Carbon dioxide adsorption on oxide nanoparticle surfaces

Carbon dioxide adsorption on oxide nanoparticle surfaces

Chemical Engineering Journal 170 (2011) 471–481 Contents lists available at ScienceDirect Chemical Engineering Journal journal homepage: www.elsevie...

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Chemical Engineering Journal 170 (2011) 471–481

Contents lists available at ScienceDirect

Chemical Engineering Journal journal homepage: www.elsevier.com/locate/cej

Carbon dioxide adsorption on oxide nanoparticle surfaces Jonas Baltrusaitis a , Jennifer Schuttlefield b , Elizabeth Zeitler a , Vicki H. Grassian a,c,∗ a

Department of Chemistry, University of Iowa, Iowa City, IA 52246, United States Department of Chemistry, University of Wisconsin Oshkosh, Oshkosh, WI 54901, United States c Department of Chemical and Biochemical Engineering, University of Iowa, Iowa City, IA 52246, United States b

a r t i c l e

i n f o

Article history: Received 20 September 2010 Received in revised form 11 December 2010 Accepted 14 December 2010 Keywords: Metal oxide nanoparticles Carbon dioxide Adsorption

a b s t r a c t In this paper, focused on environmental nanotechnology, we review some recent results for carbon dioxide (CO2 ) adsorption on hydroxylated Fe2 O3 , ␥-Al2 O3 , and TiO2 nanoparticle surfaces at 296 K as followed by transmission FTIR spectroscopy. In the absence of water vapor (<1% relative humidity, RH), following exposure to CO2 different species formed on the oxide surface due to the presence of adsorption sites with different basicities. While the majority surface species on Fe2 O3 , ␥-Al2 O3 is determined to be adsorbed bicarbonate, on TiO2 nanoparticles bidentate carbonate was more prevalent. A carboxylate species was observed on TiO2 nanoparticles under dry conditions as well. When water is present at 40% RH, the nature of the adsorbed CO2 species changed to that of solvated carbonate formation in the adsorbed water layer. Observed initial adsorption rates were calculated from time–course experiments under dry conditions and in the presence of 40% RH. When initial adsorption rates were compared between dry and wet experiments, a larger value was found for dry experiments suggesting that CO2 molecules have to compete for adsorption sites with water on these nanoparticle surfaces. As discussed here, quantum chemical calculations provide some additional insights into CO2 adsorption on hydroxylated metal oxide surfaces in the presence and absence of molecularly adsorbed water. © 2010 Elsevier B.V. All rights reserved.

1. Introduction Environmental catalysis involving oxide nanoparticles is of great interest in treating gaseous emissions and reducing waste by-products [1]. In the atmosphere, carbon dioxide concentrations are increasing which significantly impacts the Earth’s climate and presents itself as a major environmental problem [2–4]. Due to human activities, such as the combustion of fossil fuels, land-use and deforestation, global carbon dioxide annual emissions have grown between 1970 and 2004 an estimated 80% [5]. Atmospheric concentrations of CO2 , at approximately 379 ppm, currently exceed the natural range over the last 650,000 years [5]. As one of many CO2 capturing techniques at large point sources, chemical looping combustion (CLC) utilizes metal oxide particles as oxygen carriers [6]. Furthermore, carbon dioxide is important in the industrial production of several commercial products via heterogeneous catalysis including methanol and dimethyl carbonate [7,8]. Converter catalysts for these reactions are also oxide based; hence, metal oxide catalyst and CO2 interactions are of utmost importance in heterogeneous environmental catal-

∗ Corresponding author at: Department of Chemistry, University of Iowa, Iowa City, IA 52246, United States. Tel.: +1 319 335 1392; fax: +1 319 335 1270. E-mail address: [email protected] (V.H. Grassian). 1385-8947/$ – see front matter © 2010 Elsevier B.V. All rights reserved. doi:10.1016/j.cej.2010.12.041

ysis. Additionally, oxide surfaces are important environmental interfaces [9–12] and ubiquitous in nature. Thus, understanding of the chemistry that occurs at the oxide surface in the presence of carbon dioxide is necessary if chemical processes of great importance to the environment are to be fully understood. While adsorption of CO2 on hydroxylated metal oxide nanoparticle surfaces has been well studied under dry conditions, the adsorption of CO2 at the adsorbed water–metal oxide interface under ambient conditions has not. In earlier work, we have previously shown that a thin water layer on an iron oxide nanoparticle surface plays an important role in the surface chemistry of carbon dioxide [13]. Furthermore, it has been shown that presence of relative humidity greatly enhances the amount of adsorbed CO2 even at relatively low temperatures on potassium carbonate impregnated MgO support, although the exact mechanism is not well understood [14,15]. As a step towards better understanding of the molecular level details at the adsorbed water–metal oxide nanoparticle interface under ambient conditions of pressure, temperature and relative humidity, the adsorption of carbon dioxide on Fe2 O3 , ␥-Al2 O3 , and TiO2 nanoparticles is reported here at 296 K utilizing using transmission FTIR spectroscopy. Additionally, kinetic measurement data are needed to assess the effect of co-adsorbed water layer on the rate of CO2 uptake on these nanoparticle surfaces.

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2. Experimental methods 2.1. Transmission FTIR spectroscopy Transmission FTIR spectroscopy was used to investigate the adsorption of CO2 on oxide nanoparticle surfaces. Infrared spectra were collected using a single beam FTIR spectrometer, either a Mattson research series or Mattson infinity gold, equipped with a liquid nitrogen-cooled narrowband mercury cadmium telluride (MCT) detector. Typically, a total of 250 scans were acquired at an instrument resolution of 4 cm−1 over the spectral range from 800 to 4000 cm−1 . The spectrometer was purged with a commercially available air dryer (Balston 75-62) which minimized H2 O and CO2 concentrations in the purge air. An infrared cell (7 cm × 7 cm × 7 cm (length × width × height) on the outside, with a total inner volume of 197 ± 2 mL) made from a stainless steel cube with two BaF2 windows was placed in the sample compartment of the spectrometer. The stainless steel jaws were loaded into the cell with a tungsten grid for sample mounting. The half of the grid coated with the oxide powder to yield spectral features associated with the oxide surface and surface adsorbates as well as gas-phase spectral features, whereas the uncoated half of the grid only yielded information about the gas phase. Approximately 5–15 mg of oxide powder was pressed onto half of a tungsten grid. The other half was left uncoated and was used to measure the infrared spectrum of the gas phase. The infrared cell was connected by a Teflon tube to a gashandling system, which consists of a glass manifold with four ports connected to two absolute pressure transducers (MKS instruments) that operate in two different ranges from 0.001 to 10.00 Torr and from 0.1 to 1000 Torr. The gas manifold was connected to a turbo molecular pump that is used to evacuate the system down to a final pressure of around 1 × 10−6 Torr. Single beam spectra of the oxide powder in the presence of the gas or after exposure to a reactant gas are all referenced to those prior to exposure in order to provide absorbance spectra of both the surface species and the gas phase species present. All IR spectra were recorded at 296 K.

