Mechanisms and kinetics of diethanolamine degradation - Industrial

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Ind. Eng. Chem. Fundam. lQ85, 2 4 , 129-140

Mechanisms and Kinetics of Diethanolamine Degradation Malcolm L. Kennard and Axel Allelsen' Department of Chemlcal Englneerlng, The Unlverslty of Brltlsh Columbia, Vancouver, Brltlsh Columbia, Canada V6T 1 W5

Laboratory experiments with a 600-mL batch reactor were performed to investigate the degradation of aqueous dlethanolamlne (DEA) solutions under the following conditions: concentration, 0 to 100 wt % DEA; temperature, 90 to 250 O C ; total pressure, 1.5 to 6.9 MPa. Most experiments were conducted in the presence of C02, but thermal degradation using nitrogen was also examlned. The principal degradation compounds were found to be 34hydroxyethyl)-2-oxazolidone (HEOD), N,N,N-tris(hydroxyethy1)ethylenediamine(THEED), and N,N-bis(hydroxyethy1)piperazine (BHEP). I t was discovered that DEA degradation occurs by several complex routes. However, at temperatures below 175 O C and COPloadings exceedlng 0.2 g of CO,/g of DEA, the degradation is well represented by three firstorder reactions: DEA HEOD and DEA THEED BHEP. The rate constants for these reactions are given as a function of temperature. The implications of the present fundamental studiis for gas plant operations are discussed. I t is also shown that activated carbon, which is widely used industrially, is ineffective in removing the principal DEA degradation compounds.

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Introduction Diethanolamine (DEA) is extensively used in the gas processing industry for removing acid gases such as carbon dioxide and hydrogen sulfide from light hydrocarbons. DEA's popularity is based on several factors: energy savings compared with certain other solvents; high affinity for acid gases; fair resistance to degradation. Degradation is defined as the irreversible transformation of DEA into undesirable compounds. Although DEA is claimed to have good degradation characteristics, many gas planta experience problems with their DEA solutions. Degradation is undesirable because valuable DEA is lost, degradation products accumulate leading to equipment fouling, the process efficiency and throughput may be impaired, and some degradation producta may aid corrosion. The direct and indirect costa resulting from DEA degradation are considerable. Consequently, a clear understanding of degradation and methods for suppressing it are needed. Very little systematic research on DEA degradation has been reported. Plant operators have usually tried to solve the problem by trial and error, which includes changing operating conditions and/or installing activated carbon filters (Scheirman, 1973). In some cases these measures are successful, but more frequently, they are inadequate. Furthermore, satisfactory procedures for one gas plant are often ineffective for other plants. Most experimental studies on the DEA process have been limited to the absorption and desorption of acid gases, which take place in a matter of seconds. By contrast, degradation is slow and it may take several months before the degradation compounds reach significant concentrations. Degradation is a complex phenomenon and Younger (1973), Smith and Younger (1972a,b),and Nonhebel(1972) have reported that it is affected by temperature, pressure, raw gas composition, amine concentration, solution pH, and possibly the presence of metal ions. The quantitative relationship between these variables and degradation is unknown, however. It is therefore impossible to predict DEA degradation or, alternatively, to estimate improvements from changes in operating variables. The situation is further complicated by the fact that the degradation products are large organic molecules which are difficult to detect and identify. Polderman and Steele (1956) as well as Hakka et al. (1968)have reported some laboratory experiments on DEA degradation. Their work was qualitative and restricted to 0196-4313/85/ 1024-0129$01.50/0

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C02 as the acid gas. They found that DEA losses over an 8 h period ranged from 0% at 100 OC and 1257 kPa to 97% at,175 "C and 4137 kPa. NJV-Bis(hydroxyethy1)piperaine (or BHEP) was found in the degraded solutions and they suggested that it was formed by the following reactions DEA + C02 HEOD H2O (1)

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2HEOD

4

BHEP

+

+ 2C02

(2)

HEOD or 3-(hydroxyethyl)-2-oxazolidonewas regarded as an intermediate and somewhat unstable compound. If the reaction scheme represented by eq 1 and 2 is correct, the DEA degradation should be governed by a first-order kinetic equation provided C02 is present in excess. Using more sophisticated analytical procedures, Hakka et al. (1968) also detected N,N,N-tris(hydroxyethy1)ethylenediamine (or THEED) in their solutions. However, they did not propose a reaction mechanism for ita formation. In a recent study, Blanc et al. (1982) reacted DEA with C02and HEOD. Both sets of experiments were conducted in a sealed autoclave at temperatures of 90-130 "C. Blanc et al. (1982) proposed various reaction mechanisms for the production of HEOD, BHEP, and THEED and other degradation compounds. However, they provided no quantitative data in support of these mechanisms. The present study therefore had three basic objectives: (a) to elucidate a reaction mechanisms regarding the production of various degradation compounds; (b) to develop a kinetic model which can predict the degradation of DEA and the production of degradation compounds under typical industrial conditions; and (c) to propose ways of reducing DEA degradation in industrial sweetening units and to investigate ways of purifying degraded DEA solutions. Since DEA degradation results primarily from reaction with COz the present study was restricted to this compound. Experimental Section Analytical Procedure. A quantitative study of DEA degradation requires a reliable analytical technique for measuring the concentration of DEA and its degradation compounds in aqueous solutions. The analysis of amines, and especially alkanolamines, is difficult because they tend .to have fairly low vapor pressures, decompose at elevated temperatures, and are highly polar. Many methods have been investigated; however, they all suffer from various drawbacks such as lack of accuracy, specificity, reliability, speed, and cost-effectiveness. 0 1985 American Chemical Society

130 Ind. Eng. Chem. Fundam., Vol. 24, No. 2, 1985 1

Table I. Compounds Detected in Degraded DEA Solutions comDound structural formula monoethanolamine (MEA)

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N-(hydroxyethy1)ethyleneimine (HEM) N-(hydroxyethy1)ethylenediamine (HEED) diethanolamine (DEA) N-(hydroxyethy1)piperazine (HEPI

: u

0

oxazalidone (OZD)

