Low-temperature CO2 adsorption on alkali metal

International Journal of Greenhouse Gas Control 10 (2012) 191–198

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Low-temperature CO2 adsorption on alkali metal titanate nanotubes K. Upendar, A. Sri Hari Kumar, N. Lingaiah, K.S. Rama Rao, P.S. Sai Prasad ∗ Inorganic and Physical Chemistry Division, Indian Institute of Chemical Technology, Hyderabad 500 607, India

a r t i c l e

i n f o

Article history: Received 2 December 2011 Received in revised form 18 May 2012 Accepted 16 June 2012 Keywords: Adsorption Carbon dioxide Potassium titanate Sodium titanate Titanium dioxide

a b s t r a c t Na and K titanate nanotubes were prepared by hydrothermal method and characterized by BET, XRD, Raman, TEM and SEM-EDAX techniques. Low-temperature CO2 adsorption properties of the materials, along with their reusability and water tolerance, were established by generating the breakthrough curves and these properties were compared with those of the parent titania. The nanomaterials exhibited higher adsorption capacity than the parent titania. Their adsorption capacity increased in the presence of water and they also exhibited reproducibility of their breakthrough capacity. The Yoon–Nelson model used to calculate the rate constant for adsorption revealed higher values for the nanotubes than the parent titania. © 2012 Elsevier Ltd. All rights reserved.

1. Introduction The ever-increasing magnitude of greenhouse gases released into the atmosphere has been very alarming in terms of climate disasters. Particularly, carbon dioxide (CO2 ) is found to be responsible for the global warming and its effective capture and sequestration/usage is receiving considerable attention (Liu et al., 2007). Carbon dioxide capture may be achieved by several techniques. The amine-CO2 chemistry has been used for decades as a method to selectively remove CO2 from flue gas streams using aqueous solutions of monoethanolamine, diethanolamine, and methyldiethanolamine as common absorbents. Though these systems are effective, they pose several drawbacks such as corrosion, high-energy consumption and the limitation on amine concentration in the aqueous phase due to viscosity and foaming issues. Appropriate membranes are still to be identified for the satisfactory removal of CO2 from lean sources, particularly with high selectivity. Cryogenic separation becomes highly un-economic. Adsorption by solids is thus a promising technique when efficient, low cost and reusable adsorbent materials are available. Good reviews are available comparing the various methods of capture technologies (Choi et al., 2009). The carbon dioxide thus separated can go for sequestration in terrestrial ecosystems, geological formations and oceans (White et al., 2003). Alternatively, the pure gas can be used to synthesize several organic chemicals (Aresta et al., 2001), over heterogeneous catalysts (Aresta and Dibenedetto, 2004; Park et al., 2001).

∗ Corresponding author. Tel.: +91 40 2719 3163; fax: +91 40 2716 0921. E-mail address: [email protected] (P.S. Sai Prasad). 1750-5836/$ – see front matter © 2012 Elsevier Ltd. All rights reserved. http://dx.doi.org/10.1016/j.ijggc.2012.06.008

Basically there are three types of solid adsorbent materials for CO2 capture. The high temperature sorbents that include calcium based materials are applied at temperatures greater than 400 ◦ C. Quite often the carbonation temperature for CaO-based adsorbents is between 600 ◦ C and 700 ◦ C and the regeneration temperature is above 950 ◦ C (Blamey et al., 2010). The main disadvantage of these materials lies in the loss of reversible adsorption property due to sintering (Abanades and Alvarez, 2003). Even though incorporation of inert materials in CaO increases the sorption recyclability (Lu et al., 2006) capital and operating costs increase due to presence of huge amount of inert material. Alkali metal (Li, Na, K, etc.) containing ceramics also exhibit CO2 capture ability. Nakagawa and Ohashi (1998) reported Li2 ZrO3 as a high temperature material. The second class of solid adsorbents that fall in the intermediate-temperature solid sorbents (200–400 ◦ C) category includes the hydrotalcites (HTs) containing both high surface area and basic sites (Roman et al., 2008). Their desorption temperature lies around 400 ◦ C, due to strong interaction between CO2 and the basic sites (Ram Reddy et al., 2006). Magnesium oxide is also used as an intermediate temperature range CO2 adsorbent (Feng et al., 2007). However, MgO has moderate CO2 adsorption capacity and poor thermal stability during regeneration. The last class of adsorbents is the low-temperature adsorbents whose adsorption temperature falls below 200 ◦ C. Carbon based materials fall under this class and they are attractive due to low cost, high surface area, high amenability to pore structure modification and surface functionalization. But the carbon materials exhibit thermal sensitivity and poor selectivity toward CO2 due the physical and weak interaction. The CO2 capture ability of carbon materials decrease significantly at temperatures in the range 50–120 ◦ C (Arenillas et al., 2005). The zeolites adsorb CO2 at around room temperature but loose their ability beyond 200 ◦ C (Ko et al., 2003). In the presence of moisture their CO2 adsorption