2.2. Metal oxide nanomaterials: sources and characterization Commercially available metal oxide nanoparticles used in these experiments were purchased from several sources. Amorphous Fe2 O3 was purchased from Alfa Aesar whereas ␥Al2 O3 and TiO2 were purchased from Degussa, AlumOxid C and P25, respectively. These nanoparticles were characterized with several techniques including transmission electron microscopy (JEOL 2100F instrument operating at 200 kV), powder X-ray diffraction (Bruker D-5000 q – q diffractometer equipped with a Kevex energy-sensitive detector) and BET surface area measurements (Quantachrome Nova 4200e specific surface area and pore size analyzer). Average values reported for BET surface area are a result of triplicate measurements.

2.3. Sources and purity of gases and liquids Carbon dioxide was purchased from Air Products (CP grade) and distilled H2 O (Optima grade) was purchased from Fisher Scientific. Water vapor for FTIR in situ experiments was taken from the headspace of the water half-filled flask that had been degassed prior to use.

Fig. 1. Transmission electron images of (a) Fe2 O3 , (b) ␥-Al2 O3 and (c) TiO2 nanoparticles.

3. Results and discussion 3.1. Nanomaterial characterization Transmission Electron Microscopy (TEM) images of metal oxide nanoparticles are shown in Fig. 1. Analysis of 200 particles yield an average particle sizes of 5 ± 1, 13 ± 2 and 23 ± 3 nm for the Fe2 O3 , ␥-Al2 O3 and TiO2 nanoparticles. BET surface areas were determined to be 195 ± 10, 93 ± 5 and 51 ± 1 m2 /g for Fe2 O3 , ␥-Al2 O3 and TiO2 nanoparticles, respectively. X-ray diffraction patterns are shown in Fig. 2 for the three different samples. Although, Mossbauer spectroscopy indicated that Fe2 O3 nanoparticles contained some hematite and ferrihydrate [16], the X-ray diffraction results showed the amorphous nature of the iron oxide nanoparticle samples. The presence of 100% crystalline phases in the ␥-Al2 O3 sample was confirmed. X-ray diffraction pattern of TiO2 indicated that the

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Fig. 2. XRD patterns of (a) Fe2 O3 , (b) ␥-Al2 O3 and (c) TiO2 nanoparticles.

sample contained two different polymorphs, anatase (∼85%) and rutile (∼15%). Table 1 summarizes the characterization data for the three nanoparticle samples used in this study. 3.2. Analysis and assignments of the infrared spectra for adsorbed CO2 Infrared spectroscopy has long been used to investigate the nature of CO2 adsorption products on oxide surfaces [17]. In combination with metal coordination compounds, it has become customary to assign adsorbed CO2 product structures by comparison to the infrared data of inorganic or organometallic compounds. This approach has allowed for considerable amount of the data to be tabulated and reported [7,17–19]. Very recently, however, it has been also shown that this approach can be misleading and can lead to misassignments [20]. Nitrate ion, NO3 − , (isostructural with the adsorbed carbonate, CO3 2− ) adsorbed on Al2 O3 and BaO surfaces exhibited similar vibrational frequencies for both mono- and bridged configurations. In particular, nitrate monodentate formed on surface oxygen sites and two metal site bridged nitrate exhibited vibrational bands at ∼1300 and 1450 cm−1 , albeit at different doublet intensities due to the dynamic dipole moment [20,21]. Only DFT calculations allowed for an unambiguous absorption band identification which is in agreement with the work described here. Thus, a combination of experimental and theoretical methods (DFT frequency calculations combined with FTIR [21,22] or internal reflection absorption spectroscopy (IRAS) [20]) are mandatory for accurate absorption band assignments. A number of different species have been identified on oxide surfaces in the presence of gas-phase carbon dioxide including linearly adsorbed CO2 , carbonate, bicarbonate and carboxylate species. Linearly adsorbed CO2 exhibits 3 asymmetric stretch in the spectral region of 2285–2410 cm−1 , close to its fundamental gas value

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of 2349 cm−1 [17]. For adsorbed carbonate, it has been shown that the fundamental 3 asymmetric stretch at 1415 cm−1 , doubly degenerate in the gas phase, splits upon coordination to the oxide surface [17,23]. These two new absorptions are often referred to as 3 (O–C–O)a and 3 (O–C–O)s . The frequency difference upon coordination, 3 , is a measure of the interaction with the surface. Thus the measured frequencies and their differences can be used to assign coordinated carbonate to mono-, bi- and bridged configurations (3 = <100 cm−1 , 100 < 3 < 300 cm−1 and 3 > 400 cm−1 , respectively) [17]. As will be discussed, quantum chemical calculation can be used to assign absorption bands and get additional information about molecular structure of surface species (vide infra). The presence of adsorbed bicarbonate can be unambiguously identified by the presence of the absorption band around 3620 cm−1 due to the OH stretch. This is in contrast to the bicarbonate solid state dimmer form hydroxyl group stretching frequency, which is shifted to the 2450–2620 cm−1 region [17]. Additionally, three very distinct vibrational bands can be observed below 1300, 1370–1400 and 1615–1655 cm−1 due to the OH bending modes, symmetric and asymmetric, respectively [17]. Importantly, bent CO2 species or carboxylate, can form via occupation of 2␲u molecular orbital which results in O C O bending [7]. This involves electron donation from the surface and results in a set of bands in the 1200–1700 cm−1 region located close to those of bridged carbonate. This suggests those can be ambiguous in assignments which can only be resolved using quantum chemical methods. Recently, it has become possible to accurately calculate vibrational frequencies of CO2 adsorption products on metal centers using quantum chemical methods [16,21,24–28]. This route offers flexibility obtaining vibrational frequencies of structures that may form on surfaces and in the presence of co-adsorbates. Frequencies need to be calculated for optimized molecular structures to insure a minimum on the potential energy surface. Assumptions that include the use of the harmonic potential energy functions to calculated vibrational frequencies via mass weighting and diagonalization of the force constant matrix warrants the need for scaling factors that will vary with different levels of theory. Empirically determined scaling factors, calculated within a certain level of theory and basis set were shown to be accurate to two significant digits [29]. Pure and hybrid functional tend to yield calculated vibrational frequencies closer to those determined experimentally with scaling factors of 0.96–0.99, whereas introduction of Hartree Fock exchange or correlated MP2 methods yield frequencies that need to be scaled by a factor of 0.90–0.94 [29]. Baltrusaitis et al. used scaling factor of 0.9506 as determined for gas phase bicarbonate ion at B3LYP/6-31+G(d) level of theory to obtain agreement between experimental and calculated bicarbonate ion vibrational frequencies within few wavenumbers [21]. Combined theoretical and experimental studies of 18 O isotope incorporation into the adsorbed molecule can be performed thus providing additional insight on the nature and source of the oxygen atoms in adsorbed CO2 products [16]. Finally, CO2 adsorption products have adsorption overlapping bands in the 1700–1200 cm−1 region and peak fitting is often necessary to accurately distinguish the adsorption band positions and compare them with calculated ones [21].

