- Email: [email protected]
The acidity of superacidic media derived from fluorosulfuric acid and trifluoromethanesulfonic acid
Analytica Chimica Acta, 159 (1984) 149-158 Elsevier Science Publishers B.V., Amsterdam -Printed
in The Netherlands
THE ACIDITY OF SUPERACIDIC MEDIA DERIVED FROM FLUOROSULFURIC ACID AND TRIFLUOROMETHANESULFONIC ACID
BERNARD
CARRE
Laboratoire de Chimie GBnlrale et Minkrale. Facultt! des Sciences, 3 7200 Tours (France)
Part Grammont,
JACQUES DEVYNCK* Laboratoire d%lectrochimie Ecole Nationale Supkrieure Paris Cedex 05 (France)
Analytique et AppliquBe (associk au CNRS, LA 216). de Chimie de Paris, 11, rue Pierre et Marie Curie, 75231
(Received 29th July 1983)
SUMMARY The acidity level of superacidic media derived from fluorosulfuric acid and trifluoromethanesulfonic acid (TFMS) was determined by using the potentiometric indicators tetrachloroquinone and tetrafluoroquinone. A voltammetric and potentiometric study of these indicators was conducted to evaluate their pH indicator range, from the strongest acid media (SbF, solution) to the strongest base media (H,O or KF solution). The R, function in 1 M strong base medium was found to be R, = -17.2 (fluorosulfuric acid) and R, = -16 (TFMS). The autoprotolysis constant (pK3 was evaluated in both solvents; PKi = 6.1 (fluorosulfuric acid) and pKf = 6 (TFMS).
Fluorosulfuric acid has been used, alone or mixed with antimony pentafluoride, as a superacidic medium, to achieve the activation of hydrocarbons, especially the isomerization of alkanes [l-9]. The derived perfluoroalkanesulfuric acids have also been considered for such an application. Among the latter, trifluoromethanesulfonic acid (TFMS) has a high acidity level, and has been offered in the monohydrate form for making fuel cell electrolytes [ 10, 111. Although the properties of these acids are well known, very few results are available concerning their relative acidity levels; Ho functions have been determined for fluorosulfuric acid [12, 131, but only one Ho function is known for TFMS [ 141. The most important properties of these media are known to depend on their acidity level when “neat”, but when the medium also contains acidic (e.g., SbF5) or basic (e.g., HzO) solutes, this level also depends on the autoprotolysis constant [ 153. Therefore, precise correlation is necessary between these media and other superacidic media, especially those derived from hydrofluoric acid which have been extensively studied [ 151. In the work reported here, potentiometric indicators with an acidity of a type similar to those 0003-2670/84/$03.00
0 1984 Elsevier Science Publishers B.V.
150
already described were studied in order to define their R,,(H) functions [ 161 and to relate them to each other and to other superacidic media, as well as to evaluate their autoprotolysis constants. Fluorosulfuric acid Fluorosulfuric acid (HSOJF) dissociates according to the reaction [ 171 2HSOBF * H2S03F+ + SOBF-;RI = [H2S03F+] [SOBF-] = 4.8 X lo4 Although water is basic, it can react in two different ways [la], strong base at low temperature
(1) either as a
H20 + HS03F --f H30+ + FSO, or roughly by destroying the solvent when added carelessly H?O + HSOBF + HZS04 + HF No elementary protonic acids are known to be strong in fluorosulfuric acid. The strongest acids known are fluorosulfonate-accepting Lewis acids [ 121 SbFS + HSOBF * H,S03F+ + (SbF,)FSO, There is an internal dissociation of the solvent from its constituent elements: HSOBF * HF + SOa (K = [SO,] [HF] = lo*). The presence of sulfur trioxide affects the reactions because sulfur(V1) is a strong oxidizing agent. The Ho functions have been determined by Gillespie and Peel [ 121: for “neat” fluorosulfuric acid; Ho = -11.5 for HS03F + 1M HO = -15.07 KSOsF; and Ho = -18.9 for HS03F + 1 M SbF, [13]. Sommer et al. [19] have checked the protonation of the Hammet indicators through n.m.r. spectrometry. Trifluoromethanesulfonic acid The acidic properties of trifluoromethanesulfonic acid ( HSOBCFB; TFMS) have not been studied extensively and very few constants have been determined. The acid dissociation would be 2HS03CF3 * H2S03CF: + SO&&;
KI = [H$O$F;]
[SO&F,]
(2)
The solvent pKi has never been determined. Water and alkaline trifluoromethanesulfonates are strong bases [20], and like fluorosulfuric acid, the strongest acids have to be Lewis acids of the MFS type (SbF,). Only the HO function of the supposedly “neat” acid, HO = -14.1, is known [ 141. EXPERIMENTAL
The equipment used was similar to that described earlier [ 21,221. Selection of acidity indicators The use of electrochemical acidity indicators dissolved in these solvents is of considerable interest. The standard hydrogen electrode does not act
151
properly. In fluorosulfuric acid, the H+/H, system at a platinized platinum electrode is not reversible (slow system) and hydrogen is oxidized by sulfur(VI) (SO3 is produced by the internal dissociation of the solvent). The electrochemical behaviour of hydrogen is similar in TFMS. As described earlier [21, 221, quinone systems give reversible electrochemical systems which are pH indicators and are stable in these strongly acidic media. The weakly basic quinones, tetrachloro-p-benzoquinone (TCQ) and tetrafluoro-p-benzoquinone (TFQ), were used, as well as the corresponding hydroquinones. The protonation of quinones can be monitored by n.m.r. [ 231. Both quinones were studied in hydrofluoric acid: in basic medium, they reduce in a unique single step Q(O)/Q(-II): Q + 2e-+ 2I-P + QH,. One proton is exchanged by one electron. In strongly acidic medium (SbFS), quinones reduce in two oneelectron steps. A semiquinone formation Q(-I) has also been demonstrated. The reactions are as follows for Q(O)/Q(-I), for Q(-I)/Q(-II),
QH+ + e- + H++ QH;+ QH;+ + e-+ QH, or QH;++ e-+ H++ QH:
(3) (4)
depending on the pH.