I1

g

HN

I CH,

1 -CHZ

I

triethanolamine (TEA) NJ-bis(hydroxyethy1)ethylenediamine (BHEED) NJhis(hydroxyethy1)piperazine (BHEP) 0

3-(hydroxyethyl)-2-oxazalidone (HEOD)

CH~-

N-(hydroxyethy1)imidazalidone (HE11

cn2

0

I/

HOC~H.-N,~\NH

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1

-CHI (HOCzH,)2NC2H4NHC2H4OH CHI

N,N,N-tris(hydroxyethy1)ethylenediamine (THEED) N,N-bis(hydroxyethy1)imidazalidone (BHEI)

01-

N,N,N,N-tetra(hydroxyethy1)ethylenediamine (TEHEED)

A reliable and simple gas chromatographic technique was developed for detecting DEA and its degradation compounds in aqueous solutions. Details of the method are given by Kennard (1983) and Kennard and Meisen (1984). Various amines and their possible degradation compounds are summarized in Table I. Certain degradation compounds were identified by Kennard (1983) by mass spectroscopy. Equipment and Basic Procedure for Degradation Experiments. Since the degradation of DEA is complex, the experiments had to be performed under carefully controlled conditions. In particular, it was essential to keep the temperature and pressure constant for the full duration of a run. The main component of the experimental equipment was a 600-mL stainless steel, stirred autoclave (Model 4560, Parr Instrument Co., Moline, IL). The autoclave could be operated at temperatures and pressures up to 400 "C and 13.8 MPa, respectively. The autoclave was first purged with Nzto remove any trace of oxygen, and then it was filled with 250 mL of an aqueous DEA solution of known concentration. The acid gas was then admitted and the temperature and pressure were raised and maintained at the desired level. The start-up procedure took about 15 min, which is small compared with the duration of an average experiment which lasted typically more than 8 h. Liquid samples ( 5 mL) were withdrawn at predetermined intervals and analyzed chromatographically. The gas phase was not routinely sampled since no gaseous degradation products could be detected. Preliminary Experiments. Many of the experiments were performed at elevated temperatures (140-205 "C) to speed up degradation. I t is recognized that these temperatures exceed the recommended values for industrial

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Ind. Eng. Chem. Fundam., Vol. 24, No. 2, 1985

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Figure 4. THEED concentration aa a function of time and temperature (30 wt % DEA, 4137 kPa C02, 162-250 " C ) .

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Figure 2. Arrhenius plot for a 30 wt % DEA solution degraded with C02 at 4137 kPa.

TIME (hr)

Figure 5. BHEP concentration as a function of time and temperature (30 wt % DEA, 4137 kPa COz, 162-250 "C).

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Figure 3. HEOD concentration as a function of time and temperature (30 wt % DEA, 4137 kPa C02, 162-250 " C ) .

Figure 6. DEA concentration aa a function of time and initial DEA concentration (4137 kPa C02, 175 " C ) .

zero. As can be seen from Figure 2, the plot is linear for temperatures between 90 and 170 "C. However, at higher temperatures there is a departure from the straight-line behavior. This indicates once again that complex reactions are occurring. It is likely that the first-order behavior is only apparent; i.e., several consecutive reactions are taking place which are affected differently by temperature changes. What can be concluded is that the reactions are highly temperature sensitive. The initial value of KDEA increases by a factor of 3000 as the temperature is raised from 90 to 205 "C. Figures 3 to 5 show plots of HEOD, THEED, and BHEP concentrations as a function of time and temperature. The production rate of HEOD is high at first and then it levels off. The initial rate of production increases with temperature, but the total amount of HEOD produced decreases. Hence, HEOD does not appear to act as an intermediate of the type suggested by Polderman and Steele (1956). Figure 4 shows that the THEED concentration increases at a slightly lower rate than the HEOD concen-

tration and passes through a maximum. The time required to reach the maximum concentration decreases with increasing temperature. THEED appears to behave more like an intermediate than HEOD. Figure 5 shows that the concentration of BHEP steadily increases with time. The overall production of BHEP also rises rapidly with temperature. Above 185 "C,the rate of BHEP formation starts to fall slightly after several hours; this implies that the concentration of some intermediate is falling. HEOD cannot be this intermediate since its concentration remains relatively constant. THEED is more likely to be the intermediate responsible for BHEP formation since its concentration falls after reaching a maximum; this would cause the rate of BHEP formation to decline as illustrated by Figure 5. BHEP therefore appears to be formed from DEA via the formation of THEED; this mechanism differs from that proposed by Polderman and Steele (1956). Effect of Initial DEA Concentration. Figure 6 shows the change in DEA concentration with time for varying initial solution strengths at 175 O C . A t this temperature,

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Figure 9. DEA concentration as a function of time and COz pressure (30 wt 70 DEA, 195 "C).

ooolLx 10

40

60

80.

IO0

INITIAL DEA CONCENTRATION (wt 701

Figure 7. kDEA as a function of initial DEA concentration and temperature (4137 kPa eoz).

0.001 0.0308 ~

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23

1 4

15

16

1000/T ("k"]

Figure 8. Arrhenius plot for varying DEA solution strengths degraded with COz at 4137 kPa.

the data points fall on straight lines indicating a first-order reaction. However, if true first-order behavior exists, then the slope of the lines and the rate constant ( ~ D E A )would be independent of initial DEA concentration. This is clearly not the case as seen from Figure 7. Three regions appear to exist: (1)0-10 wt% DEA: k D m is constant and small; (2) 10-30 wt% DEA: kDEA rapidly increases with increasing initial DEA concentration; (3) 30-100 w t % DEA kDm is high and relatively constant, declining only slightly as the 100 w t % DEA concentration is approached. Figure 8 is the Arrhenius plot with various initial DEA concentrations as a parameter. Again the three regions can be clearly seen. It is noteworthy that the rate constant is lower at 100 w t 70than at 80 wt 70. This suggests that water plays a significant role in the overall degradation. Effect of Pressure and C 0 2 Solubility. Changes in DEA concentration as a function of pressure and time are shown in Figure 9 for a temperature of 195 "C. I t is seen that degradation increases with increasing pressure up to about 4137 kPa; above that pressure little change is noted.

