SOLVATION OF FLUORIDE IONS. 3. A REVIEW OF FLUORIDE SOLVATION THERMODYNAMICS IN NONAQUEOUS AND MIXED SOLVENTS* G.T. Hefter

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1 SOLVATION OF FLUORIDE IONS. 3. A REVIEW OF FLUORIDE SOLVATION THERMODYNAMICS IN NONAQUEOUS AND MIXED SOLVENTS* G.T. Hefter School of Mathematical and Physical Sciences, Murdoch University, Murdoch WA 6150, Australia CONTENTS Page 0. Abstract Introduction Theory and Background The Thermodynamicsa of Ion-Solvation Single Ion Thermodynamic Quantities A Note on Standard States Thermodynamic Data for Fluoride Ion-Solvation in Nonaqueous Solvents Gibbs Free Energies of Transfer Enthalpies of Transfer Entropies of Transfer Other Thermodynamic Quantities A Warning Note on the Data Systematic Trends in Fluoride Solvation Gibbs Free Energies of Transfer Enthalpies of Transfer Entropies of Transfer Fluoride Ion Solvation in Mixed Solvents Gibbs Free Energies of Transfer Other Thermodynamic Quantities Conclusions Acknowledgements 219 References 220 * Ref. 48 is Part 2. This paper was presented in preliminary form at the 3rd International Symposium on Solubility Phenomena, University of Surrey, Guildford, U.K., September

2 Vol 10, Nos. 1-3,1989 Solvation of Fluoride Ions, Abstract Thermodynamic data for fluoride salts in non-aqueous and mixed solvents are collected and reviewed. Thermodynamic parameters for the solvation of the fluoride ion have been derived from these data via the tetraphenylarsonium (or phosphonium) tetraphenylborate assumption. Gibbs free energies of transfer of F~ from water to pure nonaqueous solvents, obtained mainly from solubility data, often show large discrepancies between independent determinations. A t G(F~) H20 _>. s is positive (unfavourable) for all solvents studied so far, reflecting the strong tendency for F~ to H-bond with water. However, A t G(F - ) H2 Q_> s for dipolar aprotic solvents is very much more positive than for mildly protic solvents such as the alcohols and formamide. As expected from the coordination model of ion-solvation, A t G(F~) Hj0 - >s values correlate reasonably well with semi-empirical measures of solvent acceptor (H-bond donor) strength. Surprisingly, they also appear to correlate with the bulk solvent dipole moment, although this may be fortuitous. Data for enthalpies and entropies are scarce but there is a marked difference between the values of A t S(F~) H2 o-+s a P rot i c and mildly protic solvents. In aqueous organic mixtures, Δ,0(Ρ~) Η2 ->Η values generally show a monotonic (positive) increase which becomes steep at high organic concentrations. The corresponding enthalpy and entropy contributions for Me0H/H 2 0, the only system measured so far, show a complicated dependence on solvent composition in comparison with A t G values, reflecting the well-known complexity of aqueous organic mixtures. It is concluded that considerable scope exists for the further study of the thermodynamics of fluoride ion transfer from water to both nonaqueous and mixed solvents. 1. Introduction The solvation of fluoride ions, unlike those of the other halides, has received only scant attention. For example, in the most comprehensive compilations of Gibbs free energies of transfer to date 11 4/ fluoride data are listed for only three solvents. The situation with respect to the other thermodynamic parameters is generally worse /5/. This is unfortunate because as a group the halides provide a series of ions of well defined size and electronic structure which makes them useful for testing theories of ion-solvation. The fluoride ion is especially interesting because its bonding characteristics are relatively straightforward with only 186

3 G.T. Hefter Reviews in Inorganic Chemistry electrostatic or σ-bonding present in most interactions /6/. To this end fluoride has been used, for example, to test simple electrostatic theories of complex formation in aqueous solution /7/. Fluoride is also of interest because it is iso-electronic with the hydroxide ion (OH~) and substitutes for it very readily in many of its compounds. For example, the replacement of OH~ by F~ in the mineral apatite is the basis for the use of dietary fluoride to minimise dental caries /6/. Because of its potential interest it seemed worthwhile to collect together all the available data on fluoride solvation in order to establish exactly what has been achieved so far, and what remains to be done. 2. Theory and Background 2.1. The Thermodynamics of Ion-Solvation cycle: The thermodynamics of ion-solvation can be illustrated by the following AslnX 0 MA (s) * Μ (solv) + A~ (solv) Δ 8θ1ν χθ -Δ, nx ω Μ (g) + A"(g) where X represents any thermodynamic quantity (G, H, S, etc.) and the subscripts sin, latt and solv refer to the processes of dissolution, lattice formation (from the gas phase) and solvation respectively. For ionic substances the changes accompanying lattice formation and ion-solvation from the gas phase are generally very large, e.g. about 1 MJ mol" 1 for X = G /8/. Consequently, in discussing ion-solvation in different solvents the absolute values of the solvation processes are of less interest than the (relatively small) differences between solvents. These differences are usually expressed as transfer functions /9/: V e (0siH.s2 =[x e (i)] s2 - [X*(0] S1 (1) corresponding to the transfer of one mole of the species i under standard state conditions from solvent SI to solvent S2, i.e. 187

4 VoL 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. M + (S1) + A"(S1) M + (S2) + A~(S2) Q β Thus A t X (i) is a direct measure of the difference between X (i) in the two solvents. The choice of reference solvent SI is arbitrary but is usually H 2 0 /9/. The standard state for electrolytes is normally taken as infinite dilution in each solvent (but see below) Single Ion Thermodynamic Quantities Single ion thermodynamic quantities are not defined within the framework of thermodynamics and can only be derived from observable (whole salt) quantities by use of an extra-thermodynamic assumption /4, 9/. A great variety of these assumptions have been proposed. It is outside the scope of this review to consider the relative merits of these proposals: a task which has been performed adequately elsewhere /2,4,9/. It is widely accepted that the so-called reference electrolyte approach /9/ is the most reliable assumption currently available although it can be noted that there is good general agreement with other popular methods, especially if water is not used as the reference solvent /10/. The reference electrolyte approach assumes that the thermodynamic parameter of a suitable salt can be equally divided between its constituent cation and anion. Tetraphenylarsonium tetraphenylborate (Ph 4 AsBPh4, TATB) or its phosphorous analogue (Ph 4 PBPh 4, TPTB) are generally used for this purpose. The assumption may therefore be written A t X ö (Ph4AsBPh4) = 2A ( X ö (Ph 4 As + ) = 2A t X ö (BPh 4 ) (2) This assumption, or the virtually identical TPTB assumption /ll/, will be used to derive the thermodynamic parameters of fluoride ion solvation in this study. Typically, the fluoride data are available in the form of simple whole salt values. Thus A t X Ö (F") = A t X ö (MF) - A t X 0 (M + ) (3) A where A t X (M + ) are obtained from the literature and are based on the TATB or TPTB assumptions. 188

