CO2 capture properties of alkaline earth metal oxides and hydroxides:A combined density functional theory and lattice phonon dynamics study

Yuhua Duan1,2,a� and Dan C. Sorescu1

1National Energy Technology Laboratory, United States Department of Energy,Pittsburgh, Pennsylvania 15236, USA2URS Corp., P. O. Box 618, South Park, Pennsylvania, Pennsylvania 15219, USA

�Received 7 May 2010; accepted 9 July 2010; published online 19 August 2010�

By combining density functional theory and lattice phonon dynamics, the thermodynamic propertiesof CO2 absorption/desorption reactions with alkaline earth metal oxides MO and hydroxidesM�OH�2 �where M=Be,Mg,Ca,Sr,Ba� are analyzed. The heats of reaction and the chemicalpotential changes of these solids upon CO2 capture reactions have been calculated and used toevaluate the energy costs. Relative to CaO, a widely used system in practical applications, MgO andMg�OH�2 systems were found to be better candidates for CO2 sorbent applications due to their loweroperating temperatures �600–700 K�. In the presence of H2O, MgCO3 can be regenerated intoMg�OH�2 at low temperatures or into MgO at high temperatures. This transition temperaturedepends not only on the CO2 pressure but also on the H2O pressure. Based on our calculated resultsand by comparing with available experimental data, we propose a general computational searchmethodology which can be used as a general scheme for screening a large number of solids for useas CO2 sorbents. © 2010 American Institute of Physics. �doi:10.1063/1.3473043�

I. INTRODUCTION

Carbon dioxide is one of the major combustion productswhich, once released into the air, can contribute to the globalclimate warming effects.1 Solid sorbents containing alkaliand alkaline earth metals have been reported in several pre-vious studies to be good candidates for CO2 sorbent applica-tions due to their high CO2 absorption capacity at moderateworking temperatures.2,3 In our previous studies,4,5 we haveexplored the CO2 capture capabilities of alkali metal oxides,hydroxides, and carbonates/bicarbonates systems, and foundthat M2CO3 /MHCO3 �M=Na,K� transformation systemscan be used efficiently as CO2 sorbents. In this study, weextend our previous analysis by exploring the CO2 capturecapabilities of alkaline earth metal oxides and hydroxides,and then, based on these results, we present a general com-putational scheme that can be used to screen solid materialsfor CO2 sorbent development.

Relative to the extensive experimental studies of captur-ing CO2 with alkali and alkaline earth oxides,2,3 capturingCO2 with solid alkaline hydroxides is relatively a new ap-proach which is receiving an increasing attention from sci-entific community. For example, a novel sodium-based sor-bent that can capture CO2 from ambient to 600 °C wasdeveloped by Siriwardane’s research group at National En-ergy Technology Laboratory �NETL�.6 This sorbent consist-ing of a NaOH/CaO mixture can capture CO2 at 315 °C andcan be regenerated at 700 °C. Such characteristics make ituseful for high-temperature CO2 capture applications in coalgasification systems. Recently, the same group has reportedthe development of a Mg�OH�2 sorbent which can be used

for CO2 capture at warm gas temperatures.7 In this case, CO2

capture takes place in the 200–315 °C range while regen-eration is done above 375 °C, making it suitable for precom-bustion capture technologies. Stolaroff et al.8 explored thefeasibility of using spray-based NaOH to capture CO2 fromatmospheric air and concluded that the cost of CO2 captureusing NaOH spray in the full-scale system ranges from $53to $127 per ton CO2. From Mg-rich mineral, Lin et al.9 ex-tracted Mg�OH�2 powder, which has domain size as small as12 nm and apparent surface area of 54 m2 /g. Under 1 atm of10 vol % CO2 /N2, the carbonation of this Mg�OH�2 powdercan achieve up to 26% of the stoichiometric limit at 325 °Cin 2 h and, moreover, it was found that the amount of CO2

fixation was inversely proportional to the crystal domain sizeof the Mg�OH�2 specimens, which suggests that only amonolayer of carbonates was formed on the boundaries ofthe crystal domains in the gas-solid reaction with little pen-etration of the carbonates into the crystal domains. UsingCa�OH�2 as a CO2 sorbent in the gasification of wet biomass,Hu et al.3,10 determined that Ca�OH�2 powder can functionnot only as a CO2 sorbent but it also had a catalytic actionupon the tar cracking.

By contrast to the relative large number of experimentalstudies of CO2 reactions with alkali and alkaline earth oxidesand hydroxides, to date there are only few theoretical studiesof the same topics. Based on thermodynamic data,2,3 Feng etal.11 analyzed 11 simple metal oxides and concluded thatCaO is thermodynamically the best candidate among themfor CO2 capture in zero emission power generation systems.Similarly, Siriwardane and Stevens7 have determined basedon the analysis of thermodynamic data that Mg�OH�2 sorbentsystem is highly favorable for CO2 capture applications andcan be regenerated with high-pressure CO2. Through

a�Author to whom correspondence should be addressed. Tel.: 412-386-5771.FAX: 412-386-4542. Electronic mail: yuhua.duan@netl.doe.gov.

THE JOURNAL OF CHEMICAL PHYSICS 133, 074508 �2010�

0021-9606/2010/133�7�/074508/11/$30.00 © 2010 American Institute of Physics133, 074508-1

ab initio evolutionary simulations and high-pressure experi-ments, Oganov et al.12 investigated the high-pressure struc-tures of MgCO3, CaCO3, and CO2 and their roles in theEarth’s lower mantle. Based on density functional theory�DFT�, Jensen et al.13 investigated the CO2 adsorption onCaO and MgO surfaces. Their results showed that CO2 ad-sorbs as monodenate on edge sites and bidentate on cornersites of MgO. In contrast, CO2 adsorbs as monodenate onboth edge and corner sites of CaO. Using computational fluiddynamics �CFD�, Liu et al.14 simulated the chemical absorp-tion of CO2 from air by NaOH aqueous solution and foundthat the temperature and the concentration distributions alongthe height of the column are consistent with measured ex-perimental data.