Reference system The most suitable reference electrode used was a silver electrode ( Ag/Ag+) [24] in TFMS or a copper electrode (Cu/Cu2’) [ 251 in fluorosulfuric acid. All potentials are measured by comparison with a system supposedly independent of the solvent. It has been shown that perylene [25, 26, 271 and octafluoronaphthalene [28] may serve as such references in very acidic medium. Perylene. Perylene quickly oxidizes to a monocationic form in these acidic solvents. When yellow perylene is dissolved, the solution turns green first and purple quickly afterwards, a colour characteristic of the monocation, which can be oxidized to the dication in a reversible one-electron process (Fig. 1). With a rotating platinized platinum electrode in a 1 M strong base medium, the half-wave potentials found were Eln = 0.300 V for fluorosulfuric acid and Eln = 0.280 V for TFMS. Octufluoronaphthalene. This compound is oxidized giving three different oxidation waves: the first wave corresponds to a reversible one-electron process over the whole accessible pH range in both solvents: C8F8 + C&F: + e-. The cationic radical prepared by controlled potential electrolysis is stable and turns green (as in hydrofluoric acid [ 281). The first wave is readily reversible, while the other two are not and have a stronger intensity. With a rotating platinized platinum electrode in 1 M strong base, the half-wave potentials obtained for the first oxidation wave were Ela = 0.850 V in fluorosulfuric acid and El,? = 0.830 V in TFMS. Octafluoronaphthalene was chosen as the reference in the present work. Its oxidation potential is very high and its chemical stability is remarkable in these acidic solvents.
152
I (/A)
Fig. 1. Voltammetric curves of octafluoronaphthalene and perylene in trifluoromethanesulfonic acid: (1) residual current; (2) 0.5 mM perylene; (3) 1 mM octafluoronaphthalene; (4) 1 mM octafluoronaphthalene + 0.4 mM perylene.
Voltammetric study of quinones Several electrodes were considered (platinum, gold, glassy carbon) but the best results were obtained on platinized platinum (Table 1). Figures 2 and 3 show the evolution of the voltammograms of the quinones as a function of pH over the whole acidity range. The decrease of the limiting diffusion current in strongly basic media (Fig. 2, curves 3-8 and Fig. 3, curves 1-3) corresponds to a decrease in the solubility. Table 1 shows the E,,2 values and the slopes of E vs. log[(l, - 1)/I] plots for the reduction wave of the tetrachloroquinone and tetrafluoroquinone in basic and in neutral (neat solvent) media. The reduction proceeded in one reversible wave. The slope of the variation of Eln vs. logarithm of base concentration for both quinones in strong base (H,O) media in fluorosulfuric acid, is reported in Table 2. The same variation was observed with fluoride ion as the strong base. The slopes of the variation TABLE 1 Voltammetric data for tetrachloroquinone (TCQ) and tetrafluoroquinone fluorosulfuric acid and in trifluoromethanesulfonic acid Medium
TCQ
(TFQ)
in
TFQ
HSO,F*
H,SO,CF,
b
HSO,F*
H,SO,CF,
“Neutral”
E,, SC
(V)
0.630 55
0.360 75
0.470 80
0.220 90
Basic
E,, SC
(V)
0.510 55
0.270 65
0.330 70
0.110 70
b
‘1 M sodium fluoride. bl M water. Wlope of the linear plot for the wave: E vs. log[(~d-z)/zl*
153
Fig. 2. Voltammetric curves of 1 mM tetrachloroquinone in trifluoromethanesulfonic acid at various acidity levels: (1) 0.64 M SbF,; (2) 0.12 M SbF,; (3) “neutral” media; (4) 0.25 M H,O; (5) 0.76 M H,0;(6) 1.26 M H,0;(7) 2.5 M H,O; (8) 7 M H,O. Fig. 3. Voltammetric curves of 1 mM tetrafluoroquinone in fluorosulfuric acid at various acidity levels: (1) “neutral” media; (2) 5 x lOa M NaF; (3) 10’ M NaF; (4) 0.13 M SbF, ; (5) 0.54 M SbF,. are also close to the theoretical value (59 mV) in the case of trifluoroof El/2 methanesulfonic acid (Table 2). These variations imply an exchange of 1 electron per IT in the redox reaction Q + 2W + 2e-* QH,. The equation of the voltammetric curve is
E = E” - 0.06 pH - 0.06 log [(Id - r)/I] The results are less simple in “acidic” media and depend on the SbFs concentration. Tetrachloroquinone presents two irreversible reduction waves in both solvents; the waves shift to more positive potentials as the antimony pentafluoride concentration increases (Fig. 2, curves 2 and 3). The quinone is dark purple in acidic solution. Tetrafluoroquinone also gives two waves in dilute antimony pentafluoride solution (<0.4 M SbFs in fluorosulfuric acid or <0.2 M in trifluoromethanesulfonic acid; see Fig. 3, curve 4 and Fig. 4, curve 2) and two waves in more concentrated antimony pentafluoride solution (e.g., Fig. 3, curve 5). By analogy with the results obtained in hydrofluoric acid where the same kind of behavior has been observed [23], the twowave system indicates two one-electron steps corresponding to the semiquinone stabilization. The two reduction steps of the quinone correspond to the redox equilibria, QW + H’ * QHj+ and QHj’ + e- =+QH:. Voltammograms of a 1:l mixture of tetrafluoroquinone and tetrafluoro-
154 TABLE 2 Half-wave reduction potential shift of tetrachloroquinone (TCQ) and tetrafluoroquinone (TFQ) with base concentration (C) in fluorosulfuric acid and trifluoromethanesulfonic acid (TFMS) 10e4 M < C < 10-l M) (&I, = a log C + 5; C = H,O or F-concentration Fluorosulfuric
TCQ
LmV) (mV)
TFQ
57 f 38
a
Trifluoromethane sulfonic acid 6Oi 2a
516 f 5
291 f 4
63 f 3
68 * 3
326 i 5
103 f 3
II LmV) (mV)
acid
aConfidence limits, probability:
95%.
