Thermal Effects During the Absorption of CO 2 in Aqueous Solutions of 3-Amino-1-Propanol

Full Papers

Thermal Effects During the Absorption of CO ₂ in Aqueous Solutions of 3-Amino-1-Propanol

By F. Camacho, S. Sµnchez, and R. Pacheco*

In this work, we study the process of CO2absorption, at high partial pressures, in aqueous solutions of 3-amino-1-propanol (AP), with respect to the thermal effects of this operation. All of the experiments were performed in a stirred tank gas-liquid reactor with a flat, known interface. The variables considered were the AP concentration in the range of 0.1 to 3.0 M and the temperature within the interval of 288±313 K. From the results, we deduce that the process takes place in the instantaneous nonisothermal regime, and we propose an equation which relates the experimental results of molar flux with the initial amine concentration. At the same time, we can evaluate the temperature increase at the gas-liquid interface.

1 Introduction

In CO2absorption, it is necessary to differentiate between the use of dissolved alkanolamines in the aqueous medium and organic solvents. Currently, in industry, mostly aqueous alkanolamine solutions are used, some of the most common being monoethanolamine (MEA), diethanolamine (DEA), di-2-propanolamine (DIPA) and methyl-diethanolamine (MDEA). There is extensive literature on these amines with regard to their absorption kinetics in an isothermal regime, and fundamentally at low partial pressures of CO₂.

It should be mentioned that in the physical absorption of a highly soluble gas or in absorption with a chemical reaction, temperature of the liquid phase, especially near the gas-liquid interface, can increase due to the heat released by the solution and/or by the reaction [1±3]. In some systems, these effects are slight and the rises in temperature at the interface are hardly appreciable. Nevertheless, in certain gas-liquid systems of industrial interest, thermal effects have been noted. Among the most well-known systems are ammonia-water, sulfur trioxide-dodecylbenzene and hydrogen chloride-ethylene glycol. In addition, thermal effects have been reported in reactions of sulfonation, chlorination and oxidation of hydrocarbons [2±4].

In relation to CO2absorption, at reduced partial pressures and in aqueous solutions of alkanolamines, these effects have been considered very slight and negligible by practically all the research groups which have worked with these systems.

However, it should be pointed out that when the absorption takes place at high partial pressures of CO2the thermal effects appear to gain certain importance. In fact, there are some antecedents, such as CO2absorption in aqueous solutions of MEAatpartialpressuresnearatmosphericpressureinthework of Clarke [5]. This author used a laminar jet apparatus, MEA

concentrations within the range of 1.6 and 4.8 M, and worked at low pressures (around 80 mm Hg) and high CO₂(about 750 to 760 mm Hg) and indicated qualitatively that in this latter case, the heat of the reaction influenced the absorption rate.

In the present work, we analyse the process of CO₂ absorption, at high partial pressures, in aqueous solutions of 3-amine-1-propanol (AP), in relation to the thermal effects of this operation.

2 Materials and Methods

2.1 Experimental Device

All the experiments were made using a batch gas-liquid reactor (stirred-tank type) with respect to the liquid phase, with a flat and known interface area.

2.2 Procedure

The use of pure CO2 enables the determination of the absorption rate at different points, using a soap-film meter which allows the direct measurement of the CO2absorbed.

The operation variables considered were AP concentra- tion in the range 0.1 to 3.0 M and temperature within the interval 288 and 313 K. In all the experiments the stirring rate in the reactor was 180 rpm, maintaining a flat interface area of 35.26 cm².

2.3 Physical and Transport Properties

Under these experimental conditions, we measured the viscosity of the amine solutions and their density. The viscosity of the solutions was measured by using a capillary viscometer.

The calculation of the initial partial pressure of CO₂ is given by¹⁾:

±

[*] F. Camacho, Department of Chemical Engineering, Faculty of Science, University of Granada, 18071 Granada; S. Sµnchez (author to whom correspondence should be addressed), R. Pacheco, Department of Chemical Engineering, Faculty of Experimental Science, University of JaØn, 23071 JaØn, Spain.