Recently, we proposed a methodology to identify prom-ising solid sorbent candidates for CO2 capture by combiningDFT total energy calculations with lattice phonondynamics.15,16 It was concluded that although pure Li2O canabsorb CO2 efficiently, it is not a good solid sorbent for CO2

capture applications because the reverse reaction, corre-sponding to Li2CO3 releasing CO2, can only occur at verylow CO2 pressure or at very high temperature when Li2CO3

is in liquid phase.16 These predicted results are in very goodagreement with experimental measurements.17 The samecomputational methodology has been extended recently byus to describe the phase diagrams of M–C–O–H�M=Li,Na,K� systems.5 We have determined thatNa2CO3 /NaHCO3 and K2CO3 /KHCO3 are promising candi-dates for both precombustion and postcombustion capture inagreement with previous experimental findings.

In this study, we apply our analysis based on calculationof the thermodynamic data using first-principles densityfunctional theory and lattice phonon dynamics to screen alarger number of alkaline earth metal oxides, hydroxides, andtheir corresponding composites, and to explore their CO2

capture properties systematically. We compare our predic-tions with the available thermodynamic data to assess theaccuracy of our approach and outline a general screeningscheme that can be used to predict the CO2 capture proper-ties of the new types of materials for which thermodynamicdata might not be available.

This paper is organized as follows: Sec. II briefly de-scribes the theoretical methods employed. In Sec. III, wepresent the results of the electronic and phonon properties ofalkaline earth metal oxides, hydroxides, and carbonates, fol-lowed by the analysis of the thermodynamic properties of thecorresponding CO2 capture reactions. A general screeningmethodology that can be applied to solid materials for whichthermodynamic properties might not be available is pro-posed. Finally, a brief summary and conclusions are indi-cated in Sec. IV.

II. THEORETICAL METHODS

The complete description of our computational method-ology can be found in our previous papers.4,5,15,16 Here, welimit ourselves to provide only the main aspects relevant for

the current study. The CO2 capture reactions by solids in thepresence of water vapors can be expressed generically in theform

Solid _ A + n1CO2 ↔ Solid _ B + �Solid _ C� � n2�H2O� ,

�1�

where the terms given in �…� are optional and n1 and n2 arethe numbers of moles of CO2 and H2O involved in the cap-ture reactions. We treat the gas phase species CO2 and H2Oas ideal gases. By assuming that the difference between thechemical potentials ���0� of the solid phases of A, B �and C�can be approximated by the difference in their total energies��EDFT�, obtained directly from DFT calculations, the vibra-tional free energy of the phonons,18 and by ignoring thepressure-volume �PV� contribution terms for solids, thevariation of the chemical potential ���� for reaction �1� withtemperature and pressure can be written as15,16,19

���T,P� = ��0�T� − RT lnPCO2

n1

PH2O�n2

, �2�

where

��0�T� � �EDFT + �EZP + �FPH�T�

− n1GCO2�T� � n2GH2O�T� − �H0. �3�

Here, �EZP is the zero point energy difference betweenthe reactants and products and can be obtained directly fromphonon calculations. �H0 is an empirical correction constantand �EPH is the phonon free energy change between thesolids of products and reactants as presented in the follow-ing. If reaction �1� does not involve H2O, then the PH2O inEq. �2� �and also in the following Eq. �13�� is set to P0, whichis the standard state reference pressure of 1 bar, and the GH2O

term in Eq. �3� is not present. The “+” and “�” signs in Eqs.�2� and �3� correspond to the cases when H2O is a productand a reactant, respectively, in the general reaction �1�. Thefree energies of CO2 �GCO2

� and H2O �GH2O� can be obtainedfrom standard statistical mechanics20 as

GCO2�

7

2RT + �

i=1

4Nah�i

eh�i/kT − 1− TSCO2

�T� , �4�

GH2O � 4RT + �i=1

3Nah�i

eh�i/kT − 1− TSH2O�T� , �5�

where Na is Avogadro’s constant. The entropy of CO2

�SCO2�T�� and H2O �SH2O�T�� can be accurately calculated

from the Shomate equation.21 The vibrational frequencies��i� of CO2 molecule are taken as 673 ��u�, 1354 ��g

+�, and2397 cm−1 ��u

+�,22 and the vibrational frequencies ��i� ofH2O molecule are 3657.05 ��1�, 1594.75 ��2�, and3755.93 cm−1 ��3�.23 The zero point energies for CO2 andH2O molecules calculated using these frequencies are 0.3160and 0.5584 eV, respectively.

In Eq. �3�, �EDFT is the total energy change of the reac-tants and products calculated by DFT. In this work, theVienna ab initio simulation package24 was employed to cal-culate the electronic structures of solid oxides, hydroxides,

074508-2 Y. Duan and D. C. Sorescu J. Chem. Phys. 133, 074508 �2010�

and carbonate materials. In this case, all calculations havebeen done using the projector augmented wave �PAW�pseudopotentials with the PW91 exchange-correlationfunctional.25 Plane wave basis sets were used with a cutoffenergy of 500 eV and a kinetic energy cutoff for augmenta-tion charges of 605.4 eV. The k-point sampling grids of thegeneral form m1m2m3 were obtained using theMonkhorst–Pack method.26 In our bulk calculations, the in-dices m1, m2, and m3 were selected to corresponding consis-tently to a spacing of about 0.028 Å−1 along the axes of thereciprocal unit cells. During optimizations all atoms in thecell as well as the lattice dimensions and angles were fullyrelaxed to equilibrium. The total energies of CO2 and H2Omolecules are �22.994 09 and �14.272 67 eV, respectively,as determined from calculations of the isolated molecules ina cubic box with the length of 20 Å.

