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International Journal of Greenhouse Gas Control 5 (2011) 405–412
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Heat of absorption of CO2 in aqueous ammonia and ammonium carbonate/carbamate solutions Feng Qin a , Shujuan Wang a , Inna Kim b , Hallvard F. Svendsen b,∗ , Changhe Chen a a b
Department of Thermal Engineering, Key Laboratory for Thermal Science and Power Engineering of Ministry of Education, Tsinghua University, 100084 Beijing, China Department of Chemical Engineering, Norwegian University of Science and Technology, 7491 Trondheim, Norway
a r t i c l e
i n f o
Article history: Received 13 August 2009 Received in revised form 15 March 2010 Accepted 7 April 2010 Available online 4 May 2010 Keywords: Aqueous ammonia CO2 Heat of absorption Calorimeter
a b s t r a c t A reaction calorimeter was used to determine the enthalpies of absorption of CO2 in aqueous ammonia and in aqueous solutions of ammonium carbonate at temperatures of 35–80 ◦ C. The heat of absorption of CO2 with 2.5 wt% aqueous ammonia solution was found to be about 70 kJ/mol CO2 , which is lower than that with MEA (around 85 kJ/mol) at 35 and 40 ◦ C. The value decreases with increased loading, but not to as low a value as expected by the carbonate–bicarbonate reaction (26.88 kJ/mol). The enthalpy of absorption of CO2 in aqueous ammonia at 60 and 80 ◦ C decreases with loadings at first, then increases between 0.2 mol CO2 /mol NH3 and 0.6 mol CO2 /mol NH3 , and then decreases again. The behavior of the heat of absorption of CO2 in 10 wt% ammonium carbonate solution was found to be the same as that of aqueous ammonia at loadings above 0.6 mol CO2 /mol NH3 . The heat of absorption increases with increasing temperature. The heats of absorption are directly related to the extent of the various reactions with CO2 and can be assessed from the species variation in the liquid phase. © 2010 Elsevier Ltd. All rights reserved.
1. Introduction Chemical absorption is a promising method for CO2 capture and is also one of the most common. It is necessary, however, to reduce the high capital costs of the method and the energy requirements in the absorption and desorption processes. Power plant efficiency decreases by about 10 percentage points when the flue gas is treated for CO2 capture by chemical absorption (Rhudy et al., 2006; Wolf et al., 2006), so developing a novel solvent with low regeneration energy is significant for reducing energy penalties. Aqueous ammonia solutions are thought to have a potential as effective and economic absorbents for CO2 capture. The reactions between aqueous ammonia solutions and CO2 are very fast, and in batch or semi-batch experiments the CO2 removal efficiency can reach 90% or more (Diao et al., 2004; Kim et al., 2008; Resnik et al., 2006; Yeh et al., 2005; Yeh and Bai, 1999). This means that aqueous ammonia is very effective for CO2 absorption. Previous research has shown that the loading capacity of aqueous ammonia can be high (Yeh et al., 2005; Yeh and Bai, 1999), and much higher than that of monoethanolamine (MEA). The corrosion problems caused by aqueous ammonia are no worse than those caused by amines. Also, with aqueous ammonia there are no thermal degradation and oxidative degradation problems like those arising with amine systems, and the system can scrub SOx ,
∗ Corresponding author. Tel.: +47 73594100; fax: +47 73594080. E-mail address: [email protected] (H.F. Svendsen). 1750-5836/$ – see front matter © 2010 Elsevier Ltd. All rights reserved. doi:10.1016/j.ijggc.2010.04.005
NOx and other acidic gases simultaneously. Furthermore, ammonia is cheaper than MEA (Yeh et al., 2005). Most of the literature focuses on parametric analyses of the capture process. Liu et al. (2009) reviewed the progress of CO2 capture using aqueous ammonia, and summarized the effects of reactor type, reaction temperature, ammonia concentration, CO2 concentration, and SOx , NOx and heavy metals on the CO2 removal efficiency. Data by Yeh et al. (2005) indicate that the energy requirement for CO2 regeneration using the aqueous ammonia process could be at least 75% less than that of using MEA. The semi-continuous experiments by Yeh and Bai (1999) also indicate that the heat of absorption in aqueous ammonia is much lower than it is in MEA, and monitored temperature variations show that endothermic reactions play a role at the beginning of the absorption. It has been suggested in a number of papers that the regeneration heat requirement can be much smaller than in MEA-based processes such as the NETL and ALSTOM processes, see Black (2006), Ciferno et al. (2005), Resnik et al. (2006), Wolf et al. (2006), and Yeh et al. (2005). In all these studies on aqueous ammonia based CO2 capture system, the heats of reaction obtained or presumed were about 26–27 kJ/mol CO2 . All these evaluations, however, are based on the assumption that carbonate reacts with CO2 to form bicarbonate. It is thus very important to validate these assumptions. 2NH4 HCO3 (aq) ↔ (NH4 )2 CO3 (aq) + CO2 (g) + H2 O Hrx = 26.88 kJ/mol
(1)
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Nomenclature Symbols and constants f fugacity (Pa) Habs (Hs , Hdiff ) heat of absorption (kJ/mol CO2 ) CO2 absorbed (mol) nCO2 Qtotal heat (J) R universal gas constant (8.314 J/mol/K) T temperature (K)] Greek symbol ˛ CO2 loading (mol CO2 /mol NH3 )
NH4 HCO3 (aq) ↔ NH3 (aq) + CO2 (g) + H2 O Hrx = 64.26 kJ/mol
(2)
(NH4 )2 CO3 (aq) ↔ 2NH3 (aq) + CO2 (g) + H2 O Hrx = 101.22 kJ/mol
(3)
NH2 COONH4 (aq) ↔ 2NH3 (aq) + CO2 (g) Hrx = 72.32 kJ/mol
(4)
NH2 COONH4 (aq) + H2 O ↔ NH4 HCO3 (aq) + NH3 (aq) Hrx = 8.06 kJ/mol
(5)
When CO2 is fed into aqueous ammonia solutions, a mixture of ammonium carbonate, bicarbonate, and carbamate is produced at lower loadings. The carbamate and carbonate can further react with CO2 to produce bicarbonate, as shown in Eqs. (1)–(5) (Kohl and Nielsen, 1997; Yeh et al., 2005). Even in aqueous ammonia solutions at higher loadings (above 0.5 mol CO2 /mol NH3 ), experimental analyses and NMR results have shown that there is still a lot of free ammonia in solution (Gao et al., 1991; Holmes et al., 1998; Mani et al., 2006; Wen and Brooker, 1995). This suggests that in aqueous ammonia the reaction between carbonate and CO2 does not dominate, even when the loading exceeds 0.5 mol CO2 /mol NH3 . It is therefore difficult to predict the total heat of absorption accurately because of the complex reaction structure. The heats of reaction of CO2 in aqueous ammonia solutions are, however, very important because they are directly related to evaluation of the system. Calorimeters can be used to measure the enthalpy changes directly when a gas dissolves in a liquid solution. The enthalpy changes can also be estimated from vapor–liquid equilibrium (VLE) data using the following thermodynamic equation (Sherwood and Prausnitz, 1962):
∂ ln fˆi ∂(1/T )
= P,x
Hs R
partial pressure is always around 10% (Goppert and Maurer, 1988), which results in a big uncertainty in the heat of absorption. The heats of absorption calculated from Eq. (6) should be temperatureindependent because the plot of CO2 partial pressure versus 1/T is considered to be linear. Heats of absorption can also be estimated using a VLE model. Dardea et al. (2009) used the UNIQUAC model to simulate the whole chilled ammonia process for CO2 capture. The VLE model was fitted to solubility data, and therefore we cannot expect this simulation to give better results than those obtained using Eq. (6). To model the heats of absorption more accurately, experimental heats of absorption data need to be included with the VLE data when the model parameters are fitted. Direct calorimetric determination of the enthalpy of absorption is the best way to obtain accurate values. The enthalpy of absorption includes the heat of the physical absorption of CO2 in the solvent, and the heat of the chemical reaction between CO2 and the ammonia species. The experiments were carried out at constant temperature, so the effect of temperature on heat of absorption can be observed by carrying out several runs at different temperatures. In this work, the heats of absorption of CO2 in aqueous ammonia and in ammonium carbonate/carbamate solutions were directly measured using a reaction calorimeter with a reactor volume of 2 L (ChemiSens AB, Lund, Sweden). This is a large volume, so it is possible to feed CO2 into the reactor in batches, thereby obtaining the heat of absorption at different CO2 loadings at a given temperature. The loading spans were kept low (0.06–0.12 mol CO2 /mol NH3 ), so semi-differential measurements in loading could be obtained. Two solutions were tested: ∼2.5 wt% aqueous ammonia solution and ∼10 wt% ammonium carbonate/carbamate solution. The experiments were performed in the temperature range 35–80 ◦ C.
