Superacidity of Boron Acids H₂(B₁₂X₁₂) (X=Cl, Br)

The finding that monoprotic carborane acids of the type H(CHB₁₁X₁₁) (X=Cl, Br) are the strongest pure acids isolated to date^{[1, 2]} suggests that the analogous diprotic boron acids H₂(B₁₂X₁₂) may have comparable or even higher acidity. The hydrated acids [H(H₂O)_n⁺]₂[B₁₂X₁₂²⁻] were prepared many years ago from aqueous solutions of their alkali metal salts using an ion-exchange resin in acid form^[3] and were shown to be slightly stronger acids than H₂SO₄ in aqueous solution. However, the anhydrous acids are unknown. Can they be synthesized? Are they superacids? Of further interest is that the B₁₂H₁₂²⁻ ion starting material is considerably cheaper than the isoelectronic CB₁₁H₁₂⁻ carborane ion, so there is potential for cost savings relative to carborane acids, which are too expensive for widespread use.

An indication that the all-boron acids H₂(B₁₂X₁₂) should show comparable acid strengths to their analogous carborane acids H(CHB₁₁X₁₁) comes from the position of their anions on the ν(NH) basicity scale. In this ranking, the N–H stretching frequencies of contact-ion-pair trioctylammonium salts Oct₃N⁺–H⋯anion⁻ are compared in CCl₄ solution.^{[2, 4]} The higher the ν(NH) frequency, the lower the basicity of the anion. As shown in Table 1, the B₁₂Cl₁₂²⁻ salt has almost the same ν(NH) frequency as the CHB₁₁Cl₁₁⁻ salt. This result is surprising inasmuch as the dinegative charge on the B₁₂Cl₁₂²⁻ ion might have been expected to render it more basic than the uninegative CHB₁₁Cl₁₁⁻ ion. It suggests that chloride substituents on both anions form an effective screen for negative charge that is delocalized and buried within the icosahedral cage. Similar conclusions can be drawn from the data on the B₁₂Br₁₂²⁻ ion, although low solubility of its trioctylammonium salt in CCl₄ allow comparisons to be made only for crystalline salts. It is also clear from the data in Table 1 that the boron anions are much less basic than the (HSO₄)₂²⁻ ion, so their conjugate acids are expected to be much stronger than 100% H₂SO₄, whose acidity (H₀=−12 on the Hammett scale) defines the onset of superacidity.

Table 1.

ν(NH) frequencies of tri-n-octylammonium salts (in cm⁻¹).

Anion	ν(NH) in CCl4	ν(NH) solid
CHB11Cl11−[a]	3163	3180
B12Cl122−	3165	3167
CHB11Br11−[a]	3140	3150
B12Br122−	insoluble	3140
(HSO4)22−	3021, 2660	3080

^[a]Ref. [4].

The synthetic pathway to the anhydrous diprotic acids exploits multistep metathesis reactions similar to those used in the preparation of H(CHB₁₁Cl₁₁).^[1] Starting with the silver salts of B₁₂Br₁₂²⁻ and B₁₂Cl₁₂²⁻, their respective trityl (C-(C₆H₅)₃⁺) salts [Eq. (1)] and triethylsilylium compounds [Eq. (2)] have been prepared and characterized by X-ray crystallography.^[5]

$$\mathrm{A}{\mathrm{g}}_2[{\mathrm{B}}_{12}{\mathrm{X}}_{12}]+2{({\mathrm{C}}_6{\mathrm{H}}_5)}_3\mathrm{C}\mathrm{B}\mathrm{r}\to {[{({\mathrm{C}}_6{\mathrm{H}}_5)}_3\mathrm{C}]}_2[{\mathrm{B}}_{12}{\mathrm{X}}_{12}]+2\mathrm{A}\mathrm{g}\mathrm{B}\mathrm{r}$$ (1)

$${[{({\mathrm{C}}_6{\mathrm{H}}_5)}_3\mathrm{C}]}_2[{\mathrm{B}}_{12}{\mathrm{X}}_{12}]+2\mathrm{E}{\mathrm{t}}_3\mathrm{S}\mathrm{i}\mathrm{H}\to {(\mathrm{E}{\mathrm{t}}_3\mathrm{S}\mathrm{i})}_2({\mathrm{B}}_{12}{\mathrm{X}}_{12})+2\mathrm{H}\mathrm{C}{({\mathrm{C}}_6{\mathrm{H}}_5)}_3$$ (2)