Thus for total pressures of less than 4137 kPa (or about 0.2 g oi C02/g of DEA, Kennard and Meisen (1984)),the C02concentration becomes limiting. In some experiments (high temperature and high DEA concentration runs) the concentration of C02 in DEA fell below 0.2 g of C02/g of DEA. Under such conditions, COz is not present in excess and should be considered in the degradation mechanism. A possible reason for the deviation from straight-line behavior exhibited by certain plots (e.g., Figure 1)could therefore be changes in the C02 concentration in the reaction mixture. Since no C 0 2 is consumed during the reactions (Le., there was no change in total pressure during a run), it is possible for C02 to be converted to a form which reacts more slowly with DEA (see "ionic runs" below) or that C02 is tied up in some manner with the degradation compounds. Referring to Figure 9 it can be seen that as the pressure is increased beyond 4137 kPa, the deviation from linearity ceases although kDEAremains the same. For example, in the case of the run conducted at 6895 kPa, the C02concentration always remained above the limiting value of 0.2 g of C02/g of DEA. Effect of Solution pH. DEA treating solutions containing dissolved C02consist of various ionic and molecular compounds such as R2NH, R2NH2+,R2NCOO-, C032-, HCO;, and C02 (Dankwerts, 1979). It is likely that changes in pH will affect the various equilibria and, in turn, overall degradation behavior. DEA itself makes the solution alkaline, whereas dissolved C02 renders it acidic. Experiments were therefore carried out by adding NaOH or HC1 to the feed in order to change the solution pH. It was observed that lowering the initial solution pH reduces the degradation rate. By reducing the pH from 12.24 to 9 (a 30 w t 90DIM solution at room temperature has a pH of about 11.2), the value of kDEA decreases by more than a factor of 5. The effect of pH may be linked to the solubility of C02 in the DEA solution. The solubility of C02 increases due to the presence of hydroxyl ions. Further studies revealed that when NaOH is added to HEOD in solution, most of the HEOD is converted to DEA. This indicates that the HEOD ring is unstable and is easily attacked by OH-. Effect of Bicarbonate and Carbonate Ions ("Ionic Runs"). Since pH strongly affects degradation, it is likely that degradation involves some ionic compounds. Tests were therefore conducted in which DEA was reacted with potassium carbonate (K2C03)and potassium bicarbonate (KHCOJ under a blanket of nitrogen at 4137 kPa. Under these highly alkaline conditions, it is virtually impossible for HC03- or C032-to revert to free C02. For these runs, the molar concentrations of HC03- and C032-were made equivalent to that of C02dissolved in DEA solutions when gaseous COz was used. Using K2C03resulted in no degradation and C032- may therefore be considered to play a negligible part in the


Table 11. Comparison of k ~ Values m for Runs Conducted with KHCO$ and C02 h-' 10 wt % std 30 wt % ionic 30 wt % std temp, O C (KHCOs runs) (COz runs) (COI runs) 205 175 150

--

0.29 0.121 0.031

0.104 0.026 0.0053 I

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Table 111. Comparison of k D E A for Molecular, Ionic, and Standard Runs DEA k r m , h-' concn, wt molec + 75 temp, O C molec ionic std ionic 100 66.7 40 30 30 30

0.101 0.0242 0.0055 I

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205 205 205 205 175 150

0.195 0.175 0.168 0.14 0.075 0.0203

0.104 0.025 0.0053

0.195 0.3 0.32 0.29 0.121 0.031

0.244 0.10 0.0253

Table IV. Comparison of ~ D E Afor Thermal and Standard Runs Performed with an Initial DEA Concentration of 30 wt %

h-'

1.0

1

temp, "C

thermal

std

205 250

0.00365 0.033

0.29 0.69

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Figure 10. DEA concentration 88 a function of time and initial DEA concentration (no water present, 4137 kPa COz, 205 "C).

degradation of DEA. Using KHC03 resulted in degradation similar to that observed when using pure C02,but the rates are much slower. Table I1 summarizes the initial k D m values for the KHC03 runs and equivalent C02runs. In addition, the k D m values for runs where 10 w t % ' DEA is degraded under C02 are listed. The similarity between the kDEA values for the KHC03 (ionic) runs and the C02 runs for 10 w t % DEA is noteworthy. It may therefore be concluded that, at low DEA concentrations, degradation is primarily due to HC03- and not C02. The major difference between the ionic runs and the standard C02runs is the production of HEOD. Very little HEOD is formed by using KHC03, which may be explained by the fact that the solution is more basic than under standard conditions. As mentioned before, increasing the alkalinity causes HEOD to break down and to inhibit the production of HEOD. From these results it may be concluded that HC03- can aid DEA degradation in a similar way to C02. However, other reactions must also take place since the degradation increases when pure COz is used. Effeet of Water ("Molecular Runs"). To investigate degradation by other than ionic routes, degradation experiments were conducted with C02 in the absence of water. This was achieved by diluting DEA with methyldiethanolamine, MDEA. MDEA is similak to DEA but is not degraded by C02under the operating conditions of this study. Since water was absent from the reaction mixture, ions could not be formed and hence degradation could only be caused by C02 reacting directly with DEA. Figure 10 shows the results of some runs conducted at 205 OC. Once again, the three major degradation products were produced in the same relative amounts as in the standard runs; however, the rate is lower and kDm decreases slightly with initial DEA concentration. Thus it appears that DEA degradation occurs along two parallel paths, one involving pure COP and the other HC03-. Table 111 gives the values of kDEA for the runs without water ("molecular runs") and compares these with those for the ionic and standard runs. If DEA is degraded by two parallel, first-order reactions, then basic kinetic theory indicates that the overall degradation rate constant is the sum of the individual constants, i.e. (3) (kDEA)ovedl = (kDEA)m~lecular -k (kDEA)ionic

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Figure 11. DEA concentration as a function of time and temperature (no C02 present, 30 w t % DEA, 4137 kPa N2).

Referring to Table 111, it can be seen that the sum of the k values for the two degradation routes is close to the k values for the standard run. In all cases, the sum of the rate constants is slightly lower than that obtained in the standard runs; this is probably due to the fact that the molecular runs excluded water. In the standard runs water is always present and it seems likely that water can help the degradation, thus increasing kDEA for the molecular route. Table 111 shows that the kDEA for the molecular route decreases slightly with decreasing DEA concentration. A possible explanation is that since water is a degradation product, at high concentrations of DEA more water is produced during the degradation which can aid further degradation. Thermal Degradation. Although DEA does not break down noticeably at 205 "C over 8-h periods, it was observed to degrade over longer periods of time and at higher temperatures as shown in Figure 11. The plots are linear, indicating a first-order reaction with the value of kDEA being about one-hundredth of that obtained under standard conditions at 205 "C (see Table IV). BHEP and THEED were the major degradation products, and their corresponding concentration changes indicated that a series reaction is occurring, i.e. DEA THEED BHEP (4)

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Although the thermal degradation of DEA appears to be first order, this does not agree with the chemical equations since two molecules of DEA are required to produce one molecule of THEED or BHEP. Therefore, the reaction is not as simple as it appears. It is possible that an intermediate is slowly produced from DEA via a first-order reaction, which is then rapidly converted to THEED. For example, THEED can be produced from DEA via the

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Figure 14. DEA concentration as a function of time (reactantsdegraded DEA solution, 4137 kPa COP or N P , 175 "C). TIME

Figure 12. Typical plots of concentration as a function of time for DEA and its degradation products. I

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Figure 15. HEOD concentration 88 a function of time (reactantsdegraded DEA solution, 4137 kPa COPor Nz,175 "C).

Figure 13. Concentration of DEA, HEOD, THEED, and BHEP as a function of time (reactants-DEA and HEOD, 4137 kPa N P ,205 "C).

formation of HEM (Kennard, 1983). The results clearly show that BHEP can be produced from THEED. This does not agree with the mechanism proposed by Polderman and Steele (1956). The major degradation products (BHEP and THEED) are formed in similar amounts, albeit at quite different rates, in the thermal runs under nitrogen to the standard runs under COz. The thermal route may therefore be considered as a third degradation route with C02 acting as a catalyst. Experiments with Degradation Compounds. In general, the degradation of DEA and production of its degradation compounds can be summarized by the qualitative plot shown in Figure 12. The plot suggests that BHEP is produced in a series reaction from DEA via THEED. This hypothesis needed to be confiied and also the role of HEOD had to be understood better. A series of experiments was therefore conducted with solutions of the degradation compounds. BHEP Runs. The stability of BHEP was first tested in run a and it was found that BHEP did not degrade. In run b (see Table V), BHEP was mixed with DEA and then kept under a blanket of nitrogen at 205 "C and 4137 kPa. No reaction was observed and therefore it may be concluded that BHEP is a final degradation product. Runs c and d were performed to determine whether the presence of BHEP and other degradation products had any effect on the overall DEA degradation. Degradation appears to be inhibited by the presence of degradation products, especially BHEP (Kennard and Meisen, 1980). Run d was continued for 50 h, after which the reaction mixture was virtually identical with that of a standard run and the addition of BHEP to the reaction mixture had no effect on the overall production of BHEP. In a similar experiment (run e) where the temperature was reduced to 150 "C, no inhibiting effect due to BHEP was observed. Thus it appears that degradation products (and/or BHEP) can tie up some of the available C 0 2dissolved in the reaction mixture. If the concentration of COPis reduced below 0.2

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Figure 16. THEED concentration as a function of time (reactants-degraded DEA solution, 4137 kPa COz or N2, 175 "C).

g of CO,/g of DEA, then the reaction will be slowed down as C 0 2 becomes limiting. HEOD Runs. In run f, an aqueous solution of HEOD was heated to 205 O C for 1 h under a blanket of nitrogen. DEA, THEED, and a trace of BHEP were produced. This indicated that there is some form of equilibrium between HEOD and DEA. It is ylikely, however, that there is any equilibrium between HEOD and THEED or BHEP since neither compound was able to revert to HEOD as observed from runs a and k. Figure 13 shows the Concentration vs. time curves for the degradation of a mixture of HEOD and DEA under nitrogen (run g). HEOD is seen to be mainly converted to THEED and a trace of BHEP; the DEA loss is small. The feeds used in runs h and i were produced by degrading a 30 wt % DEA solution at 175 "C under 4137 kPa of C 0 2 for 6 h. The solution was then removed from the autoclave and gently heated to drive off any dissolved COz. Samples of the resulting mixture were then placed (at 175 "C, 4137 kPa) under a blanket of COP(run h) or N2 (run i). Figures 14 to 17 show the results of these runs. For run h it appears as if HEOD is playing little part in the degradation of DEA and its concentration remains nearly unchanged. The THEED concentration increases and then falls slightly, whereas the BHEP concentration in-


Table V. Summary of Runs to Study the Behavior of the Major Degradation Compounds feed concn, mol/cm3 temp, "C gas THEED BHEP DEA HEOD run 5 x 10-4 205 COZ 5 x lo4 3 x 10-3 205 N, 205 co, 2.7 x 10-4 3.3 x 10-4 3.55 x 10-4 3 x 10-3 4.7 x 10-4 3 x 10-3 205 COP 4.7 x 10-4 3 x 10-3 150 COZ 3 x 10-3 205 NZ 4.24 x 10-4 1.0x 10-3 205 N2 0.28 x 10-4 4 x 10-4 9.7 x 10-4 1.5 x 10-3 175 COZ 4 x 10-4 9.7 x 10-4 0.28 X lo4 1.5 x 10-3 N, 175 2.6 x 10-3 205 co, j 2.6 x 10-3 205 N2 k 205 N2 1 1.2 x 10-3 2.6 x 10-3 205 COZ m 1.2 x 10-3 2.6 x 10-3 -

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comments BHEP studies BHEP studies BHEP studies BHEP studies BHEP studies HEOD studies HEOD studies HEOD studies HEOD studies THEED studies THEED studies THEED studies THEED studies

8 8

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Scheme I HEOD

8

rapid

DEA t COz

SlDU

RzNCOO- t Ht

THEED t H 0 3 -

(5)

HCOC t R NH;

(6)

Scheme I1 HEOD -

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Figure 17. BHEP concentration as a function of time (reactantsdegraded DEA solution, 4137 kPa CO, or N,, 175 "C).

