KINETICS AND MECHANISM OF T H E THERMAL DECOMPOSITION OF SODIUM DITHIONITE
IN AQUEOUS SOLUTION R O B E R T G. R I N K E R , S C O T T L Y N N , ' D A V l D M . M A S O N , * A N D W I L L I A M H . C O R C O R A N California Ivstitute o,f Technology, Pasadena, Calif.
The thermal decomposition of sodium dithionite was studied in buffered and unbuffered aqueous solutions from 4.8 to 7.0 pH and at dithionite concentrations from 5.5 to 11.5 X 1OP3M. From 60" to 80" C. the reaction rate at zero time was found to be half order with respect to hydrogen ion and three-halves order with respect to the dithionite ion. Application of the Arrhenius equation to the rate of disappearance of dithionite gave an activation energy of 12 kcal. per mole and a frequency factor of 1.3 X 1 O8 liters/(mole) (sec.). The free radicals SOZ-. and HSOz. appeared to be part of the mechanism. After time zero; there was an induction period in which the concentration of dithionite decreased slowly; then there was an abrupt increase in rate of decomposition which was exponential with respect to time, suggesting an autocatalytic or degenerate branching chain mechanism.
N a & 0 4 . is a powerful reducing agent and the manufacture of various chemicals. Although first prepared in solution by Schonbein (7) in 1852, it was not until 1869 that Schutzenberger (8)obtained crystalline sodium dithionite and found its composition to be Na2S204. In 1911, Jellinek (3) extensively studied the properties and reactions of dithionite in aqueous solution and concluded that at 60" C. its decomposition was second order relative to its concentration. Subsequent bvork by Kolthoff and Miller (4) in polarographic studies of aqueous solutions of dithionite showed at the presence of decomposition intermediates such as S O S - ~ p H < 5.5 and the SO2--.radical at 6.5 > p H > 5.5. Eermbk's ( 7 ) polarographic studies established further the complex system of equilibria betMeen sz04C2and S o n , S 0 3 C 7 , HSOB-, SO?-%..SO2-* and S205-2. Fischer and coworkers (2) in chronopotentiometric investigations studied specifically the monomerization of S?04-' to give the SOr-. radical. Significant ivork has thus been done in establishing the nature of several intermediates in the reactions of sodium dithionite. but detailed information regarding their roles in determining the rate and mechanism of the decomposition reaction is not yet complete. 'The present work is a contribution to those details for a small p H range between 4.8 and 7.0 and is concerned with the establishment of a decomposition mechanism of aqueous dithionite for the induction period near time zero. ODIUM DITHIONITE,
S that has considerable use in vat dyeing, bleaching,
Experimental
T h e equipment used in the rate studies of the therma decomposition of sodium dithionite has been described in detail (6). '4 glass reactor with a volume of approximately 600 ml. \vas used. An effort was made to eliminate oxygen from the reaction system, and nitrogen was used to blanket the reaction. Commercial grade nitrogen containing 0.01% oxygen by volume \vas not satisfactory, and so it was passed through tw'o scrubbing towers in series. Each contained 0.2.44 1
2
Present address. DOMChemical Co.. Pittsburg. Calif. Present addrrss, Stanford University. Stanford, Calif.