5 G.T.Hefter Reviews in Inorganic Chemistry 2.3. A Note on Standard States Strictly speaking, as implied in equations (1) (3), when comparing the thermodynamic parameters of ion-solvation it is necessary to use standard state values. For an electrolyte the standard state is normally taken to be infinite dilution in the solvents of interest /9/. Such standard state values are usually obtained by correction or extrapolation to infinite dilution, by an appropriate theoretical or semi-empirical equation (e.g. the extended Debye-Huckel equation), of data obtained at finite concentrations. In this way, all ion-ion interactions are removed and the (standard state) differences in the thermodynamic parameter can be ascribed solely to differences in ion-solvent interactions. Unfortunately, few fluoride electrolyte data obtained to date are sufficiently reliable to justify such corrections. Thus, in this review the data have not been corrected to standard state conditions. Fortunately, the solubilities of most fluoride salts are sufficiently low that these corrections should in general be small (< ca. 5 kj mol" 1 ) and probably within the limits of the validity of the TATB/TPTB assumption. A more significant problem than activity coefficient effects is the degree of association (ion-pairing, complexation) of fluoride salts. Fluoride is one of the smallest of all anions and because of its relatively high charge/radius ratio, its salts are expected to associate (or ion-pair) /7/. Even the alkali metal fluorides, which are only weakly ion-paired in water /12/, are likely to be appreciably associated in solvents of lower dielectric constant /13, 14/. Since the association constants of fluoride salts in nonaqueous and mixed solvents are in general unknown /15, 16/, it is not possible to correct for such effects. This problem can, of course, also be eliminated by extrapolation to infinite dilution, where all electrolytes are fully dissociated, but the available data preclude this approach. However, because ion-pairing is strongly dependent on the cation it should be possible to detect any serious errors from this source by comparing data from different fluoride salts. For practical purposes: availability of data, degree of ion-pairing, etc., the present review has been restricted to the alkali metal fluorides. 3. Thermodynamic Data for Fluoride Ion-Solvation in Nonaqueous Solvents 3.1. Gibbs Free Energies of Transfer The Gibbs free energy of transfer of a species i, from solvent SI to solvent S2, A t G(i) sl^s2, is a measure of the overall difference of chemical reacti- 189

6 VoL 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. vity of i in SI and S2. AtG(i)sl_>S2 can in principle be obtained in a variety of ways but emf and solubility measurements are the most commonly used methods. For the fluoride ion, emf measurements have been limited by the absence of a well-characterised conventional fluoride-responsive electrode "of the second kind" /17/. Silver(I) fluoride is generally soluble in most solvents and thus an Ag AgX(s) IX" (soln.) electrode, which has been so useful in measuring free energies of transfer for the other halide ions /18, 19/, is not feasible for fluoride. In principle, the Pb(Hg) lpbf2(s) IF" (soln.) electrode developed by Ivett and De Vries /20/ could be used in spite of the relatively high solubility of PbF2. However, this electrode has been employed only rarely /21, 22/ for fluoride ion solvation studies. The single crystal lanthanum fluoride membrane (ion-selective) electrode /23/, which has become the standard method for the measurement of fluoride activities and concentrations in aqueous solution /24/, also appears to work well in many nonaqueous solvents /16, 25/ and has been used occasionally /26/ to measure AtG(MF). Unfortunately, the commercially available electrodes in which an epoxy resin is used to seal the LaF3 membrane into a plastic body are not suitable for many organic solvents and alternatives must be Not surprising therefore the overwhelming proportion of information on the free energies of transfer of the fluoride ion is available in the form of fluoride salt solubilities. For a solid salt MA in equilibrium with its saturated solution at constant temperature and pressure: MA (s) ^ M + (solv) + A~ (solv) (4) β the standard state solubility product Ks is given by Kj (MA) = am+ aa- (5) where aj is the activity of ion i in the saturated solution. The standard Gibbs free energy of solution A^G^ of MA is in turn related to the solubility product AshG 0 (MA) = -RTlnKf(MA) (6) i.e. As]nG 0 (MA) = pkf at 298 Κ (7) 190

7 G.T.Hefter Reviews in Inorganic Chemistry From equation (1) and the thermodynamic cycle given earlier it follows that: A t G ö (MA) sl -> s2 = A sh G ö (MA) s2 - A sin G 0 (MA) sl (8) A t G Ö (MA) S1^S2 = ΔpK? at 298 Κ (9) where ΔρΚ? =ρκ θ 52-ρΚ θ 81 (10) In other words, the Gibbs free energy of transfer of MA from solvent SI to solvent S2 can be directly deduced from its solubility in the two solvents, noting that the concentration scale must be specified /l, 9/. Unfortunately, few fluoride salt data have been collected or assessed systematically /28, 29/ even in the IUPAC-sponsored Solubility Data Series /30/. Because solubility data are widely scattered throughout the literature, frequently in obscure places, it has not been possible to perform a complete literature survey. For the purposes of this review the requisite data have been taken mainly from secondary sources /31, 32/ although the primary publications have been consulted wherever possible. As noted above, the solubility data have not been corrected to infinite dilution as, given the present state of the measurements, such corrections are probably smaller than the likely experimental errors. This is, in effect, an assumption that A sin G 0 (MF) = -RT In Ks (MF) where K S (MF) = expand Cmf is the observed solubility of MF (in mol/l s/n). Again for reasons given above, the present survey has been restricted to the alkali metal fluorides. The available solubility product data are summarised in Table 1. It is apparent from Table 1 that for many solvents, independent determinations of fluoride salt solubilities are in reasonable agreement [note that ±1 in pk s corresponds to about ±5 kj mol -1 in A^G at 298 K, see equation (7)]. However, for some systems the data are in serious disagreement, e.g. for NaF in AN the reported values differ by more than 10 orders of magnitude! Whilst traces of water will always create difficulties for the reliable determination of fluoride solubilities in nonaqueous solvents (see below) this is unlikely to result in such huge differences. Other experimental errors must clearly be present: a reminder that fluoride salts are often prob- 191

