ARTICLE pubs.acs.org/est
Moisture Swing Sorbent for Carbon Dioxide Capture from Ambient Air Tao Wang,*,† Klaus S. Lackner,† and Allen Wright† †
Earth Institute and Department of Earth and Environmental Engineering, Columbia University, 1038 Mudd Building, 500 West 120th Street, New York, NY 10027
bS Supporting Information ABSTRACT: An amine-based anion exchange resin dispersed in a flat sheet of polypropylene was prepared in alkaline forms so that it would capture carbon dioxide from air. The resin, with quaternary ammonium cations attached to the polymer structure and hydroxide or carbonate groups as mobile counterions, absorbs carbon dioxide when dry and releases it when wet. In ambient air, the moist resin dries spontaneously and subsequently absorbs carbon dioxide. This constitutes a moisture induced cycle, which stands in contrast to thermal pressure swing based cycles. This paper aims to determine the isothermal performance of the sorbent during such a moisture swing. Equilibrium experiments show that the absorption and desorption process can be described well by a Langmuir isothermal model. The equilibrium partial pressure of carbon dioxide over the resin at a given loading state can be increased by 2 orders of magnitude by wetting the resin.
’ INTRODUCTION Carbon dioxide sorbents, whose regeneration can be triggered by moisture, offer a new approach to CO2 separation from dilute streams including air. Such a sorbent has recently been proposed for CO2 capture from ambient air (air capture), thereby compensating for emissions from small and distributed sources.1 Current carbon capture and storage (CCS) techniques focus on capture from large point sources. According to the IPCC report,2 about 60% of global CO2 emissions from fossil-fuels are attributed to large stationary sources. Assuming 90% capture efficiency and 90% coverage of all sources, about 50% of global emissions would still be released into the atmosphere. This is far too much to allow for the stabilization of the atmospheric concentration of CO2 and insufficient to constrain the growth of atmospheric carbon dioxide concentrations as the world economy grows. Direct capture of CO2 from ambient air was first suggested by Lackner et al. in 1999 as a method to counteract global warming.3 Energy requirement1,4,5 and cost analysis studies68 claim that air capture is feasible and economically viable. At the same time, the uncertainty in economic assessments for future air capture implementation is significant, considering technique and market development.9,10 Success will depend on a more energy efficient sorbent cycle, and the use of a moisture-driven cycle offers a new opportunity to reduce energy consumption. To the best of our knowledge, this is the first time that a novel moisture swing process is used for the capture and concentrating of CO2 from a gas mixture. The moisture swing offers a new approach to regenerating CO2 sorbents: it trades input of heat in a thermal swing, or mechanical energy in a pressure-based swing against the consumption of water, whose evaporation provides the free energy that drives the cycle. Such an energy source as water is low in cost. r 2011 American Chemical Society
Compared to water consumption in biomass production,11 water consumption in a moisture swing1 is 2 orders of magnitude smaller. Moisture swing driven absorption cycles are of interest to air capture but also may prove of interest in other situations, as for example in capture from natural gas fired power plants. The objective of this paper is to systematically reveal the characteristics of the resin-based sorbent, specifically the change in thermodynamic properties in the presence of water. First, the material processes and physical properties will be introduced. A series of isothermal equilibrium experiments were conducted to determine the equilibrium constant, CO2 absorption capacity, and humidity effect during the moisture swing. The thermodynamic properties of the sorbent, such as heat of desorption, free energy change were derived from isotherms.
