10926 Chem. Commun., 2013, 49, 10926--10928 This journal is c The Royal Society of Chemistry 2013

Cite this: Chem. Commun.,2013,49, 10926

Mechanistic insight into halide oxidation by non-hemeiron complexes. Haloperoxidase versus halogenaseactivity†

Anil Kumar Vardhaman,a Prasenjit Barman,a Suresh Kumar,b Chivukula V. Sastri,*a

Devesh Kumar*b and Sam P. de Visser*c

This work presents the first detailed study on mechanistic aspects

of halide oxidation by non-heme iron complexes. We show that

while iron(III)–hydroperoxo complexes oxidise halides via oxygen

atom transfer, the corresponding iron(IV)–oxo complex reacts via

electron transfer.

A large variety of natural product biosynthesis reactions involvethe halogenation of aromatic and aliphatic carbon centres.1

Halogenating enzymes are divided into H2O2 required halo-peroxidases (heme systems) and oxygen dependent halogenases(non-heme systems).2,3 Over the years many heme containinghaloperoxidases have been discovered that utilize chloride,bromide and iodide ions and in a reaction with hydrogen peroxidelead to hypohalide products.4 The a-ketoglutarate dependenthalogenases, by contrast, form a high-valent non-heme iron(IV)–oxo (halide) intermediate that abstracts a hydrogen atom from analiphatic group and rebound the halide ligand to form ahalogenated substrate in a stepwise radical process.5 Whilehaloperoxidases catalyse electrophilic halogenation (X+) reac-tions, the non-heme iron halogenases, on the other hand,usually react via a radical process (X�).

Despite the fact that the enzymatic mechanism is well explored,surprisingly, few studies have been reported on analogous bio-mimetic models. As such, we report the first detailed mechanisticstudy on the reaction of halides with biomimetic iron(IV)–oxo com-plexes. The catalytic halogenation of organic substrates by thesemodels was reported for a number of complexes that are likelyto form a high-valent metal–oxo intermediate.6 In addition,

the reactivity of non-heme iron(III)–hydroperoxo intermediatesas effective electrophilic oxidants in substrate halogenation hasbeen reported, and it was found that under certain conditions theyare catalytically more active than iron(IV)–oxo intermediates.7 Furtherstudies on the relative activity of non-heme iron(IV)–oxo versusiron(III)–hydroperoxo in aromatic hydroxylation and sulfoxidationof substrates have been reported.8 To gain insight into the mecha-nistic pathways by non-heme iron(III)–hydroperoxo intermediates inelectrophilic reactions, we decided to explore the direct reactivity ofhalides (X� = F�, Cl�, Br� and I�) with [FeIII(OOH)(N4Py)]2+ (1a) and[FeIII(OOH)(Bn-tpen)]2+ (2a) and compare the kinetics with those ofthe corresponding iron(IV)–oxo complexes, i.e. [FeIV(O)(N4Py)]2+ (1b)and [FeIV(O)(Bn-tpen)]2+ (2b) (Fig. S1, ESI†).9

The iron(III)–hydroperoxo complexes, 1a and 2a, were generatedin situ in a methanol solution at 253 K by addition of 10 equiv. ofH2O2 following a reported procedure.10 Subsequently, tetrabutylammonium halides (TBA–X; X = Cl�, Br�, I�) were added to 1a atthe same temperature, which led to decay of the characteristic548 nm peak in the UV-vis spectra. The second order rateconstants extracted from these experiments (Fig. 1 and Table 1)showed a difference of as large as a factor of 6, and the reactivitydecreased in the order I� > Br� > Cl�. Surprisingly, the observedtrend does not correlate with the electron affinities of thesehalides. Moreover, when TBA–F was added to 1a no oxidationreaction of F� was observed, which may be due to the fact thatH2O2 lacks the thermodynamic potential to oxidize fluoride ions.11

In order to compare the reactivities of iron(III)–hydroperoxowith iron(IV)–oxo in halide oxidation, we generated 1b and 2b inCH3CN using a procedure reported previously.12 Addition ofhalides to 1b under the same conditions as those for 1a enabledus to evaluate the second order rate constants: 1b reacted withhalides (Cl�, Br�) with much lower rate constants than 1a;however, a reversal of the trend was observed in the presence of I�.In fact, for iodide oxidation, 1b was found to react about sixtimes faster than 1a. This reversal in trend may be attributed tothe steric factor influenced by the bulky I� group uponapproach of 1a. Additionally, the reactivity of 1b was found tobe sensitive to the one electron oxidation potential of halidesthus demonstrating that the reactivity of non-heme iron(IV)–oxo

a Department of Chemistry, Indian Institute of Technology Guwahati,

Assam 781039, India. E-mail: sastricv@iitg.ernet.in; Fax: +91-361-258-2349b Department of Applied Physics, School for Physical Sciences,

