Ind. Eng. Chem. Res. 1990,29, 1725-1728
phenoxide and potassium phenoxide in the absence of steam was verified in the miniflow pyrolysis unit, using procedures analogous to that used to verify the thermal stability of HCP. From the sodium and potassium phenoxide residues, anisole yields were 97% and 99% of the theoretical values, respectively. Comparison of these values to those of the starting phenoxides shows that essentially no phenol was liberated by heating to 350 OC in the absence of steam. Steam Stripping of Sodium Phenoxide and Potassium Phenoxide. Steam stripping of the alkali metal phenoxides was carried out as described above. From the steam-stripped sodium phenoxide residue, the anisole yield was 15% of the theoretical. An 85% yield of anisole was obtained upon methylation of trapped phenol and steam. This confirmed that 85% of the phenol moieties initially present in the salt were removed by steam stripping. Phenol balance was quantitative in this experiment. Comparable steam stripping of potassium phenoxide afforded an 86% phenoxide conversion and a 101% phenol balance. In another experiment, potassium phenoxide (1.35 g, 10.2 mmol) was steam stripped as described above, and the residue was dissolved in 50 mL of C02-freewater. The resulting solution was titrated with 0.50 N HC1, using a pH meter to monitor the titration. The smooth curve obtained from this titration is shown in Figure 5. Similarly,
1725
no carbonate and a 92% recovery of base were obtained from steam stripping of sodium phenoxide. The lower recovery in the latter case may be due to mechanical losses.
Acknowledgment The miniflow pyrolysis unit was constructed by R. H. Schlosberg and A. Kurs, who made it available for these studies. The advice and assistance of these workers are gratefully acknowledged. Also, very helpful discussions with Prof. H. C. Brown (Purdue University) and the late Prof. R. Pettit (University of Texas, Austin) are acknowledged with thanks. Registry No. Phenol sodium salt, 139-02-6;phenol potassium, 100-67-4;phenol, 10895-2;potassium hydroxide, 1310-58-3;phenol calcium salt, 5793-84-0.
Literature Cited Fisher, R.; Erhardt, U. Gesammelte Abh. Kennt. Kohle 1919, 4 , 237-263. Kornblum, N.; Lurie, A. P. J. Am. Chem. SOC.1959,81, 2710. Schlosberg, R. H.; Scouten, C. G. US. Patent 4 256568, 1981. Schlosberg, R. H.; Scouten, C. G. Energy Fuels 1988, 2, 582-585. Scouten, C. G. US.Patent 4595489, 1986. Scouten, C. G . US.Patent 4551637, 1987.
Received for review August 10, 1989 Accepted December 5, 1989
Reaction Mechanism and Kinetics of Aqueous Solutions of 2-Amino-2-methyl-1-propanol and Carbon Dioxide Erdogan Alper Chemical Engineering Department, Kuwait University, P.O. Box 5969, 13060 Safat, Kuwait
The mechanism and kinetics of the reaction between aqueous solutions of COz and a sterically hindered primary alkanolamine, 2-amino-2-methyl-1-propanol (AMP), were investigated a t 278-298 K by using a stopped flow technique. Experiments were also carried out with monoethanolamine (MEA) solution, which is the sterically unhindered form of AMP. The corresponding second-order rate constants a t 298 K were found to be 520 and 5545 m3/ (km01.s) for AMP and MEA with activation energies of 41.7 and 46.7 kJ/mol, respectively. On the basis of this rate constant, the reaction was considered to be formation of carbamate which subsequently hydrolyzed.