’ MATERIALS AND METHODS Material Process and Characterization. Pretreatment. A heterogeneous ion-exchange material in the form of a flat sheet is employed in this air capture study. The materials (I-200) tested were supplied by Snowpure LLC, California and were originally intended for use as electrochemical membranes. The sheet is manufactured by the company through coextrusion of a matrix polymer (polypropylene) and a resin powder comprising quaternary ammonium functional groups.12 The resin makes up about 60% of the weight of the membrane.1 The original Received: April 11, 2011 Accepted: June 20, 2011 Revised: June 15, 2011 Published: June 20, 2011 6670
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exchangeable anions are chloride ions. For the material to absorb CO2, the chloride ions must be replaced by hydroxide or carbonate ions. These ions are introduced by washing the sheet in hydroxide or carbonate solutions. The resin sheets were first cut into 5.0 cm 5.0 cm pieces with a thickness of 0.64 mm and pretreated by soaking in 8595οC deionized water (DI water) for 48 h. The hot water induces the resin swell enough to sufficiently open up the matrix structure and allow access to the ion-exchange resin particles. After soaking, the samples expanded in size to 6.0 cm 6.0 cm with a thickness of 0.8 mm. Two groups of samples were washed using a 1.0 M sodium hydroxide and a 0.5 M sodium carbonate solution respectively. The samples were immersed for 24 h and stirred to enhance the ion exchange. The samples were rinsed in DI water after every wash. The solution and wash residual were collected and titrated. The samples were washed 67 times until no chloride was detected in the solution by titration. Resin Ion Exchange Capacity. The quantity of chloride ions in the solution after the wash was titrated by the Mohr method13 to determine the effective charge density on the resin sheet. The titration results of the resin’s ion charge densities (Fc) are listed in Table 1. Upon preliminary inspection, the anion charge density determined for the carbonate is slightly higher than that for the hydroxide. The difference is much larger than the uncertainty of the titration, but variability in the pretreatment cannot be ruled out. Another reason for the discrepancy could be a preference of high-valence ions during ion exchange, possibly resulting in a ratio of inorganic carbon to cations on the resin that is slightly different from the one in the aqueous solution it is immersed in. Based on a simple charge balance, the estimated CO2 capacity for the carbonate should be half that of the hydroxide. Surface Analysis. The structures of the heterogeneous resin sheets were studied by SEM (JSM 5600LV) and are shown in Figure 1. The sheet has a porous structure with pore sizes ranging from several μm to 50 μm (Figure 1a). The sheet’s cross section presented in Figure 1b clearly shows that the material is heterogeneous. There are small spaces left between the resin and the Table 1. Absorption Performance and Fitting Results OH-
CO32-
Gc (mol/kg), resin’s ion charge density
1.78
1.90
qest (103m3/kg), CO2 at STP capacity estimated from Fc
39.8
21.3
q∞ (103m3/kg), CO2 capacity as measured
38.8
18.4
Ka (105), equilibrium constant of absorption
6.38
3.16
polypropylene matrix which can enhance gas diffusion. In order to observe the resin particles more clearly, resin-filled sheets were shaken in a high speed vibrating ball mill without a metal ball inside the vibrating chamber. The powders collected after shaking, which showed the same CO2 absorption/desorption performance as the resin-filled sheet itself, are resin particles. As shown in Figure 1c, the size of resin particles ranges from 30 to 50 μm. This is consistent with the particle size distribution shown in Figure 1b. Due to the porous structure, the air capture resin sheet has a surface area of about 2.0 m2/g (NOVA-2000 BET Surface Area Analyzer, Quantachrome). Since the weight for a 6.0 cm 6.0 cm piece of resin sheet sample is about 1.5 g, the surface area of the resin material is 400 times larger than the apparent surface area of the sheet.