Table 1 Characterization of Fe2 O3 , ␥-Al2 O3 and TiO2 nanoparticles. Nanoparticles

Source

BET surface area, m2 /ga

Average particle size, nmb

Phase

Fe2 O3 ␥-Al2 O3 TiO2

Alfa Aesar Degussa Degussa

195 ± 10 93 ± 5 51 ± 1

5±1 13 ± 2 23 ± 3

Amorphous Gamma Anatase (∼85%) and rutile (∼15%)

a b

Triplicate measurements. Measured from TEM images of 200 particles.

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Fig. 3. Transmission spectra recorded after CO2 adsorption on Fe2 O3 nanoparticles as a function of increasing pressure. Spectral regions of 2600–4000 cm−1 – hydroxyl groups – (top) and 1000–1800 cm−1 – adsorbed carbonate (bottom) – are shown. Gas-phase absorptions have been subtracted from the spectra.

Fig. 4. Transmission spectra recorded after CO2 adsorption on ␥-Al2 O3 as a function of increasing pressure. Spectral regions of 2600–4000 cm−1 – hydroxyl groups – (top) and 1000–1800 cm−1 – adsorbed carbonate (bottom) – are shown. Gas-phase absorptions have been subtracted from the spectra.

3.3. Transmission FTIR spectroscopy of CO2 adsorption on metal oxide nanoparticles under dry conditions

present is the spectra is located at 1040 cm−1 and is associated with the monodentate form of carbonate, specifically the 1 stretch. As shown in Baltrusaitis et al. [21], the monodentate form of carbonate was shown to be present in this complex region at 1383 and 1446 cm−1 through deconvolution of the FTIR spectra. These peaks could not be observed here, but are likely concealed due to spectral overlap. Surface hydroxyl group information for Fe2 O3 nanoparticles can also be obtained from Fig. 3. The absorption band at 3619 cm−1 is assigned to (OH) that occurs once the bicarbonate product has formed on the surface. The loss in intensity, as indicated by the negative peaks in the spectra at 3690 and 3665 cm−1 shown in Fig. 3, is assigned to the loss of OH groups bonded to two metal atoms and shows the likely involvement of these OH groups in the formation of surface bicarbonate [21]. A broad absorption band at 3370 cm−1 is due to the increasing hydrogen bonding between the CO2 adsorption products and hydroxyl groups on the Fe2 O3 nanoparticle surface. This shows the complexity of surface product interactions where adsorbed hydroxyl groups are not only involved in the formation of surface bicarbonate, but also in proton donating interactions via hydrogen bond induced proton delocalization. The vibrational frequencies observed are summarized in Table 2. For CO2 adsorption products on ␥-Al2 O3 shown in Fig. 4, spectra were taken at pressures of 51, 97, 185, 280, 498, and 915 mTorr. Conceptually, similar spectra were produced compared to those of the Fe2 O3 nanoparticles. The vibrational frequencies of the bands observed in the spectra are labeled and assigned in Table 2. The vibrational frequencies for the four modes of adsorbed bicarbon-

Transmission FTIR spectra of Fe2 O3 nanoparticles exposed to increasing pressures of CO2 are shown in Fig. 3. Gas-phase absorptions have been subtracted from each of the spectra. Spectra are shown for the pressures of 100, 274, 525, 691, and 891 mTorr. Upon adsorption of CO2 on the Fe2 O3 nanoparticle surface, new broad complex absorption bands appear between 1000 and 1800 cm−1 and grow in intensity as the carbon dioxide pressure increases. The formation of carbonate and bicarbonate was previously observed in this region when metal oxides were exposed to CO2 and adsorbed speciation found to be different depending on the metal oxide surface [13]. Based on recent quantum chemical calculations and FTIR isotope experiments [17,21,30–32], the absorption bands labeled in the spectra are assigned to adsorbed bicarbonate and carbonate species. For Fe2 O3 nanoparticles, the absorptions at 1222, 1410, 1629 and 3619 cm−1 are specifically assigned to the ı4 (COH), 3 (OCO)s , 2 (OCO)a , and 1 (OH) vibrational modes of the adsorbed bicarbonate. Another band can be assigned to bicarbonate at 1010 cm−1 for the 5 (C–OH) stretch. The four remaining vibrational bands; ı8 (CO2 )o.o.p. , ı6 (OCO), 9 (H torsion) and 7 (OC–(OH)), for bicarbonate are below 850 cm−1 . In addition to the bicarbonate, carbonate bands are present in the spectra Fig. 3 as broad overlapping bands observed in the spectral region from ∼1200 to 1700 cm−1 . These are observed at 1316, 1553 and 1590 cm−1 and have been assigned to the bidentate form of adsorbed carbonate [13,17,21]. Another peak observed to be

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Table 2 Assignment of vibrational frequencies adsorbed products following the adsorption of CO2 on Fe2 O3 , ␥-Al2 O3 and TiO2 nanoparticles in the spectral range 1000–4000 cm−1 . 0.274 Torr CO2 Vibrational mode assignments Bicarbonate (HCO3 − ) ␯1 (OH) ␯2 (O–C–O)a ␯3 (O–C–O)s ␦4 (COH) Carboxylate (CO2 − ) ␯3 (O–C–O)a ␯3 (O–C–O)s Carbonate (CO3 2− ) Monodentate ␯3 (O–C–O)a ␯3 (O–C–O)s ␯1 (C–O) Bidentate ␯3 (O–C–O)a ␯3 (O–C–O)s ␯1 (C–O) a b

Fe2 O3 (cm−1 )

␥-Al2 O3 (cm−1 )

TiO2 (cm−1 )