creases steadily. When the same reactions were carried out under nitrogen (run i) some major differences were noted. The rate of degradation was less than that under CO,, but still quite significant. Since it was established earlier that DEA does not degrade noticeably at 175 "C under nitrogen, it is clear that CO, and/or HCO, are provided either by the breakdown of HEOD to DEA or the formation of THEED. It can be seen from Figure 15 that HEOD does not break down completely, but its concentration levels off after an initial, sharp fall. This indicates the establishment of an equilibrium between HEOD and its breakdown products or DEA. What appears to be taking place is that HEOD reacts to form either DEA or THEED and HCO,. The bicarbonate ion then reads with DEA to produce either more HEOD or other degradation products. In the case of run i it seems that DEA just degraded to THEED once sufficient HC03- or C02were provided by the COz atmosphere. HEOD seems to be involved in a reversible, complex reaction with DEA. I t is possible that HEOD is in equilibrium with a degradation product of DEA which is, in turn, in equilibrium with DEA. It is proposed that this intermediate compound is the DEA carbamate (R,NCOO-), which is in equilibrium with both DEA and HEOD. (DEA carbamate is written as R,NCOO- since, in aqueous solution, it can only exist as an ionic complex. Carbamic acid, hNCOOH, is extremely unstable and reverts to DEA and CO,; R2NCOOH has never been isolated.) It is the carbamate (and not HEOD) that reacts with itself or more DEA and produces THEED according to Scheme I. There are problems with Scheme I since it indicates that the concentration of HEOD should increase to a maximum and then fall as R,NCOO- is slowly converted to THEED. In some cases this behavior was observed experimentally. However, generally the concentration of HEOD leveled off

r-

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1.5 -0

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TIME (hr)

Figure 18. Concentration of DEA, THEED, and BHEP as a function of time (reactants-DEA and THEED, 4137 kPa N2, 205 "C).

(see Figure 3). To explain this, it must be remembered that the CO, concentration for most high-temperature run lies between 0.2 and 0.3 g of CO,/g of DEA or about 0.45-0.7 mol of C02/mol of DEA. Therefore, DEA was initially in excess for the high-temperature runs. The formation of THEED tends to produce CO, or, depending on the experimental conditions, HC03-. These ions are able to react with excess DEA and maintain constant levels of carbamate ions and HEOD. What appears to be happening is that the excess DEA is slowly converted to THEED via the formation of the carbamate. When no more excess DEA is available, the concentrations of carbamate and HEOD fall. The modified scheme thus becomes Scheme 11. Unfortunately, the chromatographic analysis is unable to differentiate between &NH, R2NH2+, and RzNCOO-. This causes difficulties in confirming the above mechanism which therefore remains hypothetical. THEED Runs. When an aqueous solution of THEED was heated to 205 "C under 4137 kPa C02for 1h (run j), the only product was BHEP. A similar run with nitrogen (run k) also produced BHEP, but in much smaller quantities. Therefore, it appears that BHEP can be directly

136


The effect of C 0 2 causes the value of the reaction rate constant to increase about fivefold. It is likely that the COz increases the rate of reaction in a manner similar to its role in DEA degradation, i.e., via the formation of a carbamate (in this case THEED carbamate).

1 i

T

-

-

COMPOUND 0--DE* b-

HE00

0-

MEED

0- BHEP

Development of Reaction Mechanisms A model was developed by Kennard (1983) to explain the production of the main degradation products. Figure 20 shows the major reactions which are believed to be responsible for the degradation of DEA by C02. Certain reaction steps cannot be fully confirmed since they are based on the existence of DEA carbamate. Unfortunately, it was impossible to isolate and detect the carbamate, and certain aspects of the mechanism therefore remain conjectural. In general, the experimental results indicate that DEA degrades via three routes: the fast “molecular route”, the slower “ionic route”, and the very slow “thermal route”. The molecular route involves C 0 2 reacting directly with DEA to produce the carbamate, whereas the ionic route produces the carbamate from the amine salt R2NH2+HC03-. The carbamate then degrades slowly to produce THEED, which in turn loses water to form BHEP. HEOD is a side product, produced by internal dehydration of the carbamate. It is believed that the formation of HEOD does not contribute to the overall degradation of DEA. C 0 2 appears to catalyse the degradation of DEA to THEED and of THEED to BHEP via the formation of carbamate complexes. Anomalous Experimental Observations. Dependence of kDEA on DEA Concentration. Referring to Figure 7, three distinct regions are observed: (1) region &lo% DEA: in this region the main degradation route appears to be the slow ionic route; (2) region 10-30% DEA:

TIME [hrl

Figure 19. Concentration of DEA, HEOD, THEED, and BHEP as a function of time (reactanta-DEA and THEED, 4137 kPa COB,205 OC).

produced from THEED with C 0 2 acting as a catalyst. Runs 1 and m were conducted to determine the effect of D M on the reactions of THEED. Figure 18 (run 1)shows that the concentration of DEA remained unchanged for 8 h, which indicates that DEA and THEED do not react. Figure 18 also shows that as the THEED concentration decreases, the BHEP concentration increases correspondingly, This suggests a stochiometric relationship between THEED and BHEP. Figure 19 indicates that C 0 2 speeds up the conversion of THEED to BHEP (runm). DEA degrades slightly due to the presence of C 0 2 . I t is interesting to note that the degradation of the DEA in this run is much lower than it would have been if THEED were absent. It appears that THEED, like BHEP, is able to reduce the availability of COz and HC03- for DEA attack. The results also indicate that the rate of conversion of THEED to BHEP follows a first-order reaction initially. 0

I

/ c\

R-N

(MOD)

P,NH

+ coy e R,NH+COO-

-

0

I - I

CH,

11

+

H,O

COY

+

OH-

+ H,O R,NH

CHI

R,NCOO-

CO,

+ H+

__.

+

H,O

=

R,NH,+

R,NH,+

[HCO,-l

UC0,-

t U+

UC0,-

I

+ OH-

I [R,NH,+]

= = +

I

HC0,-

or H,O-

or R , N H , +

--+

I 1,

- “EH,

W-CzMr-O-CzUr-HRH

(BHEP)

