CO2 Capture from Flue Gas Using an Electrochemically Reversible

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Environmental and Carbon Dioxide Issues

CO2 Capture from Flue Gas Using an Electrochemically Reversible Hydroquinone/Quinone Solution Chuanliang Huang, Changjun Liu, Kejing Wu, Hairong Yue, Siyang Tang, Houfang Lu, and Bin Liang Energy Fuels, Just Accepted Manuscript • DOI: 10.1021/acs.energyfuels.8b04419 • Publication Date (Web): 06 Mar 2019 Downloaded from http://pubs.acs.org on March 15, 2019

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CO2 Capture from Flue Gas Using an Electrochemically Reversible Hydroquinone/Quinone Solution Chuanliang Huang,† Changjun Liu,† Kejing Wu,‡ Hairong Yue,† Siyang Tang,† Houfang Lu,†,‡ Bin Liang*,†,‡ †School of Chemical Engineering, Sichuan University, Chengdu 610065 (PR China) ‡Institute of New Energy and Low-carbon Technology, Sichuan University, Chengdu 610207 (PR China) Graphic Abstract

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KEYWORDS: CO2 capture, electrochemical method, hydroquinone and quinone, CO2 desorption.

ABSTRACT: Electrochemical methods are potentially an energy-saving way to capture CO2 from flue gas. Unlike the regeneration of amine through high-temperature thermal distillation in commercial CO2 absorption processes, the CO2 absorbent is regenerated by electrolysis at low temperature. In this work, tiron (disodium 4,5-dihydroxy-1,3benzenedisulfonate, QH2) was employed as a pH mediator since its redox reactions can change the pH of the solution. Na2Q, which was prepared by QH2 and NaOH in a molar ratio of 1:2, was used to capture CO2 because of its alkalinity. Then CO2 was desorbed by the oxidation of QH- and QH2 formed in the CO2-rich aqueous solution to quinone (Q) to release protons, and the alkalinity was recovered by the reduction of Q to the quinone dianion (Q2-). The redox performance of Na2Q in aqueous solution was investigated using cyclic voltammetry, and the CO2 capacity of Na2Q solutions at different concentrations (0.1-0.7 M) was measured. The results show that the redox behavior of Na2Q was reversible in the neutral or weakly alkaline solutions. However, the reduction of Q to Q2- by electrolysis was difficult in a high pH solution. During adsorption, the Na2Q solution absorbed CO2 at a molar ratio of about 1:1 (CO2/Q2-≈1). The CO2-saturated Na2Q solution was electrolyzed in the anode zone under constant current. The CO2 desorption rate reached 100%, and Q2-, QH-, and QH2 were oxidized to give Q. In the cathode zone, Q was reduced to Q2-, which could be used to adsorb CO2 from flue gas. Based on the potential difference between the cathode and the anode, the regeneration energy consumption was estimated to be about 2.4 GJ/t CO2. 2 ACS Paragon Plus Environment

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INTRODUCTION The increasing atmospheric CO2 concentration is considered to the major factor contributing to global warming. Global coal combustion contributed 46% of the total CO2 emissions from fossil fuel combustion, with 31% being emitted from coal-fired power plants.1 Therefore, effective capture, storage, or utilization of CO2 from the lowconcentration flue gas of power plants has become an urgent challenge. CO2 capture methods from power plants have been extensively investigated, including the mineralization and storage,2-5 solid adsorption,6-8 cryogenic separation,9 membrane absorption and separation of CO2.10-12 Most current commercialized technologies are based on organic amine absorption. Amine absorbents have fast absorption rates and high capture ratios, but their regeneration processes are very energy-intensive.13, 14 The ethanolamine (MEA) capture process, for example, generally requires heat loads as high as 4.0 GJ/t CO2 in the reboiler to desorb CO2 by distillation at 110-130 ℃; the desorption thus accounts for 57.5% of the total energy consumption.15, 16 Novel energysaving technologies17-19 and new chemical absorbent systems20-22 have been reported to cut the energy consumption of CO2-rich solution regeneration.