’ EXPERIMENTAL METHODS An experimental device with temperature and humidity control was set up to determine the absorption/desorption performance of the flat resin sheets described in the previous section. A layout of the device is shown in Figure S1. The CO2 concentration is measured online with an infrared gas analyzer (IRGA, LI-COR, LI-840). In order to avoid water condensation when the dew point in the gas is higher than room temperature, the control unit including all of its tubing was enclosed in an insulated box with a separate temperature control. Absorption Isotherms. Wet, fresh samples were flushed with either dry N2 or zero air (air without CO2 or water), and the water concentration of the outlet gas was measured by the IRGA to determine whether the samples were sufficiently dry. Drying was performed with a gas flow rate of 17 cm3/s, and it was stopped when the exhaust gas contained less than 0.2% water vapor. A fixed volume of CO2 (0.5 cm3) was repeatedly injected into the sample container by syringe, and the gas was cycled for 24 h to allow the resin to absorb the CO2 until the system reached equilibrium. The volume of the sample chamber is 1.15 L (liter), and the CO2 concentration increased 430 ppm after every injection commensurate with the volume ratios. The CO2 leakage rate was measured as 0.14 ppt/s (parts per trillion per second) at a concentration difference of 500 ppm between the inside and the outside of the device. Experiments verified that the leakage rate is indeed proportional to the concentration difference. Drift of the detector has been observed, but it is too small to account for these changes. Considering that the CO2 leakage would be slower when the concentration difference is smaller (leakage rate of 0.07 ppt/s with a concentration difference of 100 ppm), the amount of CO2 leaking into or out of
Figure 1. SEM pictures of (a) the surface structure of the resin filled sheet, (b) sheet cross section structure, and (c) resin particles removed from the polypropylene sheet. 6671
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the test system in 24 h would be less than 0.01 mL, significantly less than a single injection of 0.5 mL. At equilibrium, the CO2 absorbed can be determined as qe ¼ qinj V0 3 C
ð1Þ
where qinj, V0, and C are the total amount of injected CO2, the volume of the sample chamber, and the observed equilibrium CO2 concentration, respectively. Desorption Isotherms. The CO2 absorbed can be driven off the resin by exposing the sorbent to a high humidity atmosphere or to liquid water. This study focuses on the liquid water swing performance. In order to measure the relationship between saturation of wet sorbent and released CO2 equilibrium partial pressure, desorption experiments were conducted with different amounts of resin in fixed-size sample chambers. Because of the different resin loading in the chamber, the different experiments result in different equilibrium CO2 partial pressures. The initially dry samples were first exposed to pretreated air without water vapor for two hours to eliminate the moisture effect and get to an initial standard saturation state. Then the fully loaded samples were wetted with DI water, at which point they began to release CO2. The samples were wetted by filling the chamber completely with DI water in order to ensure that the entire resin made contact with water. Immediately after wetting, the water was drained and displaced with zero air. Because the sample chamber was filled first with DI water and then with zero air, all CO2 present in the chamber has to come from the desorption process. This method will tend to underestimate the total amount of CO2 desorbed because part of the CO2 would have been dissolved into the water and thus be excluded from the measurement. However, the water remained in the chamber for only a short time (less than two minutes). This would not have given the water enough time to reach equilibrium with the CO2 partial pressure in the chamber. Experiments also show that typically less than 10% of the measured CO2 was released by the time the water was removed from the chamber. Furthermore, we can with high confidence rule out an unmeasured, unexpected initial burst of CO2. Anzelmo et al.14 performed experiments in which the resin is immersed in water and where he measured the CO2 release rate into water. These experiments showed that the CO2 was gradually released and strengthen the observation in this paper that the CO2 undercount in the experiment is not very large. Even under the assumption that the water was saturated in CO2, the CO2 dissolved into the water would be less than 5% of the total desorbed CO2. The fully loaded sorbent, which initially is in the bicarbonate form, approaches the carbonate form as CO2 is released. The saturation of the wet sample at a certain equilibrium state can be calculated using the CO2 concentration at equilibrium θ ¼ ðm 3 q∞ V0 3 CÞ=ðm 3 q∞ Þ
ð2Þ
where m is the mass of sorbent, and q∞ is the CO2 capacity for carbonate form sorbent measured by absorption experiments. The desorption isotherm at different temperatures was then studied.
’ DATA MODEL We start from the observation that the resin when it is dry absorbs CO2 and reaches a state in which it holds a CO2 molecule for nearly every cation attached to the resin. This, we hypothesize can be approximated by a Langmuir absorption model.
