I n d . Eng. Chem. Res. 1991,30, 1077-1080
1077
KINETICS AND CATALYSIS
Kinetics of the Reaction of Chlorine with Formaldehyde in Aqueous Sulfuric Acid Bhart Indu,t M. Fazlul Hoq,t William R. Ernst,*Jand Henry M. Neumannt School of Chemical Engineering and School of Chemistry and Biochemistry, Georgia Institute of Technology, Atlanta, Georgia 30332
In sulfuric acid solution, the rate of reaction of chlorine with formaldehyde is first order with respect t o both reactants. T h e rate declines with acid concentration. A rate law previously developed to describe the rate of reaction of bromine with formaldehyde did not adequately describe the acid dependence of the chlorine-formaldehyde reaction. A two-parameter empirical model was developed to correlate the data over a wide range of conditions.
Introduction Early literature on the reaction of bromine and formaldehyde in sulfuric acid (Scheffer and van Went, 1928; Anderson, 1913; McTigue and Sime, 1963) reports that the stoichiometry involves oxidation of formaldehyde to formic acid Br2 + HzCO + H 2 0 2H+ + 2Br- + HCOOH (1) followed by the equilibrium ionization of formic acid HCOOH e+ H+ + HC02(2)
-
-
and the bromine-formate reaction Br2 + HC02- H+ + 2Br-
+ C02
(3) McTigue and Sime (1963a) studied reaction 1using commercial formaldehyde solution. They proposed a mechanism involving equilibrium formation of formaldehyde and methylene glycol H2COH+ HZCO + H+ (4) and H2CO + H2O e+ H,C(OH), (5) followed by the rate-determining step H2C(OH)2+ Br2 2H+ + 2Br- + HCOOH (6) Q
-
They derived the rate law -d[Br~l/dt = ke[F~lPr2l where
(7)
with a, = activity of water; [FT]= total "formaldehyde" concentration = [H2CO]+ [H2COH+]+ [H2C(OH)21; ho = acidity according to Ho= -log ho;Ho= Hammett acidity function; k6 = rate constant for reaction 6; k, = effective rate constant; K4 = acidity constant for protonated formaldehyde = (ho)[H2CO]/[H2COH+];and K5 = hydration
* Author to whom correspondence should be addressed. School of Chemical Engineering.
* School of Chemistry and Biochemistry. 0888-5885/91/2630-lO77$02.50/0
equilibrium constant = H2C(OH),/ [H2CO](aw). Using a value of 1300 for K5, McTigue and S h e (1963b) found that the rate law fit reaction data for an assumed value, K4 z 16000. For sulfuric acid concentrations less than 10 M, k, is not dependent upon acidity and equals ke. They showed a value of k6 = 5 X At higher acidity, k , drops sharply with sulfuric acid concentration. Cox and McTigue (1964) have also shown that hydrated aldehyde is the reactive species in the reaction of bromine with acetaldehyde. The oxidation of formaldehyde and formic acid with chlorine at moderate sulfuric acid concentrations may be important reactions in the commercial production of chlorine dioxide based upon the reduction of sodium chlorate with methanol (Masschelein, 1979; Norell, 1988; Swindells and Fredette, 1982). This process is not new; however, until recently, most chlorine dioxide processes were based upon the reduction of sodium chlorate with chloride ions (Masschelein, 1979; Ernst et al., 1988; Tenney et al., 1990). This latter process yields chlorine as a major byproduct. The methanol-chlorate process is becoming more important for environmental reasons because it produces only trace amounts of chlorine byproduct. The oxidation of formic acid with chlorine has been recently studied (Hoq et al., 1991) and was shown to follow a mechanism analogous to reactions 2 and 3 for the bromine-formic acid system. This work is an experimental study of the kinetics of the chlorine-formaldehyde reaction in sulfuric acid at concentration below 8 M, which is a region of commercial interest.
