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Ind. Eng. Cbem. Fundam. 1980, 79 2 0 7 - 2 0 9 207

Kinetics of Reaction of Sodium Hypochlorite and Sodium Sulfite byFlow Thermal Method

R, D. Srivastava, P. C. Nigam, and S. K.Goya12Depdfment 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 varioustemperatures by a flow thermal method. T h e experimental results showed that t he reaction rate was first orderwith respect to both t he hypochlorite and t h e sulfite concentrations. The second-order rate constant at 30 Owas calculated to be 6750 L/gmol s. n apparent activation energy of 15.6kcal/gmol was calculated. A mechanismhas been proposed. The oxidation of sulfite is really a reaction between hypochlorous acid and sulfite ions.

IntroductionHypochlorous acid is produced in significant amountwhen chlorine is added in power plant cooling waterstrea ms in order to keep condenser tubes free from slime.This chlorinated water containing hypochlorous acid isreturned to th e local stream or estuary where it is a threatto aquatic life.T h e reaction of sulfite ions with hypochlorous acid isuniquely s uite d to control hypochlorous acid. In a well-buffered s tream , this process only increases th e sulfate andchloride ion concentration (W hita ker , 1977). Althoughapparently unknown in the chemical engineering literature,the kinetics have been studied by Lister an d Rosenblum(1963). Th ese investigators studied the rate of reactionof sodium hypochlorite with sodium sulfite using a rapidflow technique . T he absorban ce of solution was measuredby means of a spectrop hotom eter. Th ey concluded tha tthe reaction takes place between the hypochlorite andsulfite ions and is second order. These a uthors indicateth at theirs was th e first study of this reaction and a carefulsearch of th e literature reveals th at it is currently the onlysuch s tudy.T he prese nt work was carried out (a ) to determine th erate of reaction of sodium sulfite with sodium hypochloritein aque ous solutions in wide ranges of temperature, p H,and r eac tant concentrations and (b) to clarify the mech-anism of sulfite oxidation by hypochlorite.In th is investigation, measurements of the reaction ratewere made by a flow thermal method (Hartridge andRoughton, 1923; Roughton, 1963) . Th e temperaturechange of the solu tion as a result of the he at of reactiona t a particular point along the reactor is a measure of theam oun t of reaction occurring a t tha t point.Experimental Section

T he rapid-mixing app arat us, differential measurementsof the tem perature, experimental procedure, and the me-thod of data reduction have been described elsewhere(M ishra an d Srivastava, 1975; Singh e t al., 1978).Analysis of the sodium sulfite concentration in the feedsolution was made by titratio n to the ne utral point withalkaline iodine-iodide solution. T he hypochlorite solutionswere analyzed in the following way. A standard solutionof potassium iodide was made in distilled water. An excessamount of this solution was acidified with glacial aceticacid and then it was reacted with a known amount ofDepartm ent of Chemis try, Indian Ins t i tute of Technology,Depar tm ent of Chemistry an d Chemical Engineering, Univ-Kanpur-208016, India.ersity of Saskatchewan, Saskatoon, Sask., Canada.

0196-4313/80/1019-0207 01.00/0

sample. T he reaction was allowed to proceed for 10 minin a dark room. Th e liberated iodine was then tit ratedwith a standard sodium thiosulfate solution using starchsolution as an indicator. From this ti tration the am ountof hypochlorite in the solution was readily calculated .The overall stoichiometry for the reaction in aqueoussolutions is

1)The standard enthalpy of the overall reaction is -81.3kcal/g-mol of sulfite reacted (Weast, 1978). The energyliberated by the reaction raised the tem perature of thesolution and the temperatu re rise was measured. Th echange in sulfite co ncentration, M was calculated by th eenergy balance AHRX M = AT X C . T he m ean hea tcapacity C, of th e solution was taken as th at of water sincethe c oncentration s of both th e reacting solutions were verylow. Th e experimental da ta were obtained a s potentialdifferences along the reacto r with respec t to mixing point.T he reliability of the eq uipment was checked using theN a O H - C 0 2 system. Th e average of the second-order rateconstant agreed within 5 with the value reported byPinsent et al. (1956).Results and DiscussionFigure 1 shows the plots of temperature difference asa function of residence time in two typical runs. Such plotsalways showed straigh t lines with different slopes for sho rtresidence times and th en some curvature as th e reactantsare consumed. T he rate of reaction was calculated fromth e he at of reaction an d initial slope of such lines.Rate Studies A number of experiments were per-formed at various hypochlorite concentrations to inves-tigate its effect on the ra te of reaction. All the experimentswere conducted at 30 C. T he concentration of hypo-chlorite was varied in the range 0.0006-0.006 M. T h esulfite concentration was kep t constan t a t values of 0.0025and 0.0075 M. Figure 2 shows that the ra te against hy-pochlorite concentration resu lts in one power relation withrespect to hypochlorite for both t he sulfite concentrations.In the determination of the order of reaction with re-spe ct to sulfite, th e hypochlorite concen tration was heldconstan t. T he concentration of sulfite was varied in therange 0.00124.01 M. The relation of rate vs. sulfite con-cen tratio n is shown in Figure 3. All th e dat a shown inthis figure correspond to the average hypochlorite con-centrations of 0.0012 and 0.0062 M. In both cases the rateagainst sulfite concentration resulted in one power relation.T h e reaction velocity con stan t for the overall reactionmay be calculated from the integrated form of the rateequation describing the overall eq 1. An exhaustive num-

NaClO + N a 2 S 0 3- aCl + N a 2 S 0 4

1 9 8 0 American Chemical Society

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208 Ind. Eng. Chem. Fundam., Vol. 19, No. 2 1980

3 0 31 3 2 3 3 3 6 3 5

[SO;-] 0075 M , LOC I 7 00063 MA [SO;-1 = 00025 M I LOCI-I z 0 00063 Mo5 I

_i0 6T i m e , s e c

Figure 1. Typical temperature profiles as a function of time ofreaction at 30 C .