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Figure 1. Electrochemical regeneration for the CO2 capture process The regeneration heat duty in organic amine absorbent processes includes the reaction heat (59%), sensible heat (19%), and latent heat of evaporation (22%). Electrochemical regeneration processes can significantly reduce the sensible heat and evaporation heat by operating the electrolysis desorption process at low temperature. Figure 1 demonstrates the electrochemical desorption of CO2 and regeneration of the CO2 absorbent. The oxidation reaction in the anode chamber releases protons, which raise the acidity of the anolyte and result in CO2 desorption. When the oxidized anolyte flows into the cathode chamber, the absorbent is regenerated and used to absorb CO2 from flue gas in the absorption tower.

Scovazzo et al. 23 used quinone to separate CO2; the reduced quinone (Qred) combined with CO2 to form a stable Qred-CO2 compound to absorb CO2, and was subsequently oxidized on the anode to form oxidized quinone (Qox) and release CO2. The process increased the CO2 concentration from 1% (0.425 kPa) to nearly pure (85.3 kPa). 4 ACS Paragon Plus Environment

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However, the low solubility of CO2 and quinone (Q) in the solvent, as well as the adverse effects of the O2 and H2 in the flue gas on the system, reduced the current efficiency of the overall process. Alan Hatton et al.24 utilized a quinone redox active carrier in an ionic liquid system to electrochemically separate CO2. The quinone dianion (Q2-) was used to capture CO2 as it can bind CO2 to form Q (CO2)2-. Then this Q (CO2)2is oxidized at the anode to desorb CO2 and Q is formed at the same time. Finally, Q2was regenerated by the reduction of Q at the cathode. The 1,4-napthoquinone has a high solubility in 1-ethyl-3-methylimidazolium tricyanomethanide system, which improves the efficiency of CO2 separation from the dilute gas mixture.

Stern et al. 25, 26 used copper ions to enhance the desorption rate of CO2 in the amine absorption process. The CO2 desorption rate was increased from 50% to 80-90% by the formation of a stable copper-amine complex. When the CO2-rich amine flowed through the anode, Cu2+ was produced by Cu anode electrolysis and then reacted with carbamate to form a copper-amine complex that facilitated the release of CO2. The copper-amine complex was subsequently reduced to Cu on the cathode and the organic amine was regenerated. The obvious shortcoming of the process was that both the anode and cathode were irreversible.

pH change is a very common feature of redox reactions. Electrochemical redox reactions of metal ion systems, and organic chemical systems can create a pH gradient.27

Watkins et al. 28 designed a liquid membrane device for CO2 separation.

The solution was a mixture of hydroquinone, quinone, and sodium bicarbonate. At the 5 ACS Paragon Plus Environment

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cathode, the reduction of quinone increased the pH of the catholyte, thus causing the absorption of CO2 to form HCO3-. At the anode, the oxidation of hydroquinone released H+, which then reacted with HCO3- to desorb the CO2. Nonetheless, owing to the low aqueous solubility of their quinones and the small change in the pH, their separation efficiency was not very good.

In this work, tiron, the hydroquinone compound (see Figure S1), was investigated as a pH meditator to create a pH gradient for the absorption and electrochemical desorption of CO2.Tiron is a derivative of catechol with two sulfate functional groups on the benzene ring, which greatly improve its solubility in aqueous solution.29 The solubility at 25 ℃ and redox performance of tiron in aqueous solution, as well as the CO2 absorption capacity of the Na2Q solution (Na2Q was prepared by QH2 and NaOH in a molar ratio of 1:2), were measured. The feasibility and efficiency of CO2 desorption when the CO2-saturated Na2Q was oxidized and the ability of Na2Q to be regenerated by cathodic reduction were also studied.

THEORETICAL BASIS

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Figure 2. Schematic diagram of electrochemical regeneration Because the reduced and oxidized states of tiron have different acidities in aqueous solution, the pH value changes during the redox reaction process. The Na2Q solution is alkaline and used to absorb CO2. The CO2-saturated Na2Q solution is oxidized at the anode, and the Q2-, QH-, and QH2 in the CO2-saturated Na2Q solution are oxidized to Q. This process releases H+ and promotes the desorption of CO2 from the anolyte. Figure 2 shows the electrochemical reaction process, which includes the reaction steps described below.

CO2 absorption in Na2Q solution. Na2Q is a salt that hydrolyzes to release OH- in aqueous solution, as shown in eq (1) and eq (2). Then OH- reacts with CO2 to form HCO3-(eq (3)). The CO2-saturated solution may contain the ions Na+, Q2-, QH-, QH2, and HCO3-.