Experimental Section The equipment consisted of an open 250-cm3 stirred reaction flask, partially submerged in a constant-temperature bath. The total liquid volume in each experiment was 200 cm3. Chloride concentration of the reaction liquid was monitored by a chloride sensitive electrode. The electrode was calibrated after each run by use of standard sodium chloride solutions, adjusted to temperature and acidity of the experiment. Dissolved chlorine was monitored at 322 nm by a Milton Roy Spectronic 1201 UV spectrophotometer equipped with a 0.4-cm3flow cell with a 1-cm path length. The extinction coefficient for chlorine was confirmed at a constant value (c = 75) over the range of acidities of this study, and concentrations of chlorine 0 1991 American Chemical Society
1078 Ind. Eng. Chem. Res., Vol. 30, No. 6, 1991 Table I. Production of Chloride for Several Reactant Conditions at 298 K run no. 40 41 31 27 47 25 3.25 3.25 3.25 4.5 2.0 [HZSO,], M 1.0 0.026 0.026 0.004 0.013 0.020 0.026 [Cl,], M [HCHO], M 0.23 0.23 0.10 0.10 0.10 0.23 concentration of chloride ions, M X lo3 time, s 0 60 120 180 240 300 360 420 480 540 600 660 720 780 840 900
3.24 4.30 5.35 6.61 8.05 9.61 11.10 12.60 14.30 16.00 17.60
1.53 2.55 3.67 4.92 6.22 7.49 8.95 10.30 11.80 13.20 14.70 16.10 17.50 18.90 20.30
0.25 0.30 0.35 0.40 0.46 0.52 0.59 0.65 0.71 0.77 0.84 0.90 0.97 1.04 1.11 1.17
0.32 0.43 0.56 0.69 0.84 1.00
1.17 1.35 1.53 1.71 1.88 2.07 2.25 2.45 2.62 2.82
0.38 0.56 0.76 0.99 1.24 1.49 1.75 2.05 2.34 2.64 2.92 3.22 3.52 3.83 4.12 4.42
0.22 0.55 1.11 1.71 2.30 2.91 3.53 4.15 4.76 5.33 5.91 6.47 7.02 7.60 8.15 8.68
were determined by iodometric titration. Liquid was continuously pumped a t 100 cm3/min from the reaction flask through the flow cell and back to the flask. Residence time in this recycle system was about 6 s. In each experiment, sulfuric acid solution was brought to constant reaction temperature (*OS K) in the reaction flask. Chlorine gas (99.9%) from a lecture bottle (Matheson Co.) was bubbled through the solution via a fritted glass sparger. The chlorine concentration was maintained constant by making fine adjustments to the flow so that the rate of input balanced the rate of evaporation of chlorine from the liquid surface in the open flask. After chlorine concentration achieved a constant value, formaldehyde (Fisher Scientific 37 % w/w, 10-15'70 v/v methanol) was quickly added to the flask. The concentrations of chloride and chlorine were then monitored as functions of time, usually for 15-30 min, depending upon the rate of reaction. Chlorine concentration was maintained constant as before (see Table I for typical run data). Concentration of the formaldehyde stock solution used was verified by the addition of excess solution of sodium sulfite to a known volume of the solution and the titration of the liberated alkali with standard acid (Walker, 1953). Experiments were conducted at sulfuric acid concentration between 0.1 and 8.0 M,although most of the work was done at 1 and 3.25 M. Other ranges of conditions include temperature, 280-318 K;formaldehyde concentration, 0.033-0.46 M; and chlorine concentration, 0.004-0.0335 M. Reaction rates were calculated from data recorded up to the time that the formaldehyde conversion was 3% or less.
Results and Discussion 'H NMR spectra of chlorine-formaldehyde solution showed formic acid as the only organic product of the above reaction. Because the experiments were continued to 3% (or less) conversion of formaldehyde, the reaction of formic acid with chlorine (the chlorine analogue of reactions 2 and 3) did not contribute significantly to the overall chloride formation rate. It is difficult to assess the impact of the methanol additive in the formaldehyde solution on the results. McTigue and Sime (1963a) do not discuss this point. Methanol acts as a stabilizing agent to prevent polymerization for the formaldehyde. This stabilizing action is thought to involve formation of hemiacetals which exist in a state of equilibrium with hydrated formaldehyde
Figure 1. Product concentrations during reaction experiment at 298 K, 3.25 M H2SOl, 0.1 M HzCO, and 0.020 M Cl,, (symbols: 0,C1-, +, HCOOH).