3 6

SO;-- 00025 MSO;-- 00075 M

20

-1 0 -

5 -

m -. . --x -

I I ' H M ~ l ' I ' ' ' ' 1 ' ' I I1 IO I O 2

Aver age hypoch lo r i t e concen t r a t ion x I MFigure 2. Effect of hypochlorite concentra tion on reaction rate a t30 'C.

r-

--

r OCI': 000625 MOCI-: 0 00125 Ml Lo2v e r a g e s u l f i t e concn x 10I .MFigure 3. Effect of sulfite concentra tion on reaction rate a t 30 O C .ber of runs (-40 runs) was made a t five residence timesranging from 0.011 to 0.040 s. Th e results were obtainedwith sulfite ranging from 0.001 to 0.01 M, hypochloritefrom 0.0002 to 0.007 M. Figure 4 shows a plot of[(t:oio)og A + Bo A,) A .

against time t . T he straight line relationship confirms th evalidity of the reaction between hypochlorite and sulfiteto be an irreversible second-order reaction. Error ba rs aretwice the standard deviations combined with estimated

Residence t ime , 8 , secFigure 4. Integral analysis of ex perime ntal data.

A Present workLister and Rosenblun

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Ind. Eng. Chem. Fundam., Vol. 19, No. 2, 1980 209shown in ste p (ii). Finally, the ions a +, and SOcontaining sp3 hybridized sulfur a tom a re formed.The apparent activation energy, calculated from thevariation of k with temp eratu re, was 15.6 kcal mol-'. Da taon the heats of formation of th e compoun ds (in aqueoussolution) in t he reaction

NaOCl + H 2 0- OCl + NaOH (4)mak e the h eat of hydrolysis of th e hypochlorite ion to be11.0 kcal mol-'. Hence th e activation energy for the re-action of sulfite ions and hypochlorous acid is only 4.6 kcalmol-'. Th is is perh aps to be expected, since if the newbond t o oxygen is strong, this would tend to lower theenergy of th e activated complex as it is formed.N o m e n c l a t u r eA = sulfite concentration, g-mol/LA . = initial conce ntration of sulfite, g-mol/LBo = initial concentration of hypochlorite, g-mol/LI = ionic streng th, Mk = second-order rate constant, (g-mol/L)-' s lk/ = limiting value of rate constant in Bron sted equationM = concentration, g-mol/LT = absolute temperature, K= time, s= ionic chargesL i t e r a t u r e C i t e dBosolo G., Pearson, R., Mechanism of Inorganic Reactions , 2nd ed, Wiley,Hartridge, H., Roughton, F. J. W., Roc. R . SOC.London S er . A , 104, 376Lister, M. W., Rosenblum, P., Can. J. Chem. 41, 3013 (1963).Mishra, G. C., Srivastava, R . D. Chem. Eng. Sc i . 30 1387 (1975).Moore, W. J., Physical Chemis try , 3rd ed, rentice-Hall, E n g le w d Cliffs, N.J.,

London, 1968.(1923).

1964.

Scheme I

\H

q. / -:;I-o:...s--d:OHv.1planar activated complex

GiI:]- t [HI' +tetrahedral

(iii)cH T1H') + [ : ] HHstrength is zero. If one of the reac tants is uncharged, th erate co nstant should be indep endent of ionic strength. Aninspection of Lister and R osenblum d ata shows this to bethe case. Th us, the conclusion made by these authors tha tthe reaction takes place between hypochlorite ions andsulfite ions appea rs to be incorrect. A mechanism whichsatisfies th e established order of reaction with respect tohypochlorite and sulfite is probably one of the undisso-ciated hypochlorous acid, and is given in Scheme I, wheresteps (i) and (iii) are rapid equilibria and step (ii) is ratedetermining; then

ra te = k '~HOC1] [S032- ] (3)It is possible to visualize the na ture of th e activated com-plex in the rate-determining step. As th e oxygen end ofHOCl appro aches th e planar sulfite ion, th e three oxygenatom s on the sulfur move backward to assume a tetrahe -dral desposition while the electron pair of oxygen fromHOC l accomm odates itself in th e vacant p o rbital of sulfuratom. Th e sequence of mov emen t of other electron s is also

..Pinsent, B. R., Pearson, W. L., Roughton, F. J. W., Trans. Faraday Soc., 52Roughton, F. J. W., Techniques of Organic Chemistry , Vol. VIII, Interscienc e,1512 (1956).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 fo r review October 2, 979Accepted February 8 , 1980