Formation and investigation of unstable complex hydrides from tetrabutylammonium fluoride, water and SO₂

Abstract

A phase separation occurred when SO₂ was introduced into the binary system of tetrabutylammonium fluoride (TBAF) and water system. The mechanics of this phase separation process were studied by Fourier transform-infrared spectroscopy (FT-IR), Raman spectroscopy and Karl-Fischer titration. A complex hydrides of TBAF–SO₂–H₂O was formed and the TBAF and SO₂ concentrated in the lower layer, and it could be released by heating the solution under reduced pressure (382.2 K, 10.1 kPa). The equilibrium solubility of SO₂ in the binary system of tetrabutylammonium fluoride and water was investigated. After desorption, the mixture could be reused to absorb SO₂. The effect of temperature and concentration of TBAF on the mole ratio of SO₂–TBAF and H₂O–TBAF of the triple hydrate was studied. The binary system of TBAF and water could absorb SO₂ efficiently.

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Introduction

SO₂ emissions are significant sources of atmospheric pollution. SO₂ can cause the formation of acid rain, which is a serious environmental concern. SO₂ emissions contribute to the formation of smog, which is a significant human health concern. If a practical process to remove SO₂ selectively and reversibly could be identified, these emissions, along with their corresponding effects on the environment and human health, could be minimised. The identification of a material that can selectively and reversibly capture SO₂ has proven difficult. Liquid amine aqueous solutions can react with SO₂ and trap it in the form of ammonium sulphite. However, the amines can evaporate into the gas stream due to their volatility. Furthermore, it is difficult to desorb the SO₂ from ammonium sulphite. A promising alternative is to use ionic liquids (ILs) to capture SO₂. ILs are non-flammable and have desirable chemical and physical characteristics, such as low volatilities, high thermal stabilities, and high solvation capacities.^1,2

The removal of SO₂ from flue gas with ionic liquids has been the focus of experimental attention,^3–7 but a practical method has yet to be identified. Huang found that the prepared dimethylethanolamine ILs had excellent SO₂ absorption/desorption performance. The absorbed SO₂ was stored in the IL buffers in the form of HSO₃⁻ to maximize the absorption.⁸ Tian synthesized a new functional IL 1-butyl-3-methylimidazolium lactate ([Bmim]L) and used to absorb SO₂. The comparison of [Bmim]L with other ILs based on lactate anion suggests that the cations of ILs have a significant influence on the absorption and desorption behaviors of SO₂.⁹ The imidazolium-based ILs with lactate anion is promising absorbents for the removal of SO₂ from the flue gas. Ren¹⁰ investigated the influence of sulfuric acid on absorption and desorption of SO₂ in the monoethanolammonium lactate ([MEA]L). The absorption capacity of SO₂ in [MEA]L with sulfuric acid decreased greatly. After NaOH, CaO, or CaCO₃ was added into the mixture of [MEA]L and sulfuric acid, [MEA]L could be regenerated. Lee¹¹ suggested that SO₂ was irreversibly absorbed in [BMIm][OAc]. During the absorption process, most of acetate anion dissociated from [BMIm][OAc] converted into acetic acid upon contact with SO₂ at 50, resulting in the formation of new room temperature ILs with [BMIm][HOSO₂]. Moreover, [BMIm][HOSO₂] recovered after removing acetic acid could be used as a new and reversible SO₂ absorbent. And corresponding sulfuric salts was formed and precipitated for removal. Caprolactam (CPL)–tetrabutylammonium bromide (TBAB) ionic liquid (CPL : TBAB = 1 : 1 mole ratio) has a high affinity for SO₂ (mole fraction solubility 0.680, 298.2 K) and a low affinity for diatomic gases such as N₂ and O₂, as well as H₂S.^12–14 However, the high viscosity (1500 mPa at 303.2 K) of the CPL–TBAB IL makes mass and heat transfer difficult. The removal of SO₂ from flue gas was therefore investigated using a mixture of CPL–TBAB IL and water.¹⁵ A second phase appears when SO₂ is introduced into the solution. SO₂ could be a separation switch for water and the caprolactam–tetrabutylammonium bromide ionic liquid. Liu prepared five inorganic salts based eutectic ionic liquids (EILs), acetamide–KSCN (3 : 1), caprolactam (CPL)–KSCN (3 : 1), acetamide–NH₄SCN (3 : 1), CPL–NH₄SCN (3 : 1) and urea–NH₄SCN (3 : 2), which absorb a large amount of SO₂ gas.¹⁶ In the efficient desulfurization process of [CPL][TBAB] IL and sodium humate (HA-Na) solution, humic acid (HA) can be sulfonated to some extent. Then the compound fertilizer consisting of sulfonated HA and sulfates were gained.¹⁷ Tetrabutylammonium halide amine semi-clathrate hydrate was found to be as absorbent for acidic gases such as CO₂ and H₂S.^18–20 We find that a second phase appears when SO₂ is introduced into tetrabutylammonium fluoride (TBAF) aqueous solutions, also. To get the interaction among the TBAF, water and SO₂, the reaction mechanisms between SO₂ and TBAF aqueous solutions were investigated.

