Kinetics of the reduction of the tetrahydroxyargentate(III) ion by sulphite in aqueous alkaline media

Polyhedron Vol. 8, No. 10, pp. 1299-1305, Printed in Great Britain

1989 0

KINETICS OF THE REDUCTION OF THE TETRAHYDROXYARGENTATE(IIl) ION BY SULPHITE AQUEOUS ALKALINE MEDIA LOUIS

J. KIRSCHENBAUM*

and ISSIEOU

0277-5387/89 $3.00 + .I0 1989 Pergamon Press plc

IN

KOUADIO

Department of Chemistry, University of Rhode Island, Kingston, RI 0288 1, U.S.A. and

EDOARDO

MENTASTI

Department of Analytical Chemistry, University of Torino, 10125 Torino, Italy (Received 25 October 1988 ; accepted after revision 18 January 1989) Abstract-The

reduction of Ag(OH); by sulphite ion to give silver(I) and sulphate in strong base has been studied by stopped-flow spectrophotometry. The reaction was found to fit a second-order rate law with a rate constant k = (9.1 f 0.1) x 10’ M- ’ s- ’ (T = 25°C p = 1.2 M). The estimated enthalpy and entropy of activation are 27 kJ mol- ’ and - 153 J K-’ mol- ’ respectively. Variation of rate with ionic strength is consistent with the interaction between mononegative Ag(OH); and dinegative SO:-. The redox reaction appears to proceed by a one-step oxygen atom transfer. Two kinetically indistinguishable pathways are discussed. The first involves initial attack of a sulphite oxygen at an axial silver(II1) site. The second pathway involves initial interaction between sulphur(IV) and a bound hydroxyl. Sulphite oxidation by silver(II1) is about lo4 times faster than that of phosphite, while no reaction is observed with NO;.

Silver(II1) is a spin paired d8 system like gold(III), copper(III), nickel(I1) and palladium(I1). ‘3’It forms diamagnetic square-planar complexes with ligands including hydroxide,3 triglycine, tetraglycine,4 dimethylglyoximate, ’ tellurate and periodate. Although octahedral geometry is rare, it is found in paramagnetic complexes with the general formula M3AgF67 (M = K, Na, Rb or Cs). Diamagnetic tetragonal silver(II1) fluoro complexes (MAgF4) have been successfully prepared by direct fluorination of 1: 1 mixtures of MC1 and AgN03.8 Pale yellow, diamagnetic silver(II1) complexes of tetraazamacrocyclic substrates have been extensively studied. ’ Silver(II1) porphyrin complexes have also been investigated. For instance, Kadish et al. found that oxidation of the octaethylporphynatosilver(II) leads to the corresponding silver(II1) complex.‘o Ethylene-bridged biguanide complexes of silver(III) were first prepared in 1944 by Chakravorty and Ray” as Ag(enBigH)zX3 (X = 0.5S04, N03, C104 *Author to whom correspondent

should be addressed.

or H). A thorough investigation of the kinetics of formation and decomposition of these complexes has been reported. I2 Simple, stable silver(II1) biguanide and methyl biguanide complexes also have been synthesized.‘3 Silver(II1) complexes of triglycine and tetraglycine were successfully prepared by Kirschenbaum and Rush4 by direct mixing of Ag(OH); and the ligands in aqueous basic media. There are striking spectral and kinetic similarities of these complexes with those of analogous copper(II1) polypeptide systems obtained by Margerum et a1.14 Ag(OH); has a half life of about 1.5 h in 1.2 M hydroxide at room temperature and has a UV-vis maximum at 267 nm in aqueous alkaline media. In recent years, we have studied the reaction of Ag(OH), with a number of complexing and reducing substrates.S.‘~‘7 Substitution reactions of Ag(OH): with various oxygen donor ligandeperiodate, tellurate, 3 phosphate, borate, pyrophosphate, carbonate and arsenateL8-have been investigated. In the case of

1299

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L. J. KIRSCHENBAUM

periodate and tellurate ions, the kinetically stable complexes Ag(H,IO,):and Ag(H,TeO,);, respectively, were formed. However, the substitution reactions of Ag(OH), with the other oxoanions were partial and occurred only at reduced OH-. The formation of silver(II1) complexes often provides a path for Ag(OH); reduction,‘6,‘9 but offers some increased stability under conditions where the highly reactive Ag(OH),(H,O) ion is otherwise formed [eq. (l)]‘” Ag(OH);

s

Ag(OH),(H,O)+OH-.