The structures of the trityl salts are unremarkable and similar to that of the B₁₂F₁₂²⁻ ion,^[6] but those of the triethylsilylium derivatives are informative with regard to how the dianions act as Lewis bases. As illustrated in Figure 1 for the B₁₂Cl₁₂²⁻ ion, the triethylsilyl groups adopt 1,12 or trans positions, presumably to minimize electrostatic repulsions between the electropositive silicon centers and to pack most efficiently in the crystal lattice. The sum of the C-Si-C angles, whose approach towards 360° is used as a measure of developing silylium ion character in an R₃Si^δ+ moiety,^{[7, 8]} is 348.3°. This value is close to 349.5° in Et₃Si(CHB₁₁Cl₁₁),^[9] thus indicating similar low Lewis basicities of the B₁₂Cl₁₂²⁻ and CHB₁₁Cl₁₁⁻ ions.

Figure 1.

X-ray crystal structure of (Et₃Si)₂(B₁₂Cl₁₂).

A further illustration of the comparable Lewis basicities of the B₁₂Cl₁₂²⁻ and CHB₁₁Cl₁₁⁻ ions comes from the formation of the hydride-bridged cation [Et₃Si-H-SiEt₃]_2⁻ [B₁₂Cl₁₂] when excess silane is used in Equation (2). This cation forms reversibly when silane competes with the anion for coordination to the Et₃Si⁺ moiety and is readily identified by its ν(Si-H-Si) band at approximately 1870 cm⁻¹ in the IR spectrum of the isolated solid.^[9] Similarly, halocarbon solvents such as o-dichlorobenzene (ODCB) are sufficiently basic to compete with the anion for coordination to the Et₃Si⁺ moiety.^[9] Since excess silane and solvent molecules can interfere with the subsequent formation of the unsolvated H₂(B₁₂X₁₂) acids, thus lowering their purity, considerable care must be taken to remove them from the silylium derivatives by drying under vacuum. Such contaminants include toluene that can be carried through from the trityl salt preparation.

Silylium ion character in the (Et₃Si)₂B₁₂X₁₂ derivatives is sufficiently strong that chloride abstraction from anhydrous HCl proceeds readily according to Equation (3) at sub ambient temperatures to give the desired acids H₂(B₁₂X₁₂) as nearly colorless solids in essentially quantitative yields.

$${(\mathrm{E}{\mathrm{t}}_3\mathrm{S}\mathrm{i})}_2({\mathrm{B}}_{12}{\mathrm{X}}_{12})(\mathrm{s})+2\mathrm{H}\mathrm{C}\mathrm{l}(\mathrm{l})\to {\mathrm{H}}_2({\mathrm{B}}_{\mathrm{12}}{\mathrm{X}}_{12})(\mathrm{s})+2\mathrm{E}{\mathrm{t}}_3\mathrm{S}\mathrm{i}\mathrm{C}\mathrm{l}(\mathrm{g})$$ (3)

Both H₂(B₁₂Br₁₂) and H₂(B₁₂Cl₁₂) are distinctively characterized in their solid-state IR spectra by broad, low-energy absorptions associated with short, strong, symmetrical X-H-X hydrogen bonding.^{[10, 11]} As shown in Figure 2, these bands appear at 1200 and approximately 620 cm⁻¹ for H₂(B₁₂Cl₁₂) compared to 1100 and 615 cm⁻¹ in H(CHB₁₁Cl₁₁). In the analogous bromide, they appear at 1080 and approximately 570 cm⁻¹. The lower-energy band is particularly strong, but its exact frequency is made uncertain by Evans holes.^[11] Given the overarching tendency of strongly acidic protons to attain symmetrical (or nearly symmetrical) di-solvation,^[10] and by analogy to H(CHB₁₁Cl₁₁) whose X-ray structure shows Cl-H⁺-Cl bridges in a 2D polymeric chain structure,^[11] a 3D network structure with X-H⁺-X bridges is the only reasonable structure for H₂(B₁₂X₁₂) acids. This 3D structure would also explain why these acids cannot be sublimed under vacuum up to 200°C for purification and possible single-crystal growth, even though carborane acids can be vaporized around 150°C.^[11]

Figure 2.

