J. CHEM. SOC. DALTON TRANS. 1984 103 Mechanisms of Reduction of a Nickel(1v) Oxime Complex by 1,2- and I ,4-Dihydroxybenzene Compounds in Aqueous Perchlorate Media Dona1 H. Macartney + and Alexander McAuley Department of Chemistry, University of Victoria, Victoria, B.C. V8W 2Y2, Canada The kinetics of the reduction of a nickel(iv) oxime complsx, [NiIVLI2+(H2L = 3,14-dimethyl- 4,7,10,13-tetra-azahexadeca-3,13-diene-2,15-dione dioxime), by catechol, hydroquinone, and 2-methylhydroquinone ( H2Q) have been investigated in aqueous perchlorate media in the pH range 3.5-6. Below pH 4.5 the reaction is monophasic with the NilV-Nilll step rate-determining. Above pH 4.7 a biphasic process is observed with a rapid NiIV-NilI1 reduction followed by a slower reaction of the nickel(iii) intermediate.A complex pH dependence is attributed to reaction pathways involving several reductant and oxidant species related by proton equilibria. For the principal reaction pathways involving [Ni1"LJ2+ and HQ-, rate constants fall in the range ca. 10'-I O8 dm3 mol-1 s - l . With H2Q as reductant, rate constants are ca. 105-106 times smaller. The specific rate parameters compare favourably with the rate constants determined by Marcus theory calculations. There has been considerable interest recently in the higher oxidation states of nickel.' Numerous complexes of nickel(w), and a lesser number of nickel(1v) have been reported. Nickel(ri1) complexes frequently contain nitrogen donor ligands, such as tetra-aza macrocycles. The stabilization of the nickel(iv) oxidation state has been achieved with the use of deprotonated oxime ligand~.~ The sexidentate ligand 3,14- dimethyl-4,7,10,13-tetra-azahexadeca-3,13-diene-2,15-dione dioxime ( HzL) forms pseudo-octahedral complexes with Ni", Nili1, and In aqueous solution [NiivLl2+ is an oxidizing agent (Ni'V-Nilsl, E* = 0.65 V) and its redox reactions with organic and inorganic 6.7 reductants have been reported. With metal-ion reductants in acidic solution, the one-electron reduction of [NilVLl2 + is the rate-determining step with no evidence for any long-lived nickel(iii) species.Kinetic investigations at higher pH, using ascorbic acid and [Co(phen),lz +(phen = I ,lo-phenanthroline) as reductants, have revealed separate one-electron processes involving [Nil"L]'+ and a detectable Ni"' intermediate. I n this paper, the results of a kinetic study of the reduction of [NiL]'+ by catechol (H,cat) (,I ,2-dihydroxybenzene), hydroquinone (H2quin) (I ,4-dihydroxybenzene), and 2- methylhydroquinone (H2mquin) are presented.Rate constants were determined over the range pH 3.5-6 at 25 "C to obtain further information on the nature and reactivity of the nickel(irr) and nickel( iv) species in solution. Experimental Materials.-The nickel(i1) oxime perchlorate, [Ni(H,L)]- [ClO&, was prepared by addition of H2L to a methanol solution of nickel(i1) per~hlorate.~ The nickel(1v) oxime, [NilvL]z+ + HzQ + [Ni"(H,L)IZ+ 4- Q ( I ) [NiL][CIO4],, was isolated after oxidation of [Ni(H2L)][C1O4I2 in concentrated nitric acid, as described previ~usly.