Ind. Eng. Chem. Res. 1992,31, 2425-2437
extent to which specific reactions like those listed above may be occurring is yet to be determined, these considerations suggest that a more complete depiction of the envisioned parallel pathway scheme would look like that presented below:
x CHI
c 9 COztCH4 CH30H
*
cox
Conclusions This study examined conversion, selectivity, and operating temperature trends associated with varying total pressure and O2partial pressure in noncatalytic oxidative CHI conversion; additionally, some plausible inferences about the reaction network were drawn. The resulta of the present work are consistent with those of its predecessor study: selectivity/conversion behavior in high-pressure, low-temperature noncatalytic operation resembles that of conventional catalytic OC. Such observations agree with the view that the primary role of the catalyst is that of a radical generator which initiates and sustains the reaction but does not substantially benefit product selectivity. At high enough pressures, the need for generation of radicals on the surface is reduced, and catalytic influence diminishes.
242s
Registry NO. CHI, 74-82-8; CO, 630-08-0; COZ, 124-38-9; CHSOH, 67-56-1; CH3CH3,7484-0; CH-HZ, 7485-1; CH*H, 74-86-2; CH3CHZCH3, 74-98-6.
Literature Cited Amenomiya, Y.; Birss, V. I.; Goledzinowski,M.; Galuezka, J.; Sanger, A. R. Conversion of Methane by Oxidative Coupling. Catal. Rev.-Sci. Eng. 1990, 32, 163. Asami, K.; Omata, K.; Fujimoto, K.; Tominaga, H. J. Oxidative Coupling of Methane in the Homogeneous Gas Phase under Pressure. J. Chem. SOC.,Chem. Commun. 1987, 1287. Baerns, M. Oxidative Catalytic Methane Conversion. Catal. Today 1987, 1, 357. Ekstrom, A.; Regtop, R.; Bhargava, S. Effect of Pressure on the Oxidative Coupling Reaction of Methane. Appl. Catal. 1990,62, 253. Hutchings, G. J.; Scurrell, M. S.; Woodhouse, J. R. The Role of Gas Phase Reaction in the Selective Oxidative of Methane. J. Chem. SOC.,Chem. Commun. 1988, 253. Hutchings, G. J.; Scurrell, M. S.; Woodhouse, J. R. Oxidative CouRev. 1989, pling of Methane using Oxide Catalysts. Chem. SOC. 18, 251. Onsager, 0. T.; Lodeng, R.; Soraker, P.; Anundskaaa, A.; Helleborg, B. The Homogeneous Gas Phase Oxidation of Methane and the Retarding Effect of Basic/Inert Surfaces. Catal. Today 1989,4, 355. Rytz, D. W.; Baiker, A. Partial Oxidation of Methane to Methanol in a Flow Reactor at Elevated Pressure. Znd. Eng. Chem. Res. 1991,30, 2287. Walsh, D. E.; Martenak, IA J.; Han, S.; Palermo, R. E. Direct Oxidative Methane Conversion at Elevated Pressure and Moderate Temperatures. Znd. Eng. Chem. Res. 1992,31,1259.
Received for review May 20, 1992 Accepted August 3, 1992
Selective Synthesis of a-ChlorocarboxylicAcids Erkki Paatero,?Tapio Salmi,* and Krister Fagerstolt' Laboratory of Industrial Chemistry, Abo Akademi, SF-20500 Turku, Finland
The chlorination of butanoic, hexanoic, octanoic, decanoic, and dodecanoic acids was studied in a laboratory-scale semibatch reactor operating a t temperatures ranging from 70 to 130 O C a t atmospheric pressure. The experiments showed that carboxylic acids can be selectively a-chlorinated with chlorine in the presence of molecular oxygen as a radical trapper and chlorosulfonic acid as an enolizing agent: the only products were the a-monochloro- and a,a-dichlorocarboxylic acids. High yields of a-chlorocarboxylic acids were obtained when the enolizing catalyst was added stepwise into the reaction mixture. In the case that the catalyst was introduced only at the initial stage, the reaction rate stagnated after a short period of time, which was due to the decomposition of chlorosulfonic acid. In the stepwise catalyst addition the chlorocarboxylic acid formation was autocatalytic. This effect was explained by a reaction mechanism involving the acid-catalyzed enolization of a reaction intermediate as the main rate-controlling step. The parameters of the rate equation based on the proposed reaction mechanism were determined for the carboxylic acids studied. The kinetic parameters were linearly dependent on the carboxylic acid carbon number.
Introduction We have previously studied the kinetics of the chlorination of acetic acid with molecular chlorine in the presence of acetyl chloride as a catalyst (Martikainen et al., 1987;Salmi et al., 1988). In the present investigation we extend the scope to the chlorination of straight-chain alkanoic acids with carbon numbers ranging from 4 to 12.
The resulting chlorocarboxylic acids are potential intermediates in the chemical industry due to their reactivity, which is based on the charge shift caused by the electronegative chlorine atom. The carbon atom attached to chlorine gets a positive charge and it is eager to react, e.g., with a nucleophile: CI-
I
R-C'-COOH
Present address: Lappeenranta University of Technology, SF-53851 Lappeenranta, Finland. t Present address: Neste Oy, Technology Centre, SF-06101 Porvoo, Finland. f
I
R'
+ HO-R"
-
rI
R-C'-COOH
I
+
HCI
(1)
R'
In the case of acetic acid chlorination the only reaction products are mono-, di-, and trichloroacetic acids (Grog-
0888-588519212631-2425$03.00/0 0 1992 American Chemical Society
2426 Ind. Eng. Chem. Res., Vol. 31, No. 11, 1992
gins, 1958; Martikainen et al., 1987). Monochloroacetic acid is the main product (>93%) whereas the polychlorinated acids are industrially generally regarded as undesired byproducts. In the chlorination of unsaturated fatty acids with a straight hydrocarbon chain longer than two carbon atoms, the chlorination selectivity is of central interest as aliphatic chains can undergo chlorination to a number of different positions. In the present work the aim is to synthesize monochlorocarboxylic acids with the chlorine atom attached to the a-carbon atom according to the following overall reaction: RCHzCOOH + C1,
-
RCHClCOOH
+ HC1
(2)
When the above reaction takes place through a radical mechanism, a wide range of different chlorinated compounds are produced. In the earlier studies by Ogata and his co-workers, it was shown that the radical chlorination can be suppressed with radical-trapping agents such as rn-dinitrobenzene (Ogata and Matsuyana, 1970) and oxygen (Ogata et al., 1975). In order to obtain reasonably short chlorination times, a catalyst is also needed. Strongly acidic enolizing agents, such as fuming sulfuric acid and chlorosulfonic acid, have been used as catalysts in the previous works (Ogata and Matsuyama, 1970; Ogata et al., 1975,1979). The problem is, however, that large amounts of catalysts are needed to force the reaction to high yields in reasonable reaction times as shown by Ogata et al. (1979) in the chlorination of dodecanoic acid in the presence of oxygen and chlorosulfonic acid. In spite of large amounts of the catalyst (20 mol %) the reaction stagnated after 60 min reaction time at a yield of 90%. The aim of the present work is to find a method for the selective synthesis of a-chlorocarboxylic acids by investigation of the roles of the radical-trapping and enolizing agents on the chlorination process, to study the reactivity of carboxylic acids with different chain lengths, and finally to obtain quantitative kinetic information of the chlorination reactions.
