In order to avoid the disaster brought by greenhouse effect, numerous methods have been suggested to reduce the CO2 concentration in the atmosphere. In following section, we would discuss two of the most popular methods of CO2 sequestration.
1.3.1. Absorption/Adsorption
There are many absorbents available to retain CO2. They include polymer membranes, zeolites, perovskites, magnesia, and sodalime. Pfeiffer et al. used commercial Li2O to absorb CO2 at different temperatures. Result from the thermogravimetric analyses (TGA) on the Li2O in a CO2 flux indicated an increase of weight about 14.3 wt% between 190 and 400 °C due to Li2O conversion into Li2CO3. When the reaction temperature was increased to 600 °C, more Li2O became LiCO3 and the weight increased dramatically raised to 226 wt% (86 mol%), corresponding to 1.26 g of CO2 per gram of Li2O. Figure 1.9 depicted the mechanism proposed for the CO2 absorption on Li2O and the SEM image of Li2O before and after heat treatment at 600 °C for 2 hours in a CO2 flux. [19]
Pfeiffer et al. also prepared Li2-xNaxZrO3 by a heat treatment of Li2CO3, Na2CO3, and Zr(OCH3)4 at 900 °C to absorb CO2. They determined the CO2
absorption of Na2ZrO3 was higher than that of Li2ZrO3. Results from the thermogravimetric analyses indicated that the Li2ZrO3 absorbed CO2 with an increase of weight about 4 wt% after the reaction. Similarly, the Li1.8Na0.2ZrO3 increased its weight to 6.9 wt% and Li1.4Na0.6ZrO3 increased its weight to 13.1 wt%, respectively. Comparing to other samples they prepared, the LiNaZrO3 demonstrated the highest CO2 chemisorption efficiency of 75.3
% at 600 °C.[20]
Hausler et al. estimated the influence of SO2, N-methyldiethanolamine (MDEA), and triethanolamine (TEA) on the CO2 absorption capacity by utilizing aqueous 2-(2-aminoethylamino)ethanol (AEE) solution and its blends with MDEA and TEA to absorb CO2 or CO2/SO2 mixtures at 23 °C. They
found out that the additions of 5 and 10 wt% of MDEA and TEA exerted negligible influence on the CO2 absorption in AEE solution. Furthermore, adding MDEA increased the CO2 absorption capacity of AEE slightly, whereas adding TEA decreased the CO2 absorption capacity of AEE in the absence or presence of SO2. They obtained the highest CO2 absorption capacity about 1.267 mol of CO2 per mol of amine by using 15 wt% AEE + 10 wt%
MDEA.[21]
In order to determine the best adsorbent for CO2, Snap et al. prepared nitrogen enriched carbons by urea-formaldehyde (UF) and melamine-formaldehyde (MF) in the presence of K2CO3. The K2CO3 acted as a chemical activation agent activating the reaction over a range of temperatures from 400 to 700 °C. The UF with an activation temperature of 500 °C presented the highest CO2 capacity, capturing over 8 gram CO2 per 100 gram of adsorbent at 25 °C. Higher activation temperature resulted in higher surface area, but did not improve CO2 capturing ability, suggesting that sites suitable for the adsorption of CO2 were destroyed at higher temperature.[22]
1.3.2. Electrochemical reduction
Electrochemical reduction of CO2 has attracted much attention because it might be a promising method for turning CO2 to useful materials. Some common products of CO2 from the electrochemical reduction are listed below.
The equilibrium potentials of each reaction have been reported by Sullivan et al.
