Ionic Liquids with Benzenesulfonate Anions: Non-Fluorinated, Thermally Stable Anion Options

Abstract

Thermally resistant materials have been sought after for use as lubricants, heat transfer fluids, high temperature structural materials, and other applications where thermal stability is required or desired. Herein, we present a new class of thermally robust ionic liquids containing inexpensive benezenesulfonate anions with profound long-term, high-temperature aerobic stability – i.e., no mass loss in 96 hours at 300 °C in air. A coherent correlation between melting and glass transition temperatures and the location and type of the anions was observed. Our work indicates that these systems can be designed to form thermally stable, low-melting organic salts, providing valuable design insights for engineering of their structure–property–function relationships.

Graphical Abstract

Introduction

Since 1998, when ionic liquids (ILs) arguably began to emerge as a forefront area of chemical research, more than 82,000 papers and patents have been published on the topic.¹ ILs are fundamentally understood as functional organic materials in which structure begets function, and the functionality and performance of ILs is of utmost concern. Consequently, it is hard to ignore the level of attention the field has garnered and the growing impact it is having on processes of commercial importance.^2,3 Even so, from the outset, concern was expressed about the practicality of using ILs in large volumes, usually because of their cost. In turn, the cost of many ILs is frequently driven by the anion of the ion pair. We note that many of the most widely used anions for IL formulations are fluorinated and fluorination is expensive and requires toxic reagents.

Their expense notwithstanding-fluorinated anions bring much to the table with regards to the properties they impart to the overall ion pair (IL). For example, they tend to be non-nucleophilic, non-basic, and poorly-coordinating.^4,5 As a consequence, they are inert towards the cation with which they are paired, decreasing the likelihood of it being de-quaternized, especially at elevated temperatures, thereby destroying the ionic character of the materials.⁶ In addition, their weak coordinating power means that they are less likely to compete for active sites on metal-based catalysts that might be dissolved in the ILs. Fluorinated anions also tend to impart hydrophobic character to ILs incorporating them. This is indispensable in situations where the IL is to be employed in aqueous biphasic catalysis or in extractions from an aqueous phase. Finally, several fluorinated anions are among the most thermally stable anions known.⁷

Despite the advantages provided by fluorinated anions, they also come with significant drawbacks. A major one – cost – has already been noted. Fluorination, whether to create wholly inorganic anions such as hexafluorophosphate PF₆⁻ or organic anions, e.g., bis(trifluoromethylsulfonyl)imide (CF₃SO₂)₂N⁻, is quite expensive and involves highly specialized skills and equipment. Another vital disadvantage is one which is even more important: safety. Handling F₂ for the synthesis of fluorinated building blocks poses considerable safety concerns. Furthermore, some fluorinated anions – particularly BF₄⁻ and PF₆⁻ – can thermally decompose and generate highly noxious (e.g. BF₃ and PF₅) and fatal HF.^8,9 Finally, an emerging concern involves the potential for fluorinated organics to prove toxic. For example, the perfluorooctanoate anion is the conjugate base of perfluorooctanoic acid (PFOA), which was recently classified as carcinogenic.¹⁰

With such factors in mind, one major IL manufacturer has spoken of the existence of an “anion crisis”.¹¹ This lament was a commentary on the relative scarcity of anions well-suited for IL formation as compared to cations. They further described desirable new anions as ones that would ideally form salts of “sufficiently” low melting point values, and be non-basic, non-nucleophilic, readily prepared, inexpensive, thermally, and hydrolytically stable, and have little or no halogen content.

As part of a larger, successful effort to design and prepare ILs of enhanced thermal stability,⁶ we recently screened a number of both off-the-shelf and easily synthesized anions to establish their relative thermal stabilities [all when paired with the reference cation PPN⁺, triphenyl(P,P,P-triphenylphosphine imidato-kN)-phosphorus(1+)].¹² Some of the anions in that study, such as Tf₂N⁻, are already known to be thermally stable and are widely utilized in formulating ILs; these were included to provide benchmarks against which the performance of less-widely used species might be compared. In that study, others such as acetate, benzoate, and benzenesulfonate were included not because they were thought a priori to be of particular promise, but simply because they were low-hanging fruit – inexpensive and easy to include in the study. It is noteworthy that D’Anna, F. et al., reported the synthesis and thermal studies of dicationic imidazolium-type ILs with two naphthalenedisulfonate anions. Measured by thermal gravimetric analysis (TGA), these ILs showed relatively high decomposition temperatures under the inert conditions.¹³

