摘要:

1979 735Kinetics of Oxidation of Hydrogen Azide (Hydrazoic Acid) by Tris(22'-bipyridine)nickel(Iii) Ions in Aqueous Perchlorate Media : Comparisonwith Oxidation by Aqua-cationsBy John K. Brown, David Fox, Malcolm P. Heyward, and Cecil F. Wells," Department of Chemistry,University of Birmingham, Edgbaston, P.O. Box 363, Birmingham B15 2TTE.s.r. measurements on solutions of Nil1 containing 2,2'-bipyridine (bipy) confirm that anodic oxidation producesNi"I. The stoicheiometry and kinetics of oxidation of HN, by [Ni(bipy),I3+ have been investigated. Unlike theoxidations of H202 and Br- by this cation, with HN, the rate varies with the concentration of HCIO, a t constantionic strength. Under conditions where specific cation effects are eliminated, the rate is directly proportional to[H+]-l, suggesting that the kinetically active entity of the hydrazoic acid is N3-.No removal of bipy from theNi"' occurs before the slow oxidative step, even though an outer-sphere intermediate complex may be formedbetween [Ni(bipy),I3+ and N,-. This system is compared with the oxidations of HN, by aqua-cations and withthe oxidations of other ligands by [Ni(bipy),l3fTHE oxidation of hydrazoic acid (hydrogen azide, HN,) byaqua-cations Mn+(aq) has been investigated for Mn+ =MnIII,l COIII,~J and CeIV.4 For MnIII the rate is firstorder in [MnIII] and second order in [HN,] and the rateincreases with increasing [H+] at constant ionic strength.lThis shows that the reaction proceeds through a complexformed between Mn3+(aq) and HN, and that [Mn(0H)l2+-(as) is unreactive in the 0xidation.l At high initial[HN,]/[Mn+], the reaction with CoIII resembles that withMnIII under similar conditions in being first order in[CoIII] and second order in [HN,],2 but it differs fromthe reaction with MnIII in being zero order in [H+].Atlower initial [HN,]/[Mn+] for CoIII, the reaction is firstorder in [HN,] as well as in [CoII[] and the rate is pro-portional to a + b[H+]-l, where a and b are constants a tconstant temperat~re.~ The reaction with MnIII isrelatively slow and the rates are entirely accessible toconventional spectrophotometric techniques, but for thefaster reaction with CoIII stopped-flow techniques haveto be used except at low temperatures. However, forthe much faster reaction with CeIv, only stopped-flowmeasurements are possible, and the reaction is first orderin [CeIV] and zero order in both [HN,] and [H+].4 In allcases for Mn+ (as) ,l-* the overall stoicheiometry is repre-sented by (1).2Mn+(aq) + 2HN, - ZM(?h-l)+(aq) + 3N, + 2H+(aq) (1)By comparing the rates of oxidation of various ligandsL by ColI1 at one temperature it has been argued 5 9 6 thatthey are all controlled by the rate of substitution of Linto the inner co-ordination sphere of CoIII, i.e.the raterepresents the rate of removal of a water molecule fromthis inner aqua-sphere. That this is an oversimplific-ation is shown by the very significant differences' inthe enthalpies AHT and entropies AS: of activation withvarying L and by the observation that intermediate(CoIII + L) complexes participate in the rate-deter-mining redox steps for some of the L cited: such com-plexes have been detected spectrophotometrically andkinetically for L = H202 * and kinetically for L = HN, 2and It has been pointed out 9-11 that, since bothredox processes and substitution processes of cationsinvolve considerable movement of water molecules toachieve their respective transition states, it is notsurprising that AH'.and AS'. (and therefore sometimesthe rates) for the two types of process are of similarmagnitude even if the redox process is not controlled bythe substitution process. To investigate in a moregeneral way the extent to which the water moleculesin the inner co-ordination spheres of cations are involvedin cation + ligand redox processes, the kinetics of theoxidation of H202 lo and Br- ions l1 by the tris(2,Z'-bipyridine)nickel(m) ion have been compared with thekinetics of the oxidation of these ligands by aqua-cations. The kinetics of the oxidation of hydrazoic acidby the [Ni(bipy)J3+ are now described and comparedwith the kinetics of the oxidation of HN, by aqua-cations,EXPERIMENTALMateriaZs.-Tris(2,2'-bipyridine)nickel(111) perchloratewas prepared by anodic oxidation under nitrogen.1° Theanode compartment consisted of a glass tube (length, 10cm; diameter, ca.3 cni) separated a t its base from thecathode compartment by a no. 4 glass sinter. The elec-trodes were platinum, the anode being of area ca.3.2 cmz,and the current was normally 50 mA a t 36 V. Nickel(I1)perchlorate was dissolved in 2 mol dm-, HCIO,, and thewhole electrolytic apparatus was maintained a t ca. 5 "C.Such a solution soon achieved the maximum yield of NiIIIwhich was always low: 1.33 mol dm-, NiII in 2 mol dm-3HC10, gave 2.5 x lo-, mol dm-3 NiIII after the current hadpassed for 10 min and this did not change after passage ofcurrent for 50 min (some precipitation of the bipyridine-nickel(r1) salt occurred a t this concentration); 3.2 x 10-4in01 dm-3 Nil1 in 2 mol dm-3 HC10, gave 1.2 x rnol dni-3NiIII after 10 min, and this did not change after passage ofcurrent for 40 min. The spectrum of lime-green [Ni-(bipy),13+ is shown in Figure l .AnalaR perchloric acid was used in the reaction mixtures ;solutions of sodium and lithium perchlorates were preparedby neutralization of HClO, with Na2[C0,] and Li,[CO,] (allAnalaR).G. F. Smith magnesium perchlorate was used.Solutions of sodium azide were made up by weight. Waterwas distilled once in an all-glass stillJ.C.S. DaltonProcedure.-Practically all the kinetic measurementswere made using a Durrum-Gibson stopped-flow spectro-photometer; all those at I = 2.00 mol dm-3 using addedLi[ClO,] were done on this instrument. Some of the otherrate measurements employed a Unicam SP 500 series 2spectrophotometer. All the rate measurements werecarried out a t 350 nm: the product of the oxidation in thestopped-flow experiments is [Ni(bipy),12+ and interferencefrom its absorption at this wavelength is negligible.1°Errors arising from the oxidation of water by [Ni(bipy),I3+on these time scales can be neglected.10 The concentrationsof solutions of [p\Ti(bipy),l3+ were determined as describedearlier.lOE.s.r.spectra were run on a Hilger and Watts MicrospinESR-3 spectrometer operated a t X-band frequencies.Determinations of the g value, were made by superimposingthe resonance from diphenylpicyrylhydrazyl (dpph) on thetraces from the complexes and from proton resonance metermagnetic field calibrations.X / n mSDectra of NiI1 and N W in 2 mol dmP3 HC10, forpath lengtA = 4 cm: (a) 3.0 x rnol dm-3 [Ni(bip+)J3+containing ca. 1.3 x mol dm-3 Nil' as decomposed [NP-(bipy),I2+; ( b ) 1.33 x rnol dm-3 Nilx as decomposed"i(biPY) 31 2+E.S.R.Measurements.