2. Experimental The chemicals used in this work—CO2 (AGA, ≥99.99% pure), aqueous ammonia (Merck, 25% minimum), ammonium carbonate (Sigma Aldrich, analytical purity)—were used without any further purification. Ammonia solutions were prepared by diluting 25 wt% aqueous ammonia, and the exact concentrations were found by titration with sulfuric acid after the solutions were made. The ammonium carbonate solutions were prepared by weight in distilled water. The scheme of the experimental set-up is shown in Fig. 1. The detailed experimental procedure used was identical to that described by (Kim and Svendsen, 2007) and is not repeated here. Up to 1 mol CO2 /mol NH3 in ammonia solution could be reached with the used CO2 storage cylinder pressure (3.5 bar). The total heat of absorption in every interval was the integral of heat flow with time. The total amount of CO2 added from the cylinder at each stage and the amount of CO2 left in the reactor were calculated from the pressure changes in the cylinder and the reactor using the Peng–Robinson equation of state. The heats of absorption were obtained from these data in units of kJ/mol CO2 .
(6)
When using Eq. (6), uncertainties in gas solubility data result in an increase in an uncertainty in the determination of enthalpy of about one order of magnitude. This means that with an accuracy in the solubility data of 2–3%, the accuracy of the enthalpy data would be 20–30% (Kim and Svendsen, 2007). For simplification at ambient pressures, CO2 partial pressures are always used instead of CO2 fugacity in Eq. (6). This increases the uncertainty in the estimated heats of absorption. Reviewing the NH3 –CO2 –H2 O ternary system, VLE data in the literature shows that the uncertainty in the
3. Uncertainties The main uncertainties in these experiments were, as also described by Kim and Svendsen (2007), the calculation of the amount of CO2 added from the storage cylinder, the amount of CO2 remaining in the gas phase of the reactor, and the heat flow integration curve. The molar heats of absorption were calculated from the ratio of the heat released in a loading interval (Qtotal ) to the amount of CO2 absorbed (nCO2 ). The uncertainty can be expressed as (Kim and
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Fig. 1. Experimental setup.
Svendsen, 2007):
ıHabs Habs
=
ıQtotal Qtotal
4.1. Aqueous ammonia
2
+
2
ınCO
2
nCO2
(7)
The CO2 added in each loading interval was calculated using the Peng–Robinson equation of state, but can also be calculated by integration of the mass flow controller reading. Based on the difference between these two methods, the estimated uncertainty was about 1.2%, as in Kim and Svendsen (2007). The main uncertainty in the heat flux was from determining the baseline in Fig. 2. It was assumed that the baseline changes linearly and that the curve was integrated by the trapezoidal method. The uncertainty in the total heat of absorption in every interval was estimated to be 2%, so the total uncertainty derived from Eq. (7) was about 2.3%. Another source of uncertainty is the ammonia partial pressure change with CO2 loading. To avoid this uncertainty, a low concentration of aqueous ammonia (about 2.5 wt%) was used. At lower loading levels, the total pressure in the reactor before and after CO2 addition dropped to about 0.002 bar, whereas the pressure change for pure CO2 in the storage cylinder was about 1–3 bar. Given the volume of the storage cylinder (about 4 L), and the gas volume in the reactor (about 0.8 L), the CO2 in the gas phase in the reactor could be disregarded. The added CO2 was assumed to be totally absorbed by the solution. At higher loading levels (greater than 0.5 mol CO2 /mol NH3 ), the ammonia partial pressures over the solution (about 1.1 mol NH3 /kg H2 O) are too low to be detected, even at high temperatures (80 ◦ C) according to Goppert and Maurer (1988). The total pressure change resulting from the ammonia partial pressure is thus very small, and compared with the large amount of CO2 added, the uncertainty caused by the ammonia partial pressure change with loading is negligible.
4. Results and discussion The experimental data for heats of absorption of CO2 into aqueous ammonia and ammonium carbonate/carbamate solutions at various loadings and in the temperature range 35–80 ◦ C are listed in Tables 1–7.
The possible reactions between aqueous ammonia and CO2 are shown in Eqs. (1)–(5). The first reaction is the most preferred because of the low reaction heat (enthalpy calculated under standard condition) (Kohl and Nielsen, 1997; Yeh et al., 2005). Many researchers simply assumed that aqueous ammonia reacts with CO2 to form ammonium carbonate. The first reaction takes place when the loading is over 0.5 mol CO2 /mol NH3 . The enthalpy change in Eq. (1) is the smallest, and thus this may seem certainly the preferred path to absorb and regenerate CO2 because it has the lowest Table 1 Heats of absorption of CO2 with ∼2.5 wt% aqueous ammonia at 35 ◦ C. ˛ (mol CO2 /mol NH3 )
Hdiff (kJ/mol CO2 )
35 ◦ C(1) 0.063 0.133 0.196 0.260 0.319 0.381 0.448 0.518 0.578 0.645 0.767 0.883 0.994 1.085
72.344 70.811 70.896 69.118 70.485 68.850 69.737 65.578 68.834 64.696 63.425 55.698 44.044 27.385
35 ◦ C(2) 0.055 0.109 0.166 0.223 0.278 0.334 0.390 0.446 0.501 0.554 0.664 0.765 0.863
69.367 66.398 67.755 70.789 66.612 64.881 63.105 62.950 57.195 62.105 53.259 53.951 49.860
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Fig. 2. An example of a heat flux curve for a single loading interval and a baseline for integration: — heat flow (blue); - - - CO2 flow; — baseline (purple). (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of the article.)
