.2 141.7
78.6 77.3 77.1 77.2 77.5 78.
77.6 Duplicate, 7 5 . 9 Mean, 76.7 6. c = 7 6 . 5 7 0 a= 43.55 4 3.91 102.1 IO 8 . 9 9 100.4 , 17 1 4 . 1 2 100. I 29 2 1 . 2 0 99.6 51 30.09 100.0
-
~
7 . C=93.9% a= 47.77 5 7.23 g 12.18 I j 18.48 23 25.14 28 28.52 ' 35 32.38 47 37.35
3. C = 2 8 . 3 % a = 47.52 6 4.86 12 9.08 20 14.11 30 1 9 . 5 1 59 30.74 81 36.20
Duplicate, Mean,
100.4 IO1 . g 101. I
Catalytic Decomposition. of Hydrogen Peroxide, Etc. 689 TABLE VIII-SUMMARY Catalysis in Water-Alcohol Mixtures
C
1
Methyl alcohol
c 0
ye alcohol by weight Ethyl alcohol C
Ks
24.0 34.5 46.0 70.8 93.5
=
n-Propyl alcohol
K
C
*
Ks
S
128.7 70.9 55.6 44.3 31.3 24.7
0
19.3 32.6 40.0 56. I 61 .o 93.4
128.7 77.7 54.0 50.3 44.9 45.4 45 2
0
13.0 19.8 28.3 41.8 61 .o 76.5 93.9
'
,
128.7 93.5 78. I 76.7 80.4 87.7 101. I
141.7
In Tables IX and X are given the results of experiments using as solvents a few aqueous mixtures of isobutyl and amyl alcohols, which are only partially miscible with water. I n no case were oxidation products of the alcohols detectable after the completion of the reaction. TABLE IX Catalysis in Isobutyl Alcohol NaI = 0.02053 .mols per liter C = 7 0 alcohol by weight
l x a = 44.56 5.76 5 13.46 I3 21.05 23 28.50 37 32.96 49 Duplicate, Mean,
I x
KS
a = 40.33 I20 I21 I20 I20 I 18
I20 I21
120.5
2
5 7 11
16
6.77 14.89 19.20 25.66 31.14
Duplicate, Mean,
399 400 401 399 399 399 395 397
Van L. Bohnson
690
TABLE: X Catalysis in Amyl Alcohol
NaT
=
0.0205g mols per liter 93.5% alcohol
Velocity of the Reactiovt in Glycerol-Water Mixtures.-The results of experiments using glycerol-water mixtures as solvents - are given in Table XI, and represented graphically-in-Fig.-2.
Fig.
2
Catalytic Decomposition of Hydrogen Peroxide, Etc. 69 I TABLE XI Catalysis in Mixtures of Glycerol with Water NaI = 0.02053 mols per liter C = yo glycerol by weight
l x a= 5 9 I9
4. ~ = 5 1 . i % a= 45.53 3 6.86 7 14.40 I 3 23.17 I7 27.51 24 33.31 34 38.49
Duplicate, Mean,
a= 3 7 13 I9 28
48.18 8.73 14.47 25.25
41.31 5.30 11.44 18.78 24,43 30.48
a= 51.26
a= 29.04
5
3 7
6.57 12.81
'11
17.02
IO
I4 I9
14.50 24.72 30.55 35.88
I99 201 202
204 208
_ .
283
/
23 7 227
232
SUMMARY Ks 0
13.9 28.2 41.3 51.8 75.6 95.7
Velocity of the Reaction in Pyridine-Water Mixtures.The concentration of catalyst used in the previous experiments (0.02053 mols NaI per liter) was found to accelerate the reaction in a concentrated pyridine solution t o a compara-
.
I
692
V a n L..Bohnsorz
tively small extent, as shown in Expt, 9, Table XII. A catalyst of five times that concentration was, therefore, used in determining the effect of pyridine upon the reaction. The results of these experiments are given in Table XI1 and graphically represented by Fig. 3.