In the course of the foregoing work, we unexpectedly observed that the PPN⁺ salt of the benzenesulfonate anion (BZS⁻) had remarkably stable behavior – remaining unchanged after heating in air at 300 °C for 96 hours – an apparent match to that of the PPN⁺ salt of the previously mentioned, highly utilitarian Tf₂N⁻ anion.¹⁴ We found this surprising since arenesulfonic acids, the conjugate acids of arenesulfonates, are known to have low degrees of thermal stability.¹⁵ However, it appears that the mechanism by which these species thermally decompose requires participation (hence the presence) of the sulfonyl proton.¹⁶ The upshot is that in its absence this process does not occur – thus accounting for the stability of the aforementioned benzenesulfonate salt.¹⁷

Considering the former outcome, we concluded that it is worthwhile to prepare and determine the melting points (T_m) and thermal stabilities of a wider variety of benzesulfonate salts in association with typical cations used in IL formulations. Not only is BZS⁻ very inexpensive and environmentally innocuous, it would be expected to have many of the other chemical attributes noted as being desirable for use in IL formulations. Consequently, it seemed possible that new ILs based on this anion might be identified that could take the place of salts with fluorinated anions (on an application-dependent basis). We are especially interested in this possibility because of its potential impact on the development of inexpensive new ILs well-suited for heat transfer and solar thermal energy storage applications.

Results and discussion

Using a common and straightforward anion-exchange (a.k.a anion metathesis) protocol (Scheme S1), we created new ILs that are combinations of seven cations (1–7) and three versions of the BZS⁻ anions (A–C), each beginning with a halide (Cl⁻, Br⁻, or I⁻) salt of the corresponding cations (Figure 1). We chose cations 1-ethyl-3-methylimidazolium C₂mim⁺ (1), 1-butyl-3-methylimidazolium C₄mim⁺ (2), 1-methyl-1-propylpyrrolidinium C₃mpyrr⁺ (3), and n-tetrabutylphosphonium P₄₄₄₄⁺ (4) for their wide-use. In addition to the benzesulfonate anions, we decided to synthesize BZS-based ILs using our previously known thermally stable cations: tetraphenylphosphonium Ph₄P⁺ (5), (p-phenoxyphenyl)triphenylphosphonium (Ph)₃POP⁺ (6), and PPN⁺ (7). We initially prepared six salts having the BZS⁻ anion (1A–6A, Figure 1), each beginning with a halide salt of the corresponding cation. For the aforementioned compounds, we present thermophysical data for three sets of ILs: 1) benzenesulfonate salts, 2) m-chlorobenzenesulfonate, and 3) PPN⁺ salts.

Figure 1.

Structures of benzenesulfonate anions (A–C) and the cations (1–7) used in this study. Cations 1–4 were chosen due to their ubiquity and to compare the stability of the BZS⁻ anion. Cations 5–7 were created and studied previously by our group and shown to have impressive thermal stability.¹⁴

Per our standard approach for evaluating thermal stability,¹⁴ these salts were charred in porcelain crucibles without covers and were heated in air in a muffle furnace for 96 hours at temperatures of 100 °C, 150 °C, 200 °C, 250 °C, and 300 °C in order to assess (via mass loss and NMR spectroscopic changes) their thermal stabilities. The preheating and postheating NMR spectra (¹H, ¹³C, and ³¹P) for each treated salt were obtained and displayed in the SI.

