Role of the secondary coordination sphere in metal-mediated dioxygen activation

3646 Inorg. Chem. 2010, 49, 3646–3660

DOI: 10.1021/ic901550k

Role of the Secondary Coordination Sphere in Metal-Mediated Dioxygen Activation

Ryan L. Shook and A. S. Borovik*

Department of Chemistry, University of California;Irvine, 1102 Natural Sciences II, Irvine,

California 92697-2025

Received August 4, 2009

Alfred Werner proposed nearly 100 years ago that the secondary coordination sphere has a role in determining the

physical properties of transition-metal complexes. We now know that the secondary coordination sphere impacts

nearly all aspects of transition-metal chemistry, including the reactivity and selectivity in metal-mediated processes.

These features are highlighted in the binding and activation of dioxygen by transition-metal complexes. There are clear

connections between control of the secondary coordination sphere and the ability of metal complexes to (1) reversibly

bind dioxygen or (2) bind and activate dioxygen to form highly reactive metal-oxo complexes. In this Forum Article,

several biological and synthetic examples are presented and discussed in terms of structure-function relationships.

Particular emphasis is given to systems with defined noncovalent interactions, such as intramolecular H-bonds

involving dioxygen-derived ligands. To further illustrate these effects, the homolytic cleavage of C-H bonds by

metal-oxo complexes with basic oxo ligands is described.

Introduction

The primary coordination sphere is dominated by covalent

interactions between donor atoms on ligands and metal ions.

Experimental and theoretical investigations have provided

detailed structure-function relationships that have im-

proved the understanding of key properties, such as correla-

tions between electronic and molecular structures and

chemical reactivity, and the development of organometallic

catalysts. In contrast, the secondary coordination sphere

normally utilizes noncovalent interactions and is associated

with determining chemical selectivity.

The growing body of evidence has clearly indicated that

control of both spheres is necessary to achieve highly func-

tional complexes. Yet, few synthetic systems regulate both

the primary and secondary coordination spheres in a pre-

dictable manner. The scarcity of such systems can be linked in

large part to the secondary coordination sphere and the

inability to utilize noncovalent interactions in a reliable and

reproducible manner. Because noncovalent bonds are gen-

erally weak interactions, they tend to be difficult to control

within the secondary coordination sphere, leading to the

formation of various structures and perturbation of function.

These difficulties have led to noncovalent interactions being

often overlooked in the design of metal complexes and being

relegated to solvation phenomena.²

In 1912, Alfred Werner advanced the idea that micro-

environments surrounding transition-metal complexes affect

the structure and function.^1,2This new idea invoked the

concept that metal complexes can interact with other mole-

cules in specific ways, forming secondary-sphere (or outer-

sphere) species. Until that time, molecular chemistry had

developed along the line that the primary coordination

sphere was the only contributor to the properties of metal

complexes. It is true that essential properties of metal com-

plexes are linked to the primary coordination sphere; how-

ever, that alone was not sufficient to explain all of Werner’s

experimental findings. For instance, the secondary coordina-

tion sphere effects were needed in order to explain the

association of amines with coordinatively saturated

[M(acac)₃]ⁿcomplexes and the presence of solvent molecules

in crystals of metal complexes. Werner was unaware of the

chemical forces acting within the secondary coordination

sphere and how they could be utilized to direct chemistry.

Since Werner’s seminal insights, the effects of both the

primary and secondary coordination spheres have been

probed by numerous spectroscopic, analytical, theoretical,

and kinetic methods. Fundamental differences are now

known to exist between the two coordination spheres and

can be understood within the context of chemical bonding.

Our thinking about the influences and usefulness of the

secondary coordination sphere changed with the advent of

host-guest chemistry and the emergence of structural

metallobiochemistry. The idea that metal complexes could

be “recognized” by other molecules via binding to the

coordinated ligands (and not the metal centers) was found

soon after the realization that host compounds, such as

*Towhomcorrespondence should beaddressed. E-mail:aborovik@uci.edu.

(1) (a) Werner, A. Ann Chem. 1912, 386, 1–272. (b) Werner, A. Ber. Dtsch.

Chem. Ges. 1912, 45, 121–130.

(2) Colquhoun, H. M.; Stoddart, J. F.; Williams, D. J. Angew. Chem., Int.

Ed. Engl. 1986, 25, 487–507.

r

pubs.acs.org/IC

Published on Web 04/12/2010

2010 American Chemical Society

Article

Inorganic Chemistry, Vol. 49, No. 8, 2010 3647

crown ethers, cavitands, and cyptands, could bind anions and

small molecules.³These findings led to several reports

of secondary sphere coordination complexes, examples of

which are species formed with crown ethers and metal com-

plexes containing ammine or aquo ligands (Figure 1).^4,5Note

that the intermolecular forces used to form these host-guest

species are hydrogen bonds (H bonds).

The breakthrough in spectroscopic and X-ray diffraction

(XRD) analyses of metalloproteins is arguably the most

important advance in understanding how to integrate the

primary and secondary coordination spheres in inorganic

chemistry. Information obtained from these studies illustrate

in striking clarity that control of both the primary and

secondary coordination spheres can be accomplished in an

intramolecular manner. Protein active sites have numerous

architectural features that aid in regulating function, includ-

ing (1) metal binding sites utilizing endogenous ligands from

amino acid side chains to regulate the primary coordination

spheres, (2) site isolation (when necessary) to avoid unwanted

metal ion-metal ion interactions, (3) channels or paths to

allow external reagents access to the functionally vital metal

center(s), and (4) control of the secondary coordination

sphere through protein-derived architectures. These findings

have challenged chemists to design and prepare synthetic

systems that control both the primary and secondary co-

ordination spheres. This Forum Article describes some of the

biological and synthetic systems that utilize simultaneous

regulation of both spheres to achieve dioxygen binding and

activation. In particular, we will examine how intramolecular

H bonds in the secondary coordination sphere affect func-

tion. Correlations between structure and function will be

highlighted, with emphasis on the benefits in regulating the

secondary coordination sphere during dioxygen activation.

Figure 1. Two examples of secondary coordination sphere interactions

within crystalline lattices: (A) 18-crown-6 [Cu^II(NH₃)₄] and (B) 18-

3

crown-6 [Mn^II(H₂O)₄(Cl)₂]. Only ammine and aquo H atoms are shown

for clarity.

3

Figure 2. Common structural motifs that arise from the activation of

dioxygen by iron, manganese, and copper complexes.

Figure 3. Molecular structures of the homodimeric copper-trafficking

protein Cu(Hah)₂(PDB, 1FEE) (A) and myoglobin (PDB, 1A6M) (B).

The Problem: Preventing Unwanted Multimetallic Species

Although thermodynamically favorable, dioxygen needs

to be associated with other species for O-O bond cleavage

to occur. The association with other species is a kinetic

necessity in order to overcome barriers arising from dioxy-

gen’s triplet ground state. Metal complexes have many of the

requisite properties to bind and activate dioxygen and thus

have been utilized by biological and synthetic systems.

The normal binding of dioxygen involves electron transfer

from the metal center(s) to the dioxygen, forming metal-

superoxo or -peroxo species. The reduction of dioxygen

leads to difficulties for any subsequent chemistry involving

metal-O₂complexes, the most frequent being O-O bond

cleavage and the formation of metal oxide compounds. For

instance, monomeric M-O₂complexes (M = Cu, Fe, and

Mn) are rare because of the thermodynamic propensity to

form multimetallic complexes containing oxo- and hydroxo-

bridged complexes (Figure 2).⁶These species are often

exceedingly stable and do not function as either O₂carriers

or reagents to oxidize substrates.

Site-Isolation Approach. One means to circumvent the

formation of these unwanted oxo/hydroxo-bridged spe-

cies is to physically isolate each metal center to such an

extent that dimerization is prevented. Site isolation is best

illustrated in the placement of active sites in metallopro-

teins. Comparisons of various molecular structures show

correlations between function and the location of the

active sites within the proteins. In general, the active sites

in electron transfer and metallotrafficking proteins

(Figure 3A) are found near the molecular surface, an

obvious evolutionary benefit for these proteins because of

their functional need for these active sites to closely

interact with other species. In contrast, the dioxygen-

binding proteins hemoglobin and myoglobin (Figure 3B)

have active sites somewhat buried within the protein

matrix, presumably, in part, to prevent unwanted metal-

metal interactions with neighboring molecules. This is

further demonstrated by the cytochrome P450 com-

pounds, a family of monooxygenases with internal active

sites.

(3) (a) Pederson, C. J. Angew. Chem., Int. Ed. Engl. 1988, 27, 1021–1027.

(b) Cram, D. J. Angew. Chem., Int. Ed. Engl. 1988, 27, 1009–1020. (c) Lehn,

J .M. Angew. Chem., Int. Ed. Engl. 1988, 27, 89–112.

(4) Colquhoun, H. M.; Stoddart, J. F.; Williams, D. K. J. Chem. Soc.,

Chem. Commun. 1981, 849–850.

Not surprisingly, there have been several efforts to

design synthetic systems that incorporate site-isolation

characteristics. However, few have been as successful as

metalloproteins in the reversible binding of dioxygen or

O₂activation. This is especially true of iron, manganese,

(5) