Reagents for High Temperature Aqueous Chemistry: Trifluoromethanesulfonic Acid and its Salts

LEONARD FABES A N D THOMAS WILSON SWADDLE' Department of Chemistry, The University of Calgary, Calgary, Alberta Z'N IN4

Received June 13. 1975

LEONARD FABES and THOMAS WILSON SWADDLE. Can. J. Chem. 53,3053 (1975). The rate of decomposition of trifluoromethanesulfonic acid (HTFMS) in water is first order

with respect to each of H + and TFMS-. The bimolecular rate coefficient kH can be represented by

kH = 1.7 x 1014exp (2.15 x 105/8.314T)kg mol-' s-I

at ionic strength 1.0 mol kg-I over the temperature range 571 < T < 593 K, although the decomposition occurs by two pathways of comparable importance, one leading to C(I1) and S(V1) and the other to CUV) and SUV). In alkaline solution, the rate of decomposition of TFMS- is first order in each of TFMS- and OH-, and the bimolecular rate coefficients koH at ionic strength 2.3 are given by

koH = 4.1 x lo8 exp (1.46 x 105/8.314T) kg mol-' s-'

This represents a single reaction pathway leading initially to C(I1) and S(V1) and ultimately to C0,2-, Sod2-, F- , and Hz, since C(I1) (as formate) reduces aqueous alkali to HZ under the reaction conditions. No de~o&~osition of NaTFMS could be detected after 24 h at 620 K in neutral aaueous solution. HTFMS reduces certain aaueous s~ecies at significant rates at mod- erate temperatures (e.g., FeUII) to Fe(I1) at 470 K), and metais which depend on oxide films for their corrosion resistance are attacked by HTFMS even at 294 K (e.g., Ti dissolves giving Ti(H20)63+). Otherwise, HTFMS and .its salts have good potentialities as inert electrolytes for aqueous studies at high temperatures.

LEONARD FABES et THOMAS WILSON SWADDLE. Can. J. Chem. 53,3053 (1975). La vitesse de d6composition de l'acide trifluoromethane sulfonique (HTFMS) dans I'eau est

du premier ordre par rapport a H + et par rapport a TFMS-. Le coefficient de vitesse bimolecu- laire, kH, peut &re represent6 par 1'Bquation

kH = 1.7 x 1014 exp (2.15 x 105/8.314T) kg mol-' s-'

a une force ionique de 1.0 mol kg-', a des tempkatures allant de 571 a 593 K meme si la decomposition se produit suivant deux voies d'dgale importance, une conduisant a C(I1) et S(V1) et I'autre a C(IV) et S(1V). En solution alcaline, la vitesse de dckomposition du TFMS- est du premier ordre en TFMS- et du premier ordre en OH- et le coefficient de vitesse bimolkculaire, koH, a force ionique 2.3, est donne par 1'6quation

ko, = 4.1 x lo8 exp (1.46 x 105/8.314T) kg mol-' s-'

Ceci represente une seule voie pour la reaction qui conduit originalement a C(I1) et SWI) et ulterieurement a C03'-, SO4'-, F - et Hz puisque C(I1) (SOUS forme de formate) reduit, dans les conditions de la reaction, une solution aqueuse d'alkali en Hz. A 620 K, en solution aqueuse neutre, on ne peut pas noter de d6composition du NaTFMS apres 24 h. A des temperatures moderkes, I'acide HTFMS rkduit a des vitesses significatives certaines especes en solutions aqueuses (par exemple Fe(II1) en Fe(I1) a 470 K); les metaux qui dkpendent sur la formation de films d'oxyde pour devenir resistants a la corrosion sont aussi attaques par l'acide HTFMS meme A 294 K (par exemple le Ti se dissout pour donnei du Ti(Hz0)63+). Autrement il semble que I'acide HTFMS et ses sels prksentent de bonnes potentialitks comme Blectrolytes inertes pour des etudes aqueuses a hautes temperatures.

