Aqueous Electrochemical and pH Studies of Redox-Active Guanidino Functionalized Aromatics for CO 2 Capture

: Escalating levels of carbon dioxide (CO 2 ) in the atmosphere have motivated interest in CO 2 capture and concentration from dilute streams.


Introduction
Each year, about 30 gigatons of carbon is emitted from human activity, with roughly 30% of these emissions arising from fossil fuel combustion. [1]Flue gas from fossil fuel power plants typically contains 8-10% CO2.The current state-of-the-art method for CO2 capture from flue gas uses an aqueous alkanolamine solution to capture CO2 at ambient temperature and pressure.The release of CO2 is achieved by heating the sorbent (383 K for monoethanolamine). [2,3]However, these thermal-swing methods only can achieve 5-40% of the maximum theoretical energy efficiency. [1,4]tps://doi.org/10.26434/chemrxiv-2023-fhcv4ORCID: https://orcid.org/0000-0002-9680-8260Content not peer-reviewed by ChemRxiv.License: CC BY-NC-ND 4.0 In contrast to thermal methods of CO2 capture, without Carnot limitations, electrochemical methods can achieve a theoretical energy efficiency of 100%. [3,5]One common motif for ambient electrochemical CO2 capture and concentration uses redox active molecules that either 1) directly bind to and release CO2 or 2) modify the pH of aqueous solutions (Scheme 1A). [6]In the first approach, redox carriers bind CO2 in their electron-rich reduced form (Scheme 1A, red pathway).Subsequent electrochemical oxidation decreases its electron density, lowering its CO2 binding affinity and releasing concentrated CO2 The redox carrier can then be regenerated for CO2 capture via electrochemical reduction. [7]Another approach uses redox chemistry to change the pH of an aqueous solution for CO2 capture and release (Scheme 1A, blue pathway).These pH swing cycles capitalize on the pH-dependent CO2 equilibria in aqueous solutions.At high pH values, hydroxide reacts with CO2 to form bicarbonate and carbonate, enabling capture from dilute streams.At low pH values, the equilibrium is reversed, and CO2 is released.Redox-active molecules whose proton affinity (pKa) significantly varies depending on their oxidation state are used to induce these electrochemically driven pH swings. [8]anidines are a family of neutral organic Brønsted superbases (pKa values of 12 or greater in water) with attractive CO2 capture properties. [9]They bind CO2 through two different mechanisms depending on the solvent.These mechanisms parallel the two capture routes shown in Scheme 1A. [7]In aprotic solvents, tetramethyl guanidine (TMG) binds directly to CO2 to form a zwitterionic carbamate salt (Scheme 1B, pathway A). [10] In protic solvents such as alcohols or water, sufficiently Brønsted basic guanidine deprotonate the alcohol or water, which then binds CO2 to form an alkyl carbonate or carbonate ion pair with the protonated TMG (Scheme 1B, pathway B). [11] Because of their reactivity with CO2, guanidine derivatives have been studied for their use in thermal capture and release cycles.Heldebrandt and coworkers developed a new class of CO2-binding organic liquids composed of alcohol solvents and TMG derivatives. [11]he lack of reversible redox properties in the previously studied guanidine compounds limited their use as sorbents to thermal regeneration cycles.However, Himmel and coworkers recently synthesized TMG-substituted functionalized aromatics, or guanidino functionalized aromatics (GFAs), with reversible redox properties. [15]These redox-active superbases have not been assessed for electrochemical CO2 capture and release.A variety of GFAs can be obtained through straightforward synthetic routes.Among these, 1,4-bis(tetramethylguanidino)benzene (1,4-btmgb) was selected for testing due to its mild redox potential and stability in air.In this study, we explored the use of 1,4-btmgb in organic and aqueous solutions to assess its viability as a redox carrier and pH swing mediator for electrochemical CO2 capture.