8 VoL 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. στ es on (Ν Γ- CO (Ν Γ-» I CO CO Ό νθ r^» ON ~ CS NO TT 1 Λ - >.. CO»o»o Ό ON CO CO CS «Λ ON -tf r-»o i-h v>»o NO Νώ 5r <Ν CS ON CS ON "t <o ν e S C/J Μ 3 υ 3 sr Β Ζ J5 QO V 1 3 Μ,0 Ω >. 5 3 J3 3 (Λ. tu 5 W Μ α. 3 α. Μ >. jo 3 Ζ "θ CO ON CS co Tf ' ON 00 ρ Γ ^Η ~ CO ON ON CS CO CS ON CS r-. Ö ρ ON 1 ** ^ CO CO CS 00 σ- CN (Ν CO ö Γ- CS es * ρ CO JO ON CS ON CS ON ρ od 00 CO CO CS CO CO >» - ON CS Ov 00 CO CO < Ov 1 s 2 8 Λ (Ν C-. Ό es?ϊ CO <n 00 CS _ o> CS CS CO ON CS CO 00 CO CO CO CO CS CO ON o\ On oo ON P- NO On Λ ι I 1 00 od 1 vi ON co ω -4 β < Η υ < < ζ Χ VC Χ I «^r- C5 23 2: - υ υ u Q Q Q 5 S χ 2 α α ω ω 192

9 G.T.Hefter Reviews in Inorganic Chemistry C I "o crt 00 <N JN (S Ö 5 t-; a o VO vo... OS CS Os CO CO SO c^ so <N Vi so cö 1 ri 1 00 σ\ σί c 3 3 x> 5 κ S υ C vo 5, < α υ S 5 I -S.2 β 3 8 a. ε < j CO CO Ζ Κ s I U. X Η.5 a i c 45 4J S 5 1 ρ. η * 5 w a 193

10 Vol 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. + > υ α < VC ΪΝ I I I I I m I Γ- < fs vc * S ι ι σ\ I t rn c σ\ cn 00 NC I I»λ vc m S «ι ι ι ι ι S ι + 06 < vc I I I I I "Τ I VC is <2. β! < O OV Ov I I I r- I I I I r* *r> I σν I m I I I I I U U <Γ < < + «Λ <N 2 Τ U ι ~ < "Γ 2 < ν»n vi c» Ov m ό <o * m 7 to ci r- I «<N m «η β ' ri vc»r» *rt «Λ I so «v r^ oi oo so ^ < rs + 3 < <»o? 2 Ü <3 ν vc Ö J5 Ö < ρ- (Ν pi IN VC VC w I m»λ X 3 CO Ζ I < «W u. co χ S S Σ u < Q D Q ω ω u. x 194

11 G. Τ. Hefter Reviews in Inorganic Chemistry c ε Ι 2 3 <0 = C w a 2 s 8 - ε «8* < a c t> S 3 «b. ^.5 " G - υ.i ~ M> cl Μ V. Β m Ι c <1-5 / ν - υ «w' <N Jo a 3L S ^ 3 5 S 6?r S Ό Μ cn *Ή 'ζ 00 Μ c χ: CL, «C o. cj I 5 I Η SS Μ a m «ts Η CS S 1 oo η c/> ι I ül» h a r ζ t> φ 00 Λ CN Ό ~ - s fe J 8 S > «1 I κ CS <J> V <22 < if d ε d ε - 9, ^ c Ε c g c ^ g > a> n\ iv> ω U K rs l) w D Μ Λ " " m u ö S O J S Ä Ε 8 D <3 ä i c i Ä W U Q i W ß i t t : aj Χ) υ Ό D cuj.h 195

12 VoL 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. lematic experimentally. Sodium fluoride in AN is an extreme example but quite a few of the other systems in Table 1 show unsatisfactory agreement. Clearly, much experimental work remains to be done on fluoride solubility determinations in nonaqueous solvents. Table 2 lists the thermodynamic quantities A^GCMF) and A t G(MF) Hj Q_>. s which can be derived from the solubility data in Table 1. Also included in Table 2 are A t G(M + ) H2 o_». s data taken from the literature /4/ and based wherever possible on the TATB assumption. These values are used to derive A t G(F~)H 2 o->-s via equation (3). Note that although A t G(M + )h 2 o-^s data are unavailable for some ions in some solvents, the corresponding values of A s j n G(MF) have been included in Table 2 for completeness. The single ion free energies of transfer for the fluoride ion from water to the various solvents for which both A t G(MF) Hi0^. s and A t G(M + ) H20 ->-s exist are summarised in Table 3. Discussion of these results will be deferred until the data for the other thermodynamic quantities have been presented Enthalpies of Transfer Enthalpies of transfer (0si-*-s2 are a reflection of differences in ion-solvent bond strengths. Transfer enthalpies are best obtained directly from calorimetric measurements, for example by measurement of heats of the solubility, since, assuming a negligible heat of dilution of MA /45, 46/, A t H fl (MA) sl^s2 = Δ^ΧίΜΑ)^ - A sln H 0 (MA) sl (11) or an appropriate, thermodynamically equivalent series of calorimetric measurements. Alternatively, Δ { Η can be obtained from the temperature dependence of the solubility, since /45, 46/ - -Δ Λ Η»/Τ' (12) although it is generally accepted that ΔΗ values obtained in this way often have large uncertainties. Unlike the other halides, relatively few direct calorimetric studies of heats of solution of fluoride salts have been made. For the purposes of the present survey, calorimetrically determined A t H values have been used whenever available but, on the basis that even approximate values are better 196