Introduction 2-Amino-2-methyl-1-propanol(AMP) is a commercially available primary amine which is the sterically hindered form of monoethanolamine (MEA). Sterically hindered amines cannot form stable carbamates, leading to much higher carbonation ratios (moles of C02/mole of amine). However, the steric hindrance lowers the reaction rate of C02,which may be undesirable. Although thermodynamic properties and industrial applications of sterically hindered amines have been discussed in detail in previous literature (Sartori and Savage, 1983; Sartori et al., 1987; Chakraborty et al., 1986), studies on the reaction kinetics are limited. In the case of AMP, a number of heterogeneous gas absorption studies were carried out previously. Sharma (1965) reported a second-order rate constant of 1048 m3/(kmol.s) at 298 K. Sartori et al. (1987) and Chakraborty et al. (1986) deduced a rate constant of about 100 m3/(kmol.s) at 313 K. The recent study by Yih and Shen (1988) reports a rate constant of 1270 m3/(kmol-s)at 313 K. Absorption rates of C02and H2S(but not direct kinetic data) were also reported for AMP by Zioudas and Dadach (1986). These gas absorption studies were all analyzed by OSSS-5SS5/90/2629-1725$02.50/0
using the methodology of “gas absorption with pseudo first order reaction”, and the agreement between them cannot, in general, be considered as satisfactory. Recently, Bosch et al. (1989) suggested that carbon dioxide absorption into solutions of sterically hindered amines cannot be simplified as above and the process should be considered as a case of mass transfer with parallel reversible reactions. Their rigorous numerical solution interprets the gas absorption results of Chakraborty et al. (1986) and yields a completely different conclusion. It seems, therefore, that there is a need for obtaining data from direct techniques which do not involve mass transfer. Further, there appears to be a confusion about the exact reaction mechanism since some of the authors presume no carbamate formation (Yih and Shen, 1988) while others (Sharma, 1965) consider the reported data as the forward reaction rate constant between carbon dioxide and AMP. On the other hand, Chakraborty et al. (1986) presumes the reaction to be the hydration of C 0 2 which is catalyzed by AMP as in the case of tertiary amines. The aim of this paper is to report the results obtained by a direct technique (that is, stopped flow experiments) 0 1990 American Chemical Society
1726 Ind. Eng. Chem. Res., Vol. 29, No. 8, 1990
which does not involve gas absorption so that the findings correspond to the homogeneous reaction rate between aqueous solutions of C02and AMP. A similar study with monoethanolamine was also carried out not only for the purpose of comparison with AMP (which is the sterically hindered form of MEA) but also to ascertain the validity of this direct stopped flow technique.
Reaction Kinetics Following the proposed mechanism by Danckwerts (1979),and subsequent considerable experimental evidence (for instance, Blauwhoff et al. (1983)), the general consensus for the reaction of C02with primary and secondary alkanolamines is the formation of zwitterion intermediate, rather than one-step carbamate formation. For an amine with a stable carbamate (e.g., for MEA), the following reaction takes place: RNH2 + C02 RNH,+COO(1) RNH,+COO-
+B
-
RNHCOO-
+ BH+
0.010
CMEAI
1
I
0 IW
o 500
. kmdlm'
Figure 1. Observed pseudo-first-order rate constant as a function of MEA concentration at three different temperatures.
(2)
Here B is a base which could be amine, OH-, or H 2 0 (Blauwhoff et al., 1983). If the carbamate ion is unstable, the following reaction proceeds subsequently: RNHCOO-
+ H2O
RNHz
+ HC03-
(3)
Yih and Shen (1988) presumed that AMP cannot form carbamate. Hence, the following alternative for reaction 2 was proposed by them: RNHz+COO- + H2O HCOB- + RNH3+ (4) Q
Another alternative mechanism has been proposed by Chakraborty et al. (1986): RNH2 + COa + H20 e RNH3+ HCOB(5)
+
Equation 5 is similar to reactions of tertiary amines such as methyldiethanolamine (MDEA), and the only effect of AMP could be a base catalysis which should not, however, yield such high reaction rates as reported by Sharma (1965) and Yih and Shen (1988). The evidence in support of proposed mechanisms of reactions 4 or 5, which do not involve carbamate ion formation, was that the NMR spectra of a solution of AMP containing chemically bonded C 0 2 have no noticeable carbamate peak (Chakraborty et al., 1986).
