摘要:

J. CHEM. SOC. DALTON TRANS. 1983 1029Electrochemical Reduction of the Dichloro( 3,6-diazaoctane-I ,8-diamine)-rhodium(rii) Cation, cis-.-[ Rh(trien)CI,] +Hector W. Munro, Donald M. Stevenson, and D. Huw Vaughan *Chemistry Department, Paisley College of Technology, Paisley, Renfrewshire PA I 2BEThe electrochemical reduction of cis- a- [ R h(trien) Clz] + [trien = triethylenetetramine(3,6-diazaoctane-I ,8-diamine) J has been studied using polarography and controlled potentialelectrolysis. The polarographic half-wave potential was found to depend on the pH above ca. 6.0;this was attributed to the catalysed hydrolysis of the starting complex to the corresponding chloroaquaor chlorohydroxy ion by rhodium([), which was the initial product of the electrochemical reduction.Below pH ca.6.0, no such conversion occurs because the rhodium([) is removed rapidly from thesystem by reaction with H30+ to form a rhodium(iii) hydride.There have been relatively few reports regarding the electro-chemistry of rhodium(rr1). Some of these have reportedvalues of the polarographic half-wave potential (&) for aseries of complexes, but no detailed study was carried out onthe mechanism of the electrode reaction although it wasestablished that the majority of the complexes underwent anirreversible two-electron reduction to, initially, a complex ofrhodium(1).A detailed electrochemical study of the complex ion trans-[Rh(en)2C12]+ (en = ethylenediamine) in aqueous solution 4-6using the techniques of cyclic voltammetry, controlled poten-tial electrolysis, and chronocoulometry proved that the initialproduct of the electrochemical reduction was [Rh(en)J+,but that this species underwent a series of reactions such thatthe final products of the electrode reaction were dependent onpH and on the potential at which the electrolysis was carriedout.Thus it was possible to identify amongst the products ofthe electrode reaction ~rans-[Rh(en)~Cl(OH)] + , the hydridetrans-[Rh(en),H(OH)] +, mercury adducts, and [Rh(en)2]+,which under certain circumstances was strongly adsorbed onthe mercury electrode.A brief polarographic study ’ of the tris(dithioacety1-acetonato) complex of rhodiurn(II1) in acetone again estab-lished that the complex underwent a diffusion-controlled,irreversible, two-electron reduction.It was also suggestedthat the Rh’ product was strongly adsorbed on the mercury,or that possibly a mercury adduct was formed.The only detailed polarographic studies of complexes ofrhodium(rI1) have been conducted on tran~-[Rh(py)~Cl~] +and [Rh(NH3),(OH)]2+.9 Both these were shown to undergo atwo-electron reduction with formation of a hydride by reac-tion of the initially formed Rh’ with water; in the case of thereduction of [Rh(NH3)5(OH)]2+, the hydride was identified as[Rh(NH3)5H]2+. The effect on the polarographic half-wavepotential of varying the concentrations of H+, C1-, and NH3was also investigated so that it was possible to identify thenature of the species actually reduced at the electrode.The only apparent exception to the general two-electronreduction to Rh’ is with tris- and bis-(bipyridyl) lo and -(l,lO-phenanthroline) complexes of rhodium(I1r) where it wasfound that more than two electrons (up to four) were trans-ferred to the complex.This was due to some, at least, of theelectrons being transferred to the extensively delocalisedorbitals on the ligand rather than to the rhodium itself.As part of a study of the factors influencing the behaviour ofrhodium complexes as catalysts in homogeneous reactions,we have investigated the electrochemical properties of com-plexes of rhodium(Ir1) because of their importance in thecatalytic We report here a detailed electro-chemical study of the complex cation l4 cis-a-[Rh(trien)C12]+(trien = triethylenetetramine) using the techniques of d.c.polarography and controlled potential