13 G.T. Hefter Reviews in Inorganic Chemistry TABLE 3. Summary of A t G(F~) H2 Q^. s Obtained from Solubility Data at Κ (TATB Assumption; mol L" 1 scale) Solvent A t G(F~) a /kj mol" 1 LiF NaF KF RbF CsF "Best" value* 3 AC ? AN BuOH - (14.8) (15) D 2 O d DMF c DMSO EG e EtOH e FA f MeOH (29.7) PC e ' h 1 - PrOH THF e a I. c d e Derived from equation (3) using relevant data from Table 2. Values in parentheses are probably unreliable. See text for definition. Average of values derived from KF, RbF and CsF; see text. Based on A t G(Na + ) H2 Q->D 2 o = 0 ( refe rence /44/). Average of all available data; see text. ^ Value of 24.7 kj mol -1 is given in reference /2/. g Average of LiF and NaF data only. Literature values are: 10.4 kj mol" 1 (emf data, cf. Table 2 /26/) and 18.3 kj mol" 1 (solubility data /33/). ^ Value of 54.8 kj mol" 1 given in reference /4/. than none, solubility data have been used wherever available to cover as many solvents as possible. The data obtained from the literature in these ways for the alkali metal fluorides are summarised in Table 4. It can be seen that there are somewhat fewer values available than for the Gibbs free energies of transfer. This is somewhat surprising because unlike A t G(MF) (see above), the measure- 197

14 VoL 10, Nos. 1-3,1989 Solvation of Fluoride Ions, T. = < X <3 I I I.β ös < X "7 1 1 "Τ t. ι ι ι ι χ c I I pi I I I I Χ <Γ U. U < χ + α Ζ χ < τ I I ν ft I I r- ^ Τ "Τ ι- τ i Ν CMΪΝ Ν Ν I I I <N o> VC -Η I ΓΙ I I I ι co Τ ι ι b. «Ζ X < <o I I I I I X S «Λ VC < 1 I I + J χ < U. X 3 < χ S Ζ < X 3 VC I Γ4 - D 1 I I (N. C4 I i I I ri I I VC ι S ι ι <N I I Μ I I X D X ^ Ö y < u t as s i x s U U. 2 S Ζ Ζ Ζ 198

15 G.T.Hefter Reviews in Inorganic Chemistry g - Γ 4 CD S S 2 S 5 S 3 s.s «Si (Λ 1) β ΰ * Ö - ^ 6.-ι Γ- w» <7 1 c 2 «c -β 3 ^ c d S 2 eo υ 3 κ C CT ε *>.2 «,ε ι ρ. - δ «- / Ν - 3 ε ^ Έ U«δ5 m s - -S Ζ ^ 6 CD " Ö Η ϊ, Ε Η "β Ε - g.2 α JS Τ3 C, Λ L >, «Ü "s >. 2 ε 8 ^ 2 ~ I β S «Β?! fl cd <1 ^ «^.a. Λ II 5 Ä V ph S ε ϊ ϊ 8 3 Έ Γ Β S. C Μ C 5 Ζ> 2 " π ^» Μ) 8» ε» «ί Μ Η 3 Ü C 3 Μ ^. ρ- 8 * -Η 2 S q υ 1 < 2 'S t> 1 1 q - Ό.5 < α C. ε 3 β Η < Η Β Ό α «t; Ε ~ ΐ 3 ^ σ ί Μ ε u S jt γ* 5Λ Λ.3 5 ΐ 5 δ S a SS "3 σ<ε> 8 c <1 > «ί Κ Jf < 3 - a ~ '" I -5 Β 3 " < ε - ρ -9 J3 3 a υ η α os r b 2 α> 3? < α 199

16 VoL 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. ment of A t H(MF) by calorimetry does not present any special difficulties. Also included in Table 4 are the A t H(M + ) values necessary to calculate A t H(F~) via equation (3). These values have been taken from the literature /47/ and are based wherever possible on the TATB assumption. Again, as for Table 2, data for Δ^ H(MF) have been included even if the corresponding values of Δ { Η(Μ + )h 2 o-+s are not av ailable- Table 5 summarises the single ion enthalpies of transfer of the fluoride ion for solvents for which both A t H(MF) and A t H(M + ) are available (Table 4). Most of the values for A t H(F~) Hi0 ->s should be regarded with extreme caution. As for A t G(F~), the A t H(F~) values derived from different salts TABLE 5. Summary of A t H(F )h 2 0-»-S at Solvent A t H(F") H2 o _>.g/kj mol LiF NaF KF RbF CsF "Best" value Literature values' 3 AC _ (17.9) (41.1) (46.2) (57.3) (50) AN (30) - D , -2.6 DMF EtOH FA ,21.3 d MeOH , 13.8 d, 12.3 e NB (-4) - NMA - - (42) - - (40) -42 (sic) NMF (62) - (60) 45, 28.5 a k c d e Values in parentheses are probably unreliable. Data available for NH3 at 240 Κ /47/. Marcus /47/ also gives a value of 30 kj mol -1 for DMSO but note the cited reference should be E.M. Amett and L.E. Small, J. Amer. Chem. Soc., 99, 808 (1977), (Y. Marcus, personal communication), Data from reference /47/ unless otherwise indicated. Reference /44/. Reference /2/. Reference /48/. 200

17 G.T. Hefter Reviews in Inorganic Chemistry are often quite different (formamide is an exception). Furthermore, A t H(F~) values derived from the temperature dependence of uncertain solubilities, in the absence of confirmatory studies, must be considered quite doubtful. Similarly, some of the A t H(F~) values are based on extrapolation procedures which are often unreliable cf. the TATB assumption /49/ Entropies of Transfer Although methods are available for the direct measurement of entropies of transfer, e.g. via non-isothermal emf measurements /50/, virtually all entropies of transfer obtained to date have been via the combination of free energy and enthalpy data A t S = (A t H - A t G)/T (13) The available data are summarised in Table 6. Again the values are very limited and should be regarded with extreme caution (see below) Other Thermodynamic Quantities Virtually no experimental data relating either to fluoride salts or the fluoride ion exist for other thermodynamic quantities such as partial molar volumes /51/, heat capacities /52/, etc., in either pure or mixed nonaqueous solvents. Some partial molar heat capacities can be obtained from the second derivative of the solubility-temperature curve /45, 46/ but the likely errors make their utility dubious A Warning Note on the Data It is obvious from Table 6 that A t S(F~) H2 o-»s va l ues derived here from the "best" A t G and A t H data of Tables 3 and 5 often differ from those reported in the literature even though based on (mainly) the same experimental data. This results from the use of different values of A s inx(mf), X = G or H, different values of A t X(M + ), and the attempt, embodied in Tables 1 5, to impartially systematise all the available data. Whilst this leads to relatively small differences between the present and literature values of A t G and A t H, these differences can accumulate into A t S. Consider for example the transfer of fluoride from water to methanol. The present author has reported A t X(F") to be 18.3 klmo!" 1, 12.3kJmor' and 20 J K" 1 201