Furthermore, when wet, the resin releases CO2 and falls back to a loading state which includes approximately one CO2 molecule for every two cation ions attached to the resin. We assume for our data analysis that this behavior, too, can be described by a Langmuir absorption model. Capacity. Sorbents in the hydroxide form or the carbonate form can both absorb CO2 from ambient air. Since the resin in the carbonate form already contains one CO2 per two positive charges, the CO2 capacity of the carbonate would be half of the hydroxide. The CO2 capacities for the sorbent-filled sheets were estimated (qest) from the charge density results and are listed in Table 1. Isothermal Equilibrium. For a gassolid system, the widely studied Langmuir isotherm equation can be employed to describe the relationship between sorbent coverage and CO2 partial pressure in the atmosphere. The equation is expressed as θ¼
KP 1 + KP
ð3Þ
where θ is the saturation of the sorbent, K is the equilibrium constant, and P is the equilibrium CO2 partial pressure in the gas. For a CO2 sorbent: θ is defined as qe/q∞, where qe is the CO2 absorbed at equilibrium (m3/kg), and q∞ is the CO2 absorption capacity (m3/kg). Equation 3 can be transformed into a LineweaverBurk equation15 1 1 1 1 + ¼ qe Ka q ∞ 3 P q ∞
ð4Þ
For an exact Langmuir behavior, 1/qe is a linear function of 1/P. The CO2 absorption capacity and absorption equilibrium constant, Ka, can be determined from the intercepts and slope of the straight line of 1/qe as function 1/P. In the case of this resin, one needs to distinguish between the sorbent characteristics in the dry state and in the wet state. The parameters of the Langmuir forms of the two situations differ and have both been measured. The Langmuir equation relates the partial pressure of CO2 to the equilibrium saturation of the sorbent. In order to make this relation meaningful, it is necessary to define the state of the resin when it is at zero saturation state and at a fully saturated state. For a realistic system it is always possible that other mechanisms could add or remove additional CO2. In this case, in particular one might consider that the mechanism for CO2 absorption is different for a resin that is in the hydroxide form and one that starts out in a carbonate form. As a result we consider the possibility that the starting point of the Langmuir approximation is not the hydroxide but the carbonate. Unfortunately, as it is explained in the Supporting Information, the data obtained at high loading cannot resolve between these two possibilities. This is also visible in Figure 2, where we show both fits, the one starting from the carbonate state, and the one starting from the hydroxide state. Since the desorption of the resin stops at the carbonate state, we have chosen a Langmuir fit to the data which assumes that the zero loading state is equivalent to the carbonate state, i.e., the ‘empty’ resin already contains one-half mole of CO2 per mole of cation attached to the resin.
’ EXPERIMENTAL RESULTS AND DISCUSSION Absorption Performance. Capacity and Isotherms. The first set of experiments measured the relationship between sorbent saturation and CO2 equilibrium partial pressure. Sorbents 6672
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Figure 2. The absorption isotherms of sorbents. Discrete points are experimental data.
prepared as hydroxides and carbonates were tested respectively at a constant temperature of 23 ( 0.5 °C and 0.5% H2O. The dew point of the experimental environment is about 2.7 οC which is the average weather condition during winter and spring season at a dry location, like Arizona. The measured isotherms and best fit to Langmuir models are displayed in Figure 2. We model the hydroxide sorbent by assuming that θ = 0 corresponds to the initial hydroxide, for the carbonate sorbent we assume θ = 0 corresponds to the carbonate state. In Figure 2 the left vertical axis corresponds to the carbonate case, the right one to the hydroxide case. Assuming the hydroxidebase sorbent half-filled with CO2 (θ is 0.5) is equivalent to a carbonate-based sorbent that has not yet absorbed any CO2 (θ is 0), the isotherms should overlap in Figure 2. The experimental results in Figure 2 clearly prove that the two materials are indeed identical at high saturations. In effect the material has no memory of how it was initialized. Figure 2 also shows that at the ambient CO2 concentration level of 400 ppm and water vapor concentration of 0.5%, the hydroxide and carbonate sorbents both have a very high saturation of over 99%. Even at very low CO2 concentrations, such as 100 ppm, the sorbents still have a saturation of more than 95%. Clearly, the quaternary ammonia resin is more than strong enough for CO2 absorption from air. The Langmuir curve fitted to the data provides equilibrium constants and CO2 capacity. The results are summarized in Table 1. Based on the anion densities, nearly all of the ion exchangeable functional groups were able to absorb CO2 within the range of experimental error. As reviewed by Choi et al.,16 the capacity of resin-filled sheet ranks on the high end of the spectrum for air capture sorbents. Moreover, the load factor is less important than in other sorbent cycles since the moisture swing does not involve lowering or raising the temperature of the bulk material. Humidity Effect. The data in Figure 2 were obtained at constant humidity. Figure 3 shows the CO2 saturation dependence on humidity. The graph represents a water vapor isotherm at a constant CO2 concentration of 400 ppm and a constant temperature of 23 ( 0.5 οC. The sample was prepared in a fresh carbonate state. After a known volume of CO2 was injected into the sample chamber, the humidity was adjusted until the CO2 concentration stabilized at 400 ppm. The solid dot in Figure 3 is the Langmuir model saturation data from the absorption isotherms at 400 ppm and 0.5% water vapor, presented in Figure 2. There is a strong humidity effect on absorption equilibrium at room temperature. Desorption Performance. The sorbent can readily desorb CO2 as water vapor concentration increases. It is also possible to expose the resin to liquid water to force the release of the CO2.