Literature referencesa

3619 1629 1410 1222

3620 1650 1435 1228

n.o.b 1623 1423 1222

3600–3627 1555–1671 1396–1500 1220–1269

n.o.b n.o.b

n.o.b n.o.b

1670 1251

1670 1245

n.o.b n.o.b 1040

n.o.b 1383 1040

1510 1392 1076

1446–1590 1380–1395 1040

1590,1553 1316 n.o.b

1644 1355 1010

1566 1355 n.o.b

1535–1670 1243–1355 1015

Range taken from Refs. [17,21,30–32,37,48]. n.o.: not observed

ate, ı4 (COH), 3 (OCO)s , 2 (OCO)a and 1 (OH) are observed at 1228, 1435, 1650 and 3620 cm−1 , respectively. Also present for ␥-Al2 O3 in the spectral region of ∼1200 to 1700 cm−1 are other absorption bands, located at 1355, 1383 and 1644 cm−1 . These absorptions are due to the monodentate and bidentate forms of adsorbed carbonate. As in the case of Fe2 O3 nanoparticles shown in Fig. 3, a carbonate vibrational band is also seen at 1040 cm−1 , which arises from the symmetric C–O stretches for the monodentate structures of carbonate, and thus assigned to the 1 vibrational mode. Hydrogen bonded hydroxyl group interactions for ␥-Al2 O3 are present and can be seen due to the broad band at 3400 cm−1 . These are much stronger in intensity when compared to those of Fe2 O3 nanoparticles. This indicates an increase in potential hydrogen bonding interactions between the surface hydroxyl group and the adsorbates. In the 3700 cm−1 region, losses are seen for ␥-Al2 O3 at 3700 and 3750 cm−1 due to the loss in intensity of the hydroxyl groups coordinated to metal atoms upon the adsorption of CO2 and as in the case of iron oxide, indicates an involvement in the formation of the bicarbonate product. While adsorbed carbonate absorption band assignments have been addressed extensively in other publications [17], adsorbed bicarbonate coordination modes deserve additional attention. Bidentate and monodentate bicarbonates have previously been observed in organometallic Rh complexes [33–35] and bidentate bicarbonate has been proposed as an adsorbed species on zirconia [36]. Additionally, surface bicarbonate formation in the literature has been regarded to as “CO2 insertion into hydroxyl bond” or “nucleophilic attack of OH” with two competing mechanisms and resulting final structures (see Baltrusaitis et al. [21] and references therein). Quantum chemical calculations performed to investigate different adsorbed bicarbonate structures on aluminum oxide clusters provided additional insights into the pathways and final structure of the adsorbed bicarbonate [21]. B3LYP/6-31+G* energy minimized structures of several clusters models (I–V) are shown in Fig. 5. These correspond to mono- and bidentate formed during the nucleophilic attach of OH, reorientated mono- and bidentate and bridged structures. Calculated frequencies were compared to the experimentally obtained ones after CO2 adsorption on nanoparticulate Fe2 O3 and ␥-Al2 O3 of 16 O, 18 O isotope and deuterated surfaces. Interestingly, the best agreement for the cluster model V was found with the experimental data. These data also showed that the proton of the surface hydroxyl group can migrate to an oxygen atom of the CO2 molecule. Direct proton transfer was shown to have transition state energy of 25.78 kcal/mol at MP2/6-311+G(d,p)

level with the final structure V being energetically favored by −13.57 kcal/mol with respect to the initial structure II [21]. These data provided additional insight into the molecular level mechanistic transformations of bicarbonate adsorbed CO2 product on metal oxide nanoparticles. The adsorption of CO2 on TiO2 nanoparticles was performed at CO2 pressures of 61, 100, 169, 266, 507, and 1006 mTorr and are shown in Fig. 6. The vibrational frequencies of the bands observed in the spectra are labeled and again are assigned in Table 2. Similarly, CO2 adsorption on TiO2 P25 nanoparticles has been investigated previously [37] with a focus on effects due to particle size and crystalline phase [38–40]. Upon adsorption of CO2 , a very different spectrum resulted compared to that of Fe2 O3 and ␥-Al2 O3 nanoparticles. It has been shown that different surface species can be formed during the CO2 adsorption on the surface with different acid–base character [17,41]. The reactions of carbon dioxide with surface hydroxyl groups have been shown to result in the formation of bicarbonate species but bonds due to this species is weak in the TiO2 spectrum shown in Fig. 6. The adsorption of CO2 on basic sites usually forms monodentate carbonate, whereas the adsorption of CO2 on an acidic metal ion with its neighboring basic oxygen has been shown to produce bidentate carbonate species [30,31,37,41,42]. In addition, the surface phase compositions, in this case either anatase or rutile, have been found to determine the surface acid–base properties of TiO2 and, thus, the type and number of adsorbed species formed from the adsorption of CO2 [37]. In this work, for TiO2 nanoparticle surface experiments mainly carbonate and some bicarbonate adsorbed surface species are observed. Specifically, the absorption bands at 1392 and 1510 cm−1 as well as 1355 and 1566 cm−1 are assigned to the O–C–O stretching vibrations of adsorbed carbonate. The absorption at 1076 cm−1 , is also assigned to the 1 symmetric C–O stretch for carbonate. Other bands present in the spectra from ∼1100 to 1700 cm−1 were assigned to other forms of the carbonate species. Previously, species formed by the adsorption of CO2 onto TiO2 have been characterized as either bidentate, at 1243 and 1670 cm−1 , or monodentate, ca. 1320–1370 cm−1 [17]. More recently, the bands at 1250 and 1670 cm−1 have been assigned to the carboxylate species that is adsorbed on to Ti3+ sites as a results of a charge transfer [18,37,42,43]. From the spectra shown in Fig. 6, bidentate structure of adsorbed carbonate can also be detected by the absorption bands at 1355 and 1566 cm−1 , which correspond to the 3 (O–C–O)a and 3 (O–C–O)s modes, respectively [37]. Two additional bands at 1392 and 1510 cm−1 with 3 split smaller

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Fig. 6. Transmission spectra recorded after CO2 adsorption on TiO2 as a function of increasing pressure. Spectral regions of 2600–4000 cm−1 – hydroxyl groups – (top) and 1000–1800 cm−1 – adsorbed carbonate (bottom) – are shown. Gas-phase absorptions have been subtracted from the spectra.

Fig. 5. Energy minimized structures for five cluster models [M2 (OH)4 (␮OH)(H2 O)x (HCO3 − )] (where M = Al and x = 0 or 1) optimized at B3LYP/6-31+G* level of theory. Atoms of different elements are highlighted with different colors: metal sites, green; oxygen, red; hydrogen, blue; carbon, yellow. Adapted and reproduced with permission from Journal of Physical Chemistry B 2006, 110, 12005.Copyright 2006 American Chemical Society.

than that of bidentate, can be attributed to monodentate binding configuration. Upon evacuation of the TiO2 surface the bicarbonate bands disappear, leaving only carbonate absorption thus implying that the bicarbonate species is only stable in the presence of gas-phase CO2 on the TiO2 surface. Finally, the absorption bands at 1222, 1423 and 1623 cm−1 were assigned to the ı4 (COH), 3 (OCO)s and 2 (OCO)a vibrational modes of adsorbed bicarbonate, respectively. Hydrogen bonded OH groups are evident in the broad band at approximately 3130 cm−1 and may indicate interactions between surface hydroxyl groups and the adsorbates. Similarly to the other oxide surfaces, spectral losses are seen in the 3600–3700 cm−1 region and are assigned to the loss of the hydroxyl groups coordinated to metal atoms upon the adsorption of CO2 and are likely involved in the formation of the surface products. The rate of CO2 uptake on Fe2 O3 , ␥-Al2 O3 and TiO2 nanoparticle surfaces was also investigated using FTIR spectroscopy. Time course plots are used to measure the adsorbed species by integrating absorbance in the 1200–1800 cm−1 region as a function of time under dry conditions. Spectra were collected before and after the oxides were exposed to a constant pressure of CO2 and the uptake was monitored as a function of time. FTIR spectra were recorded every 2.5 s at an instrument resolution of 4 cm−1 for a total of 1000 s for the ␥-Al2 O3 and TiO2 samples. For the Fe2 O3 nanoparticles, FTIR spectra were recorded every 2.5 s at a resolution of 4 cm−1 for a total of 400 s. The integrated absorbance is plotted versus time and shown in Fig. 7. Initial rates were calculated for these surfaces under dry conditions from the initial slopes within the first 50–100 s after gas-phase CO2 was introduced into the cell. Calculated initial adsorption rates for Fe2 O3 and ␥-Al2 O3 and TiO2 nanoparticles

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Fig. 8. Transmission spectra recorded after 0.274 Torr CO2 adsorption on (a) Fe2 O3, (b) ␥-Al2 O3 and (c) TiO2 nanoparticles in the presence of water as 40% RH. Spectral regions of 2600–4000 cm−1 – hydroxyl groups – (top) and 1000–2000 cm−1 – adsorbed carbonate (bottom) – are shown. Gas-phase absorptions have been subtracted from the spectra.

