Kinetics of bromide oxidation by peroxo complexes of molybdenum (VI) and tungsten (VI)

ELSEVIER

InorganicaChimicaActa263 (t997) 225-230

Kinetics of bromide oxidation by peroxo complexes of molybdenum(VI) and tungsten(VI) Martha S. Reynolds *, Kristen J. Babinski, Michael C. Bouteneff, Julie L. Brown, Rachel E. Campbell, Mary A. Cowan, Michelle R. Durwin, Ted Foss, Patricia O'Bfien, Heather R. Penn Departmentof Chemistry, Colgate University. Hamilton. NF 13346, USA

Received4 February1997;revised7 April 1997;accepted14 May 1997

Abstract The kinetics of bromide oxidation by several peroxo complexes of molybdenum(Vl) and tungsten(VI) were studied. All systems obey the rate law d [ Br + ] Idt = k[metal complex] [ Br- ]. At pH 1.0, we quantify the initial rate of bromine production by the spectrophotornetrie monitoring of the disappearance of fluorescein. The complexes and their rate constants at pH 1.0 and 25°(2 are MoO(O2)(nta)(nta = nitrilotriaeetate) (k = ( t.64 ___C.~2) × 10-4 M-s s-S). MoO(O,) (dipic) (H20) (dipic---pyridine-2,6-dicarboxylate) (k= (5.1 -I-0.2) × 10-4 M-I s-I), WO(O2)(dipie)(H20) (k= (2+79+0.02) × 10-2 M-I s-I), MoO(O2)2(pie)- (pic=pyridine-2-carboxylate) (k-- ( 1.633+0.001 ) × 10- i M- i s- ~) and MoO(O2)~(H20)2 (k=6.t2+0.03) × 10-, M- as- s). At pH 5.0, theinitial rates were measured spectrophotometrically by the production of bromophenol blue from phenol red. The complexes and their rate constants at pH 5.0 and 25°C are MoO(O2)2(OH)(H20)- ( k = ( 8 + l ) × 1 0 -4 M -s s - ' ) , WO(O2).~(OH)(H20)- (k=(8.3+0.7)×10 -~ M -1 s -t) and WO(O2)2(C20+) 2- (k= (1.3+0.2) × 10-2 M- ' s-'). The ~70 chemical shifts for the peroxo groups of MoO(O~)(nta) - (501 ppm), MoO(O2)(dipic)(H20) (537 ppm), WO(O2)(dipic)(H20) (364 ppm), MoO(O2)_,(pic)- (434 ppm) and WO(O2)z(CzOa) 2- (368 ppm) were measured after enrichment with H2'~O2.As expected, these values correlate with the reciprocal of the transition energies of the electronic spectra. The values i, (Mo=O) = 930 cm- s and v(O--O) = 900 era- s for MOO(O2)(nta) - are consistem with the vibrational features of similar compounds. The relationship between the spectroscopic propeaies and reactivity toward bromide oxidation follows previously observed patterns and may allow the reactivity of peroxometal complexes to be predicted from spectroscopic information. © 1997 Elsevier Science S.A. Keyword~: Kineticsand mechanism;Bromideoxidation;Molybdenumcomplexes;Tungstencomplexes;Peroxocomplexes

1. Introduction When peroxide is coordinated to early transition metal ions having high oxidation states, the reactivity of the peroxide as an oxidant is enhanced. Such an increase in rate explains the role played by peroxometal compounds in a number of biological oxidations, including those catalyzed by catalases, peroxidases and haloperoxidases [ 1,2]. Numerous reactions of commercial significance are catalyzed by synthetic peroxo complexes, including the oxidation of arches, alcohols, olefins, phosphines, sulfides and halides [3-7]. An oxidative mechanism is also thought to be involved in the success of peroxovanadium(V) complexes as mimics of insulin [8,91. The kinetics of oxidation reactions involving peroxo complexes of several early transition metals in their highest oxi* Correspondingauthor.Tel.: 315 824 7235;fax: 315 824 7935. 0020-16931971517.00© 1997ElsevierScienceS.A. All rightsreserved Pil S0020- i693 ( 97 ) 05 65 8-2

dation states have been reported. Thompson and co-workers looked at the oxidation of a cobalt-bound thiolato ligand by several peroxo complexes of vanadium(V), molybdenum(VI) and tungsten(VI) [10-13]. Espenson and coworkers examined the rate at which peroxo complexes of rhenium(VII) oxidized organic sulfides, anilines and halides [ 14-17 ]. From the laboratories of Butler [ 18,19 ] and Reynolds [ 20] came invesfgations of the oxidation of bromide by peroxo complexes of vanadium(V), molybdenum(VI) and tungsten (VI). These oxidations are believed to occur by the transfer of an oxygen atom following the attack of substrate on the peroxo ligand [7,10,11,19]. We recently reported for a series of peroxometal complexes a correlation among spectroscopic properties, namely 170 peroxo chemical shift, electronic charge transfer band energy and O-O vibrational frequency [21 ]. A further correlation was noted between these properties and the reactivity