18 VoL 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. mol" 1 for X = G, Η and S respectively /33, 48/. The corresponding values from Tables 3, 5 and 6 are 17 kj mol" 1, 20 kj mol -1 and +10 J K" 1 mol" 1. Thus, whilst the differences between A t G and Δ ( Η are relatively small, that for A t S is rather larger, changing from being moderately unfavourable /48/ to moderately favourable (Table 6). TABLE 6. Summary of A t S(F"") Hj0 _>. s at Κ Solvent A t S ( F ">H 2 0+s/ JK_1 This work 3 mol" 1 Literature' 3 AC AN D DCE DMF EtOH 10 - FA -4-4, -11 MeOH 10-8, -20 d a ^ c d Calculated via equation (13) using the "best" values from Tables 3 and 5. From Marcus /47/ unless otherwise indicated. Reference /2/. Reference /33/; see also text. It may well be that the original data are closer to the "true" values than those derived here. Certainly, use of the term "best" in Tables 3, 5 and 6 is only intended to imply "best" within the context of the available data and the present assumptions and not in any absolute sense. It is clear from the variations between different studies and different salts in Tables 1 and 3 that much of the available data merits healthy scepticism. Indeed part of the purpose of this review is to draw attention to the inadequacies of the data in order to encourage further investigations. 202

19 G.T.Hefter Reviews in Inorganic Chemistry 4. Systematic Trends in Fluoride Solvation 4.1. Gibbs Free Energies of Transfer It can be seen from the data in Table 3 that for some solvents there are marked differences among the values of A t G(F~) obtained from the different fluoride salts. Whilst the experimental uncertainties should not be overlooked, it is most likely that any systematic increases in A t G(F~) in going from LiF to CsF are due to decreasing ion-association. This is consistent with the tendency of the alkali metal fluorides and other alkali metal salts to associate /12, 53/. Thus, in deriving the "best" values in Table 3, where a systematic variation is present, the data for the heavier alkali metal fluorides have been preferred (see Footnotes to Table 3). Even so, there are some surprising variations in the data in Table 3. For example, it is puzzling that THF (e = 7.6) should show a smaller cation effect than AN (e = 37.5) given that the solvent dielectric constant is a major factor determining the degree of ion-pairing /13/. It must also be remembered for solvents where salt solubilities are high (e.g., for the heavier alkali metal fluorides such as CsF) that activity coefficients need to be corrected for, not always an easy matter for high electrolyte concentrations in nonaqueous solvents /9/. It is apparent from Table 3 that, for all solvents studied so far, A t G(F~) H2 o->-s i s strongly positive, i.e. highly unfavourable. This is a reflection of the strong stabilisation of the fluoride ion aqueous solution by hydrogen bonding. In this connection it would be interesting to have data for A t G(F~) to acidic solvents which are stronger H-bond donors than H 2 0. Figure 1 summarises graphically the single ion Gibbs free energies of transfer for all the halide ions for the solvents for which fluoride data are available (Table 3). Values for the other halides have been taken from the review by Marcus /4/ and are based wherever possible on the TATB assumption. Figure 1 shows that the transfer of the halide ions from water to both dipolar aprotic solvents and mildly protic solvents like the alcohols is always positive and increasingly so in going from iodide to fluoride. This is basically a reflection of the extent of stabilisation of the anions by hydrogen bonding. Water is a very strong Η-donor compared with most solvents 191 and the tendency of the halides to act as H-bond acceptors increases markedly up the group, i.e. with increasing charge/radius (Z/r) ratio. In- 203

20 Vol 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. Fig. 1. Selected Gibbs free energies of transfer for the halide ions at 298 K. Fluoride data from Table 3 ("best" values), other halides from ref. /4/. terestingly, although the data are too few to draw definitive conclusions, there also appears to be a difference between mildly protic (ROH, FA) and aprotic (AC, AN, etc.) solvents with A t G(X~)H 2 0-+S in the latter showing a much stronger dependence on the nature of the halide ion (Figure 1) Figure 1 is also useful for assessing the reasonableness of the values for A t G(F~)H 2 o- > S which can be derived from the solubility data included in Table 1. For example, a value of A t G(F~)H 2 o->-an = 27.8 (see also ref. 1331), which is less positive than A t G(CF)H 2 o-»-an would appear to be unlikely. Similarly, on the basis of their H-bond (electron acceptor) properties /9, 13/ it would be expected that A t G(F~)H 2 o->-s would follow the trend MeOH < EtOH < PrOH < 1 - BuOH (most positive). If the values for MeOH and EtOH (Table 3) are correct then the values for PrOH and especially BuOH appear to be too low. These uncertainties need to be resolved by further careful work. 204

21 G.T.Hefter Reviews in Inorganic Chemistry Ion solvation is a complex process and no comprehensive theory is available which can account quantitatively for all the observed phenomena. The two common models of ion-solvation are the Born approach in which ions are considered to interact electrostatically with solvent dipoles, and the donor-acceptor or coordination model in which ions are considered to interact covalently (coordinatively) with electron donor and acceptor sites on the solvent molecules. The simple Born approach /9/ predicts that A t G(ion) H20^s = - NZ 2 e 2^1 ^ ( e ^ - e' 1 ) (14) where Ν is Avogadro's number, Ze is the charge, and the radius of the ion, and e is the bulk dielectric constant. From equation (14) it follows that A t G(ion) should be proportional to. Figure 2 plots the available data for A t G(F~) H20^. s against the reciprocal of the solvent dielectric constant. Clearly little correlation exists. This is not surprising since it is well known that the simple Born model is an inadequate description of ion-solvation in nonaqueous solvents /9/. This is interesting in relation to the present data because, as noted in the Introduction, fluoride is an ion whose interactions with other species are more likely to be dominated by electrostatic forces than any other. For example, a simple electrostatic model provides a good account of fluoride complexation in aqueous solution /7/. Nevertheless, it has been recognized /7, 8/, such theories work not because they are physically realistic but rather because of a compensation of errors. It is certainly true that attempts to make the Born equation more realistic, e.g. by allowing for the change in ionic size on transfer, have not been particularly successful. The donor-acceptor approach /13/ at its present stage of development provides a semi-empirical rather than quantitative account of ion-solvation. Anion solvation is visualised in terms of the ability of solvent molecules to act as electron density acceptors F:~ * S and thus predicts that A t G(F~) H20 _> s should correlate with solvent acceptor strength. A variety of semi-empirical measures of solvent acceptor ability have been proposed over the years /9, 13/ and many of them correlate well with each other. Figure 3 plots A t G(F~) H2 o-»s a g a i nst the Dimroth-Reichardt E T parameter /9/ which was chosen because of the availability of the values for the solvents of interest. A similar relationship 205