±

1) List of symbols at the end of the paper.

p_A ˆ P ÿ p_v (1) P being the barometric pressure and p_vthe vapor pressure of the solution.

To determine the solubility of CO₂in the liquid phase, we used the relationship of Danckwerts and Sharma [6].

He ˆ 10…5:3 ÿ 1140=T† (2)

The diffusion coefficient of CO₂in the aqueous solution was measured for the relationships of Sada et al. [7] and Versteeg and van Swaaij [8].

D_A ˆ D_N2O;am D_N2O;w

!

D_CO2;w (3)

where

D_CO2;w ˆ 2:35
10^ÿ6 e^ÿ2119=T† (4)

D_N2O;w ˆ 5:07
10^ÿ6 e^ÿ2371=T† (5)

D_N2O;am ˆ D_N2O;w _w



 

(6) The diffusion coefficient of AP in aqueous solution was determined by means of the relationship of Wilke and Chang [9].

D_B ˆ 3:06
10^ÿ15 T

 (7)

2.4 Analytical Methods

The initial concentration of amine was determined by titration with HCl solutions using methyl orange as the indicator. The CO2 concentration in the liquid phase was determined by standard titration methods. Excess NaOH and the excess BaCl2solutions were added to the liquid sample and the excess NaOH was titrated with HCl solution using thymol blue as the indicator.

3 Results and Discussion

The CO₂flux (N_A) were calculated assuming that the gas behaves ideally and using the expression:

N_A ˆ nA^; ˆ P QR T A^; (8)

The value of volumetric flow (Q') coincided with the value of the slope of the straight lines resulting from the representa- tion of the CO₂volume absorbed against time (t) for each experiment, this being determined by least-squares fit.

All of the experiments were performed at air temperature practically coinciding with the operation. At the beginning of each experiment, the greatest difference between the exterior

temperature and that in the inside of the reactor (in the bulk liquid phase) was 1 C.

Initially, to determine the individual coefficient of mass transfer in the liquid phase (kL), we conducted experiments on the physical absorption of CO₂ in water, using the same reactor subsequently employed in the chemical-absorption experiments. The results obtained during the processes of physical absorption have enabled the determination of the k_L values for the range of temperatures investigated (Camacho et al. [10]).

In the absorption experiments performed with pure CO2in aqueous solutions of AP, the influence of the initial concentration of alkanolamine and the operating temperature were analyzed.

3.1 Absorption in AP Solutions

A primary alkanolamine, AP (CH2OH-CH2-CH2-NH2), in the reactions of its functional groups (-NH₂and -OH) with CO2, behaves as described by different authors for mono- ethanolamine (Astarita et al. [11]; Hikita et al. [12]).

-NH2+ CO2® -NHCOO^±H⁺ (9)

derived from carbamic acid

-OH + CO2® -OCOO^±H⁺ (10)

derived from carbonic acid

This latter reaction can take place in a basic solution (pH ³ 11.0); when the pH is lower than 10.0, even in a carbonate solution, the formation of the derivative of carbonic acid can be considered negligible [11].

The pH measurements made indicate that for all the concentrations and temperatures investigated, the pH value of the solution was less than 11.0, except in the experiments made at the highest AP concentration and during the first few moments, as shown in Fig. 1.

Therefore, it might be expected that the reaction was fundamentally the formation of a carbamic acid derivative [reaction (9)].

In addition, Fig. 2 for example provides the graphic representations of CO2absorbed per unit of surface and time for the experiments conducted at 303 K. Operating temper- ature and partial pressure of the gas were maintained constant, varying the initial concentration of AP between the values 0.155 and 2.704 M.

In all cases, the volumetric flow of CO₂absorbed increased significantly when the initial concentration was increased, as clearly shown in the steeper slopes of the straight lines. Also, a slight curvature can be detected at times exceeding those considered in the calculation of NA, fundamentally in the experiments performed at low C_Bovalues. Tab. 1 lists the N_A values corresponding for each of the initial AP concentrations investigated.