Experimental Section The experimental setup consisted of a standard stopped flow equipment with a conductivity detection system which could be used to measure directly the intrinsic rate of a rapid homogeneous reaction (in this case, that between aqueous solutions of COP and AMP). Equipment was thermostated at f O . l K, and experiments were carried out a t 278, 288, and 298 K. Reagent grade MEA and AMP were obtained from Fluka, Switzerland, and they were used as supplied. The concentration of amine was always much in excess of that of COP(usually the molar ratio was about 20 to 1)and ranged between 0.013 and 1.5 kmol/m3. Other pertinent details of the equipment and the experimental procedure can be found elsewhere (Alper, 1990). Results and Discussion Experiments with MEA showed that the conductivity of the solutions leveled to a constant value, which indicated a stable carbamate formation as the final reaction product. In the case of AMP, a constant level of conductivity was first obtained during the sweep time of the run; however, conductivity changed at a much slower rate at times much larger than sweep time. This observation could be con-
0 0003
0 0334
0 0035
111
I
0 0036
0 0037
K-'
Figure 2. Arrhenius plots for C0,-aqueous MEA and COP-aqueous AMP systems. Table I. Observed Reaction O r d e r s in Amine a n d Second-Order Rate Constants reaction order k,, m3/(kmol-s) temp,K MEA AMP MEA AMP 1.14 1435 151 278 1.04 2867 1.07 1.14 288 300 5545 1.04 1.15 502 298
sidered as a support for a mechanism where carbamate ion was first formed and hydrolyzed subsequently according to eq 3. Since the concentration of C 0 2was always kept much less than that of amine, all experimental traces were analyzed according to a simple first-order kinetics and by using an algorithm due to Marquardt (1963). Good fits were obtained, and the observed pseudo-first-order rate constant, kobs, was analyzed according to the equation = kobs[C021
(6)
COZ-MEA Results Figure 1shows the pseudo-first-order rate constant kob as function of MEA concentration a t three diffeent temperatures. The reaction orders in MEA were obtained by regression analysis, and they are shown in Table I to be very near unity. It can therefore be concluded that the reaction is first order both in C 0 2 and MEA, which agrees with previously reported data. The second-order rate constants, which were obtained by linear regression analysis, are shown in Figure 2 as an Arrhenius plot. The data yielded a second-order rate constant of 5545 m3/ (kmo1.s) at 298 K with an activation energy of 46.7 kJ/mol. A value of 5868 m3/(kmol.s) a t 298 K with an activation
Ind. Eng. Chem. Res., Vol. 29, No. 8, 1990 1727
. IO0 0 -
01
0 01
0 288 K 278 K
1
I
0.10
100
200
C A M P I , kmollm’
Figure 3. Observed pseudo-first-order rate constant as a function of AMP concentration a t three different temperatures.
energy of 41.2 kJ/mol was obtained by Hikita et al. (1977) using another direct technique, rapid mixing. Other data, which were mostly obtained from gas absorption experiments and reviewed by Blauwhoff et al. (1983), are also close, and the differences can be attributed to the uncertainties in solubility, C*, and diffusivity,D, of the dissolved COP,to determination of the exact contact area, and to possible interfacial turbulence in some types of the absorbers.
C02-AMP Results Figure 3 shows the pseudo-first-order rate constant kobs as function of AMP concentrations at three different temperatures, and the regression analysis yielded the reaction orders in AMP as shown in the second column of Table I. Although the order is near to unity, it is consistently higher than those obtained for MEA. A fractional order between 1 and 2 would be expected if the deprotonation of zwitterion is not instantaneous (Danckwerts, 1979). Therefore, there appears to be some uncertainty about the order in AMP. However, if we consider the order as unity, it is possible to obtain the second-order rate constant k2, which can then be compared with previously published results. Figure 2 shows also the k, values for AMP (which were obtained by linear regression analysis) as a function of temperature. The linear regression analysis gave the following for k2: log k2 = 10.023 - 2177.45/T (7) This resulted in an activation energy of 41.7 kJ/mol. Equation 7 predicts k2 values of 520 and 1165 m3/(kmol.s), respectively, for 298 and 313 K. Although the order of magnitude is the same, our value a t 298 K is somewhat smaller than 1048 m3/(kmol.s), which was the value reported by Sharma (1965). It should, however, be noted that all the values corresponding to different amines reported by Sharma (1965) are consistently higher than those obtained by other workers. This difference may, to some extent, be attributed to the errors in the values of solubility C* and diffusivity D, since he used those corresponding to water rather than concentrated amine solutions. (For instance, his k2 value of 7600 m3/(kmol.s) for MEA a t 298 K exceeds the 5500 m3/(kmol.s) value agreed upon by the majority of others (Blauwhoff, 1983).) Our predicted rate constant a t 313 K is 1165 m3/(kmol-s), which compares well with 1270 m3/(kmol.s) as reported by Yih and Shen (1988). The results of Chakraborty et al. (1986), who deduced for k2 a value of 100 m3/(kmol.s) at 313 K from the gas absorption data, have recently been reanalyzed by Bosch et al. (1989). It was shown by Bosch et al. (1989)
that their results should be interpreted in terms of mass transfer with parallel reversible reactions in order to prevent erroneous conclusions on the reaction mechanism and the rate constant k2. It is now possible to gain some insight into the mechanism of the reactions between C 0 2and AMP. A secondorder rate constant with a magnitude of lo3 m3/(kmol-s) at 298 K is considerably higher than values around 10 m3/ (kmol-s),which are reported for a base-catalyzed reaction (5) by tertiary amines (Versteeg and Van Swaaij, 1988). On the other hand, such a value is very reasonable for a zwitterion intermediate mechanism consisting of reactions 1 and 2. Thus it may be speculated that the reaction proceeds in the accepted way to form carbamate ion through reactions 1and 2, possibly the first step being the rate-controlling step. The carbamate ion is then hydrolyzed into bicarbonate ion through reaction 3 so that the final reaction mixture has no or negligible carbamate ion. The latter reconciles also the finding of Chakraborty et al. (1986), who found no carbamate ion peak in NMR spectra of the equilibrium reaction mixture.