~ Analysis of the product was satisfactory (Found: C, 29.25; H, 4.65; N, 15.05.C14H26C12N,Ni0,0 requires C, 29.60; H, 4.60; N, 14.80%). The concentration of aqueous [NiL]' + solutions was determined spectrophotometrically at 500 (E = 6 300) and 430 nm (E = 5 960 dm3 mol-I ~ m - ' ) . ~ Hydroquinone (Aldrich), catechol (Fisher) and 2-methyl- hydroquinone (Eastman) were recrystallized from ethanol and diethyl ether and stored at 0 "C. Acetate buffers (0.01 rnol dm-3) were used to control acid concentrations between pH 3.5 and 6, with appropriate amounts of HClO, employed at higher acidities. In some instances at higher pH's, phosphate buffers were used with no significant differences observed in rates. Sodium (lithium) perchlorate was added to maintain the ionic strength at 0.10 (1.0) rnol dm-3.Nitrogen saturated solutions of the reductants were prepared just prior to use. Kinetic Measurements.-The kinetic experiments were carried out using a stopped-flow apparatus described pre- viously.* The reduction of [NiLI2+ (ca. 5 x lo-' rnol dm-3) was followed at 500 nm under pseudo-first-order conditions of excess reductant ( I 3-5) x rnol dm-3. At pH < 4.5 plots of In ( A , - A,) against time were linear for at least three half- lives. Above pH 4.5 the decay curve was biphasic, with approximately equal decreases in absorbance at 500 nm for each step. The rate constants for the initial process were ca. 15 times greater than those of the second step. Rate data for the first reaction were derived from the first 70:< of reaction. For the slower step, rate calculations were performed on the final 80% of reaction, where there is no interference from the first step.Results The stoicheiometries of the reduction of [Ni1"Ll2+ (4.0 x mol dm-3) by H2cat, H2quin, and H2mquin [(I-10) x rnol dm-3] were determined by spectrophotometric titra- tions (500 nm) at pH 4.3. The ratios of nickel(iv) consumed to reductant added were, for H2cat 0.96 & 0.04 : 1 .OO, for Hz- quin 0.98 f 0.04 : 1.00, and for Hzmquin 1.01 k 0.04 : 1 .OO. The overall reactions may therefore be expressed by equation t Present address: Department of Chemistry, Brookhaven National Laboratory, Upton, New York, U S A . (I), where Q is the corresponding quinone.104 I. CHEM. SOC. DALTON TRANS. 1984 Table 1. Rate constants" for the reduction of the nickel(1v) (k,) and nickel(1rx) (kb) oxime complexes by benzene diols: T = 25.0 "C, I = 0.10 mol dm-3 (NaClO,) Reductan t H,ca t Hzquin Hzmquin Hzmquin PH 3.37 4.20 3.37 3.80 4.25 3.47 3.86 4.24 4.63 (4.83 3.49 3.61 3.81 3.87 4.21 4.47 4.50 4.6 I 1 O-'ko/ dm3 mol-' s-' pH 1.09 I .47 3.43 3.53 3.88 6.03 6.74 9.30 14.0 15.6)* 8.44 8.94 9.30 9.50 11.84 13.4 13.9 16.2, 4.87 5.1 I 5.30 5.58 5.78 5.96 4.75 5.25 5.55 4.69 5.12 5.60 5.64 5.83 6.08 6.19 6.445 4.98 5.30 5.62 5.92 4.89 5.05 5.36 5.85 10-3k,l dm3 mol-' s-' 1.42 3.56 6.66 10.80 15.70 0.45 I .35 3.7 1 5.80 11.0 1.99 3.30 7.10 14.1 1.90 2.45 3.80 4.80 1 O-'kb/ dm3 mol-' s-' 0.70 0.79, 0.8 1 0.91 I .05 1.22 0.76 0.84 0.9 1 1.84 1.91 2.09 2.26 2.22 2.59 2.59 3.44 5.65 5.75 5.97 6.38 6.30 6.10 6.55 ' [Ni'"] = (2.1-5.7) x rnol dm-3.Reductant concentrations in the ranges (1.784.64) x rnol d n P (H2cat), (1.95-3.81) x 1O-j rnol dm-' (H,quin), and (2.73-3.37) x 1O-j mol dm-3 (H2mquin). I = 1.0 rnol dm-3 (LCIO,). Phosphate buffer. * Though the pH > 4.7, the rate difference was such that we could not see a definite second process. Pseudo-first-order rate constants, measured using excess [HzQ], showed a first-order dependence on the reductant concentration. The overall second-order rate constants, measured at 25 "C over the entire pH range are presented in Table 1. Below pH 4.7 the decay of