Multi-site CO₂ fixation in triazolates: cooperative O,N binding enhanced by solvation and counter-ion effects

Abstract

The cooperative interaction of nitrogen and oxygen centers in aromatic heterocycles provides an effective pathway for charge-assisted CO₂ capture. Building on the pioneering work of Luo et al. on hydroxy-pyridine systems and our recent PCCP study on hydroxy-substituted N-heterocycles, this work extends the O,N-cooperative binding concept to triazolate frameworks. Density functional theory (DFT) and molecular electrostatic potential (MESP) analyses were performed on neutral, monoanionic, and dianionic 1,2,3- and 1,2,4-triazoles to elucidate how charge, counter-ion and solvation govern CO₂ adsorption. Deprotonation generates anionic and dianionic triazolates with enhanced negative potential at N and O sites, enabling O-carboxylate and N-carboxylate formation. While gas-phase dianions capture multiple CO₂ molecules through cooperative charge delocalization and carbonate-chain growth (up to three CO₂ with ΔG_ad ≈ −54 kcal mol⁻¹), the inclusion of ethanol solvation and tetramethylphosphonium counter-ions reveals that these species remain strongly exergonic. Reaction modeling shows that both mono- and dianionic triazolates undergo spontaneous CO₂ fixation in solution (ΔG ≈ 24 to −61 kcal mol⁻¹), forming ion-paired poly(carboxylate) complexes with up to six CO₂ molecules. These results demonstrate that even under polar, solvated conditions, counter-cation-stabilized triazolates preserve high-affinity, multi-site CO₂ capture, identifying them as realistic and promising building blocks for molecular CO₂-activation strategies.

---

Introduction

The continuous rise in atmospheric carbon dioxide (CO₂) concentration remains one of the major contributors to global climate change.^1–3 Developing efficient and sustainable CO₂ capture and utilization strategies is therefore a key challenge in modern physical chemistry.^4–6 Among the diverse approaches, molecular and material-based sorbents capable of reversible chemical binding of CO₂ have received increasing attention because of their tunability and potential for low-energy regeneration.^7–9

Conventional sorbents such as aqueous amines, zeolites, and metal–organic frameworks (MOFs) each have intrinsic limitations.^10–13 Amine scrubbing processes suffer from thermal degradation, volatility, and high regeneration costs, while solid sorbents such as MOFs and zeolites often display reduced uptake under humid or low-partial-pressure conditions.^14–16 These challenges have prompted the design of new molecular adsorbents that exploit charge-assisted, donor–acceptor, and cooperative electrostatic interactions to achieve strong yet reversible CO₂ binding under mild conditions.^17–19

Hardacre and co-workers demonstrated that a triazole-based superbasic ionic liquid, comprising a phosphonium cation paired with a 1,2,4-triazolate anion, enables efficient electrochemical reduction of CO₂ to formate at remarkably low overpotentials.²⁰ Another major advance in this direction was reported by Luo et al. who demonstrated extraordinarily high CO₂ capacities (up to 1.6 mol CO₂ per mol ionic liquid) in pyridine-containing, anion-functionalized ionic liquids.¹⁴ Their study established that multiple-site cooperative interactions involving simultaneous engagement of oxygen (phenolate) and nitrogen (pyridine) centers with the electrophilic carbon of CO₂ generate synergistic stabilization far exceeding that of single-site binding. Quantum-chemical analysis revealed that π-electron delocalization across the aromatic framework enhances charge transfer to CO₂, while calorimetric data confirmed the thermodynamic feasibility of the process. This discovery provided a conceptual foundation for designing multi-donor anions capable of charge-delocalized, cooperative CO₂ fixation.

Building on this breakthrough, our group recently extended the O,N-cooperative binding concept to hydroxy-substituted aromatic N-heterocycles.²¹ DFT calculations showed that hydroxy-pyridine, dihydroxy-naphthyridine, and trihydroxy-pyridonaphthyridine systems exhibit progressively stronger CO₂ adsorption through synergistic N,O interactions. The adsorption capacity scales with the number of basic and hydroxy sites, enabling one, two, and three CO₂ molecules to associate with the pyridine, naphthyridine, and pyridonaphthyridine frameworks, respectively. Anionic species form covalent-type adducts with CO₂, further stabilized by counter-cations such as Li⁺ or (CH₃)₄P⁺. Solvation slightly weakens the interaction energies but ion-paired systems remain highly exergonic. This work established multi-site, charge-assisted O,N-binding as a general molecular strategy for designing high-affinity CO₂ adsorbents.

Extending this concept, triazoles—particularly 1,2,3-triazole and 1,2,4-triazole offer a versatile molecular platform for exploring charge-delocalized CO₂ binding. Each triazole ring contains three chemically distinct nitrogen atoms—two pyridine-like and one pyrrole-like—that provide multiple Lewis-basic centers for interaction with CO₂.^22–25 The aromatic and π-excessive nature of triazoles imparts notable electronic flexibility, while hydroxy substitution and subsequent deprotonation generate anionic and dianionic triazolates with negative charge delocalized across N and O atoms. Such frameworks are expected to support cooperative O-carboxylate (–OCO₂⁻) and N-carboxylate (–NCO₂⁻) formation analogous to that in hydroxy-pyridine systems, potentially offering even stronger charge-assisted stabilization through extended conjugation and multi-site coordination.

Despite extensive studies on triazoles in catalysis, coordination chemistry, and energetic materials, their potential as molecular CO₂ capture agents have received little attention.^26,27 The influence of deprotonation, charge delocalization, and solvation on CO₂-binding strength has not been systematically investigated, leaving the underlying electronic principles of triazolate-based CO₂ fixation largely unexplored.