Conclusions COP reacts with the sterically hindered amine 2amino-2-methyl-l-propanol according to a zwitterion intermediate mechanism leading eventually to carbamate ion; the first step is possibly the rate-controlling step. When compared with the rate constant of sterically unhindered monoethanolamine, the second-order rate constant for AMP is smaller than that of MEA by a factor of about 10. The corresponding activation energies for AMP and MEA are 41.7 and 46.7 kJ/mol, respectively. The results of MEA compare well with the literature data, which were obtained mostly from gas absorption studies. Our second-orderrate constant for AMP is in good agreement with that reported by Yih and Shen (1988), but it is somehow smaller than that of Sharma (1965); the differences may, however, be a reflection of uncertainties in estimating solubility, C*, and diffusivity, D, of the dissolved C02 and the contact area of the absorber in the latter two studies. It is interesting to note that if the data of Chakraborty et al. (19%) are analyzed in terms of mass transfer with reversible parallel reactions, the reaction mechanism as proposed here in addition to a simultaneous hydration reaction are adequate to explain the observed C02absorption behavior of aqueous AMP solutions (Bosch et al., 1989). Acknowledgment The support of Kuwait University through an initiation grant (EDC-056) is gratefully acknowledged.
Nomenclature C * = solubility of dissolved C02, kmol/m3 D = diffusivity of dissolved C02,m2/s E = activation energy, kJ/mol k 2 = second-order reaction rate constant, m3/(kmold hobs= observed pseudo-first-orderreaction rate constant, s-l r = homogeneous specific reaction rate between C 0 2 and A M P in aqueous solutions, kmol/ ( m 3 4 T = temperature, K Registry No. MEA, 141-43-5; AMP, 124-68-5;COP,124-38-9.
Literature Cited Alper, E. Kinetics of reactions of carbon dioxide with diglycolamine and morpholine. Chem. Eng. J. (Lausanne) 1990, in press. Blauwhoff, P. M. M.; Versteeg, C. G.; van Swaaij, W. P. M. A study on the reactions between C 0 2 and alkanolamines in aqueous so-
I728
I n d . Eng. Chem. Res. 1990, 29, 1728-1731
lution. Chem. Eng. Sci. 1983, 38, 1411-1429. Bosch, H.; Versteeg, G . F.; van Swaaij, W. P. M. Gas-liquid mass transfer with parallel reversible reactions-I. Absorption of COz into solutions of sterically hindered amines. Chem. Eng. Sci. 1989, 44, 2723-2734. Chakraborty, A. K.; Astarita, G.; Bischoff, K. B. COz absorption in aqueous solutions of hindered amines. Chem. Eng. Sci. 1986,41, 997-1003. Danckwerts, P. V. The reactions of COz with ethanolamines. Chem. Eng. Sci. 1979, 34, 443-445. Hikita, H.; Asai, S.; Ishikawa, H.; Honda, M. The kinetics of reactions of carbon dioxide with MEA, DEA and TEA by a rapid mixing method. Chem. Eng. J. 1977, 13, 7-12. Marquardt, D. W. An algorithm for least squares determination of nonlinear parameters. J.SOC.Ind. Appl. Math. 1963, 11, No. 2. Sartori, G.; Savage, D. W. Sterically hindered amines for COz removal from gases. Ind. Eng. Chem. Fundam. 1983,22,239-249.