the [Ni"LIz+ absorbance at 500 nm is monophasic. The rate law for the stoicheiometric two-electron reduction is given by equation (2). The second- -d[NiL2+]/dr = k0[NiL2+][HrQltot (2) order rate constants, ko, increased with decreasing acid concentrations.In additional experiments carried out at much lower pH's, it was found that the hydrogen-ion dependence was not a simple inverse relationship. In some instances there was evidence for a slight induction period, indicating other possible reactions. Above pH ca. 3.5, however, good first- order decay curves were observed. Above pH ca. 4.7 the decay of the absorbance at 500 nm separates into two stages. As the pH is increased the faster initial step is more easily identified from the slower second step. The absorbance changes in the two steps are roughly 15 10-5[Ht]-'/dm3 mol-' Figure 1. Inverse acid dependence of k, for the reduction of [Ni"'LIz+ by catechol (O), hydroquinone ((3 ), and 2-methyl- hydroquinone (0) equal over the range h = 430-550 nm.Lappin and Laran- jeira ' reported the spectrum of a nickel(iri) intermediate in the reduction of [NiL]'+ by [Co(phen),]'+ with maxima at 505 and 398 nm and E values of 2 890 and 3 000 dm3 mol-' cm-', respectively. The intermediate observed in this study appears to be the same Ni"' complex. The rate laws for the consecutive (one-electron) reductions of the nickel(iv) and nickel(rii) species are given in equations (3) and (4). The k , and kb values measured between pH 4.6 -d[NiLZ+]/dt = 2k,[NiL2+][HzQllOt - d[Ni"']dt = 2kb[Ni1"][HzQ]tot (3) (4) and 6.1 at 25 "C are presented in Table 1. In this range both the [Ni'"LI2+ and [Ni"'L]+ reactions exhibit [H+]-' depend- ences (see Figures 1 and 2) although for the [Ni'"L]+ reduc- tion a hydrogen-ion independent pathway is also observed.Studies were also made at higher ionic strength (1 .O mol dm-j), the most extensive using 2-methylhydroquinone as reductant. Rates are generally ca. 10% higher than at I = 0.1 mol dm-3 but the general behaviour observed is similar to that described above (see Table 1). Discussion The results of the present kinetic study indicate that at low pH (<4.5) theone-electron reduction of [NiLIZ+ by the dihydroxy- benzene compounds is slower than the subsequent reduction of the Nil" intermediate, while at higher pH [NiL]'+ is more rapidly reduced than the Ni"' species. These observations are the result of protonation equilibria involving both oxidant and reductant species. With the reductants employed in this work the 'crossover' occurs between pH 4 and 5. In a kinetic study using [Co(phen)#+ as a reductant, Lappin and Laranjeira' reported biphasic reactions at pH > 3.For the reductions of [NiL]'+ by metal ions such as Fez+ in acidic solutions (pH < 2), no nickel(iii) intermediate was detectable using stopped-flow spectrophotometry. It is interesting to noteJ. CHEM. SOC. DALTON TRANS. 