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capacity declines due to H2 O competing with CO2 for the active adsorption sites (Janchen et al., 2007). Metal organic frameworks (MOFs) are the new family of porous materials (Bae et al., 2008) with high adsorption capacity and CO2 selectivity. However, their thermal stability and the cost of preparation limit their applicability (Bourrelly et al., 2005). Alkali metal carbonates are used for CO2 capture at temperature below 200 ◦ C (Lee et al., 2008). Interaction between alkali carbonate and CO2 in the presence of moisture is reversible and highly exothermic, but the rate of carbonation reaction is slow. In the case of Na2 CO3 /K2 CO3 dispersed adsorbents the reactivity decreases with increase in sorption/regeneration operations (Hayashi et al., 1998). Amine based solid adsorbents prepared by immobilizing organic amines on support materials (Harlick and Sayari, 2007) are also used as low temperature adsorbents for CO2 capture. However, much attention has to be focused on incorporating them in suitable mesoporous materials which are expensive. Thus, there is a strong need for the development of newer solid materials with specific requirements to enhance the CO2 sorption capacity. The metal oxide nanotubes particularly, the potassium titanate nanotubes, have attracted increasing attention (Masaki et al., 2002) in photo-cleavage of water (Uchida et al., 1997) and in fuel cell applications (Corcoran et al., 2000). The mechanism of formation and the method of regulating the morphologies of the various types of nanotubes are the focus of intense research (Kasuga et al., 1999). Among the various approaches that have been reported for the preparation of nanotubes the hydrothermal method is simple and cost-effective for large-scale production. In spite of gaining tremendous importance, the application of potassium and sodium titanate nanotubes for the adsorption of carbon dioxide has not been dealt with sufficiently in the literature, excepting a few recent reports (Li et al., 2012; Yu et al., 2008). The nanotubes, which possess highly reactive OH groups, Tin+ and basic oxygens could form important materials for CO2 adsorption (Hadjiivanov et al., 1997). The objective of the present work is to evaluate the Na and K titanate nanotubes for CO2 adsorption and to compare their ability with that of the parent titania. The nano materials were prepared by hydrothermal method and characterized by nitrogen adsorption for surface area, X-ray diffraction, TEM, FT-IR and Raman spectroscopy. The applicability of these alkali metal titanates for CO2 adsorption and water tolerance is highlighted along with the determination of the rate constant of adsorption.

2. Experimental 2.1. Chemicals and gases TiO2 (P25, Degussa AG, Germany; anatase to rutile ratio 4:1), sodium and potassium hydroxide pellets (S.D. Fine-Chem. Ltd, A.R. grade) were procured from commercial sources. Double distilled water was used in the preparation. Helium, used as purge gas during the activation and the cyclic desorption period, had a purity of 99.995%.

2.2. Synthesis of adsorbent The preparation of titanate nanotubes was based on the alkaline hydrothermal method proposed by Kasuga et al. (1998) and used by Hodos et al. (2004). Typically, 6 g commercial of TiO2 powder was dispersed in 120 ml of 10 M alkali hydroxide solution. The resulting suspension was stirred for 30 min, and transferred into a Teflon-lined stainless steel autoclave, sealed, and maintained at