( B W )

Figure 20. Schematic diagram showing the possible routes for the degradation of DEA.

major degrrdmcios routma

minor degrrdmtion route#

@

b l e c u l r r route

@

Ionic route

@

~htrculroute


Table VI. Principal DEA Degradation Routes under Various Conditions total pressure temp, "C DEA concn, w t % psi kPa route' ionic 0-10 600 4137 90-175 600 4137 ionic + molecular 10-30 90-175 600 4137 mainly molecular 30-100 90-175 600 4137 ionic + thermalb 0-10 175-250 600 4137 10-30 175-250 ionic + molecular + thermal 600 4137 ionic + molecular + thermalc 30-100 175-250

-

-

limiting compd

HZO

-

coz

-

-

137

COz+HzO

"The routes are (a) ionic: RzNHz++ HCO; RzNCOO-H+ products; (b).molecular: RzNH + COP RzNCOO-H+ products; (c) thermal: RzNH products. b A t high temperatures the thermal route will start to contribute to the degradation, although only to a small extent. CAthigh temperatures and DEA concentrations (>30 wt %), the ionic route contributes more to the degradation than at lower temperatures where the molecular route is responsible for most of the degradation.

as the concentration increases, the degradation proceeds by a combination of the ionic and molecular routes with the molecular route gaining dominance; the overall kDEA therefore becomes the sum of the k values for the two parallel degradation reactions (7) kDEA = kDEA,ionic -k k D E A , ~ ~ l e ~ ~ l a r since the thermal route is much slower than either the ionic or molecular routes, its contribution to degradation should be negligible; (3) region 30-100% DEA: as the concentration of DEA is raised (and the concentration of water falls), the reaction becomes limited by water. Arrhenius Plots. As shown in Figure 2 the data tended to deviate from the linear form at high temperatures. One reason for this deviation could be that the C02 solubility decreases and the C02 concentration becomes limiting. Another possible reason is that the ionic route becomes increasingly important with rising temperature (see Table 111) because the amount of R,NCOO- produced by the molecular route falls. Change in DEA Concentration with Time. I t was observed that, at high temperatures, the semilogarithmic plots of DEA concentration vs. time were only linear for a few hours and then began to curve. This indicates that the initial, pseudo-first-order degradation reaction was inhibited as the reaction progressed. The inhibition could be due to the following reasons. (a) A t high temperatures the concentration of C02 is very close to the critical value of 0.2 g of C02/g of DEA. Any reduction in this level causes the IZDEA to fall. (b) It has been shown that the presence of degradation compounds inhibits degradation at high temperatures by tying up some of the available COP (c) As the degradation proceeds, the C02is converted to HC03- via the formation of the degradation products and R2NCOO-. Therefore, the production of additional R2NCOO- must proceed through the ionic route which is slower than the overall rate. (d) As the reaction proceeds, the mixture becomes more acidic due to the absorption of COz and the reduction of DEA. (Although BHEP and THEED are alkaline, two molecules of DEA are required to produce one mole of BHEP and THEED; hence the number of alkaliie species falls). Experiments have shown that reducing the pH reduces the degradation rate. Summary. Table VI gives the main degradation routes for the range of operating conditions studied. Kinetic Model The purpose of the kinetic model is to predict, quantitatively, the degradation of DEA and the production of its degradation compounds. Since the degradation mechanism (Figure 20) is complex, it was simplified by invoking the following arguments. (1) Under industrial conditions it is unlikely for temperatures to exceed 150 "C. Since the Arrhenius equation was not obeyed above 175 "C, the kinetic model was made applicable only up to 175

+

"C. (2) The effect of COz could be neglected provided the C02concentration was greater than 0.2 g of C02/g of DEA. This condition is met at low temperatures and high C02 partial pressures. (3) Based on the assumption that DEA degradation is governed by a pseudo-first-order reaction, experiments showed that kDEA was not independent of initial DEA concentration (Figure 7). To account for this effect, a series of Arrhenius plots could be produced similar to Figure 8, which cover the various DEA concentration ranges. (4) The equilibrium between COz and DEA and the formation of %NCOO- are established within a matter of seconds. Therefore, these initial fast reactions may be ignored compared to the slow degradation reactions. (5) For simplicity, the model will treat the carbamate ion, R2NCOO-, as DEA. (6) The ionic and molecular routes both result in identical degradation products. Therefore they may be considered as a single route. (7) Since the thermal route is much slower than the other routes, it may be ignored. Considering these simplifications, the degradation reactions may be represented by the following set of equations DEA DEA

k

HEOD

(8)

THEED

(9)

THEED

BHEP

(10)

These reactions lead to the following rate equation d[DEA] dt

---

- (k + k'?[DEA]

- kTHEOD]

(11)

which is not of the pseudo-first-order form indicated by the experiments. To overcome this difficulty, it was decided to make the production of HEOD an irreversible reaction. This could be justified since, at low temperatures, the equilibrium between HEOD and RzNCOO- is not reached for a very long time because the HEOD concentration does not level off. Furthermore, the concentration of HEOD when compared to that of DEA is very much smaller and slight errors in the prediction of HEOD concentration should not impair the overall model performance substantially. Thus the model can be represented by the following scheme

y DEA

&% THEED

*3

BHEP

(12)

The range of conditions covered by the model is temperature, 90 to 175 "C, DEA concentration, 0 to 100 wt %,

138


I

0 01

E

I

I

I

I

I

I

I

I

I

I

2.3

2.4

2.5

1.6

2.7

L \\

l

0.00OI

-

-

1.3

lOOO/l ( O K - ' )

Figure 21. Arrhenius plots for k, a t various initial DEA concentrations (DEA HEOD).

1.1

2.5

2.6

lOOO/T (OK") Figure 22. Arrhenius plots for k, at various initial DEA concentrations (DEA THEED).

