T H E CATALYTIC DECOMPOSITION O F SODICM HYPOCHLORITE SOLCTIOW3*
I. Mechanism of the Reaction
____
BY JOHN R . LEWIS
Cntil recently no systematic study of the catalytic decomposition of sodium hypochlorite solutions had been carried out. I n 1923 and 1926, however, articles by Howell' and Chirnoaga2 were published in which appeared their results using cobalt peroxide as catalyst. Howell assumed that the decomposition rate follows the unimolecular law; but pointed out that the constant (K) increases as the reaction proceeds. He showed that the rate increases in the presence of sodium ions while hydroxyl ions have the opposite effect. Chirnoaga, on the other hand, obtained more concordant results by using the equation - dc,'dt = Kc"" (where c is the concentration, t the time in minutes, K and n constants). This is especially true in those cases where the catalyst was very active. I n other cases, i. e. where the catalyst was less active the unimolecular law held. I n general the work of the above mentioned investigators has been confirmed in this laboratory but it appeared that the mechanism of the decomposition can be explained more simply by using the equation dc/dt = K ; for when the volume of the evolved oxygen, in cubic centimeters, was plotted against the time, in minutes, a straight line was obtained. This linear relationship held in most cases from ten to twenty pCrcent of the total reaction. These results suggested a mechanism similar to that found by Xrinstrong3 and by Armstrong and Hilditch4 for the hydrolysis of sugar solutions by enzymes and for the hydrogenation of certain unsaturated organic esters using nickel as catalyst. In the pape9 on the action of enzymes on sugar it was shown that equal quantities of sugar were hydrolyzed in equal time intervals. A similar relationship was obtained in the hydrogenation experiments. While the linear portion of the curve covered approximately ten percent of the total reaction in the case of the sugar hydrolysis, the linear portion of the curve in the hydrogenations held fairly well from fifty t o eighty percent of the total reaction. Among other examples of reactions of this type may be mentioned the results of Bredig and von Berneck6 who found that dc, dt = I< for the decomposition of hydrogen peroxide by colloidal platinum, and also the experiments of Hinshelwood' on the decomposition of certain gases at the sur*Contribution from the Laboratories of General Chemistry of the Universitv 01 Kisconsin. 'Howell: Proc. Roy. SOC.,104-4, 134 (1923). * Chirnoaaa: J. Chem. SOC..1926. 168.1. 3AIIllstroig: Proc. Roy. Sdc., 73,' go0-(1904). Armstrong and Hilditch: Proc. Roy. Soc., 98.1, 27 (1920). Armstrong: Proc. Roy. Roc., 73, 500 (1904). Bredig and von Berneck: 2. hysik. Chem., 31,266 (1899). Hinshelwood: "Kinetics of C\emical Change in Gaseous Systems," 148 irg26)
2 44
JOIIS R. LEWIS
faces of hot wires. Reactions of this type are usually spoken of as zero order reactions. The conditions necessary for the realization of zero order reactions may be stated as follows I . Comparatively high concentration of the substance undergoing change 2. The catalyst should maintain its activity 3. There should be small quantities of catalyst in order that the active mass, i. e. the adsorbed reactant on the caralyst is small compared with the bulk of reactant. 4. The decomposition of the reactant-catalyst complex must be the slonr reaction, thus insuring at all times a catalyst surface completely covered with reactant. Experimental Preparation o j Materials I. Sodium hypochlorite solutions were prepared as follows: Approui mately normal sodium hydroside, prepared by the action of metallic sodiuni on distilled water, was treated with pure chlorine. The gas was allowed to bubble slowly through the sodium hydroxide solution, the container of which was kept cool by immersion in an ice bath. The progress of the reaction nae determined from time to time by pipeting out two cubic centimeters of the solution, destroying the sodium hypochlorite by means of neutralized hydrogen peroxide, and titrating the excess alkali with standard acid. The reaction was stopped when the hypochlorite solution contained but a slight excess of free sodium hydroxide. The sodium hypochlorite was determined by adding a known quantity of the solution to an acidified potassiuni iodide solution, and titrating the freed iodine with standardized sodium thiosulfate. The initial concentrations of the hypochlorite solutions prepared are as fOl1OM s : Solution I 2
3 4
Concentration of XaC10 2 9 . 7 4 gr. per L
Concentration of SaOH 0 . 1 3 Normal
3 7 . 76 gr. j 6 . 4 8 gr. 4 6 . j j gr.