Fig. 7. CO2 uptake at 0.274 Torr under dry conditions is shown for (a) Fe2 O3 , (b) ␥-Al2 O3 , and (c) TiO2 nanoparticles as a function of time. The integrated absorbance for the absorption bands of adsorbed bicarbonate and carbonate in the region of 1000–1800 cm−1 is plotted as a function of time.

were 0.0164, 0.0022 and 0.0020 s−1 , respectively. From these data it can be seen that TiO2 had the slowest observed CO2 uptake relative to Fe2 O3 and ␥-Al2 O3 nanoparticles, whereas Fe2 O3 nanoparticles were found to take up CO2 approximately 7 times faster than TiO2 and ␥-Al2 O3 . Saturation coverages were reached on all three samples at 0.430, 0.135 and 0.092 integrated absorbance units normalized for sample mass for Fe2 O3 and ␥-Al2 O3 and TiO2 nanoparticles, respectively. This would constitute a saturation coverage ratio of 4.7:1.5:1.0 which is different in Fe2 O3 nanoparticle case from BET surface area ratio of 3.8:1.8:1.0. Thus, higher surface coverage of CO2 observed on Fe2 O3 nanoparticles cannot be attributed to larger BET surface area only. We propose here that smaller Fe2 O3 nanoparticles

can have an increased amount of structural defects which would enable for surface coverage to be affected. Concentration of structural defects has recently been observed in GeO2 nanoparticles to increase with decreasing nanoparticle size due to the enhancement of surface effects [44]. Amorphous nanoparticles in that size range have been shown to be divided into two parts: the core and the surface with structures different due coordinatively unsaturated sites at the surfaces [44]. This might not be the case for slightly larger ␥-Al2 O3 (13 ± 2) and TiO2 (23 ± 3) particles with CO2 adsorption proceeding on both close to that expected from the available surface area. 3.4. Transmission FTIR spectroscopy studies of CO2 adsorption on metal oxide nanoparticles in the presence of co-adsorbed water Transmission FTIR spectroscopy was used to monitor the adsorption of CO2 on oxide surfaces in the presence of water vapor as relative humidity. Here relative humidity, %RH, is defined as %RH =

p(H2 O) × 100% p∗ (H2 O)

(1)

where p(H2 O) is water vapor pressure and p* (H2 O) is the saturation water vapor pressure. The FTIR spectra of Fe2 O3, ␥-Al2 O3 , and TiO2 nanoparticles in the presence of 0.274 Torr of CO2 and water vapor corresponding to 40% RH at 296 K are shown in Fig. 8a–c, respectively. The spectra are shown in Fig. 8 with the gas-phase absorptions subtracted. It can be seen that after exposure of carbon

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Table 3 Vibrational frequencies following adsorption of CO2 and co-adsorbed water at 40% RH on Fe2 O3 , ␥-Al2 O3 and TiO2 nanoparticles. 0.274 Torr CO2 + 40% RH Vibrational mode assignments

Fe2 O3 (cm−1 )

␥-Al2 O3 (cm−1 )

TiO2 (cm−1 )

Literature referencesa

Solvated carbonate 3 (O–C–O)a 3 (O–C–O)s 1 (C–O)s

1492 1350 n.o.b

1505 1406 1033

1428 1324 1043

1487–1520 (1484) 1335–1402 (1420) 1070 (1028)

a b

Range taken from Refs. [17,21,30,31,49]. n.o.: not observed

dioxide the resulting spectra are quite different when water vapor is present. As shown in Fig. 8 for all of the surfaces, there is an absorption band at ∼1640 cm−1 due to the bending mode of adsorbed water and two additional peaks, best discerned on Fe2 O3 nanoparticles due to the solvated carbonate. These absorption bands are assigned to the 3 (O–C–O)a and 3 (O–C–O)s vibrational modes, respectively, and have been included in Table 3. Additionally, there is also a broad OH stretching region with distinct shoulders at 3210, 3440 and 3510 cm−1 for all surfaces shown in the spectra in Fig. 8. In Baltrusaitis et al., for the Fe2 O3 nanoparticle sample, the hydroxyl stretching region showed a structured band, increasing in intensity with increasing RH, with absorptions at 3011, 3105, 3266 and 3515 cm−1 , which is seen in spectra above 40% RH [16]. For the ␥-Al2 O3 surface, the OH stretching region of the spectrum was shown less structured [16] and does not differ very much from pure water adsorption on ␥-Al2 O3 as reported by AlAbadleh and Grassian [12]. Baltrusaitis et al. showed that these surfaces were exposed to water vapor as a function of increasing relative humidity in the absence of CO2 for comparison purposes, and monitored using transmission FTIR spectroscopy [16]. Initially, the spectra showed an increase in intensity in the bending mode region, ı(H2 O), of adsorbed water at 1640 and 1646 cm−1 for Fe2 O3 and ␥-Al2 O3 nanoparticles, respectively. The spectra also showed OH stretching absorption bands centered at 3381 and 3387 cm−1 , with shoulders at 3243 and 3248 cm−1 for both Fe2 O3 and ␥-Al2 O3 nanoparticles, respectively, but with less structure than when coadsorbed with water. This comparison was provided due to the large difference in the water absorption band stretching region for the Fe2 O3 in the presence and absence of adsorbed carbonate. The smaller changes seen in the stretching region for ␥-Al2 O3 were attributed to the formation of less carbonate as was indicated by the weaker absorption bands. These shoulders on the OH stretching region are suggested to be a result of Fe–OH2+ groups and more crystalline water in the vicinity of the ions present in the thin water film [16]. It has also been shown that for both Fe2 O3 and ␥-Al2 O3 nanoparticles there was an increase in the intensity of the carbonate absorptions with an increase in relative humidity [16]. Here, the transmission FTIR spectrum of Fe2 O3 nanoparticles in the presence of 0.274 Torr of gas-phase CO2 and 40% RH is shown in Fig. 8 and shows absorption bands at 1350 and 1492 cm−1 . These carbonate bands have been shown to increase significantly relative to the decreasing intensity of the bicarbonate bands as the relative humidity is increased from 0 to 90% RH [16]. For the ␥-Al2 O3 surface following exposure to 0.274 Torr of CO2 in the presence of 40% RH water vapor shown in Fig. 8, similar features are observed and the bands at 1406 and 1505 cm−1 are assigned to the 3 (OCO)s and 3 (OCO)a modes of carbonate, respectively. The ␥-Al2 O3 carbonate absorptions show similar trends but these absorptions are not nearly as intense as in the case of the Fe2 O3 nanoparticles suggesting less carbonate formation on the ␥Al2 O3 surface as already noted. These bands, though less intense when compared to the Fe2 O3 nanoparticles, are consistent with a solvated, adsorbed carbonate ion, as inferred from previous solu-

