Introduction

There is no doubt that countless progresses have been made in controlling SOx/NOx and Hg emissions from fossil fuel fired power plants1. However, the actions on CO2 emission control have been slow2. The importance of CO2 capture in fossil fuel-fired power plants cannot be underestimated any more due to the catastrophic effect of the continuous increase in CO2 concentration in atmosphere and the ill effects it has on the environment3,4,5, as indicated in the recently reached Paris Agreement6.

Nowadays, chemical absorption by liquid solvents, including amines and carbonates, is widely considered to be the most promising method among various post-combustion CO2 capture technologies7,8,9. However, high-energy requirements in the regeneration or CO2 desorption process is the largest obstacle for preventing its industrial applications10. Conventional regenerations of spent solvents are just realized by heating spent solvents above boiling temperatures (100–150 °C). The energy consumption needed for CO2 desorption or spent sorbent regeneration accounts for ~15–30% of power plants’ electricity outputs11.

Accordingly, many other efforts have been made for lowering CO2 desorption energy consumptions and up to date, there mainly exists three important strategies10,11,12. The most popular one is to use new organic amine mixtures for CO2 sorption and desorption. It was generally believed that adsorption solvents with lower heat of absorption require less heat to be regenerated. Thus, much attention has been focused on the blending of different types of amine solutions12,13,14. The other two strategies include the application of novel equipment with superior mass and heat transfer performance as well as process optimizations12, 15. For example, stripping CO2 from aqueous potassium carbonate solutions was achieved by using two types of polymeric flat sheet microporous membrane contactors, which have been reported by Michael and co-workers16. Unfortunately, the efforts made in promoting CO2 desorption or spent sorbent regeneration are still much less than that improving CO2 absorption.

On the other hand, it should be pointed out that amine solvents, such as monoethanolamine (MEA), diethanolamine (DEA) and methyldiethanolamine (MDEA), have several serious shortcomings such as toxicity, corrosiveness, and oxidative degradation17, 18. Alternative sorbents are inorganic carbonate solutions based on Equation R1, in which M stands for Na or K.

$${{\rm{M}}}_{2}{{\rm{CO}}}_{3}+{{\rm{CO}}}_{2}+{{\rm{H}}}_{2}{\rm{O}}\leftrightarrow 2{{\rm{MHCO}}}_{3}.\,$$
(R1)

R1 based CO2 captures are not only inexpensive but also stable and thus environmentally friendly. However, like any other CO2 sorbents, its slow CO2 desorption kinetics make its wide utilization unaffordable, considering the fact that the energy consumption of the CO2 desorption step accounts for ~70–80% of those needed for the corresponding overall CO2 capture processes19.

This research was designed to fill the gap with its focus on investigating the significant effect of TiO(OH)2 on promoting aqueous NaHCO3 decomposition or CO2 desorption kinetics and the cost of Na2CO3/NaHCO3 based CO2 capture technology. The success could lead to a new pathway for future CO2 capture technology development with its focus on catalysis instead of high-capacity but expensive sorbents.

Results

Effects of several key factors on the performance of TiO(OH)2 on NaHCO3 decomposition, essentially CO2 desorption

Stirring rate

The experimental set-up for aqueous NaHCO3 decomposition or CO2 desorption test is illustrated in Fig. 1. The photo of actual experimental set-up is provided in supporting information (SI Photo 1). The quantities of CO2 desorbed under different stirring rates at the given temperature (70 °C) were first evaluated. As shown in Fig. 2, CO2 desorption amounts increase with the increase in stirring rates ranging from 400 rpm to 550 rpm while stirring rate only shows slight effect when it increases from 600 rpm to 650 rpm. This means that the mass transfer resistance of CO2 diffusing from liquid phase to gas phase was intensively reduced and thus, CO2 desorption is dominated by its reaction kinetics when stirring rate is higher than 600 rpm under experimental conditions.

Figure 1
figure 1

Schematic diagram of CO2 desorption set-up [1: N2 cylinder; 2: CO2 cylinder; 3: mass flow controller; 4: thermostatic water bath; 5: thermoelectric couple; 6: stirred tank reactor; 7; condenser; 8: desiccator; 9; gas concentration analyzer; 10: data recording unit].

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Figure 2
figure 2

NaHCO3 decomposition or CO2 desorption profiles under different stirring rates [initial NaHCO3 concentration: 0.143 mol/L; Ti/Na molar ratio: 0.78; reaction temperature: 70 °C; N2 flow rate: 500 mL/min].

