Morphologically and compositionally tuned lithium silicate nanorods as high-performance carbon dioxide sorbents

Abstract

The effective capturing of carbon dioxide using regenerable high capacity sorbents is a prerequisite for industrial applications aiming at CO₂ capture and sequestration. The removal of CO₂ directly from chemical reaction environments at high temperature is a less energy intensive method of its separation with the added benefit of improved efficiency in equilibrium limited reactions. However, the separation of CO₂ at the typical reaction temperatures of 573–1073 K is a challenging task due to the non-availability of absorbents with kinetics comparable to reaction rates. Moreover their poor durability due to sintering and particle growth on prolonged use at high temperature is also an impediment to their practical application. Herein, we demonstrate the development of an efficient CO₂ absorbent material, made of Li₄SiO₄ nanorods, with ultrafast sorption kinetics as well as remarkable durability. These nanorods enabled easier surface reaction with CO₂ due to shorter diffusion pathways for lithium from the bulk to the surface of the rods permitting extremely fast absorption of CO₂. Furthermore, the compositional tuning of the materials helped to realize absorbents with extraordinary CO₂ absorption rates of 0.72 wt% s⁻¹ at 100% CO₂/923 K. The exceptional performance of these absorbents at lower temperatures (573–823 K) as well as lower CO₂ pressures (0.15 atm) demonstrates their potential in practical CO₂ separation applications.

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1. Introduction

Global warming and its effect on climate destabilization have grown into an alarming environmental concern in recent times. The massive consumption of fossil fuels leading to vast greenhouse gas emissions since the industrial revolution is believed to be the major reason for the current rise in atmospheric temperature. Green technological alternatives based on wind, solar and hydropower plants are therefore favored for minimizing carbon in the atmosphere. However, considering the cost and time frame required for the widespread implementation of such renewable energy alternatives, an immediate solution appears remote. To reduce carbon footprints globally, a reduction in CO₂ emission by maintaining a balanced carbon cycle in the atmosphere is imperative. Selective greenhouse gas capture and separation for reuse or sequestration thus emerge as a viable choice to limit its emission to the atmosphere.¹ Among the strategies designed, CO₂ capture using sorbents with appreciable absorption capacity, enhanced kinetics, selectivity, durability and cyclic stability together with a minimum energy penalty on the process appears practically viable at least in the short term.^2–6

Removal of CO₂ at high temperatures is an effective option to reduce its emissions to the atmosphere.^7–9 In several high-temperature chemical and petrochemical processes, CO₂ is a major product and its removal therefore at the temperature and pressure of the reaction offers a less energy intensive method of separation culminating in smaller carbon footprints. In reactions such as methane reforming, the removal of CO₂ from the reaction vessel could also enhance the rate of reaction (sorption enhanced steam reforming) thereby providing additional benefits of increased productivity.^10,11 However, selective and rate enhanced sorption of CO₂ at elevated temperatures is a challenging task, especially in the temperature range of 573–1073 K where most such reactions occur. Lithium based ceramic absorbents are considered as candidate materials for selective CO₂ removal at high temperature,^12–15 although their application is thus far limited by the slow kinetics in the lower temperature range (573–823 K) and poor material durability in the higher temperature range (823–1073 K).