130 ◦ C for 48 h. The resulting material was cooled and allowed to age for 2 days. After subjecting the material to repetitive centrifuging and washing with deionized water, the material was dried in air at 100 ◦ C for 12 h and finally calcined in air at 300 ◦ C for 5 h. Close to theoretical yields of titania nanotubes were obtained. 2.3. Characterization X-ray powder diffraction patterns of the samples were recorded on a Rigaku Miniflex diffractometer using Cu K␣ radiation ˚ at 40 kV and 30 mA. The measurements were obtained (1.5406 A) in steps of 0.045◦ with account time of 0.5 s and in the 2 range of 10–80◦ . The morphological analysis was carried out using transmission electron microscopy (TEM on a JEOL 100S microscope). For the preparation of a sample for TEM, a suspension containing about 1 mg adsorbent/ml of ethanol was prepared and sonicated for 10 min. A few drops of the suspension were placed on a hollow copper grid coated with a carbon film. The FT-IR spectra were recorded on a Bio-rad (DIGILAB Excalibur series) spectrometer, with a resolution of 1 cm−1 , using KBr disk method. The specific surface areas of the samples were determined by N2 adsorption, acquiring the data on an Autosorb-1 instrument (Quanta chrome, USA) at liquid N2 temperature. The powders were first out-gassed at 200 ◦ C to ensure a clean surface prior to construction of adsorption isotherm. A cross-sectional area of 0.164 nm2 of the N2 molecule was assumed in the calculation of the surface area using the BET method. Raman spectra were acquired with a micro-Raman attachment of a Jobin-Yvon T64000 Raman spectrometer operated in triple subtractive mode. The excitation wavelength was the 514.5 nm generated by a Coherent Innova 308 argon ion lasers. Using a 20× objective, the laser beam diameter at the sample was approximately 1.5 ␮m. The power at the sample was kept below 1 mW to minimize local heating. Plasma lines were removed from the incident light using a narrow band pass laser line filter. The dispersed spectra were detected with a liquid nitrogen cooled CCD detector, and the data were recorded using Lab spec 3.03 software. 2.4. CO2 adsorption measurement CO2 adsorption experiments were carried out in a fixed-bed down-flow reactor (Stainless steel, 30 cm length, 10 mm i.d.). Helium and carbon dioxide gas flows were controlled using mass flow controllers. The temperature of the bed in the reactor was controlled using a PID controller. The reactor outlet was connected to a micro gas chromatograph (Varian CP-4900) having a thermal conductivity detector and a porapak Q column. About 1 g of adsorbent mixed with 1 g of glass beads was suspended between two quartz plugs and was initially activated in helium at 200 ◦ C for 1 h. A gas mixture of 10% CO2 balanced by He was passed through the bed for the adsorption study at three different temperatures, i.e. 50, 70 and 90 ◦ C and the outlet concentration was monitored periodically. After attaining saturation, the sample was flushed with helium for 30 min, and desorption proceeded at 140 ◦ C. Water tolerance measurements were made by bubbling helium through a saturator containing water at 40 ◦ C, and the mixture was passed over the bed for 1 h. The sample was then purged with dry helium and the subsequent adsorption was continued. After the sample had attained its CO2 saturation capacity at 70 ◦ C, it was purged with helium at the same temperature for 30 min and desorption was carried out by ramping the temperature at a rate of 5 ◦ C/min up to 140 ◦ C. The extent of desorption was calculated as the percentage of saturated adsorption.


3. Results and discussion 3.1. Structural and textual properties TiO2 is known as an n-type semiconductor, which contains donor-type defects such as oxygen vacancies and titanium interstitials. Ti O based nanotube is formed by scrolling of an exfoliated TiO2 -derived nanosheet into a hollow multiwall nanotube. The titania nanotubes produced via the alkaline hydrothermal method exist as open-ended multilayer-walled nanotubular structures (Zhou et al., 2010). The hydrothermal treatment of titania enhanced its BET surface area. The surface area of potassium titanate (K–Ti-NT) reached a maximum value of (263 m2 /g) whereas that of sodium titanate (Na–Ti-NT) was obtained as 92 m2 /g. These values are much higher than that of TiO2 (61 m2 /g) and were consistent with those reported in the literature. During hydrothermal treatment the titania transforms into titania nanotube exhibiting a layered structure with the alkali metal inside the layers. This creates porosity with a consequent increase in surface area. The reason for the high BET surface area of the potassium titanate nanotube than the sodium titanate nanotube could be due to the difference in size of the interlayer cation (Na = 0.19 nm; K = 0.26 nm) in the nanotubes. The XRD patterns of the parent titania (TiO2 ) along with those of K–Ti-NT and Na–Ti-NT samples are shown in Fig. 1. The pattern of TiO2 reveals the crystallinity of the sample containing a mixture of anatase and rutile, as reported in the literature (Sikuvhihulu et al., 2008). The pattern of K–Ti-NT can be indexed to K2 Ti8 O17 phase (JCPDS files 80-2023 and 41-1100) as reported by Yuan et al. (2004). The much broader diffraction peaks (Sun et al., 2002; Song et al., 2007) indicate the relatively amorphous nature and the smaller size of the crystals. The crystallite size calculated by the Scherrer equation is of the order of 10 nm, in the same range reported by Yuan et al. (2004). As the titania transforms itself into nanotube, a two or three layered structure with interlayer spacing of 0.7–0.8 nm between them, is formed. Thus, the original size of TiO2 reduces