-E.s.r. spectra were obtainedfrom the bipyridinenickel complexes in frozen perchloricacid-water solution at 77 K. Comparison of the spectrarecorded before and after electrochemical oxidationallowed an unambiguous identification of the mostpronounced feature of the product of anodic oxidation.This resonance had a near symmetric lineshape of maxi-mum slope width ca. lOmT and g 2.095 & 0.002. Anyother possible weak features at higher fields which wouldhave indicated g anisotropy could not be discernedbecause of interference from the resonance from un-oxidized NiII present in excess.For a number of nickel complexes there has beendiscussion as to whether their e.s.r. spectra are to beunderstood as originating from low-spin d7 NiIII or as anickel(I1)-stabilized radical-ligand s y ~ t e m .l ~ - ~ ~ Thereare a number of well authenticated examples where thestabilized radical-ligand structures are associated withhighly isotropic g values close to that for the free-electronspin (2.002 3); in contrast, the low-spin d7 complexesRESULTS AND DISCUSSIONhave anisotropic g tensors and significantly higher {g)~ a l u e s . l ~ - ~ ~ - l * Although we have no definitive evidenceof g anisotropy, on the grounds of the g value observedwe propose that the oxidized bipyridine complexesshould be formulated as being based on NiIII.Stoicheiometry .-Working with an excess of [NiIII](ca. 2.5 x lo6 mol dm4) over hydrazoic acid (1.00 x10" mol drn-,), the consumption of NiIII was measuredfor a range of acidities assuming that all the hydrazoicacid had disappeared.The observed consumptionvalues ]A[NiIII]l/lA[HN,]I in Table 1 show that thestoicheiometry conforms to equation (1), analogous to theoxidations of hydrazoic acid by the aqua-cations.TABLE 1mol [HC10,1 d n r 3 I " " i T I p H ~ 1 I1 .oo 1.102.00 0.943.00Mean = 1.01 f 0.1Rate Measurements.-The rate of reaction was firststudied at I = 5.00 rnol dm-, with [HN,] > initial[NiIII] over a range of temperature. Good linear plotsof log(optica1 density) against time were obtained withinitial [NiIII] - 1 x lo4 mol dm-3. At 25 "C a plotjofthe pseudo-first-order rate constant k,, against [HN,]over the range 0.05-0.25 mol dm-3 in 4.00 mol dm-3HClO, was linear passing through the origin, showingthat the reaction is first order in [HN,].However, intesting the dependence of rate on acidity, plots of thesecond-order rate constant k, against [H+]-, were linearand passed through the origin at 2.5, 13.0, 25.3, and32.1 "C. Since the mechanistic interpretation of thelatter variation is clearly difficult, a test for specific ioneffects was carried out. The results for equivalentadditions of Li[ClO,] and Na[C10,] at constant [HClO,]for I = 5.00 mol drn-,, summarized in Table 2, showclearly that there is a specific cation effect. Due to theTABLE 2Variation of rate with nature of added metal perchlorateat constant ionic strength; [HN,] = 2.00 x moldm-3k , - [HClO,] I dm3 - -mol mol & Added salt mol-1 kzW+ldm-3 dm-3 "C (clmol dm-3) S-1 S-15.00 5.00 25.34.00 5.00 25.3 Li[ClO,] (1.00)Na[C10,] (1.00)3.00 5.00 25.3 Li[ClO,] (2.00)Na[C104] (2.00)2.00 2.00 29.71.00 2.00 29.7 Li[C104] (1.00)Na[C104] (1 -00)1.00 1.00 29.70.50 1.00 29.7 Li[ClO,] (0.05)Na[C10,] (0.05)Mg[C1041, (0.33)0.71 3.550.78 3.121.18 4.720.98 2.942.26 6.878 156153 153203 203322 32278 78159 80153 77low solubility of Li[ClO,], rates were then measured atI = 2.00 mol dm-3 at 29.7 "C when Mg[ClO,], was alsoused as the added salt.Table 2 shows that the accelerat-ing effect of the cation is in the order Mg2+ > Na+ 1979 737Li+. However, it can also be seen that the rates at29.7 "C with added Na[C10,] or Li[ClO,] at I = 1.00 rnoldm-3 are not dependent on the cation.Since Table 2shows that k,[H+] is constant, at I = 1.00 mol dm-,,independent of the added cation H+, Li+, or Na+, and2.0 t1 2 3102[HN3]/rnol dm-3with [ H N , ] a t I = 2.00 mol dm-3 (Li[ClO,]).