[Traduit par le journal]

Introduction coefficients, etc.) which govern the chemistry of Measurements of the parameters (acid dis- aqueous electrolytes can, in principle, be made on

sociation constants, stability constants, rate pressurized solutions far above the normal boiling point of water, much as at room tem-

'Author to whom correspondence should be addressed. perature, if a thermally stable, strongly ionized,

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3054 CAN. J. CHEM. VOL. 53, 1975

poorly complexing, 'inert' electrolyte is available. The traditional electrolyte, perchloric acid, is remarkably stable kinetically towards spon- taneous thermal decomposition in dilute aqueous solution, largely as a result of a high activation energy combined with a high reaction order (I), but recent work by T. C. T. Wong in our labora- tories has shown that this decomposition can be catalyzed by aqueous metal ions. Thus, aqueous HClO, (1.0 m, i.e. 1.0 mol kg-') containing 0.01 m Fe(II1) yields 0.01 m chloride ion after 28 h at 473 K in Teflon lined autoclaves, whereas in the absence of Fe(II1) the acid would decompose only to the extent of 3 x lo-' m (1). Furthermore, perchloric acid solutions can oxidize many species of interest directly as one goes toward high temperatures and kinetic bar- riers are removed (2, 3).

A nonoxidizing alternative to perchloric acid is trifluoromethanesulfonic acid, HTFMS ('trims- ylate' (4) or 'triflic acid'). The discoverers of this acid (5) recognized its remarkable resistance to thermal decomposition and hydrolysis, and also that it is one of the strongest protonic acids known, as has subsequently been amply verified (6-13). No universal series of relative acid strength exists, since the protonating powers of various acidic solutes vary from solvent to solvent (13), but it seems safe to conclude that HTFMS will be at least as fully ionized in water at high temperatures as is HClO,, which cannot be more than slightly associated in water up to 600 K (1). The TFMS- ion is also a very weak complexing agent, being only slightly more effective in complexing aqueous Cr(II1) than is perchlorate ion (14), although TFMS- com- plexes with UO2' (15) and Mop ' (16) have been isolated. Thus, TFMS- offers an excellent alternative to C10,-, BF4-, PF,-, etc., espe- cially when reagents which are potential oxidants or are solvolytically unstable must be avoided (17-19).

The aqueous TFMS- ion has been reported to be stable in acidic solution up to at least 550 K (20) and in neutral or alkaline solution up to 620 K (21). This stability, however, is kinetic rather than thermodynamic, and we have there- fore sought to define it in terms of the appro- priate rate equations and parameters for the various hydrolysis pathways as functions of p H and temperature, thereby permitting upper limits to be set to the extent of spontaneous decom- position of the reagent in typical experimental

situations, as well as gaining information on the reaction mechanisms.

Experimental Anhydrous HTFMS, as purchased from the 3M Com-

pany, was black, presumably with organic char (20), but a simple distillation under anhydrous conditions at reduced pressure was sufficient to give a product which analyzed correctly for H + , contained no detectable HF, and gave the same kinetic results from batch to batch. A further careful fractional distillation of this colorless liquid led to no detectable change.

Kinetic Measurements The pressure vessels, previously described (I), were of

titanium - 0.2% palladium alloy or (with basic samples only) of Type 316 stainless steel, and were protected from corrosion by the samples with silica liners (acidic solu- tions only) or Teflon liners; the silica liners were slightly attacked, presumably by the H F produced in the decom- position of HTFMS, but the Teflon vessels appeared to be completely inert to both acidic and alkaline solutions. The free volume in the autoclaves was 47cm3. The TFMS- solutions were made up on a volume basis, but concentrations were expressed on the molal (mol/kg) scale using the measured densities and corrected for the evaporative loss entailed in generating the saturated vapor pressure of the solution (taken to be that for pure water) at the reaction temperatures. The derived kinetic results were the same whether a 10 or a 25 cm3 aliquot of the TFMS- solution was placed in the autoclave.

The solution samples were sealed into the autoclaves under an air or a nitrogen atmosphere at 0.1 MPa (the results were shown to be independent of the presence or absence of oxygen) and brought to the desired tempera- ture in a preheated Blue M Conwate CW-160HF-1 forced-convection oven (constancy k0.3 K during a run). It was established, by extrapolating reactant con- centrations back to their original values as functions of time, that the effective zero time of the reactions was the same at which the oven temperature (measured as de- scribed previously (1)) was seen to be stabilized, which was 40 min after placing the vessels in the oven. In general, conditions were chosen to give slow reaction rates, so as to minimize any errors associated with sample warm-up. Autoclaves were removed as required and quenched rapidly to room temperature with water.