Experimental Section Equipment Setup. The chlorination experiments were carried out in a 100-mL glass reactor equipped with a magnetic stirring bar and a electrical heater with temperature control. The chlorine gas and air were dispersed into the reaction mixture through separate glass sinters. A reflux condenser was placed on top of the reactor in order to prevent the escape of volatile liquid components. The gaseous components (including C12and HC1) which passed through the condenser were bubbled through aqueous NaOH. The chlorine feed was metered and controlled with a Wallace & Tierman Dry Feed Chlorinator (Type BA 057) with a capacity of 600 g/h. The air feed was metered with a rotameter. Experimental Procedure. The carboxylic acid (50-80 g), i.e., butanoic acid (Fluka, >99.5%), hexanoic acid (Fluka,>99%), octanoic acid (Fluka, >98%), decanoic acid (Fluka, >99%), or dodecanoic acid (Fluka, >98%), to be chlorinated was placed in the reactor. The air (AGA, medical grade) feed to the reactor was opened and the reactor content was heated to the given reaction temperature after which a given aliquot of chlorosulfonic acid (C1S03H,Fluka, >98%) was added using a l-mL syringe. The chlorine (Nokia Chemicals Oy, >99.9%) feed was started, and usually an immediate reaction was observed as a temperature elevation in the liquid phase. After about 10 min the temperature was usually maintained within 1-2 "C, In the case of stepwise catalyst addition the following
aliquots of ClS03H were manually introduced with a syringe at given time intervals. Samples of the liquid phase for analysis were withdrawn with a Pasteur pipette during the course of the reaction and placed in glass ampules. Chemical Analysis. The gas chromatograph (Varian 3300) was equipped with a single injector and a dual capillary column system: a 30-m DB-5 (J. w. Scientific, USA) connected to a flame ionization detector (FID) and a 30-m DB-1 (J. W. Scientific, USA) connected to an electron capture detector (ECD) sensitive to chlorinecontaining compounds. The carrier gas was helium with a velocity of 38 cm/s and a split ratio of 28. The temperature of the injector and detector was 270 and 300 "C, respectively. The following temperature programs were applied: C4A60 "C (5 min) + 5 "C/min, 160 "C (0 min); C6A 60 "C (0 min) + 5 "C/min, 160 "C (5 rnin); C8A 60 "C (0 min) + 8 "C/min, 180 "C (10 rnin); C l d 60 "C (0 min) + 10 OC/min, 210 "C (10 min); C12A60 "C (0 min) + 10 "C/min, 260 "C (0 min). The gas chromatographic data were computed with a Baseline 810 (Dynamic Solutions, USA) integration program. The carboxylic acids were analyzed as trimethylsilane (TMS) derivatives using N,O-bis(trimethylsily1)trifluoroacetamideas silylation agent and trimethylchlorosilane as catalyst. The FID-GC peaks were identified in separate runs on a VG Micromass 7070E gas chromatograph-mass spectrometer. The location of the chlorine substitution was determined in samples diluted with CDC13 by means of lH-NMR spectroscopy (JEOL GX400). Chemical shifts in the spectra are reported in ppm downfield from TMS. The sulfur contents of some samples of the product were determined using a sulfur analyzer (Leco Corp.). Physical Properties. The densities were determined with a 50-mL glass pycnometer, and the corresponding viscosities were measured using an Ostwald viscometer. Both measurements were carried out a t temperatures between 70 and 130 "C using a thermostated oil bath. Monochlorobutanoic acid (Fluka, >94%) was used to prepare the mixtures corresponding to different degrees of chlorination.
Results and Discussion Effect of Oxygen on the Chlorination Selectivity. Dodecanoic acid @,,A) was chlorinated at 130 "C in the presence of chlorosulfonic acid (ClSO,H), but without any radical scavenger. The chlorine feed was 40 g/h and ClS03H was added at 10-min intervals so that the fraction of C1SO3H (yCISO,H) in the solution became finally 0.0670. A broad spectrum of chlorinated products was detected by gas chromatography indicating that chlorine was added to a-,/3-, and y-positions of the carbon chain. During the course of the reaction the intermediate products were converted to final reaction products as shown in Figure 1. The dominating product was a-chlorododecanoic acid but its yield was low after 120 min reaction time the yield of a-chlorododecanoic acid was still below 20%. A t the complete conversion of dodecanoic acid the ratio between a-chlorododecanoic acid and the byproducts was approximately 1:4.3 (Figure 1). Introduction of air (10-15 mL/min) to the reactor resulted in a considerable increase in the a-chlorododecanoic acid formation rate and suppressed the byproduct generation (Figure 2). At the air flow of 15 mL/min a yield above 98% was achieved whereas at the air feed of 10 mL/min still some amounts of the byproducts were formed. Also the monochlorododecanoic acid formation rate was lower at the lower air feed (10 mL/min). The main product was a-monochlorododecanoic acid in all cases, and the amounts of a,a-dichlorododecanoic acid
Ind. Eng. Chem. Res., Vol. 31, No. 11, 1992 2427
polychlorinated
Y 0.8 -
0.4 .
o.2LL2 0.0 0
20
40
80
80
100
120mln
Figure 1. Chlorination of dodecanoicacid without oxygen; y = the yield. Conditions: temperature 130 OC; Clz feed 40 g/h; total amount of catalyst added Y C ~ O ~=H 0.0670.
0.8 -
0.8
-
0
20
40
80
20
40
60
80
100
120min
time
time
Y
0
0C.A
80
100
120mln
time
Figure 2. Effect of oxygen feed on the yield (y) of a-monochlorododecanoicacid. Conditions: temperature 130 OC; C1, feed 40 g/h; total amount of catalyst added y c ~ oH = 0.0670; air feeds ( 0 ) 0 mL/min, (I 10 )mL/min, (A) 15 rnlfmin.
formed were always very small. A similar experiment was carried out using butanoic acid. Because the reactivity of butanoic acid is higher than that of dodecanoic acid, a lower catalyst concentration was used (stepwise addition, final yCISOsH = 0.0195). The time developments of the yields of butanoic acid (C4A),monochlorobutanoic acid (MC,A), and dichlorobutanoic acid (DC4A)are presented in Figure 3. No byproduct formation was detected at the air feed of 15 mL/min. Small amounts of a,a-dichlorobutanoic acid were also formed. The above results are based on chemical analysis of the reaction producta by means of gas chromatography (GC), maw spectrometry (MS), and nuclear magnetic resonance spectroecopy (NMR). The 'H-NMR spectrum of a product mixture obtained in butanoic acid chlorination after a reaction time of 300 min at 130 OC is shown in Figure 4. The NMR spectrum clearly shows that the only reaction products were a-monochlorobutanoic and a,a-dichlorobutanoic acid. The 'H-NMR spectrum can be assigned as follows: S (400 MHz, CDC13), 4.25 (1 H, triplet, CH3CH,CHC1COOH); 2.30 (2 H, triplet, CH3CH,C&COOH); 2.00 (2 H, multiplet,
Figure 3. Chlorinationof butanoic acid in the presence of oxygen; y = the yield. Conditions: temperature 130 "C;Cl, feed 40 g/h; total amount of catalyst added yClsos~ = 0.0195; air feed 15 mL/min.
CH3C&CHC1COOH); 1.65 (2 H, sextet, CH,C&CH,COOH); 1.05 (3 H, triplet, 0.95 (3 H, triplet, C&CH,CHClCOOH); C&CH2CH2COOH). The amount of a,a-dichlorobutanoic acid was so small that it is hardly visible in the NMR spectrum (Figure 4),but obviously the small triplet at S = 1.15 ppm is caused by the protons attached to the ycarbon atom in a,a-dichlorobutanoic acid. Effect of Catalyst on the Chlorination Kinetics. At the next stage the effects of the catalyst addition policy and catalyst concentration were investigated. l b o catalyst addition methods were used; either all of the desired amount of ClS03Hwas introduced at the beginning of the chlorination or the same total amount was introduced stepwise in smaller aliquota into the reaction mixture during the course of the reaction. The final amounts of the catalyst were equal in both types of the experiment. Butanoic acid was chlorinated at 70 and 110 "C using the total amount of catalyst yew, = 0.041,chlorine mass flow 30 g/h, and air feed 15 mgfI/min. In the stepwise addition aliquota of 0.82 mol 9% catalyst were added five times at 60-min intervals. The developments of the amonochlorobutanoic acid yields are shown in Figure 5. Introduction all of the catalyst in the beginning of the reaction gives the highest initial reaction rate, but the reaction stagnated after approximately 4 h yMcd = 0.25 at 70 "C and Y M C A = 0.75 at 110 OC after 4 h and the reaction rata were practically zero. Analogous results have been obtained by Ogata et al. (1979),who chlorinated dodecanoic acid at 85 "C in the presence of chlorudfonic acid (YcBO8H = 0.20). Stepwise introduction of the catalyst gives a low initial reaction rate (Figure 5 ) , but enables forcing the reaction to ita completion. After a 5-h reaction time at 110 OC the yield of MC4A was about 0.9. Using this strategy to introduce the catalyst, it is possible to obtain high yields ale0 with low catalyst concentrations as the average amount of the catalyst in the experiment presented in Figure 5 was only about y c s O a H = 0.0082. These catalyst amounts are considerably lower than those reported by Ogata et al. (1979). Introduction of the catalyst stepwise results in an autocatalytic behavior as can be seen in Figures 3 and 5. The autocatalytic effect cannot be explained by an increase in the catalyst concentration in the reaction mixture since
2428 Ind. Eng. Chem. Res., Vol. 31, No. 11, 1992
f
I
j
I
I
"
~
?