IR spectrum of H₂(B₁₂Cl₁₂) showing the two broad absorptions at 1200 (Gaussian fit shown by dashed line) and ca. 620 cm⁻¹ (with Evans holes) characteristic of symmetrical Cl-H⁺-Cl bridge bonding. The major anion absorption is marked with an asterisk.

The acids have relatively low solubility in low-basicity solvents such as anhydrous liquid SO₂, but the appearance of a downfield peak at δ=19.3 ppm in the ¹H NMR spectrum of H₂(B₁₂Cl₁₂) is consistent with ionization to the give the disolvated H(SO₂)₂⁺ ion.^[11] The solid anhydrous acids are moisture-sensitive. As monitored by IR spectroscopy, the gradual appearance of broad ν(OH) bands in the range 3200–3600 cm⁻¹ occurs as the broad bands associated with the acid X-H⁺-X group disappear, indicating the formation of H₅O₂⁺ and H₇O₃⁺ ions.^[12]

The protonation of benzene has become a benchmark of superacidity above and beyond that of mineral acids.^[2] The strongest traditional acids, triflic acid (H₀=−14) and fluorosulfuric (H₀=−15.1), do not protonate benzene, whereas carborane acids do.^[13] This behavior indicates that the equivalent Hammett acidities of carborane acids must be −17 or greater. Like carborane acids, the Hammett acidities of the present acids cannot be measured directly because they are unavailable in liquid form.

Stirring a suspension of H₂(B₁₂X₁₂) in the smallest possible volume of benzene (to minimize water contamination) converts the acids into slurries of their corresponding benzenium ion salts, formulated as [C₆H₇]₂[B₁₂X₁₂]. The presence of the C₆H₇⁺ ion is established by IR spectroscopy, with diagnostic bands in the ν(CH) and ν(CC)/δ(CH)s regions. In particular, the ν(CH₂) frequencies of the most acidic CH₂ group at the protonated sp³ carbon atom appear as a broad absorption centered at 2733 cm⁻¹ for X=Cl and 2664 cm⁻¹ for X=Br. These frequencies are lower than those in the gas phase (2810, 2795 cm⁻¹)^[14] owing to H-bonding to halide atoms of the anions. The greater red shift for the bromide is consistent with the higher basicity of the B₁₂Br₁₂²⁻ ion relative to the B₁₂Cl₁₂²⁻ ion. As shown in Table 2, other identifying bands are all very close in frequency. IR bands attributable to X-H⁺-X bridged moieties are absent, thus indicating that H(B₁₂X₁₂)⁻ ions, which would result from monoionization of the diprotic acids, must be absent. H₂(B₁₂X₁₂) behaves as a diprotic acid: both hydrogen ions are sufficiently acidic to protonate benzene. The extraordinary stability of the B₁₂X₁₂²⁻ ions makes the isolation of the benzenium ion at room temperature straightforward. As expected for more basic arenes, the 2:1 arenium ion salts derived from protonating toluene and mesitylene can be isolated in a similar manner and characterized by IR spectroscopy (see the Supporting Information).

Table 2.

Frequencies of the C₆H₇⁺ ion as a function of counterion (in cm⁻¹).

Counterion	ν(CH2) (average)	ν(CH)aromatic	ν(CC) + δ(CCH)
none[a]	2810, 2795 (2803)	3110, 3080
CB11H6Cl6−[b]	2770, 2720 (2745)	3100, 3072, 3040, 3028	1601
B12Cl122−	2753, 2713 (2733)	3092, 3070, 3035	1600
CB11H6Br6−[b]	2757, 2714 (2736)	3095, 3073, 3066, 3023	1600
B12Br122−	2664	3084, 3060, 3028, 3015	1597

^[a]Ref. [14].

^[b]Ref. [13].

In summary, the new anhydrous diprotic acids H₂(B₁₂X₁₂) with X=Cl and Br have been prepared. They have comparable acidity to carborane acids H(CHB₁₁X₁₁). Because they protonate benzene, their acidities must be at least 10₅ greater than that of 100% H₂SO₄, that is, well into the superacidity region. These are the first examples of diprotic superacids, and the finding that both protons of H₂(B₁₂X₁₂) are able to protonate benzene is significant inasmuch as the second ionization might be expected to be of lower strength than the first. Finally, the all-boron acids are less expensive to prepare than their analogous carborane acids, although this advantage is tempered by the generally lower solubilities of B₁₂X₁₂²⁻ salts compared to those of CHB₁₁X₁₁⁻.^[15]