[(O) at 29.7, (a) a t 17.0 "C] or 1.00 mol dm-, (17.0 "C) (0)FIGURE 2 Variation of the pseudo-first-order rate constant k ,[HClO,] = 2.00J0.5 1.0 1.5 2.0[ H+]-' / dm mol-'Variation of the second-order rate constant k 2 with[H+]-l at 1 = 2.00 rnol dm-, (Li[ClO,]) and 24.7 (O), 29.7 (a),or 35.5 "C (a)FIGURE 3that k,[H+] is constant at I = 2.00 mol dm-3 for thecations Li+ and H+, a detailed kinetic investigation wascarried out at 29.7 "C at I = 2.00 mol dm-, maintainedby the addition of Li[ClO,].In all these cases linearplots were obtained for log(optica1 density) against time.Figure 2 shows that a plot of k, against [HN,] in2.00 mol dm-, HClO, at 29.7 "C is linear passing throughthe origin, and the values for k, over a range of acidityat I = 2.00 mol dm-3 (Li[ClO,]) are collected in Table 3.Figure 3 shows that a plot of k, against [H+]-l at I= 2.00 mol dm-3 (Li[ClO,]) and 29.7 "C is linear. Linearplots for log (optical density) against time were obtainedfor I = 2.00 mol dm-, with added Li[ClO,] over arange of temperature, and Figure 2 shows that plots forTABLE 3Values of K , and tz, at I = 2.00 mol dm-3 (Li[ClO,])0,"C17.0-24.729.735.5[HC104I 103[HN,] k2 - k3mol dm-3 mol dm-3 dm3 mo1-ls-l s-l2.002.002.002.001.001 .oo1 .oo1.002.002.001.601.601.001.000.500.502.002.002.002.001.601 .oo1 .oo1 .oo0.508.016.032.040.04.08.032.040.010.05.005.002.505.002.505.002.5020.016.010.01 .oo4.008.004.001.001 .oo27.425.526.222.461595047.8Mean51546868100100208200Mean :837684689014616815232955515244.861595047.853 6102108109109100100104100104 f 4166152168136144146168152165Mean = 155 f 122.00 5.00 101 2022.00 2.50 95 1901.60 5.00 124 1981.60 2.50 120 1921 .oo 2.50 196 1961.00 1.25 180 1800.50 2.50 396 198Mean = 194 f 7k, against [HN,] for 1.00 and 2.00 rnol dm-, HClO, at17.0 "C are linear passing through the origin.FromTable 3 it can be seen that values for k, are constant forvarying [HN,] at constant acidityat both 24.7 and 35.5 "C,providing. final confirmation that the reaction is firstorder in [HN,] as well as first order in [NiIII]. Plots ofk, against [H+]-l at 24.7 and 35.5 "C (Figure 3) are alsolinear passing through the origin, and Table 3 confirmsthis inverse dependence of rate on acidity at I = 2.00mol dm-3 with added Li[ClO,] for all the temperatures.Figure 4 shows that a plot of log k, against 1/T is linear(where k, = k,[H+]). From the gradient in Figure 4,determined using the least-squares procedure , theenthalpy of activation AHZ = 51 5 3 kJ mol-l; theentropy of activation A S = -36 10 J K-l mol-l.Mechanism of the Oxidation of HN, by [Ni(bipy),I3+.-Although the oxidation of HN, by [Ni(bipy),]3+ i738 J.C.S.Dalton[HN,]/[CoIII] and by MnIII(aq),l the rate-determiningstep is (6) followed by a rapid (7), and in the oxidation by[Mn+,HN,] + HN, + M(n-l)+(aq) + [H2N6]*+M(?b--l)+(aq) + 3N, + 2H+(6)(7)Mn+(aq) + -w M(n-l)+(aq) + N,* (8)M n ( 4 + [H,N,I+ -ColI1(aq) at low [HN,]/[Co'll] the rate-determining stepis (8),, which may or may not involve an intermediatecomplex and which is followed by rapid (4). The kineticsof the oxidation of hydrazoic acid by CeIV(aq) do not tellus the form of the substrate ligand in the transition state,but they do require the reaction to go through an inter-mediate (cation + ligand) complex.The