The formation of carbonate salt was also present in pure water, supported by 1 H NMR analysis.The bicarbonate (HCO3 -), respectively. [17,18]a and Electrochemical Properties with Water and Protic Solvents Since even trace water in aprotic solvents results in the formation of the protonated base and bicarbonate, we postulate that 1,4-btmgb is a better candidate for an aqueous electrochemical pH swing cycle than direct capture.To assess the suitability of 1,4-btmgb for an aqueous electrochemical pH swing, we first sought to determine the pKa values for [1,4-btmgbH2] 2+ in water.Compared to the CVs obtained on 1,4-btmgb at various stages of protonation in MeCN (Figure 1), the CVs of 1,4-btmgb in water varied with pH (Figure 3, S14-18), reflecting different protonation states.[21] Cyclic voltammograms of 1,4-btmgb were acquired at pH values between 2.5 and 13.   process. [22]Above pH 13.5, there were minimal changes to the redox potential.With differential-pulse voltammetry (DPV) experiments from pH 9.5 to 14.1, the exact breaking points in the pH-dependent and pH-independent regions allowed for the determination of pKa values (Figure S22 The redox features of 1,4-btmgb are less reversible at both higher and lower pH values.Between pH 11.3-11.9there were two additional reduction peaks negative of the reversible redox event at E1/2 = 0.2 V vs SCE (pH = 11.75)(Figure 3, blue trace).The reduction event at 0.0 V is attributed to the reduction of [1,4-btmgbH] 3+ , which is expected to form at these pH values.The more cathodic reduction event at -0.18 V vs. SCE increases in current at high pH (11.3 and higher, Figure 3, green trace) but was not observed when 1,4-btmgb was probed electrochemically in other non-aqueous protic solvents (methanol and ethanol, Figures S24-25).We attribute this reduction to an interaction with water upon oxidation, which is supported by electrolytic studies (vide infra).
An aqueous 2 mM solution of protonated [1,4-btmgbH2][Cl2] was also oxidized.The initial pH was 9.88 and then decreased to 8.72.Upon subsequent reduction, the pH of the solution increased to 9.38.(Table 1).While both solutions produced a change in pH from electrolysis, the magnitudes were smaller than expected, and the solutions could not be restored to the original pH after reduction.The solutions also  The redox event at -0.18 V vs SCE we attribute to an aniline species that is the product of the where oxidation of this compound results in the Wurster cation and undergoes further decomposition in the solution mixture due to the N-H moiety. [24]7] 1 H NMR experiments were also performed on chemically oxidized [1,4-btmgb 2+ ] in deuterated acetonitrile in the presence of water, showing the decomposition of the parent species 30 minutes after the addition of water.Over the course of five hours, additional resonances at around 2.89 ppm and in the aromatic region (6.74-6.76ppm) increase in intensity (Figures S26-27).The resonance at 2.89 ppm is assigned to the methyl proton peaks of tetramethyl urea, and the resonances between 6.74-6.76ppm are attributed to the aniline species formed post hydrolysis (tetramethylguanidinodimethyl-p-phenylenediamine has aromatic resonances at 6.49-6.69ppm while p-phenylenediamine has a resonance at 6.46 ppm in CD3CN). [28]Altogether, the new resonance peaks in the aromatic region indicate decomposition in the presence of water.Monitoring the reduced species 1,4-btmgb in D2O at high and low pH via 1 H NMR spectroscopy does not show evidence of decomposition over the span of days, indicating it is the oxidized species that is sensitive to deleterious reactions with water.4). [29]Water is also known to stabilize this radical formation. [30] verify the presence of a radical in the solution, electron paramagnetic resonance (EPR) spectra were acquired using aliquots of the oxidized 2 mM 1,4-btmgb.An EPR signal at g⊥ = 2.0066 was present for the solution after controlled potential electrolysis at 0.5 V for 10 minutes, indicating the formation of an S = ½ organic radical species (Figure S28).
Ultraviolet-visible (UV-Vis) spectroelectrochemistry was used to obtain absorption spectra of the generated oxidized species [1,4-btmgb] 2+ (Figure 6) through controlled potential electrolysis of a 0. The post-electrolysis solution obtained after oxidation in an H-cell with stirring leads to different absorption spectra than what was observed in the spectroelectrochemical experiments (Figure S29).The absorbances that correspond to the oxidized species were not present.We believe convection increases the rate of comproportionation.After one day, the absorbance at 272 nm also decreases in intensity, indicating decomposition.
Collectively, these decomposition pathways result in smaller than expected pH changes in the controlled potential electrolysis experiments (Table 1).Additionally, the side reactions hinder the regeneration of 1,4-btmgb and the overall reversibility of an electrochemical cycle.

