Ind. Eng. Cbem. Fundam. 1980, 79, 207-209
207
Kinetics of Reaction of Sodium Hypochlorite and Sodium Sulfite by Flow Thermal Method R, D. Srivastava,' P. C. Nigam,' and S.
K. Goya12
Depdfment of Chemical Engineering, Indian Institute of Technology, Kanpur-208076, India
The kinetics of the reaction of sodium hypochlorite with sodium sulfite in aqueous solutions were studied at various temperatures by a flow thermal method. The experimental results showed that the reaction rate was first order with respect to both the hypochlorite and the sulfite concentrations. The second-order rate constant at 30 O C was calculated to be 6750 L/gmol s. An apparent activation energy of 15.6 kcal/gmol was calculated. A mechanism has been proposed. The oxidation of sulfite is really a reaction between hypochlorous acid and sulfite ions.
Introduction Hypochlorous acid is produced in significant amount when chlorine is added in power plant cooling water streams in order to keep condenser tubes free from slime. This chlorinated water containing hypochlorous acid is returned to the local stream or estuary where it is a threat to aquatic life. The reaction of sulfite ions with hypochlorous acid is uniquely suited to control hypochlorous acid. In a wellbuffered stream, this process only increases the sulfate and chloride ion concentration (Whitaker, 1977). Although apparently unknown in the chemical engineering literature, the kinetics have been studied by Lister and Rosenblum (1963). These investigators studied the rate of reaction of sodium hypochlorite with sodium sulfite using a rapid flow technique. The absorbance of solution was measured by means of a spectrophotometer. They concluded that the reaction takes place between the hypochlorite and sulfite ions and is second order. These authors indicate that theirs was the first study of this reaction and a careful search of the literature reveals that it is currently the only such study. The present work was carried out (a) to determine the rate of reaction of sodium sulfite with sodium hypochlorite in aqueous solutions in wide ranges of temperature, pH, and reactant concentrations and (b) to clarify the mechanism of sulfite oxidation by hypochlorite. In this investigation, measurements of the reaction rate were made by a flow thermal method (Hartridge and Roughton, 1923; Roughton, 1963). The temperature change of the solution as a result of the heat of reaction a t a particular point along the reactor is a measure of the amount of reaction occurring at that point. Experimental Section The rapid-mixing apparatus, differential measurements of the temperature, experimental procedure, and the method of data reduction have been described elsewhere (Mishra and Srivastava, 1975; Singh et al., 1978). Analysis of the sodium sulfite concentration in the feed solution was made by titration to the neutral point with alkaline iodine-iodide solution. The hypochlorite solutions were analyzed in the following way. A standard solution of potassium iodide was made in distilled water. An excess amount of this solution was acidified with glacial acetic acid and then it was reacted with a known amount of Department of Chemistry, Indian Institute of Technology, Kanpur-208016, India. Department of Chemistry and Chemical Engineering, University of Saskatchewan, Saskatoon, Sask., Canada.
sample. The reaction was allowed to proceed for 10 min in a dark room. The liberated iodine was then titrated with a standard sodium thiosulfate solution using starch solution as an indicator. From this titration the amount of hypochlorite in the solution was readily calculated. The overall stoichiometry for the reaction in aqueous solutions is NaClO + Na2S03 NaCl + Na2S04 (1) The standard enthalpy of the overall reaction is -81.3 kcal/g-mol of sulfite reacted (Weast, 1978). The energy liberated by the reaction raised the temperature of the solution and the temperature rise was measured. The change in sulfite concentration, M , was calculated by the energy balance AHR X M = AT X C,. The mean heat capacity C, of the solution was taken as that of water since the concentrations of both the reacting solutions were very low. The experimental data were obtained as potential differences along the reactor with respect to mixing point. The reliability of the equipment was checked using the NaOH-C02 system. The average of the second-order rate constant agreed within 5% with the value reported by Pinsent et al. (1956). Results and Discussion Figure 1 shows the plots of temperature difference as a function of residence time in two typical runs. Such plots always showed straight lines with different slopes for short residence times and then some curvature as the reactants are consumed. The rate of reaction was calculated from the heat of reaction and initial slope of such lines. Rate Studies. A number of experiments were performed at various hypochlorite concentrations to investigate its effect on the rate of reaction. All the experiments were conducted at 30 "C. The concentration of hypochlorite was varied in the range 0.0006-0.006 M. The sulfite concentration was kept constant at values of 0.0025 and 0.0075 M. Figure 2 shows that the rate against hypochlorite concentration results in one power relation with respect to hypochlorite for both the sulfite concentrations. In the determination of the order of reaction with respect to sulfite, the hypochlorite concentration was held constant. The concentration of sulfite was varied in the range 0.00124.01 M. The relation of rate vs. sulfite concentration is shown in Figure 3. All the data shown in this figure correspond to the average hypochlorite concentrations of 0.0012 and 0.0062 M. In both cases the rate against sulfite concentration resulted in one power relation. The reaction velocity constant for the overall reaction may be calculated from the integrated form of the rate equation describing the overall eq 1. An exhaustive num-
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0196-4313/80/1019-0207$01.00/00 1980 American Chemical Society
208
Ind. Eng. Chem. Fundam., Vol. 19, No. 2 , 1980
o5
I
A
[SO;-] [SO;-1
0 0075 M , LOCI7 :0 00063 M z 0 00063 M
:
= 00025 M I LOCI-I
Residence time, 8 , sec
Figure 4. Integral analysis of experimental data. _i 0 06 T i m e , sec
A
Figure 1. Typical temperature profiles as a function of time of reaction at 30 "C.