In Table 1, all salts are listed with their T_onset values and their long-term mass loss data. Over the course of 96 hours at a specific temperature, percent mass loss that occurred at each temperature was recorded along with any prominent visual observations of changes to the materials. Notably, decomposition of the ILs occurs well below the T_onset, a phenomenon that has been written about extensively.¹⁸ Indeed, decomposition for the ILs with the BZS⁻ anions begins at a minimum of 50 °C below their T_onset, and their mass loss is accounted for due to the volatilization of the degraded cation or anion. While long-term isothermal studies are more representative of realistic thermal stability, T_onset still provides useful comparisons about the relative stability of the species for investigating reasons for decomposition. Interestingly, [PPN][BZS] (7A) and [PPN][Cl-BZS] (7B) exhibit almost no mass loss at 300 °C and exhibit exceptionally high T_onset temperatures in comparison to other BZS⁻ and Cl-BZS⁻ salts (possibly because of the contributions from π-stacking and unlike the melting process). In turn, the Ph₄P and Ph₃P-POP cations, which have been shown to be thermally robust and unchanged at 300 °C when paired with Tf₂N⁻,^17,19 shown somewhat diminished but still very good thermal stability when paired with BZS⁻.

Table 1.

Onset decomposition temperature with short-term TGA at a ramp rate of 10 °C/min with air. Long-term thermal stability tests were performed at their respective temperatures for 96 hours in a muffle furnace.

Temperature of degradation Tonset (±5 °C)		Percent mass loss (±1%) as a function of temperature in 96 h					Qualitative Remarks
		100 °C	150 °C	200 °C	250 °C	300 °C
Benzenesulfonate salts
1A	317	0	1	5	10	91	Charred at 300 °C
2A	319	0	0	0	7	93	Charred at 300 °C
3A	291	0	0	0	1	66	Begins to char between T = 150 °C – 200 °C+
4A	337	6	12	16	49	85	Begins to char between T = 200 °C – 250 °C+
5A	318	0	0	0	5	17	Appearance at 300 °C slightly darker than unheated material
6A	338	0	0	0	6	31	Appearance at 300 °C slightly darker than unheated material
m-chlorobenzenesulfonate salts
1B	323	0	0	5	6	91	Charred at 300 °C
2B	326	0	0	2	13	80	Charred at 300 °C
3B	309	1	3	5	6	61	Begins to char between T = 150 °C – 200 °C+
4B	344	0	0	1	24	81	Begins to char between T = 200 °C – 250 °C+
5B	347	0	0	0	3	15	Appearance at 300 °C slightly darker than unheated material
6B	359	0	0	0	2	21	Appearance at 300 °C slightly darker than unheated material
PPN+ salts
7A	403	0	0	1	1	0	Appearance at 300 °C unchanged from unheated material
7B	415	0	0	5	5	3	Appearance at 300 °C unchanged from unheated material
7C	378	0	0	1	4	90	Charred at 300 °C

Table 2 shows measurements of the melting points (T_m), glass transition temperatures (T_g), and enthalpy of fusion (ΔH_fus). Entropies of fusion (ΔS_fus) were calculated by dividing the enthalpy of fusion by the absolute temperature, as ΔG_fus for a phase transition is equal to zero (see SI). Alongside Table 2, Figures 2 and 3 are overlaid DSC curves of certain compounds.

Table 2.

Melting point, glass transition temperatures, and enthalpies of fusion reported using a ramp rate of 10 °C/min in N₂ with DSC. Entropy of fusion was calculated from the enthalpy of fusion and melting point data. Uncertainties were calculated from the standard deviation of the mean with at least nine independent trials.

Compound	Tm (°C)		ΔHfus (kJ/mol)		ΔSfus (J/mol K)
Benzenesulfonate salts
1A*	−51.0
2A*	−47.8
3A	93.0	±0.1	25.5	±0.1	69.6	±0.1
4A	56.8	±0.1	15.9	±0.2	48.3	±0.7
5A	185.2	±0.7	32.1	±0.5	70.1	±1.1
6A*	62.8
m−chlorobenzenesulfonate salts
1B*	−48.1
2B*	−45.3
3B	101.7	±0.2	30.3	±0.1	80.8	±0.2
4B*	−50.2
5B	167.4	±0.4	34.8	±0.3	79.0	±0.5
6B*	51.1
PPN+ salts
7A	189.5	±1.2	31.2	±0.3	67.5	±0.6
7B	163.6	±0.4	24	±3	64	±6
7C	214.4	±0.6	46.8	±0.2	96.0	±0.6

^*Indicates glass transition temperature (T_g) values.

Figure 2.

