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Lithium Monoxide Anion: a Ground-State Triplet with the Strongest Base to Date

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Lithium Monoxide Anion: a Ground-State Triplet with the Strongest Base to Date Lithium monoxide anion: A ground-state triplet with the strongest base to date Zhixin Tian*, Bun Chan†, Michael B. Sullivan†‡, Leo Radom†§, and Steven R. Kass*§ *Department of Chemistry, University of Minnesota, Minneapolis, MN 55455; and †School of Chemistry and Centre of Excellence for Free Radical Chemistry and Biotechnology, University of Sydney, Sydney NSW 2006, Australia Edited by W. Carl Lineberger, University of Colorado, Boulder, CO, and approved March 13, 2008 (received for review February 12, 2008) ؊ Lithium monoxide anion (LiO ) has been generated in the gas 430 ؊ phase and is found to be a stronger base than methyl anion (CH3 ). LiOH ؊ This makes LiO the strongest base currently known, and it will be CH4 a challenge to produce a singly charged or multiply charged anion LiNH2 410 that is more basic. The experimental acidity of lithium hydroxide is ؊1 NH3 kJ) and, when LiCH3 4.184 ؍ kcal⅐mol (1 kcal 6.1 ؎ 425.7 ؍ H°acid⌬ combined with results of high-level computations, leads to our LiBH2 acid ؊ o ⅐ 1 390 H2O ؎ best estimate for the acidity of 426 2 kcal mol . H computations ͉ mass spectrometry ͉ super base 370 HF he gas-phase acidities of the hydrogen halides were first Treported via the application of a thermodynamic cycle (Eqs. LiH 1–5) in 1942 by Briegleb (1, 2). 350 0.5 1.5 2.5 3.5 4.5 3 ⅐ ϩ ⅐ ͑ ͒ HX H X BDE HX [1] Electronegativity ⅐ ϩ Ϫ ⅐ H 3 H ϩ e IE͑H ͒ [2] Fig. 1. Gas-phase acidities of first-row hydrides (HX) and their lithiated analogs versus Pauling electronegativities of X [circles, experimental values ⅐ Ϫ Ϫ ⅐ X ϩ e 3 X ϪEA͑X ͒ [3] (14); squares, BD(T)/aug-cc-pVQZ theoretical values obtained in the present study]. 3 ϩ ϩ Ϫ ⌬ Њ ͑ ͒ HX H X H acid HX [4] ⌬ Њ ͑ ͒ ϭ ͑ ͒ ϩ ͑ ⅐͒Ϫ ͑ ⅐͒ Hacid HX BDE HX IE H EA X [5] A general method for producing metal-containing anions that overcomes these practical problems was developed by Bachrach, In subsequent years, the acidities of thousands of compounds Hare, and Kass (10) and subsequently exploited by O’Hair et al. have been measured by using a variety of techniques (3), and the (11–13). In this approach, metal salts of dicarboxylates are ⅐ Ϫ1 ϭ acidity scale currently spans a 125 kcal mol (1 kcal 4.184 kJ) produced by electrospray ionization (ESI) and fragmented via ⌬ ϭ Ϯ ⅐ Ϫ1 range from CH4 ( Hacid° 416.8 0.7 kcal mol ) (4, 5) to energetic collisions (CID), thereby leading to the sequential ⌬ ϭ Ϯ ⅐ Ϫ1 HN(SO2C4F9)2 ( Hacid° 291.1 2.2 kcal mol ) (6) [see expulsion of two molecules of carbon dioxide but retention of the supporting information (SI) Text]. Methyl anion is the strongest metal ion. For example, the conjugate base of phenyllithium was base currently known, which is a position it has occupied for the formed from the lithium salt of 1,2-benzenedicarboxylate, as CHEMISTRY past 30 years. This raises the question as to whether a more basic shown in Eq. 6. species can be made. In this article, we use sophisticated experimental techniques and state-of-the-art theoretical calcu- lations to show that the lithium monoxide anion (LiOϪ) is in fact [6] more basic than methyl anion, and that it will be a challenge to produce a species that is still more basic. Alkyl groups are polarizable but also are generally electron- This methodology provides a predictable and rational means releasing and, depending on which influence is larger, can for making ions that are difficult to prepare in other ways. In this destabilize anions. Kinetic measurements indicate that ethane article, we report its use to synthesize LiOϪ and determine the and the secondary position of propane [(CH3)2CH2] are 2–3 acidity of lithium hydroxide via Eq. 5 because our preliminary Ϫ kcal⅐molϪ1 less acidic than methane (7), but their conjugate bases high-level computations indicated that the LiO ion is extremely have never been observed (8, 9). This is not surprising because basic. the electron affinity of methyl radical is only 1.8 Ϯ 0.7 kcal⅐molϪ1 Ϫ Ϫ (4), and CH3CH2 and (CH3)2CH are predicted to be unbound Author contributions: Z.T., B.C., M.B.S., L.R., and S.R.K. performed research; Z.T., B.C., with respect to electron detachment (7). Electronegative