-

and C02 loading, >0.2 g of C02/g of DEA. The rate expressions corresponding to this model are d [DEA] = -(k1 dt

+ kz)[DEA]

d[HEOD] = kl[DEA] dt d[THEED] = kz[DEA] - k,[THEED] dt

(15)

d[BHEP] = kS[THEED] dt These equations can be solved to give [DEA] = [DEA]oe-(kl+k~)t

(17)

\ ::I 0 w01

2.3

2.4

IOOO/T

2s

2.6

PK')

-

Figure 23. Arrhenius plot for k3 a t various initial DEA concentrations (THEED BHEP).

Using these equations and the experimental results, it was possible to calculate the various k values. Figures 21 to 23 show the Arrhenius plots for kl, It2, and k3 As expected, the plots for k, and k2 both conform to the three concentration regimes observed in the Arrhenius plot for ,k (Figure 8). In general, the lower curve covers concentrations ranging from 0 to 10 w t % DEA and the upper curve covers concentrations between 30 and 100 wt % DEA. For concentrations in the range 10 to 30 w t % DEA there is a series of curves. The Arrhenius plot for k3

TIME [hr)

Figure 24. Comparison between the experimental and theoretical values of DEA, HEOD, THEED, and BHEP concentrations as a function of time (20 wt % DEA, 4137 kPa, 140 "C).


(Figure 23) is unaffected by DEA concentration and tends to confirm the fact that BHEP is produced from THEED. It is interesting to note that extrapolation to 205 OC gives a k value which agrees very closely with the value of 0.25 h-' calculated from the results of run m where THEED was degraded under C 0 2 to BHEP. By use of these plots (Figures 21 to 23) and eq 13 to 20, the degradation of DEA could be predicted for a given set of conditions. The model gave a very good prediction of the concentration of DEA and its main degradation compounds for various reaction times. For example, Figure 24 shows a comparison between the experimental results and predictions of the model for the degradation of 20 wt % DEA a t 140 O C . In other cases, the model tended to overpredict the concentration of HEOD. This was to be expected since no account was taken for the reversible reaction between HEOD and DEA (or more correctly RZNCOO-) .

Purification of DEA Solutions Activated carbon filters are used in many natural gas treating plants to purify degraded DEA solutions. Although it has been claimed by some that the filters are very successful, their general effectiveness has yet to be proven. In order to determine whether activated carbon can remove degradation compounds, samples from industrial filter units were tested and a series of experiments was conducted in the laboratory. Analysis of samples taken upstream and downstream of an activated carbon filter in a large gas plant showed that there was no change in the concentration of the major degradation compounds. This clearly indicates that the filter was ineffective in removing the shown degradation compounds. Similarly, samples of DEA solutions degraded in the laboratory were contacted with activated carbon for periods ranging from a few hours to a few weeks at varying temperatures. In none of the experiments was the activated carbon found to change the concentration of the degradation compounds significantly. Although none of the major degradation compounds were removed, contact with activated carbon did change the color of the degraded DEA solutions from dark brown to light yellow. The industrial effectiveness of activated carbon may lie in its ability to remove heavy hydrocarbons or surface-active compounds which may cause foaming. However, there is no evidence that activated carbon filters are able to remove any of the major degradation compounds. Practical Implications of the Present Study (a) The Effect of Temperature. To minimize DEA degradation, the design and operation of DEA units should avoid the creation of elevated temperatures throughout the plant. The heat-transfer surfaces of the stripper reboiler (especially when gas fired) are particularly prone to the formation of hot spots. To prevent such hot spots, the DEA circulation should be kept high and the steam or gas temperatures should be kept low. If, for some reason, the DEA circulation should decrease, immediate action must be taken to reduce the steam pressure or fuel gas flow to the reboiler. There are two other sites where major degradation may take place. The first is within the exchanger which heats the rich amine streams with the lean amine stream. In some cases the temperature of the rich amine stream may be as high as 125 "C. The second site is the base of the absorber. If the C02content of the raw natural gas is high, the temperature of the rich amine at the base of the absorber may rise substantially.

139

In many DEA units only the bulk solution temperatures are measured. It must be remembered that the skin temperatures of heat transfer surfaces can be very much higher, particularly during process upsets. Reliance on bulk temperatures is therefore inadequate. (b) Effect of Pressure. The partial pressures of C 0 2 should be kept as low as possible in order to minimize DEA degradation. Although it is not usual to exercise control over the COPcontent of the raw gas entering a plant, it may be possible to dilute the raw gas with a sweeter gas. This dilution would not only reduce the degradation rate, but it would also reduce the heat of absorption when C 0 2 is dissolved in the DEA. This measure would therefore help to keep the overall temperature in the absorber low. (c) Effect of DEA Concentration. From the point of minimizing degradation, plants should operate with dilute DEA solutions (if possible below concentrations of 20 wt %). However, limitations are imposed by the desired plant capacity. Future design of gas-treating plants should consider larger equipment for operation with dilute solutions of DEA. Again, using dilute solutions of DEA would minimize temperatures in the absorber. However, studies would have to be made to determine the cost effectiveness of these measures. (d) Effect of Activated Carbon Filters. Although activated carbon appears unable to remove the major degradation products, it is not recommended to remove the filters from existing plants. The filters may serve other useful functions such as removing surfactants which can cause foaming, heat-stable salts which may cause corrosion and they may also act as a means for removing fine particulates. Acknowledgment The financial support provided by the Canadian Gas Processors Association and The University of British Columbia is gratefully acknowledged. Nomenclature BHEAE = bis(hydroxyethylaminoethy1) ether BHEED = N,N-bis(hydroxyethy1)ethylenediamine BHEI = N,N-bis(hydroxyethy1)imidazolidone BHEP = N,N-bis(hydroxyethy1)piperazine DEA = diethanolamine HEED = N-(hydroxyethy1)ethylenediamine HE1 = N-(hydroxyethy1)imidazolidone HEM = N-(hydroxyethy1)ethyleneimine HEOD = 3-(hydroxyethyl)-2-oxazolidone HEP = N-(hydroxyethy1)piperazine k = reaction rate constant kDm = overall reaction rate constant for the degradation of DEA, h-I k,, k2, k3 = rate constants used in the kinetic model of the degradation of DEA, eq 13 to 20, h-' MDEA = methyldiethanolamine MEA = monoethanolamine OZD = oxazolidone R- = -CZH,OH T = absolute temperature, K t = time, h TEA = triethanolamine TEHEED = N,N,N,N-tetra(hydroxyethy1)ethylenediamine THEED = N,N,"-tris(hydroxyethy1)ethylenediamine [ ] = concentration, g-mol/cm3 [

lo = initial concentration, g-mol/cm3

Registry NO.DEA, 111-42-2;COZ, 124-389; HXOD, 3356-88-5; THEED, 60487-26-5; BHEP, 122-96-3.