226

M.S. Reynolds et al. / inorganica ChiraicaActa 263 (1997) 225-230

of the complexes toward oxidation: the more reactive complexes are those with more shielded peroxo group oxygen atoms, higher-energy charge transfer bands, and lower O-O stretching frequencies. Such a correlation suggests that the reactivity of peroxo complexes depends on the strength of the O--O bond, and therefore on the ease with which it is broken. Since the ability to predict reactivity on the basis of relatively simple spectroscopic measurements is useful in the design of new catalysts, it is of interest to verify and probe these trends through reactivity studies on an extended set of peroxo complexes. We seek. therefore, to determine how such factors as the nature of the central metal, the number ofperoxo iigands, and the nature of the chelating ligand affect the reactivity of peroxometal complexes toward the oxidation of bromide ion, and how these factors manifest themselves in spectroscopic parameters.

2. Experimental 2.1. Materials and solutions

Potassium (nitrilotriacetato) oxoperoxomolybdate(VI) hydrate, K[MoO(O~)(nta)l -H20 [ 13], hydronium oxodiperoxo (pyridine-2-carboxylato) molybdate(VI) hydrate, H[MoO(O2)2(pic)]-H20 [22] and aquaoxoperoxo(pyridine-2,6-dicarboxylato) tungsten(Vl), We(de) (dipic)(H20) [22] were prepared by literature methods. The complexes MoO(d2) (dipic) (H20) [ 13], MoO(O2)z(H20)2 [10] MoO(O2)2(OH)(H20)- [10], WO(O2)2(OH)(H20) - [ 10] and WO(O2) 2(C204) 2- were prepared in situ from solutions of potassium molybdate or tungstatc, hydrogcn peroxide and the appropriate ligand. Lithium bromide ( > 97%) was obtained from Fhka. Lithium perchlorate was prepared by the neutralization of lithium carbonate ( > 99%, Fhka) with perchlorie acid and was recrystallized three times. Potassium molybdate (Aldrich), potassium tungstate (Aldrich), pyridine-2-carboxylic acid (Aldrich), pyridine2,6-carboxylic acid (Aldrich), nitrilotriacetic acid (Aldrich), phenol red (4,4'-(3H-2,I-benzoxathiol-3-ylidene)bis(phenol) S,S-dioxide) (Fisher), and fluorescein sodium salt (Y,6'-dihydroxyspiro(isobenzofuranI (3H),9'- [9//] xanthen]-3-one disodium salt) (Aldrich) were used as received. Distilled water was purified with a Bamstead Nanopure deionizer. Dioxygen gas (28 at.% 170) was obtained from Isotec. Oxygen-17-labeled hydrogen peroxide was synthesized by a modification of Sitter and Terner [231. The total hydrogen peroxide concentration of the resulting aqueous solution was typically 0.45 M. Solutions of hydrogen peroxide were standardized by spectrophotometric measurement of chloroperoxidase-eatalyzed triiodide formation, a modification of the method of Cotton and Dunford [241. Stock solutions of bromine were prepared by dissolution of Brz in 0.t0 M NaOH and were standardized spectrephotometrically by the oxidation of iodide to triiodide

(Am~=353 nm, e=26400 M -1 cm -j for Is-) [251. A calibration curve for the dependence of fluoreseein absorbance on bromine concentration was prepared by combining bromine with fhorescein and bromide in 0.10 M HCIO4 (pH 1.0). The ionic strength was adjusted to 1.0 M with lithium perchlorate. The absorbance of each solution was measured at 438 nm, the Am.,,for fluoreseein. The value ofdA/d[Br + ] for fluorescein in the presence of bromine was obtained from the slope of a plot erA versus [Br + ] t. Essentially identical results were obtained for bromide concentrations of 25 and 500 mM. 2.2. Kinetics measurements