sub- M.B.S., L.R., and S.R.K. analyzed data; L.R. and S.R.K. designed research; and L.R. and S.R.K. stituents stabilize negative ions and increase acidities, as re- wrote the paper. Ͼ flected by the first-row hydrides [i.e., HF (most acidic) H2O The authors declare no conflict of interest. Ͼ Ͼ NH3 CH4 (least acidic)]. To decrease the acidity of a This article is a PNAS Direct Submission. compound and make a stronger base, one might employ an ‡Present address: Institute of High Performance Computing, Singapore 117528. electropositive substituent such as lithium. However, the conju- §To whom correspondence may be addressed. E-mail: [email protected] or gate bases of lithiated compounds are difficult to prepare in the [email protected]. gas phase, and almost nothing is known about them because the This article contains supporting information online at www.pnas.org/cgi/content/full/ neutral acids tend to be involatile, moisture sensitive, and 0801393105/DCSupplemental. pyrophoric. © 2008 by The National Academy of Sciences of the USA www.pnas.org͞cgi͞doi͞10.1073͞pnas.0801393105 PNAS ͉ June 3, 2008 ͉ vol. 105 ͉ no. 22 ͉ 7647–7651 Downloaded by guest on September 27, 2021 Table 1. Computed BD(T)/aug-cc-pVQZ and CAS-AQCC/aug-cc- O 1.212 pVQZ acidities of HX at 298 K Li O Li 1.214 C ⌬ ⅐ Ϫ1 1.988 H°acid, kcal mol 1.843 C 1.256 1.287 O O O O 2.032 O 2.274 HX BD(T) AQCC Experimental O C 1.279 C 1.242 1.734 LiBH2 395.1 393.7 1.746 Li Li LiCH 401.6 402.8 3 singlet triplet singlet triplet LiNH2 413.5 414.7 –1 LiOH 425.0 426.2 Erel = 0.0 4.9 3.2 4.7 kcal mol LiSH 375.8 376.0 1 (C2v) 2 (Cs) LiH 355.8 356.3 356.0 Ϯ 0.1* ϩ Ϫ Fig. 3. Computed B3-LYP/6-311 G(2df,2pd) structures for LiCO2 and W1 BeH2 393.4 395.9 relative energies. BH3 412.1 412.2 † CH4 418.8 419.2 416.8 Ϯ 0.7 Li2BH 385.3 384.8 alkali metal substituents lead to enhanced acidities in these Li2CH2 399.8 400.1 cases. However, the opposite effect can also be observed. For Li2NH 417.6 419.7 example, substitution of a lithium for a hydrogen in ammonia NaCH3 401.2 402.0 and water leads to weaker acids [i.e., ⌬H° ϭ 403.4 Ϯ 0.1 (NH ) NaOH 418.6 419.7 acid 3 vs. 413.5 [LiNH2, BD(T)] and 390.27 Ϯ 0.02 (H2O) vs. 425.0 NaSH 382.0 381.5 ⅐ Ϫ1 ‡ [LiOH, BD(T)] kcal mol ], which also are predicted to be less NH3 403.7 404.9 403.4 Ϯ 0.1 ‡ acidic than methyllithium. This reflects the acidities of lithiated H2O 390.4 394.1 390.27 Ϯ 0.02 HF 371.8 374.1 371.331 Ϯ 0.003‡ compounds, which display the opposite trend to that of first-row hydrides with Pauling electronegativity values (Fig. 1) (16). *Ref. 3. Because LiOϪ is computed to be 6.2 kcal⅐molϪ1 [BD(T)] more Ϫ †Ref. 4. basic than CH (see Table 1) and is predicted to be a ground- ‡ 3 Ref. 14. state triplet ion, this diatomic anion represents an interesting but challenging experimental target. Earlier work on metal salts of dicarboxylate anions (10) Results and Discussion Ϫ suggests that lithium carbonate (LiCO ) might be a good Electronegative substituents are well known to stabilize negative 3 precursor for synthesizing LiOϪ. The former species was readily ions and increase acidities, and this is reflected in a plot of generated by ESI, but its fragmentation under a variety of ⌬Hacid° (HX) vs. the electronegativity of X for first-row hydrides Ϫ (Fig. 1) (14). Electropositive substituents such as lithium and conditions did not afford LiO as hoped for. Instead, signal loss sodium might be expected to show the opposite behavior, but was observed. BD(T)/aug-cc-pVQZ [referred to simply as BD(T) hereafter; CID see Theoretical Procedures] calculations on methyllithium Ϫ O¡ Ϫ ϩ [7a] Ϫ1 LiOCO2 // LiO CO2 (⌬Hacid° ϭ 401.6 kcal⅐mol ) (15) and methylsodium (⌬Hacid° ϭ 401.2 kcal⅐molϪ1) indicate that both of these substrates are 15–16 ⅐ Ϫ1 kcal mol more acidic than methane. In other words, these O¡ signal loss [7b] Ϫ Lithium oxalate (LiC2O4 , m/z 95) was subsequently examined because it could lose carbon dioxide followed by carbon mon- oxide to afford the target species. This sequence indeed takes place as anticipated (Fig. 2), with the lithium salt of doubly Ϫ deprotonated formic acid (LiCO2 , m/z 51) being initially formed. OO O CID Ϫ O¡ Ϫ ϩ [8] LiO–C–C–O LiO–C CO2 m/z 95 m/z 51 This ion also can be viewed as a CO2 solvate of lithium anion, and was briefly explored. Ϫ Computations indicate two low-lying structures for LiCO2 (Fig.
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