In the harmonic approximation, the phonon free energychange ��FPH� and the entropy change ��SPH� between thesolid reactants and products can be calculated based on theHelmholtz free energy of the solids Fharm and the entropy ofthe solids �Sharm�

�FPH�T� = �solid_products

Fharm�T� − �solid_reactants

Fharm�T� ,

�6�

�SPH�T� = �solid_products

Sharm�T� − �solid_reactants

Sharm�T� ,

�7�

where the Fharm, Sharm, and the total vibrational �phonon�energy �Etot� of the solids are defined as18

Fharm = rkBT�0

g���ln2 sinh ��

2kBT��d� , �8�

Sharm = rkB�0

g��� ��

2kBT�coth ��

2kBT� − 1�

− ln�1 − e−��/2kBT��d� , �9�

Etot =1

2r�

0

g�������coth ��

2kBT�d� . �10�

In the above equations r is the number of degrees of freedomin the primitive unit cell, � is the phonon dispersion fre-quency, and g��� is the phonon density of states. It can beseen that the zero point energy �EZP� can be obtained fromEq. �10� by taking T→0,

Ezp = limT→0

�Etot�T�� . �11�

In this study, the phonon properties were evaluated using thePHONON software package27 in combination to the directmethod formalism derived by Parlinski et al.28 Specifically,supercells containing at least 222 unit cells were used inall phonon calculations. Structures with displacements of0.03 Å of the nonequivalent atoms were generated from theoptimized supercells and DFT calculations were further per-

formed to obtain the forces on each atom due to these dis-placements. These forces were input into the PHONON

package27 to fit the force constant matrix and to compute thephonon density of states. These data were subsequently usedin Eqs. �8�–�10� to calculate the thermodynamic properties.In phonon dispersion calculations, the coordinates of thehigh symmetry points in the first Brillouin zone of the crys-tals are consistent with those defined by Bradley andCracknell.29

The enthalpy change for the reaction �Hcal�T� can bederived from the above equations as

�Hcal�T� = ��0�T� + T�n1SCO2�T� n2SH2O�T�

+ �SPH�T�� . �12�

As discussed in our previous paper,5 the DFT calculatedenthalpy and the chemical potential of reaction are alwaysabout 20 kJ/mol higher than the experimental measurements.This is likely due to limitations of DFT in computing theheats of formation and similar differences were also ob-served for other systems.30,31 In order to empirically correctfor this apparent bias, we have subtracted a constant value�H0 from Eq. �3�. We have found that using a correctionterm �H0=20 kJ /mol gives reasonable performance for allsystems of oxides and hydroxides considered in this study.By taking the activities of all reactant and product phases tobe unity, the equilibrium pressure of the overall reaction canbe calculated by setting ���T,P�=0 in Eq. �2�. Hence, theequilibrium pressure is given by30

PCO2

PH2O�n2/n1

= exp��0�T�RT

� . �13�

By plotting the equilibrium pressures calculated fromEq. �13� as function of temperature T, one can obtain thevan’t Hoff plot for CO2 capture reactions.

Comparison of the calculated data to the experimentalvalues for the CO2 capture reactions has been done using theHSC CHEMISTRY package.32 In this last case, the experimentalmeasured thermodynamic data for each compound is fittedusing analytic polynomials such as the Kelly equation for thedependence of the heat capacity on temperature, and the re-sulted polynomial equations are further used to calculate thethermodynamic properties for reactions of interest.

III. RESULTS AND DISCUSSIONS

A. DFT and phonon calculated results

The optimized lattice constants and the total energies fora set of 15 alkaline earth metal oxides, hydroxides, and car-bonates considered in this work are presented in Table I,along with the corresponding experimental structural data. Itshould be pointed out that each compound might have morethan one known phase in the temperature and pressure rangeinvestigated. However, in Table I, for each compound, onlythe most stable structure with the lowest energy is listed. Wenote that we were unable to find the crystal structure ofBeCO3 although its thermodynamic data are available in theliterature. From our calculations, the energy for the MgCO3

based on the crystallographic structure of BeCO3 is about 0.5

074508-3 Theoretical study on CO2 capture properties J. Chem. Phys. 133, 074508 �2010�

eV per formula unit higher than the one for the structurebased on SrCO3 symmetry. Therefore, for BeCO3, we as-sumed that this crystal adopts the same crystallographic sym-metry as SrCO3.

As the data in Table I indicates, the agreement betweenthe optimized structural constants and the experimental datais generally very good. The calculated electronic structuralproperties �such as geometrical structure, band structure,etc.� are also in good agreement with other theoretical12,33,34

and experimental findings.35,36 The calculated total energy

�EDFT� for each compound at the optimized configuration hasbeen used to estimate the reaction enthalpy for CO2 capturereaction as described in Sec. III B.

Phonon calculations were performed for each of thecompounds listed in Table I and the finite temperature ther-modynamic properties were then computed from these datausing Eqs. �2�–�12�. As an example, the variation of the cal-culated phonon free energy and entropy for M�OH�2 andMCO3 �M=Ca,Mg� as a function of temperature are shownin Figs. 1�a� and 1�b�.

TABLE I. Comparison of the experimental and the DFT calculated structural parameters for the list of compounds involved in the CO2 capture reactionsstudied. All distances are given in angstroms and angles in degrees. Although calculations were performed on several phases of each compound, only the moststable phase, determined by DFT total energy, is listed in this table. The zero point energy and entropy calculated from phonon lattice dynamics as well as thecorresponding experimental data are also indicated.

Compound Space group and reference

Structural parametersCalculated Energy

�eV/f.u.�aEntropy

�J/mol K�

Experimental Calculated EDFT EZP Phonon �T=300 K� Expt.b �T=298.15 K�

BeO P63mc �No.186�, Ref. 41 a=2.6979 a=2.710 59 �14.366 57 0.225 59 14.713 13.770c=4.3772 c=4.400 06�=120° �=120°

MgO Fm3̄m �No.225�, Ref. 42 a=4.2198 a=4.248 88 �12.007 59 0.126 11 33.294 26.950

CaO Fm3̄m �No.225�, Ref. 43 a=4.8152 a=4.819 03 �12.987 52 0.110 88 39.374 38.10

SrO Fm3̄m �No.225�, Ref. 44 a=5.1326 a=5.193 50 �12.192 23 0.079 03 56.742 55.580

BaO Fm3̄m �No.225�, Ref. 45 a=5.539 a=5.660 43 �11.955 82 0.064 92 74.184 72.00Be�OH�2 P212121 �No.19�, Ref. 46 a=4.5301 a=4.508 25 �29.293 64 0.909 42 39.4651 53.555

b=4.621 b=4.650 19c=7.048 c=7.011 85

Mg�OH�2 P3̄m1 �No.164�, Ref. 35 a=3.1425 a=3.168 16 �27.149 38 0.775 31 65.8659 63.137c=4.7665 c=4.749 01�=120° �=120°

Ca�OH�2 P3̄m1 �No.164�, Ref. 43 a=3.582 a=3.613 77 �28.335 93 0.711 40 82.8147 83.400c=4.904 c=4.967 88�=120° �=120°