Among the lithium ceramics, lithium silicate (Li₄SiO₄) based ceramic oxides are the most efficient CO₂ sorbents by virtue of their high absorption capacity, faster kinetics and reasonable durability.^16–19 The absorption properties of some of the recently published lithium-based ceramic absorbents are shown in Table S1 in section S1 of the ESI†^20–25 for a quick evaluation of the status. We also have reported on the sol–gel based synthesis of Li₄SiO₄ particles (powders termed as sol–gel in this manuscript) with CO₂ absorption performance comparable to such published reports.²⁶ In the work reported herein, we have adapted a facile microwave assisted sol–gel synthesis route^27,28 to successfully realize Li₄SiO₄ nano-rods with significantly superior CO₂ absorption performance. A detailed synthesis protocol starting from LiNO₃ and colloidal silica is described under the experimental section. The powders thus obtained exhibited excellent CO₂ absorption capacity, cyclic stability and superior absorption/desorption rates (powders termed as microwave sol–gel in this manuscript). We believe that the enhanced kinetics for CO₂ absorption arise from the very small thickness (20 to 30 nm) of rod-shaped Li₄SiO₄ particles. Irrespective of their nano-size, the large aspect ratio of the particles provided them with better stability against aggregation at high working temperatures thereby providing improved durability for a large number of absorption/desorption cycles. Besides this, the compositional control of these morphologically tuned materials helped to realize novel eutectic compositions containing Na and K as new generation CO₂ absorbents with exceptionally high CO₂ absorption rates even at low and moderate temperatures (<723 K).

2. Experimental

2.1 Synthesis of Li₄SiO₄ through a microwave sol–gel process

Li₄SiO₄ was synthesised from LiNO₃ (Alfa Aesar, UK) and colloidal silica (Aldrich Chemicals, USA) as the starting precursors. Initially 2.39 M LiNO₃ solution was prepared by dissolving it in distilled water. Hydrolysis was carried out by the addition of NH₄OH to the LiNO₃ solution while stirring at room temperature. Colloidal silica was added dropwise under stirring for 1 hour to obtain a sol. The precursor sol was subjected to 700 W microwave radiation at 2.45 GHz for 10 min. A domestic microwave oven (Panasonic-NNGT231M) was used to carry out this experiment. The precursor was then dried at 423 K and heated to 1073 K at a ramp rate of 1 °C min⁻¹ in air atmosphere. The powder calcined at 1073 K for 3 hours was used for CO₂ absorption studies unless specified otherwise.

2.2 Processing of the mixture of Li₄SiO₄ and alkali carbonates

Na₂CO₃ (99.9%), K₂CO₃ (99.8%) and Li₂CO₃ (98%), all from Merck, India, were mixed in the weight ratio of 10 : 60 : 30 [eutectic-1], 32 : 37 : 31 [eutectic-2] and 24 : 45 : 31 [eutectic-3] and then blended with the microwave sol–gel samples in the weight ratio 1 : 5. The mixtures were then heat treated at 1073 K before the absorption studies at various temperatures.

2.3 Characterization techniques

Phase changes of the powder samples during heat treatment were traced using in situ HT-XRD (Rigaku RINT-TTR III). The morphology and crystallinity of the powders on CO₂ absorption/desorption were further characterised using TEM analysis (HRTEM, FEI Tecnai 30 G2 S-TWIN operated at an accelerating voltage of 300 kV). CO₂ absorption properties were measured using a TGA apparatus (Perkin Elmer STA 6000, Singapore), in the temperature range of 100–750 °C. N₂ adsorption based surface area analysis was performed at 77 K using a Micromeritics Tristar 111 surface area analyser after degassing the sample at 200 °C for 2 h.