1984 105 Table 2. Proton equilibrium constants and reduction potentials for the benzene diols Reductant P K a PKZ PKr, PKr, El dlv E2 "V E3 'IV H,cat ' 9.24 13.0 ca. -1 5.0 1.12 0.49 0.043 Hzquin 9.85 11.4 ca. - 1 4.1 1.09 0.45 0.023 Hzmquin 10.05 11.6 ca. -1 4.4 I .03 0.38 - 0.05 Refs. 11, 16, and 17. Ref. 11. data therein. I = 1.0 mol dmP3 (NaC104). I = 0.65 mol dm-3 (NaC104).Refs. 1 I and 16. ,I Calculated using equations (30) and (31). From ref. 16 or extrapolation using Table 3. Comparison of observed and calculated specific rate constants Oxidant Reductan t AE"/V kJ mol-' H,cat - 0.47 46.5 H,quin - 0.44 43.6 H,mquin - 0.38 37.70 Hcat - 0.15 - 12.5 Hquin - 0.20 - 17.3 Hmquin - 0.27 - 26.1 H,cat - 0.48 47.1 H,quin - 3.45 44.4 i [Ni'"L]'+ H,mquin - Hcat - [Ni' "(HL)l2 + 0.39 0.14 - 38.6 11.5 Hquin - 0.19 - 16.3 Hmquin - 0.26 -25.1 I Hcat - - 0.08 8.70 Hquin - - 0.03 3.9 Hmquin - 0.04 - 2.85 [Nil"L] + AGobr. *I kJ mol-' 52.4 49.0 47.3 20.9 20.2 17.6 22.7 17.9 14.0 34.0 31.5 29.2 k0bS.l dm3 mol-I s-I 60 270 480 4.8 x 107 6.4 x 107 1.8 x 108 1.5 x 107 1.7 x lo8 7.9 x 108 1.1 x 105 4.1 x 105 1.0 x 106 kca1c.l dm3 mol-' s-' 30 70 3 60 6.8 x lo7 1.6 x lo8 6.4 x lo8 6 13 65 1.1 x 107 2.6 x 107 1.3 x 105 3.6 x 105 1.0 x lo8 1.4 x lo6 that in cyclic voltammetry studies of the nickel-oxime system the separation of a single two-electron pattern ([Ni1"LI2+ + 2e- + 2H+ * [Ni''(H2L)l2+, E* = 0.95 V} at low pH into two one-electron processes at higher pH occurs in the region of pH 4-5.4*6*7 Reductions of [NitVLI2 + .-Kinetic and electrochemical studies have shown that the nickel(iv) species is [NiL]'+ at pH > 0 with pK < - 1 for the protonation of [NiLI2+ in aqueous solution, [NirvLI2+ + H + [NitV(HL)l3+.The dependence of k, on [H+] has been attributed, therefore, to proton equilibria involving the reductants. In neutral and acidic media the benzene diols are undis- sociated species (H2Q).In general pK1 lies in the range 9-10 and the second dissociation takes place with pK2 in the range 11-13 (Table 2).9 A reaction scheme for the reduction of [NiL]'+ by hydroquinone and the other reductants in the pH range 3.5-6 may be expressed by equations (5)-(7). H 2 Q S HQ- + H+ ( 5 ) [NiLI2+ + H2Q k,c [NiL]+ + HzQ'+ (6) [NiLI'+ + HQ- k,_ [NiL]+ + HQ' (7) If H2Q'+ and HQ' react rapidly with [NiLI2+, then the rate law for the reduction of [NiL]'+ according to the mechanism is given in equation (8). In the pH range under consideration where K1 < [H+], k, may be expressed in the form given in equation (9). (9) k, = k, + k,K,[H+]-' Above pH 4, k, exhibited a linear inverse dependence on [H+], as shown in Figure 1. The specific rate constants kl and k2 were derived from the slopes and intercepts of the [H +]-I dependence plots.The rate constants for the reaction of [NiL]'+ with the three benzene diols and their anionic forms are presented in Table 3. The greater reactivity of HQ- relative to H2Q, k2/kl CCI. 105-106, has been observed in previous kinetic studies of the oxidation of hydroquinone in neutral solutions.".12 Reduction of Ni ' I t Species.-The separation of the nickel(1rr) reduction step (above pH 4.5) from the rapid initial reaction of [NiL]'+ is sufficiently great that studies for this reaction may be made. Pseudo-first-order rate constants were measured between 4.6 and 6.1. An inverse acid dependence was again observed, attributable to the HQ- oxidation pathway. The oxime protons on the nickel(Ir1) complex are less acidic than those on nickel(1v) and a pK, of d 3.9 has been determined for the equilibrium (10).[Ni(HL)]'+ & [NiL] + + H + (10) A reaction scheme for the reduction of the nickel(rr1) inter- mediate is outlined in equations (1 1)-(14). [Ni(HL)I2+ + H2Q k,_ [Ni(HL)]+ + H2Q'+ (11)106 - - J . CHEM. SOC. DALTON TRANS. 