Sr�OH�2 Pnam �No.62�, Ref. 47 a=9.8889 a=10.000 19 �27.725 16 0.723 79 93.8210 97.069b=6.1202 b=6.046 30c=3.9184 c=3.947 27

Ba�OH�2 Pnma �No.62�, Ref. 48 a=11.03 a=11.150 59 �27.698 74 0.701 18 109.9148 107.280b=16.56 b=16.761 30c=7.11 c=7.102 24

BeCO3 Pnma �No.62� No expt. data available. a=6.762 93 �36.730 83 0.578 29 50.8681 52.007Taken same as SrCO3 b=4.851 91

c=7.012 98MgCO3 R3̄cH �No.167�, Ref. 49 a=4.6338 a=4.686 49 �35.960 46 0.532 35 69.3531 65.090

c=15.0192 c=15.137 95�=120° �=120°

CaCO3 R3̄cH �No.167�, Ref. 50 a=4.991 a=5.039 79 �37.610 11 0.484 10 95.9951 91.710c=17.068 c=17.126 72�=120° �=120°

SrCO3 Pnma �No.62�, Ref. 51 a=6.02 a=6.102 09 �37.335 26 0.460 27 104.4270 97.200b=5.093 b=5.156 53c=8.376 c=8.490 87

BaCO3 Pmcn �No.62�, Ref. 52 a=5.3293 a=5.370 61 �37.522 45 0.445 84 116.7325 112.100b=8.9275 b=9.007 70c=6.6076 c=6.552 10

CO2 molecule P1�Dh� rC–O=1.163 rC–O=1.1755 �22.994 09 0.315 98 213.388H2O molecule P1�C2v� rO–H=0.957 rO–H=0.9714 �14.272 67 0.558 41 188.832

�HOH=104.4° �HOH=104.2°

aThe acronym f.u. corresponds to one formula unit.bData taken from HSC CHEMISTRY Package �Ref. 32�.

074508-4 Y. Duan and D. C. Sorescu J. Chem. Phys. 133, 074508 �2010�

From Fig. 1�a�, one can see that the free energy differ-ences between Mg�OH�2 /MgCO3 and Ca�OH�2 /CaCO3 arequite similar with the increase in temperature, but the corre-sponding entropy differences are more pronounced. The en-tropy of Mg�OH�2 and MgCO3 are very close to each other,especially at low temperatures ��600 K� as shown in Fig.1�b�, but the entropy of CaCO3 is higher than that ofCa�OH�2. Such differences have important implications uponCO2 capture behavior as will be detailed in Sec. IV of this

paper. The calculated zero point energies obtained based onEq. �11� for each compound are also listed in Table I.

From the data in Table I, one can see that our calculatedentropies of these oxides, hydroxides, and carbonates arevery close to the experimental measured values. These find-ings indicate that reasonable agreement to experimental datafor thermodynamic properties of solids can be obtained usingthe current DFT-based computational method augmentedwith phonon modes calculations. We note that our calculatedelectronic properties and phonon dispersions curves �notshown in this paper� for oxides, hydroxides, and carbonatesare also quite close to other similar results reported in theliterature.12,33,35,36 Therefore, our results can be used reliablyfor investigation of the thermodynamic properties of the CO2

capture reactions by these oxides and hydroxides as pre-sented in the next sections.

B. Thermodynamic properties of CO2 capturereactions

1. CO2 capture by alkaline metal oxides

The reaction thermodynamics at 300 K computed fromDFT and phonon calculations are presented in Table II forthe CO2 capture reactions of alkaline metal oxides and hy-droxides. In addition, the heats and free energies of reactionscomputed from the HSC chemistry package at 300 K are alsoreported for comparison.

CaO has been widely used in industry to remove CO2

through carbonation/calcination cycle37 and for this reason,we have used the reaction of CO2 with CaO system as thereference reaction in the analysis of all the other systemsprovided in this study. The heat of reaction ��H�, the van’tHoff P-T plot, and the �� as a function of PCO2

and T for thereaction CaO+CO2=CaCO3 are plotted in Figs. 2�a�–2�c�,respectively. Three different sets of data are plotted in eachgraph. The curves labeled “DFT+phonon” were computed asdescribed above using the phonon density of states informa-tion. The data identified as “DFT-only” were computed byexcluding the phonon free energies and zero point energiesof the solids because computing the phonon contributions iscomputationally very demanding. Finally, the data labeled“Exp. From HSC” were computed from the HSC CHEMISTRY

(a)

(b)

FIG. 1. The calculated phonon related properties for Mg�OH�2, Ca�OH�2,MgCO3, and CaCO3 vs temperature: �a� Phonon free energy; �b� entropy.

TABLE II. The calculated thermodynamic properties of CO2 capture reactions by alkaline earth metal oxides and hydroxides. �EDFT and �Ezp were obtainedat 0 K, while �Hcal, �Gcal, and corresponding experimental data are shown at T=300 K. All energies values are in units of kJ/mol.

Reactions CO2wt %

Calculated thermodynamic properties Expt. dataa

�EDFT �EZP �Hcal �Gcal �H �G

BeO+CO2=BeCO3 129.35 �33.051 3.542 �77.093 �19.320 �43.090 9.565MgO+CO2=MgCO3 109.19 �92.509 8.709 �106.054 �52.666 �100.891 �48.206CaO+CO2=CaCO3 78.48 �161.745 5.523 �176.751 �129.532 �178.166 �130.127SrO+CO2=SrCO3 42.47 �207.344 6.297 �222.290 �172.391 �240.494 �188.850BaO+CO2=BaCO3 28.70 �248.216 6.266 �263.336 �211.895 �272.491 �220.394Be�OH�2+CO2=BeCO3+H2O�g� 102.28 123.911 �8.656 98.338 102.426 8.579 16.532Mg�OH�2+CO2=MgCO3+H2O�g� 75.46 �8.650 �0.051 �26.487 �19.977 �19.665 �12.762Ca�OH�2+CO2=CaCO3+H2O�g� 59.40 �57.950 1.460 �74.012 �70.410 �69.035 �64.033Sr�OH�2+CO2+SrCO3+H2O�g� 36.18 �85.746 �1.986 �104.859 �100.485 �105.424 �97.990Ba�OH�2+CO2=BaCO3+H2O�g� 25.68 �112.048 �1.246 �130.921 �125.411 �121.733 �115.686

aCalculated by HSC CHEMISTRY package �Ref. 32�.