3. Results and discussion

The morphological features of the synthesized particles at 673 K, 773 K and 1073 K are observed by TEM and are shown in Fig. 1a–f. Remarkably, the TEM images show that the particles have a nanorod morphology and are entirely different from the powders synthesized by conventional processing techniques.^29,30 The XRD patterns of the microwave sol–gel Li₄SiO₄ samples after heat treatment (calcination) at 673 K, 773 K and 1073 K are further presented in Fig. 1g. From the XRD results, it is clear that the particles exhibited an amorphous nature even after calcination at 673 K. Crystalline phases rich in the Li₂SiO₃ phase emerged only on increasing the temperature to 773 K. Heat treatment at 1073 K was essential to realize Li₄SiO₄ phase (JCPDS file no. 37-1472) rich powders. The temperature dependent phase formation of the powders during the heat treatment of the synthesized particles has also been traced using in situ HT-XRD analysis and the results are detailed in section S2 of the ESI.† The results presented in Fig. 1g and further in ESI S2† clearly revealed that the formation of the Li₄SiO₄ phase is achieved at the expense of Li₂SiO₃ nanofibers that are initially nucleated from the amorphous silica particles coated with the lithium species. From the XRD analysis (section S2, Fig. S1 of the ESI†), it was inferred that microwave treatment has influenced the hydrolysis of LiNO₃ to LiOH. We believe that the formation of the LiOH phase during hydrolysis and further, the intermediate formation of Li₂SiO₃ fibers are critical steps in the successful formation of the Li₄SiO₄ nanorods (please see section S2 of the ESI† and the figures and video in the section for a detailed explanation of the phase and morphology formation mechanism). Small peaks corresponding to the Li₂SiO₃ phase are visible in Fig. 1g as well as in Fig. S2 (ESI†). Usually, such presence of the Li₂SiO₃ phase in small quantities is considered to be due to the reaction between Li₄SiO₄ and CO₂ in air^1,26 while cooling after heat treatment or handling during XRD measurements. However, the presence of Li₂SiO₃ peaks in Fig. S2† (in situ XRD spectra) indicates the existence of some amount of the lithium meta-silicate phase in our samples.

Fig. 1 TEM images of particles calcined at 673 K (a & b), 773 K (c & d) and 1073 K (e & f). (g) XRD patterns of Li₄SiO₄ particles synthesized by a microwave sol–gel approach after calcination at temperatures of 673 K, 773 K and 1073 K (Li₄SiO₄, JCPDS 37-1472), (Li₂SiO₃, JCPDS 29-0828). (h) Dynamic absorption curves of the microwave sol–gel powders in comparison to sol–gel Li₄SO₄ powders, measured at 20 °C min⁻¹ (100% CO₂).

Based on the XRD and TEM results, 1073 K was used as the calcination temperature for obtaining samples with Li₄SiO₄ as the predominant phase in our study. Unless otherwise mentioned, the samples named as “microwave sol–gel” in this manuscript are calcined at 1073 K. The BET surface area analysis indicated a value of 7.3 m² g⁻¹ (N₂ adsorption isotherm is shown in section S3 of the ESI†) for the particles.

We have initially performed a dynamic thermogravimetric analysis of microwave sol–gel Li₄SiO₄ particles under a 100% CO₂ flow to determine the temperature range of gas absorption (Fig. 1h). For comparison, similar curves of powder samples prepared by the sol–gel method as reported elsewhere²⁶ are also shown in Fig. 1h. As is clear from the dynamic absorption curves, the microwave sol–gel sample exhibited significant CO₂ absorption capacity in the temperature range of 673–973 K. The absorption is initiated at around 673 K leading possibly to the formation of an external shell of lithium carbonate (Li₂CO₃) on the particle surface. This shell formation remained the main mechanism of absorption up to a temperature of about 823 K. It is also evident from the figure that the rate of absorption increases at around 823 K and remains high till the reversible reaction occurs (∼993 K). The faster rate of absorption from 823 K is presumed to be due to the softening of the Li₂CO₃ shell, thereby decreasing the kinetic limitations imposed by the solid shell at lower temperatures.⁷ As mentioned, at T > 993 K the Li₄SiO₄ phase becomes stable, leading to the complete desorption of CO₂. The absorption studies have thus confirmed that the microwave samples have superior CO₂ absorption performances compared to sol–gel samples in the entire range of temperatures measured.

The morphological features of the microwave sol–gel based particles after CO₂ absorption are shown in Fig. 2a and b. Compared to the original particles that exhibited rods of 20 to 30 nm width [Fig. 1e and f], the particles after CO₂ absorption appeared entirely different. As shown in Fig. 2a and b, the carbonate formation leads to the aggregation of the particles during the absorption stage. It can be inferred that the carbonate-rich particles, easily coalesce together forming a shell around a number of particles in the near vicinity, developing into an aggregate structure. However, desorption of CO₂ leads to the reversible formation of individual Li₄SiO₄ particles, morphologically similar to the original samples (Fig. 2c and d). The morphological integrity of the particles after CO₂ desorption process clearly evidences the stability of the nano-structures in the CO₂ absorption/desorption cycle. The XRD patterns of the samples before and after CO₂ absorption as well as after desorption are included in section S4 of the ESI† for further analysis of the CO₂ absorption/desorption process.