282
I&EC FUNDAMENTALS
chromous chloride solution in 1 .O.Y hydrochloric acid. Acid and water vapors were subsequently removed by passing the nitrogen through 0.1M sodium hydroxide solution and then through a calcium chloride dryer. Nitrogen was not passed through the reactor in a steady stream, but the pressure of the nitrogen above the reaction mixture was maintained slightly above atmospheric and was controlled manually by occasionally adjusting a bypass valve located upstream from the reactor. A water-filled manometer connected directly to the exit line indicated the pressure. Recrystallized sodium dithionite prepared as previously described ( 6 ) was used in the experimental work. A saturated aqueous solution of sodium dithionite prepared from the recrystallized material was used to supply dithionite to the reactor. T h e volume of the liquid mixture in the reactor was approximately 500 ml.? and the initial composition of the solution in the reaction vessel was varied by controfing the amount of saturated solution of sodium dithionite added. Most of the tests were conducted in triply distilled water. Some were performed in buffered solution wherein the pH was varied from 4.0 to 7.0 with buffering agents of KH2POd and N a O H or potassium acid phthalate and S a O H in proportions commensurate with the desired pH. Experiments at G O 0 , 70°,and 80" C. were conducted with the unbuffered solutions in order to determine the effect of temperature on the rate of decomposition. T h e buffered systems were studied only at 60" C. because only the effect of a controlled p H was of primary importance. At each temperature and for both the buffered and unbuffered systems, the initial concentration of sodium dithionite was varied from 5.5 to 11.5 x 1 0 - 3 ~ . In the unbuffered systems no products of reaction such as HS03-, SzO3-*, SOB-*,S, and H' were added to the system in order to determine their specific catalytic effects. This work had been done previously by Lynn (5). For the buffered reactions, however, HS03-, S203-*. and S O S - were ~ added in concentrations ranging from one quarter to three halves of the initial concentration of dithionite in general consideration of the stoichiometry expressed by the equation
2 S204-' f HnO
-+
Sz03-'
+ 2 HSO3-
(1)
The anion was always added while the decomposition experiment was in progress. By conducting successive tests of the reactor without removing the contents of the previous runs, the catalytic effects
of the product species !Sa03-* and sulfur \cere determined. sudlly . three dithionite solutions Lvere decomposed in each experiment. and in all cases the systems \vere buffered a t a pH of 5.00. .After the first decomposition \vhich \vas begun \cith fresh solution and after 'each successive decomposition, enough buffer solution was added to restore the volume of the reacior contents IO 500 ml., a reduction of volume having taken place as a result of the sampling procedure. .l'he initial dithionitc concentrations \\-ere varied from 5.5 to 0,0 X 1 O P . 2 1 by use of the aliquot portions of the same saturatcd solution. A pH measurement \vas taken at the beginning and end of each decomposition to ensure that slifficicnt buffer \vas present. Great care \vas taken to prevent oxygen contamination. \\lien: for example. a sample \vas transferred from the reactor to the titrarion flask. the sampling syringe \vas flushed \cith nitrogen several times prior to removal of the sample. As a further precaution to minimize subsequent contamination. ringe \vas rinsed several times Lvith oxygen-free distilled after each sample \vas taken. .l'hc concentration of li,204-2 \\\-as estimated by titration Lcith meth:,-lene blue ( ( I ) . X sample flask which received a measured volume of the reaction mixture \vas initial]>- charged \\.ith 50 mi. of 0.1.If KOIi and approximately 15 ml. of C H 3 0 H . T h e sample was titrated with methylene blue immediately in order to minimize the errors from slight oxygen contamination or further therma.: decomposition. Occasionally titrations were performed at ,approximately 5" C. to check for any significant effects of temperature upon the results. S o effect \\.as noted, so that titrarions a t room temperature \vere carried O U L Ivith confidence. .l'o determine the relative concentrations of the end products, a n unbuffered reaction .st 30" C. \vas carried out in a sealed vessel \vith a n initial dithjonite concentration of 7.32 X 10-31bf. T h e coniplerion of reaci-ion was determined by titration of a 10-mi. sample with mei.hylene blue. T h e concentration of the product S.03-* \vas estimated by iodometric titration of samples from the sealed vessel. Sampling flasks were initially filled ivith 10 ml. of formaldehyde, 30 ml. of distilled \cater: and 20 mi. of acetate buffer with a pH of about 5. T h e purpose of the formaldehyde \vas I O complex the HSO3- by the reaction
c-
CH?O
+ SaHS03
+
C H ? ( O H ) SO3 S a
that at 60" C. T h e data at 80" C. show practically no oscillations. For all three temperatures the total time of reaction may be seen to be inversely proportional to the initial concentration. T h e oscillatory behavior noted especially at 60' C. could be a result of poor analytical technique or a random catalytic effect brought about by a metallic contaminant from the syringe used to sample the reaction mixture. Relative to the former, sufficient experiments \\-ere run to show that
CI
62 zx
sg W L + J
rg Oa
0
RUN T 6 0 D 12 16 20 24 28 TIME 8 S E C O N D S X I U ~
8
4
32
Figure 1. Concentration of dithionite as a function of time for unbuffered system at 60" C.