(1)

Redox reactions of the tetrahydroxyargentate(II1) ion have been studied with reductants including ethylenediamine (en),” H202,22 Mo(CN)j-, Fe(CN)z-, Mn0$-,‘9 N- l6 S 02- l7 and arsenites. ’ ’ It was found that an3mnerfspiere two-electron transfer process was involved in the cases of azide, thiosulphate, en, and arsenite. By contrast, a silver(I1) intermediate was inferred for the other substrates. The reduction of transition metals by sulphite in aqueous alkaline medium has been studied for Mo(C!N)-, Mn(P207)!-, Fe(CN)i- 23,24and deprotonati copper tetraglycine (Cu(I-_ 3G4)-).25 Margerum et a1.25 recently showed that the kinetics of the reaction of CU(H_~G~)- with excess sulphite is first-order in sulphite at pH 9.2. In addition, the reduction reaction goes stoichiometrically (ACU(H_~G~)- : ASO:- = 2 : 1) via two reversible one-electron transfer steps, generating a sulphite radical anion in the first step. Brown and Higginson showed substitution-inert that complexes which behave as one-equivalent oxidants can give quantitative sulphate products. However, those substitution labile complexes which react with a lower non-integral stoichiometry, give a mixture of sulphate and dithionate. In the particular case of the ferricyanide-sulphite reaction, Swinehart suggested an outer-sphere mechanism involving sulphite radicals and dithionate and sulphate as major products. On the other hand, Veprek-Siska and Wagnerova favoured an inner-sphere process with sulphate the only significant sulphur-containing product. Lancaster and Murray, 24 have shown that the reaction proceeds via the intermediates Fe(CN)S(CNS03)5and Fe(CN)S(CNS03)4-, followed by hydrolysis of the latter to ferrocyanide and sulphate. They thought that the reaction pattern lies somewhere between inner- and outer-sphere, a compromise between the two contradictory explanations mentioned above. The reaction of Ag(OH); with thiosulphate at low concentrations ( < 1.Ox lo- 3 M) proceeds via the aquated species Ag(OH),(H,O), giving a mono-

et al.

thiosulphato complex, Ag(OH),S,O:-, which undergoes internal redox. I7 At thiosulphate concentrations above 1 x lo- 3 M, silver(II1) is reduced in a five-coordinate intermediate without hydroxyl replacement. In the case of the reaction of Ag(OH); with arsenite, the reaction appears to proceed by a one-step oxygen atom transfer. ’ 5We have studied the reaction of Ag(OH); with sulphite ion to see if its reduction mechanism is similar to either the thiosulphate or arsenite systems. EXPERIMENTAL Reactant solutions For all experiments, sulphite stock solutions were freshly made in doubly distilled water, or in 1.2 M base by exact weighing of reagent grade anhydrous sodium sulphite (J. T. Baker). All sulphite reactant solutions were then made by serial dilution. Sodium hydroxide solutions were prepared from 50% low carbonate NaOH (Fisher). Lithium hydroxide solutions were made from solid hydrated LiOH (Matheson Coleman and Bel). A 1.2 M NaClO, or LiC104 solution was used for ionic strength adjustment when appropriate. NaClO, was prepared from NaOH and HC104 (Fisher), adjusted to pH 7, followed by filtration. Similarly, LiC104 was obtained from LiOH and HClO.+ Ag(OH); solutions were prepared by anodic oxidation as previously described.4 About 200 cm3 of 1.2 M NaOH or LiOH solution were placed in a 250 cm3 polyethylene beaker containing a silver foil anode. Argon was bubbled through the silver(II1) solution during electrolysis from a capillary placed near the anode to provide constant stirring of the solution. In order to minimize spontaneous oxidation of sulphite to sulphate2’ by atmospheric oxygen, solutions used to prepare SO:- solutions were also degassed with argon. In addition, great care was taken to avoid trace metal and multivalent cation catalytic auto-oxidation ” of sulphite by using very clean glassware. For the purpose of comparison, we also studied the reduction of silver(II1) by NO; and HPO$-. The nitrite stock solutions were prepared by direct weighing of reagent grade NaN02 (Mallinckrodt). The phosphite solutions were made from a 50% solution of phosphorous acid (Allied Chemical) by potentiometric titration with 1.2 M NaOH. An Aminco-Morrow stopped-flow apparatus was used in the kinetic runs.4 Large excesses of reductant were used for all experiments. The reactions were monitored as Ag(OH); disappearance at267nm(s= 1.17~10~M-‘cm-~).Atthiswavelength, neither sulphite, nitrite, nor phosphite