Fig. 2 TEM images (a and b) of microwave sol–gel Li₄SiO₄ particles after the CO₂ absorption process; (c and d) particles after desorption process. (e) CO₂ absorption curves at various temperatures. (f) Graph of ln K versus 1/T for the two different processes of chemisorption (K₁) and diffusion (K₂) observed in the microwave sol–gel sample; (g) absorption–desorption performance of microwave sol–gel Li₄SiO₄ powders for 10 cycles. The first cycle of the run was done at 973 K and further runs at 873 K. In all cases, desorption was carried out by switching 100% CO₂ gas to 100% N₂ gas.

The CO₂ absorption capacity values of the particles in the temperature range of 673–973 K are compared in Fig. 2e. For collecting the absorption data, the samples were heated to the absorption temperature at 10 °C min⁻¹ in 100% N₂ gas and kept for 2 h under a 100% CO₂ gas flow. As expected, the amount of CO₂ absorbed was found to increase with increase in the absorption temperature. This relationship between temperature and absorption capacity as well as the mechanism of absorption are better expressed using kinetic constants and activation energy values calculated from the Arrhenius plots of kinetic constants as shown in Fig. 2f. Further details on the calculation of kinetic constants are included in section S5 of the ESI.† It should be noted that, contrary to the general trend, Fig. 2f indicates that the K₂ values are approximately 10 times larger than those of K₁. The larger K₂ values obtained here signifies faster lithium ion diffusion to the reaction interface compared to the chemisorption reaction.²⁹ The activation energy values calculated from the plots were 22.70 kJ mol⁻¹ for the chemisorption process (corresponding to K₁) and 61.10 kJ mol⁻¹ for the diffusion process (corresponding to K₂). The higher activation energy value for the diffusion process substantiated the extremely fast diffusion of lithium ions from the core of the material to the reactive interface at absorption temperatures.

Absorption rates were also calculated from the first two minutes of absorption curves at different temperatures. The particles synthesised through a microwave sol–gel method displayed an enhanced absorption rate of 0.093 wt% s⁻¹ at 973 K. This value is much higher than the values reported in the recent literature (see table in ESI S1† for a comparison of some of the reported values). This enhanced absorption rate should mainly be attributed to the nano-rod morphology characterized by a very small thickness/width of the particle, facilitating a rapid surface carbonate layer formation over the entire length during the first stage of absorption. Moreover, the rod morphology should also have enabled easier surface reaction providing shorter diffusion pathways for lithium from the bulk to the surface of the particle.