Given that for all the concentrations and temperatures investigatedwithaqueoussolutionsofAP,thecarbonationratio g(molesofCO₂permoleofamineinsolution)waslessthan0.5,

in agreement with Astarita et al. [11], the main product of the reactionshouldbecarbamate.Inthissense,themechanismmay be that indicated by the reactions (11) and (12), comparable to that described for MIPA by Hikita et al. [13]:

CO₂+ RNH₂® RNHCOO^±+ H⁺ fast (11) RNH₂+ H⁺® R-NH₃⁺ virtually instantaneous (12) Though there is less information available concerning this alkanolamine in comparison to the other primary ones, the reaction mechanisms proposed, and the results of g for AP (g<0.5), nevertheless indicate that this is a reaction of an overall order of two ± one with respect to CO2and one with respect to AP (Penny and Ritter [14]) ± and the three possible regimes that could take place, according to Astarita et al. [11], should be considered:

a) If CBo/2C^*A<< 1 the absorption rate being given by:

N_A ˆ kL C_A^ (13)

which corresponds to the hydrodynamic regime or that of physical absorption

b) If 1 < < C^Bo

2 C^_A < < 

k CBo 

p (14)

N_A ˆ kL C_Bo

2 (15)

Figure 1. Variation in pH of the AP solution with time, over the course of the absorption process at 293 K at the corresponding concentrations; (s) 0.157 M;

(*) 2.534 M.

Figure 2. CO2absorbed per unit of surface and time in the experiments with AP at 303 K; (*) 0.155 M; (D) 0.414 M; (*) 0.850 M; (&) 1.812 M; (s) 2.704 M.

Table 1. Molar fluxes in the absorption of CO2in 3-amine-1-propanol solutions.

which corresponds to the instantaneous-reaction regime c) If



k CBo 

p < < C^Bo

2 C^_A (16)

NA ˆ C^_A 

DA k CBo

p (17)

which corresponds to the fast-reaction regime

When g> 0.5, in addition to reactions (11) and (12), the following reaction would occur:

RNHCOO^±+ 2H₂O + CO₂® R-NH₃⁺+ 2HCO₃^± (18) the kinetics of which fit a pseudofirst-order reaction.

3.2 Reaction Regime

To determine the reaction regime during the absorption process with this alkanolamine, we have used the criteria proposed by Astarita et al. [11] when g<0.5.

According to the values of the initial amine concentrations used, the quotient CBo/2C^*A is much greater than one and consequently Eq. (13) cannot be applied to correlate the experimental values for CO2 absorption by aqueous AP solution, as, on the other hand, it indicates the appreciable influence of C_Boon N_A.

If we assume as an initial hypothesis the fast-reaction regime of the order ªmº with respect to the CO₂and ªnº with respect to AP, the molar flux is given by the expression (Charpentier [15]):

NA ˆ



2

m‡ 1 D^A km ; n …C^_A†^m‡1 Cⁿ_Bo r

(19) It has been amply demonstrated that the order of the reaction with respect to CO₂in the absorption with aqueous solutions of primary amines is 1. If we admit this fact for the case of AP, then Eq. (19) is reduced to

N_A ˆ C^_A 

D_A k_1;n Cⁿ_Bo

p (20)

This equation coincides with that of Astarita et al. [11] when n = 1, (Eq. (17)) and it can be linearized in the form:

log N²_A He² p²

A D_A

" #

ˆ log k1;n ‡ n log CBo (21) For this equation, we considered that C^*A is the CO2

concentration in equilibrium with the gaseous phase and this can be evaluated by Henry's law (p_A= He C^*_A).

From the experimental results, we represented the first term of Eq. (21) against the log C_Bo. The results did not adjust well to Eq. (21), and the slope ªnº (order of reaction with respect to AP) was greater than 1.5 and close to 2. This finding completely disagrees with the literature, which states that

the order of the reaction is 1 with respect to AP. In addition, from the ordinate at the origin (log k1,n) we should have constant kinetics, close to those found in the literature, but in fact we found k values of two orders of magnitude lower. For all these reasons, these results induce us to reject the idea that the process of absorption was operated in a fast-reaction regime.