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Quantity of TiO(OH)2 added

The effect of the quantity of TiO(OH)2 or the Ti/Na molar ratio in solution on the decomposition of NaHCO3 is shown in Fig. 3. NaHCO3 decomposition or CO2 desorption significantly increases with the quantity of added TiO(OH)2 or the Ti/Na ratio. However, the improvement sensitivity decreases with Ti/Na ratio. Under the test conditions, the highest improvement in NaHCO3 decomposition or CO2 desorption reaches 800% at 110 s of reaction time. Considering the fact that the quantity of H2O in reactor is much more than that stoichiometrically needed, the promotional effect of TiO(OH)2 on NaHCO3 decomposition or CO2 desorption is exceptional.

Figure 3
figure 3

Effect of Ti/Na molar ratio on NaHCO3 decomposition or CO2 desorption [initial NaHCO3 concentration: 0.238 mol/L; stirring rate: 600 rpm; reaction temperature: 65 °C; N2 flow rate: 500 mL/min].

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Temperature

The effects of TiO(OH)2 on NaHCO3 decomposition or CO2 desorption kinetics at temperatures of 40 °C, 50 °C, 60 °C and 70 °C are presented in Fig. 4. The quantity of CO2 released within the first 180 seconds with the use of TiO(OH)2 at 40 °C is 2.79 mmol, ~510% higher than that obtained without use of TiO(OH)2 under the same condition. At 70 °C, the catalyst can improve NaHCO3 decomposition by about 490% within the first 180 s. The catalytic effects gradually decrease with time as dictated with thermodynamic theories, although they are still considerably large after 600 s of NaHCO3 decomposition. For example, 5.63 mmol of CO2 was released with the use of TiO(OH)2 at 40 °C after 600 s, while only 1.61 mmol of CO2 was released without presence of TiO(OH)2under the same conditions, a 71.4% decrease.

Figure 4
figure 4

Effect of temperature on catalytic NaHCO3 decomposition or CO2 desorption [initial NaHCO3 concentration: 0.143 mol/L; Ti/Na molar ratio: 0.78; stirring rate: 600 rpm; N2 flow rate: 500 mL/min].

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Cyclic performance of TiO(OH)2 on NaHCO3 decomposition

In addition, the regeneration performance of NaHCO3 solution decomposition using TiO(OH)2 plays a key role in practical applications and cannot be neglected, which has also been examined in detail in this work. The fresh TiO(OH)2 and CO2-treated TiO(OH)2 or cycled TiO(OH)2 are hereafter denoted as F-TiO(OH)2 and C-TiO(OH)2, respectively. The C-TiO(OH)2 promoted CO2 release curves obtained from aqueous NaHCO3 are almost overlap and no remarkable decrease in the total CO2 desorption amounts are noticed in five continuous cycles as shown in Fig. 5. Furthermore, considering the resulting pore structure data summarized in Table 1, the specific surface area and pore volume of C-TiO(OH)2 after 5 cycles of uses are 693.1 m2/g and 0.658 cm3/g, respectively, are comparable to those of F-TiO(OH)2. This suggests that TiO(OH)2 is a stable catalyst for NaHCO3 decomposition or CO2 desorption.

Figure 5
figure 5

Cyclic performance of TiO(OH)2 on NaHCO3 decomposition [initial NaHCO3 concentration: 0.238 mol/L; Ti/Na molar ratio: 0.78, stirring rate: 600 rpm; reaction temperature: 60 °C, N2 flow rate: 500 mL/min].

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Table 1 BET characteristics of F-TiO(OH)2 and C-TiO(OH)2.
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Discussion

N2 adsorption/desorption isotherms and pore-size distributions were conducted to investigate the pore structures of F-TiO(OH)2 and C-TiO(OH)2 with 5 cycles of CO2 desorption and sorption. The resulting pore structure data were summarized in Table 1. Both samples are nanoporous with the average diameter close to 1.7 nm. The specific surface area of F-TiO(OH)2 could reach to as high as 807.4 m2/g, while that of C-TiO(OH)2 is 693.1 m2/g, about a 14% drop.