The TFMS- content of the solutions could not easily be determined directly with good precision, but it was found that an Orion 92-81 perchlorate ion selective elec- trode, used in conjunction with an Orion 801 p H meter and 90-02 double-junction reference electrode, gave a linear response to TFMS- when the p H was adjusted to about 5 with sodium acetate, and could be used to measure [TFMS-] in the range 0.0024.05 m without significant interference from the decomposition products F- , HC03-, HS03-, or SO4'-. The Gran method (22) enabled TFMS- concentrations to be measured with a standard deviation of 7%, which, though poor, was acceptable for measurement of the rate of disappearance of TFMS- in acidic solutions, since the usual alternative of following the formation of the reaction products was not feasible; the S and C products were distributed amongst various oxidation states, while the F product,

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FABES AND SWADDLE: TRIFI JJOROMETHANESULFONIC ACID 3055

present as volatile HF, was partially depleted by reaction with the autoclave walls or possibly by vapor loss on opening the pressure vessels.

The fluoride contents of the. solutions were measured with about 1% precision using an Orion 94-09 ion selec- tive electrode, as previously described (23). This method was effective in measuring the rate of hydrolysis of TFMS- in alkaline solution.

Reaction Products Dissolved fluoride was measured with the Orion 94-09

electrode. Sulfate was measured gravimetrically as BaSO,. Sulfite was estimated by redox titration. Car- bonate was measured by a p H titration procedure, which also gave rough indications of the F- and S032- con- tents through inflexions in the titration curves at lower p H than for C03'-.

The gaseous decomposition products of aqueous HTFMS were determined by mass spectroscopy, the samples of the aqueous acid having been sealed into quartz ampoules which were heated in autoclaves, partially filled with water to pressurize the tubes exter- nally. This procedure was not followed with alkaline TFMS- solutions, as these would have attacked any available material other than Teflon, which would have been too soft to form a gastight ampoule at high tem- peratures. Nevertheless, when NaTFMS was decomposed in 9 m NaOH for 19 h in a Teflon lined autoclave at 572 K, a sample drawn from the gas space of the cooled autoclave immediately on opening was found by mass spectrometry to contain much hydrogen gas; in the same way, sodium formate was shown to yield HZ under these conditions.

Results Reaction Products

The gaseous products of the partial thermal decomposition of aqueous HTFMS were pri- marily CO,, CO, and SO,, with traces of H,S, F,CO, CS,, and some trivial components with unassignable mass numbers. The recorded 44 : 28 (CO, : CO) mass ratio rose from 1.6 for samples decomposed at 563 K to 2.9 for 593 K, as com- pared with 5.7 for pure CO, (CO,' undergoes partial f~agmentation to CO+); these data indi- cate that two parallel paths of comparable im- portance were operating in the decomposition of HTFMS in acidic solution, one of relatively low activation energy Ea leading to CO and one of higher Ea leading to CO,. Some elemental sulfur was recovered, especially at the higher temperatures, so that this product may be associated with the CO, producing reaction. The aqueous phase contained HF, HC0,-, S 0 2 - , and HS0,-. It was shown in separate experi- ments that S(IV) disproportionates to S(0) and S(V1) in aqueous acid at high temperatures (e.g., in 3 m HCl or H,SO, at 584 K over 19 h), thus accounting for the presence of elemental sulfur

in the decomposition products of HTFMS. The disproportionation of S(1V) is thermodynam- ically possible in aqueous systems at all p H values even at 298 K (24), but mechanistic limitations evidently prevent its occurrence except at high temperatures ( - 570 K) and high acidities, on the usual experimental time scale. The dispro- portionation reaction was insufficient to account for all the S(V1) produced in the acidic decom- position of HTFMS, and in any event the decomposition path leading to C(I1) (carbon monoxide) must produce S(V1) directly. The main reactions in the decomposition of TFMS- in acidic solutions are therefore 1, 2a, and 2b.

[l] CF3S03- + 2Hz0 = 3HF + HS04- + CO

In alkaline solutions at about 570 K, TFMS- was hydrolyzed quantitatively to 3F-, SO,'-, and CO,'-, and it was shown qualitatively that Hz was a major product. It was also established that SO,'- quantitatively survives heating at about 580 K in 1 m NaOH for several days, so that no pathway analogous to reaction 2 could have been operative in the alkaline decom- position of TFMS-. In alkaline solution, a S(V1) producing reaction analogous to reaction 1 would yield formate rather than CO (ref. 24, p. 278), and formate ion is thermodynamically capable of reducing alkaline water to H, even at 298 K (25) :

[3] HCOz- + OH- = Hz + CO3'- AGa = -35 kJ mol-I

Thus, the final products of an alkaline hydrol- ysis of TFMS- producing initially S(V1) and C(I1) should be F-, SO2-, CO,'-, and Hz, if reaction 3 is sufficiently rapid at the reaction temperature, and this is indeed what is observed.