Reactions of peroxometal complexes with bromide were monitored by the production of bromine as manifested by the disappearance of fluorescein (pH 1.0) or by the production of bromophenol blue from phenol red (pH 5.0). Spectral changes were followed at 25.0 + 0.5°C with a HewletbPackard 8452A diode array spectrophotometer. Conventional mixing techniques were employed for the slower reactions; a stopped-flow mixing device (Hi-Teeh SFA-12 Rapid Kinetics Accessory) was used in conjunction with the diode array spectrophotometer for the faster reactions. In a typical experiment, a solution of the metal complex, excess ligand, hydrogen peroxide (where appropriate) and fluorescein or phenol red was mixed with an equal volume of a solution of lithium bromide, Both solutions contained 0.10 M HCIO+ (pH 1.0) or acetate or oxalate buffer (pH 5.0), and sufficient lithium perchlorate to maintain the ionic strength at 1.0 M. Initial concentrations after mixing are listed in Table 1. Reactions were monitored speetrophotometrically over the wavelength range 350-550 nm (pH 1.0, fluorescein) or 350-750 nm (pH 5.0, phenol red). Data analysis was performed at 438 nm, the Am~ for fluorescein, or at 588 nm, the Am~ for bromophenol blue. Initial rates dA/dt for each reaction were measured from the slope of the steady-state portion of the absorbanee versus time trace. Two or three replicate determinations were averaged for each concentration. Absolute rates d[Br + ]/dt were obtained by the application of the reciprocal of the factor dA/d[Br + ] to the experimental rates dAldt. Control experiments revealed no reaction between metal complex and either fluoreseein or phenol red in the absence of bromide. Least-squares fits were performed with the program Origin [26]. 2.3. ;70 N M R spectroscopy

Oxygen-17-1abeled peroxo complexes MoO(d2)(nta)-, MoO(d2) (dipic) (H20), We(d2) (dipic) (H20) and MoO(O2)2(pic)- were prepared from the unlabeled solid compounds; enrichment was effeeted by exchange with H2tTO,.. The pH 1.0 (0.10 M HCIO4) solutions contained t Theprimarybromineproductfromtheoxidationofbroraideisunknown becausean equilibriummixtureof HOBr,OBr-, Br~and Br3- is rapidly obtained.Wethereforedesignatebrominein its oxidizedformas 'Br+'.

M.S. Reynoldset al. / inorganica ChimicaActa 263 (1997)225-230

227

Table 1 Initialconcentrationsand rateconstantsforbromideoxidationby peroxomolybdenum(VI)and peroxotungsten(Vl)complexes Complex MoO(O~)(nta)-" Moo(e2) {dipic)(HzO) c we(o:) (dipic)(HzO) i MoO(O2)z(pic)= " MoO(O=)z(HzOh" MoO(O2)2(OH)(H20)- ~ WO(O2)z(OH)(HzO) - ¢ WO(Oz)z(C~O,)2- ¢

pH 1.0t' 1.0b !.0 b 1.0~ 1.0b 5.0 d 5.0 d 5.0 e

[Metal]o (raM) 0.12-1.1 0.1-0.4 0~05-0.5 0.1-t.0 0.1-1.0 0.5-1.0 0.01-0.1 0.015-0.15

[H202]o (M)

[Ligand]o (raM)

(I-4)× I0-4 0.001-0.010 0.001-0.010 0.I00 0.040 0.100

!.2-11 4 0.5 0. I-I.0

|0

[Br-]o

k

(M)

(M-I S-')

0.045-0.450 0.045-0.450 0.045-.0.450 0.045--0.450 0.015-0.!50 0.050--0.¢51 0.050-0.500 0.050-0.500

(1.64d:0.02) × 10 - 4

(5.1:1:0.2)× 10-4 (2.79+0.02)×10 -2 ( 1.633+0.001) X 10-t (6.12:1:0.03)X 10-I (8+I)XI0 -4 (8.3+0.7)×!0 -~ (I.3:i:0.2)× 10-z