𝑄2 ― + 𝐻2𝑂 = 𝑄𝐻 ― + 𝑂𝐻 ―

(1)

𝑄𝐻 ― + 𝐻2𝑂 = 𝑄𝐻2 + 𝑂𝐻 ―

(2)

𝑂𝐻 ― + 𝐶𝑂2 = 𝐻𝐶𝑂3―

(3)

Anodic reactions for the desorption of CO2 by oxidizing the CO2-saturated Na2Q solution. Using the CO2-saturated solution as the anolyte, Q2-, QH-, and QH2 are oxidized to Q, as shown in eq (4), eq (5), and eq (6). This process releases H+, which reacts with HCO3- to release CO2, and the pH of solution drops as a result of the oxidation. After the separation of CO2, Q solution is pumped to cathode chamber to act as the catholyte. 7 ACS Paragon Plus Environment

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𝑄2 ― ―2𝑒 ― = 𝑄

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(4)

𝑄𝐻 ― ―2𝑒 ― = 𝑄 + 𝐻 +

(5)

𝑄𝐻2 ―2𝑒 ― = 𝑄 + 2𝐻 +

(6)

𝐻 + + 𝐻𝐶𝑂3― = 𝐻2𝑂 + 𝐶𝑂2

(7)

Cathodic reactions to regenerate Na2Q. At the cathode, Q is reduced to Q2-, and Na+ permeates through the cation exchange membrane (CEM) to balance the electrical charges.

𝑄 + 2𝑒 ― = 𝑄2 ―

(8)

𝑄 + 2𝐻2𝑂 + 2𝑒 ― = 𝑄𝐻2 +2𝑂𝐻 ―

(9)

The mechanism of the proton coupled electron transfer (PCET) reaction of Q changes in different solutions.30 The reduction of Q to hydroquinone (QH2) involves reaction with 2H+ and 2e- in buffered solution, but it is quite different in non-buffered solutions. Quan et al.

31, 32

suggested that Q accepts 2e- to form a strongly hydrogen-bonded

quinone dianion (Q2-) in non-buffered aqueous solutions, as shown in eq (8). Because of the hydrolysis effect, Q2-, QH-, and QH2 co-existed in the aqueous solution. Wang et al.

33-35

suggested that water is the proton donor in Q reduction in non-buffered

solutions, as shown in eq (9). Although the Q reduction process is complicated in nonbuffered solutions, it results in an increase of the solution pH value, and therefore the reduced solution can be used to capture CO2. In our work, the Q2- formed by the reduction of Q at the cathode was balanced with Na+, which permeated through the CEM from the anode chamber, and thus, the Na2Q solution was regenerated.

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Figure 3. Electrode potential-pH relationship. It is assumed that the ion activity of the solution is 1, and the gas pressure is 100 kPa. 𝑅𝑇

𝑎𝑄𝐻2

φ𝑄 = 𝜑𝜃𝑄,𝑄𝐻2 ― 2𝐹𝑙𝑛𝑎

2 𝑄𝑎𝐻 +

φ𝑄 = 0.72 ― 0.05916𝑝𝐻

（10） （11）

The standard electrode potential of tiron is 0.72 V, and its Nernst equation is shown in eq (10).29 If the ion activity in solution is 1 and the gas pressure is 100 kPa, the theoretical potential and pH can be related by the expression in eq (11). Hydrogen evolution and oxygen evolution are possible side reactions in the aqueous solution electrolysis process. According to the Nernst equation, the electrode potential is a function of the concentration and acidity. For a given concentration, the electrode potential depends on the pH value. The potential-pH relationship for the reduction of Q and for the side reactions are compared in Figure 3. The potential of the redox reaction of tiron is always between that of the oxygen evolution reaction and the hydrogen evolution reaction. Thus, oxygen and hydrogen evolution can be avoided as side reactions. 9 ACS Paragon Plus Environment

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EXPERIMENTAL SECTION Materials. All chemicals were purchased and were of analytical grade or higher. Tiron (disodium 4,5-dihydroxy-1,3-benzenedisulfonate) of AR grade was purchased from Shanghai Macklin Biochemical Co., Ltd. China.