The cyclic stability and regenerability of the powder samples were evaluated through cyclic absorption–desorption measurements and the results are shown in Fig. 2g [absorption with 100% CO₂ and desorption with 100% N₂ gases]. The primary aim of this cyclic loading experiment was to examine whether the sintering of the particles or the segregation of the carbonate phase due to continuous use of the absorbent at high temperatures induced any decay in absorption performance. The initial absorption run was done at 973 K; thereafter the temperature of absorption was switched to 873 K and cyclic absorption–desorption performance for 9 consecutive cycles was recorded. The initial run at high temperature leading to more or less full conversion of the material to Li₂CO₃ and further to the Li₄SiO₄ phase helped to realize higher absorption values at 873 K (Fig. 2e and g). As shown in Fig. 2g the samples displayed consistent absorption–desorption performance for all the 10 cycles measured indicating high durability and cyclic stability of the materials. Furthermore, it should be noted that the desorption rate was better than the absorption rate through all the measurements. The large desorption rate obtained without thermal cycling would allow the application of the materials for CO₂ capture in pressure swing mode. Although, further cycling studies of thousands of cycles may be necessary before considering the material for real life applications, initial results as reported in Fig. 2g indicated that the microwave sol–gel powders may be considered as promising materials for industrial absorption applications. Schematic illustrations comparing the carbon dioxide sorption and desorption mechanisms of Li₄SiO₄ nanoparticles and the newly developed nano rods, as well the influence of particle morphology on their durability are provided in Fig. 3a and b. The small thickness of the nanorods enables fast absorption kinetics just as in the case of nanoparticles. The better durability of the nanorod sample is highlighted based on ceramic sintering and particle coalescence mechanisms. It is well known that in liquid phase assisted sintering, the densification is achieved by the rearrangement and shape changes of the particles. Rearrangement is strongly affected by the size and morphology of the particles. Spherical and mono-dispersed particles are advantageous for particle rearrangement and sintering while particles like nanorods, having a high aspect ratio, are difficult to sinter. The Ostwald ripening/particle coalescence process invariably results in a reduction of the total surface area of the particles.^31,32 Hence the particle coarsening process as shown in Fig. 3b of the schematic always leads to a reduction in absorption kinetics owing to the reduction of the absorbent–gas interface available for chemisorption. In contrast, the use of higher aspect ratio Li₄SiO₄ nanorods impedes coalescence and agglomerate formation due to which it is possible to maintain high absorption kinetics even after several absorption/desorption cycles. The effects of morphological tuning on the kinetics of the carbon dioxide absorption process as well as the durability of the resulting particles are clearly depicted in the illustration.

Fig. 3 Schematic illustrations comparing the carbon dioxide absorption and desorption mechanisms and the durability of (a) the newly developed Li₄SiO₄ nanorods and (b) commonly found Li₄SiO₄ nanoparticles.

Further, we tried to enhance the rate of absorption as well as the absorption capacity of the materials in the low-temperature range by modifying the materials with a eutectic composition of mixed alkali carbonates. Carbonates of sodium, potassium and lithium were mixed in the weight ratio of 10 : 60 : 30 [eutectic-1], 32 : 37 : 31 [eutectic-2] and 24 : 45 : 31 [eutectic-3] and further blended with the microwave sol–gel samples in the weight ratio of 20 : 80. The mixtures were then heated to 1073 K before absorption studies. Fig. 4a and b present the TEM images of the eutectic-3 sample obtained after the heat treatment process. The rod morphology of the sample is well maintained with the thickness varying from 40 nm to 100 nm. The XRD pattern presented in Fig. 4c confirmed the presence of ortho- and meta-silicates of lithium in addition to the alkali silicate phases of Na and K. A comparative evaluation of the dynamic absorption isotherms of the sol–gel,²⁶ unmodified microwave sol–gel and microwave sol–gel samples modified with eutectic powder mixtures are presented in Fig. 4d. It is clear that the addition of eutectic mixtures increased the absorption capacity in the low-temperature range of 623–823 K significantly. In the unmodified samples, the rate of absorption was very low in the initial absorption step (<823 K) compared to the second absorption step. However, in the samples containing the eutectic mixtures the absorption rate in the initial step was comparable to that of the second absorption step. In the samples eutectic-1 and eutectic-2, two absorption steps are clearly visible as in the case of the unmodified sample. This is attributed to the carbonate shell formation, although a softer one compared to the unmodified sample, in these two cases. However, in the case of the eutectic-3 sample the two steps have combined to form a more or less single step absorption curve.

Fig. 4 TEM images (a & b) of the eutectic-3 particles. (c) X-ray diffraction patterns of eutectic-3 [Li₄SiO₄, JCPDS 37-1472], (Li₂SiO₃, JCPDS 29-0828), (Na₄SiO₄, JCPDS 072-0185), (KLi₄SiO₄, JCPDS 038-0015), (K₄SiO₄, JCPDS 046-0601). (d) Dynamic absorption curves of the sol–gel and microwave sol gel as well as eutectic 1, 2 and 3 samples.