under the standard conditions against NHE.[23]
2CO2 + 2H+ + 2e- → H2C2O4 E0 = -0.475 (1.1)
Fujishima et al. used various metal wires (Ti, W, Ni, Pd, Pt, Cu, Ag, Zn, Sn, and Pb) as the working electrode, Pt wire as the counter electrode, and 0.3 M tetrabutylamonium perchloride (TBAP) in methanol as a supporting electrolyte to conduct electrochemical reduction of CO2 galvanostatically at 25
°C, 41 atm, mainly at 200 mAcm-2. The electrochemical reduction products were analyzed by flame ionization detector (FID) and thermal conductivity detector (TCD). They found out that W, Ti, and Pt electrodes did not possess the required ability in electrochemical reduction of CO2. Formate was produced at Sn and Pb electrodes, but much more CO was also observed. In contrast, electrolysis at Ag, Zn, and Pd electrodes yielded CO mostly. The Cu electrode revealed better ability to form hydrocarbons in aqueous electrolyte than in methanol system. However, the hydrocarbon formation at Ni electrode was more efficient in methanol than that of the aqueous system.[24]
Fujishima et al. also used RuO2 deposited on boron-doped diamond (BDD) as the working electrode, the Pt used as the counter electrode, and SCE as the reference electrode to reduce CO2. They adjusted the pH value of the solution by NaOH and they obtained the optimized efficiency for CO2 reduction (almost 80 %) at pH of 4, which was close to saturated aqueous solution of carbonic
acid. The applied potential was -0.55 V and the current density was about -0.45 mAcm-2. In their study, their found out the use of BDD as a substrate for the RuO2 layers resulted in much lower Faradic efficiency for CO2
reduction to methanol as compared to using the TiO2 as a substrate.[25]
Hori et al. prepared the silver-coated ion exchange membrane (solid polymer electrolyte, SPE) electrodes through the electroless method to deposit silver onto the ion exchange membrane. The SPE electrode was used as the working electrode, a Pt plate was used as the counter electrode, and 0.2 M K2SO4 was adopted as the electrolyte. Ag/SPE prepared from an anion exchange membrane (AEM) reduced CO2 to CO and HCOOH for more than 2 hours. They controlled the current density at 50 mAcm-2 and obtained average electrode potential about -1.8 V (vs. SHE) when the reduction process was performed at the AEM electrode with silver coated two layers. However, the SPE electrode system prepared from the cation exchange membrane (CEM) was not suitable for CO2 reduction, since OH-, HCO3-, and CO32- formed in the CO2 reduction could not be removed from the metal membrane interface.[26]
Ogura et al. prepared copper(I) halide-confined Cu-mesh electrodes by electrochemical oxidation (applied potential: 0.2 to 0.4 V) of HCl, KBr, and KI, respectively. The electrolysis potential for the electrochemical reduction of CO2 was -2.4 V (Ag/AgCl). When they used the CuCl as working electrode, they chose 3 M KCl as electrolyte and recorded current density about 46 mAcm-2. When they used the CuBr as working electrode, they selected 3 M KBr as electrolyte and obtained current density about 37 mAcm-2. Although the CuBr electrode presented lower current density, it conversed 24.3 % of CO2
more than that of CuCl electrode (17.1 %). Table 1.1 presents the Faradaic efficiencies for the products obtained in the electrochemical reduction of CO2
at -2.4 V (Ag/AgCl) on a copper(I) halide-confined Cu-mesh electrode. [27]
Köleli et al. used the electrodeposition method to deposit polypyrrole on the platinum as working electrode for electrochemical reduction of CO2. The electrolyte was MeOH/0.1 M LiClO4/H+/H2O, the applied potential was -0.4 V (Ag/AgCl) and the obtained current density was smaller than 13 mAcm-2 under ambient condition. When the electrochemical process was operated under ambient condition, only a minute amount of CO2 became HCOOH, CH3COOH, and HCHO. However, they obtained much more HCOOH, CH3COOH, and HCHO under high pressure (20 bar).[28]
Kaneco et al. used copper the electrode in methanol with sodium supporting salts to reduce CO2 by electrochemical method. The reduction process was investigated with various sodium supporting salts, such as acetate, chloride, bromide, iodide, thiocyanate, and perchlorate, at a low temperature (-30 °C). The best results they obtained were utilizing 0.5 M NaClO4
(methanol-based) electrolyte at -3.0 V (Ag/AgCl). The current density was 27 mAcm-2 and the faradic efficiency of methane was 70.5 %.[29]
Koleli et al. studied the electrochemical reduction of CO2 on Pb and Sn electrodes in aqueous KHCO3 and K2CO3 electrolyte in a fixed-bed reactor.