energy requirement. However, the temperature sensitivity in the CO2 partial pressure will be low according to Eq. (6), and in addition it seems impossible to control the main reaction in the absorption and desorption process so that it follows only Eq. (1). This has also been pointed out by Mathias in his chilled ammonia process evaluation (Mathiasa et al., 2009). We measured the differential enthalpies of absorption of CO2 in ∼2.5 wt% aqueous ammonia over the temperature range 35–80 ◦ C. The results are given in Figs. 2–5. The experimental results suggest that the first reaction does not occur even when the loading reaches 0.5 mol CO2 /mol NH3 . It seems that this reaction only takes place at much higher CO2 loadings. Figs. 3 and 4 show that, at low temperatures (35 and 40 ◦ C), the heat of absorption seems to be constant up to loadings of about 0.6 mol CO2 /mol NH3 , with the value remaining at about 70 kJ/mol Table 2 Heats of absorption of CO2 with ∼2.5 wt% aqueous ammonia at 40 ◦ C. ˛ (mol CO2 /mol NH3 )
Hdiff (kJ/mol CO2 )
40 ◦ C(1) 0.060 0.124 0.184 0.247 0.303 0.363 0.419 0.483 0.538 0.597 0.715 0.825 0.938
72.631 68.441 70.707 67.014 78.002 70.649 70.631 67.390 75.361 72.281 64.214 64.312 57.574
40 ◦ C(2) 0.055 0.109 0.162 0.214 0.266 0.318 0.371 0.424 0.480 0.522 0.571 0.677 0.771 0.860 0.944
64.285 67.103 65.606 67.838 69.913 67.441 68.860 73.015 70.824 77.900 72.793 67.434 63.318 58.431 46.289
CO2 . The heat of absorption goes down slowly when the loading is higher than 0.6 mol CO2 /mol NH3 . The enthalpies of absorption can be estimated using the speciation data of Mani et al. (2006), measured by NMR, as shown in Fig. 5. At the first point in Fig. 5, the estimated enthalpy using Eqs. (2)–(4) should be in the range 64–101 kJ/mol CO2 . In Fig. 5 the carbamate, carbonate, and bicarbonate fractions of total CO2 in the solution at the first point are about 43%, 34%, and 23%, respectively. Using this species distribution and the enthalpies in Eqs. (2)–(4) to estimate the heat of absorption of CO2 with aqueous ammonia gives a value of about 80 kJ/mol CO2 . The uncertainty in NMR species determination is about 10% (Holmes et al., 1998), so the results in this work are in reasonable agreement with the estimates obtained using Mani’s data. The reason for the high enthalpy of absorption is that, even with higher loadings than 0.5 mol CO2 /mol NH3, there is still a lot of free ammonia in the solution (Holmes et al., 1998), which means
Table 3 Heats of absorption of CO2 with ∼2.5 wt% aqueous ammonia at 60 ◦ C. ˛ (mol CO2 /mol NH3 )
Hdiff (kJ/mol CO2 )
60 ◦ C(1) 0.066 0.124 0.182 0.240 0.295 0.356 0.410 0.473 0.526 0.582 0.692 0.803 0.904
71.837 64.683 71.262 58.250 67.235 66.262 82.734 72.574 93.934 97.752 87.898 88.783 72.889
60 ◦ C(2) 0.058 0.117 0.174 0.234 0.291 0.351 0.409 0.466 0.523 0.581 0.688 0.795
81.851 65.572 60.348 60.628 66.017 76.442 81.273 83.812 86.058 96.914 88.043 88.421
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409
Table 4 Heats of absorption of CO2 with ∼2.5 wt% aqueous ammonia at 80 ◦ C.
Fig. 3. Differential enthalpies of absorption of CO2 in ∼2.5 wt% aqueous ammonia in the reactor at 35 ◦ C: (䊉) enthalpies of run 1; () enthalpies of run 2.
Fig. 4. Differential enthalpies of absorption of CO2 in ∼2.5 wt% aqueous ammonia in the reactor at 40 ◦ C: (䊉) enthalpies of run 1; () enthalpies of run 2.
˛ (mol CO2 /mol NH3 )
Hdiff (kJ/mol CO2 )
80 ◦ C(1) 0.052 0.110 0.167 0.223 0.279 0.333 0.385 0.439 0.493 0.544 0.641
65.386 46.490 43.859 57.217 74.804 87.464 112.467 114.080 113.309 122.090 134.507
80 ◦ C(2) 0.060 0.121 0.181 0.239 0.300 0.361 0.417 0.473 0.533 0.587 0.641 0.691 0.742 0.838 0.932
54.709 53.148 48.789 54.133 59.933 81.250 108.197 96.006 130.150 120.880 114.608 127.611 102.315 100.874 86.940
that the reaction between ammonium carbonate and CO2 is not the dominant one. Actually, free ammonia reacting with CO2 is one of the main sources of heat of absorption. If the loading approaches 1 mol CO2 /mol NH3 , the enthalpy of absorption decreases very quickly because carbonate/carbamate starts reacting with CO2 with a lower reaction enthalpy. Fig. 5 shows the percentages, but not the amounts, by which carbamate and carbonate concentration decrease with increased loading. These decreases occur because of the rapid increase in the amount of bicarbonate in the solution. This implies that carbamate/carbonate react with CO2 at the very high loading end, not at the lower or medium loading end of the absorption process. In Fig. 5, the loading of the last point is more than 1 mol CO2 /mol NH3 because the high pressure (3.5 bar) results in CO2 physically dissolving into the solution. The heat of absorption is seen to be about 27 kJ/mol CO2 for a loading span of Table 5 Heats of absorption of CO2 with 10 wt% ammonium carbonate solution at 35 ◦ C.
Fig. 5. Variations in carbamate, carbonate, and bicarbonate as a function of pH: (䊉) carbamate; () carbonate; () bicarbonate (Mani et al., 2006).
˛ (mol CO2 /mol (NH4 )2 CO3 )
Hdiff (kJ/mol CO2 )
35 ◦ C(1) 0.174 0.348 0.510 0.664 0.788
56.848 50.397 41.827 40.754 30.419
35 ◦ C(2) 0.081 0.247 0.557 0.701 0.845
69.403 58.996 45.039 35.348 25.784
35 ◦ C(3) 0.121 0.241 0.406 0.560 0.710 0.835
48.323 50.510 44.659 37.366 32.548 31.811
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Table 6 Heats of absorption of CO2 with 10 wt% ammonium carbonate solution at 40 ◦ C. ˛ (mol CO2 /mol (NH4 )2 CO3 )
Hdiff (kJ/mol CO2 )
40 ◦ C(1) 0.178 0.371 0.682 0.797
49.907 40.710 38.399 34.891
40 ◦ C(2) 0.078 0.234 0.388 0.573 0.763 0.920
55.871 53.309 51.324 43.962 31.608 11.879
40 ◦ C(3) 0.186 0.361 0.528 0.683 0.825
65.737 57.137 50.948 48.946 37.735
Table 7 Heats of absorption of CO2 with 10 wt% ammonium carbonate solution at 50–80 ◦ C. ˛ (mol CO2 /mol (NH4 )2 CO3 )
Hdiff (kJ/mol CO2 )
50 ◦ C 0.167 0.327 0.480 0.628 0.755
76.756 66.716 56.006 51.634 33.962
60 ◦ C 0.167 0.355 0.511
92.943 81.811 72.673
80 ◦ C 0.156 0.296 0.426 0.509
101.096 92.035 85.392 80.444
Fig. 7. Differential enthalpies of absorption of CO2 in ∼2.5 wt% aqueous ammonia in the reactor at 80 ◦ C: (䊉) enthalpies of run 1; () enthalpies of run 2.
slightly and reaches a minimum at a loading of around 0.2 mol CO2 /mol NH3 . The enthalpy then increases with increased loading. When the loading is around 0.6 mol CO2 /mol NH3 , the enthalpy of absorption of CO2 with aqueous ammonia reaches a maximum. The enthalpy of absorption then starts to decrease again, but remains high even at high loadings. The shapes of the enthalpy curves are different from those at lower temperatures, and they are also different from those of most amine curves at the same temperature. Similar shapes were, however, found in MAPA solutions by Kim (2009). A possible explanation for this may be changes in the activities of species in the solution with loading (Kim, 2009). More detailed studies are needed to explain this unusual behavior. The data for heats of absorption of CO2 in aqueous ammonia solutions are scattered, especially at high temperatures. 30 wt% MEA solutions were tested at 40 ◦ C for validation in this paper and compared with data from Kim and Svendsen (2007); they were in good agreement, as shown in Fig. 8. 4.2. Ammonium carbonate
0.994–1.084 mol CO2 /mol NH3 . Comparison with the physical heat of absorption of CO2 in water, around 20 kJ/mol CO2 (Austgen, 1989; Chen et al., 1979), confirms the reliability of the calorimetric data in this work. Figs. 6 and 7 show the enthalpy changes in the reactor at 60 and 80 ◦ C. At the beginning, the enthalpy of absorption decreases
To examine specifically the reaction given by Eq. (1), we tested 10 wt% ammonium carbonate solution (this weight percentage is comparable to 2.5 wt% aqueous ammonia at a loading of 0.5 mol CO2 /mol NH3 ) over the temperature range 35–80 ◦ C.
Fig. 6. Differential enthalpies of absorption of CO2 in ∼2.5 wt% aqueous ammonia in the reactor at 60 ◦ C: (䊉) enthalpies of run 1; () enthalpies of run 2.
Fig. 8. Calorimeter validation with 30 wt% MEA at 40 ◦ C: (䊉) data from Kim and Svendsen; () data in this work.