According to Astarita et al. [11], under the inequality in (14), the reaction regime is instantaneous, satisfying Eq. (15). The value of h (average residence time of the elements at the interface), in agreement with the penetration theory, could be evaluated by the equation,

 ˆ D^A

k²_L (22)

On the other hand, if we consider the values of the kinetic constant by Penny and Ritter [14], we can, by linear regression, arrive at the equation:

ln k₂= 26.6 ± 5226:3

T (23)

This expression is valid within the temperature interval 293 to 303 K. However, given that outside this range no information is available from the literature, we made an approximate extrapolation to 288 K and a higher temperature of 303 K. On the basis of these considerations, and if in principle it caused no interface temperature rise, the calcula- tion of the third term of the inequality (14) is immediate, and can be put into the form:

1 < < C^Bo He

2 p_A < < 1 k_L



k CBo D_A

p (24)

From the experimental results and the k values deduced from Eq. (23), we determined the second and third term of the inequality, values given in Tab. 1. In general, the inequality proved to hold for practically the entire range of concentra- tions and temperatures investigated, and therefore, in principle, the absorption should take place in the instanta- neous reaction regime.

3.3 Influence of the AP Concentration and Partial Pressure of CO₂

In general, the molar flux in a process with chemical reaction can be expressed in the form:

NA ˆ E kL C^_A (25)

If the reaction regime is instantaneous (E = Ei), and given the definition of E_i (instantaneous-enhancement factor) according to the film theory, it would result that:

N_A ˆ 1 ‡ D^B z D_A C_Bo

C^_A

" #

k_L C^_A (26)

In relation to Eq. (26), it should be indicated that according to the stoichiometry of the reaction which takes place, z = 2. In addition, this equation is reduced to Eq. (15) in the cases in which it can be considered that DA» DBand CBo/2C^*A>>1.

To determine whether Eq. (15) is applicable, we prepared the graphic representations of NAvs CBo, as shown in Fig. 3.

We find that the fit is not appropriate for the entire range of concentrations of alkanolamine studied, clearly detecting that the highest concentrations show no linear dependence between the molar flux values (N_A) and the amine concentra- tions (CBo).

Figure 3. Variation in NAwith the AP concentration at the temperatures indicated; (&) 288 K; (s) 293 K; (*) 298 K; (H) 303 K; (n) 308 K; (D) 313 K.

An analysis of Fig. 3 indicates that at high alkanolamine concentrations, there could be a certain approximation to the transition regime (so called for being intermediate between the fast and instantaneous reaction regimes). In addition, the fit of Eq. (15) within the concentration intervals indicated in the figure is not adequate, as there is an ordinate at the origin which cannot be disregarded and the value of the slope is different from the corresponding experimental value kL/2.

This fact may indicate that there is an appreciable temperature rise at the interface.

3.4 Temperature Increase

The repetition of certain experiments at the highest concentrations, in which temperature was monitored in the center of the AP solution over the process has demonstrated that in fact an appreciable rise in the temperature should occur

at the interface, since, in the entire volume of the reactor, the temperature increased by 1 to 2 C. Considering that the reaction takes place in a thin film of the gas-liquid interface, we might deduce that the temperature increase should be substantial.

Therefore, the most apt hypothesis appears to be that the absorption process occurs during the instantaneous non- isothermal regime. In this sense, the temperature in the bulk liquid would be TB, while that in the interfacial film would be T_s, (T_s ± T_B) the increase of surface temperature being generated by the dissolution and reaction process.

After assaying the application to our results of the different models proposed for the nonisothermal absorption processes, the greatest consistency was found with a modification of Eq. (26) (Eq. (27)), which enables us to relate the experi- mental results NA and CBo, and in turn evaluate the temperature increase at the interface:

N_A ˆ kL C_As^ ‡ D^B k_L

2 D_As C_Bo (27)

In this equation, it is accepted that AP diffuses (DB) from the bulk liquid phase towards the interface at temperature T_B, and that the individual coefficient of mass transfer (k_L) also to be determined at temperature TB, a fact that contradicts the model of Mann and Moyes [2], given that the thickness of the thermal film near the interface is appreciably greater than the thickness of the mass-transfer film. However, this equation fits better the results of the present research.