In this paper, the two phases that separate when SO₂ is introduced into TBAF and water mixtures were characterised by Fourier transform-infrared spectroscopy (FT-IR), Raman spectroscopy and Karl-Fischer titration, and the mole ratio of SO₂, water and TBAF in each phase were measured and calculated at different temperature. The equilibrium solubility of SO₂ in various compositions of TBAF and water was determined and correlated at different temperature. The viscosity of the TBAF and water mixtures was measured and correlated, also. The mechanics of phase separation and the possibility of recycling the solutions were explored.

Experimental section

Materials

The TBAF aqueous solutions were prepared by mixing TBAF (CAS no. 429-41-4, C₁₆H₃₆NF, 99.0% purity) and de-ionised water. The individual TBAF and water components were weighed using a Brookfield balance with an uncertainty of ±0.0001 g. SO₂ with a purity of 99.9% was provided by Beijing Analytical Instrument Factory, Beijing, China.

Viscosity measurement

Viscosities of the binary mixtures of TBAF and water were measured using LVDV-II + Pro viscometer, which was obtained from Brookfield-Instruments, USA. Ultra Low Adapter (ULA) was used to increase accuracy. And the cover was used to prevent the evaporation of water. The temperature was controlled with a temperature accuracy of ±0.1 K.

Absorption of SO₂ in mixtures of TBAF and water

The absorption experimental apparatus in the mixture of TBAF and water was similar to others reports.¹⁶ The temperature was controlled by the water bath with a temperature accuracy of ±0.1 K. SO₂ gas was bubbled through the predetermined amounts of TBAF aqueous solutions (about 10 g) in glass vessels at a rate of 10 mL min⁻¹. Unabsorbed SO₂ was neutralized with sodium hydroxide. The glass vessel was weighed using a balance with an uncertainty of ±0.0001 g to get the mass of SO₂ absorbed, through which the solubility of SO₂ in the mixture can be calculated. Considering the absorption rate, the equilibrium can be considered to be reached after 2 h in the thermodynamic experiments.

Recycling the TBAF

The two phases mixtures of TBAF, water, and SO₂ (Scheme 1, L₁ and L₂) were separated using a separator funnel. L₁ and L₂ (the SO₂-enriched phase) were gained. SO₂ was recycled from L₂. An oil bath, fitted with a temperature controller and a vacuum indicator, was used to heat L₂ to release the SO₂ and the little amount of water. During the desorption process, the weight of L₂ was periodically determined; experiments showed that SO₂ in L₂ was desorbed completely after heating for 2 h at 383.2 K and 10.1 kPa vacuum. The binary mixture, TBAF and water, of L₁ and the residue of L₂ recycled SO₂ and little water was reformed. After added the appropriate amount of water, the TBAF aqueous solution was again subjected to SO₂ absorption (308.2 K, 101.3 kPa). When the mixture was saturated with SO₂, it was separated again.

Scheme 1 Schematic of TBAF–water–SO₂ phase behaviour.

Results and discussion

Equilibrium solubility of SO₂ in the TBAF aqueous solution

TBAF and water can be miscible at special proportions at ambient conditions. When SO₂ was introduced into the mixture of the TBAF and water, a second liquid phase appeared; this process was represented in Scheme 1. This phenomenon is similar to what happens when CO₂ is added to 1-n-hexyl-3-methylimidazolium bis(trifluoromethylsulfonyl)imide ([hmim]-[Tf₂N]) + acetonitrile (10% mol IL) and when SO₂ is added to the caprolactam tetrabutylammonium halide ionic liquids.^15,21 The densest phase was rich in TBAF and SO₂ (labelled L₂), whereas the least dense phase was rich in water (labelled L₁). The amounts of densest phase decreased clearly with the temperature increase and the compositions of TBAF decrease (Schemes 2 and 3). The content of SO₂ and TBAF in the L₂ phase decreased with the concentration of TBAF increasing. The content of SO₂ in the L₂ phase decreased with temperature increasing, whereas the content of TBAF in the L₂ phase increase with temperature increasing.

Scheme 2 Schematic of TBAF–water–SO₂ phase behaviour with respect to temperature (a) 283.2 K, (b) 288.2 K, (c) 293.2 K, (d) 298.2 K, (e) 303.2 K.

Scheme 3 Schematic of TBAF–water–SO₂ phase behaviour at different mole fraction of TBAF and water solution at 283.2 K; (a) 0.05 mol, (b) 0.02 mol, (c) 0.01 mol, (d) 0.005 mol.

The equilibrium solubility (S, g per 100 g solutions) of SO₂ (obtained in mixture of L₁ and L₂) in different compositions of TBAF and water was presented in Fig. 1. The solubility of SO₂ in TBAF and water solutions decreased when the temperature increased and the concentration of TBAF decreased. The reason for SO₂ solubility relationship with temperature is very similar to the reason that vapor pressure increases with temperature. Meanwhile, the effect of temperature on the solubility of SO₂ in the TBAF and water solutions was very clear because of the weak intermolecular bonds, also. The kinetic energy increases with increasing temperature. As the temperature is raised, gases usually become less soluble in TBAF aqueous solutions. For instance, the solubility of SO₂ in the mixture of TBAF and water (0.05 mol L⁻¹) was 34.11 g at 278.2 K and decreased to 16.22 g at 308.2 K, which was similar with the absorbtion capacity for SO₂ by the mixture of CPL–TBAB IL and water.

Fig. 1 Solubility (S) of SO₂ in different concentrations of TBAF and water solutions as a function of temperature: 0.001 mol L⁻¹; 0.01 mol L⁻¹; 0.02 mol L⁻¹; 0.05 mol L⁻¹.

The viscosity of the TBAF and water (mole fraction 0.05, 11.3 mPa s at 303.2 K) solution (Fig. 2) was far lower than that of the CPL–TBAB IL (1500 mPa s at 303.2 K) and the mixture of CPL–TBAB L and water (0.29, 191.4 mPa s at 303.2 K).¹⁵ And the mixture of TBAF and water could be reused easily and economically. Therefore, the TBAF and water solutions (mole fraction 0.05) have excellent properties to efficiently capture SO₂.

Fig. 2 Viscosity (η) of the TBAF and water solutions as a function of temperature: 0.001 mol L⁻¹; 0.01 mol L⁻¹; 0.02 mol L⁻¹; 0.05 mol L⁻¹.

The solubility of SO₂ in different compositions of TBAF and water solutions appeared to be non-linear with respect to temperature (Fig. 1). For the purposes of application, the SO₂ solubility in the binary system of TBAF and water was expressed as a function of temperature (eqn (1)):¹⁵

S = a + b(T/K) + c(T/K)² + d(T/K)³ (1)

In this equation, S represents the mass fraction solubility value; T is the absolute temperature; and a, b, c and d refers to the fit coefficients. The values of the parameters a, b, c, d and R-square (R²) are listed in Table 1.

Table 1 Parameters for eqn (1) and R-square values for the equilibrium solubility of SO₂ in the binary system of TBAF and water with respect to temperature

S	a (× 10^−3)	b (× 10^−1)	c (× 10)	d (× 10^4)	R^2
0.001	13.61	−13.42	4.416	−4.850	0.9989
0.01	8.136	−7.796	2.495	−2.666	0.9972
0.02	0.2119	0.3389	−0.2848	0.4986	0.9995
0.05	−16.84	17.52	−6.041	6.913	0.9997

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The viscosities of TBAF and water solutions were measured with respect to temperature. These data were fit to an equation, which was first proposed by Vogel–Fulcher–Tammann (VFT) and subsequently used by other researchers (eqn (2)).^15,22 The fitted data are shown in Fig. 2.

η/mPa s = A exp{B/[(T/K) − C]} (2)

In this equation, η represents the viscosity value; T is the absolute temperature; and A, B and C refer to the fit coefficients. The values of the parameters A, B, C and R-square (R²) are listed in Table 2.