hydroquinone were obtained in trifluoromethanesulfonic acid (Fig. 4); the equilibrium potential of a platinum electrode at which the system is reversible is directly linked to the pH of the solution by E = E" -0.06 pH. The voltammograms are well defined in basic and in neutral media (curves) but distortion appears in acidic media, probably because of the oxidizing power of antimony(V) fluoride.
Fig. 4. Voltammetric curves of a mixture of 5 mM tetrafluoroquinone and 5 mM tetrafluorohydroquinone in trifluoromethanesulfonic acid at various acidity levels: (1) < lOa MH,0;(2)0.13MSbF,;(3)0.4MSbF,;(4)0.7MH,0;(5)5MH,O.
155 RESULTS AND DISCUSSION
De termination of the Ro (H) acidity functions The evaluation of the reduction half-wave potential of quinones (or equilibrium potential of a mixture of the quinone and hydroquinone) allows calculation of the&(H) acidity function. The expression for R(H) is R(H) = F(2.3 BT)-’
(&n(ao)
-h!d
w&m J-b caq) refers to the potential of the pH indicator in water and El12(,, refers to the potential of the same pH indicator in the solvent, both potentials being expressed by reference to the half-wave potential of a reversible system supposed to be independent of the solvent [ 161. The solvent-independent system is octafluoronaphthalene (OFN/OFN+) which was previously studied in hydrofluoric acid [28]. Because this compound is insoluble in water, its half-wave potential must be evaluated by using an intermediate compound which is soluble in the solvent and in water. Perylene/perylene+ can be used for this purpose. Tetrachloroqumone was used to evaluate theR(H) function of the strongly basic solution. The reduction of this quinone proceeds by the same scheme in water and in TFMS or fluorosulfuric acid solution. The El,2 values for the reduction of tetrachloroquinone in aqueous solution are as follows: 0.260 V vs. ferrocene (Fe/Fe+), 0.840 V vs. perylene (Pe’/Pe’+) and 1.390 V vs. octafluoronaphthalene (OFN/OFN+). The corresponding calculated values of I&,(H) in sodium fluoride or water solutions are reported in Table 3. These values correlate with the R,(H) function obtained with the same quinone in hydrofluoric acid, which wasR,(H) = 14.2 (1 M KF) [22,29]. The tetrafluoroquinone reduction presents the same one-step, twoelectron reversible reduction in strongly basic solution and in dilute strong acid solution (SbF5 < 0.4 M in fluorosulfuric acid and <0.2 M in TFMS).
TABLE 3 R,(H) acidity function of basic (Hz0 or NaF) and acidic (SbF,) media in fiuorosulfuric and trifluoromethanesulfonic acid. octafluoronaphthaiene (OFN/OFN+). Acidity indicators: tetraPotential reference: chloroquinone (TCQ) and tetrafluoroquinone (TFQ) Medium
J%
HSO,F + 1 MNaF HSO,F + 1 M SbF, HSO,CF, + 1 M H,O HSO,CF, + 1 M SbF,
0.360 0.330 0.694
TCQ TFQ TFQ
-17.2 -17.2 -23.4
0.430 0.210 0.572
TCQ TFQ TFQ
-16 -16 -22.1
0’)
Indicator
R,(H)
156
Thus, the same system can be used to calculate the acidity function of the strongest acid solution. Extrapolated values (corresponding to a 1 M SbF, solution) are reported in Table 3. Calculation 0 f apparent au topro tolysis constants According to the autoprotolysis equilibria (Eqns. 1 and 2) reported above, the evaluation of the acidity function of a strong acid and a strong base (NaF or H@) with the same electrochemical pH indicator leads to the determination of the autoprotolysis constant, pK1. The values obtained with tetrafluoroquinone are pKi = R,(H) (1 M SbFS) - R,(H) (1 M NaF) = 6.2 k 0.5 for fluorosulfuric acid, if SbFs is considered as a strong acid as in hydrofluoric acid, and pKI = R,(H) (1 M SbFs) --R,(H) (1 M HzO) = 6.1 f 0.5 for trifluoromethanesulfonic acid. The value for pKi in fluorosulfuric acid is one unit lower than that obtained by Barr et al. [ 171 by conductimetric measurements because they have considered SbFS as only partially ionized (pK = 2.15). There is no publishedvalue in the case of trifluoromethanesulfonic acid. Comparison between acidity ranges Figure 5 shows the three fluorinated acid solvents, hydrofluoric, fluorosulfuric and trifluoromethanesulfonic acids, related to the R0 scale. The correspondence with the I& scale is represented, and the results confirm the 6-unit deviation between R0 and Ho functions already observed in
-% - _ _ _ ____
IF --
16.:
11.5 10.3
_
---
-_
-
-
--
-
-_-__
..
