Alkali atoms possess the lowest ionization energy (IE), ranging from 5.39 to 3.89 eV, among all the elements in the periodic table. However, superalkalis are clusters whose ionization energies are even lower than this range. These clusters were originally introduced by Gutsev and Boldyrev in 1982 using sp-block elements [15]. They proposed species like Li₂F, Li₃O, and Li₄N as superalkalis. In the form of a superatom, these clusters impersonate the behavior of alkali atoms. There have been several studies on the design of various kinds of superalkalis [17–24]. For instance, the binuclear superalkalis including F₂Li_3, have been widely studied [17]. Hou et al. [18] described non-metallic binuclear cations such as F₂H₃ ⁺ and O₂H₅ ⁺. Zhao et al. [19] proposed some special superalkalis like N₄Mg₆Li, Al₃, Mn(B₃N₃H₆)₂, B₉C₃H₁₂, Al₁₂P, and C₅NH₆, which were designed by using different schemes like jellium rule, 18-electron rule, Wade–Mingos rule, and Huckle’s rule, respectively. Al₁₂P was reported to be an alkali–metal-like superatom [20, 21]. Recently, Sikorska and Gaston [22] reported the superalkali behavior of polynuclear N₄Mg₆M for M = Li, Na, and K. Srivastava [25] noticed that the IE of C₆Li₆, being lower than that of Li, makes it a closed-shell superalkali. The IE scale of these superalkalis is depicted in Figure 1.

Superalkalis find applications in the design of supersalts [26–28], superbases [29–31], alkaloids [32–34], and so forth. Due to their low IE, superalkalis play an important role in chemical industries as reducing agents. Here, we provide an account of how superalkalis are exploited to activate the CO₂ molecule.

Srivastava [35] studied the interaction of CO₂ with FLi₂, OLi₃, and NLi₄ superalkalis using the second-order Møller–Plesset perturbative (MP2) method [39] and the 6–311+G(d) basis set in the Gaussian 09 program [40]. Such interaction leads to the formation of complexes, as shown in Figure 3, and corresponding parameters can be found in Table 1. It is evident that the minimum energy of these complexes corresponds to the structure in which the interaction between CO₂, and superalkalis is mediated by both the O atoms of CO₂. The bond length of Li-O lies between 1.865 and 1.892 Å. The low-lying isomers of FLi₂-CO₂ and OLi₃-CO₂ are of higher energy in which CO₂ interacts via a single atom, whereas in the case of NLi₄-CO₂, there are no competing isomers obtained.

The binding energy (BE) of superalkali–CO₂ complexes is calculated and listed in Table 1. The BE of these complexes monotonically decreases with the increase in the size of superalkalis. This can be explained on the basis of a more delocalized electron cloud that is generally associated with the larger superalkalis. The natural population analysis (NPA) [45] charges (∆q) on CO₂ have also been listed. The most stable structure of superalkali–CO₂ complexes takes the values of ∆q as −0.90e for FLi₂, −0.88e for OLi₃, and −0.85e for NLi₄. In isomer (b) of FLi₂-CO₂, ∆q has a very small magnitude (−0.17e), which is consistent with an almost linear CO₂ moiety just as in a neutral CO₂ molecule. On the contrary, the ∆q in isomer (b) of OLi₃-CO₂ is, albeit smaller than that in its lowest energy structure (a), large enough to bend the CO₂ moiety. It should be noted that the size of the superalkalis is a more important factor than their IE in CO₂ activation. As per calculated binding energy and charge transfer, FLi₂ is more effective for CO₂ reduction. In these complexes, the CO₂ moiety is bent by 133° and the bond length C-O becomes 1.246 Å which is comparable to the bond lengths of 1.237 Å and bond angle 137° in the CO₂ ⁻ anion, obtained at the MP2/6-311+G(d) level. This study suggests simple and catalyst-free single-electron reduction of CO₂ by using typical superalkalis such as FLi₂, OLi₃, and NLi₄.

Park and Meloni [41] reported the interaction of CO₂ and superalkali species Li₃F₂ using the CBS-OB3 composite model [46] through the Gaussian 09 program. They obtained three isomers, two planar (a) and (b), as well as one non-planar (c), of the Li₃F₂-CO₂ complex, as shown in Figure 4. There was no appreciable change in bond lengths between Li and F on interaction with CO₂. Despite some structural changes in the superalkalis, the structure of CO₂ changes from linear to bending. Therefore, it is clear that the strongly bound CO₂ is activated upon interaction with the superalkali. From Table 1, the BE of Li₃F₂-CO₂ isomers is found to be in the range of 1.06–1.63 eV (106–163 kJ/mol). The lowest BE was obtained for the isomer (c) in which one oxygen interacts with both the terminal Li atoms and the other oxygen with the central Li atom. The isomer (a) possesses greater BE and therefore stronger intermolecular interaction than isomer (b) because the electron density is more localized between the two terminal Li atoms and two oxygen atoms.