In this work, we present a density functional theory (DFT) investigation of neutral, monoanionic, and dianionic derivatives of 1,2,3- and 1,2,4-triazoles to elucidate how molecular charge, electronic structure, and solvation govern CO₂ adsorption. Molecular electrostatic potential (MESP) analysis is employed to locate reactive sites and visualize charge redistribution upon CO₂ coordination. Solvent and counter-cation effects are modeled using the SMD solvation framework (ethanol) and the tetramethylphosphonium cation, analogous to the phosphonium ionic liquids studied by Luo et al. The results reveal a systematic enhancement in CO₂ adsorption energy from neutral to dianionic species, driven by cooperative charge delocalization and orbital overlap. These insights identify triazolate frameworks as promising multi-site, charge-assisted CO₂ capture motifs, extending the O,N-cooperative mechanism established for hydroxy-pyridine systems to a new class of nitrogen-rich heterocycles. The present work is not intended to model a complete catalytic cycle or bulk-phase CO₂ fixation. Instead, it focuses on establishing intrinsic structure–energy relationships governing CO₂ binding to neutral, anionic, and dianionic triazole-based motifs. This molecular-level approach enables isolation of the fundamental electronic factors controlling CO₂ activation, which would be obscured in fully dynamical or bulk-phase simulations.

Methodology

All density functional theory (DFT) calculations were carried out using the M06-2X hybrid meta-GGA functional,²⁸ which provides a reliable description of noncovalent and charge-assisted interactions. The M06-2X functional is particularly effective for accurately modelling intermolecular non-covalent interactions, which are critical for understanding CO₂–adsorbent complexes.²⁸ The cc-pVQZ basis set²⁹ was employed for all geometry optimizations and single-point energy calculations after benchmarking against 6-311++G(d,p), cc-pVDZ, cc-pVTZ, and cc-pV5Z (Table S1). The most extended basis set, cc-pV5Z, produced interaction energies that differ by less than 4% from those obtained using cc-pVQZ. The widely used 6-311++G(d,p) basis set performed reasonably well for many systems, but deviations of up to 18% were observed for some species (Table S1). Therefore, considering both accuracy and computational efficiency, the M06-2X/cc-pVQZ level of theory was adopted for all subsequent calculations. To further validate the robustness of the chosen computational protocol, a benchmark assessment was performed on representative neutral, monoanionic, and dianionic triazole–CO₂ complexes using correlation-consistent basis sets of increasing flexibility, namely aug-cc-pVTZ, and aug-cc-pVQZ, both with and without empirical dispersion corrections (GD3). This analysis was carried out to evaluate the influence of diffuse functions and dispersion on adsorption energetics, which can be particularly relevant for anionic and weakly bound systems. The benchmark results (Table S2) indicate that basis-set augmentation and dispersion inclusion lead only minor structural changes and yield negligible differences in adsorption energies. These findings confirm that the interaction energies are effectively converged at the M06-2X/cc-pVQZ level, justifying its use for the full set of calculations reported in this study. All calculations were performed with Gaussian 16 (revision C.01).³⁰

Neutral, monoanionic, and dianionic forms of 1,2,3- and 1,2,4-triazoles were examined along with their hydroxy, deprotonated (anionic), and double-deprotonated (dianionic) derivatives. In this work, the deprotonated species are formally referred to as triazolates; the terms anionic and dianionic triazoles are used for brevity. Geometry optimizations were followed by vibrational-frequency analyses to confirm that all stationary points correspond to minima (no imaginary frequencies). Thermal corrections at 298 K were included in the reported Gibbs free energies.

To eliminate basis set superposition error (BSSE) arising from basis set incompleteness in the weakly bound complexes, all interaction energies were corrected using the Boys–Bernardi counterpoise method.³¹ The adsorption energy (ΔE_ad) was evaluated as:

ΔE_ad = E_complex − (E_{triazole/triazolate} + E_CO2) + E_BSSE (1)

The corresponding Gibbs free energy of adsorption,

ΔG_ad = G_complex − (G_{triazole/triazolate} + G_CO2) + E_BSSE (2)

Molecular electrostatic potential (MESP)^32,33 topographical analysis was performed on all complexes at the M06-2X/cc-pVQZ level to understand intermolecular interactions and electron density delocalization from the molecule to adsorbed CO₂.^34–36 The MESP minimum or the most negative MESP value is observed at electron-rich sites such as π-regions and lone-pair regions. This minimum (V_min) is a (3, +3) critical point (CP) in the topography of MESP.^37–39

Solvent effects⁴⁰ were modeled with the SMD (solvation model based on density) within the self-consistent reaction field (SCRF) implicit solvation model^41,42 using ethanol as the representative polar medium. To examine the influence of counter-ions, the tetramethylphosphonium cation (CH₃)₄P⁺ was used as a computationally tractable analogue of bulky phosphonium ionic–liquid cations employed experimentally. All solvation calculations were carried out using fully optimized geometries at the M06-2X/SMD/cc-pVQZ level, except for the larger cation–anion paired systems, which were optimized at the M06-2X/SMD/cc-pVTZ level. All structures and energies are provided in the SI.

Results and discussion

Structural and energetic benchmarking

Triazole exists in two structural isomers, 1,2,3-triazole (Tz3) and 1,2,4-triazole (Tz4), each of which can adopt two N–H tautomers depending on the proton position within the ring. This leads to four parent forms, denoted Tz3-1H, Tz3-2H, Tz4-1H, and Tz4-4H. From these, representative mono- and dihydroxy derivatives were generated by substitution at the ring carbon positions, followed by single or double deprotonation to obtain the corresponding monoanionic and dianionic species. The complete set of structures, nomenclature, and substitution patterns considered in this work are summarized in Fig. S1. In the following discussion, emphasis is placed not on individual tautomeric distinctions but on general structure–property relationships, particularly the effects of charge state, substitution, and available donor sites on CO₂ binding behavior. Since the study involves anionic and dianionic species, reliable structural and energetic characterization requires the use of large multiple-ζ basis sets.

Analysis of total electronic energies identified the most stable species within each category (Fig. 1). Among the neutral forms, Tz3-2H (1) and Tz4-1H (2) are the most stable tautomers. For the hydroxy derivatives, Tz3-2H-4OH (1_OH) and Tz4-1H-5OH (2_OH) are the lowest-energy forms, while the most stable dihydroxy derivatives are Tz3-2H-4,5OH (1_OH2) and Tz4-1H-3,5OH (2_OH2). The corresponding monoanions with the lowest energies are Tz3-2H-4O⁻ (1_O−) and Tz4-1H-3O⁻ (2_O−), and the most stable dianions are Tz3-2H-4,5O²⁻ (1_O2−) and Tz4-4H-3,5O²⁻ (2_O2−). These species were therefore selected for detailed investigation of CO₂-binding interactions.

Fig. 1 Neutral, anionic, and dianionic systems considered for CO₂ capture.

MESP analysis of neutral triazole systems and their CO₂ complexes

The MESP topography of representative neutral 1,2,3-triazole systems (Fig. 2) reveals distinct (3, +3) critical points corresponding to the lone-pair regions of nitrogen and oxygen atoms (see Fig. S2 for 1,2,4-triazole systems). These local potential minima (V_min) quantitatively describe the electron-rich zones responsible for electrophilic interactions with CO₂. Hydroxy substitution leads to a systematic decrease in electron density at the nitrogen centers, manifested as less negative V_min values, indicating reduced basicity and nucleophilicity of these sites.

Fig. 2 MESP isosurfaces of 1, 1_OH, and 1_OH2 computed at the M06-2X/cc-pVQZ level of theory. Black dots mark V_min positions. Isosurface value: −12.0 kcal mol⁻¹. Color key: gray – H, blue – N, green – C, red – O.

MESP maps projected on the 0.005 a.u. electron density isosurfaces (Fig. 3) further elucidate the electrostatic complementarity between CO₂ and the triazole ring (see Fig. S3 for 1,2,4-triazole systems). The carbon center of CO₂ exhibits a pronounced electropositive potential (red region), while its oxygens display moderately negative potentials (V_min ≈ −13.2 kcal mol⁻¹). In the 1⋯CO₂ complex, the carbon atom of CO₂ aligns toward the nitrogen lone pair of the triazole, forming a weak N → C interaction that partially transfers electron density to CO₂. This results in a reduction in the negative potential around the interacting nitrogen and a corresponding increase in electron density on the CO₂ oxygen atoms.

Fig. 3 MESP on the 0.005 a.u. isodensity surface for CO₂, 1, 1_OH, 1_OH2, and their CO₂ complexes. V_min values (kcal mol⁻¹) and N⋯C, O–H⋯O distances (Å) at the M06-2X/cc-pVQZ level of theory. Color gradient: blue (electron-rich) to red (electron-deficient).

Similar patterns of potential redistribution are observed for 1_OH⋯CO₂ and 1_OH2⋯CO₂ complexes, where the interacting nitrogen exhibits a pronounced decrease in V_min upon CO₂ coordination (see Fig. S3 for 1,2,4-triazole systems). In addition, a stabilizing hydrogen bond is often formed between the hydroxy group and the oxygen of CO₂, further reducing the negative potential at the involved oxygen site. Overall, CO₂ complexation leads to a more uniform MESP distribution, reflecting enhanced charge delocalization and electrostatic stabilization in the complex.

Structures and interaction energies of triazole–CO₂ complexes

The most stable CO₂ complexes of the neutral triazoles (1 and 2), their hydroxy (1_OH and 2_OH), and dihydroxy (1_OH2 and 2_OH2) derivatives are illustrated in Fig. 4. In all cases, the CO₂ molecule interacts primarily with one of the two nitrogen atoms of the triazole ring through its electrophilic carbon center. The optimized N⋯C distances lie in the range of 2.75–2.85 Å, accompanied by a slight deviation of the O–C–O angle from linearity (177–179°), indicative of a weak donor–acceptor interaction.

Fig. 4 Optimized geometries of the neutral triazole–CO₂ complexes computed at the M06-2X/cc-pVQZ level of theory. Bond lengths are in Å and bond angles in degrees. Color key: gray – H, blue – N, green – C, red – O.

The computed adsorption energies (ΔE_ad) for these complexes range from −4.4 to −6.6 kcal mol⁻¹, while the corresponding Gibbs free energies (ΔG_ad) are positive, between +2.6 and +3.7 kcal mol⁻¹ (Table 1). These values confirm that the interaction of CO₂ with neutral triazole derivatives is thermodynamically unfavorable in the gas phase, consistent with their weakly polar and nonionic nature. Among all studied systems, 2_OH2⋯CO₂ exhibits the largest stabilization energy (ΔE_ad = −6.6 kcal mol⁻¹), reflecting the cooperative contribution of electron-donating hydroxy groups and possible OH⋯O hydrogen bonding with CO₂.