(Walker, 1953). We measured the rate of reaction of chlorine with methanol at various sulfuric acid concentrations. We conclude that, in the chlorine-formaldehyde experiments, if none of the methanol were converted to the hemiacetal, the reaction of chlorine with methanol would contribute less than 4% of the overall rate of chloride formation. We also ran several chlorine-formaldehyde experiments in which acidified formaldehyde was introduced into the chlorine solution. The rates of chloride formation were identical within experimental accuracy with those rates found by the usual procedure in which formaldehyde is introduced into acidified chlorine solution. This finding shows that equilibria among formaldehyde species and sulfuric acid are established rapidly. Figure 1 (upper curve) shows a graph of the measured chloride ion concentrate versus time for one experiment. For some experiments these graphs displayed a slight curvature. The graph of formic acid concentration (lower curve) was calculated from the chloride graph. In general for all of the experiments these graphs were linear. To compute the relationship between chloride and formic acid concentration, we assumed the stoichiometry is determined by the reaction of chlorine and methylene glycol (McTigue and Sime, 1963) H,C(OH), + C12 2H+ + 2C1- + HCOOH (9) and the equilibrium hydrolysis of chlorine C1, + H 2 0 HOC1 + H++ C1(10)
-
Q
The equilibrium formation of C13-
c1, + c1- c1,-
(11)
was neglected, because it has a reported equilibrium constant (Zimmerman and Strong, 1957) of 0.191 at 25 "C. At the chlorine concentration of this study, less than 0.6% of chloride would be converted to CI,-. For experiments conducted at constant chlorine, formaldehyde, and sulfuric acid concentration, the relationship between formic acid and chloride concentration is (see the Appendix) [HCOOH] = 0.5 ([Cl-] - [C1-]02/[C1-]) (12) In determining reaction orders, one reactant concentration was varied while others were kept constant in a series of experiments. Orders were calculated from the slopes of appropriate logarithmic plots of formic acid formation rate versus concentration. These series of experiments were conducted at several sulfuric acid concentrations. Figure 2 shows the log rate versus log chlorine
Ind. Eng. Chem. Res., Vol. 30,No. 6,1991 1079
,
-5 ,
0.W4
O.-0.003
0.0025
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-
-5.5
-
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0.002 0.0015
41
n\o
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-5.8
1
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-Le
-2
-22
-18
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-14
2
4
Figure 2. Chlorine reaction order (298 K, 1 M H2SOI, and 0.1 M HZCO). -4.7 -4.8
-51
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am
-
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-
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-
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Figure 4. Dependence of effective rate constant on sulfuric acid concentration at 298 K, 0.23 M H&O, and 0.026 M Clz ( 0 ,experimental data; solid line, calculated).
1
-
-5
6
IHZS041. M
Lnglc12L M
-5
, -13
,
, -11
,
-
0.9
0.7
-05
0.3
l o d H x O 1M
Figure 3. Formaldehyde reaction order (298 K, 3.25 M H&304,and 0.026 M Clz).
concentration at 0.1 M formaldehyde, 1 M sulfuric acid, and 298 K. The order in chlorine is 0.94f 0.06 at 3.25M H2S04and 0.94 f 0.03 at 1 M H2SOI. Figure 3 shows the log rate versus log formaldehyde concentration at 0.026M chlorine, 3.25 M sulfuric acid, and 298 K. The order in formaldehyde is 0.95 f 0.04 at 3.25 M H2S04 and 0.96 f 0.06 at 1 M H2S04. The rate is not dependent upon chloride concentration, as indicated by the linearity of the formic acid concentration-time graphs, such as in Figure 1. Therefore, at constant acidity, the rate law for this system is d[HCOOH]/dt = k,'[Cl,l[F~] (13) where k,' is a function of acidity. Effect of Acidity. The model proposed by McTigue and Sime (1963)(eqs 7 and 8) did not fit the chlorineformaldehyde reaction data. Using their values of K4 and K5, the model predicts that k,' is not dependent upon acidity for sulfuric acid concentration less than 10 M. We attempted to use this model by setting K5 at a currently accepted value, 2000 (Bell, 1966),and letting K4 be an adjustable parameter. The agreement between model prediction and data was poor for all assumed values of KI. The model gives a poor representation because of the large increase in the value of ho with acidity. An empirical equation kKa, k,' = M + Ka,
\
1
-; -8
-8.50.003
0.0032
I/T.np.ratlr.. 0.0034 K
0.0038
Figure 5. Temperature dependence of effective rate constant at 1 M H2S04,0.23 M HzCO, and 0.026 M Clz (0, experimental data, solid line, calculated by wing eq 14).
obtained by substituting molarity of sulfuric acid, M,for ho and by substituting k = k6 and K = K4Kb(assuming K5a, >> 1)in eq 8 provides an excellent fit of the data (see Figure 4). The optimal values of the parameters k and K a t 298 K are 3.73 X s-' M-' and 1.93 M, respectively. This provides a practical basis for rate calculation. Effect of Temperature. Values of parameters k and K were likewise determined at other temperatures. Temperature coefficients were determined assuming the parameters follow the usual exponential form k = kOe-E/RT (15) and K = KoeQ/RT (16) To account for the temperature variation of a,, we used the general expression reported by Giauque et al. (1960) a, = aWOe-LIRT (17) where L is a function of acidity. The following values were found for the constants in eqs 15 and 16: ko = 1.0 X 10" s-l M-l, E = 76.1 kJ/mol, KO = 0.0103 M, and Q = 12.72 kJ/mol. A model for predicting effective rate constants at all temperatures and acid concentrations can therefore be developed by utilizing the above constants and inserting the temperature-dependent equations 15-17 into eq 14. To illustrate the predictability of this model, Figures 5 and
1080 Ind. Eng. Chem. Res., Vol. 30,No. 6, 1991 -4
7
~
I
I -4 5
-7 5 C
-8
-8 5 -9
~-
~~~~~~
0 003
00032
00034
experimental data; M HzS04,0.23 M HzCO, and 0.026 M ClP. (0, solid line, calculated by using eq 14).