"
'
~
'
I
'
"
'
'
"
~
5
6
'
~
'
"
~
"
/
'
"
"
"
"
I
"
'
"
'
~
'
'
I
'
'
"
'
'
2
3
4
"
'
l
1
Figure 4. NMR spectra of the product from the chlorination of butanoic acid. Conditions: reaction time 300 min; temperature 130 "C; Clp = 0,0195; air feed 15 mL/min. feed 40 g/h; total amount of catalyst added yCISOsH
40.10
/I
0.8.
I
I
/
..
0
50
100
150
200
250
300mln
time Figure 5. Effect of catalyst addition on the yield (y) of a-monochlorobutanoic acid. Conditions: temperature (0, m) 70 OC and (0, 0 ) 110 "C;
Clz feed 30 g/h; air feed 15 mL/min; total amount of catalyst added yCBOsH = 0.041. The catalyst w a introduced ~ either all in the beginning of the experiment ( 0 , O ) or in five aliquota at 1-h intervals (m, 0 ) .
chlorosulfonic acid is decomposed in the liquid phase: sulfur analysis of the reaction product obtained from butanoic acid chlorination showed that 75% of the initial sulfur was present in the liquid phase after a reaction time of 2 h. We therefore regard the autocatalytic effect as real and c a d by the reaction products MCA and DCA. This effect has not been reported previously for the long-chain carboxylic acids, but we have observed autocatalytic formation of monochloroacetic acid in using acetyl chloride as the catalyst (Martikainen et al., 1987). The catalyst concentration effect on the reaction rate was investigated in the chlorination of butanoic acid at 130 OC. The catalyst was added at 10-min time intervals so that the final amounts became y = 0.0099 and 0.0195. The chlorine feed was 40 g/h=the air feed was 15 mL/min. The yields of monochlorobutanoic acid and
0.0 0
20
-
0.00
40
60
80
100
120 min
time Figure 6. Effect of catalyst concentration on the yields b)of am o w and a,a-dichlorobutanoic acids. Conditione: temperature 130 OC; Clz feed 40 g/h; air feed 15 mL/min; total amount of catalyst added ycBoaH= 0.0099 (m, 0 ) and yc~ofl=0.0195 ( 0 , O ) .
dichlorobutanoic acid are shown in Figure 6. The results clearly show that the rates of both MC,A and DC,A formation increase with increasing catalyst concentration and similar autocatalytic formation kinetics can be observed for both these products, although the generation rate of MCIA is much higher. As a matter of fact, the ratio yDCIA/yMctA is independent of time as well an catalyst concentration, which can be seen in Figure 7. The simple relationship YDCA = ~ Y M C A (3) is satisfied by the experimental data and in the case of butanoic acid chlorination a ia approximately 0.06 (Figure 7). Similar dependence of the reaction rate on the catalyst concentration as for butanoic acid was also observed in the chlorination of dodecanoic acid. The chlorination took place at 130 "C using the chlorine feed 40 g/h, the air feed
Ind. Eng. Chem. Res., Vol. 31, No. 11, 1992 2429 6.0 0.04
1
In(p,,lkPa)
r
5.01
i
\
\
'
%C,A
\
0.03
4.0 0.02
3.0
I
0.01
J
0.00 0.0
0.4
0.2
0.6
0.8
1.0
2.0
-
1.o
%C,A
Figure 7. Interdependence of the yields
(y) of mono- and dichlorobutanoic acids. Conditione: temperature 130 OC; Cl, feed 40 g/h; air feed 15 mL/L; total amount of catayst added y c ~ =o ~ ~ 0.0099 (B) and YCWSH 0.0195 (0).
O.'
7 I 1 I
0.6
oo---
-
-
0
1
2
3
4
5
6
number of carbon atoms
Figure 9. Estimation of the vapor pressure p (), with different carbon numbers at 130 OC.
of carboxylic acids
was confiied by a duplicate experiment. The result can, however, be simply explained by consideration of the vapor pressures of the carboxylic acids. The vapor prwures (pd of C1-C5 carboxylic acids at 130 "C can be estimated using the Antoine equation:
Bi l n p w i = A i - -T + Ci
(4)
where the Antoine constants are given by Reid et aL (1987). The plot of the logarithm of the vapor pressure vs the acid carbon number is shown in Figure 9. The relationship is a straight line for the longer carbon chains (n 1 2) and the linear equation In pw = a. + a,n (5) 0
15
30
45
60 min
time
Figure 8. Yields (y) in a-chlorination of carboxylic acids (C,A) with different carbon numbers, n: (A) 4; (0)6; (+) 8; (m) 10; (v)12. Conditione: temperature 130 O C ; Clz feed 40 g/h; air feed, 15 mL/ min; total amount of catalyst added y c B 0 ~= 0.0178.
15 mL/min, and the catalyst amounts y w = 0.0349 and 0.0670. The reaction rate increased with increasing catalyst concentration, and the chlorination was autocatalytic up to MClzA yields about 0.7,after which the reaction rate decreased. After a reaction time of 120 min the yield of M C l d was above 0.98 (Figure 2). Effect of Chain Length on the Chlorination Kinetics. In order to investigate the effect of the alkyl chain length on the chlorination kinetics, butanoic, hexanoic, octaaoic, decanoic, and dodecanoic acids were chlorinated at 130 "C using the catalyst amount ycwa = 0.0178,the chlorine feed of 40 g/h, and the air feed of 15 mL/min. The time evolutions of the yields of respective monochlorocarboxylic acids are shown in Figure 8. The results show that all these acids are chlorinated autocatalytically; the yield curves have very similar parabolic forms for all monochlorocarboxylic acids. The reaction rate decreases with an increasing carbon number. Thia effect is expeded as the concentration of active chlorine-containing sites is lower for the long-chain carboxylic acids and the viscosity of the medium increases with increasing carbon number. One exception appears, however, in Figure 8. The chlorination rate of hexanoic acid is higher than that of butanoic acid. This anomaly of the butanoic acid curve
is valid at 130 "C at which temperature the parameters are a, = 6.654 and a, = -0.7905. Equation 5 predicts the vapor pressures of butanoic, hexanoic, and octanoic acids at 130 "C as 32.75,6.70,and 1.38kPa, respectively. These values clearly show that butanoic acid is volatilized during the experiments and it may have a partial pressure in the gas phase, whereas the partial pressure of hexanoic acid in the gas phase is probably very low and the partial pressures of octanoic, decanoic, and dodecanoic acids are practically zero at the actual conditions. The partial pressure of chlorine in the chlorination of butanoic acid is less than 100 kPa because of the coexistence of butanoic acid in the gas phase, and therefore also the chlorine concentration in the liquid phase is lower in the chlorination of butanoic acid than the chlorine concentration of hexanoic acid. We thus conclude that the lower chlorination rate obse~edfor butanoic acid is not a real kinetic effect, but caused by the lower liquid-phase chlorine concentration. Reaction Mechanism. Diverging opinions exist on the formation mechanism of chlorinated carboxylic acids. In the earlier studies of Ogata et al. (Ogata et al., 1977,1979; Ogata and Watanabe, 1979)it is proposed that a ketene might be the reactive intermediate in chlorination. The existence of a ketene in the reaction medium was proved by spectroscopic methods (Ogata et al., 1977): carboxylic acids and fuming H8O4gave a ketene which was thought to add chlorine to the double bond. The chlorine addition step was supposed to be rate determining. The ketene (RCH-4) formation was proposed to be acid catalyzed. The reaction mechanism is summarized as follows:
2430 Ind. Eng. Chem. Res., Vol. 31, No. 11, 1992
RCH *C/OH O ‘H
-H’ -H#
RCH=C=O
Cl2
00
RCHCICC (6) CI