Table 1 BSSE-corrected adsorption energies (ΔE_ad) and free energies (ΔG_ad) of neutral triazole–CO₂ complexes computed at the M06-2X/cc-pVQZ level of theory (in kcal mol⁻¹)

Triazole–CO2 complex	ΔEad	ΔGad
1⋯CO2	−4.4	3.6
2⋯CO2	−4.8	3.7
1OH⋯CO2	−5.2	3.5
2OH⋯CO2	−6.4	2.6
1OH2⋯CO2	−5.2	3.5
2OH2⋯CO2	−6.6	2.7

---

Despite these modest binding energies, the presence of substituents such as hydroxy groups significantly enhances the CO₂ affinity of triazoles by locally increasing the electron density around the nitrogen sites and by enabling additional electrostatic or hydrogen-bonding contacts. Overall, the results for the neutral systems establish a baseline for comparison with their anionic and dianionic counterparts, where deprotonation and charge delocalization are expected to greatly strengthen the CO₂ binding interaction.

MESP analysis of CO₂ binding in anionic triazoles

The MESP maps of the anions 1_O− and 2_O−, along with their corresponding CO₂ adducts, are shown in Fig. 5. Compared with their neutral precursors (1_OH and 2_OH) both anions exhibit substantially more negative electrostatic potentials around the oxygen and nitrogen centers, signifying enhanced nucleophilicity and stronger electrostatic attraction toward the electrophilic carbon of CO₂. The V_min values at the oxygen atoms (V_min/O) are consistently more negative than those at the nitrogen sites (V_min/N), identifying the O centers as the most favorable sites for CO₂ coordination.

Fig. 5 MESP topography of anions 1_O−, 2_O− and their most stable CO₂ adducts computed at the M06-2X/cc-pVQZ level of theory. Black dots indicate V_min value in kcal mol⁻¹. Isosurface value: −125 kcal mol⁻¹. Color key: gray – H, blue – N, green – C, red – O.

Upon complexation, the most stable CO₂ adduct of 1_O− arises when CO₂ binds through its carbon atom to the anionic oxygen site, yielding an O-carboxylate species denoted 1_OCO2−. Among the possible N-carboxylate forms involving N₁ and N₃ sites, the N₃-bound adduct (1_NCO2−) is the more stable configuration (Table S3 and Fig. S4). For the 2_O− anion, however, the preferred adduct is the N-carboxylate (2_NCO2−) rather than the O-bound form (2_OCO2−). This reversal of preference suggests that while electrostatic attraction governs the initial approach, the overall stability of the adduct is dictated by charge delocalization and orbital overlap between the N lone pair and the CO₂ antibonding orbitals.

In both O- and N-carboxylate complexes, significant charge transfer from the anionic triazolate to the CO₂ moiety is reflected by highly negative V_min values on the carboxylate oxygens, confirming formation of a delocalized –OCO₂⁻ fragment. The adsorption of CO₂ on both anions is strongly exothermic (ΔE_ad ≈ −18 kcal mol⁻¹) and exergonic (ΔG_ad ≈ −6 kcal mol⁻¹) (Table 2). Between the two systems, 2_NCO2− is marginally more stable than 1_OCO2−, indicating that N-centered coordination can be energetically comparable or even superior to O-centered binding when favorable orbital interactions operate.

Table 2 Adsorption energies (ΔE_ad) and Gibbs free energies (ΔG_ad) for anionic triazolate–CO₂ adducts at the M06-2X/cc-pVQZ level of theory (kcal mol⁻¹)

Anionic systems	ΔEad	ΔGad
1OCO2−	−17.8	−5.7
2NCO2−	−18.7	−6.6
1OCO2−⋯CO2	−28.0	−7.1
2NCO2−⋯CO2	−28.3	−7.3

---

Further adsorption of a second CO₂ molecule on 1_OCO2− and 2_NCO2− was also explored (Table S4). In both cases, the additional CO₂ interacts noncovalently with the carboxylate group (Fig. 6), producing weakly bound but thermodynamically stable species 1_OCO2−⋯CO₂ and 2_NCO2−⋯CO₂. Although the second adsorption step is less exothermic (≈10 kcal mol⁻¹ stabilization), the overall process remains exergonic by up to 7 kcal mol⁻¹, suggesting that sequential CO₂ capture is feasible. The reduction in binding energy for the second CO₂ is consistent with charge delocalization over the first carboxylate unit, which diminishes localized nucleophilicity.

Fig. 6 Optimized geometries of the most stable anionic CO₂ adduct at the M06-2X/cc-pVQZ level of theory. Bond lengths in Å, bond angles in degrees and ΔE_ad in kcal mol⁻¹. Color key: gray – H, blue – N, green – C, red – O.

Interaction of dianionic triazoles with CO₂

The dianionic species 1_O2− and 2_O2−, generated via double deprotonation of the dihydroxy derivatives 1_OH2 and 2_OH2, exhibit markedly enhanced electron density compared to their monoanionic counterparts. Their MESP maps (Fig. 7) display V_min values that are approximately 98–99 kcal mol⁻¹ more negative than those of the corresponding monoanions, underscoring their exceptional nucleophilic character. Both oxygen and nitrogen centers exhibit high electron density, offering multiple reactive sites for CO₂ binding and thereby increasing the overall CO₂ capture potential.

Fig. 7 MESP topography of dianions 1_O2− and 2_O2− at the M06-2X/cc-pVQZ level of theory. Black dots represent V_min positions. Isosurface value: −250 kcal mol⁻¹. Color key: gray – H, blue – N, green – C, red – O.

The first CO₂ adsorption occurs strongly and results in the formation of N-carboxylate dianions: 1_{(NCO2−)O−} and 2_{(NCO2−)O−}, with adsorption energies (ΔE_ad) of −43.9 and −44.0 kcal mol⁻¹, respectively. The corresponding free energies (ΔG_ad) of −31.7 and −32.8 kcal mol⁻¹ confirm that CO₂ fixation on these dianions is highly exergonic and thermodynamically favorable. These adducts feature short covalent N–C bond distances (≈1.47–1.51 Å), indicative of strong donor–acceptor bonding between the triazolate nitrogen and the carbon center of CO₂.