tion phase experiments [16,27,32]. For these surfaces, it has been shown that there is a decrease in the intensity of the adsorbed bicarbonate bands for the Fe2 O3 and ␥-Al2 O3 nanoparticles as the relative humidity is increased and a simultaneous increase in intensity of the absorption bands near 1333, 1511 cm−1 and 1394, 1525 cm−1 for Fe2 O3 and ␥-Al2 O3 nanoparticles, respectively [16]. As discussed above, these bands have been attributed to adsorbed solvated surface carbonate. For the TiO2 surface, very weak absorption bands are also observed at 1324 and 1428 cm−1 and assigned to the 3 (OCO)s and 3 (OCO)a modes of carbonate respectively. These features have very low intensity and are broader and shifted from that of the iron and alumina surfaces, suggesting that the adsorption on this surface may be different from that of the other surfaces. The exact solvated carbonate binding mode on metal oxide clusters is still under debate. Several groups using quantum chemical calculations proposed various structures as most stable energetically and/or calculated frequencies agreed best with the experimental ones [16,25,27]. Inner bridged and outer hydrogen bonded structures were proposed in a study of surface carbonate complexes on hematite [27], monodentate and outer sphere hydrogen bonded on ferrihydrate [25], and bridged on both nanoparticulate Fe2 O3 and ␥-Al2 O3 [16]. Particular care must be paid to the data obtained where there is no bulk water present, as is case of the experiments under controlled relative humidity. Recent molecular dynamics (MD) calculations have suggested that under ambient conditions adsorbed water does not cover surfaces uniformly even at 100% RH thus allowing for the gaseous molecules to bind (a) directly to the surfaces sites, (b) via hydrogen bonded interactions with water and/or partially or fully solvated surface [45,46]. Similarly, solvated nitrate ion, isoatomic with carbonate, adsorption on metal oxide surfaces in the presence of relative humidity has been shown to result in several species bonded either directly with the surface of with a different degree of hydrogen bonding [22]. A non-uniform water layer will have even more profound impact on the adsorbed CO2 product structure on metal oxide nanoparticles due to the high surface are-to-volume ratio and atoms being located at corner and edge sites [46]. In solution phase, adsorption from carbonate solutions result in carbonate ion coordinated to oxide nanoparticle surfaces. Adsorption of CO2 from the gas phase in the presence of relative humidity on metal oxide nanoparticles also results in structure attributed to the solvated carbonate, [13,16]. In fact, FTIR spectra of both adsorbates, e.g. solution phase carbonate and CO2 in the presence of relative humidity, are almost identical [13]. In this case, CO2 reaction mechanisms to form adsorbed carbonate are of interest, as it is unclear whether carbonate is formed in the adsorbed water layer or via surface assisted reactivity. Possible metal oxide nanoparticle surface enhanced adsorbed carbonate formation mechanisms from CO2 via bicarbonate intermediate were recently investigated computationally [47]. Cluster models, corresponding to the reaction of CO2 with adsorbed H2 O molecule to form bicarbonate followed by formation of carbonic acid are shown in Fig. 9. The CO2 and H2 O reaction proposed proceeded via coordination of both reac-

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Fig. 9. RI-SCS-MP2/aug-cc-pVTZ calculated relative energy levels for CO2 and nH2 O reaction (where n = 1, 2, 3) coordinated to Al(OH)3 (I and II) together with the corresponding intermediates (III and IV) and transition states (IV). Final reaction structures of H2 CO3 coordinated to Al(OH)3 are also shown (VI). B3LYP/6-311++G(d,p) optimized structures for CO2 and H2 O reaction are also shown. Bond lengths and interatomic distances are given in Angstroms. Bond angles are given in degrees. Adapted and reproduced with permission from Journal of Physical Chemistry A 2010, 114, 235. Copyright 2010 American Chemical Society.

tants to Al(OH)3 cluster (I and II), followed by the formation of the corresponding bicarbonate intermediates (III and IV) and transition states (IV). Final reaction structures of H2 CO3 coordinated to Al(OH)3 are also shown (VI). Almost all of the reaction path structures were energetically favorable with respect to the initial reactants with the main transition state of hydrogen transfer IV decreasing in energy with two or three water molecules involved. In three water molecule mechanism transition state barrier was relatively small, 8.3 kcal/mol, thus showing possible bicarbonate intermediate in adsorbed solvated carbonate formation. Collectively, these data showed possible routes and intermediates of adsorbed carbonate formation as facilitated by hydroxylated metal oxide surfaces. The rate of CO2 uptake on metal oxides nanoparticle surfaces was further investigated in the presence of co-adsorbed water layer. Kinetic experiments were done on Fe2 O3 , ␥-Al2 O3 and TiO2 nanoparticles to calculate initial CO2 rates in the presence of water as relative humidity. Time course plots were used to illustrate the growth of the integrated bicarbonate and carbonate region as a function of time. Spectra were collected after the oxides were exposed to 0.274 Torr of CO2 and 40% RH and the uptake was monitored as a function of time. FTIR spectra were recorded every 150 and 6 s, at an instrument resolution of 4 cm−1 for a total of 3500 and 1200 s for the ␥-Al2 O3 and TiO2 samples, respectively. For the Fe2 O3 nanoparticles, FTIR spectra were recorded every 4 s, at a resolution of 4 cm−1 for a total of 500 s. The absorption bands located in the 1440–1560 cm−1 region for Fe2 O3 and ␥-Al2 O3 , and in the 1350–1490 cm−1 region for TiO2 nanoparticles, respectively, were integrated over the course of the experiment. The integrated

absorbance of these absorption bands are plotted versus time and have been normalized for mass. These plots are shown in Fig. 10. The initial adsorption rates were calculated for these surfaces under humid conditions from the initial slopes within the first 100 s after gas-phase CO2 and H2 O was simultaneously introduced into the cell. From these data, at 0.274 Torr of CO2 and 40% RH H2 O, ␥-Al2 O3 has the slowest uptake relative to Fe2 O3 and TiO2 nanoparticles, whereas Fe2 O3 nanoparticles were found to take up CO2 approximately 100 times faster than ␥-Al2 O3 and TiO2 with calculated initial adsorption rates if 0.0016, 0.00002 and 0.0002 s−1 for Fe2 O3, ␥-Al2 O3 and TiO2 nanoparticles, respectively. Importantly, adsorbed CO2 products on Fe2 O3 nanoparticles continued to increase in spectral intensity even after the fast initial growth as well as on TiO2 nanoparticles, whereas on ␥-Al2 O3 no growth was observed after the initial change in adsorbed CO2 amount. Amorphous nature of Fe2 O3 nanoparticles combined with a very high surface area enables for the enhanced reactivity towards CO2 . The relative initial rates for CO2 uptake at 40% RH was also compared to the initial uptake rates determined for the oxide surfaces under dry conditions as discussed in the previous section. The initial rates for all surfaces under dry conditions were found to be faster than under wet conditions. For Fe2 O3 nanoparticles, the rate under dry conditions was found to proceed approximately 10 times faster than under wet conditions. For the ␥-Al2 O3 surface, the rate under dry conditions was found to proceed more that 100 times faster than in the presence of relative humidity. Finally, TiO2 under dry conditions was found to take up CO2 approximately 10 times faster than under humid conditions. This can be explained by the fast formation of the adsorbed water layer followed by a slower diffusion of