The FT-IR absorption spectra of two TiO(OH)2 samples, Na2CO3 and NaHCO3 are shown in Fig. 6. Both TiO(OH)2 samples show broad peaks in the 400–900 cm−1 range corresponding to Ti-O bending. The additional peaks in the 3,000–3,600 cm−1 range existing in C-TiO(OH)2 sample are likely due to the partial hydration during the reaction process. Typically, two small peaks at 1,124 and 1,072 cm−1 in the F-TiO(OH)2 likely result from Ti-O-C, the ending and bridging isopropyl groups, considering that TiO(OH)2 are directly synthesized from the hydrolysis of TTIP. Similar phenomena were also reported by Sui and coworkers20. Small peaks at 1,132, 1,117 and 1,022 cm−1 were associated with Ti-O-C and butoxyl groups. In addition, it should be noted that C-O-H bending vibrations could appear as a broad and weak peak at 1,440–1,220 cm−1 in alcohols and phenols21, 22. Accordingly, the two small peaks observed at 1,124, 1,072 cm−1 for F-TiO(OH)2 and C-TiO(OH)2 may also be related to Ti-O-H group, which may be the major player in accelerating NaHCO3 decomposition or CO2 desorption.

Figure 6
figure 6

FT-IR spectra of F-TiO(OH)2, C-TiO(OH)2, Na2CO3, and NaHCO3.

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Moreover, the FT-IR spectra of the NaCO3 and NaHCO3 obtained in experiments are consistent with those in references23,24,25. The noteworthy peaks at 1,300 cm−1 and 1,400 cm−1 for NaHCO3 and Na2CO3, respectively, are due to carbonate asymmetric stretching26. In addition, a peak at 1,605 cm−1 was observed for NaHCO3, which can be attributed to symmetric stretching of CO2. Therefore, the peak at 1,636 cm−1 for C-TiO(OH)2 may be due to the presence of CO3 2− or HCO3 −. This indicates that NaHCO3 or NaCO3 may be sorbed on the inner pore of F-TiO(OH)2, which could be responsible for the decrease in specific surface area of C-TiO(OH)2 as observed in Table 1.

The TGA data of F-TiO(OH)2 and C-TiO(OH)2 are shown in Fig. 7. Both samples started to lose absorbed water at ~120 °C, and then decomposed to TiO2 above 300 °C. The final masses of F-TiO(OH)2 and C-TiO(OH)2 account for 80.1% and 81.2% of their initial masses, respectively, indicating that no titanium loss took place during the NaHCO3 decomposition or CO2 desorption process in this research. F-TiO(OH)2 shows slower mass loss rates at lower temperatures, which could be due to loss of organic groups27 from TTIP. This agrees with what was observed in the FTIR results.

Figure 7
figure 7

TGA data of F-TiO(OH)2 and C-TiO(OH)2 [heating rate: 20 °C/min; N2 flow rate: 100 mL/min].

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Furthermore, Raman spectroscopy analyses of F-TiO(OH)2, C-TiO(OH)2, pure Na2CO3 and NaHCO3 were also conducted and the results are shown in Fig. 8. All samples have a similar broad band in the range of 3,000–3,200 cm−1, which resulted from H-O vibration in the water absorbed from air28. Two very strong peaks at ca 1,043 cm−1 for NaHCO3 and at ca 1,079 cm−1 for Na2CO3 were observed, which has been reported by others29, 30 and can both be assigned to the ν 1 symmetric stretch.

Figure 8
figure 8

Raman spectra of F-TiO(OH)2, C-TiO(OH)2, pure Na2CO3 and NaHCO3.

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Both TiO(OH)2 samples show relatively intense Raman bands at ca 446 cm−1 and ca 636 cm−1 (ca 584 cm−1 for reacted TiO(OH)2). Typically, the frequencies of the Raman bands observed are 513 cm−1 and 636 cm−1 for anatase and 446 cm−1 and 609 cm−1 for rutile, respectively29, 31, 32. The TiO(OH)2 sample in this work is amorphous and the position and intensity of its characteristic peaks can be changed by modifying its material structures. It should be pointed out that the resulting Raman spectrum for C-TiO(OH)2 is quite similar to that of F-TiO(OH)2, implying that TiO(OH)2 is stable or can be cyclically used.