6 have been constructed. These figures show the close agreement between experimental and calculated values of effective rate constants over a range of temperatures at two sulfuric acid concentrations, 1 and 3.25 M. In these figures, the data points represent experimentally determined values whereas the solid line represents the model prediction.
Acknowledgment We gratefully acknowledge financial support from Eka Nobel, Inc.
Appendix. Derivation of Eq 12. From the stoichiometry of the reaction system, eqs 9 and 10, chloride concentration can be described by the expression [Cl-] = 2[HCOOH] + [HOC11 (A.1) Assuming equilibrium is established quickly for chlorine hydrolysis [Cl,l
Combining these two expressions
HCOOH, 64-18-6.
Literature Cited
l/Tsmparalure K
[HOC11[H+][C1-]
Both proton and chlorine concentrations remain constant during an experiment. Combining eqs A.3 and A.4, we obtain [Cl-1 = 2[HCOOH] + [C1-],2/[Cl-] (A.5) which rearranges to yield eq 12. Registry No. Clz, 7782-50-5; HCHO, 50-00-0; C1,16887-00-6;
0 0036
Figure 6. Temperature dependence of effective rate constant at 3.25
Keq =
At the beginning of an experiment pure chlorine is dissolved in water to form a solution establishing equilibrium concentrations of chloride and hypochlorous acid in equimolar amounts. Therefore Kw can be expressed as
(A.2)
Anderson, E. The Oxidation of Aldehydes by an Aqueous Solution of Bromine. Am. Chem. J. 1913,49, 182. Bell, R. P. The Reversible Hydration of Carbonyl Compounds. Ado. Phys. Org. Chem. 1966, 4, 1. Cox, B. G.; McTigue, P. T. The Oxidation of Aldehydes with Bromine. Aust. J. Chem. 1964, 17, 1210-1216. Ernst, W. R.; Shoaei, M.; Forney, L. Selectivity Behavior of the Chloride-Chlorate Reaction Svstem in Various Reactor Types. -_ AIChE J. 1988, 34, 1927. Giauque, W. F.; Hornung, E. W.; Kunzler, T. R.; Rubin, T. R. The Thermodvnamic ProDerties of Aaueous Sulfuric Acid Solutions and Hydrates from 15 to 300K. j.Am. Chem. SOC.1960,82,62. Hoq, M. F.; Indu, B.; Ernst, W. R.; Neumann, H.M. Kinetics of the Chlorine-Formic Acid Reaction. J. Phys. Chem. 1991, 95, 681-683. Masschelein, W. J. Chlorine Dioxide; Ann Arbor Science: Ann Arbor, MI, 1979. McTigue, P. T.; Sime, J. M. The Basic Strength of Formaldehyde. A u t . J. Chem. 1963,16,592-595. (a) The authors did not specify the composition of the commercial formaldehyde solution which they used. (b) The authors did not report a value for k, in eq 8; it is apparent from Figure 2 of their article that an approximate best fits their data. value, 5 X Norell, M. US. Patent No. 4,770,868, Sept 13, 1988. Scheffer, F. E. C.; van Went, N. B. L'action du Brome Sur la Formaldehyde. Recl. Trav. Chim. 1928,47, 406-414. Swindells, R.; Fredette, M. C. J. US.Patent No. 4,325,934, April 20, 1982. Tenney, J.; Shoaei, M.; Obijeski, T.; Ernst, W. R.; Lindstroem, R.; Sundblad, B.; Wanngard, J. An Experimental Investigation of the Chloride-Chlorate Reaction System. Znd. Eng. Chem. Res. 1990, 29, 916. Walker, J. Formaldehyde, 2nd ed.; Reinhold New York, 1953. Zimmerman, G.; Strong, F. C. Equilibria and Spectra of Aqueous Chlorine Solutions. J. Am. Chem. SOC.1967, 79, 2063.
Receiued for review September 12, 1990 Accepted January 14, 1991