074508-5 Theoretical study on CO2 capture properties J. Chem. Phys. 133, 074508 �2010�

package32 and are assumed to be essentially equal to experi-mental data.

As described in the previous sections, due to the limita-tions of the DFT method for predicting the heats of forma-tion, the calculated energy change for each reaction ��H� hasan offset of about 20 kJ/mol relative to the experimentaldata, as shown in Fig. 2�a�. This is a common issue in many

gas-solid reactions30 and is unlikely to be solved within tra-ditional DFT scheme. In order to obtain good agreement withexperimental data, we found that a quantity of 20 kJ/mol isnecessary to be subtracted from �H for all reactions ana-lyzed here.

From Fig. 2�a�, one can see that DFT-only values repre-sent a poor prediction of the �H values and they show aqualitatively wrong dependence with temperature. In contra-distinction, the DFT+phonon data has about 20 kJ/mol dif-ference and the right curvature along the whole temperaturerange relative to the experimental HSC data. After makingthe empirical correction described above, the predicted re-sults are in good agreement with experimental data and thisfinding is also valid for the other systems described later inthis paper �see Figs. 3 and 5�. Based on these results, it canbe concluded that including the phonon contribution is cru-cial for predicting the correct trends in �H, especially athigher temperatures.

Similarly, in Fig. 2�b�, it is shown that the DFT-onlycalculated van’t Hoff plot has about the right slope, but thereis an offset relative to the HSC data by several orders ofmagnitude in the relative pressure. The DFT+phonon data,however, are in good agreement with the corresponding HSCvalues.

The dependence of the chemical potential for the CO2

capture reaction on the CO2 pressure and temperature isshown in Fig. 2�c�. In this figure, only the equilibrium curve,defined by ��=0 equation, is explicitly plotted. From Fig.2�c�, it is convenient to identify the temperature and pressureranges, where CO2 will be captured by CaO ����0�, re-spectively, where CO2 is released ����0� from CaCO3 andCaO is regenerated. We define the maximum temperature�turnover T� at which CO2 can be captured at a given pres-sure of CO2 as Tmax. For example, at PCO2

=1 bar, Tmax

�1200 K. This value agrees reasonably well with the re-ported experimental operating conditions which indicate thatCaO absorbs CO2 up to 1123 K and is regenerated above1173 K.37,38

The heats of reactions as a function of temperature for

(b)

(a)

(c)

FIG. 2. Comparison of the calculated and experimental thermoproperties forour reference reaction CaO+CO2↔CaCO3. �a� The calculated reactionheats vs temperature without empirical correction; �b� van’t Hoff plot; and�c� counterplotting the variation of the chemical potential vs CO2 pressureand temperature. For a clear representation, only ��=0 curve is plottedexplicitly. On the line ��=0, the reaction is reversible; above the line ���0, CaO absorbs CO2 and reaction goes forward; below the line ���0,CaCO3 releases CO2, and reaction proceeds backward. The pressure valuesin panels �b� and �c� are plotted in logarithmic scale.

FIG. 3. The calculated heats of reactions for the alkaline metal oxides cap-turing CO2. The calculated results and corresponding data from the HSC

package for each reaction were plotted in the same color.

074508-6 Y. Duan and D. C. Sorescu J. Chem. Phys. 133, 074508 �2010�

the CO2 capture reaction by the set of oxides BeO, MgO,CaO, SrO, and BaO, are plotted in Fig. 3, together with thecorresponding experimental data from the HSC package.32

From Fig. 3, it can be seen that the calculated �H values forthe CO2 capture reactions by MgO, CaO, and BaO are veryclose to the experimental data, with differences of less than�5 kJ/mol.4,16 SrO sorbent shows somewhat larger discrep-ancies of about 12 kJ/mol. Among the set of five oxides, BeOhas the largest difference of 15 kJ/mol from experimentaldata. This is probably due to the fact that for BeCO3, wehave not used a crystal structure consistent to the experimen-tal one. As indicated above, this is due to the lack of avail-able experimental data for this structure; instead, we used theSrCO3-based structure for BeCO3, as shown in Table I. Dif-ferences between the SrCO3-based structure of BeCO3 as-sumed here and the real experimental structure might be thereason for the observed larger errors in �H values.

The dependence of the chemical potentials ���� for theCO2 capture reactions by the set of five oxides with the CO2

pressure and temperature are shown in Fig. 4. Overall, it canbe seen that by combining DFT with phonon dynamics cal-culations, very good predictions relative to the available ex-perimental data can be obtained. The major exception is,again, the case of BeO sorbent for which experimental crys-tallographic data for BeCO3 is not available. Our calculatedresults �not indicated here� show that for all cases analyzedin this study, the DFT-only approach predicted higher CO2

pressures at each temperature than the DFT+phonon ap-proach and the experimental data.

From Fig. 4, one can see that up to about 1200 K, SrOand BaO can still absorb CO2. The CO2 pressure can rangedramatically from very low ��10−30 bar� to very high��1010 bar� values despite of the reaction kinetics. Theseresults indicate that during the first half cycle, these oxidesare very good for CO2 absorption �in the ���0 region�.However, as Fig. 4 demonstrates, in the second half cycle ofsorbent regeneration where CO2 is released, the correspond-ing carbonates are very stable and can dissociate only at very

high temperatures where ���0. Overall, these strong sor-bents are not efficient candidate systems for CO2 captureapplications.

From Fig. 4, one can see that BeO can absorb CO2 justabove room temperature with lowest pressure of 10−2 bar,which corresponds well to the �T, P� requirements of post-combustion capture technologies as will be presented later inSec. III C. However, although BeO apparently has a favor-able thermodynamics, this system is not a good candidate forCO2 capture because the health issues associated to beryl-lium powder or dust.