22 VoL 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. Fig. 2. Gibbs free energies of transfer of fluoride (Table 3, "best" values) as a function of the reciprocal of solvent dielectric constant 19/. exists between A t G(F~) H2 Q^. s and the Gutmann-Mayer acceptor numbers /131 for these solvents. Plots of A t G(F~) H2 o-»s a gainst a variety of other bulk solvent properties do not yield any significant correlations /5/ but surprisingly, there seems to be an apparent, albeit approximate, correlation between A t G(F~) H q-^s and the bulk solvent dipole moment (Figure 4). Although at least partial correlations are known between solvent dipole moment and H-bond formation /54/ such a relationship is unexpected in terms of ion-solvation /9/. (It may be noted that there is virtually no correlation between acceptor strength and dipole moment). It is, furthermore, intuitively surprising that A t G(F") H20 _ >s should become more positive, i.e. less favourable, with increasing dipole moment given the "electrostatic" character of the fluoride ion. However, it must be remembered that Figure 4 is based on limited and inadequately substantiated data. It will be interesting to see if future measurements confirm this possibly fortuitous relationship. 206

23 G.T.Hefter Reviews in Inorganic Chemistry <E t )/k cal mol"' Fig. 3. Gibbs free energies of transfer of fluoride (Table 3, "best" values) as a function of solvent acceptor strength parameter Ε χ /9/ Enthalpies of Transfer Unlike the Gibbs free energies (Table 3), the enthalpies of transfer (Table 5) for the fluoride ion do not show any systematic dependence on the particular fluoride salt (i.e. on the cation) for the few solvents for which data are available. This is probably a reflection of the fact that calorimetric data are routinely extrapolated to infinite dilution and also that ΔΗ for ionpairing reactions in general /55/ and fluoride in particular /7/ are small [i.e., ΔΗ (MF ) as 0]. The available enthalpies of transfer of the fluoride ion from water to nonaqueous solvents (Table 5) are summarised along with the corresponding values for the other halides in Figure 5. As for the Gibbs free energies (Fig- 207

24 Vol 10, Nos. 1-3,1989 Solvation of Fluoride Ions, DIPOLE MOMENT (D) Fig. 4. Possible correlation between Gibbs free energy of transfer of fluoride and solvent dipole moment. ure 1) the enthalpies of transfer become increasingly positive (endothermic, unfavourable) in going from iodide to fluoride, consistent with their increased Η-bonding (electron density donation) abilities. However, unlike A t G, there does not appear, on the basis of the limited data available, any difference between aprotic (AN, DMF etc.) and mildly protic (ROH, FA) solvents (Figure 5). In general, A t H(F~) H20 _». s is less positive than A t G(F~) H2 o-*s indicating that entropies play a significant role in determining A t G (see below). Figure 5 can, of course, be used to estimate A t H(F~) Hj0 -»-s f rom data for the other halides. For example, a value of ~40kJmol~ 1 would be predicted for A t H(F~)n o-»-pc However, great care should be exercised in such extrapolations; a point which has been well made, albeit in a different context, by Bhattacharya et al. /21/. 208

25 G.T. Hefter Reviews in Inorganic Chemistry F" CI" Br" Γ ANION Fig. 5. Summary of enthalpies of transfer of the halide ions at 298 K. Fluoride data from Table 5, other halides from ref. /47/ based on TATB assumption Entropies of Transfer Figure 6 plots the entropies of transfer of the halide ions from water to the limited range of nonaqueous solvents for which fluoride data are available (Table 6). As for the Gibbs free energies (Figure 1) and enthalpies (Figure 5), the entropies of transfer for the fluoride ion are in general unfavourable. As for A t G, there appears to be a significant difference between A t S(X) H2 Q_». s for mildly protic solvents (ROH, FA) which, remembering the sign, show the general order A t S(F~) > (CI") > (Br~) > (I~) whereas for the aprotics (AN, DMF), A t S(F") «(Cl~) > (Br") > (Γ). It is also noteworthy that for all the halide ions A t S(X") H20 -+s is ' ess favourable (more negative) for the dipolar aprotic solvents than for the mildly protic solvents. Bearing in mind the limitations of the data it is interesting to speculate on these differences. A simplistic representation of the solvation of an ion 209

26 Vol. 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. ANION Fig. 6. Summary of entropies of transfer of the halide ions at 298 K. Fluoride data from Table 6, other halides from ref. /47/, TATB assumption. in notional dipolar aprotic and protic solvents is given in Figure 7 (for more sophisticated models see for example Refs. /2/ and 191). It follows from this idealised model that the entropy of transfer from water to an aprotic solvent would in general be large and negative. The magnitude and sign of A t S(X~) from H 2 0 to another protic solvent will, on the other hand, depend on the relative structuredness (degree of disorder) of the solvated ion in each solvent and of the solvents themselves and is therefore likely to be only mildly positive or negative. 210