On the basis of these assumptions the rate would be proportional to the unchlorinated carboxylic acid, to the acid catalysts, and to the chlorine concentration. In more recent studies Ogata and his mworkers suggest, however, that the reactive intermediate is a monoacyl sulfate (RR’CHC02S03H),which is formed in a reaction between carboxylic acid and chlorosulfonic acid. The existence of monoacyl sulfate was proved by Raman spectroscopy: the spectrum of a mixture of acetic acid and C1SO3H was identical with that of an authentic sample of CH3C02S03H(Ogata and Agachi, 1982). In a mixture of propionic acid and ClS03Ha quantitative evolution of HC1 was observed, indicating the formation of an addition compound consisting of propionic acid and a sulfate group. Monoacyl sulfate was assumed to be in equilibrium with ita enolic form, and the addition of chlorine to the double bond was proposed to be the rate-determining step. The reaction scheme can be summarized as follows:
According to both mechanisms 6 and 7 the chlorinated acid chloride is the intermediate that gives the a-monochlorocarboxylic acid, probably via a rapid chlorine exchange reaction with the unchlorinated carboxylic acid:
It could also be presumed that the unchlorinated carboxylic acid is enolized by acid catalysis, and chlorine is added to the enol form (Ogata et al., 1970,1975). The reaction steps would then be given by
-
RCH2Ceo OH O ,H ACH ‘C O ‘H
H‘
-H’ ‘-c
NOH RCHZC O ‘H
CI~ -t
RCHCICCo OH
+ HCI
(9)
The formation of a,a-dichlorocarboxylic acids is not discussed in the works of Ogata et al. In case of the chlorination of acetic acid it has been proposed (Sioli et al., 1979;Martikainen et al., 1987)that the reactive intermediate in chlorination is acetyl chloride undergoing acid-catalyzed enolization and chlorine is thus added to the enol to form chloroacetyl chloride, which reacts in rapid chlorine exchange reactions with acetic acid forming monochloroacetic acid and regenerating acetyl chloride. According to Sioli et al. (1979)the chlorine exchange reaction proceeds through an intermediate anhydride formed from chloroacetyl chloride and acetic acid molecules. The rate-determining step was proposed by Martikainen et al. (1987)and Salmi et al. (1988)to be the acid-catalyzed enolization of acetyl chloride. That reaction mechanism also explains the autocatalytic formation of monochloroacetic acid. It is proposed to depend on the fact that monochloroacetic acid is a stronger acid than acetic acid and thus it is also a more efficient catalyst in the enol formation reaction-therefore the reaction rate accelerates when the concentration of monochloroacetic acid increases in the reaction mixture. The formation of dichloroacetic acid has not been discusaed in the previous
works concerning the kinetics of the chlorination of acetic acid. Any consistent reaction mechanism for the chlorination of carboxylic acids has to include a catalyst regeneration step and to provide an explanation of the autocatalytic parallel formation of mono- and dichlorocarboxylic acids. Our results clearly show that rather small amounts of chlorosulfonic acid are sufficient to force the chlorination reaction to ita completion, and stoichiometric amounts of chlorosulfonic acid predicted by reaction step 7 are not necessary. We thus conclude that the essential reaction step might be the acid-catalyzed formation of an intermediate with a carbon-carbon double bond. The intermediate can in principle be a ketene, an acid chloride, or the enol form of the unchlorinated carboxylic acid. Chlorination of the enolic intermediate proceeds via a parallel path giving mono- and dichlorocarboxylic acids or corresponding acid chlorides simultaneously. If we assume that acid chloride is the reaction intermediate, a plausible reaction mechanism can be written as follows:
RCH=COOH + 2CI2 C ‘I
-
RCC12CeZI + 2HCI
(13)
The first reaction step (10)initiates the process giving acid chloride, which is enolized in step 11, and chlorinated in parallel in steps 12 and 13. The exchange reactions (14 and 15)regenerate the acid chloride and give the chlorinated carboxylic acids. In principle it could also be supposed that chloroacid chloride formed in step 12 could undergo an enolization analogous to the acid chloride enolization step (11)forming the enol form of chloroacid chloride. This enol could then react with the chlorine molecule to give dichloroacid chloride which through a reaction with the unchlorinated carboxylic acid gives dichlorocarboxylic acid and regenerates acid chloride. This assumption would predict consecutive formation of dichloroacid chloride from monochloroacid chloride. In a recent work concerning the selectivity of chlorination of acetic acid it is, however, shown that mono- and dichloroacetyl chlorides are formed parallel from acetyl chloride in the presence of chlorosulfonic acid (Miikih e l a , 1990). Therefore, we prefer the parallel chlorination path (12and 13). If we assume that the enol form of the unchlorinated carboxylic acid is the reactive intermediate, the following reaction mechanism can be written: R C H ,0 C ~= ~E H’~
RCH=COOH ‘OH
RCH * cO ,H O ‘H
+ 2C12
-
-
RCH=C O ,H O ‘H
(16)
RCC12Ceo + 2HCI OH
(18)
-H+
The first step (16) gives the enol form of the unchlorinated
Ind. Eng. Chem. Res., Vol. 31, No. 11, 1992 2431 Table I. Determination of the Empirical Parameters for the Chlorination of Carboxylic Acids with Different Carbon Numbers at T = 130 "C. Ea 21 acid bl b2 WSRS 0.495 X C,A 0.183 X 0.116 X 0.851 X C& 0.177 X 0.142 X 0.550 X lo-' C& 0.137 X 0.110 X 0.107 X 0.137 X lo4 Cl& 0.834 X CI2A 0.120 X lo-' 0.780 X lo4 0.230 X ~~
carboxylic acid, which is chlorinated in the subsequent steps (17 and 18). The catalytic acid in steps 11and 16 can be any acid in the system, but the main contributions to the acid catalysis may originate from chlorosulfonic acid and monochlorocarboxylic acid. Steps 11and 16 are initiated with the aid of C1S03H,but as the reaction proceeds the acid catalytic contribution of monochlorocarboxylic acid increases since ita concentration in the reaction mixture increases. The contributions of the unchlorinated carboxylic acid and the dichlorocarboxylic acids to the acid catalysis are probably minor since, on the other hand, the unchlorinated acid behaves like a very weak acid in the actual milieu and, on the other hand, the concentration of the dichlorocarboxylic acid is low compared to that of the monochlorocarboxylic acid. It could also be assumed that monochlorocarboxylic acid generated in step 17 might undergo acid-catalyzed enolization and add one chlorine atom to its carbon to carbon double bond. This hypothesis would, however, imply a consecutive formation route of dichlorocarboxylic acid from monochlorocarboxylic acid, which is in clear conflict with the experimental observations obtained in the chlorination of butanoic acid (Figure 7). Chloroeulfonic acid decomposes during the chlorination reaction causing the stagnation of the reaction as shown in the experiments presented in Figure 5. The chlorosulfonic acid decomposition reaction is somewhat unclear. One plausible decomposition step might be the acid chloride generation step (10). Furthermore, it is known (Kirk and Othmer, 1949) that ClS03H undergoes monomolecular decomposition to HC1 and SO3: ClS03H
-
HC1+ SO3
(19)
In the presence of hydrogen iodide C1S03H decomposes to SO2 and H2S04 (Ogata et al., 1979): (20) 2HI + 2C1S03H 12 + 2HC1+ SO2 + H2S04
-
Analogously it could be presumed that HC1-like HI does in reaction 20-catalyzes the decomposition of C1S03H. Reaction Kinetics at Constant Catalyst Concentration: Diffusional Effects. In stepwise addition of C1S03H to the reaction mixture, the yield of monochlorocarboxylic acid as a function of time has approximately a parabolic form as can be seen in Figures 6 and 8. The empirical equation YMCA
= bit