F. Qin et al. / International Journal of Greenhouse Gas Control 5 (2011) 405–412
Fig. 9. Differential enthalpies of absorption of CO2 in 10 wt% ammonium carbonate solution in the reactor at 35 ◦ C: (䊉) enthalpies of run 1; () enthalpies of run 2; () enthalpies of run 3.
The trends in enthalpies of absorption at different temperatures were similar to those found for aqueous ammonia, decreasing with increased loading (Figs. 9–11). The enthalpies increased with increasing temperature, as shown in Fig. 11 for 50, 60, and 80 ◦ C. It was noticed that at 35 and 40 ◦ C the initial enthalpy was 60–70 kJ/mol CO2 , much higher than 26.88 kJ/mol CO2 . Therefore, the reaction in Eq. (1) cannot be expected to dominate in neither ammonium carbonate solutions nor in aqueous ammonia. The reason is that when ammonium carbonate dissolves in water it can produce free ammonia (Wen and Brooker, 1995). The data reproducibility of heats of absorption of CO2 in the ammonium carbonate solutions was not very good. The low purity, and bicarbonate content, of the ammonium carbonate salts could contribute to the observed scatter (Mani et al., 2006). The calculations of the loadings in Figs. 9–11 were based on the weight balance and the molecular weight of ammonium carbonate and with the carbonate also containing bicarbonate, this may not be valid. If the data were reliable, and if the ammonium carbonate purity was high, then the enthalpies of absorption of CO2 with aqueous ammonia and with ammonium carbonate should be in very good agreement and theoretically should have the same loading
Fig. 10. Differential enthalpies of absorption of CO2 in 10 wt% ammonium carbonate solution in the reactor at 40 ◦ C: (䊉) enthalpies of run 1; () enthalpies of run 2; () enthalpies of run 3.
411
Fig. 11. Differential enthalpies of absorption of CO2 in 10 wt% ammonium carbonate solution in the reactor at 50, 60, and 80 ◦ C: (䊉) enthalpies at 50 ◦ C; () enthalpies at 60 ◦ C; () enthalpies at 80 ◦ C.
coordinate (mol CO2 /mol equivalent NH3 ). The loading in ammonium carbonate solution can be calculated from following equation, based on the formula of ammonium carbonate: loading (mol CO2 /mol equiv NH3 ) = 0.5 + loading (mol CO2 /mol (NH4 )2 CO3 ) × 0.5
(8)
Comparative results are plotted in Fig. 12. It can be seen to be a large deviation between the enthalpies of absorption of CO2 in aqueous ammonia (䊉) and in ammonium carbonate () solution. BaCl2 precipitation, which is used to determine the total inorganic CO2 in unloaded ammonium carbonate solution, found around 20% more CO2 in the ammonium carbonate solution than the calculated value based on the formula of ammonium carbonate. This is because commercially available ammonium carbonate is a mixture of ammonium bicarbonate and ammonium carbamate (Mani et al., 2006). This means that the initial loading in Eq. (8) was not 0.5 mol CO2 /mol NH3 , but around 0.6 mol CO2 /mol NH3 . The enthalpies of absorption of CO2 in aqueous ammonia (䊉) and in ammonium car-
Fig. 12. Comparison of differential enthalpies of absorption of CO2 with ∼2.5 wt% aqueous ammonia and 10 wt% ammonium carbonate solution at 40 ◦ C: (䊉) enthalpies of absorption of CO2 with ∼2.5 wt% aqueous ammonia; () enthalpies of absorption of CO2 with 10 wt% ammonium carbonate solution (loading calculation based on Eq. (8)); () enthalpies of absorption of CO2 with 10 wt% ammonium carbonate solution (loading calculation based on the total CO2 analysis).
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bonate () solution were in good agreement after changing the loading coordinate according to this. 5. Conclusions A reaction calorimeter was used to determine the enthalpies of absorption of CO2 in aqueous ammonia and in aqueous solutions of ammonium carbonate at temperatures of 35–80 ◦ C. The heat of absorption of CO2 with 2.5 wt% aqueous ammonia solution was found to be about 70 kJ/mol CO2 , which is lower than that with MEA (around 85 kJ/mol) at 35 and 40 ◦ C. The value slightly decreases with increased loading, but not to as low a value as expected by the carbonate–bicarbonate reaction (26.88 kJ/mol). The enthalpy of absorption of CO2 in aqueous ammonia at 60 and 80 ◦ C decreases at first, then increases with loadings between 0.2 mol CO2 /mol NH3 and 0.6 mol CO2 /mol NH3 , and then decreases again. The behavior of the heat of absorption of CO2 in 10 wt% ammonium carbonate solution was found to be the same as that for aqueous ammonia at loadings above 0.6 mol CO2 /mol NH3 . The heat of absorption increases with increasing temperature. The heats of absorption are directly related to the extent of the various reactions with CO2 and can be assessed from the species variation in the liquid phase. Future work will look at developing a heat of absorption model based on the NH3 –CO2 –H2 O ternary system VLE. Acknowledgments We greatly appreciate financial support from the National Natural Science Foundation of China (Project 50876051) and from the NTNU–Tsinghua University bilateral collaboration sponsored by the Norwegian Research Council (Project 175110/D15). References Austgen, D.M., 1989. A model of vapour–liquid equilibria for acid gas–alkanolamine–water system. Ph.D. Thesis. University of Texas at Austin, Austin, TX, USA. Black, S., 2006. Chilled ammonia scrubber for CO2 capture, MIT Carbon Sequestiation Forum VII. Cambridge, MA. Chen, C.C., Brit, H.I., Boston, J.F., Evans, L.B., 1979. Extension an dapplication of the Pitzer equation for vapour–liquid equilibrium of aqueous electrolyte systems with molecular solutes. AIChE Journal 25, 820–831. Ciferno, J.P., DiPietro, P., Tarka, T., 2005. An economic scoping study for CO2 capture using aqueous ammonia. NETL final report.
Dardea, V., Thomsen, K., van Well, W., Stenby, E., 2009. Chilled ammonia process for CO2 capture. International Journal of Greenhouse Gas Control 4 (2), 131– 136. Diao, Y.F., Zheng, X.Y., He, B.S., Chen, C.H., Xu, X.C., 2004. Experimental study on capturing CO2 greenhouse gas by ammonia scrubbing. Energy Conversion and Management 45, 2283–2296. Gao, W., Zhang, C.F., Zheng, Z.S., Qin, S.J., 1991. Hydrolysis kinetics of carbomide and its effect on carbonation process. Journal of East China Institute of Chemical Technology 17 (2), 139–146 (in Chinese). Goppert, U., Maurer, G., 1988. Vapor–liquid equilibria in aqueous solutions of ammonia and carbon dioxide at temperatures between 333 and 393 K and pressure up to 7 MPa. Fluid Phase Equilibria 41, 153–185. Holmes II, P.E., Naaz, M., Poling, B.E., 1998. Ion concentrations in the CO2 –NH3 –H2 O system from 13 C NMR spectroscopy. Industrial & Engineering Chemistry Research 37, 3281–3287. Kim I., 2009. Heat of reaction and VLE of post combustion CO2 absorbents. Ph.D. Thesis. Norwegian University of Science and Technology, Trondheim, Norway. Kim, I., Svendsen, H.F., 2007. Heat of absorption of carbon dioxide (CO2 ) in monoethanolamine (MEA) and 2-(aminoethyl)ethanolamine (AEEA) solutions. Industrial & Engineering Chemistry Research 46, 5803–5809. Kim, Y.J., You, J.K., Hong, W.H., et al., 2008. Characteristics of CO2 absorption into aqueous ammonia. Separation Science and Technology 43, 766–777. Kohl, A.L., Nielsen, R.B., 1997. Gas Purification. Gulf Publishing Company, Houston, Texas, USA. Liu, F., Wang, S.J., Chen, C.H., Xu, X.C., 2009. Re search progress of CO2 capture by using ammonia from flue gas of power plant. CIESC Journal 60 (2), 269– 278. Mani, F., Peruzzini, M., Stoppioni, P., 2006. CO2 absorption by aqueous NH3 solutions: Speciation of ammonium carbamate, bicarbonate and carbonate by a 13C NMR study. Green Chemistry 8, 995–1000. Mathiasa, P.M., Reddy, S., O’Connell, J.P., 2009. Quantitative evaluation of the chilled-ammonia process for CO2 capture using thermodynamic analysis and process simulation. International Journal of Greenhouse Gas Control 4 (2), 174– 179. Resnik, K.P., Garber, W., Hreha, D.C., Yeh, J.T., Pennline, H.W., 2006. A parametric scan for regenerative ammonia-based scrubbing for the capture of CO2 . In: Proceedings of the 23rd Annual International Pittsburgh Coal Conference, Pittsburgh, PA. Rhudy, D., Gal, E., Robert, H., 2006. CO2 testing program. An industry/EPRI initiative to develop CO2 capture and storage test capabilities. Presentation at the CO2 capture test network meeting. Sherwood, A.E., Prausnitz, J.M., 1962. The heat of solution of gases at high pressure. AIChE Journal 8 (4), 519. Wen, N.P., Brooker, M.H., 1995. Ammonium carbonate, ammonium bicarbonate, and ammonium carbamate equilibria: a Raman study. Journal of Physical Chemistry 99, 359–368. Wolf, M., Braustch, A., Gernert, J., Kaefer, G., Pfeffer, A., Winkler, D., 2006. Outlook on post combustion CO2 capture in ALSTOM gas turbine plants. Paper presented at PowerGen Europe, May 30–June 1, Cologne, Germany. Yeh, A.C., Bai, H., 1999. Comparison of ammonia and monoethanolamine solvents to reduce CO2 greenhouse gas emissions. The Science of the Total Environment 228, 121–133. Yeh, J.T., Resnik, K.P., Rygle, K., Pennline, H.W., 2005. Semi-batch absorption and regeneration studies for CO2 capture by aqueous ammonia. Fuel Processing Technology 86, 1533–1546.