22
172
- -__.
-
HSosCF,
8.1
2L.6
----r .
-----A-
- -_-.
229
_ _ _..
16
IL.2
HF Fig. 5. Relative position of the acidity scales of fluorosulfuric, and hydrofluoric acids, and the relationship between the R,(H)
trifluoromethanesulfonic and H, acidity functions.
157
hydrofluoric acid [14]. It can be seen that hydrofluoric acid has the widest acidity range, which enables the most acidic and the most basic media to be reached (p.& = 13.7). Conclusion and tetrafluoroquinone As in hydrogen fluoride, tetrachloroquinone appear to be suitable pH indicators; they are sufficiently stable and their protonation allows the pH to be evaluated in the whole acidity range of the solvents. These quinones are the two electrochemical pH indicators useful in very acidic media. Comparison between the values determined for RO and HO shows that the difference between these two functions is almost constant. The difference found is the same as in hydrofluoric acid. The pK, values determined for fluorosulfuric and trifluoromethanesulfonic acid through RO measurements are of the same magnitude. The accessible pH range covers 6 units. Thus superacid media based on these two acids makes it possible to reach very acidic media by addition of antimony(V) fluoride (R,, = -25), but the acidity range is not as wide as that of hydrofluoric acid. The presence of S(VI), from the sulfur trioxide produced secondarily to their internal dissociation, differentiates these two acids from hydrofluoric acid. Sulfur trioxide is a potent oxidizing agent which participates directly in oxidations. It oxidizes hydrogen so that it becomes impossible, in practice, to use a hydrogen electrode in sulfur-containing superacids, as in sulfuric acid. REFERENCES 1 2 3 4
G. A. Olah and J. Lukas, J. Am. Chem. Sot., 89 (1967) 2227. H. Hogeveen and J. Gaasbeek, Rec. Trav. Chhn., 88 (1969) 703. G. A. Olah, Y. Haipern, J. Shen and Y. K. MO, J. Am. Chem. Sot., 95 (1973) 4960. J. W. Larsen, P. A. Bovis, C. R. Watson and R. M. Pagni, J. Am. Chem. Sot., 96 (1974) 2284. 5 C. Pitti, F. Bobillart, A. Thiebault and M. Herlem, Anal. Lett., 8 (1975) 241. 6 M. Herlem, F. Bobillart, A. Thiebault and A. Jobert, Anal. Lett., 10 (1978) 767. 7 C. Pitti, M. Cerles, A. Thiebault and M. Herlem, J. Electroanal. Chem., 126 (1981) 163. 8 G. M. Kramer, J. Org. Chem., 40 (1975) 5173. 9 D. Brunel, A. Germain and A. Commeyras, Nouv. J. Chim., 2 (1978) 275. 10 V. B. Hughes, B. D. McNicol, M. R. Andrew, R. B. Jones and R. T. Short, J. Appl. Electrochem., 7 (1977) 161. 11 A. A. Adams and R. T. Foley, J. Electrochem. Sot., 126 (1979) 775. 12 R. J. Gillespie and T. E. Peel, J. Am. Chem. Sot., 95 (1973) 5173. 13 H. H. Hyman, M. Kiipatrick and J. J. Katz, J. Am. Chem. Sot., 79 (1957) 3668. 14 J. Grondin, R. Sagnes and A. Commeyras, Bull. Sot. Chim. Fr., 11 (1976) 1779. 15 P. L. Fabre, J. Devynck and B. Tremiilon, Chem. Rev., 82 (1982) 591. 16 H. Strehlow and H. Wendt, Z. Phys. Chem. (Frankfurt am Main), 30 (1960) 141. 17 J. Barr, R. J. Gillespie and R. C. Thompson, Inorg. Chem., 3 (1964) 1149. 18 G. Adhami, These de doctorat d’&at, Paris VI (1972). 19 J. Sommer, P. Canivet, S. Schwartz and P. Rimmelin, Nouv. J. Chhn., 5 (1981) 45.
158 20 D. C. Russel and J. D. Senior, Can. J. Chem., 52 (1974) 2975. 21 J. Devynck, A. Ben Hadid, P. L. Fabre and B. TrBmillon, Anal. Chim. Acta, 100 (1978) 343. 22 J. P. Masson, J. Devynck and B. Trdmillon, J. Electroanal. Chem., 64 (1975) 193. 23 A. Ben Hadid, P. Rimmelin, J. Sommer and J. Devynck, J. Chem. Sot. Perkin Trans. 2, (1982) 269. 24 J. Verastegui, These 38me cycle, Universitd Pierre et Marie Curie, Paris VI (1973). 25 A. Jobert, These docteur ing&ieur, Universite Pierre et Marie Curie, Paris VI (1979). 26 D. Bauer and M. Bouchet, C. R. Acad. Sci., Sbr. C, (1972) 274. 27 D. Bauer, J. P. Beck and P. Texier, Collect. Czech. Chem. Commun., 36 (1971) 360. 28 A. Ben Hadid, Thbse d’Qtat, Universite Pierre et Marie Curie, Paris VI (1980). 29 J. P. Masson, J. Devynck and B. Tremillon, J. Electroanal. Chem., 64 (1975) 175.