+ b2t2
(21)
was fitted to the yield vs time data for the C4-C12 acids presented in Figure 8. The corresponding parameter values for bl and bz are listed in Table I. The fit of the empirical model (21) to the experimental data is good. The reaction rate could then be estimated based on (21) according to dYMCA/dt bl + 2b2t (22) The reaction discussed gives indications concerning the concentration dependencies of the reaction rate. The protonation steps in the ketene mechanism (6), in the acid
chloride mechanism (10-151, and in the carboxylic acid enolization mechanism (16-18) is likely a rapid step, whereas the subsequent double bond formation and chlorination steps might be rate determining. The carboxylic acids are dominantly dimers in the actual milieu (Popovych and Tompkins, 1981). In the chlorination reaction steps the active form of the acid is, however, the monomer. Assuming that the dimer would strongly dominate in the monomer-dimer equilibria implies that the monomer concentration is proportional to the square root of the dimer concentration: [HA] KHA'/2~HA1/2 (23) where KHA is the equilibrium constant. In the ketene mechanism (6) and in the carboxylic acid enolization mechanism (1618) the concentration of the carbonium ion RCH2+C(OH)2would thus be proportional to the square root of the unchlorinated carboxylic acid concentration. In the subsequent acid catalytic steps in mechanisms 6 and 1618, any acid in the system can in principle contribute, but as previously discussed the main contribution is expected to come from monochlorocarboxylic acid and from ClS03H. Application of the monomer-dimer equilibria also for monochlorocarboxylic acid and chlorosulfonic acid predicta that the concentrations of the active forms of the acids are proportional to c M c A ~ / and ~ C C ~ S O & ~ according to eq 23. The ketene mechanism (6) and the carboxylic acid enolization mechanism (1618)would thus predict that the reaction rate is proportional to the product cCA'/~CMCA'/~. In terms of the yield of the monochlorocarboxylic acid (YMCA), this implies that the rate is proportional to (1- (1 + c~)YMcA)'/~YMcA~/~, where the contribution CYYMCA comes from dichlorocarboxylic acid according to eq 3. In practice the amounts of dichlorocarboxylic acids were small and thus the rate would be approximately proportional to the product (1- Y M C ~ ) ' / ~ Y M C A ' / ~according to the ketene formation and enolization mechanisms. When the reaction rates obtained from the empirical model (21) were plotted against (1- YMCA)'/~YMCA'/~, the conclusion was clear: the reaction rate is not proportional to (1 - Y ~ ~ A ) ' / ~ ~ M C A ' / ~ . This proportionality would predict a declining reaction rate after a certain MCA yield. In linear regression analysis, where the experimental rate (dyMcA/dt)was the dependent variable and the quantity (1 - YMCA)'/~YMCA'/~ was the independent variable, a systematic deviation of the calculated rates from the experimental rates was always observed for all the acids studied. The predicted reaction rates were too low at the lowest yields, too high at the intermediate yields, and too low at the highest yields. The correlation coefficients in linear regression for rate data obtained for C4-CiO carboxylic acids (Figure 8) were only 0.90.96, which is low taking into account that the experimental scattering in the data is very small. The acid chloride mechanism (10-15) would predict the rate to be proportional to the concentration of the carbonium species RCH2+CC10H. Application of the quasiequilibrium approximation on the proton donation step in (11)implies that the concentration of the carbonium species is proportional to the acid chloride and the acid catalyst concentrations. Acid chloride is present in the system only as very small amounts and it is regenerated in the reaction steps 14 and 15. Thus the acid chloride concentration is considered to be virtually constant in the reaction mixture-at least at low and moderate conversions. At very high conversions the acid chloride concentration decreases because the concentration of the unchlorinated acid approaches zero and the regeneration of acid chloride in steps 14 and 15 is retarded.
2432 Ind. Eng. Chem. Res., Vol. 31, No. 11, 1992 :2:.:
r
7
-
A
Table 11. Determination of the Lumped Parameters for the Chlorination of Carboxylic Acids with Different Carbon Numbers at T = 130 OC; Eq 32 acid p1 s(%) p' s(%) WSRS 43.4 0.423 x 10-3 CIA 0.198 X lo-' 8.1 0.909 X 61.1 0.745 X CaA 0.225 X lo-' 8.4 0.760 X C ~ A 0.201 x 10-1 2.3 0.562 x 10-3 17.8 0.356 x io+ 18.8 0.157 X lo4 Cl& 0.202 X lo-' 1.5 0.293 X ClzA 0.180 X lo-' 6.4 0.324 X loJ 368 0.196 X
0.010-
0
0.2
0.4
0.8
0.8
d YMCA Figure 10. Rates (r)of a-chlorination of carboxylic acids (C,A) with different carbon numbers (n: (A)4;(0)6; (*) 8; (m) 10;(V)12)as a function of the product yield. Conditions: temperature 130 "C; Clz feed 40 g/h; air feed 15 mL/min; total amount of catalyst added YCEO~H = 0.0178.
Also in the acid chloride mechanism (10-15) the dominant acid catalysts are probably monochlorocarboxylic acid and chlorosulfonic acid. The rate is thus proportional to the product cACCMCA'/', where cAc denotes the acid chloride concentration. At a constant acid chloride concentration the rate is proportional to cMcA'l2. Thus a plot Of the rate (dyMcA/dt) a g 8 b t YMCA'/~should give a Straight line. This test was applied on the kinetic data presented in Figure 8. The results are shown in Figure 10, and it is obvious that the rate is indeed proportional to YMCA'/~. Linear regression analysis, where dyMcA/dtwas the dependent variable and YMCA'/' the independent variable, was applied to the data for C4-Clo acids (Figure 10). The correlation coefficient was always larger than 0.994, and the deviations of the predicted reaction rates from the observed rates were purely stochastic. We can therefore conclude that the reaction rate is proportional to YMCA'/', which supports the validity of the acid chloride mechanism (10-15). The formation rate of monochlorocarboxylic acid in the liquid phase (rMCA) is related to the mass balance of the acid in the semibatch reactor: d n ~ c ~ / d= t~ M C A V L (24) where nMCA denotes the amount of monochlorocarboxylic acid and V, is the liquid volume. Assuming the acid chloride mechanism (10-15) to be valid and the acid-catalyzed enolization step (11) and the chlorination steps (12 and 13) to be rate limiting gives the following rate equation for monochlorocarboxylic acid formation: ~ M C A=
K'XC1S03H'12~ci~03H1/2C k+iKi*K)12Ci1/2 (25) 1+ CY
where K"is the fmt-order rate constant for the formation of acid chloride in the initiation step (lo), Kcao is the dimer dissociation equilibrium constant for C1Sb3H, k+i is the rate constant of the acid i catalyzed enolization step (ll),Ki* is the equilibrium constant of the rapid protonation step in enolization, and Ki is the dimer dissociation constant of the acid. The details of the derivation of the rate equation (25) are given elsewhere (Salmi et al., 1992). Parameter a relates the rates of mono- and dichlorocarboxylic acid formations. According to eq 3 we thus obtain (26) Neglecting all other acid catalytic effects than those caused ~DCA
Q'~MCA
by mono- and dichlorocarboxylic acids and chlorosulfonic acid and introducing the yields YMCA, YNA and yclsoa YMCA = ~ M C A / ~ O ,YDCA = ~ D C A / ~ O , YCISOaH = nCISO#/nO (27) and the concentration definitions CMCA = ~ M C A / V L CDCA , = ~DCA/VL, CCISO,H nClSOsH/VL (28) transform the rate equation (25) to