By the application of Henry's law, we can determine the concentration of CO₂dissolved at the temperature T_s, C^*_As, whereupon Eq. (27) can be transformed to:

N_A ˆ kL

p_As

He_s ‡ D^B k_L

2 D_As C_Bo (28)

The values of p_As, He_sand D_Asat the temperature T_scan be determined by the Eqs. (1), (2) and (3), respectively, so that, by means of these three equations and expression (28), the value of Tscan be calculated.

Following the calculation sequence in Fig. 4a, we deter- mined the temperature rises within the concentration inter- vals where the reaction regime is fully instantaneous, coinciding with the good fit provided by Eq. (28), represented with a solid line in Fig. 3. Tab. 2 shows all of the experimental results used in the calculation sequence, which has enabled us to determine the corresponding temperature increases. In addition, these increases (DT = Ts±TB) determined in each experimental series are represented against the AP concen- tration in Fig. 5.

Thus, DT increases with the AP concentration, and in the series performed, T_B= 313 K, at the highest concentrations, there was decline in DT, possibly due to an elimination of heat caused by transport by convection or even evaporation, towards the gaseous phase and to the heat transfer from the interface to the bulk liquid phase.

Finally, to formulate an empirical expression which relates the values of NA and CBo throughout the concentration

interval studied, we have tried to find an equation as close as possible to Eqs. (21) and (28). Among the different equations

assayed, the two that best fit the molar flux values and the amine concentration, for any temperature series, took the form:

NA ˆ 1 ‡ 2 C_Bo¹⁼² ‡ 3 CBo (29)

log N_A² D_A C^_A ²

!

ˆ ₁ ‡ ₂ log CBo (30)

In Eq. (29), the parameters k1, k2 and k3 have been determined by nonlinear regression from the experimental values NAand CBo, while in Eq. (30) the parameters û1and û2

have been calculated by the method of least squares from the values NA, DA, C^*A(pAand He) and CBo. The values of these five parameters are listed in Tab. 3, as well as the goodness of fit (r²) provided by the two equations.

Undoubtedly, of the two equations the more useful one is (30), since, requiring fewer parameters and relating a larger number of variables, it in general gives a better fit in all cases.

In addition, in this equation, the values of the parameters (û₁ and û₂) show a definite tendency, both increasing with temperature. In the fast-reaction regime û2® 1 and û1 ® log k; the value of û₂decreases with T tending towards one, a result in agreement with the form of Eq. (30), which would correspond to a fast-reaction regime. However, the û₁value differs appreciably from the log k, given that, as mentioned above, the absorption takes place in the instantaneous- reaction regime or in any case in the transition between the two regimes.

Figure 4. a) Calculation sequence for the determination of Ts. b) Calculation sequence for the determination of E and Ha.

Table 2. Interfacial temperature, enhancement factors and Hatta number in the experiments on AP.

Figure 5. Variation of the temperature increase at the interface with the concentration in the AP experiments at the TBvalues considered; (&) 288 K;

(s) 293 K; (*) 298 K; (H) 303 K; (n) 308 K; (D) 313 K.

3.5 Enhancement Factors Once the value of T_s is calcu- lated, in the experiments in which the reaction regime was fully instantaneous, the determination of the Hatta number (Ha) is immediate by the equation:

Ha ˆ



k_2s C_Bo D_As k²_L s

(31) Itshouldbeindicatedthatforthe calculation of the kinetic constant (k2s), Eq. (23) was used with the considerations mentioned above.

On the other hand, accepting the hypothesis described in the previous section, we can evaluate the instantaneous- enhancement factor at temperature T_s, E¢_is,

E^;_is ˆ 1 ‡ D_B 2 DAs

C_Bo He_s

p_As (32)

With the E¢_isand Ha values, and using the expression of DeCoursey [16], for maximum calculating ease, we can finally determine the enhancement factor (E):

E ˆ



Ha⁴

4 …E⁰_isÿ1 †² ‡ E⁰_is Ha² E⁰_isÿ 1 ‡ 1 vu

ut ÿ Ha²

2 …E⁰_isÿ1† (33) The calculation sequence to determine the Hatta number and the enhancement factor as well as instantaneous- enhancement factors is summarized in the diagram in Fig.