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International Journal of Greenhouse Gas Control journal homepage: www.elsevier.com/locate/ijggc
Heat of absorption of CO2 in aqueous ammonia and ammonium carbonate/carbamate solutions Feng Qin a , Shujuan Wang a , Inna Kim b , Hallvard F. Svendsen b,∗ , Changhe Chen a a b
Department of Thermal Engineering, Key Laboratory for Thermal Science and Power Engineering of Ministry of Education, Tsinghua University, 100084 Beijing, China Department of Chemical Engineering, Norwegian University of Science and Technology, 7491 Trondheim, Norway
a r t i c l e
i n f o
Article history: Received 13 August 2009 Received in revised form 15 March 2010 Accepted 7 April 2010 Available online 4 May 2010 Keywords: Aqueous ammonia CO2 Heat of absorption Calorimeter
a b s t r a c t A reaction calorimeter was used to determine the enthalpies of absorption of CO2 in aqueous ammonia and in aqueous solutions of ammonium carbonate at temperatures of 35–80 ◦ C. The heat of absorption of CO2 with 2.5 wt% aqueous ammonia solution was found to be about 70 kJ/mol CO2 , which is lower than that with MEA (around 85 kJ/mol) at 35 and 40 ◦ C. The value decreases with increased loading, but not to as low a value as expected by the carbonate–bicarbonate reaction (26.88 kJ/mol). The enthalpy of absorption of CO2 in aqueous ammonia at 60 and 80 ◦ C decreases with loadings at first, then increases between 0.2 mol CO2 /mol NH3 and 0.6 mol CO2 /mol NH3 , and then decreases again. The behavior of the heat of absorption of CO2 in 10 wt% ammonium carbonate solution was found to be the same as that of aqueous ammonia at loadings above 0.6 mol CO2 /mol NH3 . The heat of absorption increases with increasing temperature. The heats of absorption are directly related to the extent of the various reactions with CO2 and can be assessed from the species variation in the liquid phase. © 2010 Elsevier Ltd. All rights reserved.
1. Introduction Chemical absorption is a promising method for CO2 capture and is also one of the most common. It is necessary, however, to reduce the high capital costs of the method and the energy requirements in the absorption and desorption processes. Power plant efficiency decreases by about 10 percentage points when the flue gas is treated for CO2 capture by chemical absorption (Rhudy et al., 2006; Wolf et al., 2006), so developing a novel solvent with low regeneration energy is significant for reducing energy penalties. Aqueous ammonia solutions are thought to have a potential as effective and economic absorbents for CO2 capture. The reactions between aqueous ammonia solutions and CO2 are very fast, and in batch or semi-batch experiments the CO2 removal efficiency can reach 90% or more (Diao et al., 2004; Kim et al., 2008; Resnik et al., 2006; Yeh et al., 2005; Yeh and Bai, 1999). This means that aqueous ammonia is very effective for CO2 absorption. Previous research has shown that the loading capacity of aqueous ammonia can be high (Yeh et al., 2005; Yeh and Bai, 1999), and much higher than that of monoethanolamine (MEA). The corrosion problems caused by aqueous ammonia are no worse than those caused by amines. Also, with aqueous ammonia there are no thermal degradation and oxidative degradation problems like those arising with amine systems, and the system can scrub SOx ,
∗ Corresponding author. Tel.: +47 73594100; fax: +47 73594080. E-mail address: [email protected] (H.F. Svendsen). 1750-5836/$ – see front matter © 2010 Elsevier Ltd. All rights reserved. doi:10.1016/j.ijggc.2010.04.005
NOx and other acidic gases simultaneously. Furthermore, ammonia is cheaper than MEA (Yeh et al., 2005). Most of the literature focuses on parametric analyses of the capture process. Liu et al. (2009) reviewed the progress of CO2 capture using aqueous ammonia, and summarized the effects of reactor type, reaction temperature, ammonia concentration, CO2 concentration, and SOx , NOx and heavy metals on the CO2 removal efficiency. Data by Yeh et al. (2005) indicate that the energy requirement for CO2 regeneration using the aqueous ammonia process could be at least 75% less than that of using MEA. The semi-continuous experiments by Yeh and Bai (1999) also indicate that the heat of absorption in aqueous ammonia is much lower than it is in MEA, and monitored temperature variations show that endothermic reactions play a role at the beginning of the absorption. It has been suggested in a number of papers that the regeneration heat requirement can be much smaller than in MEA-based processes such as the NETL and ALSTOM processes, see Black (2006), Ciferno et al. (2005), Resnik et al. (2006), Wolf et al. (2006), and Yeh et al. (2005). In all these studies on aqueous ammonia based CO2 capture system, the heats of reaction obtained or presumed were about 26–27 kJ/mol CO2 . All these evaluations, however, are based on the assumption that carbonate reacts with CO2 to form bicarbonate. It is thus very important to validate these assumptions. 2NH4 HCO3 (aq) ↔ (NH4 )2 CO3 (aq) + CO2 (g) + H2 O Hrx = 26.88 kJ/mol
(1)
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Nomenclature Symbols and constants f fugacity (Pa) Habs (Hs , Hdiff ) heat of absorption (kJ/mol CO2 ) CO2 absorbed (mol) nCO2 Qtotal heat (J) R universal gas constant (8.314 J/mol/K) T temperature (K)] Greek symbol ˛ CO2 loading (mol CO2 /mol NH3 )
NH4 HCO3 (aq) ↔ NH3 (aq) + CO2 (g) + H2 O Hrx = 64.26 kJ/mol
(2)
(NH4 )2 CO3 (aq) ↔ 2NH3 (aq) + CO2 (g) + H2 O Hrx = 101.22 kJ/mol
(3)
NH2 COONH4 (aq) ↔ 2NH3 (aq) + CO2 (g) Hrx = 72.32 kJ/mol
(4)
NH2 COONH4 (aq) + H2 O ↔ NH4 HCO3 (aq) + NH3 (aq) Hrx = 8.06 kJ/mol
(5)
When CO2 is fed into aqueous ammonia solutions, a mixture of ammonium carbonate, bicarbonate, and carbamate is produced at lower loadings. The carbamate and carbonate can further react with CO2 to produce bicarbonate, as shown in Eqs. (1)–(5) (Kohl and Nielsen, 1997; Yeh et al., 2005). Even in aqueous ammonia solutions at higher loadings (above 0.5 mol CO2 /mol NH3 ), experimental analyses and NMR results have shown that there is still a lot of free ammonia in solution (Gao et al., 1991; Holmes et al., 1998; Mani et al., 2006; Wen and Brooker, 1995). This suggests that in aqueous ammonia the reaction between carbonate and CO2 does not dominate, even when the loading exceeds 0.5 mol CO2 /mol NH3 . It is therefore difficult to predict the total heat of absorption accurately because of the complex reaction structure. The heats of reaction of CO2 in aqueous ammonia solutions are, however, very important because they are directly related to evaluation of the system. Calorimeters can be used to measure the enthalpy changes directly when a gas dissolves in a liquid solution. The enthalpy changes can also be estimated from vapor–liquid equilibrium (VLE) data using the following thermodynamic equation (Sherwood and Prausnitz, 1962):