Present work Lister and Rosenblun
r
20
-
0
0
SO;-- 00025 M SO;-- 00075 M
Em
/
L
10-
.- . L
%x Y
-
-
5 -
-
2L 13 0
I1 1
I
'
H
M
~
l
'
'
I
''''1''
IO'
I
31
32
33
36
35
36
I
IO2
Average hypochlorite concentration x IO',
M
Figure 2. Effect of hypochlorite concentration on reaction rate at 30 'C.
l,Lo2 rn
OCI': 0 00625 OCI-: 0 00125
M M
A v e r a g e s u l f i t e concn x 10I.M
Figure 3. Effect of sulfite concentration on reaction rate at 30 O C .
ber of runs (-40 runs) was made at five residence times ranging from 0.011 to 0.040 s. The results were obtained with sulfite ranging from 0.001 to 0.01 M, hypochlorite from 0.0002 to 0.007 M. Figure 4 shows a plot of -
[ (t:oio) + log A
(Bo- A,)
A.
against time t. The straight line relationship confirms the validity of the reaction between hypochlorite and sulfite to be an irreversible second-order reaction. Error bars are twice the standard deviations combined with estimated
Ind. Eng. Chem. Fundam., Vol. 19, No. 2, 1980 209
a,
Scheme I
shown in step (ii). Finally, the ions H+, and SO-: containing sp3 hybridized sulfur atom are formed. The apparent activation energy, calculated from the variation of k"with temperature, was 15.6 kcal mol-'. Data on the heats of formation of the compounds (in aqueous solution) in the reaction NaOCl + H 2 0 HOCl + NaOH (4)
\
H
-
..q.. /-
v.1 -
:;I-o:..s--'d:
OH
activated complex
planar
GiI:]- t
[HI'
+ tetrahedral
[H')
+ [:3]
I H
cH-T1
(iii)
H
strength is zero. If one of the reactants is uncharged, the rate constant should be independent of ionic strength. An inspection of Lister and Rosenblum data shows this to be the case. Thus, the conclusion made by these authors that the reaction takes place between hypochlorite ions and sulfite ions appears to be incorrect. A mechanism which satisfies the established order of reaction with respect to hypochlorite and sulfite is probably one of the undissociated hypochlorous acid, and is given in Scheme I, where steps (i) and (iii) are rapid equilibria and step (ii) is rate determining; then rate = k'~HOC1][S032-] (3) It is possible to visualize the nature of the activated complex in the rate-determining step. As the oxygen end of HOCl approaches the planar sulfite ion, the three oxygen atoms on the sulfur move backward to assume a tetrahedral desposition while the electron pair of oxygen from HOCl accommodates itself in the vacant p orbital of sulfur atom. The sequence of movement of other electrons is also
make the heat of hydrolysis of the hypochlorite ion to be 11.0 kcal mol-'. Hence the activation energy for the reaction of sulfite ions and hypochlorous acid is only 4.6 kcal mol-'. This is perhaps to be expected, since if the new bond to oxygen is strong, this would tend to lower the energy of the activated complex as it is formed. Nomenclature A = sulfite concentration, g-mol/L A. = initial concentration of sulfite, g-mol/L Bo = initial concentration of hypochlorite, g-mol/L I = ionic strength, M k" = second-order rate constant, (g-mol/L)-' s-l k/ = limiting value of rate constant in Bronsted equation M = concentration, g-mol/L T = absolute temperature, K t = time, s 2 = ionic charges Literature Cited Bosolo, G., Pearson, R., "Mechanism of Inorganic Reactions", 2nd ed, Wiley, London, 1968. Hartridge, H., Roughton, F. J. W., Roc. R . SOC.London, Ser. A , 104, 376 (1923). Lister, M. W., Rosenblum, P., Can. J. Chem., 41, 3013 (1963). Mishra, G. C., Srivastava, R. D., Chem. Eng. Sci., 30, 1387 (1975). Moore, W. J., "Physical Chemistry", 3rd ed,Prentice-Hall, E n g l e w d Cliffs, N.J., 1964. .. Pinsent, B. R., Pearson, W. L., Roughton, F. J. W., Trans. Faraday Soc., 52, 1512 (1956). Roughton, F. J. W., "Techniques of Organic Chemistry", Vol. VIII, Interscience, New York. N.Y.. 1963. Singh, D. K., Sharma, R. N., Srivastava, R. D., AIChE J., 24, 232 (1978). Weast, R. C., "CRC Handbook of Chemlstry and Physics", 59th ed,CRC, 1978. Whkaker, S., Ind. Eng. Chem. Fundam., 16, 391 (1977).
Received for review October 2, 1979 Accepted February 8 , 1980