electrolysis.No com-parable study of a rhodium(rr1) complex with this geometryappears to have been reported.ExperimentalBuffer solutions were made from Britton-Robinson stocksolution (0.04 mol dm-j acetic acid, 0.04 rnol dm-3 boricacid, and 0.04 mol dm-3 phosphoric acid). A measuredvolume (12.5 cm3) of this solution was run into a 25 cm3flask, the required amount of 0.2 mol dm-3 sodium hydroxidesolution was added, and the volume was made up with 0.2rnol dm-3 potassium nitrate solution, thus maintaining aconstant concentration of supporting electrolyte throughoutthe pH range. The chloride concentration was varied in asimilar manner by adding KCl.The complex cis-a-[Rh(trien)C12]Cl*H20 was prepared asdescribed previously; l4 its electronic spectrum was in agree-ment with the previous report [A,,,.352 and 288 nm (& 240and 210 dm3 mol-I cm-’)I. The complex was stable in acidand neutral solution but above pH ca. 10 the rate of hydroly-sis became significant and so fresh stock solutions of thecomplex (2.5 x rnol dm-j) in deionised water wereprepared every 2 days. Suitable volumes of the stock solutionwere then diluted with a buffer solution to give the requiredfinal solution.Classical d.c. polarography was carried out using a Radio-meter Polariter P04(d) instrument; the capillary employedhad the following characteristics: M = 3.074 mg s-l, t = 3.3s when h(corr.) = 64.8 cm.For the controlled potentialelectrolysis experiments a vessel of about 50 cm3 volume wasfilled with doubly distilled mercury to a depth of about 1 cm.Electrical contact with the mercury was made by a platinumwire electrode set in glass. The complex solution (25 cm3 of0.001 mol dm”) was then added, and the cell completed witha reference electrode and platinum anode both separatedfrom the solution by sintered glass. A pH combined glassreference electrode and nitrogen inlet were also included. Themercury cathode, anode, and reference electrodes wereconnected to a Witton T6 Tutorial Potentiostat and the cur-rent was passed through a 2.2 ohm resistor. The voltage dropacross this resistor was recorded on a Servoscribe recorderas a means of following the current changes during theelectrolysis. The quantity of electricity passed was calculatedfrom the area under the voltage-time curve.Electronic spectra were recorded on a Pye-Unicam SP 8-100spectrophotometer using silica cells.Changes in the electroni1030 J. CHEM. SOC. DALTON TRANS. 19830.7 I Table 2. Variation of E+ with [Cl-] for the complex ion cis-a-[Rh(trien)CIz]+ at fixed values of pH (complex concentration lo-'mol d ~ n - ~ ; E+ values relative to s.c.e.)PHFigure. Variation of the polarographic half-wave potential (E,)with pH for the complex ion cis-a-[Rh(trien)CI,] +Table 1. Variation of E+ with pH for the complex ion cis-a-[Rh(trien)CI,]+ (complex concentration mol dm?; E., valuesrelative to s.c.e.)- PH3.464.024.706.006.757.207.507.848 .oo8.509.5410.8' Main wave.Pre-wave.E,IV '0.740.740.740.7650.800.8350.840.8850.880.9051.011.06- E*/V0.6850.6850.6850.700.720.730.740.72spectra during the controlled potential electrolysis experi-ment were followed in a separate run by using 75 cm3 ofa 0.001 mol dm-3 solution of the complex and pumping thesolution via plastic tubing through a ' flow through ' silicacell in the sample compartment of the spectrophotometer.Results and DiscussionPo1urugraphy.