∂ ln fˆi ∂(1/T )
= P,x
Hs R
partial pressure is always around 10% (Goppert and Maurer, 1988), which results in a big uncertainty in the heat of absorption. The heats of absorption calculated from Eq. (6) should be temperatureindependent because the plot of CO2 partial pressure versus 1/T is considered to be linear. Heats of absorption can also be estimated using a VLE model. Dardea et al. (2009) used the UNIQUAC model to simulate the whole chilled ammonia process for CO2 capture. The VLE model was fitted to solubility data, and therefore we cannot expect this simulation to give better results than those obtained using Eq. (6). To model the heats of absorption more accurately, experimental heats of absorption data need to be included with the VLE data when the model parameters are fitted. Direct calorimetric determination of the enthalpy of absorption is the best way to obtain accurate values. The enthalpy of absorption includes the heat of the physical absorption of CO2 in the solvent, and the heat of the chemical reaction between CO2 and the ammonia species. The experiments were carried out at constant temperature, so the effect of temperature on heat of absorption can be observed by carrying out several runs at different temperatures. In this work, the heats of absorption of CO2 in aqueous ammonia and in ammonium carbonate/carbamate solutions were directly measured using a reaction calorimeter with a reactor volume of 2 L (ChemiSens AB, Lund, Sweden). This is a large volume, so it is possible to feed CO2 into the reactor in batches, thereby obtaining the heat of absorption at different CO2 loadings at a given temperature. The loading spans were kept low (0.06–0.12 mol CO2 /mol NH3 ), so semi-differential measurements in loading could be obtained. Two solutions were tested: ∼2.5 wt% aqueous ammonia solution and ∼10 wt% ammonium carbonate/carbamate solution. The experiments were performed in the temperature range 35–80 ◦ C.
2. Experimental The chemicals used in this work—CO2 (AGA, ≥99.99% pure), aqueous ammonia (Merck, 25% minimum), ammonium carbonate (Sigma Aldrich, analytical purity)—were used without any further purification. Ammonia solutions were prepared by diluting 25 wt% aqueous ammonia, and the exact concentrations were found by titration with sulfuric acid after the solutions were made. The ammonium carbonate solutions were prepared by weight in distilled water. The scheme of the experimental set-up is shown in Fig. 1. The detailed experimental procedure used was identical to that described by (Kim and Svendsen, 2007) and is not repeated here. Up to 1 mol CO2 /mol NH3 in ammonia solution could be reached with the used CO2 storage cylinder pressure (3.5 bar). The total heat of absorption in every interval was the integral of heat flow with time. The total amount of CO2 added from the cylinder at each stage and the amount of CO2 left in the reactor were calculated from the pressure changes in the cylinder and the reactor using the Peng–Robinson equation of state. The heats of absorption were obtained from these data in units of kJ/mol CO2 .
(6)
When using Eq. (6), uncertainties in gas solubility data result in an increase in an uncertainty in the determination of enthalpy of about one order of magnitude. This means that with an accuracy in the solubility data of 2–3%, the accuracy of the enthalpy data would be 20–30% (Kim and Svendsen, 2007). For simplification at ambient pressures, CO2 partial pressures are always used instead of CO2 fugacity in Eq. (6). This increases the uncertainty in the estimated heats of absorption. Reviewing the NH3 –CO2 –H2 O ternary system, VLE data in the literature shows that the uncertainty in the
3. Uncertainties The main uncertainties in these experiments were, as also described by Kim and Svendsen (2007), the calculation of the amount of CO2 added from the storage cylinder, the amount of CO2 remaining in the gas phase of the reactor, and the heat flow integration curve. The molar heats of absorption were calculated from the ratio of the heat released in a loading interval (Qtotal ) to the amount of CO2 absorbed (nCO2 ). The uncertainty can be expressed as (Kim and
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407
Fig. 1. Experimental setup.
Svendsen, 2007):
ıHabs Habs
=
ıQtotal Qtotal
4.1. Aqueous ammonia
2
+
2
ınCO
2
nCO2
(7)
The CO2 added in each loading interval was calculated using the Peng–Robinson equation of state, but can also be calculated by integration of the mass flow controller reading. Based on the difference between these two methods, the estimated uncertainty was about 1.2%, as in Kim and Svendsen (2007). The main uncertainty in the heat flux was from determining the baseline in Fig. 2. It was assumed that the baseline changes linearly and that the curve was integrated by the trapezoidal method. The uncertainty in the total heat of absorption in every interval was estimated to be 2%, so the total uncertainty derived from Eq. (7) was about 2.3%. Another source of uncertainty is the ammonia partial pressure change with CO2 loading. To avoid this uncertainty, a low concentration of aqueous ammonia (about 2.5 wt%) was used. At lower loading levels, the total pressure in the reactor before and after CO2 addition dropped to about 0.002 bar, whereas the pressure change for pure CO2 in the storage cylinder was about 1–3 bar. Given the volume of the storage cylinder (about 4 L), and the gas volume in the reactor (about 0.8 L), the CO2 in the gas phase in the reactor could be disregarded. The added CO2 was assumed to be totally absorbed by the solution. At higher loading levels (greater than 0.5 mol CO2 /mol NH3 ), the ammonia partial pressures over the solution (about 1.1 mol NH3 /kg H2 O) are too low to be detected, even at high temperatures (80 ◦ C) according to Goppert and Maurer (1988). The total pressure change resulting from the ammonia partial pressure is thus very small, and compared with the large amount of CO2 added, the uncertainty caused by the ammonia partial pressure change with loading is negligible.
4. Results and discussion The experimental data for heats of absorption of CO2 into aqueous ammonia and ammonium carbonate/carbamate solutions at various loadings and in the temperature range 35–80 ◦ C are listed in Tables 1–7.