” ”
0.03
”
” ”
0.11
”
” ”
0.032
”
The catalysts solutions were made from recrystallized salts; standard 2. stock solutions of copper sulfate, ferric sulfate and cobalt sulfate being made. The cobalt peroxide, copper oxide, and iron oxide suspensions were made according t3 the method of Howell’ and Chirnoaga2 except that the preeipitated oxides were washed by decantation in every case rather than by filtering. This process mis continued until the wash water was practically free from alkali, requiridg in some cases three or four weeks. 1 Howell: Proc. Roy. Soc., 104A, 134 (1923) ‘Chirnoaga: J. Chem. SOC., 1926, 1693.
DECOMPOSITION O F SODIC&: HYPOCHLORITE SOLUTIONS
245
Apparatus and Method of Experimentation The apparatus used for the first experiments was essentially the same as that devised and used by Walton' and his students for their studies in chemical kinetics. Twenty-five cubic centimeters of the sodium hypochlorite and the catalysts (usually I c.c.) were placed in especially designed reaction flasks supported by a shaking device in a thermostat. The catalyst was either placed in a glass capsul held in the neck of the flask until the desired moment, when it was allowed to fall into the hypochlorite solution; or pipeted directly into the reaction flask through a short side-neck which was immediately closed. At the instant the catalyst was mixed with the hypochlorite. a stopwatch and the shaker were started Readings of the evolved oxygen, collected in water-jacketed burets, were taken at suitable intervals of time until the reaction was complete. Calculations The data for tabulation in the tables which follow were obtained by the following equations. Equation ( I ) is the so-called zero-order equation, equation ( 2 ) the ordinary unimolecular expression, and (3) which is one form of the well known Freundlich Adsorption equation used by Chirnoaga' in his hypochlorite studies. (1)
where x is the C . C .
K = dx/dt or more simply K = x t of oxygen evolved in time (t) expressed in minutes. Ki
=
2.303 ~
t
loge,
co - Ct
where t is the time in minutes, C, is t.he concentration of hypochlorite espressed in cubic centimeters of oxygen a t zero time, and C, is the concentration after time t .
K?
(3)
-
ctT-
I
N
t
h- -
= -
where C, and C, and t are the same as in
(2)
i
I
and N is a constant.
Preliminary Results The effect of the rate of shaking of the reaction mixture on the rate of oxygen evolution was determined. The following table gives the results.
TABLE I Experiment Sumber
1
Sumber of shakes per minute
I
400
of oxygen minutes 4.3
2
630
6.6
3 4
800
8.1j 8.I
I 190
Walton: Z. physik. Chern., 47, 185 (1904).
C.C.
after
Ii
20
0 . 2 I j
0.33 0.40; 0.40j
,
246
JOHK R. LEWIS
I t \Till be observed that the values for the last two agree well, showing that rates of 800 or above give concordant results. All subsequent experiments were carried out with the rate of shaking at 800 or more per minute. I t was found that glass or Bakelite beads placed in the reaction flasks to prevent supersaturation, retarded the reaction rate. This was caused in the
FIG. I Decomposition of sodium hypochlorite solutions using various catalysts Scale A = I . Copper oxide 2. Copper sulfate 3 and 4. Mixtures of copper and ferric oxides 5 . Cobalt sulfate 6. Cobalt peroxide a t 4 5 T . Scale B = 7 . Mixed copper and ferric oxides 8. Cobalt peroxide a t 35OC.
case of the Bakelite by a chemical reaction between the hypochlorite and the Bakelite. The glass beads cut down the rate of hypochlorite decomposition by holding the catalyst particles mechanically on the bead surfaces.