Babasaheb Bhimrao Ambedkar University, Vidya Vihar, Rai Bareilly Road,

Lucknow 226 025, India. E-mail: dkclcre@yahoo.comc Manchester Institute of Biotechnology and School of Chemical Engineering and

Analytical Science, The University of Manchester, 131 Princess Street,

Manchester M1 7DN, UK. E-mail: sam.devisser@manchester.ac.uk

† Electronic supplementary information (ESI) available: Experimental and com-putational methods, and experimental and computational raw data in the form oftables and figures. See DOI: 10.1039/c3cc46792a

Received 5th September 2013,Accepted 8th October 2013

DOI: 10.1039/c3cc46792a

www.rsc.org/chemcomm

ChemComm

COMMUNICATION

Publ

ishe

d on

08

Oct

ober

201

3. D

ownl

oade

d by

FO

RD

HA

M U

NIV

ER

SIT

Y o

n 31

/10/

2013

10:

55:3

0.

View Article OnlineView Journal | View Issue

This journal is c The Royal Society of Chemistry 2013 Chem. Commun., 2013, 49, 10926--10928 10927

complexes can be significantly affected by the identity ofsubstrates. The difference in the reactivity between 1a and 1bwas in the order of B103 in the oxidation reaction of Br� andincreased to B105 for Cl�. A similar trend was observed whenthe halide oxidation reactions of 2a and 2b were compared.

To quantify the difference in the oxidation pathways foriron(III)–hydroperoxo and iron(IV)–oxo intermediates, we thenmade a plot of the second order rate constants log(k2) againstthe one electron oxidation potential (E0

ox) of these halides. A linearcorrelation was observed with a slope of –0.88 and �6.62 for 1aand 1b respectively (Fig. 2). Similar slopes were obtained for 2a(�1.23) and 2b (�5.16) (ESI†). Such large negative values of slopesfrom a plot of log(k2) versus one electron oxidation potentialimplicates a rate determining electron transfer step in the reactionfrom halides to oxidants, i.e. 1b and 2b, rather than a grouptransfer.13 The ease of a plausible electron transfer from I� to

iron(IV)–oxo can now account for the higher oxidation rates observed.The difference in the reaction mechanism for the oxidation ofhalides by iron(IV)–oxo as compared to iron(III)–hydroperoxo com-plexes can be attributed to the interactions of the negativelycharged substrates with the electrophilic iron centre.

Further evidence of the mechanistic insights was obtainedfrom the catalytic oxidation of anisole by 1a and 1b in thepresence of TBA-Cl (see ESI† for details). These two chlorinatingspecies yield different products with anisole. The oxidation ofchloride involves either a one-electron oxidation (Cl�- Cl� + e�)to form a chloride radical or an overall two electron oxidation(Cl�- Cl+ + 2e�) to give the chloronium ion. Chlorination by 1agave only monochlorinated products ( para and ortho isomers).In stark contrast, the chlorination of anisole by 1b gave a mixture ofmono- and di-substituted chlorinated products, Scheme 1. Theisomer distribution and product ratios indicate that the chlorinatingagent in the presence of 1a is Cl+. The formation of di-substitutedanisole and phenoxymethyl chloride is characteristic of free radicalchlorination of anisole in the presence of 1b.14

To gain further insight into the difference in reactivity of 1aversus 1b in a reaction with halides we performed a set of DFTstudies on their reactivity with Br�, as explained in detail inthe ESI.†15 Bromide oxidation by either 1a or 1b proceeds in a

Fig. 1 Comparison of the second order rate constants for halide oxidation by 1a( ) and 1b ( ) in the presence of (a) TBA-Cl, (b) TBA-Br and (c) TBA-I at 253 K.

Table 1 Second order rate constants for the oxidation of halides

X�k2 (M�1 s�1) (253 K)

1a 1b 2a 2b

Cl� 1.46 � 10�1 6.33 � 10�7 4.00 � 10�2 6.00 � 10�5

Br� 5.50 � 10�1 6.41 � 10�4 2.70 � 10�1 1.00 � 10�2

I� 1.20 7.03 8.01 � 10�1 19.01

Fig. 2 Plot of log(k2) against the one-electron oxidation potentials (E0ox) of

halides of 1a ( , blue) and 1b ( , red) at 253 K.