The possible reactions between aqueous ammonia and CO2 are shown in Eqs. (1)–(5). The first reaction is the most preferred because of the low reaction heat (enthalpy calculated under standard condition) (Kohl and Nielsen, 1997; Yeh et al., 2005). Many researchers simply assumed that aqueous ammonia reacts with CO2 to form ammonium carbonate. The first reaction takes place when the loading is over 0.5 mol CO2 /mol NH3 . The enthalpy change in Eq. (1) is the smallest, and thus this may seem certainly the preferred path to absorb and regenerate CO2 because it has the lowest Table 1 Heats of absorption of CO2 with ∼2.5 wt% aqueous ammonia at 35 ◦ C. ˛ (mol CO2 /mol NH3 )
Hdiff (kJ/mol CO2 )
35 ◦ C(1) 0.063 0.133 0.196 0.260 0.319 0.381 0.448 0.518 0.578 0.645 0.767 0.883 0.994 1.085
72.344 70.811 70.896 69.118 70.485 68.850 69.737 65.578 68.834 64.696 63.425 55.698 44.044 27.385
35 ◦ C(2) 0.055 0.109 0.166 0.223 0.278 0.334 0.390 0.446 0.501 0.554 0.664 0.765 0.863
69.367 66.398 67.755 70.789 66.612 64.881 63.105 62.950 57.195 62.105 53.259 53.951 49.860
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Fig. 2. An example of a heat flux curve for a single loading interval and a baseline for integration: — heat flow (blue); - - - CO2 flow; — baseline (purple). (For interpretation of the references to colour in this figure legend, the reader is referred to the web version of the article.)
energy requirement. However, the temperature sensitivity in the CO2 partial pressure will be low according to Eq. (6), and in addition it seems impossible to control the main reaction in the absorption and desorption process so that it follows only Eq. (1). This has also been pointed out by Mathias in his chilled ammonia process evaluation (Mathiasa et al., 2009). We measured the differential enthalpies of absorption of CO2 in ∼2.5 wt% aqueous ammonia over the temperature range 35–80 ◦ C. The results are given in Figs. 2–5. The experimental results suggest that the first reaction does not occur even when the loading reaches 0.5 mol CO2 /mol NH3 . It seems that this reaction only takes place at much higher CO2 loadings. Figs. 3 and 4 show that, at low temperatures (35 and 40 ◦ C), the heat of absorption seems to be constant up to loadings of about 0.6 mol CO2 /mol NH3 , with the value remaining at about 70 kJ/mol Table 2 Heats of absorption of CO2 with ∼2.5 wt% aqueous ammonia at 40 ◦ C. ˛ (mol CO2 /mol NH3 )
Hdiff (kJ/mol CO2 )
40 ◦ C(1) 0.060 0.124 0.184 0.247 0.303 0.363 0.419 0.483 0.538 0.597 0.715 0.825 0.938
72.631 68.441 70.707 67.014 78.002 70.649 70.631 67.390 75.361 72.281 64.214 64.312 57.574
40 ◦ C(2) 0.055 0.109 0.162 0.214 0.266 0.318 0.371 0.424 0.480 0.522 0.571 0.677 0.771 0.860 0.944
64.285 67.103 65.606 67.838 69.913 67.441 68.860 73.015 70.824 77.900 72.793 67.434 63.318 58.431 46.289
CO2 . The heat of absorption goes down slowly when the loading is higher than 0.6 mol CO2 /mol NH3 . The enthalpies of absorption can be estimated using the speciation data of Mani et al. (2006), measured by NMR, as shown in Fig. 5. At the first point in Fig. 5, the estimated enthalpy using Eqs. (2)–(4) should be in the range 64–101 kJ/mol CO2 . In Fig. 5 the carbamate, carbonate, and bicarbonate fractions of total CO2 in the solution at the first point are about 43%, 34%, and 23%, respectively. Using this species distribution and the enthalpies in Eqs. (2)–(4) to estimate the heat of absorption of CO2 with aqueous ammonia gives a value of about 80 kJ/mol CO2 . The uncertainty in NMR species determination is about 10% (Holmes et al., 1998), so the results in this work are in reasonable agreement with the estimates obtained using Mani’s data. The reason for the high enthalpy of absorption is that, even with higher loadings than 0.5 mol CO2 /mol NH3, there is still a lot of free ammonia in the solution (Holmes et al., 1998), which means
Table 3 Heats of absorption of CO2 with ∼2.5 wt% aqueous ammonia at 60 ◦ C. ˛ (mol CO2 /mol NH3 )
Hdiff (kJ/mol CO2 )
60 ◦ C(1) 0.066 0.124 0.182 0.240 0.295 0.356 0.410 0.473 0.526 0.582 0.692 0.803 0.904
71.837 64.683 71.262 58.250 67.235 66.262 82.734 72.574 93.934 97.752 87.898 88.783 72.889
60 ◦ C(2) 0.058 0.117 0.174 0.234 0.291 0.351 0.409 0.466 0.523 0.581 0.688 0.795
81.851 65.572 60.348 60.628 66.017 76.442 81.273 83.812 86.058 96.914 88.043 88.421
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409
Table 4 Heats of absorption of CO2 with ∼2.5 wt% aqueous ammonia at 80 ◦ C.
Fig. 3. Differential enthalpies of absorption of CO2 in ∼2.5 wt% aqueous ammonia in the reactor at 35 ◦ C: (䊉) enthalpies of run 1; () enthalpies of run 2.
Fig. 4. Differential enthalpies of absorption of CO2 in ∼2.5 wt% aqueous ammonia in the reactor at 40 ◦ C: (䊉) enthalpies of run 1; () enthalpies of run 2.
˛ (mol CO2 /mol NH3 )
Hdiff (kJ/mol CO2 )
80 ◦ C(1) 0.052 0.110 0.167 0.223 0.279 0.333 0.385 0.439 0.493 0.544 0.641
65.386 46.490 43.859 57.217 74.804 87.464 112.467 114.080 113.309 122.090 134.507
80 ◦ C(2) 0.060 0.121 0.181 0.239 0.300 0.361 0.417 0.473 0.533 0.587 0.641 0.691 0.742 0.838 0.932
54.709 53.148 48.789 54.133 59.933 81.250 108.197 96.006 130.150 120.880 114.608 127.611 102.315 100.874 86.940
that the reaction between ammonium carbonate and CO2 is not the dominant one. Actually, free ammonia reacting with CO2 is one of the main sources of heat of absorption. If the loading approaches 1 mol CO2 /mol NH3 , the enthalpy of absorption decreases very quickly because carbonate/carbamate starts reacting with CO2 with a lower reaction enthalpy. Fig. 5 shows the percentages, but not the amounts, by which carbamate and carbonate concentration decrease with increased loading. These decreases occur because of the rapid increase in the amount of bicarbonate in the solution. This implies that carbamate/carbonate react with CO2 at the very high loading end, not at the lower or medium loading end of the absorption process. In Fig. 5, the loading of the last point is more than 1 mol CO2 /mol NH3 because the high pressure (3.5 bar) results in CO2 physically dissolving into the solution. The heat of absorption is seen to be about 27 kJ/mol CO2 for a loading span of Table 5 Heats of absorption of CO2 with 10 wt% ammonium carbonate solution at 35 ◦ C.
Fig. 5. Variations in carbamate, carbonate, and bicarbonate as a function of pH: (䊉) carbamate; () carbonate; () bicarbonate (Mani et al., 2006).
˛ (mol CO2 /mol (NH4 )2 CO3 )
Hdiff (kJ/mol CO2 )
35 ◦ C(1) 0.174 0.348 0.510 0.664 0.788
56.848 50.397 41.827 40.754 30.419
35 ◦ C(2) 0.081 0.247 0.557 0.701 0.845
69.403 58.996 45.039 35.348 25.784
35 ◦ C(3) 0.121 0.241 0.406 0.560 0.710 0.835
48.323 50.510 44.659 37.366 32.548 31.811
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Table 6 Heats of absorption of CO2 with 10 wt% ammonium carbonate solution at 40 ◦ C. ˛ (mol CO2 /mol (NH4 )2 CO3 )
Hdiff (kJ/mol CO2 )
40 ◦ C(1) 0.178 0.371 0.682 0.797
49.907 40.710 38.399 34.891
40 ◦ C(2) 0.078 0.234 0.388 0.573 0.763 0.920
55.871 53.309 51.324 43.962 31.608 11.879
40 ◦ C(3) 0.186 0.361 0.528 0.683 0.825
65.737 57.137 50.948 48.946 37.735
Table 7 Heats of absorption of CO2 with 10 wt% ammonium carbonate solution at 50–80 ◦ C. ˛ (mol CO2 /mol (NH4 )2 CO3 )
Hdiff (kJ/mol CO2 )
50 ◦ C 0.167 0.327 0.480 0.628 0.755
76.756 66.716 56.006 51.634 33.962
60 ◦ C 0.167 0.355 0.511
92.943 81.811 72.673
80 ◦ C 0.156 0.296 0.426 0.509
101.096 92.035 85.392 80.444
Fig. 7. Differential enthalpies of absorption of CO2 in ∼2.5 wt% aqueous ammonia in the reactor at 80 ◦ C: (䊉) enthalpies of run 1; () enthalpies of run 2.