-The poiaroyraphic waves for lo-* tnoldm-j solutions of the complex ion cis-u-[Rh(trien)CIz]+ inbuffer solutions between pH 3.46 and 10.8 fell into twodistinct groups (Figure and Table 1). From pH 3.46 to aboutpH 5.5 the waves were single steps of nearly constant half-wave potential.At higher values of pH there was a marked pHdependence: the half-wave potentials varied linearly with pHand with a slope of -0.059 V per pH unit. At the same time apre-wave became apparent, beginning as a small disturbance atpH ca. 6.45 and increasing with increasing pH until by pH 8it accounted for about 20% of the limiting current. At highervalues of pH the pre-wave remained of constant height.At all values of pH the height of the pre-wave was directlyproportional to the height of the mercury column (correctedfor back pressure) and it was therefore concluded that thiswave was due to adsorption of the reduction product.For all waves the total heights (adsorption plus main wave)were directly proportional to the concentration of complexover the pH range 3.46-9.5 though the gradients of the wave-height-concentration plot decreased at higher pH.Forexample, at pH 7.6 the gradient was 13.8 PA I mmol-' while atpH 9.5 it had decreased to 6.2 PA 1 mmol-'.To establish the extent to which the reductions werePH [Cl-]/mol dm-33.5 10-47.2 a 10-40.07840.2960.01 60.07842.000.07840.1540.2960.07840.2967.84 10-410.8 10-4- EtlV0.740.7550.7450.8350.950.9651.040.8851.01 51.031.0451.061.091.11' dE,/d(pCI) = 0.040 V per pCI unit.pC1 unit.dE,/d(pCI) = 0.045 V perdiffusion-controlled the total wave-height was plotted againstthe square root of the corrected height of mercury.Up toabout pH 7.6 these plots were linear, of gradient 0.17 pAcm-*, and passed through the origin. At high pH howeverthe gradients were markedly lower, though the plots werestill linear, and significant intercepts on the current axis werenoted. For example, at pH 9.3 the gradient was 0.04 FA cm-*and the intercept 0.48 PA.From these observations it is clear that there is a changefrom diffusion control of the polarographic current at lowpH to a mixture of diffusion and kinetic (or catalytic) controlof the current as the pH increases.In the region of diffusion control the llkovic equation canbe used to estimate the number of electrons (n) involved inreduction of the complex ion. For this purpose a value of4.4 x lo-" cmz s-I was taken for the diffusion coefficient.2The results ranged from 3.5 to 4.2 over the pH range 3.5-7.6.Since the simple reduction of a rhodium(rr1) complex torhodium(1) should require n to have the integral value of 2,there appears to be some additional reaction.Further evidencefor this view comes from the logarithmic analysis of thewaves. According to the theory of irreversible waves l5 theplots of potential (E) against log (i/il - i) should be linear andof gradient -0.054/anA where i = current at potential E,il = limiting current, u = transfer coefficient, and nA =number of electrons involved in the rate-determining step.The results obtained gave unA values which varied with pH(in parentheses):0.68 (3.46), 0.76 (4.70), 1.28 (7.20), 1.52 (9.54)Since a is most likely to have a value between 0.3 and 0.7,and this may vary with potential it is reasonable to proceedon the assumption that f l A has the value 2.The effect of an excess of chloride ion on half-wave poten-tial at various values of pH was investigated (Table 2 showstypical results).In solutions of pH 3.5 and pH 10.8 the excessof chloride had no effect, but between these values E+ movedto more negative values as the chloride ion concentrationwas increased. The variation of E* with pC1 was linear, thegradient being about 0.04 V per pC1 unit, and the value ofanA at constant pH remained constant as the chloride ionconcentration changed.The dependence of E+ on pH and on pC1 can only beexplained if the reduction of the electroactive specieJ.CHEM. SOC. DALTON TRANS. 