Conclusion
A guanidine functionalized aromatic (GFA) compound, 1,4-btmgb, was examined as a redox-active candidate for CO2 capture.Reactivity with CO2 in the presence of trace water forms protonated guanidine and carbonate.In addition, the stability of 1,4-btmgb and the preliminary redox behavior in water indicate it could potentially mediate a pH swing cycle.The aqueous pKa values of [1,4-btmgbH2] 2+ were measured to be 11.0 and 13.5.However, controlled potential electrolysis indicates more complicated reactivity in water, with hydrolysis and radical formation decomposition pathways.The use of GFAs in electrochemical CO2 capture and concentration will require addressing these decomposition pathways.

Experimental Section
General Methods: Synthetic work was carried out in either ambient air or out under a dinitrogen (N2) atmosphere in a glovebox where noted.All solvents and reagents were purchased from commercial vendors and used without further purification unless otherwise noted.Deuterated acetonitrile used for NMR characterization were purchased from Cambridge Isotope Laboratories, Inc., and were degassed via free-pump-thaw and stored over activated 3 Å molecular sieves.Tetrabutylammonium hexafluorophosphate (TBAPF6) was purified via recrystallization from ethanol and dried under a heated (80 °C) vacuum (1 x 10 -3 Torr) and stored in a glovebox.Experiments using carbon dioxide (CO2) atmospheres were performed using ultra-high purity (99.999%)CO2 which was additionally passed through a purification column to eliminate residual H2O, O2, CO, halocarbons, and sulfur compounds Physical Methods: Electrochemical experiments were carried out under atmospheres of N2 or CO2, where indicated.Solutions for electrochemical measurements were recorded in acetonitrile or water using 100 mM TBAPF6 (or TMAPF6) or 100 mM KCl as supporting electrolytes, respectively.Cyclic voltammetry was performed with a Pine instrument WaveDriver 10 potentiostat.When specified, ferrocene (Fc) was used as an internal standard, and potentials are referenced to the ferrocenium/ferrocene (Fc +/0 ) couple.In water, a saturated calomel electrode (SCE) was used as the internal reference electrode.UV-Visible spectroelectrochemistry (UV-Vis SEC) was performed using a UV-Vis SEC kit from Pine Instruments with a Pt working/counter electrode and Ag wire as a pseudo reference electrode.UV-Vis absorption spectra were recorded using a 1 cm quartz cuvette with an Agilent Cary 60 UV-Vis spectrophotometer.CO2 monitoring was performed on a CM-0111 gas sensor kit CM-0111 by CO2Meter with the software GasLab.A Bruker EMX spectrometer equipped with an ER041XG microwave bridge was used to collect X-band EPR spectra.Fourier Transform Infrared (FTIR) spectra were collected using a Thermo Scientific Nicolet iS5 FTIR Spectrometer with iD5 diamond ATR. ) program package was used to determine the unit-cell parameters and for data collection (2 sec/frame scan time).The raw frame data was processed using SAINT 2 and SADABS 3 to yield the reflection data file.Subsequent calculations were carried out using the SHELXTL 4 program package.There were no systematic absences nor any diffraction symmetry other than the Friedel condition.The structure was solved by direct methods and refined on F 2 by full-matrix least-squares techniques.The analytical scattering factors 5 for neutral atoms were used throughout the analysis.Hydrogen atoms were located from a difference-Fourier map and refined (x,y,z and Uiso).The molecule was located about an inversion center.1,4-btmgb: Synthesized following previously described prep by Himmel and coworkers. [15]

Scheme 1 .
Scheme 1.A (top): Two modes of CO2 capture by a redox-active molecule (R) via direct capture or pH swing methods.B (bottom): Two modes of CO2 capture by guanidines.
8. From these experiments, a partial Pourbaix diagram (Figure 4) was constructed by plotting the pH measured to the observed E1/2 (or Epa and Epc, S19-21).The reduction potential remained relatively constant below pH values of 11.The oxidation potential shifted cathodically as the pH increased from 11 to 13.5 (Figures 4, S19-21).The slope of the shift in E1/2 versus pH was observed to be -40.4mV/pH, indicating a two-electron, one-proton
turned a deep red color during oxidation.The color faded in intensity over time but did not completely disappear.The CVs of the oxidized 1,4-btmgb solution are shown in Figure 5 (red trace).Compared to the trace of the solution before oxidation, a new reductive event appears at -0.18 V vs. SCE.This feature and the color change indicate the formation of new species after controlled potential oxidative electrolysis.
decomposition of 1,4-btmgb via the loss of a TMG group via base-catalyzed hydrolysis (Scheme 3).This reduction feature is only present in CVs of 1,4-btmgb in water among the protic solvents surveyed, indicating it is related to the [1,4-btmgb 2+ ] generated in situ interacting with water/hydroxide.Prior studies have likened the redox behavior of the TMG substituted aniline to that of 1,4-bis(dimethylamino)-benzene