The highest current efficiency for formic acid, the predominant product, obtained in 0.5 M KHCO3 at -1.5 V (SCE) after 30 min electrolysis was found to be 90 % for Pb electrode and 74 % for Sn electrode. Meanwhile, the current efficiency for the formic acid in 0.1 M K2CO3 at -1.5 V (SCE) after 30 min electrolysis was found to be 39 % for Pt electrode. Figure 1.10 depicted the image of the electrochemical fixed-bed reactor, the current-potential diagram and the Faradaic current efficiency–potential diagram for formic acid formation on Pb electrode in 0.1 M K2CO3 at different time intervals.[30]
1.4. Properties and applications of photocatalysts
Photochemistry is the subject studying the relationship between light and molecules. Light is the common name for electromagnetic (EM) radiation in the visible, near-ultraviolet, and near-infrared spectral range. The electromagnetic spectrum includes a variety of radiations from very long radio waves with the dimension of buildings to very short gamma rays which are much smaller than an atom nucleus. In the wave model, the frequency (λ) is inversely proportional to the wavelength (ν) according to the equation:
c = λν (1-7)
The value of c is constant (2.998 × 108 ms-1 in vacuum).[31]
In the quantum model, the photon is used to describe the quantized energy of an electromagnetic wave. A photon has no mass but it has a specific energy (E) directly proportional to the frequency (ν) of the radiation, according to the Planck relation:
E = hν (1-8)
Where h is the Pranck constant (6.626 × 10-34 J．s).[31]
The reactions induced by light are defined as photochemical reactions.
The first step of a photochemical process is the photoexcitation (the mechanism of electron excitation by photon absorption), where the reactant is elevated to an excited state possessing a higher energy than that of ground state.
M + hν → M* (1-9)
where M is the molecule at ground state, M* is the molecule at excited state, and hν is the photon energy. [31]
The molecule in the excited states could return to the ground state by various processes.
M* → M + hν’ (1-10)
M* → M + heat (1-11)
M* + Q → M+ Q’ (1-12)
where Q is the molecule that absorbed the excited energy of M*.
The photochemical reaction could be categorized by the usage of photosensitizing materials (photocatalyst). If the initial reactant could not absorb light energy or the light energy could not derive sufficient energy for photochemical reaction, the photocatalysts were added to absorb light and promote the desirable photochemical reaction. In principle, the photochemical reaction with photocatalysts added may proceed on the surface of a semiconductor through several steps. First, electron-hole pairs are created by exciting the semiconductor with light or suitable energy. Second, isolation of the electrons and holes on the semiconductor surface takes place.
Third, the separated electrons and holes would initiate individual redox process with the reactants adsorbed on the surface. Finally, the products are released and the surface reconstructed.[31]
Yang et al. prepared the Cu2O nanoparticles with diameter of 35 nm via the electrochemical method in alkaline NaCl solutions with copper as electrode
and K2Cr2O7 as additive. Electrolysis was performed under stable current densities (50, 70, 90, 100, and 110 mAcm-2) at 70 °C for 1 hour and the Cu2O nanocrystal prepared under 100 mAcm-2 was chosen as the catalyst. They observed thtat 97 % of 50 mgL-1 methyl orange (MO) was decomposed under a 125 W high-pressure mercury lamp for 2 hours or under sunlight for 3 hours when Cu2O in 2gL-1 was added. In contrast, pure TiO2 and CdS photocatalysts were effectively only under UV irradiation.[32]
Andronic et al. prepared the TiO2 film by Spray Pyrolysis Deposition (SPD) in order to study the influence of the TiO2 in specific surface (powder, film) on the photocatalytic degradation of MO. The photcatalysis process was operated under an 18 W fluorescence lamp by adding 1 g TiO2 powder per 1 L solutions with different MO concentrations. At higher MO concentrations, the light penetration was reduced which was due to heavy MO adsorption on the TiO2 and fewer photons were able to reach the catalyst surface. When the film and powder were 0.004 gL-1 and the MO was 7.8125 mgL-1, the film of TiO2 demonstrated photocatalytic efficiency about 5.10 % after 6 hours of reaction, this value was lower than that of TiO2 (efficiency: 7.12 %).[33]
Liu et al. synthesized the BiFeO3 (BFO) nanoparticles via a sol-gel method and studied their photocatalytic abilities through decomposition of MO.
They dissolved bismuth nitrate and iron nitrate within 2-methoxyethanol and added polyethylene glycol as a dispersant. The mixture was calcinated under 500 °C for 2 hours to form perovskite-type BFO. The initial concentration for MO was 15 mgL-1 with a catalyst loading of 30 mmolL-1 (11.545 gL-1) and more than 90 % of MO was decolorized after 8 hours of irradiation under a 300 W Xe lamp. Figure 1.11 presents the photocatalytic ability of bulk and nanoparticlate BFO on degradation of MO under UV-vis light irradiation and
visible light irradiation. [34]
Parida et al. prepared the hydrated titanium oxide by a sol-gel approach, adopting titanium isopropoxide as starting material. The as-prepared TiO2
nanoparticles were made into a series of sulfated TiO2 samples via an aqueous wetness impregnation method with various weight percentages of SO42-. The photocatalytic degradation of MO was carried out under sunlight with a solar intensity about 800 Wm-2 by taking 20 mL of 150 mgL-1 MO solution with 1.0 gL-1 catalyst. The samples loaded with 2.5 wt % SO42- indicated higher degradation ability than without or less loading and its behavior may be due to the addition of sulfate that effectively decreased the crystal size of the TiO2. By adjusting the pH value of the solution from 8 to 2, they found out the percentage of degradation was increased with decreasing pH values. Since the surface of the sulfate-modified TiO2 became positively charged at pH lower than 4.5 to 5.0 and MO was an anionic dye, the photocatalytic reaction was likely to be faster at acidic pH. Figure 1.12 depicts the mechanisms occurring on the TiO2 surfaces exposed to light for the photodegradation of organic pollutants.[35]
Figure 1.1. The crystal structure of Cu2O.[1]
Figure 1.2. SEM images of Cu2O deposited on ITO substrates from electrolyte containing 0.02 M Cu(Ac)2, 0.1 M NaAc and CTAB with different concentrations: (a) 0, (b) 0.4 mM, (c) 0.8 mM, and (d) 2.8 mM.[4]
Figure 1.3. TEM images for Cu2O nanocubes fabricated by adding the mixture of AA and NaOH into the solution containing Cu2+ and (a) 4 mL, (b) 2 mL, and (c) 1 mL of 0.05 M PEG.[6]
Figure 1.4. SEM images (a), (b), and XRD pattern (c) of hollow Cu2O spheres produced with a NaOH titration rate of 0.25 mLmin−1.[7]
Figure 1.5. SEM (columns 1 and 2) and TEM (columns 3 and 4) images of the
Cu
2O nanocubes for samples A to F (seed to transfer 5 times).
[8]Figure 1.6. SEM images ((a) to (e)) and EDX spectra (d
1and d
2) for the samples isolated in the time-dependent experiments with glutamic acid.
[9]Figure 1.7. SEM images of Cu2O particles prepared under different pH values:
(a) pH of 6.5, (b) pH of 6.7, (c) pH of 7.5, and (d) pH of 8.0.[10]
Figure 1.8. (a) TEM image (inset is the SAED patterns) and (b) HRTEM image of the Cu2O nanorods via solvothermal treatment of CuSO4·5H2O and NaOH in a mixed solution of ethanol and deionized water at 140 °C for 10 hours.[11]
Figure 1.9. SEM image of Li2O before treatment (A), heat treated at 600 °C for 2 hours in a CO2 flux (B), and scheme of the mechanism proposed for CO2
absorption on Li2O (C).[19]
Table 1.1. Faradaic efficiencies for the products obtained in the electrochemical reduction of CO2 at -2.4 V (Ag/AgCl) on a copper(I) halide-confined Cu-mesh electrode.[27]
Figure 1.10. (a) The IV curves for CO2 reduction on Pb electrode in 0.1 M K2CO3 at various time periods. (b) Faradaic current efficiency–potential diagram for formic acid formation on the Pb electrode in 0.1 M K2CO3 at different time intervals; (■) 30, (●) 60, (▲) 90, and (▼) 120 min. (c) The image of the electrochemical fixed-bed reactor.[30]
Figure 1.11. Photocatalytic ability of bulk and nanoparticulate BFO on degradation of methyl orange under irradiation of UV-vis and visible light.[34]
Figure 1.12. The mechanisms occurring on TiO2 surfaces exposed to light for the photodegradation of organic pollutants.[35]