27 G. Τ. Hefter Reviews in Inorganic Chemistry O o O A solv s<< 0 ION APROTIC UNSTRUCTURED SOLVENT w SOLVENT DIPOLES ORIENTED ABOUT ΙΝ (a) Large decrease In entropy when ion is solvated "or; ex"'. w. : α., α Μ ' 0 0 PROTIC STRUCTURED SOLVENT (b) Relatively small change In entropy when Ion Is solvated Fig. 7. The effect of solvent structuredness on the entropy change during ion solvation. The general, but relatively small, trend of A t S(Cl") > (Br") > (I") observed in all solvents (Figure 6) is possibly a size effect: an increasingly unfavourable entropy as a result of steric interactions or "solvophobic" solvation (cf. hydrophobic hydration /9/) as the ion-size increases. Table 7 summarises the data for those solvents for which all three thermodynamic transfer quantities are available. In general it has been considered /2, 561 that the enthalpy contribution to the overall solvation of ions far outweighs the entropy effect. However, the present data (Table 7) are too limited and uncertain to decide definitely if this is true for fluoride ion transfer. There is a tantalising suggestion in Table 7 and in a comparison of Figures 1 and 6 that it may not be, but clarification on this point will have to await further measurements. 211

28 Vol 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. values at K in TABLE 7. Summary of A t X(F") H Q -»-S kj/mol; TATB assumption Solvent A t G a A t H b A t S c ' d 298 A t S c AC AN DMF EtOH FA MeOH a "Best" values from Table 3. k "Best" values from Table 5. c d Table 6, present values. Units: J K" 1 mol" 1 5. Fluoride Ion Solvation in Mixed Solvents 5.1. Gibbs Free Energies of Transfer Whilst the thermodynamics of ion-solvation in pure solvents are of considerable interest /9, 56/, mixed solvents are more attractive for many practical applications /9, 19, 57, 58/. A typical example is the use of mixed organic solvents to optimise the performance of lithium batteries /59/. Furthermore, the study of thermodynamic properties as a function of solvent composition provides a greater insight into ion-solvent interactions than is available from pure solvent data alone /3, 9/. The only data available to date on fluoride ion solvation in mixed solvents refer to aqueous-organic mixtures. The published values for the Gibbs free energies of transfer of fluoride from water to aqueous-organic mixtures are summarised in Table 8. These data were derived from the authors' original values by appropriate numerical and graphical conversions. Note therefore that the values for the pure organic solvents may differ somewhat from the values in Table 3 but have been included in Table 8 for comparison. Very few studies have covered the entire composition range and most of the data are restricted to the water-rich region (Table 8). Many have been 212

29 G.T.Hefter Reviews in Inorganic Chemistry ν) m 3 Γ«Ι r; m SO 00 Γ"" ι f-^ ι Ζ S Ό > c «I to 8 * c I ^ s m co V ι m 8 c tu /-n + χ- Ι «Ν I I b C 1 6 υ 00 < I I I I I I <ν I I I I «I I I I υ O R <*ΐ, Γ--' 00 «I I Μ «Ν VO 00 «Ν»-< I \ ι^ ρ ««ρ «ν* θ\ «CO CO CO ON TT (Ν ei \ö η ό \ <η t «η «-«CS νό CO w> sc Ό Ό - Μ tju s s 5 5 Ω Ω Ω Ι 00 C^ Tf CS 00 co σ> on ρ es r^ «-ί --ί \ «««W) CS CS CO CS «r^ o\ SO CO CO to r*^ t^ oo vq ^t co *τ -t' vo ^t es es' es' oo no in o\ «o «o «n co cn -*r ri η «η η»o es co r-- r- r- es es O o α «- w ö s a rt W Ü c ~ 8 <1 Β ι 0 1 u c S ε '55 «Β S ü α si I 4-> > c 0 o. * ~ < & * Λ /-s ^ c U ε I, JO I Γ 8 E t S o«> S Ä «~ < c I s 2 JS a &. Ε «s 5 3 'S Η Μ Β g -Ό α 1 * * Ό & 1 S s α 3.Ε kl G ι! C s«ί ι 2 1 S S 'S Μ Μ <4 II. Ε t S 5 J2 ~ - a Μ s J2 3 ha Ε ώ Ό ι Λ cd S β Ό S S ί? «w S > u Ό υ 213

30 VoL 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. obtained by emf rather than solubility measurements. Where comparison is possible (i.e. MeOH), data from independent studies are in excellent agreement especially in the water-rich region. It should be noted that Table 8, unlike Tables 1 and 4, does not attempt to be exhaustive. Some information on fluoride salt solubilities (mainly NaF and KF) available /31/ for aqueous/organic, chiefly alcohol, mixtures have been excluded. This is because the solubilities are in general high, at least in the water-rich regions, and are likely to differ significantly from the standard state values. The interested reader is referred to the relevant literature /31/ for these data. Whole salt (MF) data are also available in H 2 0/H mixtures /26/. Figure 8 plots selected data from Table 8 to better illustrate the variation of A t G(F~) with solvent composition in aqueous-organic mixtures. All curves show a steady increase in A t G(F~) with increasing organic content, with the curves becoming especially steep at high organic concentrations. This behaviour is indicative /3, 9/ of the selective solvation of the fluoride ion by water molecules in the water-organic mixtures /33/. As water is a very much better Η-bonder than most other solvents A t G(F") becomes increasingly less favourable as the water content decreases. Not surprisingly, A t G(F") H2 Q->-H Me0H correlates very well with the (mixed) solvent acceptor strength, as measured by the Gutmann-Mayer acceptor numbers /33/. The often sharp increase in A t G(F~) at high organic concentrations affords a partial explanation for the lack of consistency among independent determinations of A t G(MF) H2 o->s (Table 1) as even slight traces of water will have a significant effect on the observed values in "pure" nonaqueous solvents. The variation of A t G(F~) with solvent composition closely parallels those of the other halide ions as shown in Figure 9 for acetonitrile-water mixtures. Similar curves are obtained for other aqueous solvent mixtures (e.g., /48/). The increase in A t G(X~) in Figure 9 is, of course, consistent with the relative Η-bonding strengths of the halide ions. The slightly negative values for A t G(I~) at low organic concentrations and the inflection in the A t G(F~) curve, are a reminder of the complexity of the solute-solventsolvent interactions which are present in aqueous organic mixtures /9, 64/. Comparison of A t G(F~) for the various alcohols is interesting. The data in Table 8 clearly show that A t G(F~) H20 ->.H R0H becomes increasingly positive in going from MeOH to EtOH to t-buoh consistent with their decreasing ability to stabilise the fluoride ion by Η-bonding as a result of the inductive effect of the alkyl chain 214