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of adsorption was found to be faster on the iron oxide surfaces than for the alumina or titanium oxide surfaces. In the presence of adsorbed water, the reaction between carbon dioxide and water proceeds to yield adsorbed carbonate and protonated hydroxyl groups. In the absence and presence of coadsorbed water, iron oxide was found to show the fastest initial CO2 adsorption kinetics followed by alumina or titania surfaces. These results show the importance of co-adsorbed water and the reaction of CO2 under ambient conditions, which has potential implications in heterogeneous processes on metal oxide nanoparticles ranging from catalysis to atmospheric chemistry. Acknowledgements This material is based upon work partially supported by the National Science Foundation under Grant No. CHE-0952605. Any opinions, findings, and conclusions or recommendations expressed in this material are those of the author(s) and do not necessarily reflect the views of the National Science Foundation. References

Fig. 10. CO2 uptake at 0.274 Torr and 40% RH is shown for (a) Fe2 O3 , (b) ␥-Al2 O3 , and (c) TiO2 nanoparticles as a function of time. The integrated absorbance for the absorption bands of adsorbed solvated carbonate is plotted as a function of time.

CO2 molecules to form solvated carbonate species. Adsorbing water molecules form strongly bonded network with surface hydroxyl groups and the formation of solvated carbonate can proceed via the surface enhanced protonation reactions in the adsorbed water layer rather than direct reaction with the surface hydroxyl groups. 4. Conclusions Adsorption and coordination of CO2 to metal oxide nanoparticle surfaces have been reviewed with particular emphasis on infrared spectroscopy based detection and complementary theoretical methods to aid in observed frequency assignments. Major adsorption products were bicarbonate and mono- and bidentate carbonates. Adsorbed water plays an important role in the reaction chemistry of carbon dioxide on metal oxide nanoparticle surfaces. In the absence of co-adsorbed water, surface hydroxyl groups react with carbon dioxide to form adsorbed bicarbonate. The initial rate

[1] V.H. Grassian (Ed.), Environmental Catalysis, CRC Press, Boca Raton, FL, 2005. [2] T.R. Knutson, R.E. Tuleya, Y. Kurihara, Simulated increase of hurricane intensities in a CO2 -warmed climate, Science 279 (1998) 1018–1020. [3] D. Raynaud, J. Jouzel, J.M. Barnola, J. Chappellaz, R.J. Delmas, C. Lorius, The ice record of greenhouse gases, Science 259 (1993) 926–934. [4] M.I. Hoffert, K. Caldeira, A.K. Jain, E.F. Haites, L.D.D. Harvey, S.D. Potter, M.E. Schlesinger, S.H. Schneider, R.G. Watts, T.M.L. Wigley, D.J. Wuebbles, Energy implications of future stabilization of atmospheric CO2 content, Nature 395 (1998) 881–884. [5] IPCC The Fourth Assessment Report, Climate Change 2007: Synthesis Report, in: R.B. Abdelkader Allali, S. Diaz, I. Elgizouli, D. Griggs, D. Hawkins, O. Hohmeyer, L.k.K.-B. Bubu Pateh Jallow, N. Leary, H. Lee, D. Wratt (Eds.), IPCC The Fourth Assessment Report, Intergovernmental Panel on Climate Change, 2007. [6] M. Ishida, H. Jin, CO2 recovery in a power plant with chemical looping combustion, Energy Convers. Manage. 38 (1997) S187–S192. [7] H.J. Freund, M.W. Roberts, Surface chemistry of carbon dioxide, Surf. Sci. Rep. 25 (1996) 227. [8] K. Weissermel, H.J. Arpe, Industrial Organic Chemistry. Important Initial and Intermediate Products, Verlag Chemie, Weinheim, Germany, 1976. [9] W. Stumm, Chemistry of the solid–water interface: processes at the mineralwater and particle-water interface in natural systems, John Wiley & Sons, Inc., New York, 1992. [10] B.J. Finlayson-Pitts, L.M. Wingen, A.L. Sumner, D. Syomin, K.A. Ramazan, The heterogeneous hydrolysis of NO2 in laboratory systems and in outdoor and indoor atmospheres: An integrated mechanism, Phys. Chem. Chem. Phys. 5 (2003) 223. [11] G.E. Brown Jr., V.E. Henrich, W.H. Casey, D.L. Clark, C. Eggleston, A. Felmy, D.W. Goodman, M. Graetzel, G. Maciel, M.I. McCarthy, K.H. Nealson, D.A. Sverjensky, M.F. Toney, J.M. Zachara, Metal oxide surfaces and their interactions with aqueous solutions and microbial organisms, Chem. Rev. 99 (1999) 77–174. [12] H.A. Al-Abadleh, V.H. Grassian, Oxide surfaces as environmental interfaces, Surf. Sci. Rep. 52 (2003) 63–161. [13] J. Baltrusaitis, V.H. Grassian, Surface reactions of carbon dioxide at the adsorbed water–iron oxide interface, J. Phys. Chem. B 109 (2005) 12227–12230. [14] S.C. Lee, B.Y. Choi, T.J. Lee, C.K. Ryu, Y.S. Ahn, J.C. Kim, CO2 absorption and regeneration of alkali metal-based solid sorbents, Catal. Today 111 (2006) 385–390. [15] S.C. Lee, H.J. Chae, S.J. Lee, B.Y. Choi, C.K. Yi, J.B. Lee, C.K. Ryu, J.C. Kim, Development of regenerable MgO-based sorbent promoted with K2 CO3 for CO2 capture at low temperatures, Environ. Sci. Technol. 42 (2008) 2736–2741. [16] J. Baltrusaitis, J.D. Schuttlefield, E. Zeitler, J.H. Jensen, V.H. Grassian, Surface reactions of carbon dioxide at the adsorbed water–oxide interface, J. Phys. Chem. C 111 (2007) 14870–14880. [17] G. Busca, V. Lorenzelli, Infrared spectroscopic identification of species arising from reactive adsorption of carbon oxides on metal oxide surfaces, Mater. Chem. 7 (1982) 89. [18] D.H. Gibson, Carbon dioxide coordination chemistry: metal complexes and surface-bound species. What relationships? Coord. Chem. Rev. 185–186 (1999) 335–355. [19] M. Mikkelsen, M. Jorgensen, F.C. Krebs, The teraton challenge. A review of fixation and transformation of carbon dioxide, Energy Environ. Sci. 3 (2009) 43–81. [20] A. Desikusumastuti, T. Staudt, H. Grönbeck, J. Libuda, Identifying surface species by vibrational spectroscopy: bridging vs. monodentate nitrates, J. Catal. 255 (2008) 127–133. [21] J. Baltrusaitis, J.H. Jensen, V.H. Grassian, FTIR spectroscopy combined with isotope labeling and quantum chemical calculations to investigate adsorbed bicarbonate formation following reaction of carbon dioxide with surface