The BE of Li₃F₂-CO₂ complexes is comparable to or smaller than the BE of superalkali (FLi₂, OLi₃, NLi₄)-CO₂ complexes reported by Srivastava [35]. The charge transfer to CO₂ in Li₃F₂-CO₂ isomers ranges from −0.63e to −0.88e (see Table 1). Thus, the charge on the CO₂ moiety along with its structure suggests that CO₂ is reduced to a CO₂ ⁻ anion.

Zhao et al. [19] presented the rational design of superalkalis and studied the activation of CO₂ by these special superalkalis using MP2/6-311+G(d) level in the Gaussian 09 program. They analyzed the interaction of CO₂ with special superalkalis like Al₃, Mn (B₃N₃H₆)₂, B₉C₃H₁₂, and C₅NH₆, which leads to the complexes shown in Figure 5. The distance between the CO₂ moiety and superalkali clusters has been calculated as 1.950 Å, 1.730 Å, 2.320 Å, and 1.040 Å for Al₃, Mn (B₃N₃H₆)₂, B₉C₃H₁₂, and C₅NH₆, respectively. In the case of Al₃, O-C bonds extend to 1.290 and 1.510 Å, about 4.4–22% longer than those in the CO₂ ⁻anion, whereas in B₉C₃H₁₂, the O-C bonds are extended to 1.290 and 1.280 Å, about 3.3–4.4% longer than those in CO₂ ⁻. The bond extension in Mn (B₃N₃H₆)₂ and C₅NH₆ is observed to be 1.250 Å which is slightly longer than that of 1.240 Å in CO₂ ⁻ and 1.35 Å which is about 9.3% longer than that in CO₂ ⁻, respectively. The bond angle of O-C-O in Al₃CO₂, Mn (B₃N₃H₆)₂CO₂, B₉C₃H₁₂CO₂, and C₅NH₆CO₂ is 126°, 133°, 131°, and 129° making the bond bend by 8°, 3, 4, and 7% more than the corresponding value in CO₂ ⁻. Thus, both the stretching of the O-C bonds and the bending of the O-C-O angle weaken the O-C bonds of CO₂, making it easy to activate.

The NPA charge has been listed in Table 1 to show how much charge is transferred to CO₂. This transfer of charge results in the bending of CO₂ and weakens the CO₂ bond, and therefore making it easier to break. The amount of charge transferred from Al₃, Mn (B₃N₃H₆)₂, B₉C₃H₁₂, and C₅NH₆ to CO₂ are −1.26e, −0.90e, −0.91e, and −0.63e, respectively. Note that this amount of charge transferred from Al₃, Mn (B₃N₃H₆)₂, B₉C₃H₁₂ are greater than that of noble gas (0.77e) [47] being very close to unity, whereas in the case of C₅NH₆, the amount of charge transferred is less. From this analysis, one may note that although the IE of Al₃ is not the lowest among these four superalkalis (see Figure 1), the charge transferred is the most and it is capable of bending the CO₂ molecule the most. This indicates that the quantitative nature of the activation of CO₂ depends on the electronic structure and size of the superalkalis, as seen in an earlier section.

Kumar et al. [42] explored the scope of non-metallic superalkalis in the activation of CO₂. They studied the interaction of CO₂ with non-metallic superalkalis such as O₂H₅ and N₂H_7, employing the MP2/6-311++G (d,p) level via the Gaussian 09 program. The equilibrium structures of O₂H₅-CO₂ and N₂H₇-CO₂ are shown in Figure 6, and related parameters are listed in Table 1. It was noticed that in O₂H₅-CO₂ complexes, O atoms of CO₂ interact with the H-atom of superalkali, unlike in N₂H₇-CO₂, in which the C-atom of CO₂ interacts with the H-atom. This may be due to the repulsion between excess electrons of N and O atoms.

The (b) isomers of O₂H₅-CO₂ and N₂H₇-CO₂ are 0.24 and 0.78 eV, higher in energy in which CO₂ interacts through the O-atom as well as the C-atom in the N₂H₇-CO₂ isomer. The relative stability of isomers can be explained on the basis of H-bond interactions. For instance, the bond lengths of O-H and C-H are 2.060 Å and 3.710 Å, respectively. The BE of these complexes provides relative strength through the interaction of CO₂ with non-metallic superalkalis. From Table 1, the BE suggests that O₂H₅-CO₂ isomers are stable, whereas N₂H₇-CO₂ is slightly destabilized due to the negative value of BE.