with the formation of nanotube. Antonio et al. (2007) also reported synthesis of nanotubes with 10 nm outer diameter. In the case of Na–Ti-NT the pattern exhibited four broad peaks at 2 = 9.7◦ , 24.6◦ , 28.6◦ and 48◦ . The line positions revealed a good match to the pattern given in JCPDS file 72-0148, exhibiting a monoclinic structure of Na2 Ti3 O7 (Song et al., 2007; Weng et al., 2006). These peaks do not correspond to the anatase phase implying complete transformation into an orthorhombic structure of layered sodium titanates (Rivera et al., 2011). The titania nanotube retains its structure up to 400 ◦ C (Antonio et al., 2007). Only after heating beyond this temperature the main peak at 9.7◦ disappears with simultaneous reappearance of peaks due to tetragonal anatase. In the present case since the calcination was done at 300 ◦ C, the anatase phase in nanotube structure was not detected. Fig. 2 shows the FT-IR spectra of TiO2 and the nanotubes. K–Ti-NT showed broad and intense bands located at 3390 and 1621 cm−1 due to the O H stretching vibrations and the H O H bending vibrations of water, respectively ascribed to the presence of water either in the interparticle space or in the tunnel like interlayers, as also reported by Yuan et al. (2004). Titania holds adsorbed water more tightly due to its strong Lewis acidity, which is again due to the presence of coordinately unsaturated surface (cus) Ti4+ atoms (Morrow, 1990). A comparison of the areas of the peaks under these two bands reiterates the fact that the cus sites are more in number in the case of the nanotubes than in the parent titania. A characteristic difference between the spectra of nanotubes and the non-nanotubes was that the former displayed a broad band at 768 cm−1 representing the symmetric stretching vibrations of the Ti O bonds of TiO4 tetrahedra (Yurchenko et al., 1981) involving non-bridging oxygen coordinated with metal ions which is not present in the titania. In Fig. 2c the band at about 900 cm−1 can be attributed to the four coordinated oxygen atoms interacting with Na+ .

# Anatase * Rutile K2Ti8O17 Na2Ti3O7

Intensity (a.u.)

(c)

*

*

*

*

(b)

#

# #

* 10

20

* 30

* 40

#

*

*

50

60

70

(a)

80

2Theta(degree) Fig. 1. XRD patterns of (a) TiO2 , (b) K–Ti-NT, and (c) Na–Ti-NT.

193

Fig. 2. FT-IR spectra of (a) TiO2 , (b) K–Ti-NT, and (c) Na–Ti-NT.

194


(c)

Intensity (a.u.)

(b)

(a)

100

200

300

400

500

600

700

800

900

1000

Raman shift (cm-1) Fig. 3. Raman spectra of (a) TiO2 , (b) K–Ti-NT, and (c) Na–Ti-NT.