Table 2 Parameters for eqn (2) and R-square values for the viscosity measurements of different concentrations of TBAF and water solutions with respect to temperature

η	A	B	C	R^2
0.001	2.519	3487	−19.67	0.858
0.01	6.089	2988	46.08	0.9960
0.02	3.739	2988	46.67	0.9833
0.05	0.001922	1623	116.1	0.9898

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Characterisation of L₁ and L₂

L₁ and L₂ were characterised by FT-IR spectroscopy (Fig. 3) at different compositions of TBAF–H₂O–SO₂. Unionized SO₂ molecular stretches (1330 cm⁻¹), as stretch S–O (1175 cm⁻¹) and stretch S–O (1089 cm⁻¹) were clearly observed in L₂, which were obscure in L₁. There is a large amount of H₂O (about 90.59% of the total amount of water) and small quantities of TBAF (about 30.12% of the total amount of TBAF) and SO₂ (about 0.91% of the total amount of SO₂) in L₁. SO₂ was captured and enriched in the lower layer (L₂). In L₂, SO₂ and TBAF concentrations were high (about 99.09% and 69.88% of the total amount of SO₂ and TBAF), and the concentration of H₂O was low (about 9.41% of the total amount of water). The SO₂ was removed from L₂ by heating under reduced pressure. During the recycle process of SO₂, less energy was used to recycle SO₂ from L₂ than that of the mixtures of SO₂ and water, ammonia or organic amines, because of little of water evaporating. Moreover, the interaction forces between TBAF, water and SO₂ was low. So little energy was needed during the reuse of SO₂. The intensity of unionized SO₂ molecular stretches (1330 cm⁻¹) increased with the increasing of the composition of SO₂. Some of the H₂SO₃ from SO₂ and H₂O were ionized due to the S–O. The F⋯H–O asymmetric stretch is also an intense IR vibration, red shifted by 739 cm⁻¹ from the O–H local mode stretch for H₂O and H₂SO₃.

Fig. 3 FT-IR spectra of the layers from the TBAF–H₂O–SO₂ system at different temperature: L₂₁: 283.15 K; L₂₂: 293.15 K; L₂₃: 303.15 K; L₁: 293.15 K.

From Fig. 4 we could know that the peak of L₂ with greater intensities is almost not found that of L₁. The peak of HO–S (1150 cm⁻¹) in L₂ is blue shift. The intermolecular force of HO–S in L₂ is higher than that of in L₁.

Fig. 4 Raman spectra of the L₁ and L₂ (the mole fraction of TBAF aqueous solution is 2.54%, and absorbed SO₂ at 293.15 K).

The SO₂-saturated TBAF and water solution at different temperature and concentration of TBAF aqueous solution was characterised by Raman spectroscopy and Karl-Fischer titration. The mole ratios of SO₂–TBAF and H₂O–TBAF in the L₂ phase are shown in Fig. 5. From the data summarised in Fig. 5, it is clear that the mole ratio of SO₂–TBAF increases clearly from (3.5 to 4.7) with decreasing concentration of TBAF from (0.05 to 0.02 mol). And the mole ratio of SO₂–TBAF increases clearly from (1.9 to 3.5) with decreasing temperature from (308.2 K to 283.2 K). However the mole ratio of H₂O–TBAF increases sharply with decreasing concentration of TBAF and increasing temperature. As the temperature increases, the interactive forces between the molecule of SO₂ and water and TBAF become weaker and weaker, this is very similar to the reason that vapor pressure increases with temperature also. As TBAF is such a strong hydrogen bond acceptor it is near impossible to dry hydrated samples. So, the planar water in strong hygroscopic TBAF molecular increase with the increasing temperature.

Fig. 5 Effect of temperature and concentration of TBAF on the mole ratio of SO₂–TBAF and H₂O–TBAF: the mole ratio of SO₂–TBAF at concentration of TBAF of 0.05 mol; the mole ratio of SO₂–TBAF at concentration of TBAF of 0.02 mol; the mole ratio of H₂O–TBAF at concentration of TBAF of 0.05 mol; the mole ratio of H₂O–TBAF at concentration of TBAF of 0.02 mol.

Reusability of TBAF

After the two layers (L₁ and L₂) were separated, the SO₂ was removed from L₂ by heating under reduced pressure (383.2 K, 10.1 kPa). The composition of the regenerated mixture was almost identical to that of the original solution, except that a small amount of water was lost. The results of five consecutive absorption/desorption cycles with the TBAF–water solution are shown in Fig. 6. The performance of the TBAF and water mixture did not significantly change after repeated absorption/desorption cycles.

Fig. 6 SO₂ adsorption (mass percent) in TBAF and water (0.05 mol) (a) and SO₂ desorption rate (mass percent) in L₂ (b) over five adsorption/desorption cycles.