Detailed results of the CO₂ absorption performance of the sample (eutectic-3) at temperatures of 623–923 K are shown in Fig. 5a. As shown, absorption capacity as high as 35% could be observed at the temperature of 923 K. It should be noted that the amount of CO₂ adsorbed was higher than the value expected based on the stoichiometry of the reaction of one CO₂ molecule with one orthosilicate available in the eutectic-3 sample (see Fig. 4c and 5a). It is possible that the reaction might have continued till silica is formed (instead of Li₂SiO₃) at least in some of the fractions of the powder mixture. Further studies are required to fully understand the chemisorption mechanism of these samples.

Fig. 5 (a) CO₂ absorption curves of the eutectic mixture (eutectic-3) at various temperatures. (b) Absorption rates of eutectic-3 at different temperatures (rate values calculated from the initial two minutes of the absorption curve) in comparison to those of the microwave sol–gel and sol–gel samples. (c) Graph of ln K versus 1/T for the two different processes of chemisorption (K₁) and diffusion (K₂) observed for the eutectic-3 sample.

An enhanced absorption rate was observed at temperatures as low as 623 K. Fig. 5b shows a good comparison of the absorption rates of the eutectic-3 sample with the unmodified microwave and sol–gel²⁶ samples (calculated for the first 2 minutes of absorption). The rate of absorption observed for the eutectic-3 sample, 0.28 wt% s⁻¹ at 923 K, was significantly higher than the absorption rates measured with the unmodified microwave sol–gel and sol–gel samples. A comparison of the reported results (as shown in ESI S1†) including the most recent literature also shows that the absorption performance of the sample is exceptional. The CO₂ absorption value of the sample, in fact, reached >30 wt% within the first 42 seconds, based on which the absorption rate could be calculated to be 0.72 wt% s⁻¹.

In order to find out the rate determining steps, the absorption curves (Fig. 5a) were fitted to the double exponential model and the resulting kinetic parameters were used to derive the Arrhenius plots as shown in Fig. 4c (further details are included in section S5 of the ESI†). It should be noted that the K₁ values are larger in this case compared to the K₂ values and this is contrary to the behaviour of the microwave unmodified samples. However, this is attributed to the presence of eutectic phases in the modified samples, which are supposed to enhance significantly the chemisorption rates. As a result, the modified and unmodified samples have different reaction limitations in the entire absorption process despite having more or less similar particle sizes. The activation energy values calculated from Fig. 5c were 82.3 kJ mol⁻¹ and 205.18 kJ mol⁻¹ for the chemisorption and diffusion processes, respectively. The activation energy for the chemisorption process was found to be much lower than that for the diffusion process. The sizes of sodium and potassium ions are larger compared to the lithium ions and, therefore, the whole diffusion process could be restricted in this case.

Cyclic absorption–desorption performances for 15 cycles at 948 K using 100% CO₂ and 100% N₂ gases were recorded and the results are shown in Fig. 6a. As shown, the sample retained >95% (33.6 wt%) of its original absorption capacity even after 15 cycles at this very high temperature. Nevertheless, some structural and morphological changes are suspected as the shape of the absorption curve changed from a near perfect rectangle to smoothed edges as the number of cycles increased. Further improvement in the stability of the powders may be necessary and techniques like the addition of rare-earth second phases as reported by us recently may be required.³⁰ As shown here and in Fig. 5a, the sample showed an absorption capacity of ∼35% for absorption temperatures >823 K when the partial pressure of CO₂ gas was 1 atm (100% CO₂). Fig. 6b shows the comparative absorption curves of eutectic-3 at 873 K for the CO₂ partial pressure values of 1 atm (100% CO₂) and 0.15 atm (15% CO₂/85% N₂). It is clear that the sample is capable of absorbing CO₂ to >25 wt% within the first few minutes of absorption even at the reduced CO₂ partial pressure of 0.15 atm.