in The Netherlands
THE ACIDITY OF SUPERACIDIC MEDIA DERIVED FROM FLUOROSULFURIC ACID AND TRIFLUOROMETHANESULFONIC ACID
BERNARD
CARRE
Laboratoire de Chimie GBnlrale et Minkrale. Facultt! des Sciences, 3 7200 Tours (France)
Part Grammont,
JACQUES DEVYNCK* Laboratoire d%lectrochimie Ecole Nationale Supkrieure Paris Cedex 05 (France)
Analytique et AppliquBe (associk au CNRS, LA 216). de Chimie de Paris, 11, rue Pierre et Marie Curie, 75231
(Received 29th July 1983)
SUMMARY The acidity level of superacidic media derived from fluorosulfuric acid and trifluoromethanesulfonic acid (TFMS) was determined by using the potentiometric indicators tetrachloroquinone and tetrafluoroquinone. A voltammetric and potentiometric study of these indicators was conducted to evaluate their pH indicator range, from the strongest acid media (SbF, solution) to the strongest base media (H,O or KF solution). The R, function in 1 M strong base medium was found to be R, = -17.2 (fluorosulfuric acid) and R, = -16 (TFMS). The autoprotolysis constant (pK3 was evaluated in both solvents; PKi = 6.1 (fluorosulfuric acid) and pKf = 6 (TFMS).
Fluorosulfuric acid has been used, alone or mixed with antimony pentafluoride, as a superacidic medium, to achieve the activation of hydrocarbons, especially the isomerization of alkanes [l-9]. The derived perfluoroalkanesulfuric acids have also been considered for such an application. Among the latter, trifluoromethanesulfonic acid (TFMS) has a high acidity level, and has been offered in the monohydrate form for making fuel cell electrolytes [ 10, 111. Although the properties of these acids are well known, very few results are available concerning their relative acidity levels; Ho functions have been determined for fluorosulfuric acid [12, 131, but only one Ho function is known for TFMS [ 141. The most important properties of these media are known to depend on their acidity level when “neat”, but when the medium also contains acidic (e.g., SbF5) or basic (e.g., HzO) solutes, this level also depends on the autoprotolysis constant [ 153. Therefore, precise correlation is necessary between these media and other superacidic media, especially those derived from hydrofluoric acid which have been extensively studied [ 151. In the work reported here, potentiometric indicators with an acidity of a type similar to those 0003-2670/84/$03.00
0 1984 Elsevier Science Publishers B.V.
150
already described were studied in order to define their R,,(H) functions [ 161 and to relate them to each other and to other superacidic media, as well as to evaluate their autoprotolysis constants. Fluorosulfuric acid Fluorosulfuric acid (HSOJF) dissociates according to the reaction [ 171 2HSOBF * H2S03F+ + SOBF-;RI = [H2S03F+] [SOBF-] = 4.8 X lo4 Although water is basic, it can react in two different ways [la], strong base at low temperature
(1) either as a
H20 + HS03F --f H30+ + FSO, or roughly by destroying the solvent when added carelessly H?O + HSOBF + HZS04 + HF No elementary protonic acids are known to be strong in fluorosulfuric acid. The strongest acids known are fluorosulfonate-accepting Lewis acids [ 121 SbFS + HSOBF * H,S03F+ + (SbF,)FSO, There is an internal dissociation of the solvent from its constituent elements: HSOBF * HF + SOa (K = [SO,] [HF] = lo*). The presence of sulfur trioxide affects the reactions because sulfur(V1) is a strong oxidizing agent. The Ho functions have been determined by Gillespie and Peel [ 121: for “neat” fluorosulfuric acid; Ho = -11.5 for HS03F + 1M HO = -15.07 KSOsF; and Ho = -18.9 for HS03F + 1 M SbF, [13]. Sommer et al. [19] have checked the protonation of the Hammet indicators through n.m.r. spectrometry. Trifluoromethanesulfonic acid The acidic properties of trifluoromethanesulfonic acid ( HSOBCFB; TFMS) have not been studied extensively and very few constants have been determined. The acid dissociation would be 2HS03CF3 * H2S03CF: + SO&&;
KI = [H$O$F;]
[SO&F,]
(2)
The solvent pKi has never been determined. Water and alkaline trifluoromethanesulfonates are strong bases [20], and like fluorosulfuric acid, the strongest acids have to be Lewis acids of the MFS type (SbF,). Only the HO function of the supposedly “neat” acid, HO = -14.1, is known [ 141. EXPERIMENTAL
The equipment used was similar to that described earlier [ 21,221. Selection of acidity indicators The use of electrochemical acidity indicators dissolved in these solvents is of considerable interest. The standard hydrogen electrode does not act
151
properly. In fluorosulfuric acid, the H+/H, system at a platinized platinum electrode is not reversible (slow system) and hydrogen is oxidized by sulfur(VI) (SO3 is produced by the internal dissociation of the solvent). The electrochemical behaviour of hydrogen is similar in TFMS. As described earlier [21, 221, quinone systems give reversible electrochemical systems which are pH indicators and are stable in these strongly acidic media. The weakly basic quinones, tetrachloro-p-benzoquinone (TCQ) and tetrafluoro-p-benzoquinone (TFQ), were used, as well as the corresponding hydroquinones. The protonation of quinones can be monitored by n.m.r. [ 231. Both quinones were studied in hydrofluoric acid: in basic medium, they reduce in a unique single step Q(O)/Q(-II): Q + 2e-+ 2I-P + QH,. One proton is exchanged by one electron. In strongly acidic medium (SbFS), quinones reduce in two oneelectron steps. A semiquinone formation Q(-I) has also been demonstrated. The reactions are as follows for Q(O)/Q(-I), for Q(-I)/Q(-II),
QH+ + e- + H++ QH;+ QH;+ + e-+ QH, or QH;++ e-+ H++ QH:
(3) (4)
depending on the pH.