The Raman spectra of TiO2 and that of the nanotubular materials are presented in Fig. 3. The titania exhibited five bands at 146, 195, 395, 514, and 638 cm−1 . Evidently, the peaks at 145, 195 and 640 cm−1 correspond to the active modes of anatase, and those of rutile phase at 450 and 614 cm−1 . The spectrum for K–Ti-NT showed three major bands at 284, 453, and 661 cm−1 and a few weak bands at 87, 107, 201, and 371 cm−1 . As reported by Ma et al. (2005), all these bands are different from those of anatase phase. The bands at 450 and 668 cm−1 could be assigned to Ti O Ti vibrations (Bavykin et al., 2006). Recent studies have also reported Raman bands at 280, 448, 668 and 917 cm−1 assignable to titanate nanotubes (Sikhwivhilu et al., 2009). Thus the Raman information corroborated the data obtained from that of the FT-IR. Raman spectroscopy was used to investigate further the local structure of as-prepared samples of Na2 Ti3 O7 . The Raman spectrum of asprepared Na–Ti-NT nano-sample obtained at room temperature revealed broad bands at 148, 170, 276, 374, 446, 661 cm−1 (Fig. 3c). The Raman bands at ∼276 and ∼446 cm−1 can be assigned to the Na O Ti stretching and framework Ti O Ti vibrations, respectively. The band at ∼910 cm−1 (broad band) can be assigned to the symmetric stretch of short Ti O bonds involving non-bridging oxygen atoms associated with sodium ions, characteristic of the parent bulk sodium compound Na2 Ti3 O7 . It is interesting to note that the bands of nanotubes of Na2 Ti3 O7 were broader than those of bulk Na2 Ti3 O7 (Papp et al., 2005). Fig. 4 presents the TEM photographs of the two types of nanotubes synthesized hydrothermally. The tubes are of 10–20 nm diameters with a length of more than 120–200 nm. During the interaction of TiO2 with the concentrated solution of the alkali some of the Ti O Ti bonds are broken, forming an intermediate product containing Ti O A (A = alkali) and Ti OH, thus leading to the formation of lamellar fragments that are the intermediate phases in the formation process of the nanotube materials (Kasuga et al.,

Fig. 4. TEM image of (a) K–Ti-NT and (b) Na–Ti-NT.

1999). The elemental Na, and K contents were measured with EDX (energy dispersive X-ray) spectroscopy attached to a scanning electron microscope (SEM S-3000N). The EDX analysis revealed the Ti/K and the Ti/Na atomic ratios of 4:1 and 2:1 which agree closely with the reported values (Sun et al., 2002; Yuan et al., 2004; Liu et al., 2010). 3.2. Carbon dioxide sorption capacity After confirming the structure of the nanotubes, their CO2 adsorption capacities were estimated by generating the breakthrough curves (Fig. 5). The CO2 adsorption capacities of the sorbents are shown in Table 1. For the sake of comparison the result of the analysis obtained on TiO2 is also included. It may be observed that the breakthrough adsorption capacity increased considerably by the formation of nanotubes, as also the saturation capacity. However, with increase in temperature there has been a gradual decrease in the adsorption characteristics. This is expected because of the exothermic nature of adsorption. Thus, the nanotube materials were proved to be superior to their parent titania. The nanotubes contain three types of active sites, highly active OH groups, Ti4+ and huge amount of interlayer water molecules (Hadjiivanov et al., 1997). At lower temperatures, one or more of these species may be responsible for the adsorption capacity of the nanotubes. The nanotubular structure also seems to be more significant in dictating the adsorption capacity. In Ti O based nanostructures, metals/protons occupy the cavities between the layers of the TiO6 octahedra. Potassium titanates K2 O·nTiO2 (n = 2–8), with a low potassium content (n = 6, 8), particularly, potassium octatitanate K2 Ti8 O17 , possesses a stable tunnel-like structure, with a framework formed by edgeand corner-shared TiO2 octahedra (Anderson and Wadsley, 1961).


195

Table 1 CO2 adsorption capacities at different temperatures. Material

Adsorption temp. (◦ C)

TiO2

50 70 (wet CO2 ) 70 (dry CO2 ) 90 50 70 (wet CO2 ) 70 (dry CO2 ) 90 50 70 (wet CO2 ) 70 (dry CO2 ) 90

K–Ti-NT

Na–Ti-NT

kYN (min−1 )

Adsorption capacity (mmol/g) BTC

SAC

1.486 0.784 0.938 0.465 1.550 1.556 1.016 0.488 2.111 2.122 1.943 1.904

1.970 1.109 1.372 1.206 2.334 2.190 1.757 1.320 4.415 4.302 3.020 2.815

0.139 0.120 0.123 0.119 0.305 0.381 0.293 0.184 0.268 0.203 0.193 0.139

BTC, breakthrough capacity; SAC, saturation adsorption capacity.