From Table 1, the value of NPA charges of CO₂ is calculated to be −0.75e for O₂H₅-CO₂ and −0.77e for N₂H₇-CO₂ lowest energy structures (a). Thus, the NPA charge values are very close to each other. In isomer (b) of O₂H₅-CO₂ and N₂H₇-CO₂, the CO₂ moiety is bent, similar to that in its lowest energy structure. Therefore, the activation and the consequential reduction of CO₂ can also be possible by non-metallic superalkalis such as O₂H₅, if not by N₂H₇.

The compact (quasi) icosahedral Al₁₂X (X = Be, Al, C, and P) clusters have been employed to analyze the dissociation and absorption of small gas molecules [48–54]. Zhang et al. [43] studied the interaction of Al₁₂P superalkali with CO₂ using Minnesota density functional (M06-2X) [55] and 6-311+G(d) basis set in the Gaussian 09 program. They obtained three isomers of the Al₁₂P-CO₂ complex as shown in Figure 7. The lowest energy corresponds to the isomer (a) in which the interaction is mediated by both O atoms. The isomers (b) and (c) are found to have a high energy of 0.11 eV (2.64 kcal/mol) and 0.20 eV (4.52 kcal/mol), respectively. Obviously, the chemisorbed CO₂ molecule undergoes structural changes from linear to bending in each Al₁₂P-CO₂ isomer.

They confined their analyses to the global minimum structure, that is, isomer (a). The bond distance of C-O in isomer (a) is 1.24 Å, which is found to be 7.6% larger than free CO₂ (1.15 Å), consequently, weakening the C-O bond. Moreover, the variation of the C-O-C angle from 180° to 130.4° in isomer (a) advocates the change in hybridization of carbon in CO₂ from sp to quasi-sp ² after activation by Al₁₂P. Furthermore, the C-O bond is marginally larger (1.23 Å) as compared to isolated CO₂ ⁻, whereas bond bending is also larger, about 4.7%, than that in isolated CO₂ ⁻, which clearly supports the activation of CO₂ assimilated on the Al₁₂P cluster. The computed total NPA charge on the CO₂ subunit is −0.707e, which shows the transfer of almost one electron charge from Al₁₂P to CO₂ in the complex. Thus, CO₂ is successfully reduced to CO₂ ⁻ anion. The low IE of Al₁₂P superatom is the main source of CO₂ reduction as it facilitates the transfer of charge to CO_2, which ultimately results in the contraction of the O-C-O angle and the weakening of the C-O bond of the CO₂ moiety. The small activation barrier of 23 kcal/mol is calculated for the chemisorption of CO₂ on Al₁₂P to form the Al₁₂P-CO₂ complex (a), which further suggests the application of Al₁₂P as a potential catalyst for CO₂ conversion. It has been found that the Al₁₂P complex shows high adsorption intensities in the visible region and, hence, promotes photocatalysis or photothermal catalysis of CO₂ and its transformation by absorbing sunlight.

Recently, Sikorska and Gaston [22] explored new superalkali species, N₄Mg₆M (M = Li, Na, K) by performing the MP2/6-311+G(d) and single-point CCSD(T)/6-311+G (3df) calculations Ref. 56 in the Gaussian 09 program. They studied the catalytic behavior of N₄Mg₆Li, N₄Mg₆Na, and N₄Mg₆K for CO₂ activation. For the sake of brevity, we will discuss the interaction of CO₂ with N₄Mg₆Li superalkali. CO₂ interacts with N₄Mg₆M and the resultant N₄Mg₆M-CO₂ complexes are displayed in Figure 8 for M = Li. In all isomers, the interaction between CO₂ and N₄Mg₆M takes place via the C-N bond of 1.452–1.454 Å, which supports the evolution of a single bond between carbon and nitrogen.

The isomers (b) and (c) are found to be 0.04 eV (0.97 kcal/mol) and 0.11 eV (2.48 kcal/mol) higher in energy than the lowest energy isomer (a) in the case of N₄Mg₆Li-CO₂. The binding energy of these isomers lies in the range of 1.46–1.57 eV (see Table 1). The C-O bond lengths, being in the range of 1.223–1.224 Å and 1.335–1.336 Å are indeed 8% larger than that of the CO₂ anion. Furthermore, the bending of the angle O-C-O, 123–126°, was 9% more than the angle of isolated CO₂ ⁻. The NPA charge transferred from N₄Mg₆M superalkalis to CO₂ varies from −0.799e to −0.806e. Ionization energy plays a major role in the activation process of CO_2, as the amount of transferred charge increases with the decrease in ionization energy because it is effortless to transfer charge from species with low ionization energy. Therefore, the extension of the bond distance of C-O along with the bending of angle O-C-O results in the weakening of C-O bonds of CO₂. Thus, the newly designed N₄Mg₆M superalkalis could be used as a catalyst for CO₂ activation.