1301

Reduction of Ag(OH); by sulphite showed any significant absorbance. The reactions were followed on a Tektronix 564B storage oscilloscope coupled to a Radio Shack TRS-80 computer which recorded and processed the kinetic data. All observed pseudo-Grst-order rate constants are reproducible within 5%. Spectral analysis was done on a Cary 15 W-vis spectrophotometer. Product analysis

Known concentrations of sulphite and Ag(OH): were mixed together. Immediately upon completion of the reaction, a saturated solution of barium nitrate was added to precipitate the product and excess sulphite where excess reductant was used. After filtration, the precipitates were dried and their IR spectra recorded on a Perkin-Elmer 1310 spectrophotometer. Blank sulphite and sulphate solutions were also analysed to check for possible contamination. To test for reactions between reactants and products, parallel experiments were done for the systems silver(&-sulphite and silver(II)-sulphate. RESULTS Identljication

of the reaction product

At excess Ag(OH):, only sulphate IR peaks were observed. When excess reductant was used, sulphate and sulphite peaks were obtained. We therefore conclude that the reaction product is sulphate.

When an excess of sulphite is mixed with yellow Ag(OH);, the colour is immediately discharged. A rapid first-order decrease in absorbance in the stopped-flow apparatus over the region 230-400 nm was observed, with no indication of intermediate complex formation. The dependence of the reaction rate on the sulphite ion concentration (5 x lo- 4 M < [SO:-] < 5 x lop3 M) was studied at 25°C and p = 1.2 M, with initial silver(II1) concentrations ranging from 1 x lo-’ to 1 x lop4 M. As illustrated in Fig. 1, the reaction shows a firstorder dependence in excess sulphite as given in eq. (3) :

The stoichiometry of the reaction was determined at excess silver(II1) by measuring the difference between initial and final absorbances in the stopped-flow apparatus. The value for A[Ag”‘] thus obtained when compared with the amount of sulphite present, indicated a stoichiometry of 1: 1. Attempts to determine the reaction stoichiometry at excess sulphite by iodometric back titration were not conclusive. Indeed, scattered ratio values were obtained. The difficulty in getting good stoichiometric data at excess sulphite is probably due to air-oxidation of sulphite or sulphate.” The detection of sulphate at moderate and large excess sulphite concentrations, however, leads us to conclude that the 1 : 1 stoichiometry is also operative under the kinetic conditions. Based on the 1 : 1 stoichiometric ratio and the product analysis at excess SO:-, the reaction goes as follows [eq. (211: Ag(OH); +SO:- +HzO. (2)

= kdAgKW;1 = k[SO:-][Ag(OH);].

- WtdOW;lldt

(3)

A least-squares treatment of the data gives a straight line with an intercept that is close to zero (Fig. 1). The second-order rate constant obtained fromthistreatmentisk = (9.1&0.1)x 102M-Is-‘. Agreement between observed and calculated pseudofirst-order rate constants is generally better than 5% (Table 1). Temperature

dependence

The reaction was carried out over the range 1&53”C with a 1 x low3 M sulphite solution at [OH-] = 1.2 M and p = 1.2 M. A plot of the data using eq. (4)30 is linear (Fig. 2). In (k,,,NhlRT)

Reaction stoichiometry

Ag(OH); + SO:- L

Kinetics

= -AHr/RT+

AStIR.

(4)

The enthalpy and entropy of activation calculated from the slope and intercept of the leastsquares treatment of eq. (4) are 27 kJ mol- ’ and - 153 J K- ’ mol- ‘, respectively. The recalculated temperature dependent rate constants are in good agreement with the observed rate constants.

', i!

-

q /

. 0.12M * 0.3 M q 0.6 M t 0.9 M " 1.2 M

Fig. 1. Variation of the observed pseudo-first-order rate constants with sulphite at different hydroxide concentrations.