Membrane and electrode pretreatments. The Nafion 117 (Du Pont) membrane used in this work was pretreated with a 5% hydrogen peroxide solution at 80 ℃ for 1 hour, and then immersed in deionized water at 80 ℃ for 1 hour. The pre-treated membrane was kept in a 0.1 M NaOH solution at room temperature. A glassy carbon electrode with a diameter of 3 mm (Tianjin Aida Co., Ltd. China) was polished with a 0.3 µm alumina suspension on a polishing pad, rinsed with water and ethanol, polished again with a 0.05 µm alumina suspension, and finally rinsed with water and ethanol. A graphite felt electrode (China Shenhe Carbon Fiber Material Co., Ltd.) with a thickness of 5 mm was first washed with deionized water and then calcined in air at 400 ℃ for 6 hours.

Cyclic voltammetry test. Cyclic voltammetry (CV) curves were measured at room temperature using a CHI 660E potentiostat (CHI Corporation, USA) with a threeelectrode cell. The working electrode was a glassy carbon electrode, the counter electrode was a platinum mesh electrode, and the reference electrode was a Ag/AgCl (saturated KCl solution) electrode. A mixture of a 10 mM Na2Q solution and a 1 M aqueous Na2SO4 was used as the electrolyte. The Na2Q solution was composed of tiron (QH2) and sodium hydroxide in a 1:2 molar ratio. The electrolyte was adjusted to 10 ACS Paragon Plus Environment

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different pH values by absorbing CO2, which was introduced by bubbling a gas stream containing 12% CO2 through the solution for different lengths of time.

Solubility test. The solubility of the tiron solutions was measured using an ultraviolet-visible spectrometer (UV-1500PC, Macy China Instruments Inc.). Absorbance was measured at 290 nm and compared to an absorbance-concentration calibration curve.

CO2 capacity of the Na2Q solution. Measurement of the CO2 capacity of the Na2Q solution was conducted in a round bottom flask that was immersed in a water bath. A Na2Q solution of a given concentration was stirred in the flask and a gas stream containing 12% CO2 was bubbled through at a flow rate of 200 mL/min. The CO2 concentration of the outlet gas was measured using an OmniStar Gas Mass Spectrometer (Pfeiffer, Germany). When the outlet CO2 concentration returned to 12%, the CO2 absorption had stopped, and the amount of CO2 absorbed was calculated.36

Structure of the electrolytic cell. The electrolytic cell was composed of a cathode chamber and an anode chamber, separated by a CEM (Nafion 117) with a membrane area of about 2.14 cm2, as shown in Figure S2. The chambers were approximately 40 mL, and the electrode distance was around 6 cm. Both chambers were equipped with agitation rotors.

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RESULTS AND DISCUSSION

Figure 4. (a) CV curves of 10 mM Na2Q in a CO2-saturated 1 M Na2SO4 solution (pH= 5.8) at scan rates of 20 mV/s, 50 mV/s, 100 mV/s, and 200 mV/s, respectively. (b) Relationship between peak current density (Jp) and the square root of the scanning speed (ν1/2). Redox performance of Na2Q. As shown in Figure 4(a), Na2Q exhibited quasireversibility with an obvious peak potential separation; the peak potential difference increased with the scanning rate. The peak current density (Jp) was proportional to the square root of the scanning rate (ν1/2) (see Figure 4(b)); this linearity indicates that the redox process of tiron was diffusion-controlled. Na2Q also was relatively stable in the presence of trace O2 (See Figure S3).

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Figure 5. (a) CV curves of 10 mM Na2Q in a 1M Na2SO4 solution at various pH values. (b) CV curves of 1M Na2SO4 solution at various pH values. The scanning speed was 100 mV/s. The pH value of the electrolyte greatly influenced the redox performance of Na2Q. As shown in Figure 5 (a), the reductive current at 0.24 V gradually decreased with the pH value, and at different pH values the oxidative current at 0.7 V also changed, and new oxidative peaks appeared. At pH = 11.2, multiple oxidative peaks were observed. The drop in the reductive current at 0.24 V indicated that the reduction of Q was difficult in strongly alkaline solution. Comparing Fig. 5(a) and (b), it can be found that the redox potential of Na2Q at different pH values is always within the potential window of water spitting.

CO2 capacity of the Na2Q solution. Unlike many hydroquinones, tiron has very good solubility in aqueous solution; its solubility at 25 °C in water can reach 1 M. The measured CO2 capacity of Na2Q solutions at different concentrations are shown in Figure 6.