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Figure 3. Humidity effect on CO2 saturation of sorbent at a constant temperature of 23 ( 0.5 °C.
Figure 4. The CO2 desorption isotherms at different temperatures for wet sorbent material.
Table 2. Desorption Isothermal and Thermodynamic Parameters 24 οC
35 οC
45 οC
Kd (103)
2.30
3.78
5.37
r
0.96
0.98
0.96
ΔGd,T (kJ/mol) ΔHo (kJ/mol)
14.98 31.81
14.35
13.79
ΔSo (J/(mol 3 K))
56.65
Unlike the humidity swing, the water swing would require clean water and would most likely incur a higher water loss. However, the liquid water swing has the potential to improve desorption kinetics and to simplify the sorbent regeneration process. A water swing may be advantageous over a humidity swing from a heat management point of view because the humidity swing may require elevated temperatures to produce water vapor, which would not be necessary for a water swing. We, therefore, studied the desorption characteristics of a wet resin. Desorption Isotherms. The sorbent only converts from bicarbonate state back to carbonate state during desorption. The relationship between sorbent saturation and CO2 equilibrium partial pressure during desorption is illustrated in Figure 4. Via desorption the CO2 partial pressure can be easily elevated to several kPa. Given the CO2 concentration in ambient air is about 400 ppm (0.04 kPa), the partial pressure is increased by 2 orders of magnitude simply by making the sorbent wet. Higher temperature favors desorption, especially at lower CO2 partial pressure. The values of the equilibrium constants and correlation coefficients (r) for the data sets at different temperatures are tabulated in Table 2. Langmuir isotherms using the best fit for desorption equilibrium constant, Kd, are plotted in Figure 4, 6673
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which clearly shows that desorption processes are also well described by the Langmuir model. As shown in Table 2, Kd decreases with increasing temperature. For reaction with an ideal gas, the relationship between the equilibrium constant and temperature is related to the standard Gibbs energy change of reaction as ΔGo ¼ RTln K
ð5Þ
and can be further expressed as ln K ¼
ΔH o 1 ΔSo + R 3T R
ð6Þ
where ΔGo, ΔHo, and ΔSo are standard change of free energy, enthalpy, and entropy, respectively. R is the universal gas constant with the value of 8.314 J/(mol 3 K), and T is the absolute temperature. By fitting the curve of ln Kd vs 1/T, the thermodynamic parameters during desorption can be determined using the slope and intercept. Since the resin temperature change is not large, the change of enthalpy and entropy with temperature can be neglected. The free energy change of desorption at different temperature can be estimated as ΔGd, T ¼ ΔH o TΔSo
ð7Þ
All of the above thermodynamic parameters were obtained and are listed in Table 2. The positive values of enthalpy and entropy indicate the endothermic nature of the desorption process and the increase in the degree of freedom of the absorbed species, i.e., the releasing of CO2. The positive value of free energy indicates that the reaction is not spontaneous and would happen only if enough free energy is provided. The water not only provides an aqueous environment for free proton and ion exchange during desorption but also directly or indirectly provides the necessary free energy for CO2 release. A Hypothetical Mechanism. The alkaline anion groups of hydroxide or carbonate on the resin surface, when exposed to dry air, can capture CO2 and produce bicarbonates, they can also form carbonates, and there has to be some equilibrium between hydroxide, carbonate, and bicarbonate OH R + + CO2 f HCO3 R +
ð8Þ
CO3 2 ðR + Þ2 + H2 O + CO2 ðgaseousÞ f 2ðHCO3 R + Þ ð9Þ OH- R + + HCO3 R + T CO3 2 ðR + Þ2 + H2 O
ð10Þ
2ðHCO3 R + Þ f CO3 2 ðR + Þ2 + H2 O + CO2 ðgaseousÞ ð11Þ +