Fig. 6 (a) Cyclic absorption–desorption performance of eutectic-3 (absorption was carried out at 948 K and desorption at the same temperature by changing 100% CO₂ to 100% N₂ gas). (b) Comparison of the absorption curves of eutectic-3 with 100% CO₂ flow and with 15% CO₂ flow. Absorption was carried out at 873 K (15% CO₂ or 100% CO₂).

It should be noted that the product streams of several high-temperature chemical reactions as well as the flue gases from power plants contain 15% CO₂. Therefore, the results indicate the possibility of using the developed materials for practical applications of CO₂ capture. This exceptional absorption behavior of eutectic-3 should be due to the morphological as well as compositional features of this sample. With regard to the kinetics and CO₂ absorption capacity, these morphologically and compositionally tailored particles outperform other structured lithium silicate particles reported in the recent literature (data from the literature is shown in the section S1 of the ESI†).

4. Conclusions

In conclusion, a microwave assisted sol–gel synthetic approach is demonstrated for the synthesis of Li₄SiO₄ particles with a nanorod morphology. The nanorods exhibited a dramatically enhanced absorption rate for CO₂ along with exceptional durability. The nanorod morphology of the Li₄SiO₄ particles allowed ultrafast sorption kinetics due primarily to an easier surface reaction with CO₂ by virtue of shorter diffusion pathways for lithium from the bulk to the surface of the rods. In addition, the large aspect ratio of the nanorods helped to enhance the durability of the particles by limiting their Ostwald ripening upon high-temperature cyclic absorption/desorption loading. Further, we have modified the chemical composition of the Li₄SiO₄ particles by mixing them with a eutectic mixture of K and Na. This compositional control of the materials helped to realize absorbents with extraordinary CO₂ absorption rates of 0.72 wt% s⁻¹ at 100% CO₂/923 K. Cycling absorption–desorption studies of these powders revealed that the materials remain durable up to 15 cycles without any significant reduction in CO₂ absorption capacity. Furthermore, the modified samples showed remarkable absorption performance at lower temperatures (573–823 K) as well as lower CO₂ pressures (0.15 atm) demonstrating their potential in practical CO₂ separation applications.

Acknowledgements

The authors acknowledge the Council of Scientific and Industrial Research (CSIR), New Delhi, India, & Noritake Co. Limited, Aichi, Japan, for providing research facilities and financial support. The authors acknowledge Mr Kiran Mohan for TEM analyses and Dr Bhoje Gowd and Mr N. Prithviraj for XRD analyses.