Compared to K–Ti-NT, its Na analog showed higher adsorption capacity. The role of Na ions in Na–Ti-NT was investigated by Suetake et al. (2008). There exist two states of Na ion in the Na–TiNT, strongly and weakly interacting with the titania nano tube (TNT) system (Suetake et al., 2008). The Na ions that are strongly bound to the TNT interlayer play an important role in stabilizing the tube structure. This observation is in accordance with the data reported by Walton et al. (2006) wherein it was demonstrated that the Na containing nanotube had higher adsorption capacity than that of the K containing nanotube, i.e. the smaller size of Na cation has greater ion-quadruple interaction with CO2 than the larger size K cation. Antonio et al. (2007) reported FTIR spectra of adsorbed CO2 on TiO2 nanotubes. The peaks at 1406 and 1617 cm−1 were assigned to CO3 −2 and H2 O, respectively. It was also reported that even after evacuation at 0 ◦ C these peaks were visible. An FTIR spectroscopic analysis of CO2 adsorbed on TiO2 and K–Ti-NT nanotube, collected after passing CO2 at 70 ◦ C, also showed a broad peak at 1630 cm−1 , due to H2 O and a peak at 1406 cm−1 in the case of the nanotube due to the formation of CO3 −2 (Fig. 6). The peak at 1406 cm−1 was not seen in the case of parent titania. As proposed by Antonio et

al, the following reaction could have occurred more favorably on nanotubes, leading to the formation of carbonate species. CO2 + 2OH− → CO3 −2 + H2 O The main difference that can be observed from the FTIR and Raman spectra was that the anatase phase in the parent titania disappears in case of nanotubes. Normally, the layered structure of nanotubes, rearrange to tetrahedral anatase after calcination at higher temperatures above 400 ◦ C. In the present case the sample were calcined at 300 ◦ C. Hence we feel that the difficult desorption on titania could be due to strong adsorption of CO2 on the anatase phase. Table 2 compares the adsorption capacities of different materials. It can be observed that many of the adsorbents developed thus far suffer from one or more drawbacks such as low capacity, poor selectivity, poor tolerance to water and high temperature regeneration or activation. A close look at the values indicates considerable promise for the titania nanotubes. 3.3. Recyclability and water tolerance of the materials

1.0

The nanotubes have been further examined 4 times for their reproducibility of adsorption capacity (Table 3). There is only about 5% drop in the capacity in the case of Na–Ti-NT and K–Ti-NT. TiO2 has shown more than 50% drop. Nanotubes are known to exhibit high chemical stability (Varghese et al., 2003).

C/C0

0.8 0.6 0.4 0.2

(c)

0.0 1.0

0

5

10

15

20

25

30

(b)

35

0.6 0.4 0.2

(b)

0.0 0

5

10

15

20

25

30

0

5

10

15

20

25

30

35

1.0

C/C0

0.8 0.6

(a)

Transmittance (%)

C/C0

0.8

0.4 0.2

(a)

0.0

35

Time (min) Fig. 5. Adsorption capacities (breakthrough curves) of adsorbents at 1 atm pressure, and 70 ◦ C temperatures under wet and dry conditions: (a) (*) wet TiO2 , (♦) dry TiO2 ; (b) (*) wet Na–Ti-NT, (♦) dry Na–Ti-NT; and (c) (*) wet K–Ti-NT, (♦) dry K–Ti-NT.

1800

1700

1600

1500

1400

1300

Wave number (cm-1) Fig. 6. FTIR spectra of CO2 adsorbed: (a) TiO2 and (b) TNT.