Subsequent CO₂ adsorption yields distinct bis(carboxylate) complexes depending on the isomer (Table S5). In 1_{(NCO2−)O−}, a second CO₂ molecule binds at the remaining N site to form the N,N′-bis(carboxylate) adduct 1_(NCO2−)2, while in 2_{(NCO2−)O−}, the second CO₂ attaches to the oxygen center, producing an N,O-bis(carboxylate) complex 2_{(NCO2−)OCO2−}. Both species show large cumulative adsorption energies (ΔE_ad = −71.9 and −73.6 kcal mol⁻¹, respectively), indicating substantial stabilization upon the second CO₂ capture. The incremental binding energies of −28.0 and −29.6 kcal mol⁻¹ demonstrate that even after the first CO₂ attachment, additional CO₂ adsorption remains highly favorable.

A third CO₂ molecule can also associate with these bis(carboxylate) dianions. In 1_(NCO2−)2, the third CO₂ interacts noncovalently with the carboxylate group (O⋯C ≈ 2.60 Å), giving 1_(NCO2−)2⋯CO₂, while in 2_{(NCO2−)OCO2−}, the third CO₂ forms a partially covalent O–C bond (1.61 Å), producing 2_{(NCO2−)(OCO2−)}−CO₂. These tertiary adducts exhibit additional exothermic stabilization (ΔE_ad = −85.6 and −87.8 kcal mol⁻¹, respectively), and the corresponding free energies remain negative (ΔG_ad ≈ −54 kcal mol⁻¹), demonstrating the thermodynamic viability of sequential CO₂ capture (Table 3 and Fig. 8).

Table 3 Adsorption energies (ΔE_ad) and Gibbs free energies (ΔG_ad) for CO₂ adducts of dianionic triazolates computed at the M06-2X/cc-pVQZ level (kcal mol⁻¹)

System	ΔEad	ΔGad
1(NCO2−)O−	−43.9	−31.7
2(NCO2−)O−	−44.0	−32.8
1(NCO2−)2	−71.9	−59.1
2(NCO2−)OCO2−	−73.6	−51.1
1(NCO2−)2⋯CO2	−85.6	−68.8
2(NCO2−)(OCO2−)−CO2	−87.8	−53.5

---

Fig. 8 Optimized structures of CO₂ adduct with dianions at the M06-2X/cc-pVQZ level of theory. Distances in Å, angles in degrees and ΔE_ad in kcal mol⁻¹. Color key: gray – H, blue – N, green – C, red – O.

The systematic deepening of the adsorption energy through successive CO₂ additions reflects strong electronic cooperativity within the dianionic framework. Each CO₂ binding event delocalizes charge across multiple carboxylate units, resulting in enhanced electrostatic stabilization and partial activation of the CO₂ molecule-an essential prerequisite for further chemical transformation.

Influence of solvent and counter-cations on CO₂ adsorption

To achieve a realistic description of CO₂ capture behavior, solvent and counter-cation effects were incorporated into the theoretical model. The implicit solvation model (SMD) with ethanol as solvent and the tetramethylphosphonium (CH₃)₄P⁺ cation as the counter-ion were employed to simulate the experimental environment. This cation, although smaller than bulky trialkylphosphonium analogues, effectively reproduces their electronic characteristics while maintaining computational efficiency.²¹

In ethanol solution, both the O-carboxylate 1_OCO2− and N-carboxylate 2_NCO2− exhibit shorter adduct bond distances (without cation) compared with the gas-phase structures (O–C from 1.51 to 1.42 Å and N–C from 1.52 to 1.46 Å) (Fig. 9). These shortened bonds indicate an increase in covalent character within the adducts. Despite this, the adsorption energies in solution (ΔE_ad(s)) are reduced by about 5 kcal mol⁻¹ relative to gas-phase values due to the stabilizing solvation of the isolated anions. The adsorption process remains exergonic, confirming that CO₂ binding is thermodynamically feasible in solution (Table 4). A second CO₂ molecule binds more weakly through noncovalent interactions, leading to a reduced incremental stabilization.

Fig. 9 Optimized geometries of the most stable anionic and dianionic CO₂ adduct in ethanol solvent at the M06-2X/cc-pVQZ level of theory. Bond lengths in Å, bond angles in degrees and ΔE_ad in kcal mol⁻¹. Color key: gray – H, blue – N, green – C, red – O.

Table 4 BSSE-corrected adsorption energies (ΔE_ad(s)) and Gibbs free energies (ΔG_ad(s)) of selected anionic and dianionic triazolate–CO₂ complexes in ethanol solution computed at the M06-2X/SMD/cc-pVQZ level (kcal mol⁻¹)

System	ΔEad(s)	ΔGad(s)
1OCO2−	−12.7	0.1
2NCO2−	−13.7	−2.0
1(NCO2−)O−	−20.3	−8.6
2(NCO2−)O−	−29.1	−17.7
1(NCO2−)2	−30.4	−7.2
2(NCO2−)OCO2−	−31.8	−8.6

---

For the dianionic species, similar patterns are observed. In ethanol, 1_{(NCO2−)O−} and 2_{(NCO2−)O−} exhibit ΔE_ad(s) values of −20.3 and −29.1 kcal mol⁻¹, respectively, corresponding to decreases of 15–24 kcal mol⁻¹ relative to the gas-phase data. The bis(carboxylate) adducts 1_(NCO2−)2 and 2_{(NCO2−)OCO2−} show ΔE_ad(s) = −30.4 and −31.8 kcal mol⁻¹, with respective ΔG_ad(s) = −7.2 and −8.6 kcal mol⁻¹, demonstrating that the complexes remain exergonic even under solvated conditions (Fig. 9). Thus, although solvation weakens the intrinsic adsorption strength, charge delocalization and cooperative binding preserve favorable thermodynamics.