Weak glass transition DSC curves of [C₂mim][BZS] 1A, [C₄mim][BZS] 2A, then [Ph₃P-POP][BZS] 6A (Left). Normal melting point behavior of [C₃mpyrr][BZS] 3A, [P₄₄₄₄][BZS] 4A, then [Ph₄P][BZS] 5A (Right).

Figure 3.

Notable phase behavior from PPN⁺ salts: [PPN][BZS] 7A, [PPN][Cl-BZS] 7B, [PPN][Me-BZS] 7C.

As may be gathered from Table 2, the new salts of BZS⁻ fall into two subcategories: those having distinct T_m and those for which only a T_g was observed. For purposes of the subsequent discussion, it must be stressed that the failure to observe a T_m for a given compound does not mean that the compound does not have one; rather, the organization of the ions into a crystal lattice may simply be kinetically frustrated, forcing the IL to become an amorphous solid. Considerable effort was made to crystallize the ‘glassy’ ILs through traditional benchtop methods as well as utilizing the DSC to thermally encourage crystallization; however, these efforts proved to be unsuccessful.

Somewhat surprising to us (in several cases) is that the T_m values of the BZS⁻ salts compare favorably with those of same-cation fluoroanion counterparts. For example, the BZS⁻ salt of the P₄₄₄₄⁺ cation melts at 55.6 °C; in turn, the salts of all three fluoroanions used as benchmarks – Tf₂N⁻, BF₄⁻, and PF₆⁻ – melt at higher temperatures (65 °C, 67 °C, and 219 °C, respectively, shown in Table 3). On the other hand, C₃mpyrr⁺ (3) cation with the BZS fluoroanion produces a significantly higher melting point compared to the 12 °C and 64 °C values found for its Tf₂N⁻ and BF₄⁻ salts, respectively. Indeed, it seems apparent that simple structural effects cannot alone account for the melting point increases or decreases. Therefore, we sought to understand the melting behavior via enthalpic and/or entropic driving forces.

Table 3.

Compilation of melting point data between BZS⁻, Cl-BZS⁻ and fluoroanion salts of the C₂mim⁺ (1), C₄mim⁺ (2), C₃mpyrr⁺ (3), P₄₄₄₄⁺ (4), and Ph₄P⁺ (5) cations. Data presented about the BZS⁻ and Cl-BZS⁻ anions is compared to the common fluoroanions traditionally used.

Cation	Tm (°C)
	BZS−	Cl-BZS−	Tf2N−	BF4−	PF6−
1	−51.0*	−48.1*	−1719	1521	68–7021
2	−47.8*	−45.3*	−4.922	−7123	10.3622
3	93.0	101.7	1224	6424	11324
4	56.8	−50.2*	6525	6724	21923
5	185.2	167.4	135.021	350.524	391–39321

^*Indicates glass transition temperature (T_g) values.

We now point out several particular comparisons between the T_m values of the BZS⁻ ILs and insights into their behavior that can be gleaned from their entropy and enthalpy of fusion values. Table 4 lists comparisons that were made when the cation structure was fixed with only the anion structure being altered; for example, 3B is C₃mpyrr⁺ with the meta-chlorinated benzenesulfonate compared to 3A which has no substituent groups on the phenyl ring, so a relationship between the melting point change can be explained by examining the driving forces caused by the addition of the halogen. We first note that in each case, the changes in enthalpy and entropy are very similar and in the same direction, thus having mutually negating impacts on melting point differences. Although we assign a driving force to denote whether the enthalpy or entropy was the dominate factor in the melting point change, we acknowledge that the differences are so small as to make generalizations difficult, even those resulting from relatively large melting point changes of ~25 °C. This is consistent with the complexities associated with developing structure–property relationships for species with complex molecular-level interactions.

Table 4.

Comparison of thermodynamic properties and driving forces between ion pairs that share a common cation.

Comparisons	ΔTm (°C)	ΔHfus Ratio	ΔSfus Ratio	Driving force
3B/3A	8.7	118.82%	116%	Enthalpic
5B/5A	−17.8	108.41%	113%	Entropic
7B/7A	−25.9	76.92%	80%	Enthalpic

In addition, in previous work, it was noted that an increase in the dipole moment of the ion generally led to a decrease in the T_m.²⁰ Structural changes of the benzene ring (e.g. incorporation of asymmetrical meta-Cl substituent) increases the dipole moments of the anion, leading to a clear pattern in T_m reduction. In fact, the electron withdrawing property of Cl group and position of this substitution modifies the orientation of the dipole moment, altering the intermolecular interactions that influence the phase transition behavior of ILs.