"Preparedfromsolidcompound. b0. I0 M HCIO4. c Preparedin site. d0.010 M acetatebuffer. e 0.010M oxalatebuffer. 0.020 M metal complex, 0.020 M additional heteroligand, and 0.020 M (MoO(O2)(nta)-, MoO(O2)(dipic)(H20), WO(O2)(dipie)(H20)) or 0.200 M (MoO(O2)2(pic)-) H21~O2. A solution of I~O-enriched we(e2)2(C204) 2- was prepared in sire from 0.010 M potassium tungstate, 0.100 M H21:O2 and 0.070 M oxalate barfer at pH 5.0. Oxygen-17 NMR spectra were measured in 10 mm tubes on a Braker AC-250 (5.9 I', 33.94 MHz) spectrometer against distilled water as the external reference (0 ppm). Spectra were acquired without a field lock. The spectral width was 50 000 Hz, and the number of accumulations ranged from 1750 to ! 22 000. Because spin-spin relaxation of'70 in peroxo complexes is extremely rapid, a short acquisition delay ( 10 p,s) was employed. The severely rolling baseline resulting from pulse breakthrough was diminished by subtraction (left shift) of the first two to four points from the FID. A line-broadening factor of 50 Hz was applied in the exponential multiplication. 2.4. UV-Vis and IR spectroscopy

Electronic spectra of the peroxo complexes were measured on a Hewlett-Packard 8452a diode array spectrophotometer. IR spectra were measured on a Nicolet 740 FTIR spectrometer as potassium bromide pellets or Nujol mulls.

3. Results and discussion

The rates of bromide oxidation for MoO(O2)2(OH)(H20) -, WO(O2)2(OH) (H20) - and w e ( e 2 ) 2(C204) zwere followed at pH 5.0 by spectrophotometric monitoring of the bromination of phenol red (A,,~--432 nm) to bromophenol blue (Am~=588 nm); this method has been described previously [20]. Initial rates were taken from the linear (steady-state) portion of the absorbance versus time curve at Aam=588 nm. The complexes MoO(Oz)(nta)-, MoO(e2) (dipic) (H20), WO(Oz) (dipic) (H=,O), MoO(O2)2(pie) - and MoO(O2)2(HzO)2 required more acidic condi-

tions for optimal stability. At pH 1.0, the yellow-to-blue color change is not observed for the bromination of phenol red because both phenol red and bromophenoi blue are yellow at this pH. A different bromine indicator was therefore developed for the kinetics studies at low pH. F l ~ e i n undergoes bromination to the eosin Y tetrabromination product, as shown in Eq. (1). Bf

.4e,"-----

FLLJOFIF.SCEIN

a¢

..e

(1)

ECISIN Y

This reaction was suggested as an assay for haloperoxidases [27] and was reported as an assay for hypocb.lomus acid and hypochlorite ion in seawater [28]. The absorption maximum for fluorescein (438 tun at pH 1.0) disappears upon bromination to yield the spectral f e a ~ of eosin Y (~=528 nm at pH !.0). Because the interntediate bmmination products also absorb near 528 nm, the rate of binruination was obtained by following the disappearance of fluorescein instead of the appearance of eosin Y. The initial rate was taken as the slope of the linear (steady-state) portion of the absorbance venus time curve for each reaction. For all complexes examined, the variation of dAIdt with initial concentration was linear for both the peroxo con-q~ex and bromide, indicating a first-order dependence of the initial rate on these reactants. The second-order rate law for the formation of bromine is given in Eq. (2). d[Br + ]/dr= k[complex] [Br- ]

(2)

The absolute reaction rate d[Br + ]/(it was obtained from the experimental quantity dAsss/dt (bromophenol blue) or dA43s/dt (fluorescein) by use of Eq. (3). d[ Br + ] ~dr= (dA/dt) (d[ Br + ]/dA)

(3)

228

M.S. Reynoldset al. / inorganica ChimicaActa 263 (1997) 225-230

The value dA/d[Br ÷ ] = 15 800 M-~ for bromophenol blue was obtained previously [20]. The value dA/d[Br + ] = - 3206 M- ~for fluorescein was obtained at pH !.0 and ionic strength !.0 M at 0.025 and 0.500 M bromide. This factor and Eqs. (2) and (3) permit the calculation of the secondorder rate constant k. Table 1 lists the values ofk for each of the peroxo complexes examined. For MoO(O2)2(OH)(H20)- and WO(O2)2(OH)(HzO)-, the dependence of the rate on the initial concentration of metal complex was examined at a constant initial bromide concentration, and the dependence of the rate on the initial bromide concentration was examined at a constant initial concentration of metal complex. Initial concentrations, initial rates dA/dt, and k for each of the experiments inv oiv ing MoO (O2) 2(OH) (H20) and WO(O2)2(OH) (H20) - are available in Tables St and $2, respectively (see Section 4). Table 1 reports the average value of k and its standard deviation for both complexes. For the remaining compounds, the concentrations of metal complex and bromide were varied simultaneously. Initial concentrations and initial rates - d A I d t for MoO(Oz)(nta)-, MoO(O2) (dipic) (H~O), WO(O2) (dipic) (H20), MoO(O2)2(pie)-, MoO(O2)2(H20)2 and W O ( O 2 ) 2 ( C 2 0 4 ) 2- are available in Tables $3--$8, respectively (see Section 4). For each complex, a two-variable, least-squares fit to the surface described by Eq. (4) was performed, and the rate constant k was obtained from k' by dividing the latter value by dA/ d[Br÷l. d A / d t = k ' [ m e t a l l [Br- l