Literature Cited Blanc, C.; Grail, M.; Demarais, 0. Roc. Gas Cond. Conf. March 1982,32. Dankwerts, P. V. Chem. Eng. Sci. 1979. 3 4 , 443. Hakka, L. E.; Singh, K. P.; Batta, G. L.; Testart, A. C.; Andrejchyshyn, W. M. Gas Process.lCanada 1P68,67(1), 32.

140

Ind. Eng. Chem. Fundam. 1985, 2 4 , 140-147

Kennard, M. L. Ph.D. Thesis, Unlversity of British Coiumbla, Vancouver, BC, 1983. Kennard, M. L.; Meisen. A. J . Chromtcgr. W83, 267,373. Kennard, M. L.; W e n , A. l-@drocarbon Process. 1980, 59(4), 103. Kennard, M. L. Melsen, A. J . Chem. Eng. Data 1984, 29(3),309. Nonhebel, G. “Gas Pwlficatlon Processes for Air Pollution Control”, 2nd ed.; Newness-Butterworths: London, 1972. Polderman, L. D.; Steele, A. B. Oil Gas J . 1958, 54(65), 206.

Scheirman, W. L. Hydrocarbon Process. 1073, 53(7),95. Smlth, R. F.;Younger, A. H. Roc. Gas Cod. Conf. 1972, 22(E),1. Smlth, R. F.; Younger, A. H. Hydrocarbon Process. 1972, 57(7),98. Younger, A. H. Roc. Gas Cond. Conf. 1973, 23(E),1.

Received for review September 13, 1984 Accepted July 9, 1984

Nonideal Behavior in Liquid Metal Solutions. 1 Physical Theory Model Thomas Stolcos and Charles A. Eckert’ Department of Chemical Engineering, University of Illinois, Urbana, Iilinois 6 180 7

A “physical” model of liquid metal solutions, based on the concept of electron transfer, has been used to calculate mixture propertles for systems with positive or small negative deviations from Raouit’s law. I t gives very good quantitative results for excess Oibbs energies and excess enthalpies, and qualitatively correct resuits for excess volumes. I t has been used successfully to predict solid-liquid and liquid-liquid equilibria, and the extension to multicomponent systems is discussed. The physical significance of the charge transfer parameter is described.

Introduction The thermodynamics of liquid metal mixtures is rapidly becoming more important as an increasing number of metallurgical processes depend heavily on solution chemistry. For example, in some cases aluminum metal is used to deoxidize steel and often pig iron is desulfurized with caustics; in both cases the process depends on activities in liquid solutions. Some processes depend on very large positive deviations from Raoult’s law, to give separations by liquid-liquid equilibria. These include the lead-copper system, which exhibits a relatively small miscibility gap, and the lead-zinc system, where the gap is substantially larger. There are also a number of processes which depend on large negative deviations from ideality, such as the carbothermic and carbonitrothermic solvent metal reductions of reactive metal oxides (Anderson and Parlee, 1974, 1976) or the liquid tin repurification of spent uranium fuel elements (Anderson and Parlee, 1971). Liquid metal solutions tend to have relatively larger nonidealities than the mixtures of organic chemicals with which chemical engineers are familiar. Despite the higher temperatures, the interatomic forces are a great deal stronger for metals, and few mixtures exhibit anything like ideality. Mixtures can show such varied and complex behavior that there exists no physical explanation which is extensive enough for the general description of all the actual deviations from ideality. The utility of different expressions based on semiempirical interpretations such as the Wohl expansion, the Scatchard formula, the van Laar equation, the Wilson equation, and the NRTL equation are discussed by Tomiska (1980) and Eckert et al. (1982) and were shown to be generally not applicable for regression of experimental data. Depending on the nature and the degree of the nonideality, the metallic solutions can be classified into two main categories. First, there are systems which show large or small positive or small negative deviations from Raoult’s law, resulting from relatively weaker, nonspecific interactions between the components, which are characterized as physical forces. This paper deals with the physical model for such systems. Second, there are solutions which 0196-431 3/85/1024-0 140$01.50/0

show pronounced negative deviations from ideality and very abrupt changes of transport properties such as viscosity and density at specific compositions. These deviations are attributed to strong specific forces which lead to the formation of new molecular species; such forces are called chemical forces for which a more extensive discussion follows in part 2 (Stoicos and Eckert, 1985). Attempts to calculate the excess thermodynamic properties of simple liquid alloys have been made from firstprinciple methods. Christman (1967),Tamaki and Shiota (1968),Umar et al. (1974), and Hafner (1976) have all used some form of the pseudopotential approach, coupled with an energy minimization procedure for calculating the necessary parameters. These calculations have been applied to a few systems only, mainly liquid alkali alloys, and have not been extended to a wide category of metallic solutions. Paulaitis and Eckert (1981) proposed a physical model which combined corresponding-states theory with a perturbed hard-sphere representation for predicting mixture properties. Another model based on an electron theory which includes both structure-dependent and structureindependent contributions has been developed by Cox (1979). Both of these models contain one parameter which was obtained from mixture data. The models have been successfully used for describing the thermodynamic behavior of many binaries and have some generality. The physical model developed and presented here is based on an electron theory coupled with an empty-core pseudopotential form. As in the two previously described approaches, it contains one model parameter. Positive and small negative deviations from Raoult’s law are attributed to a certain amount of charge transferred between dissimilar metals upon mixing. The important elements in the proposed model are that all the equations have a theoretical significance, and the charge-transfer parameter is based on a phenomenon of plausible physical importance. By and large, experimental data on liquid metal solutions are very difficult to take, because of the high temperature, or they are fragmentary, because of freezing or 0 1985 American Chemlcal Society

Mechanisms and kinetics of diethanolamine degradation - Industrial

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