1983 1031involves the participation of the hydrogen ion and chlorideion. The gradient of the E+-pCl plot points to the reduciblespecies having one chloride ion less than the original complex,and as the dichloro ion is known to be hydrolysed in aqueous~ 0 1 u t i o n , ~ ~ ~ ~ ~ * ~ * the ion which is most likely to be reducedunder the conditions in which chloride ion concentrationaffects E-, is the chloroaqua ion formed by hydrolysis [re-action (i)]. Even if the chloroaqua ion is not the actual speciesreduced, this ion must be the source of the reducible species.cis-a-[Rh(trien)C12]+ + H20 +~is-a-[Rh(trien)Cl(H~O)]~+ + C1- (i)A difficulty with this mechanism is that hydrolysis of thedichloro ion is generally s~ow,~~*'~*'* and the establishment ofequilibrium (i) demands that both forward and backwardreactions be reasonably fast.This anomaly can be explainedif some rhodium(r) is present, because it is well established 43*19that in neutral or basic solution rhodium(1) can catalyse thesubstitution reaction of similar complexes of rhodium(r1r)by an inner-sphere redox reaction. Rhodium(1) is of courseproduced in the electrode reaction, and so the equilibriumcan readily be set up at the surface of the mercury as soon asreduction commences. The requirement of neutral or basicconditions for the catalytic action also explains why E* isindependent of pC1 at low pH. The fact that a polarographicwave is observed in acid solution shows that the dichloroion can be reduced directly.The half-wave potential for thereduction of the dichloro ion must be close to that for reduc-tion of the chloroaqua ion (otherwise two waves would beobserved) and equilibrium (i) must lie well to the right whenit is set up (otherwise the chloroaqua ion would not be thechief electroactive species). Just as the dependence of E+ onpC1 points to the participation of the chloride ion, so thedependence on pH in solutions of pH greater than about 5.5is indicative of an electrode reaction involving the hydrogenion. The direction and rate of variation of E* (-0.059 V perpH unit) is evidence of a simple acid-base reaction in which theelectroactive species acts as an acid.15 In the present case themost probable equilibrium, in view of the well knowntendencies of aqua cations to act as acids in is (ii).cis-a-[Rh(trien)C1(H20)]'+ + H20 +cis-a-[Rh(trien)Cl(OH)]+ + H30+ (ii)Here the acid species is the chloroaqua ion.While the effects of chloride and hydrogen ion on the half-wave potential provide extremely strong evidence that theion which is electroactive in neutral or basic solution is thechloroaqua ion, this ion can account for the consumptionof only two electrons in its reduction to rhodium(r), and theheight of the polarographic waves requires that four electronsbe used per rhodiurn(111) complex ion.However, controlledpotential electrolysis experiments provide evidence that thereduction of chloroaqua ion is accompanied by reduction ofhydrogen ion, except at very high pH.Electrolysis at Controlled Potentid-Solutions of 1 0-3mol dm-3 cis-a-[Rh(trien)C12]+ were electrolysed using amercury cathode at - 1.2 V us.the saturated calomel electrode(s.c.e.). (This potential corresponded to the top of the waves.)All solutions were stirred with a magnetic stirring bar on topof the mercury pool, and both pH and the electronic absorp-tion spectrum were monitored during electrolysis. In all casesnitrogen gas was passed for 10 rnin before and during electro-lysis.With unbuffered 0.2 mol dm-3 KCl as supporting electro-lyte, the pH rose very rapidly, reaching pH 8.9 after 2 minand 9.96 after 10 min. As the current approached its back-ground value of about 0.50 mA after about 50 min, the pHlevelled off near pH 11.2. During the electrolysis spectralpeaks at 352 and 288 nm due to the dichloro ion were replacedby a broad shoulder at 295 nm (E 582 dm3 mol-' cm-I).Thetotal electricity passed during complete electrolysis was 5.610coulomb [about 16% more than necessary for the completereduction of all rhodium(111) to rhodium(~)]. When 0.2 moldm-3 KN03 was used as supporting electrolyte, the totalquantity of electricity passed was 8.035 coulomb, which issome 66% more than necessary for reduction of rhodium(n1)to rhodium(r). However, the final spectrum and pH wereidentical with those obtained