Fig. 3
C 0
33.5 57.3 67. I 76.8 81.7 88. j 90.2 94.7
K, 620 278 I20
72
48 39.3 35 38.6 71.2
Catalytic Decomposition of Hydrogen Peroxide, Etc. 693
TABLE XI1 Catalysis in Water-Pyridine Mixtures
NaI = 0.1026 mols per liter C = yopyridine by weight I
!
a = 25.83 4 9
5
11.22
IO
a= 45.56 5.21 10.46 23 39 15.81 21.21
27.77 32.31
III
;.
Duplicate, Mean, C=90.2,% 14 ,?O
51
-., 11
100
127
5.57
1 1 . I5
17.28 23.41 27.75 31.83
Duplicate, Mean,
_ _
I5 23 36
5.56 10.39 14.63 20.24
I22 I22 I22 I22
27.20
123 _ _
I1
57 85
KS
a = 42.56
5.71
Duplicate, Mean,
I
51 73 102
17.67 23.18 28.93
Duplicate, Mean,
13.18 19.57 2'6.40 32.22
20
33 51 73
0.9
117 I20
3.2
.o. I .o . 3 .o . 4 .o . 7 11. I
-
Duplicate, $3. I Mean, i9.3
4.16 9.26 14.88 19.05 24.88 29.13
a = 46.90 5.86 8
I22
-0.5
7 17 3n 42 63 84
,
2.0
a = 48.34 5.39
15
30 62 83 108
10.24 18.76 23.31 27.89
1.2
4.5 4.4 4.4 4.6
__
4.4 Duplicate, 5 . 7 Mean, 5 . 0
9. N a I = 0.02053 rnds/liter a= 47.13 '0.5 1.71 25 i9.8 49 3.39 6.04 '0.I 90 70.0 7.61 120 '2.4 9.41 151 72 .o
6.4 6.6 6.6 6.4 5.7
-
_ .
10.8
7.1 6.7
71
.6
'1.2
6.3
Van L . Bohizsoia
694 I
4. Discussion of Results
By extrapolation, the following constants (Ks = K X 104) are obtained for the velocity of the reaction in the pure solvents, using in each the same concentration of sodium iodide as catalyst. Water Methyl Alcohol Ethyl Alcohol %-PropylAlcohol Isobutyl Alcohol Amyl Alcohol Glycerol Pyridine
128.7 23 45 164 above 397 above 537 376
7
Pyridine, the solvent having the lowest dielectric constant, appears to exert the greatest retarding effect on the reaction, while the latter is accelerated most in amyl alcohol, the dielectric constant of which is but little greater than that of pyridine. If the alcohols alone are considered, the velocity of reaction decreases as the dielectric constant of the solvent increases, a result entirely a t variance with the usual assumptions. Hence there appears to be no direct relation between the velocity of this reaction and the dielectric constant of the pure solvent. Drudel has shown that the dielectric constant of mixtures of water and methyl alcohol varies directly with the concentration of the alcohol. We have found no such direct proportionality in aqueous alcohol mixtures between the reaction velocity and the concentration of the alcohol. It is obvious, therefore, that if there is a relation between the dielectric constant of a medium and the velocity of a reaction in that medium, it is in most cases entirely offset by some other effect. Again, of all the solvents used, pyridine has the lowest viscosity, and in it the reaction goes slowest, while it is much faster in the medium having the greatest viscosity, pure glycerol. Considering the alcohols (with the exception of I
1
Drude: Zeit. phys. Chem.,
23, 267 (1897).