Conclusion

TBAF and water can be miscible at special proportions at ambient conditions. A complex hydrides of TBAF–SO₂–H₂O was formed and the TBAF and SO₂ concentrated in the lower layer, and it could be released by heating the solution under reduced pressure (382.2 K, 10.1 kPa). The mole ratio of SO₂–TBAF increases clearly with decreasing concentration of TBAF and decreasing temperature. However the mole ratio of H₂O–TBAF increases sharply with decreasing concentration and increasing temperature. The solubility (g per 100 g solutions) of SO₂ in TBAF and water increased when the concentration of TBAF was increased from 0.001 to 0.05 mol at (278.2 K to 308.2 K). The viscosity of L₁ and L₂ is lower than that of TBAF aqueous solutions. Moreover, the absorption of excessive SO₂ by TBAF and water is physisorption. So it is benefit to recycle TBAF and SO₂. The binary system of TBAF and water could absorb SO₂ efficiently.

Acknowledgements

This research is supported by Natural Science Foundation of China (no. 21106033), Science and Technology Department of Hebei (no. B2012208037), Hebei Education Department (BJ2014024, Y2011107 and 2011167) and Hebei University of Science and Technology (no. XL201227).

Notes and references

1. J. D. Holbrey, W. M. Reichert, R. G. Reddy and R. D. Rogers, Ionic Liquids as Green Solvents: Progress and Prospects, American Chemical Society, Washington DC, 2003.
2. E. D. Bates, R. D. Mayton, I. Ntai and J. H. Davis, J. Am. Chem. Soc., 2002, 124, 926.
3. J. Huang, A. Riisager and W. B. Rolf, J. Mol. Catal. A: Chem., 2008, 2, 170.
4. J. Huang, A. Riisager, P. Wasserscheid and R. Fehrmann, Chem. Commun., 2006, 38, 4027.
5. Y. Wang, C. Wang, L. Q. Zhang and H. R. Li, Phys. Chem. Chem. Phys., 2008, 10, 5976.
6. D. An, L. B. Wu and S. P. Zhu, Macromolecules, 2007, 40, 3388.
7. W. Z. Wu, B. X. Han, H. X. Gao, Z. M. Liu, T. Jiang and J. Huang, Angew. Chem., Int. Ed., 2004, 43, 2415.
8. Y. Wang, H. H. Pan, H. R. Li and C. Wang, J. Phys. Chem. B, 2007, 111, 10461.
9. K. Huang, J. F. Lu, Y. T. Wu, X. B. Hu and Z. B. Zhang, Chem. Eng. J., 2013, 215–216, 36.
10. S. H. Ren, Y. C. Hou, S. D. Tian, W. Z. Wu and W. N. Liu, Ind. Eng. Chem. Res., 2012, 51, 3425.
11. K. Y. Lee, C. S. Kim, H. Kim, M. Cheong, D. K. Mukherjee and K. D. Jung, Bull. Korean Chem. Soc., 2010, 31(7), 1937–1940.
12. B. Guo, E. H. Duan, A. L. Ren, Y. Wang and H. Y. Liu, J. Chem. Eng. Data, 2010, 55, 1398.
13. P. Luis, L. A. Neves, C. A. M. Afonso, I. M. Coelhoso, J. G. Crespo, A. Garea and A. Irabien, Desalination, 2009, 245, 485.
14. M. B. Shiflett and A. Yokozeki, Ind. Eng. Chem. Res., 2010, 49, 1370.
15. E. H. Duan, B. Guo, M. M. Zhang, Y. N. Guan, H. Sun and J. Han, J. Hazard. Mater., 2011, 194, 48.
16. B. Y. Liu, F. X. Wei, J. J. Zhao and Y. Y. Wang, RSC Adv., 2013, 3, 2470.
17. Y. Zhao and G. X. Hu, RSC Adv., 2013, 3, 2234.
18. W. Lin, A. Delahaye and L. Fournaison, Fluid Phase Equilib., 2008, 264, 220.
19. K. Yasushi, Y. Yukiyasu, E. Takao, O. Hiroyuki, S. Wataru and N. Hideo, Energy Fuels, 2005, 19, 1717.
20. S. Fan, S. Li, J. Wang, X. Lang and Y. Wang, Energy Fuels, 2009, 23, 4202.
21. B. R. Mellein and J. F. Brennecke, J. Phys. Chem. B, 2007, 111, 4837.
22. K. R. Harris, M. Kanakubo and L. A. Woolf, J. Chem. Eng. Data, 2007, 52, 1080.

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