To further elucidate the combined influence of solvation and counter-cations, the energetics of the following solution-phase reactions were examined:

(1) [1_OH]_sol + n[CO₂]_sol + [(CH₃)₄P⁺(OH⁻)]_sol → [(CH₃)₄P⁺(CO₂)_n(1_O−)]_sol + [H₂O]_sol

(2) [2_OH]_sol + n[CO₂]_sol + [(CH₃)₄P⁺(OH⁻)]_sol → [(CH₃)₄P⁺(CO₂)_n(2_O−)]_sol + [H₂O]_sol

(3) [1_OH2]_sol + n[CO₂]_sol + [2(CH₃)₄P⁺(OH⁻)]_sol → [((CH₃)₄P⁺)₂(CO₂)_n(1_O2−)]_sol + 2[H₂O]_sol

(4) [2_OH2]_sol + n[CO₂]_sol + [2(CH₃)₄P⁺(OH⁻)]_sol → [((CH₃)₄P⁺)₂(CO₂)_n(2_O2−)]_sol + 2[H₂O]_sol

These reactions describe the interaction of hydroxy-substituted triazoles with phosphonium hydroxide in the presence of CO₂ in solution (sol), leading to the formation of the corresponding anionic or dianionic CO₂ adducts accompanied by the elimination of water. The computed reaction energetics, summarized in Table 5, cover the formation of both anionic complexes (n = 1–3) and dianionic complexes (n = 1–6), providing insight into the stability and CO₂-binding capability of these species in solution.

Table 5 Reaction energies (ΔE_rea(s)) and free energies (ΔG_rea(s)) for CO₂ capture by hydroxy and dihydroxy triazoles in ethanol in the presence of (CH₃)₄P⁺(OH), computed at the M06-2X/SMD/cc-pVTZ level (kcal mol⁻¹)

No.	Reactions in ethanol	ΔErea(s)	ΔGrea(s)
1	1OH + CO2 + (CH3)4P^+(OH)^− → (CH3)4P^+(CO2)1(1O−) + H2O	−36.0	−23.6
2	1OH + 2CO2 + (CH3)4P^+(OH)^− → (CH3)4P^+(CO2)2(1O−) + H2O	−34.4	−10.5
3	1OH + 3CO2 + (CH3)4P^+(OH)^− → (CH3)4P^+(CO2)3(1O−) + H2O	−25.5	9.0
4	2OH + CO2 + (CH3)4P^+(OH)^− → (CH3)4P^+(CO2)1(2O−) + H2O	−38.0	−25.5
5	2OH + 2CO2 + (CH3)4P^+(OH)^− → (CH3)4P^+(CO2)2(2O−) + H2O	−33.7	−10.3
6	2OH + 3CO2 + (CH3)4P^+(OH)^− → (CH3)4P^+(CO2)3(2O−) + H2O	−25.2	9.1
7	1OH2 + 2CO2 + 2 (CH3)4P^+(OH)^− → ((CH3)4P^+)2(CO2)2(1O2−) + 2H2O	−78.2	−49.4
8	1OH2 + 3CO2 + 2 (CH3)4P^+(OH)^− → ((CH3)4P^+)2(CO2)3(1O2−) + 2H2O	−72.5	−33.0
9	1OH2 + 4CO2 + 2 (CH3)4P^+(OH)^− → ((CH3)4P^+)2(CO2)4(1O2−) + 2H2O	−67.1	−17.0
10	1OH2 + 5CO2 + 2 (CH3)4P^+(OH)^− → ((CH3)4P^+)2(CO2)5(1O2−) + 2H2O	−56.6	2.9
11	2OH2 + 2CO2 + 2 (CH3)4P^+(OH)^− → ((CH3)4P^+)2(CO2)2(2O2−) + 2H2O	−86.7	−61.0
12	2OH2 + 3CO2 + 2 (CH3)4P^+(OH)^− → ((CH3)4P^+)2(CO2)3(2O2−) + 2H2O	−86.7	−49.4
13	2OH2 + 4CO2 + 2 (CH3)4P^+(OH)^− → ((CH3)4P^+)2(CO2)4(2O2−) + 2H2O	−80.1	−31.1
14	2OH2 + 5CO2 + 2 (CH3)4P^+(OH)^− → ((CH3)4P^+)2(CO2)5(2O2−) + 2H2O	−85.4	−27.0
15	2OH2 + 6CO2 + 2 (CH3)4P^+(OH)^− → ((CH3)4P^+)2(CO2)6(2O2−) + 2H2O	−79.9	−10.0

---

The formation of the anionic complexes is both exothermic and exergonic for n = 1 and n = 2, but becomes endergonic at n = 3 (Table 5). The dianionic complexes exhibit an even stronger thermodynamic driving force: for 1_O2−, CO₂ addition remains exothermic and exergonic up to n = 4, with the reaction becoming endergonic only upon adsorption of the fifth CO₂ molecule. In comparison, CO₂ binding to 2_O2− is even more favorable, showing consistently higher exothermic and exergonic character, and the onset of endergonicity appears only at the adsorption of the seventh CO₂ molecule.

It is noteworthy that for the monoanionic species 1_O−, the first CO₂ molecule forms an O-carboxylate adduct, whereas for 2_O−, the initial binding results in an N-carboxylate structure. Subsequent CO₂ molecules are sequentially adsorbed, with the second attaching to the first and the third attaching to the second, giving rise to an extended “carbonate-chain” motif (Fig. 10). This chain-like arrangement enables delocalization of the negative charge along the connected CO₂ units. A similar carbonate-chain topology is observed for the dianionic systems 1_O2− and 2_O2− (Fig. 10). In 1_O2−, the chain grows outward from an initial N-carboxylate site, with additional CO₂ molecules forming a continuous, connected framework. In contrast, in 2_O2−, the two arms of the dianion extend differently: one side develops through an N-carboxylate linkage, while the other propagates via an O-carboxylate linkage.