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Figure 6. CO2 capacity of Na2Q solutions at different concentrations. The CO2 capacity increased linearly with the Na2Q concentration, and the CO2/Q2molar ratio was close to 1 (see Figure 6). With increasing concentration, the CO2/Q2molar ratio decreased slightly. When the concentration of Na2Q in the solution reached 0.8 M, the CO2-saturated solution became over-saturated and a precipitate formed. This was very likely due to the concentration of QH- exceeding its maximum solubility.

In the subsequent measurements, 0.7 M Na2Q solution was used as the absorption solution, as this concentration showed the maximum CO2 capacity. The calculated Kb1 and Kb2 of Na2Q are 𝐾𝑏1 = 5.66 × 10 ―3,𝐾𝑏2 = 1.81 × 10 ―12(see the calculation process in supporting information). Na2Q mainly undergoes primary hydrolysis reaction as shown

in eq (1) because Kb2 is very small compared with Kb1, so the ions in the solution mainly included Na+, QH-, HCO3-, and Q2-.

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Figure 7. Attenuated total reflectance infrared spectra of the different solutions The attenuated total reflection infrared spectroscopy (IR-ATR) results are shown in Figure 7. After CO2 absorption, the IR spectra of the CO2-saturated Na2Q solutions showed a new peak at 1350 cm-1, which corresponded to NaHCO3. This indicated that CO2 had dissolved in the solution and reacted with OH- to form HCO3- according to Eq (3). From the intensities of the HCO3- peaks, the concentration of HCO3- was observed to increase with the concentration of Na2Q.

Anode half-cell. In order to investigate the desorption of CO2 when the CO2saturated Na2Q was oxidized, an anode half-cell experiment was conducted using 35 ml of a CO2-saturated 0.7 M Na2Q solution as the anolyte and graphite felt as the anode. The cathode chamber was composed of a platinum plate electrode and 1 M Na2SO4 solution as the catholyte, so that the cathode reaction would reduce protons to hydrogen. As the hydrogen evolution reaction occurred, the pH value of the catholyte rose and 15 ACS Paragon Plus Environment

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Na+ ions permeated through the CEM from the anolyte to the catholyte, resulting in the formation of NaOH in the catholyte. The anode potential versus the Ag/AgCl electrode was measured using a potentiostat under a constant current of 0.25 A (current density: 18mA/cm-2). As an electrolysis reaction nears the terminal point, the anode potential may increase sharply, so the experiment was stopped when a steep potential rise was observed. In order to detect the gaseous products on the anode, 50 mL/min of N2 was used as a purging gas in the anode chamber, and the outlet gas was analyzed using a mass spectrometer.

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Figure 8. (a) Anode potential (vs. Ag/AgCl) at a constant current of 0.25 A. (b) Composition of the anode gas products. (c) pH values of the anolyte and the catholyte. The CO2-saturated Na2Q solution mainly contained the ions Na+, QH-, HCO3-, and Q2-, so the anode oxidation included reactions (4), (5), and (7). The sharp rise of the anode potential (vs. Ag/AgCl) from 230 minutes to 260 minutes (see Figure 8(a)) indicated that the terminal point of the electrolysis reaction was approaching. The electrolysis experiment was conducted for approximately 260 minutes, and the final reaction Coulomb number was about 4000 C.

Analysis of the gas from the anode chamber showed that a significant amount of CO2 was released during electrolysis (see Figure 8(b)), while no O2 was produced. The CO2 desorption involved an induction period, which indicated that the protons accumulated around the electrode because of the mass transfer limitation (see Figure S4). And it takes time for CO2 to mix well and for N2 to purge CO2 to the mass spectrometry. This would be another reason for the slow change in the concentration of CO2 at the initial stage of electrolysis. The acidity of the anolyte increased as the anodic oxidation 17 ACS Paragon Plus Environment

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proceeded, which led to the decomposition of carbonate and a sharp increase in the concentration of CO2. As the HCO3- around the electrode was gradually depleted, the CO2 concentration began to decline gradually due to the limited mass transfer of HCO3-. The pH value also dropped gradually with the decomposition of carbonate and the desorption of CO2 (see Figure 8(c)). Without the presence of Na2Q, no current was generated at the anode potential (vs. Ag/AgCl) of 0.3 V, 0.5 V, respectively and the pH of the anolyte did not change (see Figure S5 (a)). Hydrogen was produced on the cathode, and the electroreduction of protons on the cathode also produced OH-, which led to the rise in the catholyte pH value.