31 G. Τ. Hefter Reviews in Inorganic Chemistry R Η -ι : F" the Εχ values for these solvents being respectively 55.4, 51.9 and Fig. 8. Gibbs free energies of transfer of fluoride as a function of solvent composition for selected aqueous-organic solvent mixtures. Data from Table

32 VoL 10, Nos. 1-3,1989 Solvation of Fluoride Ions, I I SOLVENT COMPOSITION (100x AN ) Fig. 9. Gibbs free energies of transfer of the halide ions as a function of solvent composition in AN/H 2 0 mixtures. Fluoride data from Table 8, other halides from refs. /67, 68/ (smoothed data). Although no data are available for F~ in 1-BU0H/H 2 0 mixtures, with Ej (1-BuOH)) = 50.2 the values of A t G(F') Hj0 _,. H _ Bu 0H would be expected to be somewhat more positive than the Et0H/H 2 0 values (Table 8). The good correlation between Εχ and A t G(F~) H20^. Ho0 + ROH Ρ Γ " vides another reason to suspect the value of A t G(F~) to pure l-buoh in Table

33 G.T.Hefter Reviews in Inorganic Chemistry The values of A t G(F~) in ethyleneglycol (EG)/H 2 0 mixtures are also noteworthy. It has been suggested that EG undergoes a strong interaction with F", possibly via an Η-bonded chelate complex /33/. Certainly, NMR evidence for strong Η-bonding between EG (and other polyols) and F~ has been reported /65/. However, the data in Table 8 /33/ indicate that these bonds are not thermodynamically predominant as A t G(F~) is more positive for EG than for MeOH despite the fact that EG is a slightly better electron acceptor (E T = 56.3 for EG cf for MeOH /9/) Other Thermodynamic Quantities To date only one study /48/ has reported enthalpies and entropies of transfer of fluoride in a mixed solvent system and then only for one salt in Me0H/H 2 0. These data are summarised in Table 9. TABLE 9. Mixtures at K a Thermodynamic Transfer Properties for Fluoride in Aqueous Methanol MeOH composition (100 x) A t G(F") A t H(F") A t S(F~) TA t S(F") a Graphically interpolated from original data /48/. Units: kj mol mol -1 ). TATB or TPTB assumption, mol L -1 scale. 1 (A t S, J K

34 Vol 10, Nos. 1-3,1989 Solvation of Fluoride Ions, 3. As is often found to be the case the monotonic change in A t G (Table 8 and Figure 8) disguises dramatic, but largely compensating, changes in A t H and A t S /3/. The enthalpy data for the fluoride ion along with those of the other halides are plotted in Figure 10. A similar family of curves is obtained for Δ { 8(Χ~) Η2 0-*Η Μβ0Η / 48 / SOLVENT COMPOSITION <v/v% MeOH) Fig. 10. Enthalpies of transfer of the halide ions as a function of solvent composition for Me0H/H 2 0 mixtures (after ref. /48/). The shape of these curves can be rationalised as follows /48/. The addition of a small amount of MeOH to water enhances the latter's 3-D structure with the MeOH molecules mainly being accommodated within the voids of the water structure. A t H(F") is therefore initially unfavourable (positive) because of increased competition for H 2 0 molecules between F~ and other solvent molecules (mainly H 2 0 at these compositions). As the MeOH 218

35 G.T. Hefter Reviews in Inorganic Chemistry concentration increases, the three-dimensional water structure begins to break down. Volume of mixing data suggest this occurs at about 10-20% MeOH. Thus, further addition of MeOH frees up more H 2 0 molecules to bond to (solvate) the fluoride ions and A t H(F~) becomes more negative. The final increase in A t H(F~) at high MeOH concentrations is a reflection of H 2 0 molecules being replaced in the coordination shell of the fluoride ion by the less strongly coordinating (weaker Lewis acceptor or hydrogenbond donor) MeOH molecules. Not surprisingly this effect is somewhat greater for F~ than for the other halide ions. Similar arguments apply to the entropies of transfer /48/. Virtually no data exist for the other thermodynamic parameters for transfer from water to mixed solvents. However, some limited information on apparent molar heat capacities, A t C p ^ and volumes, have been reported for TBA/HjO and AN/H 2 0 mixtures at low organic concentrations /52, 66/. 6. Conclusions Relatively few data, many of them of dubious quality, have been reported for the thermodynamic transfer properties of the fluoride ion from water to nonaqueous or mixed solvents. Although measurement of A t X(F~), where X = G, H, S, etc. is somewhat more difficult for the fluoride ion than for the other halides, with appropriate attention to experimental detail it should be possible to obtain reliable results and considerable scope exists for further work in this area. Future measurements could be profitably directed towards deciding whether there is really a relationship between A t G(F~)H 2 o->s an d μ δ, the relative importance of entropy and enthalpy in determining the magnitude of A t G(F~)H 2 o->-s and whether there is a significant difference in A t X(F~)H 2 o-»-s between protic and aprotic solvents. 7. Acknowledgements The author thanks Profs. C. Kaiidas, D. Feakins and Y. Marcus, and Drs. M. Salomon, and B.G. Cox for access to their information on fluoride salt solubilities, Dr. B.W. Clare for assistance with calculations and Ms. P. McLay for assistance with data collection and correlation. 219

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39 G.T. Hefter Reviews in Inorganic Chemistry APPENDIX List of Solvent Abbreviations AC - acetone; propanone AN - acetonitrile; ethanenitrile t-buoh - tertiary-butanol; 2-methyl-2-propanol DCE 1,2-dichloroethane DG - diglyme; 1,r-oxybis(2-methoxy)ethane DME - 1,2-dimethoxyethane; monoglyme DMF - Ν,Ν-dimethylformamide; N,N-dimethylmethanamide DMSO - dimethylsulphoxide DX - 1,4-dioxane EG - ethylene glycol; 1,2-dihydroxyethane EtOH - ethanol FA formamide; methanamide GBL - γ-butyrolactone; 2-oxotetrahydrofuran ME - 2-methoxyethanol MeOD - O-deuteromethanol MeOH - methanol NB nitrobenzene NMA - N-methylacetamide; N-methylethanamide NMF - N-methylformamide; N-methylmethanamide PC - propylene carbonate; 4-methyl-l,3-dioxolan-2-one 1-PrOH - 1-propanol TG - tetraglyme; 2,5,8,11,14-pentaoxapentadecane THF - tetrahydrofuran 223