Results The results obtained using the above equations as a basis for calculating the rate of decomposition are given in the following tables. I n order to conserve space, duplicate results (obtained in all cases) are not given. Tables I1 and I11 are for cobalt peroxide at 45' C and 35OC,while Table IV is for cobalt sulfate solution. Tables V and VI give the results obtained for
DECOMPOSITION O F SODIUM HYPOCHLORITE SOLUTIONS
247
copper sulfate solution and for precipitated copper oxide. Finally in Tables VII, VI11 and IX are given the results for mixed oxides of iron and copper. In Tables IV, V, VII, VI11 and I X the catalysts were added in the solution form. Due to the alkalinity of the hypochlorite solutions, the catalyst solutions reacted, forming the desired oxides. Data for the decomposition of the hypochlorite using precipitated iron oxide only are not given since the decomposition rate is so slow that it may be considered as without effect In the tables (t), (x), (K), (K1), and (K2) have the same significance as in the equations. The data given in the tables are also presented graphically in Fig. I .
TABLE I1 C, = 160 Catalyst: I C.C. cobalt peroxide i / n = 0.8 peroxide (0.2 I 5 gr. per L) X K K t KI 7.05 ' 705 ,00462 ,0125 IO ,0128 10.06 '755 ,0046I 14 13.2 '734 ,00486 I8 ,0133 18.5 .770 .005I2 24 .0'39 . or40 27.6 ,745 ,005'3 37 . or40 32.8 '745 ,00521 44 ,0142 37.1 ,743 ,0053I 50 41.2 '747 ,00543 .0145 55 66 48.0 ,728 ,00543 ,0149 52.0 '723 .00546 ,0145 72 ,0146 59.8 ' 705 ,00552 85 62.6 ,696 .00553 90 ,0145 66.3 ,697 ,00563 ,0148 95 IO0 69.0 ,696 ,00565 ,0147 ,0146 104 70.8 ,681 .00564 TABLE I11 Temp. 35°C. C, = 160. Catalyst: I C.C. cobalt peroxide I /n = 0.8 X K t K, KP IO .oor8z 2.95 '295 .0055 20 .322 6.46 .00207 ,00587 IO.56 ,0022 5 ,00616 ,351 30 ,362 .00240 ,00624 40 14.5 18.2 .00244 .00620 ,364 50 22.8 62 ,368 ,00252 ,00685 26.03 ,00663 , 0 0 250 71 .367 ,368 82 30.2 , 0 0 2 58 ,00654 .00257 ,00678 34.6 95 ,354 108 .00260 ,00686 39.4 ,365 ,00263 ,00697 42.5 '363 "7 48.6 ,00266 ,00692 136 ,357 50.I .ooz66 ,007 I 2 142 .353 .ooz70 ,00709 I55 54.9 ,354 Temp. 45'C.
J O H S R. LEWIS
2-18
TABLE IV C, = 112. I C.C. CoS04solution containing 1 . 7 2 gr. per L.
Temp. 3 5 T . t I4
I / n = 0.8 lir
Y
i.13 IO.7
. O IIO
,0135
33
15.35 18.1
41
22.
5' 55 61 65
27.j
,0139
29.8 32.61 34.6
,0140 '
io
37.0
,528
20 28
,0127
,0136 ,0138
j
,535
,00568
,0141
.533
. 0 0 j68 .ooj;2
,0146
,0141
TABLE Temp. 3 j ° C ' .
('o
2.2
20
3
30
5.46 7.5
'
I 3
13.1 15.3 18.9
L1-
12;
148 I94
11.5
0.8
I