Catalytic Decomposition of Hydrogeiz Peroxide, Etc. 695
IO1
5c
a Fig. 4
,
Van L. Bohnson
696
amyl alcohol) as a group, the greater the viscosity, the greater is the velocity constant. This result is contrary to expectations, if the viscosity of 'the medium is the major influence on the velocity of the reaction, and we must, therefore, conclude,% as in the case of the dielectric constant, that whatever relation there may be between the viscosity of the solvent and the velocity of a reaction is quite overbalanced by another and greater influence. Such an influence might be due in the case of pyridine, for example, to the formation of an addition compound of the pyridine with the hydrogen peroxide, since the former is well known to be basic while the latter has been shown to have weak acidic pr0perties.l In Table XIV are given the viscosities and fluidities of aqueous mixtures of all the solvents used. The fluidities are
Glycerol
0
+
-1.0089 111.1 0.013; 25 30 0.01j 3 40 0.0157 50 0.0153 60 0.0139 7 0 0.0123 75 20
72.90.0170 58.8 0.0183 54.6 0.0135 74.1
1 . 0 2 0 0 50.0 62.5 9 54.4 49.c 1.0615 16.3 45.5 46.1 1.3203 3.1 80 0 . 0 1 0 2 98.00.0174 57.5 0.0244 41.00.0189 52.5 go 0.0077 129.7 0.0143 69.90.0222 45.0o.orgg 74.1 0.16 IOO 0.0056 178.50.0111 90.1 0.0196 51.00.0088 113.t 5.330
65.3 0.0213 63.7 0.0234 65.3 0.0234 71.90.0224 81.3 0.0199
46.9 0.0224 42.7 0.0253 42.7 0.0266 44.60.0270 50.2 0.0259
44.6 0.0160 39.5 0.0184 37.6 0.0204 37.00.0219 38.60.0217
The data in this table were obtained by interpolating the results of: Dunstan: Jour. Chem. SOC., 85, 817 (1904); 87, I Z (1905) (alcohols); Hartley, Thomas and Appleby: Jour. Chem. SOC.,93, 544 (1908) (pyridine); Jotles and Schmidt: Am. Chem. Jour., 42, 37 (1909) (glycerol).
Catalytic Decomposition of Hydrogen Pcroxide, Etc. 697 It will be noted by comparing these curves with Pigs. I to 3 that in no case is the solution of minimum fluidity the one in which there is a corresponding minimum in the reaction velocity. It has been suggested by Dr. Bancroftl that the reaction velocity in the mixture may be a resultant of the velocities of two different reactions, one in each of the solvents, with the effect due to viscosity superimposed upon them. I n order to further investigate the possible effect of the viscosity of the medium, the reaction velocity was determined in a gelatine solution of marked viscosity (of such a concentration that it solidified a t a temperature only a few degrees below that of the thermostat. The results obtained are given in Table XIII. TABLE XIIT Catalysis in Gelatine Solution NaI = 0.02053 mols per liter K, for water = 128.7 t
a = 5
I3 22
34 62
44.01 4.53 10.79 16.71 22 ,99 32.65
94.3 94.0 94.2 94.4 94.5
94.3
The velocity in this glycerine solution is shown to be about 28% slower than in water with the same concentration of catalyst. This result indicates that the viscosity of the medium does exert an effect upon the decomposition of hydrogen peroxide, in spite of the fact that it exerts no effect on the hydrolysis of methyl acetate.? The two reactions differ, however, in that one of the products of the former is a gas, and in such a case i t is to be expected that the influence of viscosity would be greater. In the case of solvents other than 1
In a personal letter to Professor Walton. Reformatsky: LOC.cit.