when KCl was used as support-ing electrolyte.Different behaviour was observed on electrolysing usingbuffer solutions of pH 3.5, 8.0, and 10.0 as supportingelectrolytes. In all these cases the electrolysis current did notdecay to the expected background of about 0.50 mA (as withthe unbuffered solutions) but fell only to between 2 and 4 mA.During electrolysis at pH 3.5 and 8.0, the spectrum of thereactant ion was replaced by a peak at 300 nm (E 460 dm3mol-' cm-').At pH 10.0 the final spectrum was identicalwith that obtained in the unbuffered solution, i.e. a shoulderat 295 nm (E 582 dm3 mol-l cm-I).When the electrolysis was conducted at pH 11 -4 the currentdecayed to a small value and the quantity of electricity passedcorresponded to n = 1.95. The final spectrum was againidentical with that obtained in the unbuffered solution. In allthese experiments no spectroscopic evidence was obtainedfor any intermediates, i.e.the spectrum of the reactant ionwas replaced gradually by that of the final product as theelectrolysis proceeded.Previous work 4-623*9 on the electrochemistry of rhodium(rr1)has shown that the initial product of electrochemical reduc-tion is a rhodium(r) species which may undergo an acid-basetype reaction to form a hydride according to (iii). At higherRh+ + H30+ + [RhH(H2O)l2+values of pH equilibrium (iii) will lie to the left, so that theconcentration of Rh' increases as the pH increases. It wouldthus be expected that the final product of the controlledpotential electrolysis of cis-a-[Rh(trien)C12]+ at an acid pHwould be a hydride of probable formula cis-a-[Rh(trien)H-(H20)l2+. The electronic spectrum of this species has beenreported 21 as showing a peak at 300 nm, and this is consistentwith the position of the final peak observed in controlledpotential electrolysis at pH 3.5 and 8.0. Unfortunately themolar absorptivity found here (460 dm3 mol-' cm-') is muchgreater than that reported previously (108 dm3 mol-I cm--');this discrepancy will be discussed in the next section.If weaccept that the hydride is the final product in buffer solutions,the spectral change during electrolysis is accounted for bythe production of the hydride in equilibrium with a rhodium(1)species according to (iv). In unbuffered solutions in which[Rh(trien)]+ + H30+(iii)cis- a- [ Rh(trien)H(H20)] 2+ (iv)the pH becomes high, or in buffered solutions of high pH,the favoured species will be rhodium(1) rather than thehydride, and so the different spectroscopic changes observedwith unbuffered solutions are explicable.The difference in current consumption between bufferedand unbuffered solution can be explained if the hydride iselectroactive according to (v).The participation of such acis-a-[Rh(trien)H(H20)]'+ + e- +[Rh(trien)]+ + 3H2 + H20 (v1032 J. CHEM. SOC. DALTON TRANS. 1983reaction can also account for the consumption of more thantwo electrons per complex ion in the polarographic reduction,as (iv) and (v) together are equivalent to the catalytic reduc-tion of hydrogen ion. A similar sequence has been observedpreviously in the case of trans-[Rh(en)2H(H20)]2 +, which iseffective in catalysing the reduction of hydrogen ion to di-hydrogen gas at a mercury cathode in the pH range 3-11.The extent to which these two reactions participate willdepend on pH through the control exerted by (iv).Underconditions of high pH the catalytic reduction is inhibitedbecause equilibrium (iv) lies well to the left, so that at pH11.4 virtually no catalytic reduction of hydrogen ion occursand the quantity of electricity consumed corresponds exactlyto reduction of the rhodium(@ reactant to rhodium(1). Atmore acid pH, the catalytic cycle is operative so that inbuffered solutions below ca. pH 10.00 the limiting electrolysiscurrent is high. The increase of pH during the electrolysis ofthe unbuffered solution arises from the consumption ofhydrogen ion by reaction (iv) followed by (v).The eventualcessation of the electrolysis is due to the shift of equilibrium(iv) to the left as the pH increases.General Discussion.