1984 640 600 I 4 8 12 16 20 24 ' 2 8 32 [ H']-'/drn 3mol-' Figure 2. Plots of k, against [H+]-' in the reduction of [NiI'IL]+ by catechol (O), hydroquinone ( (3 ), and 2-methylhydroquinone (0) [equation (17)] [NiL]+ + HzQ k,_ NiL + HzQ" (13) [NiL]+ t HQ- k,_ NiL 1- HQ' (14) The semiquinone radicals produced in equations (6) and (7) and (11)-(14) react rapidly with a second Ni'" or Nil" ion. If the radicals decay in part via disproportionation *l [equation (1 5)] then the simple stoicheiometry observed would not be expected. The rate law for the disappearance of Ni"' is given in equation (16).HQ' + HQ' * H2Q + Q (15) The work terms wlZ and w21 correspond to the energy associ- ated with bringing the reactants and products, respectively, to a separation distance r in the activated complex. The terms wll and w22 similarly apply to the self-exchange reactions. In aqueous solution at 25 "C and an ionic strength of 0.10 mol dm-3, the Debye-Huckel expression for the work term l 4 is given by equation (22). 4.242122 r(1 + 0.104rdZ) w1z = In order to calculate the specific cross reaction rate con- stants for the reactions a knowledge is necessary of the reactant reduction potentials and rate constants for the individual self-exchange processes.Both the reductants and oxidants in this study have pH dependent redox equilibria. The electrochemistry of the Ni",Ni"',Ni'"-oxime system has been investigated by means of cyclic voltammetry tech- n i q u e ~ . ~ * ~ - ~ Below pH 5 a single two-electron wave is observed with a potential of 0.95 V (uersus normal hydrogen electrode) [equation (23)]. [NiiVLl2+ + 2e- + 2H+ =G= [Ni1I(HZL)l2+ (23) Above pH 5, two one-electron processes are identified. The equilibria of interest in this study [equations (24)-(26)] have [Ni111(HL)]2+ + e- + [Ni"(HL)]+ (25) [Ni'''L]+ + e- [NiIIL] (26) potentials of 0.65,4 0.64,' and 0.42 V,4 respectively. At higher acidities (0.1-1.0 mol dm-7 the couple [Ni(H2L)]3+/2+ For the evaluation of the specific rate constants it was assumed that k4/k3 and k6/k5 were in the region of lo5-lo6 as observed previously." In the reduction of [Nil"(HL)IZ + and [Ni"'L] + by [Co(phen)J2 + and ascorbic acid, Lappin and Laranjeira observed that [Ni(HL)]'+, a stronger oxidant (Ee = 0.64 V), was 20-40 times more reactive than [NiL]+ (Ee = 0.42 V). If only k4 and k6 contribute to the reaction rate, the rate law may be expressed in the form (17) under conditions of higher pH where Ka > [H+].The hydrogen-ion dependence of the rates is shown in Figure 2. The specific rate constants (Table 3) are again large (ca. 105-108 dm3 mol-' s-I). The rate constant for an outer-sphere electron-transfer reaction may be predicted by use of Marcus theory equation^.'^ This correlation establishes a relationship between the free energy of reaction (AGlze) and the free energies of activation for the cross reactions (AGI2*) and reductant and oxidant self-exchange reactions (AGll* and AG22*).(E* = 1.23 V) has been shown to be important in redox processes.6 Contributions at pH - 1 from this reactant may be a reason for the more rapid Ni"'-Ni'' reduction under these conditions leading to the observation of only one process. The redox equilibria involved in the hydroquinone/semi- quinone self-exchange reactions are outlined in equations (27)-( 29). H2Q El"_ H2Q'+ + e- (27) +H+ 11 pKl +H+ 11 pKr, HQ- EI"_ HQ' + e- + H e 11 PK, +H+ 11 PKr2 QZ- E3O T- Q'- + e-J . CHEM. SOC. DALTON TRANS. 