(4)

The rate constants k and their standard errors are summarized in Table 1. Table 2 summarizes the rate constants for bromide oxidation reactions examined in our laboratory. These results confirm earlier observations that tungsten(VI) peroxo complexes ate more reactive than their molybdenum(Vl) counterparts [ 10,12,20], a trend that holds here for monoperoxo and diperoxo complexes; for oxalate, pyridine-2,6dicarboxylat~ and H20/OH- ligands; and for pH 1.0 and 5.0. The del:¢ndence of the rate on pH is clearly seen in the case of MoO(Oz)z(OH) (H20)- and MoO(O2)~,(H20)2, where a gr,~ater than 700-fold increase in rate occurs on

decreasing the pH from 5.0 to 1.0. That the oxalate-containing complexes react more quickly than do those with water and hydroxide ligands at pH 5.0 has been interpreted as indicative of attack of the bromide ion at the peroxo ligand; a reverse trend in reactivity might be expected if the mechanism w~re to involve direct binding of bromide to the metal center, inasmuch as the metal ion is more accessible in the case of MoO(O2)2(OH)(H20 ) - and WO(O2)2(OH)(H20)[ 11,20]. The faster rate of the oxalate-containing complexes suggests that the nature ofthe chelating ligand is more important than the charge on the metal complex in determining reactivity with bromide, since the dianionic oxalato complexes react more quickly than the monoanionie aquahydroxo complexes. Such an argument is not easily made, however, at pH 1.0 for the monoperoxo complexes MoO(O2)(nta)and MoO(O2)(dipic)(H20) or the diperoxo complexes MoO(O2)2(pic)- and MoO(O2)2(H20)2, where the neutrally charged complexes react about three times faster with bromide than do the monoanions. In terms of the number of peroxo groups bound to the metal, the diperoxo MoO(O2)2(pic)- and MoO(O2)2(H20)2 are significantly more reactive than the monoperoxo MoO(O2)(nta)- and MoO (02) (dipic) (H20), consistent with earlier reports for other complexes [11,12], but the effect of the chelating ligand must be understood more fully before a generalization can be made about the relative reactivity of mono- and diperoxo complexes. The '70 NMR, electronic, and vibrational properties of the complexes under consideration are compiled in Table 2. We seek to determine the generality of the correlations that exist among the spectroscopic properties of peroxo complexes, correlations that may extend to the reactivity of the compounds as oxidizing agents [21]. The 170 NMR chemical shifts measured here for the peroxo groups of MoO(O.~)~, (nta) -, MoO(O~) (dipic) (H20), WO(O2) (dipic) (H20), MoO(Oz)2(pic)- and WO(O2)2(C204) 2- correlate fairly well with the ligand-to-metal charge transfer (LMCT) wavelengths of the electronic spectra, indicating that the shielding in these complexes is consistent with a formalism developed by Ramsey, as was previously observed for a larger set of complexes [21]. The correlation between O-O stretching

Table2 Spectroscopicpropertiesand bromideoxidationrateconstantsfor peroxomolybdenum(VI)and pero~otungsten(Vl)complexesa Complex

pH

MoO(Oz)(nta) MoO(O~)(dipic)(H20) WO(O2)(dipic)(H20) MoO(O2)z( pic ) -

1.0 1.0 1.0

MoO(O2) ,( ]'120)~,

|.0 5.0 5.0 5.0 5.0

MOO(OD,(OH)(H20) MoO(O2)2(C204)"~WO(Oz):(OH)(H20) WO(O2)2(C204) 2-

1.0

3(0-0) (ppm)

'~ma~ (nm)

S01 537 364 434 435 412 433 346. 361 368

364 364 272 328 32b 310 320 234 202

u(O-O)

k

(on-')

(M-t s-

900 900 875 850

(1,64 + 0.02) × 10- 4 (5.1 + 0.2) × 10- 4 (2.79 ± 0.02) × 10- a

872 871. 855

"Data fromthis workaxein boldfacetype;remainingdataaxefromthe referencesindicated.