Finally, relative to CaO, MgO appears to be a good can-didate for warm temperature CO2 capture applications be-cause its turnover temperature is around 600 K. Such char-acteristics make this system a potential sorbent candidate forprecombustion technologies.13

2. CO2 capture by solid alkaline earth metalhydroxides

Usually, the alkaline earth metal hydroxides can absorbCO2 and form carbonates according to reaction N�OH�2

+CO2=NCO3+H2O�g� �N=Be,Mg,Ca,Sr,Ba�. The calcu-lated and the experimentally fitted values for �H and �G at300 K for the reactions of alkaline earth metal hydroxidescapturing CO2 are listed in Table II. As one can see, overall,our calculated results are close to the experimental measuredvalues and, in most cases, the corresponding differences areless that 10 kJ/mol. The only exception is Be�OH�2, wherethe calculated results for �H and �G are much higher thanthe fitted experimental values. However, we note that eithercalculated data or the experimental values for �H and �Gare positive within the interested temperature range, whichindicates that BeCO3 is unstable even at room temperatureand dissociates easily to BeO �Refs. 4 and 5� or Be�OH�2

when H2O is available. The relative significant differencesobserved for the case BeCO3 reaction might be due to theunstable nature of this crystal for which no known experi-mental crystal structure is currently available.

The calculated reaction heats for the alkaline earth metalhydroxides absorbing CO2 at different temperatures are de-picted in Fig. 5. For comparison, we also indicate in thesame figure the experimental fitted enthalpy changes ��H� ofthese reactions based on HSC CHEMISTRY data.32 Similar tothe case of alkali hydroxide systems, the observed disconti-nuities in the fitted curves indicate the existence of differentphase transitions which we have not included in the theoretictreatment. From Fig. 5, one can see that overall, our calcu-lated heats of reaction are comparable with the fitted valueswithin a range of about 10 kJ/mol. When moving from Be toBa, it is found that �H values for CO2 capture reactionsbecome more negative, i.e., heat is released, which indicatesthat the reverse reactions become more difficult to take placedue to larger heat requirements.

Finally, the chemical potential ���� for the reactions ofalkaline earth metal hydroxides capturing CO2 can be evalu-ated based on Eqs. �2� and �3�. The dependence of �� on Tand PCO2

/PH2O variables is depicted in Fig. 6, in which �� isrepresented as a contour plot and only ��=0 curve is shownexplicitly. Similar to Fig. 5, the results calculated from HSC

FIG. 4. The contour plots of the calculated chemical potentials vs CO2

pressures and temperatures for the reactions of alkaline earth metal oxidescapturing CO2. Y-axis is plotted in logarithm scale. Only ��=0 curve isshown explicitly. For each reaction, above the ��=0 curve, the hydroxidecan absorb CO2 and the reaction goes forward ����0 region�, whereasbelow the ��=0 curve, CO2 is released and the reaction goes backward toregenerate the hydroxide ����0 region�.

074508-7 Theoretical study on CO2 capture properties J. Chem. Phys. 133, 074508 �2010�

CHEMISTRY package32 are also plotted in Fig. 6. By compar-ing Figs. 5 and 6, one can see that similar to the trendsobserved for the heats of reaction, the calculated �� curvesare very close to the fitted experimental results. The onlymain exception is Be�OH�2 which, as shown above, is aproblematic case due to the unstable nature of the productBeCO3 even at ambient conditions.

From Fig. 6, it can be seen that N�OH�2 �N=Ca,Sr,Ba� still can absorb CO2 at very high temperaturesto form carbonates and the reverse reactions correspondingto N�OH�2 regeneration can occur only at very low CO2

pressures ��10−10 if PH2O=1 bar�. These characteristics in-dicate that these materials do not represent good sorbent can-didates for CO2 capture due to the difficult conditions re-quired for sorbent regeneration. As discussed above,Be�OH�2 is also not a good candidate due to the unstablenature of the BeCO3. Based on this analysis, it can be con-cluded that only Mg�OH�2 represents a good candidate forCO2 capture applications. As indicated in Fig. 6, at

PCO2/PH2O=1, this system can absorb CO2 up to 750 K.

Above 750 K and in the presence of H2O, MgCO3 dissoci-ates and leads to regeneration of Mg�OH�2. This prediction isin good agreement with experimental findings.7,9 From Fig.6, one can also see that regenerating Mg�OH�2 during therelease of CO2 cycle can be done at higher temperature andhigh pressures of CO2 �30 bar or even higher�. This pressurerange is efficient for practical sequestration applications be-cause using heat to compress CO2 is a much cheaper alter-native than using electricity. These energy savings have beenalso pointed out before by Siriwardane and Stevens.7

Similar to alkali metal hydroxides, when alkaline earthmetal hydroxides absorb CO2 and form carbonates, the ac-tual CO2 pressure depends on the steam pressure in the re-actor �see Fig. 6�. When the steam pressure is increased��1 bar, for example�, the PCO2

will be decreased as theratio of PCO2

and PH2O shown in Fig. 6 is fixed, and viceversa.

By comparing the reactions of the alkaline earth metaloxides to capture CO2 with those of their corresponding hy-droxides, it can be seen that at low temperatures, the �Hvalues of the former reactions are more negative than the �Hof the hydroxide reactions, which means the oxides are muchstronger to bind with CO2 than the corresponding hydrox-ides. When oxides absorb CO2, the entropies are decreased��S�0� rapidly. As for hydroxides, when they absorb CO2

to form carbonates, the entropy changes of the reactions donot change much because they also release H2O gas. There-fore, the entropy changes ��S� of the CO2 capture reactionsby oxides are much larger than the cases of the correspond-ing hydroxides. When the temperature is increased, thechemical potential change �which in this case is the same asthe Gibbs free energy� for CO2 capture reactions by oxideschange rapidly relative to the case of the corresponding hy-droxides as shown in Figs. 4 and 6, resulting in lower turn-over temperatures of the oxide sorbents than the correspond-ing hydroxides.

3. Applications to precombustion and postcombustionCO2 capture technologies

Under precombustion conditions, after water-gas shift-ing, the gas stream mainly contains CO2, H2O, and H2. Thepartial CO2 pressure in this case is around 20–25 bar and thetemperature is around 300–350 °C. For precombustion tech-nologies, the programmatic goal put forward by the DOE isto capture at least 90% CO2 with an increase in the cost ofelectricity of no more than 10%.39 In order to minimize theenergy consumption, the ideal sorbents should work in theabove indicated pressure and temperature ranges to separateCO2 from H2. In this work, we have estimated the turnovertemperatures for precombustion conditions of ten capture re-actions. These temperatures, denoted with T1 and listed inTable III, represent the values above which the sorbents can-not absorb CO2 and the release of CO2 starts. This indicatesthat during the first half cycle of the CO2 capture process, theoperating temperature should be lower than T1, whereas theoperating temperature may be higher than T1 �depending on

FIG. 5. The calculated heats of reactions ��H� for the alkaline earth metalhydroxides capturing CO2. The calculated results and corresponding datafrom the HSC package for each reaction were plotted in the same color.