Reference system The most suitable reference electrode used was a silver electrode ( Ag/Ag+) [24] in TFMS or a copper electrode (Cu/Cu2’) [ 251 in fluorosulfuric acid. All potentials are measured by comparison with a system supposedly independent of the solvent. It has been shown that perylene [25, 26, 271 and octafluoronaphthalene [28] may serve as such references in very acidic medium. Perylene. Perylene quickly oxidizes to a monocationic form in these acidic solvents. When yellow perylene is dissolved, the solution turns green first and purple quickly afterwards, a colour characteristic of the monocation, which can be oxidized to the dication in a reversible one-electron process (Fig. 1). With a rotating platinized platinum electrode in a 1 M strong base medium, the half-wave potentials found were Eln = 0.300 V for fluorosulfuric acid and Eln = 0.280 V for TFMS. Octufluoronaphthalene. This compound is oxidized giving three different oxidation waves: the first wave corresponds to a reversible one-electron process over the whole accessible pH range in both solvents: C8F8 + C&F: + e-. The cationic radical prepared by controlled potential electrolysis is stable and turns green (as in hydrofluoric acid [ 281). The first wave is readily reversible, while the other two are not and have a stronger intensity. With a rotating platinized platinum electrode in 1 M strong base, the half-wave potentials obtained for the first oxidation wave were Ela = 0.850 V in fluorosulfuric acid and El,? = 0.830 V in TFMS. Octafluoronaphthalene was chosen as the reference in the present work. Its oxidation potential is very high and its chemical stability is remarkable in these acidic solvents.
152
I (/A)
Fig. 1. Voltammetric curves of octafluoronaphthalene and perylene in trifluoromethanesulfonic acid: (1) residual current; (2) 0.5 mM perylene; (3) 1 mM octafluoronaphthalene; (4) 1 mM octafluoronaphthalene + 0.4 mM perylene.
Voltammetric study of quinones Several electrodes were considered (platinum, gold, glassy carbon) but the best results were obtained on platinized platinum (Table 1). Figures 2 and 3 show the evolution of the voltammograms of the quinones as a function of pH over the whole acidity range. The decrease of the limiting diffusion current in strongly basic media (Fig. 2, curves 3-8 and Fig. 3, curves 1-3) corresponds to a decrease in the solubility. Table 1 shows the E,,2 values and the slopes of E vs. log[(l, - 1)/I] plots for the reduction wave of the tetrachloroquinone and tetrafluoroquinone in basic and in neutral (neat solvent) media. The reduction proceeded in one reversible wave. The slope of the variation of Eln vs. logarithm of base concentration for both quinones in strong base (H,O) media in fluorosulfuric acid, is reported in Table 2. The same variation was observed with fluoride ion as the strong base. The slopes of the variation TABLE 1 Voltammetric data for tetrachloroquinone (TCQ) and tetrafluoroquinone fluorosulfuric acid and in trifluoromethanesulfonic acid Medium
TCQ
(TFQ)
in
TFQ
HSO,F*
H,SO,CF,
b
HSO,F*
H,SO,CF,
“Neutral”
E,, SC
(V)
0.630 55
0.360 75
0.470 80
0.220 90
Basic
E,, SC
(V)
0.510 55
0.270 65
0.330 70
0.110 70
b
‘1 M sodium fluoride. bl M water. Wlope of the linear plot for the wave: E vs. log[(~d-z)/zl*
153
Fig. 2. Voltammetric curves of 1 mM tetrachloroquinone in trifluoromethanesulfonic acid at various acidity levels: (1) 0.64 M SbF,; (2) 0.12 M SbF,; (3) “neutral” media; (4) 0.25 M H,O; (5) 0.76 M H,0;(6) 1.26 M H,0;(7) 2.5 M H,O; (8) 7 M H,O. Fig. 3. Voltammetric curves of 1 mM tetrafluoroquinone in fluorosulfuric acid at various acidity levels: (1) “neutral” media; (2) 5 x lOa M NaF; (3) 10’ M NaF; (4) 0.13 M SbF, ; (5) 0.54 M SbF,. are also close to the theoretical value (59 mV) in the case of trifluoroof El/2 methanesulfonic acid (Table 2). These variations imply an exchange of 1 electron per IT in the redox reaction Q + 2W + 2e-* QH,. The equation of the voltammetric curve is
E = E” - 0.06 pH - 0.06 log [(Id - r)/I] The results are less simple in “acidic” media and depend on the SbFs concentration. Tetrachloroquinone presents two irreversible reduction waves in both solvents; the waves shift to more positive potentials as the antimony pentafluoride concentration increases (Fig. 2, curves 2 and 3). The quinone is dark purple in acidic solution. Tetrafluoroquinone also gives two waves in dilute antimony pentafluoride solution (<0.4 M SbFs in fluorosulfuric acid or <0.2 M in trifluoromethanesulfonic acid; see Fig. 3, curve 4 and Fig. 4, curve 2) and two waves in more concentrated antimony pentafluoride solution (e.g., Fig. 3, curve 5). By analogy with the results obtained in hydrofluoric acid where the same kind of behavior has been observed [23], the twowave system indicates two one-electron steps corresponding to the semiquinone stabilization. The two reduction steps of the quinone correspond to the redox equilibria, QW + H’ * QHj+ and QHj’ + e- =+QH:. Voltammograms of a 1:l mixture of tetrafluoroquinone and tetrafluoro-
154 TABLE 2 Half-wave reduction potential shift of tetrachloroquinone (TCQ) and tetrafluoroquinone (TFQ) with base concentration (C) in fluorosulfuric acid and trifluoromethanesulfonic acid (TFMS) 10e4 M < C < 10-l M) (&I, = a log C + 5; C = H,O or F-concentration Fluorosulfuric
TCQ
LmV) (mV)
TFQ
57 f 38
a
Trifluoromethane sulfonic acid 6Oi 2a
516 f 5
291 f 4
63 f 3
68 * 3
326 i 5
103 f 3
II LmV) (mV)
acid
aConfidence limits, probability:
95%.