Finding a cost-effective method for carrying the released or desorbed CO2 out of reaction system is important. A pump was used to compare its performance on carrying the released CO2 out of its desorption system to that achieved with N2 as a carrier gas, and the results obtained within 55–65 °C are shown in Fig. 9. It can be seen that the quantity of CO2 carried out with N2 was high than that with a pump when NaHCO3 decomposition was not catalyzed with TiO(OH)2. However, the pump carried out 25% more CO2 than the same amount of N2 did when the catalyst was used. The quantity of the CO2 generated from non-catalytic NaHCO3 decomposition with 200 ml/min N2 being a carrier gas is only ~1/8 of that from catalytic NaHCO3. Furthermore, when a 12 W pump was used for carrying the generated CO2 out of NaHCO3 decomposition reactor, the CO2 releasing improvement due to the help of the catalyst is about 25 times or 2,500%.

Figure 9
figure 9

(a) Effect of different CO2 carrying out methods on CO2 release [initial NaHCO3 concentration: 0.238 mol/L; Ti/Na molar ratio: 0.78; stirring rate: 600 rpm; reaction temperature: 65 °C, N2 flow rate: 200 mL/min; power of pump: 12 W]. (b) Ratio of CO2 released by using N2 gas and the pump with and without uses of catalyst for each 10 seconds.

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Figures 3, 4 and 9 clearly show that noticeable CO2 release from catalytic NaHCO3 decomposition was observed earlier than that from non-catalytic NaHCO3 decomposition. This is not only confirmed with the experimental set-up shown in Fig. 1 but also with the data collected with the FTIR function of the available in-situ FTIR-MS (SI Photo 2) for solid NaHCO3 decomposition and shown in Fig. 10. The time needed for the appearing of the peak of C=O in the released CO2 of catalytic NaHCO3 decomposition is longer than that of non-catalytic NaHCO3 decomposition.

Figure 10
figure 10

FT-IR- characterization of catalytic and non-catalytic solid NaHCO3 decomposition [NaHCO3: TiO(OH)2 mole ratio: 1:2; temperature: 150 °C; He flow rate: 5 ml/min].

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Conclusion

Nanoporous TiO(OH)2 is a very effective catalyst for aqueous NaHCO3 decomposition or CO2 desorption, and thus CO2 capture. The finding could be significantly helpful for reducing the global concern about the high energy consumption required for CO2 emission control. Further works should focus on understanding the associated mechanism and extending the new concept to other CO2 emission and utilization technologies development.

Methods

TiO(OH)2 preparation

Titanium isopropoxide (TTIP), ethanol (EtOH) and sodium bicarbonate (NaHCO3) were purchased from Sigma-Aldrich and used without any further purification prior to preparing TiO(OH)2. The first TiO(OH)2 preparation step was to add 25 mL titanium isopropoxide to 350 mL deionized water at room temperature, followed by 4 hours of vigorous stirring. The resulting white precipitate was separated by filtration, and then washed several times with deionized water and anhydrous ethanol sequentially. The wet TiO(OH)2 was dried under 120 °C in an oven for 12 h.

TiO(OH)2 characterization

The Brunauer-Emmett-Teller (BET) surface areas and pore size distribution of TiO(OH)2 samples were measured by using a Quantachrome Autosorb-iQ pore structure analyzer. Pore volumes were estimated from the adsorbed amount of N2 at a relative pressure of P/P0 = 0.99. Fourier transformed-infrared (FTIR) spectroscopy data were collected with a Thermo Nicolet Magna-IR 760 spectrometer (SI Photo 1). Thermal gravity analysis (TGA) tests were conducted on a TA Instruments SDT Q600 thermogravimetric analyzer. Raman spectra of the TiO(OH)2 samples were collected on a Raman Sierra IM-52 instrument from Snowy Range Instruments with a 532 nm laser and 3 mW power.

NaHCO3 decomposition or CO2 desorption tests

NaHCO3 decomposition tests were started by adding predetermined amounts of TiO(OH)2, NaHCO3 and deionized water into the reactor stirred at the rate of 600 rpm. Reaction temperatures were regulated using a digital temperature controller. During NaHCO3 decomposition, pure N2 or vacuum pump (Karlsson Robotics, 12 W) was used to carry the desorbed CO2 out of reactor. The change of CO2 outlet concentration with time was recorded in a data acquisition unit, and the concentration-time profile was used to calculate the quantity of CO2 desorbed and evaluated to observe the difference in CO2 desorption due to the use of TiO(OH)2. When TiO(OH)2 was used for cyclic NaHCO3 decomposition or CO2 desorption tests, the used TiO(OH)2 and deionized water were mixed in a closed tank with pure CO2 at 200 psi followed by vigorous stirring under room temperature to study the effect of CO2 sorption or activation on TiO(OH)2.