R represents the quaternary ammonium ion in the resin. Because the resin is never entirely dry under ambient conditions with vapor pressure present, the water consumed in eq 9 can be provided by the hygroscopic resin. The net reaction of eq 10 that converts bicarbonate and hydroxide into carbonate and water may occur through states in aqueous solution. Our working hypothesis for the resin absorption/desorption mechanism is that in the dry state, reaction of eq 8 proceeds without reaction of eq 10. In other words, in a dry system the equilibrium allows for the coexistence of bicarbonate ions and hydroxide ions in far larger concentrations than in an aqueous environment. The resin will absorb CO2 until all hydroxide is consumed, reaching the bicarbonate state. Adding moisture will shift the bicarbonate-carbonate equilibrium to what is expected in water, and CO2 is released via eq 11. This is consistent with the
behavior of a 1-molar sodium bicarbonate solution. At 25 °C, the equilibrium partial pressure of CO2 above a one molar NaHCO3 solution is 6.0 kPa when 20% of the bicarbonate is decomposed to carbonate.17 The driving force toward a higher CO2 release may be further amplified because it is known that strong base resins often have a preference for divalent ions, thereby raising the effective concentration of bicarbonate ions in the solution.18 As the resin dries, the equilibrium represented by eq 10 shifts back toward the left. This shift in equilibrium is counterintuitive as there is water on the product side of the equation. However, the equation ignores the associated water molecules that are bound into the hydration clouds of the various ions, and it is reasonable to assume that their numbers are reduced as the system dries out. Thus the net balance of water in the equation is more complicated than it appears at first. While our results do not confirm this hypothesis, they are in agreement. The dry (with hydrated water equilibrium with the ambient vapor pressure rather than anhydrous) resin’s ability to absorb CO2 is likely related to its inherent structure and geometry. The equilibrium in eq 10 may move to the left because two monovalent positive quaternary ammonium groups are needed to balance one divalent carbonate anion. If the spacing of cationic groups far exceeds the size of the carbonate ion or the similarly sized bicarbonate ion,19 then it could become energetically unfavorable for the two ammonium ions to hold onto a single carbonate ion. With the cationic groups far apart, they could not hold the polyvalent anions in a stable state and would greatly prefer monovalent anions. Previous research20,21 on the behavior of resin has proven these ideas. Kimura et al.20 discovered that the proton was liberated at neutral pH when bicarbonates interact with polyamines with proper geometry. Horng et al.21 also reported the strong-based resins could convert HCO3 to CO32- and H2AsO4 to HAsO42- with the expulsion of a proton. They further discussed that several resin characteristics, such as appropriate charge spacing and presence of hydrophilic groups, were related with the ability to produce high-valent anions.
’ ASSOCIATED CONTENT
bS
Supporting Information. Schematic of experimental system; absorption models; error analysis. This material is available free of charge via the Internet at http://pubs.acs.org.
’ AUTHOR INFORMATION Corresponding Author
*Phone: (212)851-0243. Fax: (212)854-7081. E-mail: tw2296@ columbia.edu.
’ ACKNOWLEDGMENT The authors thank the support from the Lenfest Center for Sustainable Energy, Columbia University. ’ NOMENCLATURE C = measured equilibrium CO2 concentration [ppm]; Go = standard change of free energy at standard conditions (STP, 100 kPa, 298.15 K) [kJ/mol]; K = equilibrium constant; Ka = absorption equilibrium constant; Kd = desorption equilibrium constant; m = mass of sorbent [kg]; P = equilibrium CO2 partial pressure [Pa]; qest = CO2 capacity estimated from resin’s ion charge density [103 m3/kg]; q∞ = CO2 capacity [103 m3/kg]; 6674
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Environmental Science & Technology qe = CO2 absorbed at equilibrium [m3]; qinj = total amount of injected CO2 [m3]; r = correlation coefficient; R = universal gas constant [8.314 J/mol 3 K]; T = temperature (K); V0 = volume of the sample chamber [m3]; Fc = resin’s ion charge densities [mol/kg]; θ = saturation of the sorbent; ΔHo = standard change of enthalpy and [kJ/mol]; ΔSo = are standard change of entropy [J/(mol 3 K)]
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