Exploratory measurements showed that the rate of decomposition of formate ion in 1.9 m NaOH at 586 K, as followed by KMnO, titra- tion, was very close to that calculated by Arrhenius extrapolation (see below) for the dis- appearance of TFMS- under the same con- ditions, so that a substantial buildup of formate ion during the alkaline hydrolysis of TFMS- was neither expected nor observed.

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TABLE 1. Initial rates of decomposition of trifluoromethane sulfonate ion in acidic aqueous solution

d Nominal Actual Initial Mean Initial Ionic - z[CF,SO,-l

temperature temperature [CF,SO, -1 [CF,SO,-l [H+l strength x lo6 kH x lo6 (K) (K) (mol kg-') (mol kg-') (mol kg-') (mol kg-') (mol kg-' s-') (kg mol-' s-I)

'HTFMS + NaTFMS.

TABLE 2. Initial rates of decomposition of trifluoromethane sulfonate ion in alkaline aqueous solution

Nominal Actual Initial Initial Ionic d[CF3S03-1 106 temperature temperature [CF,SO,-l [OH-] strength - dt koH x lo6

(K) (K) (mol kg-') (mol kg-') (mol kg-') (mol kg-' s-I) (kg mol-' s-')

'Ionic strength maintained with NaCI.

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FABES AND SWADDLE: TRIFLUOROMETHANESULFONIC ACID 3057

Kinetics Because the formal reaction 2a (at least) con-

sumes H f whereas reaction 1 does not, the time dependence of the HTFMS concentration could not be expected to follow a simple integrated rate expression over a substantial fraction of the decomposition reaction. Instead, the concentra- tion of TFMS- was monitored over the first 20% of the reaction in acidic solution, and the slope of the crudely linear plot of [TFMS-] against time t was taken to be -d[TFMSP]/dt at the mean [TFMS-1. These data are collected in Table 1. The reproducibility of these rates was as poor as + 14%, but this is as good as could be expected in view of the 7% uncertainty in the [TFMS-] analyses and the difficulty of presetting the nominal reaction temperatures accurately.

Log-log plots of the rate data of Table 1 at a given nominal reaction temperature against stoi- chiometric [HTFMS] when [H + ] = [TFMS-] were linear, within the experimental uncertainty, with a slope of about 1.8, which suggests a second-order reaction with a retarding ionic strength effect (cf. the decomposition of aqueous HClO,, which is actually 4th order with an apparent reaction order of 3.5). A plot of log(rate) us. log [H+] for 583 K and nearly constant ionic strength (2.6-2.7 m) had slope -0.9, showing the reaction to be essentially first order with respect to each of H f and TFMS-, and accor- dingly bimolecular rate coefficients k, are given in Table 1. At the nominal temperatures 582.6 and 592.9 K, k, values generally decreased as the ionic strength I increased, and can be repre- sented to within the rather large experimental error by equations 5 and 6.

For I = l.Om, we have k, - 4.3 x 8.5 x and 2.1 x lo-' m-' s-' at 571.6, 582.6, and 592.9 K respectively. These data give a somewhat curved Arrhenius plot (concave upwards, as expected for two parallel reactions having the same rate law) but can be reproduced satisfactorily by an activation energy of 215 kJ mol-' and a preexponential factor of 1.7 x lo1, ,,,-I s - ~ , the uncertainty in these parameters

being about 20%. The rate of disappearance of TFMS- in

alkaline solution is much more readily measured

as one-third of the rate of formation of F-, low concentrations of which were determined with 1% precision. This permitted rate measurements to be made with OH- in large excess over TFMS-, and the initial rates of TFMS- col- lected in Table 2 were reproducible to 5%, most of this margin representing the difficulty of repro- ducing the reaction temperature from one run to another. Examination of the data OF Table 2 shows that the reaction is accurately first order in each of OH- and TFMS- at a given ionic strength, and the corresponding bimolecular rate coefficients k,, listed for 552 K, I = 2.2 + 0.1 m are obviously constant within 5% standard deviation. The data for I = 2.28 m as a function of temperature gave a good Arrhenius plot with E, = 146 kJ mol-' and a preexponential factor of 4.1 x 10' m-' s-'. The alkaline hydrolysis rate is evidently accelerated by increasing the ionic strength; this is opposite to the ionic strength effect on the acidic hydrolysis reaction, as expected for reactions of opposite charge types.