Fig. 10 Representative structures of triazolate–CO₂ adducts in ethanol with tetramethylphosphonium counter-ions, computed at the M06-2X/SMD/cc-pVTZ level. Distances in Å, angles in degrees; Color key: gray – H, blue – N, green – C, red – O, yellow – P.

These results demonstrate that even in polar solvents, charge-assisted and counter-cation-stabilized binding modes maintain strong thermodynamic favorability. The synergy between solvation, ion pairing, and electronic delocalization underscores the potential of triazolate-based ionic systems for practical CO₂ fixation and activation.

Conclusions

This comprehensive DFT investigation elucidates the structural, electrostatic, and thermodynamic factors governing CO₂ adsorption on 1,2,3- and 1,2,4-triazole derivatives in their neutral, monoanionic, and dianionic forms. Benchmark calculations confirmed that M06-2X/cc-pVQZ provides reliable interaction energies across the series. Molecular electrostatic potential (MESP) analyses established that CO₂ binds preferentially to electron-rich N and O sites, with deprotonation greatly enhancing nucleophilicity and charge-assisted stabilization.

Neutral triazoles exhibit only weak physisorption of CO₂ (ΔE_ad ≈ −4 to −7 kcal mol⁻¹), while the corresponding anions form strong covalent O-carboxylate (–OCO₂⁻) and N-carboxylate (–NCO₂⁻) adducts, releasing 17–19 kcal mol⁻¹ of binding energy and showing exergonic adsorption. Successive CO₂ binding on these anions remains energetically favorable through secondary noncovalent interactions. In the dianions, sequential adsorption yields bis(carboxylate) and tris(carboxylate) complexes with stepwise adsorption energies up to −88 kcal mol⁻¹, demonstrating the cooperative electronic effects of multiple charge centers and extensive delocalization.

Crucially, the inclusion of solvation (ethanol) and tetramethylphosphonium counter-cations provides a realistic assessment of CO₂ capture under experimental conditions. Although solvation reduces intrinsic binding strengths, both anionic and dianionic triazolates retain strongly exergonic CO₂ fixation. Solution-phase reaction modeling reveals spontaneous formation of mono- and poly(carboxylate) complexes, with dianions capable of exergonically binding up to six CO₂ molecules. The counter-cation plays a decisive stabilizing role by dispersing charge, modulating local electrostatics, and enabling cooperative adsorption sequences.

Thus, triazolate-based ionic systems combine multi-site O,N cooperativity, charge-delocalized activation, and counter-ion stabilization to deliver robust CO₂ capture even in polar media. These features position triazolates as promising molecular platforms for realistic CO₂ fixation, activation, and potential downstream conversion.

Author contributions

CHS conceptualized the project, supervised the work, and contributed to the writing of the manuscript. PKA and DS conducted research, curated the data, and contributed to the writing of the manuscript.

Conflicts of interest

There are no conflicts to declare.

Data availability

The data underlying this study are available in the published article and its supplementary information (SI). Supplementary information: SCF energies, zero-point correction, thermal correction to Gibbs free energies and BSSE correction (Table S6) and Cartesian coordinates of all the systems (Table S7). See DOI: https://doi.org/10.1039/d5cp04586j.

Acknowledgements

PKA is thankful to DST, Inspire, for a Senior Research Fellowship. CHS acknowledges CSIR-NIIST, Trivandrum for his deputation assignment as Director to SRIBS, Kottayam.