Sartori, G . ; Ho, W. S.; Savage, D. W.; Chludzinski, G . R.; Wiechert, S. Sterically-hindered amines for acid-gas absorption. Sep. Purif. Methods 1987, 16(2), 171-200. Sharma, M. M. Kinetics of reactions of carbonyl sulphide and carbon dioxide with amines. Trans. Faraday SOC. 1965, 61, 681-688. Versteeg, G. F.; van Swaaij, W. P. M. On the kinetics between C 0 2 and alkanolamines both in aqueous and non aqueous solutions11. Tertiary amines. Chem. Eng. Sci. 1988, 43, 587-591. Yih, S. M.; Shen, K. P. Kinetics of carbon dioxide reaction with sterically hindered AMP aqueous solutions. Znd. Eng. Chem. Res. 1988, 27, 2237-2241. Zioudas, A. P.; Dadach, Z. Absorption of COz and HzS in sterically hindered amines. Chem. Eng. Sci. 1986, 41, 405-408.
Received for review August 4, 1989 Revised manuscript received March 15, 1990 Accepted March 23, 1990
Natural Convection Mass Transfer in Slender Conical Cavities Gomaa H. Sedahmed*tt and Abdel-Moneum M. Ahmed* Chemical Engineering Department, Faculty of Engineering, and Chemistry Department, Faculty of Science, Alexandria University, Alexandria, Egypt
The rates of free convection mass transfer inside a conical cavity were studied by measuring the limiting current of the cathodic deposition of copper from acidified copper sulfate solution. T h e variables studied were the physical properties of the solution and the slant height of the cavity. The data were correlated in the range 1.46 X 1O'O < ScCr < 8.3 X 10" with eq 3. Comparison with previous mass-transfer studies on hemispherical, rectangular, and cylindrical cavities shows that cavity geometry plays a significant role in determining the nature of flow and the rate of mass transfer inside the cavity.
Introduction Although much work has been done on natural convection heat transfer in enclosures (Ostrach, 19721, little has been done on natural convection mass transfer in these geometries despite their practical importance. Previous studies of natural convection mass transfer in enclosures include geometries such as horizontal (Sedahmed and Shemilt, 1981) and vertical annuli (Sedahmed and Shemilt, 1982), horizontal tube (Sedahmed and Shemilt, 1983), a vertical narrow gap (Bohm et al., 1966), and cylindrical (Somerscales and Kassemi, 1985) and rectangular cavities (Kamotani et al., 1985). The present work deals with natural convection mass transfer in conical cavities. Conical cavities are encountered frequently in chemical engineering practice either alone or as a part of a more complex geometry; for instance, in the design of storage tanks, chemical reactors, and crystallizers,a conical bottom is preferred to other geometries because it allows complete drainage of the solution. Other equipment such as cyclones, centrifuges, and hoppers also use conical cavities. The present study assists in predicting the rate of diffusion-controlled reactions that might take place in conical cavities under free convection or weak forced convection with free convection contributing, e.g., electroplating,
f
Chemical Engineering Department. Chemistry Department.
electropolishing, corrosion, electroless plating, electroforming, and electrochemical machining. In view of the analogy between heat and mass transfer, it is hoped that the present work may be useful to heat transfer since only heat transfer to cylindrical cavities was studied (Japkise, 1973). The present case corresponds to a cavity with a constant-temperature heated wall where the fluid density at the wall is less than that in the bulk. Cavities filled with liquids are used in cooling equipment such as nuclear reactors, transformer cores, turbine blades, and internal combustion engines under the name thermosyphon (Japkise, 1973). The rates of mass transfer were measured by the well-known electrochemical technique that involves measuring the limiting current of the cathodic deposition of copper from acidified copper sulfate solution (Selman and Tobias, 1978).
Experimental Technique The apparatus (Figure 1) consisted of a cell and electrical circuit. The cell consisted of a vertical conical cavity cathode machined in a copper block. The cavity was filled with electrolyte to different heights. Six cavities were machined with apex angles 32,36,42,48,56, and 68'. The slant height ranged from 5 to 9.1 cm. A 2-mm-diameter vertical copper wire was placed in the center of the cavity to act as the anode. The advantage of placing the anode inside the cavity is that the primary current distribution
0888-5885/90/2629-1728$02.50/0 @Z1990 American Chemical Society