(2)
Twenty minutes \cere allo\\.ed for completion of the reaction. I n the cornpiexed form. the HS03- \vas unaffected by Is-. so that the concentration O F S,Oa-* alone could be determined. 'I'he HS03.- ion \vas analyzed by a n iodometric oxidation in which the excess iodine \cas titrated ivith a standardized 0.01 .\- S,03-2 solution. T h e sampling flasks \\ere initially filled with 50 nil. of standardized 0.01.1- iodine solution and 25 ml. of acetate buffer with a pH of about 5. T h e resulting suni of the concentrations of SeOa-? and H S 0 3 - allo\ced the concentration of HSOB- i o be calculated bl- difference. 'l'he pH of the reacting mixtures \vas determined n i t h a standard Beckman 5lodi.l G pH meter which \vas connected to rlectrodes and inserted into the reactor through standardtaper glass fittings. Readings of the pH \vere taken approximately ever)- 100 seconds? so that, depending upon the duration of reaction, between 20 and 40 measurements of pH were obtained for each r..in. For the buffered systems the pH \vas measured only beforii and after reaction to ensure that the hydrogen ion concentration had been held constant.
I
1
2
4
I
I
6 8 IO 12 14 TIME 8 , SECONDS x
16
18
Figure 2. Concentration of dithionite as a function of time for unbuffered system at 70" C.
m
0 X
-0 RUN T 80 A 0 RUN T 8 0 0 9 RUN T 8 0 C
8 IO cJ d
awe v)
W
Res u Its
Unbuffered Reactions. Rate data for both buffered and unbuffered systems Lvere collected at 60'. 70". and EO" C. Figures 1. 2. and 3 she\-. In each case there is an induction period in \vhich the concentration of dithionite ion decreases slo\vly, follcnvcd b:, an abrupt decrease in dithionite concentration. At 60" C . the experimental points she\\- marked oscillations about the smoothed curves Lcith variations as high as *l5% from the smoothed values,. the average being closer to =k5707,. At 70' C . therc is also an oscillatory effect but less marked than
d2 6 V
z
94 f-
z 02 I
k n
2
4
6
e
TIME
e,
SECONDS x
10
I2
Figure 3. Concentration of dithionite as a function of time for unbuffered system at 80" C. VOL. 4
NO. 3
AUGUST
1965
283
the oscillation was not poor analytical procedure, and for the latter an operation in a glass sampling system did not change the degree of oscillation. T h e concentration of hydrogen ion in the unbuffered systems also was oscillatory. and it was measured by an electrode mounted directly in the reaction mixture so that no problems in transfer of a sample were involved. T h e oscillations are believed to be significant, but their meaning is not yet fully understood. Figure 4 shows the variation of hydrogen ion as compared to dithionite ion as a function of time for a run at 70" C. in an unbuffered system. Buffered Reactions. Experimental data for the concentration of dithionite ion as'a function of time for buffered systems are presented in Figure 5. T h e curves for concentration as a function of time are practically the same as those obtained for the unbuffered systems. Again there is an induction period follo\ved by a rapid decrease in the concentration of the dithionite ion. T h e characteristic oscillations which were present in the unbuffered systems were also found in the buffered sbstems. At pH 7 the concentration of dithionite decreased so slo\vly that the rapid decomposition period was not reached during the time in Lvhich the reaction was studied. Significant oscillations Lvere still present. It \vas found experimentally that the induction time and likeivise the total time of reaction decreased with increasing concentration of hydrogen ion. For the pH range of 4.80 to 6.00 and at an initial dithionite concentration of 11.0 X 1O - 3 M , time of the induction period \vas inversely proportional to the first power of the concentration of hydrogen ion. At a pH lower than 4.5, the decomposition \vas very rapid and could not be studied with the methods used. T h e addition of HSOa-, s203-*,S04P. and S O S - ~ had no measurable effect on the rate of decomposition of the dithionite at p H 7. Figure 6 sho\vs the results of a typical run in which HS03- and S 2 0 3 - * were added to the reacting system. In Figures 7 and 8 are presented the results of two identical experiments in solutions buffered at p H 5. Three successive decomposition reactions \