1302

L. J. KIRSCHENBAUM

et al.

Table 1. Values of the observed pseudo-first-order rate constants for the reaction of Ag(OH); with excess sulphite in different hydroxide concentrations @= 1.2 M)

[OH-] (M) [SO:-] x lo3 (M)

NOHI 0.5 1.0 2.5 5.0

0.50 1.03 2.34 4.72

0.52 0.98 2.38 4.71

0.5 1.0 2.5 5.0

0.48 1.10 2.31 4.42

0.52 0.98 2.38 4.71

0.6

0.5 1.0 2.5 5.0

0.42 1.12 2.60 4.50

0.52 0.98 2.38 4.71

0.3

0.25 0.5 1.0 2.5 5.0

0.27 0.44 0.95 2.00 4.60

0.28 0.52 0.98 2.38 4.71

0.12

0.5 1.0 2.5 5.0

0.49 0.99 2.00 4.66

0.52 0.98 2.38 4.71

0.5 1.0 2.5 5.0

0.45 1.04 2.34 4.67

0.52 0.98 2.38 4.71

1.2

0.9

[LiOH] 1.2

“Calculated from the least-squares line of k = 0.07+ 9.10 x loz[so:-]. Ionic strength dependence

The kinetics were followed in the ionic strength range from 0.12 to 1.2 M in both LiOH and NaOH at [SO:-] = 5 x 1O-3 M. A hydroxide concentration of 0.12 M was maintained for all reactions. A net increase of the pseudo-fist-order rate constant with ionic strength was observed. A DebyeHiickel-Brtinsted-Davies3 ’ equation plot of log kobsvs ~‘/~/(p~‘~+ 1) - 0.05~ was linear in the *The decomposition half-life of Ag(OH); in electrolysed 1.2 M NaOH is about 90 min. However, when mixed in the stopped-flow apparatus with solutions which have not been pre-electrolysed to remove trace impurities, the decomposition rate of this complex is significantly faster. Therefore, higher phosphite concentrations and higher ionic strength were required in order to separate the reaction of phosphite with Ag(OH); from spontaneous decomposition.

-30 3

3.2

3.1

3.3 l/7;

3.4

35

3.6

I/KX IO3

Fig. 2. Temperature dependence for the reaction between sulphite ([SO:-] = 1 x lo-’ M) and Ag(OH);.

case of LiOH, with a slope of 1.9 (Fig. 3). This is close to the value of two predicted for the interaction of mononegative Ag(OH); with dinegative SO:-. However, at high ionic strengths (3 0.6 M), a small deviation from this line was observed for the NaOH data. Reaction of siIver(II1) with nitrite andphosphite

When a nitrite solution is mixed with silver(II1) in 1.2 M NaOH, the rate of decrease of the silver(II1) absorbance cannot be distinguished from the expected spontaneous decomposition of Ag(OH);,3*6 even at [NO;] = 1 M. For the system silver(III)-phosphite, the reaction was first carried out over the range 0.01408 M [HPO:-1, [OH-] = 0.96 M, p = 1.2 M at 25°C. A slight acceleration in the decomposition of Ag(OH); was observed. Attempts to account for this change by kinetic means were unsuccessful. There seems to be a competition between the spontaneous decomposition of Ag(OH); and the oxidation of phosphite, which we were unable to separate because of the complex product-catalysed kinetics of the former reaction.* However, when

‘c

0.9 0.6

2

Q7

5 0

0.6

. .

I/

.

0.5 0.4

q

in LiOH

.

hNaOH

l

t-p' 0,302 I

0.25

0.3

0.36

/P/Q'%

0.4

0.45

0.5

0.55

I,-0.0%

Fig. 3. Ionic strength dependence in aqueous perchlorate media at [SO:-] = 5 x lo-’ M and [OH-] = 0.12 M.