References

1. F. Pulido and P. Berthold, Adv. Ecol. Res., 2004, 35, 151–183.
2. K. Caldeira and J. F. Kasting, Nature, 1993, 366, 251–253.
3. D. A. Lashof and D. R. Ahuja, Nature, 1990, 344, 529–531.
4. D. M. D'Alessandro, B. Smit and J. R. Long, Angew. Chem., Int. Ed., 2010, 49, 6058–6082.
5. J. M. Nordbotten and M. A. Celia, Geological storage of CO₂: modeling approaches for large-scale simulation, New Jersey, 2011.
6. J. D. Figueroa, T. Fout, S. Plasynski, H. McIlvried and R. D. Srivastava, Int. J. Greenhouse Gas Control, 2008, 2, 9–20.
7. C. Dhoke, A. Zaabout, S. Cloete and S. Amini, Ind. Eng. Chem. Res., 2021, 60, 3779–3798.
8. J. Yu, L.-H. Xie, J.-R. Li, Y. Ma, J. M. Seminario and P. B. Balbuena, Chem. Rev., 2017, 117, 9674–9754.
9. P. Nugent, Y. Belmabkhout, S. D. Burd, A. J. Cairns, R. Luebke, K. Forrest, T. Pham, S. Ma, B. Space and L. Wojtas, Nature, 2013, 495, 80–84.
10. R. Poloni, B. Smit and J. B. Neaton, J. Am. Chem. Soc., 2012, 134, 6714–6719.
11. L. Du, Z. Lu, K. Zheng, J. Wang, X. Zheng, Y. Pan, X. You and J. Bai, J. Am. Chem. Soc., 2013, 135, 562–565.
12. J. A. Thompson, N. A. Brunelli, R. P. Lively, J. Johnson, C. W. Jones and S. Nair, J. Phys. Chem. C, 2013, 117, 8198–8207.
13. R. Babarao, R. Custelcean, B. P. Hay and D.-E. Jiang, Cryst. Growth Des., 2012, 12, 5349–5356.
14. X. Luo, Y. Guo, F. Ding, H. Zhao, G. Cui, H. Li and C. Wang, Angew. Chem., 2014, 126, 7173–7177.
15. I. Niedermaier, M. Bahlmann, C. Papp, C. Kolbeck, W. Wei, S. Krick Calderon, M. Grabau, P. S. Schulz, P. Wasserscheid and H.-P. Steinrück, J. Am. Chem. Soc., 2014, 136, 436–441.
16. S. Wang, W.-C. Li, L. Zhang, Z.-Y. Jin and A.-H. Lu, J. Mater. Chem. A, 2014, 2, 4406–4412.
17. H. Tang and C. Wu, ChemSusChem, 2013, 6, 1050–1056.
18. H. M. Lee, I. S. Youn, M. Saleh, J. W. Lee and K. S. Kim, Phys. Chem. Chem. Phys., 2015, 17, 10925–10933.
19. M. Juríček, P. H. Kouwer and A. E. Rowan, Chem. Commun., 2011, 47, 8740–8749.
20. N. Hollingsworth, S. R. Taylor, M. T. Galante, J. Jacquemin, C. Longo, K. B. Holt, N. H. De Leeuw and C. Hardacre, Angew. Chem., Int. Ed., 2015, 54, 14164–14168.
21. P. K. Arathi and C. H. Suresh, Phys. Chem. Chem. Phys., 2025, 27, 14630–14644.
22. W. Li, X. Zhao, F. Liu and B. Hou, Corros. Sci., 2008, 50, 3261–3266.
23. V. P. Krivopalov and O. P. Shkurko, Russ. Chem. Rev., 2005, 74, 339.
24. R. Kharb, P. C. Sharma and M. S. Yar, J. Enzyme Inhib. Med. Chem., 2011, 26, 1–21.
25. Y. Naito, F. Akahoshi, S. Takeda, T. Okada, M. Kajii, H. Nishimura, M. Sugiura, C. Fukaya and Y. Kagitani, J. Med. Chem., 1996, 39, 3019–3029.
26. V. V. Chaban and N. A. Andreeva, J. Mol. Liq., 2023, 378, 121623.
27. R. Boulmène, M. Prakash and M. Hochlaf, Phys. Chem. Chem. Phys., 2016, 18, 29709–29720.
28. Y. Zhao and D. G. Truhlar, Theor. Chem. Acc., 2008, 120, 215–241.
29. T. H. Dunning Jr, J. Chem. Phys., 1989, 90, 1007–1023.
30. M. J. Frisch, G. W. Trucks, H. B. Schlegel, G. E. Scuseria, M. A. Robb, J. R. Cheeseman, G. GScalmani, V. Barone, G. A. Petersson, H. Nakatsuji, X. Li, M. Caricato, A. V. Marenich, J. Bloino, B. G. Janesko, R. Gomperts, B. Mennucci, H. P. Hratchian, J. V. Ortiz, A. F. Izmaylov, J. L. Sonnenberg, D. Williams-Young, F. Ding, F. Lipparini, F. Egidi, J. Goings, B. Peng, A. Petrone, T. Henderson, D. Ranasinghe, V. G. Zakrzewski, J. Gao, N. Rega, G. Zheng, W. Liang, M. Hada, M. Ehara, K. Toyota, R. Fukuda, J. Hasegawa, M. Ishida, T. Nakajima, Y. Honda, O. Kitao, H. Nakai, T. Vreven, K. Throssell, J. A. Montgomery, J. E. Peralta, F. Ogliaro, M. J. Bearpark, J. J. Heyd, E. N. Brothers, K. N. Kudin, V. N. Staroverov, T. A. Keith, R. Kobayashi, J. Normand, K. Raghavachari, A. P. Rendell, J. C. Burant, S. S. Iyengar, J. Tomasi, M. Cossi, J. M. Millam, M. Klene, C. Adamo, R. Cammi, J. W. Ochterski, R. L. Martin, K. Morokuma, O. Farkas, J. B. Foresman and D. J. Fox, Gaussian 16, Revision A.03, Gaussian, Wallingford, CT, 2016.
31. S. F. Boys and F. Bernardi, Mol. Phys., 1970, 19, 553–566.
32. P. Politzer and D. G. Truhlar, Chemical applications of atomic and molecular electrostatic potentials: reactivity, structure, scattering, and energetics of organic, inorganic, and biological systems, Boston, 2013.
33. S. R. Gadre and R. N. Shirsat, Electrostatics of atoms and molecules, Hyderabad, 2000.
34. N. Mohan and C. H. Suresh, J. Phys. Chem. A, 2014, 118, 1697–1705.
35. P. V. Bijina and C. H. Suresh, J. Chem. Sci., 2016, 128, 1677–1686.
36. C. H. Suresh and S. Anila, Acc. Chem. Res., 2023, 56, 1884–1895.
37. C. Suresh and S. R. Gadre, J. Am. Chem. Soc., 1998, 120, 7049–7055.
38. S. R. Gadre and P. K. Bhadane, J. Phys. Chem. A, 1999, 103, 3512–3517.
39. F. B. Sayyed and C. H. Suresh, J. Phys. Chem. A, 2012, 116, 5723–5732.
40. C. Reichardt, Org. Process Res. Dev., 2007, 11, 105–113.
41. A. J. Reynolds, T. V. Verheyen, S. B. Adeloju, E. Meuleman and P. Feron, Environ. Sci. Technol., 2012, 46, 3643–3654.
42. A. V. Marenich, C. J. Cramer and D. G. Truhlar, J. Phys. Chem. B, 2009, 113, 6378–6396.

---