Van L.Bohizson
698
gelatine solution, it is apparent that the effect of viscosity is masked by one of greater influence. An inspection of the surface tension data in Table I1 will reveal that there is no important relation between that property of the solvent and the velocity constant, the values of surfaLe tensions of the alcohols being quite similar, while the velocity constants in the sevQal alcohols differ widely. Some measurements were made of the surface tensions of the various pyridine solutions used in these experiments, after the hydrogen peroxide was completely decomposed. The apparatus used was of the torsion wire type which records the force necessary to detach a platinum ring from the surface of the solution. The results follow: 3 3 . 5 57.3 67.1 76.8 81.7 8 8 . 5 90.2 94.7 55 53 52 50 49 48 47 47
Percent pyridine Dial reading
The slight difference in the surface tension of concentrated and dilute pyridine solutions, and the absence of a maximum or minimum, make it clear that the relation between surface tension of the solvent and reaction velocity is negligible. Waltonl has shown that the velocity of decomposition of hydrogen peroxide by sodium iodide in water solution is directly proportional to the concentration of the iodide ion. It is, therefore, of interest to consider the conductivity of sodium iodide in such aqueous mixtures as are used in these experiments. The following values were obtained by Jones and his co-workers. ?(ABLE
NaI
= 0.0312
XV mols per liter Molecular Conductivity
Percent alcohol Methyl alcohol2
I
Ethyl alcohol3
I
0
25 50 75 IO0 1
106
70.6 57.2 56.4 72 .o
I 06
49 35.5 30.7 26.8
Walton: LOC.cit. Jones: Carnegie Institution Publication No. 80, 47. Wightman, Davis, Holmes and Jones: Jour. Chim. phys..
12,3 8 j
(1914).
Catalytic Decompositiovt of Hydrogen Peroxide, Etc. 699 These values are shown graphically in Fig. 5 . It will be noted a t once that the conductivities of sodium iodide in water mixtures of these alcohols are not in accord with an as- 100 sumption that the velocities of reaction in the same solvents are due entirely to the concentration of iodide ion. The van’t Hoff t h e o r y 50 heretofore mentioned is obviously not applicable to homogeneous systems of mixed solvents such as the one studied, o 0 25 50 75 the solubility of the catalyst or Fig. 5 hydrogen peroxide in the separate solvents not being a factor. The final equilibrium is the same in all of these experiments, complete decomposition of the hydrogen peroxide being attained. The foregoing considerations must lead to the conclusion that there is no outstanding relation between the velocity of decomposition of hydrogen peroxide as influenced by sodium iodide and the dielectric constant, viscosity, or surface tension of the solvent. Whatever influence these properties may have, if any, must be counterbalanced by other effects. It is noticeable that in the case of the alcohols, the reaction proceeds fastest in the solvent of highest molecular weight. It is known that alcohols of higher molecular weight are less associated than their lower homologues and it is possible that this property is a deciding factor in influencing the velocity of the reaction. In general, it may be said that the solvent exerts a specific effect on the velocity of a reaction taking place in it, this effect being probably the resultant of several effects, among others those due to association of the solvent, its viscosity, its surface tension, the dissociation of the catalyst, and by no means least, a chemical reaction between solvent and dis solved substance.
-
700
.
Van L . Bohnson
The author desires to take this opportunity for expressing his appreciation to Professor J. H. Walton, a t whose suggestion this research was undertaken, and under whose direction it was carried out, for his kindly interest and guidance during its progress. 5 . Summary I . The velocity of the catalytic decomposition of hydrogen peroxide by sodium iodide has been measured in different aqueous mixtures of methyl, ethyl, normal propyl, isobutyl and amyl alcohols, glycerol, and pyridine. The solvents apparently do not undergo oxidation ; the decomposition of the hydrogen peroxide is complete. 2. The relakive velocity of reaction in the pure solvents may be expressed by, the following constants : methyl alcohol, 23 ; ethyl alcohol, 45; n-propyl alcohol, 164; isobutyl alcohol, 397; amyl alcohol, 537; glycerol, 376; pyridine, 7; water, I 28.7; gelatine solution (of marked viscosity), 94.3. There is apparently no relation between the velocity of the reaction and the dielectric constant, viscosity, or surface tension of the pure solvent. 3. Curves for the reaction velocity in aqueous mixtures in some cases exhibit minima', but these do not correspond with the minima in the fluidity curves. No relation has been found to hold between the reaction velocity and any physical property of the mixed solvent. 4. It appears that the solvent exerts a specific effect on the velocity of a reaction, this effect being probably the resultant of a number of other effects due to association of the solvent, its viscosity and surface tension, the dissociation of the catalyst, and a possible reaction between solvent and dissolved substances. Laboratory of General Chemistry Universzty of Wisconsin Aiadzson, July 1920