In the first example, 3A/3B, a T_m increase of 8.7 °C is driven by the slightly higher enthalpy of fusion of 3B compared to the enthalpy of fusion of 3A. Similar changes in the enthalpy and entropy are observed for the 5B/5A comparison, however, in this case the entropic change is slightly larger, resulting in a ~18 °C decrease in the melting point. Each of these changes, although subtle, can be understood by recalling that the change in the anion between the A and B variants is that in B, a chlorine has replaced a hydrogen on the benzene sulfonate. This change increases the polarity of the anion which could increase the molecular level attraction between the species, resulting in a larger enthalpy of fusion. We note that if the chlorine were in the ortho or para positions, the charge separation would likely be decreased through resonance. Additionally, the substitution of a chlorine for a hydrogen at the meta position increases the bulk and asymmetry of the anion, potentially frustrating packing and increasing the entropy of fusion. The smaller change in the enthalpy of fusion for cation 5 may be due to dominance of the π–π interaction which exist between this cation and either anion, allowing the entropic term to dominate, resulting in a melting point decrease. Interestingly, opposite trends in changes in the enthalpy and entropy of fusion are observed for cation 7, where both the enthalpy and entropy of fusion decrease (7B/7A) by ~20%. Again, very similar changes for the two properties and in this instance, the enthalpy decrease appears to slightly dominate resulting in a melting point decrease.

Because of the broad use of bistriflimide anion in our group’s work developing thermally robust ILs, thermodynamic comparison between it and the A is merited, with the n-tetraphenylphosphonium cation being the most widely studied cation for which the thermodynamic data is available (6A and [6][Tf₂N] would also be an interesting comparison, however the 6A only exhibits a glass transition). The relevant thermodynamic properties of n-tetraphenylphosphonium bistriflimide¹⁹ (called 5D here) are T_m = 135.0 °C, ΔH_fus = 32.2 kJ/mol, ΔS_fus = 79.2 J/mol K and comparisons between 5A/5D are shown in Table 5 (using the bistriflimide as the reference compound).

Table 5.

Comparison of thermodynamic properties and driving forces between n−tetraphenylphosphonium benzenesulfonate (5A) and n−tetraphenylphosphonium bistriflimide (5D).

Comparisons	ΔTm (°C)	ΔHfus Ratio	ΔSfus Ratio	Driving force
5A/5D	50.2	99.7%	88.5%	Entropic

Replacing the bistriflimide ion with the benzenesulfonate ion results in a significant increase in melting point (50 °C) that is driven entirely by a decrease in the entropy of fusion. Most likely, this is due to loss in degrees of freedom expected when changing the flat, stiff BZS⁻ system for bistriflimide, which is known to often switch between cisoid and transoid configurations, a phenomenon the importance of which was recently dissected in elegant detail by Toda et al.²⁶

Experimental

The detailed synthetic procedure and the DSC methods used in this work were described in the Supporting Information.

Conclusions

In conclusion, we assessed the aerobic thermal stability of fifteen salts of imidazolium, pyrrolidinium, and phosphonium cations paired with a benzenesulfonate, meta-chloro benzenesulfonate or para-tolylsufonate anions (Table 1). These experiments were carried out at temperatures of 200 °C, 250 °C, and 300 °C for 96-hour periods. Mass loss data, post-heating appearance and NMR spectroscopy results showed the following ordering of the relative thermal stability of BZS⁻ anions (A > B > C), with the salt [PNP][BZS] (7A) showing the highest stability without mass loss for 96 hours at 300 °C. Future work will focus on adding increased flexibility to the benzenesulfonate base structure to increase the liquid phase entropy, while maintaining functional groups which preserve thermal stability. Therefore, this work demonstrates unique insights into IL structure–property relationships that aligns with and expands upon our previous discoveries, providing a framework for rational design to enhance thermal stability of these functional organic materials.