Neutral solutions of NaTFMS (1 m) under- went no detectable decomposition, even at 620 K over 24 h.

Reaction of HTFMS with Metals and Metal Ions

At 294 K, type 316 stainless steel tubing was visibly etched within one day by deaerated HTFMS (1.0 m), and commercial titanium dis- solved in it to give a purple solution of Ti(H20)63 + over a few hours; Ti - 0.2% Pd alloy was more resistant, but nevertheless gave an appreciably yellow-orange solution over 24 h. By contrast, these metals were not detectably attacked by 1 m HClO, or H2S04 over several days.

The divalent metal ions Mn, Fe, Co, Ni, and Cu were unchanged after 24 h in 1 .O m HTFMS at 570-580 K. Further studies in our laboratories by T. C. T. Wong, however, have shown that HTFMS reduces aqueous iron(II1) to iron(I1) under relatively mild conditions (e.g., 0.05 m Fe(II1) in 1.5 m HTFMS is reduced to the extent of at least 7% in 24 h at 473 K, with the con- comitant formation of HF).

Discussion The rate determining step in the alkaline

hydrolysis of TFMS- evidently involves nucleo- philic attack of OH- on the S atom of TFMS-

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3058 CAN. J. CHEM. VOL. 53, 1975

The actual initial products of this step may be S 0 2 - and HCF,, or SO:-, CF,, and F- , but in any event the fluorocarbon moiety will hydrolyze to HC0,- in alkali (26). We have demonstrated that formate ion is oxidized to carbonate with H, evolution at a rate sufficiently rapid at the temperatures involved here to to account for the observed reaction products, sulfate, carbonate, fluoride, and hydrogen.

In acidic solution, the reactions 1 and 2a occur simultaneously, with reaction 1 accounting for some 33% of the decomposition rate at 563 K and about 1697, at 593 K, on the basis of the mass spectroscopic yields of CO and CO,, corrected for the difference in the ionization cross-sections of these molecules (27). Both acidic decom- position pathways are evidently first order in each of TFMS- and H + , if ionic strength effects are allowed for, and accordingly, neutral solu- tions of TFMS- were found not to hydrolyze significantly even at temperatures approaching the critical point of water. The decomposition mechanisms presumably involve the undissociated HTFMS molecule (in which case HTFMS must still be a fairly strong acid in water at 600 K. in order to give the observed reaction orders).

Assuming, for argument's sake, a bimolecular mechanism for the hydrolysis of HTFMS, nucleophilic attack of water on the S atom of HTFMS will be seen to give S(V1) and C(II), much as in equation 7, whereas attack of the water molecule on the C atom will yield C(1V) and S(IV),

the postulated trifluoromethanol intermediate hydrolyzing rapidly to carbonic acid and HF.

Whether the decomposition of TFMS- pro- ceeds via reaction 1, 2a, or 4, the potential reducing power of this anion is clearly seen, since a familiar reducing agent (COY SO,, H,) is pro- duced in each case. This reducing power can be exercised directly at relatively low temperatures. Thus, Fe(II1) is reduced to Fe(I1) in acidic aqueous solution around 470 K, and metals (e.g., titanium) or alloys (e.g., stainless steels) which are corrosion resistant by virtue of having a tenacious oxide film are attacked quite rapidly by aqueous HTFMS even at 290 K, presumably because of reduction of the oxide coat. The latter

property has heen utilized to activate passive metals prior to chromium plating (28). These considerations limit the usefulness of HTFMS as a high-temperature aqueous reagent, but nevertheless it can be of great value in situations where a strong electrolyte of good thermal stability is required and reducing conditions can be tolerated.

Finally, this study serves to illustrate the fact that much familiar 'room temperature chemistry' exists only by virtue of high kinetic barriers to the realization of thermodynamic exigency. The resistance of TFMS- to hydrolysis is due in part to the high activation energies of the various hydrolysis pathways, and in part to the need to involve either H + or OH- in the decomposition mechanisms. Similarly, SO, should dispropor- tionate in water, and formate ion (i.e., CO) should reduce aqueous alkali to hydrogen, at 298 K ; these reactions do occur on a relatively short time-scale at 500-600 K, where many kinetic barriers are overcome.

We thank J. W. Cobble for first drawing our attention to the potential uses of HTFMS, and Atomic Energy of Canada, Ltd., for financial support of this work.

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