conclusionfrom the above, combined with the new informationfrom the oxidation by [Ni(bipy),I3+, appears to be thatthe oxidation of HN, requires two HN, molecules in thetransition state, whereas oxidation of the azide ionrequires only one. Since the kinetic requirements a thigh [HN,]/[Mfi+] for MnIII(aq) and CoIII(aq) is that, forIs Mn+(aq) + HN, * [Mn+,HN,l(aq) (9)(9), p[HN,] < 1, then at low [HN,]/[Mn+] for CoIII(aq)any intermediate complex must have a very low concen-tration indeed, suggesting that, under the latter condi-tions, such an intermediate is not involved. However,the path involving the intermediate complex of HN, isenergetically preferred with CoII1( aq).2 Since [N,] -exists in such low concentrations in these strongly acidicsolutions, the possibility of concerted electron and protontransfers taking place in a solvated outer-sphere { [Ni-(bipy),I3+ + HN,) complex cannot be excluded.AH: and AS: are concerned with the rearrangement ofthe redox system and solvation in the transition stateto allow electron tunnelling to occur between the ligandand the cation.The overall AH: (kJ mol-l) for HN,with aqua-cations, CeIV (zero) * < CoIII (53)2 < MnIII(SO),l does not follow the reverse order of their redoxpotentials (V), CoIII (1.95) 23 > CeIV (1.76) 24 > MnIII(1.56),25 presumably because of the very high p for(9) with Mn+ = CeIV. The overall A S reflects thebalance between the lowering of restriction on the solventin the transition state (AS$,) to accommodate the lower-ing of charge on the cation and the increase in restrictionimposed by any proton ejected (ASXp).10 Thus, al-though the overall AH1 for HN, with CoIII(aq) and [Ni-(bipy),13+ are very close (53 and 51 k J mol-l), the transferof the proton to the solvent in the latter case producesgreater restriction and a negative overall A S comparedwith an overall AS: = 7 J K-l mol-l for the protoniccharge developing on [H2N6]+ with ColI1(aq).Comparison with Oxidation of Other Ligands by[Ni(bipy),l3+.--For oxidation by [Ni(bipy),I3+, AS: =-3 J K-1 mol-1 for the rate-determining steps withboth [N3]- and Br- without proton ejection beinginvolved, but AS: = -126 J K-l mol-l for H202 withlA.SIp1 > IASfel when a proton is transferred to solvent.sensitive to specific cation effects at high ionic strengths,unlike the oxidation by CoIII(aq),3 it seems clear that therate is inversely dependent on [H+].For such a depen-dence in an oxidation by an aqua-cation there are anumber of possibilities to account for the variation with[H+] : acid dissociation of HN,, hydrolysis of the aqua-cation, or dissociation of a proton from an intermediatecomplex. As in the oxidation of H202 lo and Br- l1 by[Ni(bipy)J3+, intermediate complexes are not detectedwith HN,: for H,0210 and Br-,ll the oxidations by[Ni(bipy),I3+ are independent of [H+], confirming thatthe acid dependence observed for the oxidations ofthese ligands by aqua-cations must arise from either ofthe last two possible causes given above.With HN,such a proton dependence operating on [Ni(bi~y),]~+ isnot possible since there are no water molecules in thefirst co-ordination sphere and the rate of removal ofbipy ligand is directly dependent on [H+].20 However,12 . 