-The polarographic results require thatthe nature of the rhodium(rr1) species which is reduced at themercury electrode depends on the pH of the solution. Belowabout pH 6.0 the reducible species is the dichloro ion as shownin (vi). At higher pH the chloroaqua ion becomes reduciblecis-a-[Rh(trien)Cl,]+ + 2e-[Rh(trien)]+ + 2C1- (vi)~is-cc-[Rh(trien)Cl(H~O)]~+ + 2e- ---t[Rh(trien)]+ + CI- + H20 (vii)[reaction (vii)]. The change from (vi) to (vii) as the chiefelectrode reaction occurs because the chloroaqua ion becomesmore abundant than the dichloro ion in the region near theelectrode, and this is caused by the aquation equilibrium (i)being set up under the influence of the catalyst rhodium(1)which is produced at the electrode.The current consumed in the controlled potential electro-lysis experiments and the heights of the polarographic wavesdemand that other reactions participate in the overall reduc-tion process.The involvement of the hydride cis-a-[Rh(trien)-H(HzO)l2+, for the presence of which there is strong evidence,in reactions (iv) and (v) not only explains the amount ofelectricity used but can also account for the change in reduciblespecies which occurs near pH 6.0. Under acid conditionsrhodium(1) is converted into the hydride [through reaction(iv)] and so the concentration of rhodiurn(1) is reduced to alevel which prevents it functioning as a catalyst for the aqu-ation of the dichloro ion.If the reduction mechanism proposed is correct thencontrolled potential reduction should result in the dichloro-chloroaqua equilibrium being established in the bulk solutionas well as in the diffusion layer at the electrode surface becauseof the production of large amounts of rhodium(1) to act ascatalyst.No spectroscopic evidence of a change from dichloroto chloroaqua ion was obtained, but it is likely that thechanges would not be observable because of the close similar-ity of the two spectra [di~hloro,'~ hmX. 352 and 288 nm(E 240 and 210 dm3 mol-' cm-l); chl~roaqua,'~ 347 and295 nm (E 212 and 157 dm3 mol-' cm-')I and by the spectrumof the reduction product being much more intense than thatof the reactants below 330 nm. Spectral changes due to con-version of dichloro into chloroaqua ion would therefore bedisguised by formation of product whether the product ischiefly rhodium(1) at high pH or the rhodium(1)-rhodium(@hydride mixture at pH less than about 10.0.The similarity ofthe spectra of the products of electrolysis at pH 3.5 and 8.0is surprising, as it might have been expected than an acid-baseequilibrium involving the species ci~-a-[Rh(trien)H(H~O)]~ +and cis-a-[Rh(trien)H(OH)] + would be set up, and beaccompanied by a change in the electronic spectrum as thepH was altered. However, it was noted previously duringthe electrolysis of trans-[Rh(en),H(H20)l2 + that no changeoccurred in the electronic spectrum of this rhodium complexas the pH increased from 3 to 1 1.It was observed earlier that the absorbance due to thehydride is different from that previously reported for thecomplex prepared by reaction between sodium tetrahydrobor-ate and cis-a-[Rh(trien)C12] + and characterised as the hydrideby its n.m.r.spectrum. It is possible that this preparation didnot result in quantitative conversion of starting material, whichwould result in a low value for E. The higher value (460 dm3mol-' cm-' at 300 nm) is consistent with known trends 22in the spectra of species of type [RhN4XY], where N4 repre-sents four nitrogen ligands [e.g. (NH3)4, (en)z, trien] and X andY are the same or different.For e ~ a m p l e , ~ ~ ~ trans-[Rh(en),H-(HZ0)l2+ has Lmx. 295 and 257 nm (E 323 and 272 dm3mol-' cm-') and on pzssing from trans to cis