Scheme 1 Products obtained from the reaction of anisole with 1a and 1b in thepresence of TBA-Cl at 298 K. The values in parentheses represent relative yieldsin mole%.

Communication ChemComm

Publ

ishe

d on

08

Oct

ober

201

3. D

ownl

oade

d by

FO

RD

HA

M U

NIV

ER

SIT

Y o

n 31

/10/

2013

10:

55:3

0.

View Article Online

10928 Chem. Commun., 2013, 49, 10926--10928 This journal is c The Royal Society of Chemistry 2013

concerted manner via a single transition state (TSBr) to form aproduct complex of hypobromide bound to iron(II)(N4Py). Fig. 3displays the rate determining transition states for the reactionsof 1a/1b with Br�. Geometrically, the reaction of 1a with Br�

gives early transition states with longer O–Br distances ascompared to the reaction of 1b with Br�. The large differencesbetween the transition states, however, relate to the spindensities and charges. Thus, the iron(IV)–oxo complex reactsvia an initial one-electron transfer to the substrate to produce atransition state with a large degree of radical character on theapproaching halide: rBr = �0.49 (�0.51) for 5TSBr (3TSBr),similar to what was observed before for electrophilic reactionmechanisms by non-heme iron(IV)–oxo complexes.16 In addi-tion, we calculated a large amount of charge transfer (QCT) fromthe oxidant to the substrate as obtained from the difference incharge of the halide in the reactant complex with respect to TSBr.This way, we obtained values of QCT to be 0.61 (0.64) for 5TSBr

(3TSBr), respectively. On the other hand, the bromide oxidationtransition state of 1a, i.e. 6TSBr,1a, shows little changes in groupspin densities with respect to the reactant complex and only asmall degree of charge transfer (QCT = 0.13),7 hence the iron(III)–hydroperoxo species reacts via group transfer rather than one-electron transfer.

These computational studies show that the electron transferprocesses in electrophilic reaction mechanisms are dependent onthe nature of the oxidant, whereby iron(IV)–oxo abstracts electronsearly, whereas the electron transfer by iron(III)–hydroperoxowill happen after the transition state, on the way to productformation. The electron affinity difference of 1a versus 1b will,therefore, determine the electrophilic reaction rates with halides.We calculated electron affinities for 1a and 1b of DE + ZPE + Esolv =102.6 and 108.7 kcal mol�1, respectively. Thus, the larger electronaffinity of the iron(IV)–oxo as compared to the iron(III)–hydroperoxospecies will enable a much quicker and efficient electron transferby the iron(IV)–oxo complexes. Indeed, the transition statesshow much earlier electron transfer from the iron(IV)–oxo ascompared to the iron(III)–hydroperoxo species. These computa-tional results are also in perfect agreement with the plot inFig. 2 that shows a correlation of the reaction rate with theelectron oxidation potential.

In summary, we reported a comparative study of the reac-tivity of non-heme iron(III)–hydroperoxo versus iron(IV)–oxo inhalide oxidation reactions. We identified dramatic differences

in the electron transfer processes for the two reactions, wherebyiron(IV)–oxo reacts via a one-electron transfer and iron(III)–hydroperoxo by group transfer.

The authors acknowledge support from the Department ofScience and Technology, India, for the Ramanujan Fellowship(SR/S2/RJN-11/2008) to DK and grants to CVS: SR/S1/IC-02/2009and support from the Council for Scientific & IndustrialResearch (01(2527)/11/EMR-II). CPU time was provided by theNational Service of Computational Chemistry Software (NSCCS)to SdV.

Notes and references1 (a) G. W. Gribble, J. Chem. Educ., 2004, 81, 1441; (b) A. Butler and

M. Sandy, Nature, 2009, 460, 848.2 (a) F. H. Vaillancourt, E. Yeh, D. A. Vosburg, S. Garneau-Tsodikova

and C. T. Walsh, Chem. Rev., 2006, 106, 3364; (b) K. H. Van Pee,C. Dong, S. Flecks, J. Naismith, E. P. Patallo and T. Wage, Adv. Appl.Microbiol., 2006, 59, 127; (c) J. L. Anderson and S. K. Chapman,Mol. Biosyst., 2006, 2, 350; (d) D. G. Fujimori and C. T. Walsh,Curr. Opin. Chem. Biol., 2007, 11, 553.

3 L. P. Hager, D. R. Morris, F. S. Brown and H. Eberwein, J. Biol.Chem., 1966, 241, 1769.

4 (a) J. A. Manthey and L. P. Hager, Biochemistry, 1989, 28, 3052;(b) J. A. Manthey and L. P. Hager, J. Biol. Chem., 1985, 260, 9654;(c) M. P. Roach, Y. P. Chen, S. A. Woodin, D. E. Lincoln, C. R. Lovelland J. H. Dawson, Biochemistry, 1997, 36, 2197.