4b, and the values are listed in Tab. 2, where the instantaneous regime is confirmed by:

E ^ E¢is and

Ha > 15 E¢_is

In addition, Tab. 2 contains a column corresponding to the expression CBoHes/2 pAs, which, if compared with the Hatta number (Eq. (31)), confirms that we are really applying the criterion of Astarita et al. [11] to the temperature T_s, which is the one found in the reaction. These two columns can also be compared with the last two in Tab. 1, where initially the temperature increase at the interface was not considered.

Thus, the assumed instantaneous regime deduced with a certain approximation to the temperature T_B, remains valid, since in all cases it holds that:

1 k_L



k_2s C_Bo D_As

p > > C^Bo He_s

2 p_As (34)

In addition, in Fig. 6, on representing E vs Ha on logarithm coordinates, we again confirm the instantaneous regime, as we find that the experimental results are situated in the first

quadrant, acceptably fit the same curve and are far from the straight line that would appear if the experiment had been operated in the fast-reaction regime.

Received: January 7, 2000 [CET 1190]

Symbols used

A [m²] interfacial area

C^*_A [kmol m^±3] CO₂concentration in equilibrium with gaseous phase

C^*As [kmol m^±3] CO2concentration in equilibrium with gaseous phase at temperature T_s CBo [kmol m^±3] initial concentration of

amine in aqueous phase DA [m²s^±1] diffusion coefficient of

component A (CO₂) in aqueous alkanolamine solution

Table 3. Parameters for equations (29) and (30).

Figure 6. Variation of the enhancement factor (E) with the Hatta number (Ha) in the AP experiments at the temperatures indicated; (&) 288 K; (s) 293 K; (*) 298 K; (H) 303 K; (n) 308 K; (D) 313 K.

D_As [m²s^±1] diffusion coefficient of CO₂ in liquid phase at

temperature T_s DB [m²s^±1] diffusion coefficient of

alkanolamine in liquid phase

DCO2,w [m²s^±1] diffusion coefficient of carbon dioxide in water DN2O,am [m²s^±1] diffusion coefficient of

nitrous oxide in amine solution

D_N2O,w [m²s^±1] diffusion coefficient of nitrous oxide in water

E [±] enhancement factor

E_i [±] instantaneous-enhance-

ment factor

E¢is [±] modified instantaneous-

enhancement factor at temperature Ts

Ha [±] Hatta number

He [kPa m³kmol^±1] Henry's law constant He_s [kPa m³kmol^±1] Henry's law constant at

temperature T_s

k reaction rate constant

k_L [m s^±1] liquid-phase mass-transfer coefficient

k₂ [m³kmol^±1s^±1] second-order reaction-rate constant

k2s [m³kmol^±1s^±1] second-order reaction-rate constant at temperature T_s

n [±] order of reaction with

respect to amine

n' [kmol s^±1] rate of absorption of CO2

NA [kmol m^±2s^±1] rate of absorption per unit interfacial area

pA [kPa] partial pressure of CO2

p_As [kPa] partial pressure of CO₂at the temperature Ts

P [kPa] total pressure

p_v [kPa] vapor pressure of water

Q' [m³s^±1] volumetric flow rate of absorbed CO₂ R [kPa m³kmol^±1K^±1] gas constant

T [K] temperature

TB [K] temperature in bulk liquid

phase

Ts [K] temperature in interfacial

film

z [±] stoichiometric coefficient

Greek symbols

b1, b2 [±] constants defined in Eq. (30)

c [±] constant defined in Eq. (6)

g [(mol of CO₂)

(mol of amine)^±1] carbonation ratio

h [s] average life of surface

elements

k₁, k₂, k₃ [±] constants defined in Eq. (29)

l [Pa s] viscosity of solution amine

lw [Pa s] viscosity of pure water

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