Figure 9. CO2 desorption ratio and Faraday efficiency during electrolysis at a constant current of 0.25 A 𝑡ℎ𝑒 𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝐶𝑂2 𝑑𝑒𝑠𝑜𝑟𝑏𝑒𝑑

desorption ratio = 𝑡ℎ𝑒 𝑎𝑚𝑜𝑢𝑛𝑡 𝑜𝑓 𝐶𝑂2 𝑖𝑛 𝑡ℎ𝑒 𝑠𝑎𝑡𝑢𝑟𝑎𝑡𝑒𝑑 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛 × 100% (12) 𝑄𝐻 ― + 𝐻𝐶𝑂3― ―2𝑒 ― →Q + 𝐻2𝑂 + 𝐶𝑂2

(13)

The CO2 desorption ratio during electrolysis was estimated using Eq (12), in which ‘the amount of CO2 desorbed’ was calculated as the difference between the CO2 content of the initial solution and that of the electrolyzed solution. The Faraday efficiency was 18 ACS Paragon Plus Environment

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calculated by assuming that the ratio the e-/CO2 was equal to 2 in the electrochemical reaction, and the oxidation of QH- involves 2-electron transfer, as shown in Eq (13). Figure 9 shows the CO2 desorption ratio and the Faraday efficiencies at different times. The desorption ratio increased almost linearly with time. When the Coulomb number reached about 4000 C in 260 minutes, the CO2 desorption ratio reached 100%. The Faraday efficiencies fluctuated between 1 and 1.1; these efficiencies may have been influenced by the physical adsorption of CO2 in the supersaturated state. Assuming that the true Faraday efficiency is 90%, the amount of CO2 in the supersaturated state released after 4 hours of electrolysis accounts for about 10% of the total CO2. However, 18% CO2 in 0.7 M CO2-saturated solution can be released only by purging with N2 (Figure S6). This means that the supersaturated CO2 was released into the N2 purge gas during the desorption process, resulting in Faraday efficiencies of greater than 1.

Cathode half-cell. After the desorption of CO2, the oxidized anolyte was delivered to the cathode chamber to regenerate the Na2Q solution. The cathode half-cell measurements were conducted in the electrolytic cell. 35 ml of 0.7 M Q solution, which was obtained by oxidizing of CO2-saturated Na2Q solution, was used as the catholyte, and graphite felt was used as the electrode. Na+ permeated through the CEM from anode chamber, balancing the charge of the Q2- to form Na2Q. The anolyte was 1 M NaOH + 1 M Na2SO4 solution, and the electrode was a platinum plate. The oxygen evolution reaction occurred on the anode during the electrolysis process. The experiment was conducted at a constant current of 0.25 A (current density: 18 mA/cm-2). The cathode potential (vs. Ag/AgCl) was recorded using a potentiostat. Similarly, the cathode 19 ACS Paragon Plus Environment

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potential sharply increased in the final period as the terminal point of the electrolysis approached. The gaseous products of the cathode were purged with a gas stream of N2 at a flow rate of 50 mL/min, and analyzed using a mass spectrometer.

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Figure 10. (a) Change in the cathode potential (vs. Ag/AgCl) at a constant current of 0.25 A. (b) Gas composition of the cathode chamber at a constant current of 0.25 A. (c) pH values of the anolyte and the catholyte Figure 10 (a) shows that the cathode potential (vs. Ag / AgCl) declined gradually within 100 minutes and decreased dramatically at around 100 minutes. These results demonstrated that the reduction time was much shorter than the oxidation time, which in turn indicated that Q could not be completely reduced to Q2-. From the CV curves, the reduction of Q in strongly alkaline solution was shown to be difficult. During the reduction process, the pH value of the catholyte rose with the reduction of Q, preventing more complete reduction of catholyte, and thus, the final pH value of the catholyte reached only 9.2 (See Figure 10 (c)). Cannan et al. [34] found that a pH swing near the electrode surface occurred during a quinone reduction using confocal laser scanning microscopy. Therefore, the local pH around the electrode may be much greater than 9.2, making the reduction current decrease and then the cathode potential drop sharply. H+ permeated the Nafion membrane in the form of H3O+ from anode side to cathode, 21 ACS Paragon Plus Environment

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which diluted the catholyte and would cause a decrease in the catholyte pH (Figure S7). No hydrogen evolution was observed during this process (see Figure 10 (b)).