5 (a) C. Krebs, D. Galonic Fujimori, C. T. Walsh and J. M. Bollinger Jr,Acc. Chem. Res., 2007, 40, 484; (b) Iron-containing enzymes: Versatilecatalysts of hydroxylation reactions in nature, ed. S. P. de Visser andD. Kumar, Royal Society of Chemistry Publishing, Cambridge, UK,2011.

6 (a) T. Kojima, R. A. Leising, S. Yan and L. Que Jr., J. Am. Chem. Soc.,1993, 115, 11328; (b) A. Bakac, C. Shi and O. Pestovsky, Inorg. Chem.,2004, 43, 5416; (c) J. Y. Kravitz and V. L. Pecoraro, Pure Appl. Chem.,2005, 77, 1595; (d) P. Comba and S. Wunderlich, Chem.–Eur. J., 2010,16, 7293; (e) W. Liu and J. T. Groves, J. Am. Chem. Soc., 2010,132, 12847; ( f ) W. Liu, X. Huang, M.-J. Cheng, R. J. Nielsen, W. A.Goddard III and J. T. Groves, Science, 2012, 337, 1322; (g) P. Galloni,M. Mancini, B. Floris and V. Conte, Dalton Trans., 2013, 42,11963.

7 A. K. Vardhaman, C. V. Sastri, D. Kumar and S. P. de Visser,Chem. Commun., 2011, 47, 11044.

8 (a) A. Thibon, V. Jollet, C. Ribal, K. Senechal-David, L. Billon, A. B.Sorokin and F. Banse, Chem.–Eur. J., 2012, 18, 2715; (b) J. Cho,S. Jeon, S. A. Wilson, L. V. Liu, E. A. Kang, J. J. Braymer, M. H. Lim,B. Hedman, K. O. Hodgson, J. S. Valentine, E. I. Solomon andW. Nam, Nature, 2011, 478, 502.

9 Abbreviations used: N4Py, N,N-bis(2-pyridylmethyl)-N-bis(2-pyridyl)-methylamine; Bn-tpen, N-benzyl-N,N0,N0-tris(2-pyridylmethyl)ethane-1,2-diamine.

10 (a) A. Hazell, C. J. McKenzie, L. P. Nielsen, S. Schindler andM. Weitzer, J. Chem. Soc., Dalton Trans., 2002, 310; (b) K. D.Koehntop, J.-U. Rohde, M. Costas and L. Que Jr, Dalton Trans.,2004, 3191.

11 (a) O. Nashiru, D. L. Zechel, D. Stoll, T. Mohammadzadeh, R. A. J.Warren and S. G. Withers, Angew. Chem., Int. Ed., 2001, 40, 417;(b) D. L. Zechel, S. P. Reid, O. Nashiru, C. Mayer, D. Stoll, D. L.Jakeman, R. A. J. Warren and S. G. Withers, J. Am. Chem. Soc., 2001,123, 4350.

12 J. Kaizer, E. J. Klinker, N. Y. Oh, J.-U. Rohde, W. J. Song, A. Stubna,J. Kim, E. Munck, W. Nam and L. Que Jr., J. Am. Chem. Soc., 2004,126, 472.

13 Y. Goto, T. Matsui, S.-I. Ozaki, Y. Watanabe and S. Fukuzumi,J. Am. Chem. Soc., 1999, 121, 9497.

14 F. S. Brown and L. P. Hager, J. Am. Chem. Soc., 1967, 89, 719.15 (a) D. Kumar, W. Thiel and S. P. de Visser, J. Am. Chem. Soc., 2011,

133, 3869; (b) R. Latifi, M. A. Sainna, E. V. Rybak-Akimova andS. P. de Visser, Chem.–Eur. J., 2013, 19, 4058.

16 (a) D. Kumar, B. Karamzadeh, G. N. Sastry and S. P. de Visser,J. Am. Chem. Soc., 2010, 132, 7656; (b) D. Kumar, G. N. Sastry andS. P. de Visser, Chem.–Eur. J., 2011, 17, 6196.

Fig. 3 Optimized geometries of bromide oxidation transition states by 1a and1b with bond lengths (in Å) and spin densities and charges (in au).

ChemComm Communication

Publ

ishe

d on

08

Oct

ober

201

3. D

ownl

oade

d by

FO

RD

HA

M U

NIV

ER

SIT

Y o

n 31

/10/

2013

10:

55:3

0.

View Article Online