Introduction of the lumped parameters in rate equation 29 and insertion of the rate equation in the mass balance equation (24) of MCA gives finally (32) dYMcA/dt = PIYMCA"'+ P' Integration of (32) gives the time dependence of YMCA:
The parameters p1 and p' can now be determined by nonlinear regrwion analysis (Vajda and Valk6,1985). The data for the carboxylic acids presented in Figure 8 were used in regression analysis. The parameter valuea obtained are listed in Table 11,and the yields calculated with these parameter values are depicted in Figure 8. The standard deviation of parameter p1 is small (below 10%) for all carboxylic acids, whereas the standard deviation of parameter p'is clearly larger (Table 111, since this parameter is less identifiable than parameter pl. The main contribution to the reaction rate originates from parameter p1 (i.e., from the acid catalytic contribution of MCA), whereas the role of parameter p' (Le., the acid catalytic contribution of chlorosulfonic acid) is dominant only at the initial stage of the reaction. Rate equation 32 predicts correctly the proportionality of the reaction rate to YMCA'/'. A comparison of the calculated yields of monochlorocarboxylic acids with the experimental yields reveals that the fit of the model to the experimental data is very good (Figure 8). The deviation between experimental and calculated yields is usually below 55%. In spite of the good fit of the experimental data to rate equation 32, there remains the doubt whether the reaction
Ind. Eng. Chem. Res., Vol. 31, No. 11, 1992 2433 Table 111. Reaction Rates ( r )in the Chlorination of Carboxylic Acids with Different Carbon Numbers at Different Yields (T= 130 "C) r at YMCA = r at YMCA = r at YMCA= acid 0 (mi&) 0.6 (mm-') 1.0 (min-') 0.207 X lo-' C,A 0.909 X 0.149 X lo-' 0.233 X lo-' C& 0.760 X 0.167 X lo-' 0.206 X lo-' C& 0.562 X 0.147 X lo-' C1& 0.292 X 0.146 X lo-' 0.204 X lo-' 0.181 X lo-' C12A 0.324 X lo4 0.128 X lo-'
rate is affected by mass-transfer limitation of chlorine and whether the reaction rate is rapid enough compared to film diffusion to justify the use of balance equation 24 as it is based on the assumption that the reaction in the liquid film at the gas-liquid interface is negligible and all reaa tions occur in the liquid bulk phase solely. The importance of the fYm reactions can be estimated using the zero-order reaction concept. As can be seen in Figure 8, the reaction rates are approximately constant at the initial state close to zero conversion and also at higher conversions between reaction times 40 and 60 min. The reaction rates calculated using rate model 32 at yields YMCA = 0, YMCA = 0.5, and YMCA = 1.0 are listed in Table 111. If chemical reactions proceed in the liquid film, the following balance equation can be written for chlorine: (34)
According to the reaction stoichiometry the reaction of chlorine is related to the reaction rates of mono- and dichlorocarboxylic acids: (35) rciz = +MCA + ~ ~ D C A ) Relationship 35 gives rcl, = -(I
+ 2a)rMcA = 4"
(36)
where kttdenotes the local apparent zero-order rate constant. Solution of the differential equation (37)
with the boundary conditions cclp= ccl,b at z = 0
(38)
E = NCl2/Nc1;
Differentiation of the concentration profile expression (40) with respect to z and insertion of the derivative in the flux definition (41) give finally the enhancement factor
The chlorine concentration can be expressed with ycl,: CClZ = YCl@O/ VL) (45) and the rate constant can be expressed as follows: k" = k'(no/VL) (46) Thus k'= dyMcA/dtis valid locally at a constant reaction rate. Introduction of definitions 45 and 46 in eq 44 transforms the enhancement factor M E=l+ (47) Ycl; - YC1,b where M,the Hatta number, is here defined as kR.2
The values of the reaction rates (4') for different carboxylic acids at different conditions (at yields YMCA = 0, YMCA = 0.5, and YMCA = 1.0) are listed in Table 111. The chlorine diffusion coefficients were estimated using the Wilkdhang equation valid for solvent mixtures (Reid et al., 1987): (CXiI$iMi)'/2T
The association factor (&) was assumed to be 2 for all carboxylic acids, and the molar volume of chlorine in the liquid phase, V,(Cl2,L), was assumed to be 45.50 cm3/mol (Reid et al., 1987). The viscosities of pure carboxylic acids were calculated from the Andrade equation ln(Ti/cP) = A[ - B [ / T (50) and the viscosity of the mixture was estimated from the rule of ideal mixing In Vmix = Cri In Ti
and
(43)
(51)
I
cclz = ccl,S at z = SL
(39)
gives the chlorine concentration profile in the liquid fih cr!,.(z) --' =
The chlorine flux at the gas-liquid interface is given by Ncl, = -Dc12(
2)
r=O
In the case of pure physical diffusion the chlorine flux is given by
The enhancement factor is defined as the ratio between chemical and physical absorption:
Viscosity data for long-chain fatty acids are available in the literature (Reid et al., 1987; Liew et al., 1991), but not for the chlorinated acids. Therefore we measured the viscosity for C4A and MC4A as well as for their mixtures at 70,90,110, and 130 OC (Figure 11). The corresponding Andrade parameters are listed in Table IV, where also the temperature dependencies of the densities of these acids are given. The calculated viscosities at different temperatures are shown in Figure 11. The deviations between experimental and calculated viscosities were below 1 ?% in the temperature interval 70-130 OC. Estimation of the chlorine diffusion coefficients in butanoic and monochlorobutanoic acid revealed that essential changes in the chlorine diffusion coefficients occur during the course of the reaction: the diffusion coefficient in butanoic acid is Dc, = 1.04 X lo4 m2/s at 130 "C whereas the diffusion coefficient in a-monochlorobutanoicacid is Dell = 0.67 X mz/s (Table V). Viscosity data were not available for other chlorinated acids (C6-C12),but a rough estimation of the chlorine diffusion coefficient in the monochlorocarboxylic acids can be performed sup-
2434 Ind. Eng. Chem. Res., Vol. 31, No. 11, 1992 Table IV. Experimentally Determined Andrade Parameters (In q/cP = A'- B'/T) and Density Function Parameters (p = B"T) for Butanoic Acid at Different Degrees of Chlorination"
A"
+
Andrade param A' B' (K) 1514 -4.706 1860 -5.260 2028 -6.552 2137 -5.637
deg of chlorination (5%) 0 33 67 100 O R 2
RZ 0.99999 0.99349 0.99977 0.99960
density function param A"(ke dm-? B" (ke dm+ K-9 1.265 -0.00101 1.416 4.00130 1.526 4.00116 1.508 4.00105
N 4 4 4 4
R2 ~~
N 4 3 4 4
~
0.99959 0.997 17 0.99833 0.99396
= the square of the correlation coefficient. N = number of observations. The validity region: 5" E [343,403]K.
Table V. Viscosities ( q ) , Chlorine Diffusion Coefficients (Dee), and Hatta Numbers (M) for Unchlorinated, Partly Chlorinated, and Fully Chlorinated Carboxylic Acids at 130 O C deg of chlorinaacid tion (%) C4A 0 100 50 Cl2A 0 100 50
n (cP) Dpl*(m2sd) 0.3866 1.037 X 0.7100 0.666 X lo+' 0.5238 0.836 X 1.4055 0.365 X 2.5810b 0.221 X 1.9047b 0.284 X
M
E-la
5 OE-02
~
p 2'0::03
Pl
I
3 OE-02
1.OE-03
0.7308X 1.6 X 0.2596 X 5.5 X 0.1487 X 3.2 X 0.7414X lo4 1.6 X 0.6817 X 1.4 X 1O-I 0.3745 X 8.0 X 1 .OE-02
"Estimated values based on ClZ solubility in CH,COOH and CH2C1COOH. bEstimatedvalue based on the ratio of viscosities of CIA and ita chlorinated producta.
0
2
4
8
8
10
'*
12
'
O.OE*OO
number of carbon atoms Figure 12. Dependence of the rate parameters p1 (m) and p' ( 0 )on the carboxylic acid carbon number.
i
I
0.0' 60
80
100
120
140 'C
temperature
Figure 11. Viscosity (7) of butanoic acid at different degrees of chlorination as a function of temperature.