The interaction of CO₂ with the binuclear Li₃F₂ superalkali was reported by Park and Meloni as discussed in an earlier section. Recently, Meloni et al. [44] investigated the interaction of CO₂ with Li₃F₂ inside fullerene (C₆₀) by using B3LYP [57, 58] with the 6-31G (d) basis set in the Gaussian 09 program. They noticed two important features. When CO₂ is encapsulated within C₆₀, it gets destabilized as its binding energy is −147 kJ/mol. The Li₃F₂ inside C₆₀ assumes trigonal bipyramidal (D_3h) geometry with a binding energy of 119 kJ/mol, unlike the free Li₃F₂ linear cluster (see Figure 1). Thus, C₆₀ strongly interacts with Li₃F₂ and there is no reduction of C₆₀.

The unforeseen result has been noticed on insertion of CO₂ inside Li₃F₂(D₃h)@C₆₀ endofullerene, as displayed in Figure 9. On the inspection of the geometry inside the C₆₀, it was found that CO₂ has been activated by making a ∠OCO 132° and that the bond length of C-O has been increased to 1.20 Å. The activation of CO₂ has been attained by the transfer of F atom from Li₃F₂ to CO₂, due to the F-C interaction with the bond distance of 1.38 Å. Thus, the endo-reaction simulates a non-planar (trigonal pyramidal) FCO₂ interacting with the FLi₃-like species (also shown in Figure 9). There have been several studies [59, 60] in which the structures and interactions between species are greatly modified by encapsulation within C₆₀.

Thus, the strong reducing power enables superalkalis to reduce CO₂ as well as several other molecules [61, 62], So far, it has been found that the superalkalis are restricted to activate one CO₂ molecule per unit, that is, only one CO₂ molecule is reduced by a superalkali. This may limit the capture and storage of superalkalis for practical applications. Srivastava [63] reported that a single molecule of hexalithiobenzene (C₆Li₆) is not only capable of reducing but also capturing up to six CO₂ molecules sequentially using the ωB97xD functional [64] and 6-311+G(d) basis set in the Gaussian 09 program. The planar C₆Li₆ molecule has equal ring bond lengths of 1.418 Å, whose IE is reported to be lower than that of Li, thereby characterizing it as a superalkali molecule [25]. Note that planar star-like C₆Li₆ has been previously studied by several groups [65–67]. The sequential interaction of CO₂ molecules with C₆Li₆ results in the C₆Li₆-nCO₂ complexes as displayed in Figure 10.

The interaction between CO₂ and C₆Li₆ is mediated by one C-C and two Li-O bonds in these complexes. The C₆Li₆ ring moiety in C₆Li₆-nCO₂is deviated from planarity due to the out-of-plane displacement of Li atoms because of the Coulomb-repulsion between neighboring Li-atoms for n = 1, 2, 4, and 5. However, both C₆Li₆-3CO₂ and C₆Li₆-6CO₂ possess the perfect planar ring moiety, having equal bond lengths of 1.42 Å and 1.40 Å, respectively. The CO₂ moiety in C₆Li₆-CO₂ has a bond length and an angle of 1.26 Å and 122.5°, respectively. With an increase in the number of CO₂, the average bond length and the average bond angle of CO₂ decrease and increase continuously up to 1.25 Å and 125.2°, respectively, for C₆Li₆-6CO₂.

The activation of CO₂ leads to an increase in the bond length and a decrease in the bond angle by bending. Table 2 lists the NBO charge (∆q) located at CO₂ moieties, adsorption (binding) energy (E _ad) per CO₂ molecule, and consecutive adsorption energy (ΔE _ad). It is clear that the Δq of C₆Li₆-nCO₂ becomes −0.83e, −0.81e, −0.79e, and −0.77e for n = 1–3, n = 4, n = 5, and n = 6, respectively. This suggests that C₆Li₆ can be employed in the activation of all CO₂ molecules and, consequently, their adsorption. One can note that the E _ad values are fairly large, ranging from 3.18 to 2.79 eV per CO₂, which predicts the stability of these C₆Li₆-nCO₂ complexes. It is to be noticed that the adsorption of molecules is not feasible with negative ΔE _ad values, whereas sequential adsorption becomes feasible with positive ΔE _ad values. Thus, the study suggests that C₆Li₆ is not only capable of activation but also effective in the sequential adsorption of six CO₂ molecules.