J. Baltrusaitis et al. / Chemical Engineering Journal 170 (2011) 471–481

[22]

[23] [24]

[25]

[26]

[27]

[28]

[29] [30] [31]

[32] [33]

[34] [35]

hydroxyl groups on Fe2 O3 and Al2 O3 , J. Phys. Chem. B 110 (2006) 12005– 12016. J. Baltrusaitis, J. Schuttlefield, J.H. Jensen, V.H. Grassian, FTIR spectroscopy combined with quantum chemical calculations to investigate adsorbed nitrate on aluminum oxide surfaces in the presence and absence of co-adsorbed water, Phys. Chem. Chem. Phys. 9 (2007) 4970–4980. G. Herzberg, Infrared and Raman Spectra of Polyatomic Molecules, 1945. V.P. Indrakanti, J.D. Kubicki, H.H. Schobert, Quantum chemical modeling of ground states of CO2 chemisorbed on anatase (0 0 1), (1 0 1), and (0 1 0) TiO2 surfaces, Energy Fuels 22 (2008) 2611–2618. D.B Hausner, N. Bhandari, A.-M. Pierre-Louis, J.D. Kubicki, D.R. Strongin, Ferrihydrite reactivity toward carbon dioxide, J. Colloid Interface Sci. 337 (2009) 492–500. J.D. Kubicki, G.P. Halada, P. Jha, B.L. Phillips, Quantum mechanical calculation of aqueous uranium complexes: carbonate, phosphate, organic and biomolecular species, Chem. Cent. J. 3 (2009), doi:10.1186/1752-1153X-1183-1110. J.R. Bargar, J.D. Kubicki, R. Reitmeyer, J.A. Davis, ATR-FTIR spectroscopic characterization of coexisting carbonate surface complexes on hematite, Geochim. Cosmochim. Acta 69 (2005) 1527–1542. J.R. Rustad, P. Zarzycki, Calculation of site-specific carbon-isotope fractionation in pedogenic oxide minerals, Proc. Natl. Acad. Sci. U.S.A. 105 (2008) 10297–10301. K.K. Irikura, R.D. Johnson III, R.N. Kacker, Uncertainties in scaling factors for ab initio vibrational frequencies, J. Phys. Chem. A 109 (2005) 8430–8437. L. Ferretto, A. Glisenti, Study of the surface acidity of an hematite powder, J. Mol. Catal. A: Chem. 187 (2002) 119–128. A.M. Turek, I.E. Wachs, E. DeCanio, Acidic properties of alumina-supported metal oxide catalysts: an infrared spectroscopy study, J. Phys. Chem. 96 (1992) 5000. C. Su, D.L. Suarez, In situ infrared speciation of adsorbed carbonate on aluminum and iron oxides, Clays Clay Miner. 45 (1997) 814–825. T. Yoshida, D.L. Thorn, T. Okano, J.A. Ibers, S. Otsuka, Hydration and reduction of carbon dioxide by rhodium hydride compounds. Preparation and reactions of rhodium bicarbonate and formate complexes, and the molecular structure of RhH2 (O2 COH)(P(i-Pr)3)2, J. Am. Chem. Soc. 101 (1979) 4212–4221. B.R. Flynn, L. Vaska, Carbon dioxide fixation leading to stable molecular bicarbonato complexes of d8 metals, J. Am. Chem. Soc. 95 (1973) 5081–5083. R.J. Crutchley, J. Powell, R. Faggiani, C.J.L. Lock, The formation and molecular structure of a monodentate bicarbonate complex of palladium(II), Inorg. Chim. Acta 24 (1977) L15–L16.

481

[36] K.T. Jung, A.T. Bell, An in situ infrared study of dimethyl carbonate synthesis from carbon dioxide and methanol over zirconia, J. Catal. 204 (2001) 339–347. [37] W. Su, J. Zhang, Z. Feng, T. Chen, P. Ying, C. Li, Surface phases of TiO2 nanoparticles studied by UV raman spectroscopy and FT-IR spectroscopy, J. Phys. Chem. C 112 (2008) 7710–7716. [38] D.J.C. Yates, Infrared studies of the surface hydroxyl groups on titanium dioxide, and of the chemisorption of carbon monoxide and carbon dioxide, J. Phys. Chem. 65 (1961) 746–753. [39] G. Ramis, G. Busca, V. Lorenzelli, Low-temperature carbon dioxide adsorption on metal oxides: spectroscopic characterization of some weakly adsorbed species, Mater. Chem. Phys. 29 (1991) 425–435. [40] L.F. Liao, C.F. Lien, D.L. Shieh, M.T. Chen, J.L. Lin, FTIR study of adsorption and photoassisted oxygen isotopic exchange of carbon monoxide, carbon dioxide, carbonate, and formate on TiO2 , J. Phys. Chem. B 106 (2002) 11240–11245. [41] A. Auroux, A. Gervasini, Microcalorimetric study of the acidity and basicity of metal-oxide surfaces, J. Phys. Chem. 94 (1990) 6371–6379. [42] G. Ramis, G. Busca, V. Lorenzelli, Low-temperature CO2 adsorption on metaloxides – spectroscopic characterization of some weakly adsorbed species, Mater. Chem. Phys. 29 (1991) 425–435. [43] F. Boccuzzi, A. Chiorino, M. Manzoli, D. Andreeva, T. Tabakova, FTIR study of the low-temperature water-gas shift reaction on Au/Fe2 O3 and Au/TiO2 catalysts, J. Catal. 188 (1999) 176–185. [44] V.V. Hoang, T.P. Duy, Structural defects and thermodynamics of vitreous GeO2 nanoparticles, Curr. Appl. Phys. (2010), doi:10.1016/j.cap.2010.07.023. [45] A. Rahaman, V.H. Grassian, C.J. Margulis, Dynamics of water adsorption onto a calcite surface as a function of relative humidity, J. Phys. Chem. C 112 (2008) 2109–2115. [46] V.H. Grassian, Surface science of complex environmental interfaces: oxide and carbonate surfaces in dynamic equilibrium with water vapor, Surf. Sci. 602 (2008) 2955–2962. [47] J. Baltrusaitis, V.H. Grassian, Carbonic acid formation from reaction of carbon dioxide and water coordinated to Al(OH)3: A quantum chemical study, J. Phys. Chem. A 114 (2010) 2350–2356. [48] G. Lefevre, In situ Fourier-transform infrared spectroscopy studies of inorganic ions adsorption on metal oxides and hydroxides, Adv. Colloid Interface Sci. 107 (2004) 109–123. [49] G. Magnacca, G. Cerrato, C. Morterra, M. Signoretto, F. Somma, F. Pinna, Structural, Surface characterization of pure and sulfated iron oxides, Chem. Mater. 15 (2003) 675–687.