FIG. 6. The contour plots of the calculated chemical potentials vs CO2 andH2O pressures and temperatures for alkaline earth metal hydroxides captur-ing CO2. Y-axis is given in logarithmic scale. Only ��=0 curve is shownexplicitly. For each reaction, above the ��=0 curve, the hydroxide absorbCO2 and the reaction goes forward ����0 region�, whereas below the��=0 curve, CO2 is released and the reaction goes backward to regeneratethe hydroxide ����0 region�.

074508-8 Y. Duan and D. C. Sorescu J. Chem. Phys. 133, 074508 �2010�

the desired CO2 pressure� during the second half cycle whenthe sorbents are regenerated and CO2 is released.

For postcombustion conditions, the gas stream mainlycontains CO2 and N2. In this case, the partial pressure of CO2

is around 0.1–0.2 bar, while the temperature range can varysignificantly. In the current analysis, however, we will con-sider the case of low-temperature capture ��473 K�. In thiscase, the DOE programmatic goal for postcombustion tech-nologies is to capture at least 90% CO2 with an increase inelectricity cost of no more than 35%.39 The correspondingturnover temperatures �denoted as T2� for postcombustioncapture by alkaline earth metal oxides and hydroxides arealso listed in Table III.

From Table III, it can be seen that the turnover tempera-tures �T1,T2� for SrO, BaO, Ca�OH�2, Sr�OH�2, andBa�OH�2 are all over 1000 K. This means that these oxidesand hydroxides are very good for absorbing CO2, but theircorresponding carbonates are very stable and can be dissoci-ated to release CO2 only at very high temperatures. There-fore, none of these systems represent good sorbents for CO2

capture under either precombustion or postcombustion con-ditions. CaO can work at high temperatures �around 1030 K�for CO2 capture under postcombustion conditions �PCO2=0.1 bar�. These findings agree with the experimental re-sults, which indicate that at PCO2

=1 bar, CaO absorbs CO2

up to 1123 K and is regenerated above 1173 K.37,38 Compar-ing with CaO, from the list of alkaline earth metal oxides andhydroxides analyzed here �see Table III�, only BeO, MgO,Be�OH�2, and Mg�OH�2 can work at lower temperatures forCO2 capture and release. As mentioned in Sec. III B, al-though BeO apparently has a favorable thermodynamics, thissystem is not a viable candidate for CO2 capture because ofthe health issues associated to beryllium powder or dust.Similarly, Be�OH�2 is not a good candidate for CO2 capturedue to its low operating temperature ��300 K�.

Our results show that MgO could be used for both pre-combustion and postcombustion capture technologies due toits low regenerating temperature �T2=560 K for postcom-bustion conditions and T1=720 K for precombustion condi-tions� which are close to experimental findings.40 However,Mg�OH�2 can only be used for postcombustion capture tech-

nologies with a turnover T2=720 K because its turnovertemperature �T1� is very high, outside the temperature rangeof interest for precombustion applications.

Among the list of alkaline earth metal oxides and hy-droxides analyzed in this study, comparing with CaO, onlyMgO and Mg�OH�2 are found to be good sorbents for CO2

capture. Upon absorption of CO2, both of these two systemscan form MgCO3. However, the regeneration conditions ofthe original systems can take place at different conditions asindicated in Fig. 7. In this case, we present the calculatedphase diagram of MgO–Mg�OH�2–MgCO3 system at differ-ent CO2 pressures and under several fixed PH2O values �0.01,0.1, 1.0, and 10.0 bar, respectively�. From Fig. 7, it can beseen that when H2O is present and at low temperatures,MgCO3 can release CO2 to form Mg�OH�2 instead of form-ing MgO. For example, at PH2O=0.01 bar, only for tempera-tures under the transition temperature �Ttr� 420 K, MgCO3

can be regenerated to form Mg�OH�2. By the increase in theH2O pressure, the transition temperature is increased. Asshown in Fig. 7, when PH2O is increased to 10 bar, the cor-responding Ttr=600 K. Above Ttr, MgCO3 is regenerated toMgO. Therefore, when water is present in the sorption/desorption cycle, no matter whether the initial sorbent isMgO or Mg�OH�2, and for temperatures below Ttr, the CO2

capture reaction is dominated by the process Mg�OH�2

+CO2↔MgCO3+H2O�g�, whereas above Ttr the CO2 cap-ture reaction is given by MgO+CO2↔MgCO3. The reasonis that between MgO and Mg�OH�2, there is a phase transi-tion reaction MgO+H2O�g�=Mg�OH�2 happening at thetransition temperature Ttr. Obviously, by controlling the pres-sure of H2O as shown in Fig. 7, the capture CO2 temperature�T swing� can be adjusted. However, more water in the sor-bent system will cost more energy due to its sensible heat;there should be a trade-off to balance them in the practicaltechnology.

C. A hierarchical computational approach for rapidscreening CO2 solid sorbents

In this section, we outline a computational hierarchy forscreening a large number of candidates for suitability as CO2

TABLE III. The turnover temperatures �T1,T2� of the alkaline metal basedsorbents capturing CO2 under precombustion and, respectively, postcombus-tion conditions, assuming PH2O=1 bar.

ReactionsPrecombustion T1

�K�Postcombustion T2

�K�

BeO+CO2=BeCO3 450 370MgO+CO2=MgCO3 690 540CaO+CO2=CaCO3 1350 1010SrO+CO2=SrCO3 hTa 1200BaO+CO2=BaCO3 hT 1400Be�OH�2+CO2=BeCO3+H2O�g� �300 �300Mg�OH�2+CO2=MgCO3+H2O�g� hT 600Ca�OH�2+CO2=CaCO3+H2O�g� hT hTSr�OH�2+CO2+SrCO3+H2O�g� hT hTBa�OH�2+CO2=BaCO3+H2O�g� hT hT

ahT acronym means that the turnover temperature exceeds the temperaturerange of practical interest �1500 K�.