Figure 10 (c) depicts the pH change of the anolyte and catholyte. In the cathodic reduction process, the pH of the cathodic electrolyte increased from 5.7 to 9.2, which indicated that Q became Q2- and that Q2- increased the pH of the cathodic electrolyte. Without the presence of Q, no current was generated at the cathode potential (vs. Ag/AgCl) of -0.2 V, -0.4 V, respectively and the pH of the anolyte did not change (see Figure S5 (b)). During the experiment, a large amount of bubbles appeared in the anode chamber, and the pH value of the anolyte decreased slightly, indicating that oxygen was produced by the oxidation of OH-. The catholyte, which had a pH of 9.2, was used to absorb 12% CO2, and the CO2/Q2- ratio of the CO2-saturated solution was 0.3. Although this was lower than the 0.9 ratio of CO2/Q2- in the initial CO2-saturated 0.7 M Na2Q solution, its CO2 absorption capacity was still relatively good.

Intermittent cell electrolysis. Based on the results obtained from the two half-cells, the electrolysis was conducted in an intermittent cell, in which 35 mL of CO2-saturated Na2Q solution was used as the anolyte and 35 mL of Q solution obtained by oxidizing saturated Na2Q solution was used as the catholyte. Both electrodes were graphite felt, and a constant current of 0.25 A (current density: 18 mA/cm-2) was used. Because the composition of the gas in the cathode chamber and the anode chamber could not be measured at the same time, the gas composition of the cathode chamber was measured during the initial 30 minutes, and then the gas in the anode chamber was measured. 22 ACS Paragon Plus Environment

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During the electrolysis process, 50 mL/min of N2 was used as a purge gas to carry the gas to the mass spectrometer for detection of the gas composition and concentration.

Figure 11. Gas composition of the cathode chamber and anode chamber during the 0.25 A constant-current electrolysis process Figure 11 shows the gas composition during the electrolysis, which demonstrated the possibility for the electrochemical desorption of CO2. The potential difference between the cathode and anode ranged between 0.35 V and 0.75 V during the 75 minutes of electrolysis and the potentials of both the anode and the cathode were always within the potential window of water spitting (Figure S8). The pH swing of the anode and cathode in the batch cell is consistent with that of the anode half-cell and the cathode half-cell (Figure S9). Based on the potential difference between the cathode and the anode, the energy consumption of CO2 desorption was estimated to be about 2.4 GJ/t CO2 (The calculation process was shown in supporting information). the energy consumption of the most amine-based technologies is between 3.11 GJ/t CO2 and 4.0 GJ/t CO2 [15-17], 23 ACS Paragon Plus Environment

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so the energy consumption of 2,4 GJ/t CO2 is quite competitive. The energy consumption could be further reduced by optimizations of the electrolysis cell, such as reduction of the inner resistance and improvement of the electrodes.

Figure 12. (a) The linear sweep voltammetry (LSV) curves of the anode reaction and the cathode reaction in each half-cell at 12 ℃. (b) The Tafel curves of the anode reaction and the cathode reaction. The anolyte is a 0.7 M CO2-saturated solution and the catholyte is a 0.7 M Q solution. Both electrodes were graphite felt (1cm × 1.5 cm × 5mm thinkness). According to the Tafel curves shown in Figure 12, the transfer coefficient (α) and exchange current density (𝑖0) can be calculated. The results are shown in Table 1. Table 1. The kinetic parameters in each half-cell The half-cell

α

𝒊𝟎 (mA/cm-2)

Anode

0.90

1.72 x 10-3

Cathode

0.18

0.26

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Cycle tests. The catholyte, which had a pH of 9.2 after the first cycle, was used to

absorb 12% CO2 for the cycle tests. This CO2-saturated catholyte was used as the anolyte in the second cycle in the anode half-cell. The cathode chamber was composed of a platinum plate electrode and 1 M Na2SO4 solution as the catholyte. After the oxidation reaction was over, the anolyte in the anode half-cell was used as the catholyte in the cathode half-cell. And graphite felt was used as the electrode. The anolyte was 1 M Na2SO4 solution, and the electrode was a platinum plate.