5 1I I 13 *2 3.3 3 4 3.5I O ~ K 11FIGURE 4 Plot of the logarithm of k3( = k,[H+]) against 1/Tthe experimental inverse dependence shows that aproton is removed from the reactants in the transitionstate. Although the azide must be present as HN, inthe bulk solution at these acidities, since the acid-dissociation constant of HN, at I = 2.00 mol dm-3 isca. I x lo4 mol dm-, at 20 0C,21 the form in the transi-tion must be [N,]- as in (2)-(4) with k'> k.TheRate = kK[Ni(bipy)33+][HN3][H+]-1 (5)rate is then given by (5), conforming to the observedkinetic orders, with k, = kK. Taking the experimentalvalues for AH: and AS$ above for k,, andusing AHe =15.1 k J mol-l and AS* = -32.6 J K-l mol-l found 22for the acid dissociation (2) , for the rate-determiningstep (3) we obtain AHX = 36 & 3 kJ mol-l and AS$ =-3 -+ 10 J K-1 mol-I.Comparison with Oxidation of HN, by Aqua-cations.-For the oxidation of hydrazoic acid by ColI1(aq) at hig1979AH: (kJ mol-l) for these ligands with [Ni(bipy),I3+,H202 (38) lo < HN, (51) < Br- (60),11 does not follow thereverse order of their redox potentials (V),2s Br-(1.087) > H202 (0.68) > HN, (-3.1); presumably, thelow AH$ for H,02 arises from a large negative contri-bution from AH:,.[7/1707 Received, 28th September, 19771REFERENCESC. F.Wells and D. Mays, J . Inorg. Nuclear Chent. Letters,1968, 4, 61; J . Chem. SOC. ( A ) , 1968, 1622.C . F. Wells and D. Mays, J . Chem. SOC. ( A ) , 1969, 2175; J .Inorg. Nuclear Chem., 1971, 33, 3855.R. K. Murmann, J . C. Sullivan, and R. C. Thompson, Inovg.Chem., 1968, 7, 1876; R. C. Thompson and J . C. Sullivan, ibid.,1970, 9, 1590.C. F. Wells and M. Husain, J . Chem. SOC. ( A ) , 1969, 2981.J. Hill and A. McAuley, J . Chem. SOC. ( A ) , 1968, 2405; M. N.G. Davies and K. 0. Watkins, J. Phys. Chem., 1970,74, 3388.A. McAuley and R. Shanker, J.C.S. Dalton, 1973, 2321. * J . H. Baxendale and C. F. Wells, Trans. Faraduy SOC., 1957,53, 800; C.F. Wells and M. Husain, ibid., 1971, 67, 760; C. F.Wells and D. Fox, J . Inorg. Nuclear Chem., 1976, 38, 107.C. F. Wells, A. F. M. Nazer, and D . Mays, J . Inorg. NuclearChem., 1977, 39, 2001.lo C. F. Wells and D. Fox, J.C.S. Dalton, 1977, 1498.Malik, J . Hill, and A. McAuley, ibid., 1970, 643.11 C. F. Wells and D. Fox, J.C.S. Dalton, 1977, 1502.12 A. H. Maki, N. Edelstein, A. Davidson, and R. H. Holm, J .13 E. I. Steifel, J. H. Waters, E. Billig, and H. B. Gray, J .l4 A. H. Maki, T. E. Berry, A. Davidson, R. H. Holm, and15 P. Kreisman, R. Marsh, J. R. Preer, and H. B. Gray, J . Amer.l6 A. Wolbergand J. Mannassen, J . Amer. Chem. SOC., 1970,92,1' A. Wolberg and J . Mannassen, Inorg. Chem., 1970, 9, 2365.l8 I?. V. Lovecchio, E. S. Gore, and D. H. Bush, J . Amer.l9 D. Sen and C. Saha, J.C.S. Dalton, 1976, 776.2o P. Krumholz, Nature, 1949, 163, 724; J. H. Baxendale andP. George, ibid., vol. 162, p. 777; ibid., vol. 163, p. 725; Trans.Faraday Soc., 1950, 46, 736; F. Basolo, J . C. Haynes, and H. M.Neumann, J . Amer. Chem. SOC., 1953, 75, 5102; 1954, 76, 3807;F. Basolo and R. G. Pearson, ' Mechanisms of Inorganic Reac-tions,' Wiley, New York, 1967, pp. 218-219 and 311-313.21 M. Quinton, Compt. rend., 1940, 210, 625.22 P. Gray and T. C. Waddington, Proc. Roy. Soc., 1956, A235,23 D. A. Johnson and A. G. Sharp, J . Chem. SOC., 1964, 3490.24 E. Wadsworth, F. R. Duke, and G. A. Goetz, Analyt. Chem.,25 H. Diebler and N. Sutin, J . Phys. Chem., 1964,68, 174.26 Handbook of Chemistry and Physics,' 56th edn., CRCAmer. Chem. Soc., 1964, 86, 4580.Amer. Chem. Soc., 1965, 87, 3016.A. L. Ralch, J . Amer. Chem. SOC., 1966, 88, 1080.Chem. Soc., 1968, 90, 1067.2982.Chem. Soc., 1974, 96, 3109.106.1957, 29, 1824.Press, Cleveland, Ohio, 1976, pp. D141-142

ISSN:1477-9226    

DOI:10.1039/DT9790000735

出版商:RSC    

年代:1979

数据来源: RSC