1200

196


Table 2 Comparison of the adsorption capacities of various materials. Name of adsorbent

Adsorption temperature (K), pressure (bar)

CO2 adsorption capacity (mmol/g)

References

Limitations

CaO/Ca12 Al14 O33

923, 1

11.6

Li et al. (2006)

Ca(OH)2

923, 0.41

10.7

Wu et al. (2007)

Li2 ZrO3 or Li4 SiO4

848, 1 or 773, 0.02

6.1

Hydrotalcites Carbon

481, 1 275, 1

0.9 3.5

Fernandez et al. (2006) or Kato et al. (2002) Ding and Alpay (2001) Wang et al. (2008)

High CO2 partial pressure; high desorption temperature High CO2 partial pressure; high desorption temperature High desorption temperature

Na-Y Na-X HKUST-1 (4 wt% H2 O)

273, 0.1 298, 0.4 303, 1

4.9 3.9 8

Michelena et al. (1977) Cavenati et al. (2004) Yazaydin et al. (2009)

Alkali metal carbonate based TEPA on MCM-41

333, 1 348, 0.05

9.4 4.5

Samanta et al. (2012) Yue et al. (2008)

Titania nanotubes

323, 1

4.4

Present work

High temperature desorption High pressure, low temperature adsorption Poor water tolerance Poor water tolerance High pressure; poor thermal stability in the presence water or acidic gases Highly exothermic Poor water tolerance, week hydrothermal stability Better water tolerant, higher adsorption capacity, Better desorption at low temperature

Table 3 Reproducibility of CO2 adsorption capacities (70 ◦ C, dry conditions) and the extent of desorption. No. of runs

CO2 adsorption capacity BTC (SAC) (mmol/g)

1st run 2nd run 3rd run 4th run

Amount of adsorbed CO2 desorbed (%)

TiO2

Na–Ti-NT

K–Ti-NT

TiO2

Na–Ti-NT

K–Ti-NT

0.938 (1.373) 0.420 (0.595) 0.404 (0.581) 0.388 (0.575)

1.943 (3.020) 1.888 (2.768) 1.879 (2.693) 1.852 (2.781)

1.015 (1.756) 0.972 (1.484) 0.941 (1.445) 0.991 (1.434)

56 62.8 52.4 47.9

92.7 96.8 97.5 98.3

89.5 93.1 95.6 95.2

The water tolerance capacities of the three samples were examined by generating the breakthrough curves before and after treating the samples with water vapor (Fig. 5). In the case TiO2 , there was a decrease in the adsorption capacity, whereas the nanotubes showed the reverse trend. The increase in the adsorption capacity could be attributed to the formation of alkali carbonates in the case of nanotubes. Thus, the metal titanate nanomaterials appear to be promising adsorbents for the low temperature capture of carbon dioxide. The influence of adsorbed water on CO2 uptake was studied by first saturating the adsorbent surface with water at 70 ◦ C (partial pressure of water: 67 mmol Hg). It was observed that the pre-adsorption of water had positive effect on CO2 uptake. As reported by Raupp and Dumesic (1985), CO2 and water have similar effect on the adsorption capacity of TiO2 .

3.5. Extent of desorption Table 3 also gives the percentage recovery of CO2 during desorption. It can be observed that the desorption was only about 60% in the case of titania whereas the nanotubes displayed almost complete recovery. The desorption curves obtained on TiO2 and TNTs are shown in Fig. 7. The extent of desorption was higher in the case of nanotubes than in the parent titania.

3.4. The kinetics of CO2 adsorption A relatively simple Yoon–Nelson model (Yoon and Nelson, 1984), addressing the adsorption and breakthrough of adsorbate vapors or gases was selected for the estimation of the adsorption rate constant. The linear form of Yoon–Nelson equation for a singlecomponent system can be expressed as follows:

ln

C C0 − C

= kYN t − t0.5 kYN

where kYN is the Yoon–Nelson rate constant in min−1 ; t is the breakthrough time in min. The parameter kYN was determined from a plot of ln(C/(C0 − C)) versus sampling time, t and the corresponding data are presented in Table 2. As can be seen from the table, higher values were obtained for the rate constant in the case of the nanotubes, than in the case of titania itself.

Fig. 7. The CO2 desorption curves of TiO2 and TNTs at 140 ◦ C.


4. Conclusion Na and K titanate nanotubes are promising as CO2 adsorbents showing higher efficiency than the parent TiO2 under low temperature and atmospheric pressure conditions. The materials possess good tolerance to the presence of water and show good cyclability during the adsorption–desorption cycles. The Yoon–Nelson kinetic parameters substantiate these observations.

Acknowledgment The authors thank Department of Science and Technology (DST), India, for financial support.

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