1303

Reduction of Ag(OH); by sulphite l-l

cl

O\p7

O\pI-

O\p

\

&_ Ho&w_

/O

0' I

\

I

67 I

.’

cm

/

Ap'

/ //’

A

C

0

Scheme 1. three concentrations of phosphite in the range 0.20.4 M @ = 1.88 M, [OH-] = 0.72 M) were used, the decrease in absorbance at 267 nm was indeed first-order. The plot of kobs vs [HPO:-] over this limited range was linear with a zero intercept. The slope of this line corresponds to a second-order rate constant of 1.0 x 10-l M-’ s-l. DISCUSSION Sulphite has been found to be an effective complexing agent for transition metal ions such as palladium(II),32 cobalt(III),33 and tellurium(I).34 In the reduction of ferricyanide by sulphite, two iron sulphite intermediate complexes were reported. 4 In the present study, however, there is no kinetic or spectral evidence for a silver(III)sulphite transient complex. In previous studies, it has been reported that the product of the reduction of transition metals by sulphite is either dithionate or sulphate, According to the product analysis, or both. 23*26,35 sulphate appears to be the only oxidation product of this reaction. The possibility of sulphate resulting from the disproportionation of the dithionate in base can be excluded. Although dithionate is thermodynamically unstable36 in base, it is kinetically inert with respect to disproportionation.37

In blank experiments with Ag”‘+SO:and Ag’+ SO:-, no reactions were observed, indicating that reactants and products do not react with one another. Furthermore, the stoichiometry observed in this study supports that sulphate is the only product of the reaction. As is inferred in the two-electron transfer reaction of arsenite with Ag(OH);,” the oxidation of sulphite by silver(III) probably proceeds via a one-step oxygen atom transfer. Two kinetically indistinguishable pathways that account for our results are in Schemes 1 and 2. Scheme 1 involves attack by a sulphite oxygen on an axial site of the squareplanar silver(II1) complex. In Scheme 2, three-coordinate sulphur(IV) undergoes a nucleophilic attack by a hydroxyl ligand followed by oxygen transfer in either a singly-bridged or chelated (as in Scheme 1) intermediate. Both these schemes depend on the expansion of the sulphur(IV) coordination sphere with the fourth oxygen in the sulphate product deriving from Ag(OH);. The insignificant intercept observed in Fig. 1 indicates that there is no participation by Ag(OH),(H,O) in the reduction process as is in the thiosulphate reaction. I7 The lack of any hydroxide dependence on the rate, is as expected since SO:- is the only sulphurous acid species in alkaline media (pKa, = 1.91,

“\

PI

OAS

I

bH

HO

0’ Scheme 2.

’

0

1304

L. J. KIRSCHENBAUM

pKaz = 7.18). 38 The possibility of a rapid transfer

of a hydroxyl ligand to sulphite, concurrent with donation of its electron pair to silver(II1) via the empty d,z_,,z orbital (oriented in the equatorial plane of the complex), appears reasonable. Support of the possibility that redox occurs through a fivecoordinate intermediate as in Scheme 1 is provided by the results of other Ag(OH); reactions. For instance, the presence of third-order terms in the reaction of en2’ HO; ” and thiourea3’ implies the formation of five-coordinate transients. Activation parameters for silver(II1) complexation reactions have been interpreted in terms of an associative mechanism for ligand exchange. The activation parameters reported here are similar to those obtained in the reduction of silver(II1) by arsenite15 species, H*AsO; and HAsO:-, where a five-coordinate silver(II1) intermediate was postulated. The reaction is likely to proceed by (1) ion association, (2) axial attack by negative sulphite on the positive silver(II1) centre, (3) changes in bond lengths of the two reactants and reconstruction of their solvation spheres and (4) a fast electron transfer step involving the rupture of a silver-oxygen bond to yield the products. The observed rate seems to be due to the formation of a bridged precursor intermediate (A or A‘) which is then followed by a fast redox process. This is consistent with a large negative entropy as is observed. The good agreement for the ionic strength dependence in LiOH-LiC104 with expectations indicates the absence of substantial ion pairing in the intimate step of the reaction, as is reported for similar reactions.26,40 We note that at 1.2 M OH-, there is no difference in rate between NaOH and LiOH (Table 1). Thus sodium (but not lithium) ion pairing is probably not responsible for the deviation in Fig. 3. If we presume that the reduction of silver(II1) by phosphite is, like the sulphite reaction, independent of [OH-] and has the expected ionic strength variation, we may correct the second-order rate constant obtained at p = 1.88 M3’ and compare the rates of the two reactions. The rate constant correctedtop = 1.2Mforphosphiteis9.5x 10-2M-’ s- ’ which is about four orders of magnitude smaller than the value of 9.1 x 10’ M-’ s-’ obtained for the reaction of silver(II1) with sulphite (25°C). The difference in reactivity between sulphite and phosphite, which are isoelectronic, may stem from the fact that sulphite has a lone pair of electrons on sulphur that is readily available in the electron transfer process. On the other hand, there is no such pair of electrons on the phosphite phosphorus which is bound to a non-labile hydrogen.4’ Ghosh et a1.42 reported that although phosphite is a power-

et al.