geometry thevalue of E for the lowest energy d-d transition usually in-creases. It would therefore be expected that E at 300 nm wouldbe greater than 323 dm3 mol-' cm-I for the complex cis-a-[ Rh( t rien)H(H20)]* + .Reaction (iv) can also explain how the polarographicadsorption current changes with pH. If [Rh(trien)] + wereadsorbed on mercury, this would give rise to an adsorptionpre-wave. But in acid solution reaction (iv) could effectivelyremove the adsorbed material to form the rhodium(m) aquahydride and so the pre-wave should be more apparent insolutions of high pH, as is indeed the case. The occurrenceof adsorption of such species has previously been reportedin the case of [Rh1(en)2]+,5 and in the electrochemical reduc-tion of tris(dit hioacet y lacet onato)rhodium(m) and trans-dichlorotetrapyridinerhodium(rIr).* The catalytic currentsobserved in the controlled potential experiments usingbuffered solutions are largely unaffected by adsorption effectssince stirring the mercury interface will continually produce afresh mercury surface.One observation that we are unable to account for fully isthat at high pH the polarographic current becomes limited bythe rate of a slow, i.e.kinetically controlled, reaction. Re-action (ii), at high pH, will lie well to the right, and thekinetic control of the polarographic current probably arisesthrough the decrease in the rate of the reverse of this reactionbecause of the diminished hydrogen ion concentration.Thekinetic control does not arise through any change in the rateof formation of the chloroaqua or chlorohydroxy species bythe rhodium(1)-catalysed inner-sphere redox reaction. Infact, it would be expected that the rate of this catalysedreaction would increase as the pH increases because of thegreater concentration of rhodium@ present through reaction(iv).References1 D. R. Crow, Inorg. Nucl. Chem. Lett., 1969, 5, 291.2 R. D. Gillard, J. A. Osborn, and G. Wilkinson, J. Chem. SOC.,3 A. W. Addison, R. D. Gillard, and D. H. Vaughan, J . Chern.4 R. D. Gillard, B. T. Heaton, and D. H. Vaughan, J . Chem. SOC.5 J. Gulens, D. Konrad, and F. C. Anson, J. Electrochem. SOC.,1965,4107.Soc., Dalton Trans., 1973, 1187.A, 1970, 3126.1974,121, 1421J. CHEM. SOC. DALTON TRANS. 1983 10336 J. Gulens and F. C. Anson, Znorg. Chem., 1973, 12, 2568.7 A. M. Bond, G. A. Heath, and R. L. Martin, J. Electrochem.8 L. E. Johnston and J. A. Page, Can. J. Chem., 1969,47, 2123.9 L. E. Johnston and J. A. Page, Can. J. Chem., 1969, 47, 4241.10 G. Kew, K. DeArmond, and K. Hanck, J. Phys. Chem., 1974,11 G. Kew, K. DeArmond, and K. Hanck, J. Phys. Chem., 1975,12 C . Masters, ‘ Homogeneous Transition-metal Catalysis,’13 B. R. James, Adv. Organomet. Chem., 1979,17, 319.14 P. M. Gidney, R. D. Gillard, B. T. Heaton, P. S. Sheridan, and15 L. Meites, + Polarographic Techniques,’ 2nd edn., Interscience,16 H. Matsuda and V. Ayabe, Bid. Chem. SOC. Jpn., 1956, 134,SOC., 1970, 117, 1362.78, 727.79, 1828.Chapman and Hall, London, 1981.D. H. Vaughan, J. Chem. SOC., Dalton Trans., 1973, 1462.New York, 1965.29, quoted by J. Heyrovsky and J. Kuta in ‘ Principles ofPolarography,’ Academic Press, New York, 1966, p. 222.17 E. Martins and P. S . Sheridan, Znorg. Chem., 1978, 17, 3631.18 E. Martins, E. B. Kaplan, and P. S. Sheridan, Znorg. Chem., 1979,18, 2195.19 R. D. Gillard, B. T. Heaton, and D. H. Vaughan, J. Chem.SOC. A , 1971, 1840; J. D. Miller and F. D. Oliver, J. Chem.SOC., Dalton Trans., 1972, 2469, 2473.20 J. Burgess, ‘ Metal Ions in Solution,’ Ellis Harwood, Cliichester,1978, ch. 9.21 R. D. Gillard and G. Wilkinson, J. Chem. SOC., 1963, 3594.22 E. J. Bounsall and S. R. Koprich, Can. J . Chem., 1970, 48,1481 ; L. H. Skibsted and P. C. Ford, Acta Chem. Scand., Ser.A, 1980, 34, 109.Received 2 1 st June 1982 ; Paper 2/ 104

ISSN:1477-9226    

DOI:10.1039/DT9830001029

出版商:RSC    

年代:1983

数据来源: RSC