1984 107 The potentials for the H2Q'-/H2Q and HQ'/HQ- couples may be determined from the proton equilibrium constants and the Q'- /Q'- reduction potentials using equations (30) and (31). The E," and pK values 10.16 employed in equations (30) and (3 1) were determined experimentally or extrapolated from values for similar compounds using a linear relationship between E" and pK.1°.17 The equilibrium constants and potentials for the reductants are listed in Table 2.Previous studies of the oxidation of hydroquinone and related complexes have resulted in an estimate of AGll* for the H2Q'+/H2Q self-exchange reaction of 18.8 kJ mol-' (k,l = 5 x lo7 dm3 mol-' s-').'O A similar value was assumed for the HQ'/HQ- couple. The self-exchange rate constants for the NiIV-Ni"' and Ni'"-Ni'' reactions have been determined independently in this laboratory and by Lappin and Laran- jeira with good agreement.For the [NiL]'+'+, [Ni(HL)]'+'+, and [NIL]+'' couples, exchange rate constants of 6 x lo4, 2 x lo3, and 1 x lo3 dm3 mo1-l s-', respectively, have been employed. The calculated specific rate constants for reductions of the Ni111 and Nil" oximes by the benzene diols are presented in Table 3. Reasonable agreement, within an order of magnitude in most cases, is observed between the experimental and calculated values. The discrepancies may be due to un- certainties in the reduction potentials and self-exchange paranieters of the species involved. Calculations have been made of the rate for reaction (k,) [equation (1 l ) ] between the protonated complex [Ni"'(HL)]'+ and the neutral sub- strates H2Q.The values obtained are generally <lo2 dm3 mol-' s and inspection of the data reveals that inclusion of pathway ( I 1 ) would contribute <lo% towards the acid- independent parameter in equation (17). It is of interest to note that, as with the nickel(1v) systems, the rates of reaction of the anionic form of the reductant are ca. lo6 times greater than for the neutral species. Substrates of this type are well known as electron-transfer agents in naturally occurring systems.1B The rapid rates of redox reactions involving the anions in the present study confirm the importance of dissociation pheno- mena, not only of the negatively charged reductants but also of the anion radicals formed as the immediate reaction products. Further studies are in progress on the reactions of hydroquinone and related compounds with other Nil", oxime complexes and Ni"' monoxime species.These investigations may provide more information on the nature and reactivity of the NilV and Ni"' species in solution. Acknowledgements One of us (D. H. M.) thanks the University of Victoria for the award of a graduate Fellowship. Financial support from N.S.E.R.C. (Canada) is acknowledged. We thank Dr. Graham Lappin for making available to us data prior to publication and for continuing discussions in this area. We are grateful to Lee Spencer for experimental assistance. References I R. I . Haines and A. McAuley, Coord. Chem. Rev., 1981, 39, 2 K. Nag and A. Chakravorty, Coord. Chem. Rev., 1980, 33, 3 J. G. Mohanty, R. P. Singh, and A. Chakravorty, Inorg. Chem., 4 J. G. Mohanty and A. Chakravorty, Inorg. Chem., 1976, 15, 5 A. G. Lappin, M. C. M. Laranjeira, and L. Youde-Owei, J. 6 D. €1. Macartney and A. McAuley, Can. J. 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Received 15th December 1982; Paper 212092