posing that the ratios between the viscosities of the unchlorinated and the chlorinated acids are equal to the corresponding ratio calculated for butanoic and monochlorobutanoic acid. Obviously the ratio between the viscosities is closer to 1 for longer chain (X4) carboxylic acids, but the use of the ratio obtained for the butanoic acid-monochlorobutanoic acid system provides a conservative estimate of the diffusion coefficients. The chlorine diffusion Coefficient in dodecanoic acid is Da, = 0.36 X 10-8 mz/s at 130 "C, and the corresponding value in monochlorododecanoic acid would thus be approximately Dclp = 0.22 X m2/s at 130 "C (Table V). The diffusion coefficients enable the calculation of the Hatta number (M). The values of M for the chlorination of butanoic and dodecanoic acids are presented in Table V. They were obtained using the rate constants given in Table 111, the diffusion coefficients given in Table V, and the liquid film thickness $. = 10-4 m. The Hatta numbers are very small, especially in the beginning of the reaction: e.g., for butanoic and dodecanoic acids M is 7.3 X lo+ and
7.4 X lo-', respectively, at the initial stage of the reaction. At conversions of 0.5, corresponding approximately to the end conditions of the experiments, the Hatta numbers are larger than 10-4 for the chlorination of C4Aand CIA(Table V). An exact estimation of the enhancement factors was not possible since chlorine solubility data in the actual carboxylic acids were not available, but crude estimates can be obtained using the chlorine solubility data of Martikainen et al. (1987) in acetic acid and monochloroacetic acid. These experimental results revealed that chlorine is equally soluble in acetic acid and monochloroacetic acid, and a temperature function of Henry's constant was determined. Using the value of Henry's constant given by Martikainen et al. (1987) at 130 "C (H= 212.4 bar) suggests that the mole fraction of chlorine at the gas-liquid interface would be yc1," = 4.7 X at the chlorine pressure of 1 bar. This mole fraction at the interface would predict and 5.5 X that E - 1 would vary between 1.6 X in butanoic acid chlorination conversions ranging from 0 to 1. The corresponding variation of E - 1 in dodecanoic acid chlorination would be 1.6 X 10-"1.45 X lo-' (Table VI. In the present case, however, the dodecanoic acid conversion was less than 0.3 at the end of the experiment. E - 1 would be 0.08at the conversion of 0.5, and it can thus safely be concluded that the enhancement factor is close to 1 in the experimental conditions for all the studied carboxylic acid chlorinations. This implies that pure physical diffusion dominates in the liquid film, and the reaction proceeds mainly in the liquid bulk. Thus the use of the mass balance equation (24) is justified. A t the following stage the dependence of the kinetic parameters p1 and p' on the carbon number of the carboxylic acids was examined. The parameters given in Table I proved to depend linearly on the chain length according to p = co - cln (52) where n denotes the carbon number and p is p1or p '. The dependence of p1 and p ' on the carbon number is illustrated in Figure 12. The values of co and c1 were obtained
Ind. Eng. Chem. Res., Vol. 31, No. 11, 1992 2435 Table VI. Determination of the Dependence of the Rate Parameters on the Carboxylic Acid Carbon Number (Eq 52); p = c,, cln, Where p Is either p1or p'and I ) Is the Carbon Number
+
p
;'
co (min-')
0.02622 0.00156
c1 (min-') -0.6687 X -0.1227 X
scl (%)
R2
26.1 4.62
0.8796 0.9957
Estimated Rate Parameters for the Chlorination of the C,-C,I Carboxylic Acids acid CIA C& C& C,A C&
p1
0.0236 0.0229 0.0222 0.0215 0.0209 O.'
0.8
P' 0.102 X 0.903 X 0.780 X 0.657 X 0.535 X
acid C& Cl& CllA C12A
p1
0.0202 0.0195 0.0189 0.0182
P' 0.412 X 0.290 X 0.167 X 0.438 X lo-'
* i
r
A
0.4
0.3
0
15
30
45
60 min
time Figure 13. Simulated yields (y) of a-chlorocarboxylic acids with different carbon numbers (n = 6-12) using the rate parameters in Table 11. Symbols as in Figure 8.
by linear regression using the parameter values listed in Table I. Parameters co and c1 for p1 and p' are given in Table VI. The parameters for C&-C12A were taken to regression analysis, but those for CIA were omitted due to the volatilization effect. The simple linear dependence (52) enables an eatimation of the rate parameters not only for the acids studied experimentally but also for C a , C,A, C&, and CllA as listed in Table VI. Generally it can be concluded that the dependence of parameter p' on the chain length is much stronger than that of parameter p1 (Table VI). This implies that the initial rates of the different carboxylic acids are rather different whereas the differences in final reaction rates are much smaller. Finally, the yields of the monochlorocarboxylic acids MC&-MC12A were computed as a function of reaction time at the experimental conditions at 130 "C using the parameter vduea obtained from eq 52. The calculated and experimental yields are compared in Figure 13. The comparison reveals that the rate model (32) with the carbon number dependent parameters p1and p'b able to predict the semibatch reactor behavior at thb experimental conditions; the deviations of the predictad yields from the experimental yields are slightly higher than when the specific parameters p1and p'are used (Table 111,but are still within a few percent. This observation leads us to believe that the rate model might be interpolated also to the acids not studied in this work with carbon numbers within the range from 4 to 12. Reaction Kinetics at Decreasing Catalyst Concentration. The plausible catalyst decomposition reactions
Table VII. Parameters of the Catalyst (C180,H) Decomposition Kinetics, Eq 58 temp ("CI yo,clsosH a ( m i d ) s (%) WSRS 70 0.041 0.0144 3.97 0.1291 x 110 0.041 0.0230 4.95 0.2904 x 10-l
19 and 20 indicate that the decomposition kinetics could be described by rCISOsH -kcCIS08Hm (53) where m 1 1 as discussed earlier. The mass balance of the catalyst in the semibatch reactor is dnckx@i/dt = ~ C ~ S O ~ H V L (54) Introduction of the dimensionless quantity yCISOsH = nCIS03H/n0(27) transforms the mass balance to
The quantity no/ VL is assumed to be approximately constant during the reaction and the definition a = k(nO/VL)"'-l (56) is introduced. Integration of the balance equation (55) gives Y0,ClSOIH
YClSOsH
=
[1
+ a(m - l)Yo,CISOsHm-lt]
l'(m-l)
(57)
where Y0,ClS03H is the initial mole fraction of the catalyst. For a firaborder (rn = 1) decomposition eq 57 is converted to YClSOsH = Y0,C1S03Heqt (58) The data gathered from the experiments with butanoic acid (Figure 5 ) were treated as follows: first parameters p1and p' were determined according to rate equation 32 based on the experiments with stepwise catalyst addition. These parameters p1 and p' corrected with the catalyst concentration were then used in estimation of the decomposition parameter a according to eq 57 using the experimental data obtained when the catalyst was introduced only in the beginning of the chlorination (Figure 5). Parameter estimation revealed that the best fit to the experimental data was provided by the first-order decomposition model (m= 1). The second-order decomposition kinetics gave a poor fit to the experimental data. The parameter values obtained for first-order catalyst decomposition at 70 and 110 OC are listed in Table VII. Finally, the yields of monochlorobutanoic acid were computed using the experimental conditions and the parameter values listed in Table 11. The development of the MCIA yields as a function of time is illustrated in Figure 5. The figure shows that the first-order catalyst decomposition kinetics can provide an explanation of the stagnation of the chlorination rate in the case that the catalyst is introduced only in the beginning.
Conclusions The experimental results showed that the carboxylic acids with carbon numbers ranging from 4 to 12 can be a-chlorinated selectively if a radical scavenger (e.g., oxygen) and a strong enolizing agent (e.g., chlorosulfonic acid) are continuously introduced to the reaction mixture. In the absence of the radical scavenger, unselective radical chlorination takes place, giving a broad spectrum of products (Figure 1). The key point in achieving a high yield of monochlorocarboxylic acid is the continuous or stepwise addition of the enolizing agent instead of intro-
2436 Ind. Eng. Chem, Res., Vol. 31, No. 11,1992
duction of the catalyst only in the beginning of the reaction. Continuous addition during the come of the reaction compensates the catalyst decomposition, and the chlorination reaction is forced to ita completion (Figures 2 and 5)* In stepwise addition of the enolizing catalyst the reaction rate in the semibatch reactor is autocatalytic for the C4-CI2 carboxylic acids (Figure 8). The autocatalytic effect is supposed to be due to the acid catalytic enolization of the active reaction intermediate, acid chlorode: the monochlorocarboxylic acid created in the reaction is a much stronger acid than the unchlorinated carboxylic acid giving thus a higher rate to the enolization step (ll), which is likely the major rate-controlling step in chlorination. Mono- and dichlorocarboxylic acids are formed via a parallel path as shown in the chlorination of butanoic acid (Figure8 3 and 5). The estimation of the Hatta numbers and the enhancement factors (Table V) using the experimentally determined chlorination rates (Table III) and the diffusion coefficients of chlorine estimated from the viscosity data (Table IV) revealed that the chlorination reactions can be treated as slow reactions, which mainly take place in the liquid bulk phase. Thus the carboxylic acid mass balance equation (24) written for the liquid bulk phase was used in the quantitative treatment of the experimental data. The kinetic behaviors in the chlorinations of the C4-C12 carboxylic acids are principally similar for all these dcids (Figure 8). The rate equation (29) based on the assumption of the validity of the chlorination mechanism (1&15) is able to explain the reaction kinetics in the semibatch reactor. The kinetic parameters included in rate equation 32 proved to be linearly dependent on the chain lengths of the carboxylic acids.