hydroquinone were obtained in trifluoromethanesulfonic acid (Fig. 4); the equilibrium potential of a platinum electrode at which the system is reversible is directly linked to the pH of the solution by E = E" -0.06 pH. The voltammograms are well defined in basic and in neutral media (curves) but distortion appears in acidic media, probably because of the oxidizing power of antimony(V) fluoride.
Fig. 4. Voltammetric curves of a mixture of 5 mM tetrafluoroquinone and 5 mM tetrafluorohydroquinone in trifluoromethanesulfonic acid at various acidity levels: (1) < lOa MH,0;(2)0.13MSbF,;(3)0.4MSbF,;(4)0.7MH,0;(5)5MH,O.
155 RESULTS AND DISCUSSION
De termination of the Ro (H) acidity functions The evaluation of the reduction half-wave potential of quinones (or equilibrium potential of a mixture of the quinone and hydroquinone) allows calculation of the&(H) acidity function. The expression for R(H) is R(H) = F(2.3 BT)-’
(&n(ao)
-h!d
w&m J-b caq) refers to the potential of the pH indicator in water and El12(,, refers to the potential of the same pH indicator in the solvent, both potentials being expressed by reference to the half-wave potential of a reversible system supposed to be independent of the solvent [ 161. The solvent-independent system is octafluoronaphthalene (OFN/OFN+) which was previously studied in hydrofluoric acid [28]. Because this compound is insoluble in water, its half-wave potential must be evaluated by using an intermediate compound which is soluble in the solvent and in water. Perylene/perylene+ can be used for this purpose. Tetrachloroqumone was used to evaluate theR(H) function of the strongly basic solution. The reduction of this quinone proceeds by the same scheme in water and in TFMS or fluorosulfuric acid solution. The El,2 values for the reduction of tetrachloroquinone in aqueous solution are as follows: 0.260 V vs. ferrocene (Fe/Fe+), 0.840 V vs. perylene (Pe’/Pe’+) and 1.390 V vs. octafluoronaphthalene (OFN/OFN+). The corresponding calculated values of I&,(H) in sodium fluoride or water solutions are reported in Table 3. These values correlate with the R,(H) function obtained with the same quinone in hydrofluoric acid, which wasR,(H) = 14.2 (1 M KF) [22,29]. The tetrafluoroquinone reduction presents the same one-step, twoelectron reversible reduction in strongly basic solution and in dilute strong acid solution (SbF5 < 0.4 M in fluorosulfuric acid and <0.2 M in TFMS).
TABLE 3 R,(H) acidity function of basic (Hz0 or NaF) and acidic (SbF,) media in fiuorosulfuric and trifluoromethanesulfonic acid. octafluoronaphthaiene (OFN/OFN+). Acidity indicators: tetraPotential reference: chloroquinone (TCQ) and tetrafluoroquinone (TFQ) Medium
J%
HSO,F + 1 MNaF HSO,F + 1 M SbF, HSO,CF, + 1 M H,O HSO,CF, + 1 M SbF,
0.360 0.330 0.694
TCQ TFQ TFQ
-17.2 -17.2 -23.4
0.430 0.210 0.572
TCQ TFQ TFQ
-16 -16 -22.1
0’)
Indicator
R,(H)
156
Thus, the same system can be used to calculate the acidity function of the strongest acid solution. Extrapolated values (corresponding to a 1 M SbF, solution) are reported in Table 3. Calculation 0 f apparent au topro tolysis constants According to the autoprotolysis equilibria (Eqns. 1 and 2) reported above, the evaluation of the acidity function of a strong acid and a strong base (NaF or H@) with the same electrochemical pH indicator leads to the determination of the autoprotolysis constant, pK1. The values obtained with tetrafluoroquinone are pKi = R,(H) (1 M SbFS) - R,(H) (1 M NaF) = 6.2 k 0.5 for fluorosulfuric acid, if SbFs is considered as a strong acid as in hydrofluoric acid, and pKI = R,(H) (1 M SbFs) --R,(H) (1 M HzO) = 6.1 f 0.5 for trifluoromethanesulfonic acid. The value for pKi in fluorosulfuric acid is one unit lower than that obtained by Barr et al. [ 171 by conductimetric measurements because they have considered SbFS as only partially ionized (pK = 2.15). There is no publishedvalue in the case of trifluoromethanesulfonic acid. Comparison between acidity ranges Figure 5 shows the three fluorinated acid solvents, hydrofluoric, fluorosulfuric and trifluoromethanesulfonic acids, related to the R0 scale. The correspondence with the I& scale is represented, and the results confirm the 6-unit deviation between R0 and Ho functions already observed in
-% - _ _ _ ____
IF --
16.:
11.5 10.3
_
---
-_
-
-
--
-
-_-__
..