slightly and reaches a minimum at a loading of around 0.2 mol CO2 /mol NH3 . The enthalpy then increases with increased loading. When the loading is around 0.6 mol CO2 /mol NH3 , the enthalpy of absorption of CO2 with aqueous ammonia reaches a maximum. The enthalpy of absorption then starts to decrease again, but remains high even at high loadings. The shapes of the enthalpy curves are different from those at lower temperatures, and they are also different from those of most amine curves at the same temperature. Similar shapes were, however, found in MAPA solutions by Kim (2009). A possible explanation for this may be changes in the activities of species in the solution with loading (Kim, 2009). More detailed studies are needed to explain this unusual behavior. The data for heats of absorption of CO2 in aqueous ammonia solutions are scattered, especially at high temperatures. 30 wt% MEA solutions were tested at 40 ◦ C for validation in this paper and compared with data from Kim and Svendsen (2007); they were in good agreement, as shown in Fig. 8. 4.2. Ammonium carbonate
0.994–1.084 mol CO2 /mol NH3 . Comparison with the physical heat of absorption of CO2 in water, around 20 kJ/mol CO2 (Austgen, 1989; Chen et al., 1979), confirms the reliability of the calorimetric data in this work. Figs. 6 and 7 show the enthalpy changes in the reactor at 60 and 80 ◦ C. At the beginning, the enthalpy of absorption decreases
To examine specifically the reaction given by Eq. (1), we tested 10 wt% ammonium carbonate solution (this weight percentage is comparable to 2.5 wt% aqueous ammonia at a loading of 0.5 mol CO2 /mol NH3 ) over the temperature range 35–80 ◦ C.
Fig. 6. Differential enthalpies of absorption of CO2 in ∼2.5 wt% aqueous ammonia in the reactor at 60 ◦ C: (䊉) enthalpies of run 1; () enthalpies of run 2.
Fig. 8. Calorimeter validation with 30 wt% MEA at 40 ◦ C: (䊉) data from Kim and Svendsen; () data in this work.
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Fig. 9. Differential enthalpies of absorption of CO2 in 10 wt% ammonium carbonate solution in the reactor at 35 ◦ C: (䊉) enthalpies of run 1; () enthalpies of run 2; () enthalpies of run 3.
The trends in enthalpies of absorption at different temperatures were similar to those found for aqueous ammonia, decreasing with increased loading (Figs. 9–11). The enthalpies increased with increasing temperature, as shown in Fig. 11 for 50, 60, and 80 ◦ C. It was noticed that at 35 and 40 ◦ C the initial enthalpy was 60–70 kJ/mol CO2 , much higher than 26.88 kJ/mol CO2 . Therefore, the reaction in Eq. (1) cannot be expected to dominate in neither ammonium carbonate solutions nor in aqueous ammonia. The reason is that when ammonium carbonate dissolves in water it can produce free ammonia (Wen and Brooker, 1995). The data reproducibility of heats of absorption of CO2 in the ammonium carbonate solutions was not very good. The low purity, and bicarbonate content, of the ammonium carbonate salts could contribute to the observed scatter (Mani et al., 2006). The calculations of the loadings in Figs. 9–11 were based on the weight balance and the molecular weight of ammonium carbonate and with the carbonate also containing bicarbonate, this may not be valid. If the data were reliable, and if the ammonium carbonate purity was high, then the enthalpies of absorption of CO2 with aqueous ammonia and with ammonium carbonate should be in very good agreement and theoretically should have the same loading
Fig. 10. Differential enthalpies of absorption of CO2 in 10 wt% ammonium carbonate solution in the reactor at 40 ◦ C: (䊉) enthalpies of run 1; () enthalpies of run 2; () enthalpies of run 3.
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Fig. 11. Differential enthalpies of absorption of CO2 in 10 wt% ammonium carbonate solution in the reactor at 50, 60, and 80 ◦ C: (䊉) enthalpies at 50 ◦ C; () enthalpies at 60 ◦ C; () enthalpies at 80 ◦ C.
coordinate (mol CO2 /mol equivalent NH3 ). The loading in ammonium carbonate solution can be calculated from following equation, based on the formula of ammonium carbonate: loading (mol CO2 /mol equiv NH3 ) = 0.5 + loading (mol CO2 /mol (NH4 )2 CO3 ) × 0.5
(8)
Comparative results are plotted in Fig. 12. It can be seen to be a large deviation between the enthalpies of absorption of CO2 in aqueous ammonia (䊉) and in ammonium carbonate () solution. BaCl2 precipitation, which is used to determine the total inorganic CO2 in unloaded ammonium carbonate solution, found around 20% more CO2 in the ammonium carbonate solution than the calculated value based on the formula of ammonium carbonate. This is because commercially available ammonium carbonate is a mixture of ammonium bicarbonate and ammonium carbamate (Mani et al., 2006). This means that the initial loading in Eq. (8) was not 0.5 mol CO2 /mol NH3 , but around 0.6 mol CO2 /mol NH3 . The enthalpies of absorption of CO2 in aqueous ammonia (䊉) and in ammonium car-
Fig. 12. Comparison of differential enthalpies of absorption of CO2 with ∼2.5 wt% aqueous ammonia and 10 wt% ammonium carbonate solution at 40 ◦ C: (䊉) enthalpies of absorption of CO2 with ∼2.5 wt% aqueous ammonia; () enthalpies of absorption of CO2 with 10 wt% ammonium carbonate solution (loading calculation based on Eq. (8)); () enthalpies of absorption of CO2 with 10 wt% ammonium carbonate solution (loading calculation based on the total CO2 analysis).
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bonate () solution were in good agreement after changing the loading coordinate according to this. 5. Conclusions A reaction calorimeter was used to determine the enthalpies of absorption of CO2 in aqueous ammonia and in aqueous solutions of ammonium carbonate at temperatures of 35–80 ◦ C. The heat of absorption of CO2 with 2.5 wt% aqueous ammonia solution was found to be about 70 kJ/mol CO2 , which is lower than that with MEA (around 85 kJ/mol) at 35 and 40 ◦ C. The value slightly decreases with increased loading, but not to as low a value as expected by the carbonate–bicarbonate reaction (26.88 kJ/mol). The enthalpy of absorption of CO2 in aqueous ammonia at 60 and 80 ◦ C decreases at first, then increases with loadings between 0.2 mol CO2 /mol NH3 and 0.6 mol CO2 /mol NH3 , and then decreases again. The behavior of the heat of absorption of CO2 in 10 wt% ammonium carbonate solution was found to be the same as that for aqueous ammonia at loadings above 0.6 mol CO2 /mol NH3 . The heat of absorption increases with increasing temperature. The heats of absorption are directly related to the extent of the various reactions with CO2 and can be assessed from the species variation in the liquid phase. Future work will look at developing a heat of absorption model based on the NH3 –CO2 –H2 O ternary system VLE. Acknowledgments We greatly appreciate financial support from the National Natural Science Foundation of China (Project 50876051) and from the NTNU–Tsinghua University bilateral collaboration sponsored by the Norwegian Research Council (Project 175110/D15). References Austgen, D.M., 1989. A model of vapour–liquid equilibria for acid gas–alkanolamine–water system. Ph.D. Thesis. University of Texas at Austin, Austin, TX, USA. Black, S., 2006. Chilled ammonia scrubber for CO2 capture, MIT Carbon Sequestiation Forum VII. Cambridge, MA. Chen, C.C., Brit, H.I., Boston, J.F., Evans, L.B., 1979. Extension an dapplication of the Pitzer equation for vapour–liquid equilibrium of aqueous electrolyte systems with molecular solutes. AIChE Journal 25, 820–831. Ciferno, J.P., DiPietro, P., Tarka, T., 2005. An economic scoping study for CO2 capture using aqueous ammonia. NETL final report.
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