Acknowledgment The authors are grateful to Dr. R. Sjoholm, who performed the NMR analysis, Mr. M. Reunanen, who carried out the mas8 spectroscopic analysis, and to Mr. S. Lindberg, who determined the densities and viscosities of the carboxylic acids.
Nomenclature a = kinetic perameter, eq 56 ao, a1 = parameters in eq 5 A, B, C = parameters in Antoine vapor pressure equation (4) A ’, B‘ = parameters in Andrade viscosity equation (50) A”, B” = density function parameters, Table IV bl, b2 = parameters in eq 21 c = analytical concentration co, c1 = parameters in eq 52 D = diffusion coefficient E = enhancement factor H = Henry’s constant k = rate constant k’ = zero-order rate constant on mole fraction basis, eq 46 k ” = zero-order rate constant, eq 36 K = equilibrium constant K” = lumped constant, eq 25 m = order of the catalyst decomposition kinetics M = molar mass, eq 49 M = Hatta number, eqs 47 and 48 n = amount of substance no = initial amount of the unchlorinated carboxylic acid N = number of observations in regression N = component flux in chemical absorption N‘ = component flux in physical absorption p = kinetic parameter, eq 52 p l , p ’ = kinetic parameters, eqs 30 and 31
p v p = vapor pressure r = reaction rate R = correlation coefficient s = standard deviation (% ) t = time
T = temperature
VL = liquid volume V,(CI,,L) = molar volume of chlorine in liquid phase 3t = mole fraction, eqs 49 and 51 y = yield or normalized amount of substance, eq 27 z = length coordinate of the liquid film a = ratio between rate parameters, eq 3 6 = chemical shift bL = liquid film thickness 7 = viscosity p = density 9 = association factor [ ] = concentration Subscripts and Superscripts b = liquid bulk property calc = calculated value
i = component index L = liquid phase mix = mixture property n = number of carbon atoms obs = observed value 0 = initial value + = forward reaction step Abbreviations AC = acid chloride CA = unchlorinated carboxylic acid
C,A = unchlorinated carboxylic acid containing n carbon atoms DCA = a,a-dichlorocarboxylic acid HA = acid MCA = a-monochlorocarboxylicacid WSRS = weighted sum of residual squares
Literature Cited Groeeins. P. H. Unit Processes in Oraanic - Synthesis: McGraw-Hill: Giw York, 1958. Kirk, R., Othmer, D., (Eds. Encyclopedia of Chemical Technology; Mac Printinn ComDanv: New York. 1949 Vol. 3. D 885. Liew, K. Y.; Seng, C. E.; i a u , E. K. Viscositi& of Som; Long-chain Fatty Acids and Their Relationship with Chainlength. J. Am. Oil Chem. SOC.1991,68,488. MU-Arvela, P. Unders8kning av metoder f6r framstUning av ren monoklorattiksyra [Investigationof methods for synthesis of pwe monochlorpacetic acid (in Swedish).] Lic. Technol. Thesis, Ab0 Akademi, Abo, Finland, 1990, p 50. Martikainen, P.; Salmi, T.; Paatero, E.; Hummelstadt, L.; Klein, P.; DamBn, H.; Lindroos, T. Kinetics of Homogeneous Catalytic Chlorination of Acetic Acid. J. Chem. Tech. Biotechnol. 1987,40, 259.
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Receiued for reuiew April 28,1992 Revised manuscript received July 15, 1992 Accepted August 7, 1992
Reaction Kinetics of Ca-Based Sorbents with HCl Brian K. Gullett' Air and Energy Engineering Research Laboratory, US.Enuironmental Protection Agency, Research Triangle Park, North Carolina 27711
Wojciech Jozewicz Acurex Enuironmental Corporation, P.O. Box 13109, Research Triangle Park, North Carolina 27709
Leonard A. Stefanski Department of Statistics, North Carolina State Uniuersity, Raleigh, North Carolina 27695
The kinetics of the reaction between CaO and HC1 were investigated under conditions that minimize bulk mass transfer and pore diffusion limitations. Reactivity data from 0.2-to 1-sexposure to 5000 ppm HC1 in a fixed bed reactor were analyzed by a shrinking core model of diffusion and chemical reaction control, either singly or in combination. Between temperatures of 150 and 350 OC, the reaction is controlled by gaseous diffusion through the developing product layer. The apparent activation energy is about 28.1 kJ/mol(6.7 kcal/mol), and the reaction is first order with respect to HC1 concentration. Reactivity is a minor function of the measured particle size and surface area, likely due to the agglomerative nature of the individual grains that comprise the particle structure and complicate the interpretation of these measured values. Extrapolation of these results to the high-temperature, furnace sorbent injection process provides preliminary agreement with pilot-scale tests.
Introduction The reaction of Ca-based sorbents with hydrogen chloride (HC1) is of interest for control of acid gas emissions from combustion procesees, most notably municipal waste combustion (MWC) and hazardous waste incineration ("I It ) is also .likely that high-temperature removal of HC1 will limit the downstream formation of chlorinated dioxins and furans (Gullett, 1991). Current methods of HCl removal often use dry Ca-based sorbent injection into the post-flame region. This process closely parallels the work that has been done to develop sulfur dioxide (SO,) control technology by furnace injection of Ca-based sorbents. However, unlike the Ca/S02 reaction, fundamental information concerning the kinetics, the controlling mechanism, and the reaction products of the Ca/HC1 reaction i~limitsd. In the case of the Ca/SOz reaction, this information has proved crucial toward optimizing sorbent reactivity with SO, through optimizing design/operation parameters such as injection temperature, injector location, and sorbent loading and sorbent parameters such as particle size, porosity, and surface area. This work examines the kinetica of the reaction between Ca-based sorbents and HC1 with the intent of providing information that will be useful toward optimizing the removal of HC1 from flue gas. Injection of calcium hydroxide [Ca(OH),] into hot (>400 "C) flue gas results in the loss of H20 through Ca(OH), F? CaO + H20 (1) Similarly, injection of calcium carbonate (CaC03) (>650 "C) results in calcination: CaC03 s CaO + COz (2) The resulting calcium oxide (CaO) is more porous and of
higher surface area than the original Ca(OH), or CaC03, unless prolonged heating results in sintering. CaO will react with HC1 to presumably form calcium chloride (CaC1,) as such CaO + 2HC1 e CaCl, + HzO (3) Free energy calculations suggest that (3) is favorable over the full range of temperatures in a combustor. &uilibrium concentrations of about 500 ppm HC1(10% HzO) likely impose an upper limit on the reaction at around 800 "C. Since 782 OC is the melting point of CaCl,, this also may constrain the upper practical temperature range of the reaction. The reaction can be considered irreversible since C0.1 ppm HC1 is in equilibrium with the assumed CaCl, product at the peak teat temperature, 350 OC (from 1 to 10% HZO). Various types of Ca-based systems for HC1 removal are discussed in the literature (Buekens et al., 1984;MayerSchwinning and Laibold, 1989;Ellison, 1989;Schmal et al., 1989),and a few experimental results are reported, covering a wide range of operating conditions and systems (Verbeek et al., 1987;Gullett et al., 1989;Karlsson et al., 1981; Walters and Daoudi, 1987; Weinell et al., 1992). Numerous waste combustion facilities use some type of sorbent injection for acid gas control (White and Vancil, 1989),yet very little fundamental kinetic information is available to assist design efforts and enhance large-system performance. A few reported results elucidate some of the important parameters regarding the Ca/HC1 reaction. Karlsson et al. (1981)determined a fmt-order rate dependency upon HC1 concentration when running breakthrough analyses with HCl and Ca(OH), from 150 to 400 "C. Similarly,
Q8SS-~sS5/92/2631-2437$03.00/0 0 1992 American Chemical Society