Notes and references

1. M. Kato, S. Yoshikawa and K. Nakagawa, J. Mater. Sci. Lett., 2002, 21, 485–487.
2. J. Wang, L. Huang, R. Yang, Z. Zhang, J. Wu, Y. Gao, Q. Wang, D. O'Hare and Z. Zhong, Energy Environ. Sci., 2014, 7, 3478–3518.
3. B. N. Nair, R. P. Burwood, V. J. Goh, K. Nakagawa and T. Yamaguchi, Prog. Mater. Sci., 2009, 54, 511–541.
4. D. M. D'Alessandro, B. Smit and J. R. Long, Angew. Chem., Int. Ed., 2010, 49, 6058–6082.
5. L. Joos, K. Lejaeghere, J. M. Huck, V. Van Speybroeck and B. Smit, Energy Environ. Sci., 2015, 8, 2480–2491.
6. M. Zaman and J. H. Lee, Korean J. Chem. Eng., 2013, 30, 1497–1526.
7. A. Samanta, A. Zhao, G. K. H. Shimizu, P. Sarkar and R. Gupta, Ind. Eng. Chem. Res., 2012, 51, 1438–1463.
8. J. Ortiz-Landeros, I. C. Romero-Ibarra, C. Gomez-Yanez, E. Lima and H. Pfeiffer, J. Phys. Chem. C, 2013, 117, 6303–6311.
9. B. N. Nair, T. Yamaguchi, H. Kawamura, S. I. Nakao and K. Nakagawa, J. Am. Ceram. Soc., 2004, 87, 68–74.
10. C. Wang, Y. Chen, Z. Cheng, X. Luo, L. Jia, M. Song, B. Jiang and B. Dou, Energy Fuels, 2015, 29, 7408–7418.
11. M. S. Yancheshmeh, H. R. Radfarnia and M. C. Iliuta, Chem. Eng. J., 2016, 283, 420–444.
12. Y. Duan, H. Pfeiffer, B. Li, I. C. Romero-Ibar, D. C. Sorescu, D. R. Luebke and J. W. Halley, Phys. Chem. Chem. Phys., 2013, 15, 13538–13558.
13. H. Xu, W. Cheng, X. Jin, G. Wang, H. Lu, H. Wang, D. Chen, B. Fan, T. Hou and R. Zhang, Ind. Eng. Chem. Res., 2013, 52, 1886–1891.
14. T. Yamaguchi, T. Niitsuma, B. N. Nair and K. Nakagawa, J. Membr. Sci., 2007, 294, 16–21.
15. R. Rodriguez-Mosqueda and H. Pfeiffer, J. Phys. Chem. A, 2010, 114, 4535–4541.
16. K. Wang, X. Guo, P. Zhao, F. Wang and C. Zheng, J. Hazard. Mater., 2011, 189, 301–307.
17. J. Du and L. R. Corrales, J. Mater. Chem. B, 2006, 110, 22346–22352.
18. K. Essaki, M. Kato and H. Uemoto, J. Mater. Sci., 2005, 40, 5017–5019.
19. M. Seggiani, M. Puccini and S. Vitolo, Int. J. Greenhouse Gas Control, 2011, 5, 741–748.
20. H. Kim, H. D. Jang and M. Choi, Chem. Eng. J., 2015, 280, 132–137.
21. J. H. Lee, B. Moon, T. K. Kim, S. Jeoung and H. R. Moon, Dalton Trans., 2015, 44, 15130–15134.
22. K. Wang, X. Y. Wang, P. F. Zhao and X. Guo, Chem. Eng. Technol., 2014, 37, 1552–1558.
23. S. Shan, S. Li, Q. Jia, L. Jiang, Y. Wang and J. Peng, Ind. Eng. Chem. Res., 2013, 52, 6941–6945.
24. V. L. Mejía-Trejo and E. Fregoso-Israel, Chem. Mater., 2008, 20, 7171–7176.
25. R. Quinn, R. J. Kitzhoffer, J. R. Hufton and T. C. Golden, Ind. Eng. Chem. Res., 2012, 51, 9320–9327.
26. P. V. Subha, B. N. Nair, P. Hareesh, A. Peer Mohamed, T. Yamaguchi, K. G. K. Warrier and U. S. Hareesh, J. Mater. Chem. A, 2014, 2, 12792–12798.
27. H. Pfeiffer, P. Bosch and S. Bulbulian, Synthesis of lithium silicates, J. Nucl. Mater., 1998, 257, 309–317.
28. B. N. Nair, K. Keizer, T. Okubo and S. I. Nakao, Adv. Mater., 1998, 10, 249–252.
29. M. J. Venegas, E. Fregoso-Israel, R. Escamilla and H. Pfeiffer, Ind. Eng. Chem. Res., 2007, 46, 2407–2412.
30. P. V. Subha, B. N. Nair, P. Hareesh, A. Peer Mohamed, T. Yamaguchi, K. G. K. Warrier and U. S. Hareesh, J. Phys. Chem. C, 2015, 119, 5319–5326.
31. G. Cao, Nanostructures & Nanomaterials. Synthesis, Properties and Applications, Imperial College Press London, 2004.
32. M. Prokesova and Z. Panek, Ceram. Int., 1989, 15, 369–374.

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Footnote

† Electronic supplementary information (ESI) available: Experimental and supplementary figures and tables included. See DOI: 10.1039/c6ta06133h

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