FIG. 7. The calculated phase diagram of MgO–Mg�OH�2–MgCO3 systemvs the CO2 pressure at fixed PH2O=0.01, 0.1, 1.0 and 10.0 bar. For eachPH2O, only Mg�OH�2 can be regenerated from MgCO3 for temperatures un-der the transition values �Ttr�. Above Ttr values, only MgO can be obtained.

074508-9 Theoretical study on CO2 capture properties J. Chem. Phys. 133, 074508 �2010�

sorbents for either precombustion or postcombustion captureapplications. We have shown that the predicted heats of re-actions from DFT-only calculations are higher than the ex-perimental measured values by about 20–30 kJ/mol at lowtemperatures and that the curvature of the correspondingcurves as a function of temperature is often not consistentwith the experimental one �see Fig. 2�a��. From Fig. 2�c�,one can see that along a wide range of temperatures, theDFT-only predicted chemical potentials of the reactions arealso higher than either the experimental data or the resultspredicted using both the DFT and the phonon free energycalculations. As described in Sec. II, phonon dispersion cal-culations have to be done with large supercells and this stepis computationally intensive. In order to speed up the screen-ing process, it is reasonable to use the DFT-only energiesonly as a prescreening tool with appropriate error bars. Inparticular, we have taken the CaO system �Fig. 2� as ourreference capture reaction. If the calculated heat of reactionfor a test material is lower than the heat of reaction for thereference system, then we include this solid in the list ofso-called “good” candidates. In this way we can rapidlyeliminate a large number of reactions at minimal computa-tional cost. We then further refine the screening process byperforming phonon calculations for materials belonging onlyto the list of good candidates. A schematic of our searchscheme is presented in Fig. 8. This scheme is based on a setof computational filters summarized below.

Filter �0�: For each solid, we first conduct basic screen-ing based on acquisition of general data, such as the weightpercent of absorbed CO2 in the assumption of the completereaction, materials safety, materials cost, etc. We also includewhere available the thermodynamic data from literature andfrom general software package, such as HSC CHEMISTRY,FACTSAGE, etc. If the necessary data for evaluation of thethermodynamic properties exists, then the use of DFT calcu-lations is not necessary. Otherwise, if the material passesbasic screening, continue to the next step.

Filter �i�: Perform DFT calculations for all compounds inthe candidate reaction with this solid. If ��EDFT−�Eref� /n1

�20 kJ /mol, where n1 is CO2 molar number in reaction �1�and �Eref is the DFT energy change for the reference capturereaction �e.g., CaO+CO2=CaCO3�, we add this compoundto the list of good candidates. Otherwise, we go back to filter�0� and pick another solid.

Filter �ii�: Perform phonon calculations for reactant andproduct solids to obtain the corresponding zero point ener-gies and the phonon free energies for the list of good candi-dates. Specify the target operating conditions �temperature,partial pressures of CO2 and H2O� and compute the changein chemical potential for the reaction, namely, ���T,P� fromEqs. �2� and �3�. If ���T,P� is close to zero �e.g.,����T,P���5 kJ /mol� at the operating conditions, then weselect this reaction as a member of the “better” list. Only ashort list of compounds will likely be left after application offilter �ii�.

Filter �iii�: Additional modeling could be performed torank the remaining short list of better candidates. One is thekinetics of the capture reactions, which could be done bytransport and diffusion calculations as well as experimentalmeasurements. Another necessary and doable modeling taskis the behavior of the solid in the reactor, which can be doneby CFD methods based on finite element method approach.These simulations are currently underway. Application ofthese screening filters will ensure that only the most promis-ing candidates will be identified for the final experimentaltesting.

This screening methodology provides a path for evaluat-ing materials for which experimental thermodynamic dataare unavailable. One area where this approach could be usedto great advantage is in evaluating mixtures and doped ma-terials, where thermodynamic data are generally not avail-able but for which the crystallographic structure is known orcan be easily determined.

IV. CONCLUSIONS

We have demonstrated that first-principles density func-tional theory can be used in combination with phonon den-sity of states calculations to provide a reasonable estimate ofthe thermodynamics of CO2 capture reactions involving al-kaline earth metal oxides, hydroxides, and carbonates. Dueto the limitations of the DFT method used, our calculatedheats of reaction ��H� for each solid capturing CO2 is about20 kJ/mol higher than the corresponding experimental datafrom the HSC CHEMISTRY software package.

Comparing with CaO, a widely used industrial sorbent,among the alkaline earth metal oxides and hydroxides, wefound that MgO and Mg�OH�2 are better sorbents for bothpostcombustion and precombustion CO2 capture becausethey require less regenerating amounts of heat and can workat lower temperature. When MgO is used as CO2 sorbent andin the presence of steam, or when Mg�OH�2 is used as a CO2

sorbent, both of these systems will lead upon CO2 absorptionto formation of MgCO3. However, the regenerated sorbentscan be different under different experimental conditions.Generally, with high steam pressure, the Mg�OH�2 will be

FIG. 8. Our modeling scheme for screening solid CO2 sorbents.

074508-10 Y. Duan and D. C. Sorescu J. Chem. Phys. 133, 074508 �2010�

mainly regenerated at low temperatures. Above certain tem-peratures �depending on the pressures of steam and CO2�,only MgO could be regenerated.

We have outlined a general screening method that can beuseful for assessing potential sorbents for CO2 capture. Thisscreening approach is particularly useful in those cases whenno thermodynamic data exist. This methodology involves ap-plying a series of screening filters of increasing computa-tional complexity in order to identify the optimal systemsthat can perform under specific operating conditions, consis-tent to either precombustion or postcombustion CO2 capturetechnologies. Such a scheme has the potential to identify anumber of promising solid sorbent candidates that can beexplored experimentally for reversible CO2 capture applica-tions.

ACKNOWLEDGMENTS

This work was performed in support of the National En-ergy Technology Laboratory’s Office of Research and Devel-opment under Contract No. DE-FE-0004000 with activitynumber 4000.2.660.241.001. One of us �Y.D.� thanks Dr. S.Chen, Dr. J. K. Johnson, Dr. Y. Soong, Dr. H. W. Pennline,and Dr. R. Siriwardane for fruitful discussions.

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