ful reducing agent, the reduction of the chromium(V) chelate bis(2-ethyl-2-hydroxybutyrato) oxochromate(V) by phosphite is immeasurably slow. This is consistent with the sluggish reactivity of phosphite. Kinetics of nitrite ion oxidation by metal ions have been recently reported.42-44 For example, the reduction of chromium(V) by nitrite was inferred to proceed by an NO; radical intermediate.42 Wilmarth et aZ.44reported rate constants for oxidations of Fe(bpy)i+ by nitrite and sulphite of 3.3 x lo4 and 4.6x lo6 M-’ s-‘, respectively. However, there is no evidence of a nitrite-silver(II1) reaction even at [NO;] = 1 M. An explanation for the non-reactivity of nitrite may be found in the inability of nitrogen to accept electrons into d orbitals. Before the bound hydroxyl oxygen can be transferred, an additional electron pair must coordinate to the central atom of the oxoanion (vide spa). While the availability of d orbitals permits an expanded number of valence electrons in arsenite, sulphite and phosphite, nitrite cannot coordinate to oxygen while it is also bound to silver. One-electron transfer to form NO; and a silver(I1) intermediate, while a potential alternative to oxygen atom transfer, would have a high activation energy. The NO;/NO, potential is just over 1 V.45 We have estimated the Ag(OH);/Ag(OH):one-electron reduction to be about 3/10 V less than this,46 thus posing a thermodynamic barrier to the electron transfer. Therefore, the non-reactivity of NO; with Ag(OH); is entirely consistent with the oxidation of the other oxoanions via an atom transfer mechanism.

Acknowle&ements--We wish to thank the Donors of the Petroleum Research Fund administered by the American Chemical Society for support of this Research.

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by sulphite

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28. H. N. Alyea and H. L. T. Backstron, J. Am. Chem. sec. 1929,51, 90. 29. G. Nickless, Inorganic Sulphur Chemistry. Elsevier, Amsterdam (1968). 30. F. Basoto and R. G. Pearson, peckish of Inorganic Reactions, 2nd edn. John Wiley, New York (1967). 31. B. Perlmutter-Hayman, in Progress in Reaction Kinet&s, Vol. 6, p. 239. Pergamon Press, Oxford (1972). 32. G. Mahal and R. Van Eldik, fnorg. Chem. 1987,26, 2838. 33. W. G. Jackson, Inorg. Chem. 1988,27,777. 34. L. Peter and M. Beat, fnorg. Chem. 1985,19,3071. 35. D. W. Carlyle and 0. Zeck Jr, Inorg. Chem. 1973, 12, 2978. 36. A. J. Bard, R. Parsons and J. Jordan, Standard Potentials in Aqueous Solution, p. 106. Marcel Decker, New York (1985). 37. D. W. Carlyle, J. Am. Chem. Sot. 1972,94,4525. 38. D. C. Harris, Quantitative Chemical Anulysis, 2nd edn, p. 737. W. H. Freeman and Cie, New York (1987). 39. L. J. Kirschenbaum and R. K. Panda, Patyhedron 19X8,7,2753. 40. A. R. Olson and T. R. Simonson, J. Chem. Phys. 1949,17,348. 41. H. S. Gutowski, D. W. McCall and C. P. Slichter, J. Chem. Phys. 1953,21,279. 42. S. K. Ghosh, R. N. Bose and E. S. Gould, Inorg. Chem. 1987,26,2688. 43. G. Stedman, A&. Inorg. Chem. ~adioe~m. 1979, 22, 113. 44. W. K. Wihnarth, D. M. Stanbury, J. E. Byrd, H. N. PO and C.-P. Chua, Coard. Chem. Rev. 1983,51,155. 45. M. S. Ram and D. M. Stanbury, J. Am. Chem. Sot. 1984,106,8136. 46. E. T. Borish, L. J. Kirschenbaum and E. Mentasti, J. Chem. Sot., Da&on Trans. 1985, 1789.