22
172
- -__.
-
HSosCF,
8.1
2L.6
----r .
-----A-
- -_-.
229
_ _ _..
16
IL.2
HF Fig. 5. Relative position of the acidity scales of fluorosulfuric, and hydrofluoric acids, and the relationship between the R,(H)
trifluoromethanesulfonic and H, acidity functions.
157
hydrofluoric acid [14]. It can be seen that hydrofluoric acid has the widest acidity range, which enables the most acidic and the most basic media to be reached (p.& = 13.7). Conclusion and tetrafluoroquinone As in hydrogen fluoride, tetrachloroquinone appear to be suitable pH indicators; they are sufficiently stable and their protonation allows the pH to be evaluated in the whole acidity range of the solvents. These quinones are the two electrochemical pH indicators useful in very acidic media. Comparison between the values determined for RO and HO shows that the difference between these two functions is almost constant. The difference found is the same as in hydrofluoric acid. The pK, values determined for fluorosulfuric and trifluoromethanesulfonic acid through RO measurements are of the same magnitude. The accessible pH range covers 6 units. Thus superacid media based on these two acids makes it possible to reach very acidic media by addition of antimony(V) fluoride (R,, = -25), but the acidity range is not as wide as that of hydrofluoric acid. The presence of S(VI), from the sulfur trioxide produced secondarily to their internal dissociation, differentiates these two acids from hydrofluoric acid. Sulfur trioxide is a potent oxidizing agent which participates directly in oxidations. It oxidizes hydrogen so that it becomes impossible, in practice, to use a hydrogen electrode in sulfur-containing superacids, as in sulfuric acid. REFERENCES 1 2 3 4
G. A. Olah and J. Lukas, J. Am. Chem. Sot., 89 (1967) 2227. H. Hogeveen and J. Gaasbeek, Rec. Trav. Chhn., 88 (1969) 703. G. A. Olah, Y. Haipern, J. Shen and Y. K. MO, J. Am. Chem. Sot., 95 (1973) 4960. J. W. Larsen, P. A. Bovis, C. R. Watson and R. M. Pagni, J. Am. Chem. Sot., 96 (1974) 2284. 5 C. Pitti, F. Bobillart, A. Thiebault and M. Herlem, Anal. Lett., 8 (1975) 241. 6 M. Herlem, F. Bobillart, A. Thiebault and A. Jobert, Anal. Lett., 10 (1978) 767. 7 C. Pitti, M. Cerles, A. Thiebault and M. Herlem, J. Electroanal. Chem., 126 (1981) 163. 8 G. M. Kramer, J. Org. Chem., 40 (1975) 5173. 9 D. Brunel, A. Germain and A. Commeyras, Nouv. J. Chim., 2 (1978) 275. 10 V. B. Hughes, B. D. McNicol, M. R. Andrew, R. B. Jones and R. T. Short, J. Appl. Electrochem., 7 (1977) 161. 11 A. A. Adams and R. T. Foley, J. Electrochem. Sot., 126 (1979) 775. 12 R. J. Gillespie and T. E. Peel, J. Am. Chem. Sot., 95 (1973) 5173. 13 H. H. Hyman, M. Kiipatrick and J. J. Katz, J. Am. Chem. Sot., 79 (1957) 3668. 14 J. Grondin, R. Sagnes and A. Commeyras, Bull. Sot. Chim. Fr., 11 (1976) 1779. 15 P. L. Fabre, J. Devynck and B. Tremiilon, Chem. Rev., 82 (1982) 591. 16 H. Strehlow and H. Wendt, Z. Phys. Chem. (Frankfurt am Main), 30 (1960) 141. 17 J. Barr, R. J. Gillespie and R. C. Thompson, Inorg. Chem., 3 (1964) 1149. 18 G. Adhami, These de doctorat d’&at, Paris VI (1972). 19 J. Sommer, P. Canivet, S. Schwartz and P. Rimmelin, Nouv. J. Chhn., 5 (1981) 45.
158 20 D. C. Russel and J. D. Senior, Can. J. Chem., 52 (1974) 2975. 21 J. Devynck, A. Ben Hadid, P. L. Fabre and B. TrBmillon, Anal. Chim. Acta, 100 (1978) 343. 22 J. P. Masson, J. Devynck and B. Trdmillon, J. Electroanal. Chem., 64 (1975) 193. 23 A. Ben Hadid, P. Rimmelin, J. Sommer and J. Devynck, J. Chem. Sot. Perkin Trans. 2, (1982) 269. 24 J. Verastegui, These 38me cycle, Universitd Pierre et Marie Curie, Paris VI (1973). 25 A. Jobert, These docteur ing&ieur, Universite Pierre et Marie Curie, Paris VI (1979). 26 D. Bauer and M. Bouchet, C. R. Acad. Sci., Sbr. C, (1972) 274. 27 D. Bauer, J. P. Beck and P. Texier, Collect. Czech. Chem. Commun., 36 (1971) 360. 28 A. Ben Hadid, Thbse d’Qtat, Universite Pierre et Marie Curie, Paris VI (1980). 29 J. P. Masson, J. Devynck and B. Tremillon, J. Electroanal. Chem., 64 (1975) 175.