Ref. ') [ 13] [ 13,22] [22]

( 1.633 + 0.001 ) × 10 - '

[22 ]

(6.12+0.03) × 10-I (85:1)×10 -4 (4.8 ± 0.4) × I0 - 3 (8.3+0.7) × 10-a (!.3±0,2) × 10-z

[21 ] [21] [20,21] [21 l [21 ]

M.$. Reynolds et ai. I lnorganica Chimica Acta 263 (1997) 225-230

229

Wavelength (rim) 0.015

250 0.800

o 0.600

O.OLO.

0.400 o.~, 0.200-

0.000 1.25

o.ooo 1.oo

t.~o

1.~s

i.0o

1,',,,Iz (lO~1 k..r~)

Fig. I. Plot of bromide oxidation rate constant vs, the reciprocal oftboLMCT energy (lowerx axis) and wavelength (upperxaxis) for the datain Table 2 at pH 1.0. Symbol shapes correspond to the following central metal ions: molybdenum ( O ) , tungsten (&).

frequency and LMCT wavelength for the set of compounds presented in Table 2 is not well defined. The vibrational spectrum of K [ MoO( 02 ) (nta) ] exhibited strong bands at 930 and 900 cm- m.Assignment of these features to the stretches of the Mo=O and O-O bonds, respectively, was made by analogy with the spectra of MoO(O~)(dipic)(H20) (u(Mo=O)=970, u(O-O)=900 cm -I) [221 and MoO(e2) (CN)42- (v(Mo=O) = %0, u(O-O) = 910 cm- i) [291. In our earlier report of arelationship between spectroscopic properties and the oxidizing reactivity of the peroxo group, it was observed that the most reactive l~roxo complexes were those with A,,~ values below 400 nm, 0-43 stretching frequencies below 900 cm-i, and t~O chemical shifts below 600 nm [21], criteria which are fulfilled for all of the complexes studied here. Indeed, the complexes examined in our laboratory that have lower energy LMCT bands, higher energy O-O stretching frequenciesor more deshielded chemical shifts than those described above showed no significant reactivity toward bromide, namely VOz(O2)(C204)22[21 ], VO(Oz) (dipic) (H20) - [21 ] andVO(O2) (nta) 2- 2. A plot of the bromide oxidation rate constant as a function of the reciprocal of the LMCT energy (k versus I/AE) for the reactive complexes studied at pH 1.0 is displayed in Fig. 1. The reciprocal of the LMCT energy was employed because it was shown earlier to correlate reasonably well with both '~O chemical shift and O-O stretching frequency for complexes of this type [21]. With the exception of We(C2) (dipic) (H20), there is a linear trend toward greater reactivity as the energy of the LMCT transition increases. A similar but less dramatic change in reactivity is observed for the compounds studied at pH 5.0, as shown in the plot of k versus I/AE in Fig. 2. That strict linearity is not observed in these plots is perhaps not surprising since the charge transfer 2 The complex (NH~)2IVO(O2) (nta)l was prepared according to Ref. [30], and exhibited ,~(O-O) =603 ppm (pH 1.0), u(O-O) =925 cm -I. A mixture of 0.010 M (Nl'~)z [VO(O:) (nla) ] and 0.450 M LiBr at pH 5.0 showed no formation of bromophenol blue after 30 rain.

l~s

1~o

I~

zoo

11~ (lo"mx~-I) Fig. 2. Plot ofl~omide oxidation m e constant vs. the r~iprocal o f ~ LMCT energy (lowetx axis) and wavelength (upper x axis) for file daminTablc 2 at pH 5.0. Symbol shapes correspond to the following central metal ions: molybdenum (IlL tungsten ( • ).

process is a one-electron event, whereas oxygen atom transfer involves two electrons. The observation that reaction rate and LMCT wavelength correlate rather generally, however, suggests that spectroscopic properties might be used to predict the relative reactivity of peroxo complexes, a capability that can be very useful in the design of new catalysts, haloperoxidase analogues and insulin mimics. Efforts to determine the influence of the chelating ligand on complex reactivity by a systematic variation of tl~ electronic properties of the ligand are in progress, as are further studies of the effect of pH on these reactions.

4. Supplementary material Kinetics data (Tables SI-$8) arc available from the authors on request.

Acknowledgements

Support of this work by a Joseph H. DeFrees Grant from the Research Corporation; by the donors of the Petrok:um Research Fund, administered by tbe American Chemical Society; and by NSF-REU Grant CHE-9300589 is gratefully acknowledged.

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