Figure 13. (a) Anode potential (vs. Ag/AgCl) at a constant current of 0.25A (current density: 18mA/cm-2) in the second redox cycle of Q/Q2- in the anode half-cell. (b) Cathode potential (vs. Ag/AgCl) at a constant current of 0.25A (current density: 18mA/cm-2) in the second redox cycle of Q/Q2- in the cathode half-cell. Figure13 shows the anode potential and cathode potential change with time at a constant current of 0.25A in the second redox cycle of Q/Q2-. According to the anode and cathode potential swing, tiron still has good redox performance in the second redox cycle. The absorbed CO2 also could be 100% released during CO2 desorption and the final pH is 6.1. And then this anolyte was reduced at the cathode. However, the pH of the catholyte is only 7.0 when the electrolysis was stopped at the cathode potential of -1 (vs. Ag/AgCl), which could not absorb CO2 again because of the low pH. It can be found that the catholyte formed in the second redox cycle still has good redox performance via investigating its redox behavior in the third cycle, and the final pH of the catholyte is 6.7. Although the Q reduction reactions at the cathode may continue to occur by increasing the cathode potential and the pH of the solution may also rise, the 25 ACS Paragon Plus Environment

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side reactions of hydrogen evolution would also occur. The surface area of the electrode has a great influence on the reduction of Q. The Q reduction performance improved with increasing electrode surface area (Figure S10). Therefore, the next step is to improve the reduction performance of Q by improving the cathode electrode to increase the solution pH. CONCLUSION In this paper, the aqueous tiron was used as the solvent to capture CO2. The feasibility of absorbing and desorbing CO2 and of partial Na2Q regeneration during the CO2 capture process was confirmed.

Tiron has good aqueous solubility, reaching 1 M at 25 °C. The pH value greatly influences the redox performance of the tiron solution. Under neutral and weakly alkaline conditions, tiron exhibits quasi-reversible properties. However, under strongly alkaline conditions, the oxidation reaction of tiron becomes very complex and the reduction reaction is difficult to carry out. The CO2/Q2- ratio of CO2-saturated Na2Q solution (0.1-0.7 M) varies at around 1.

CO2 can be completely desorbed by the electrolysis oxidation of the CO2-saturated Na2Q solution. However, it is difficult to completely regenerate Na2Q via Q regeneration under strongly alkaline conditions. The results show that CO2 can be desorbed without side reactions in a batch electrolysis cell. The energy consumption of CO2 desorption was about 2.4 GJ/t CO2. Tiron has good redox performance under multiple cycles, but the alkalinity could not be recovered.

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AUTHOR INFORMATION Corresponding Author *E-mail: [email protected] Author Contributions The manuscript was written through contributions of all authors. All authors have given approval to the final version of the manuscript. Notes The authors declare no competing financial interest. ACKNOWLEDGMENT We are thankful for the financial support provided by the National Natural Science Foundation of China Project (No. 21878190). REFERENCES (1. Jos G.J. Olivier (PBL), G. J.-M. E.-J.; Marilena Muntean (EC-JRC), J. A. H. W. P. P. Trends in Global CO2 emissions:2016 report; PBL Nehterlands Environmental Assessment Agency: 2016; p 5. 2. Ye, L.; Yue, H.; Wang, Y.; Sheng, H.; Yuan, B.; Lv, L.; Li, C.; Liang, B.; Zhu, J.; Xie, H., CO2 Mineralization of Activated K-Feldspar + CaCl2 Slag To Fix Carbon and Produce Soluble Potash Salt. Industrial & Engineering Chemistry Research 2014, 53, (26), 10557–10565. 3. Gan, Z.; Cui, Z.; Yue, H.; Tang, S.; Liu, C.; Li, C.; Liang, B.; Xie, H., An efficient methodology for utilization of K-feldspar and phosphogypsum with reduced energy consumption and CO 2 emissions. Chinese Journal of Chemical Engineering 2016, 24, (11), 1541-1551. 4. Shangguan, W.; Song, J.; Yue, H.; Tang, S.; Liu, C.; Li, C.; Liang, B.; Xie, H., An efficient milling-assisted technology for K-feldspar processing, industrial waste treatment and CO 2 mineralization. Chemical Engineering Journal 2016, 292, 255-263. 27 ACS Paragon Plus Environment

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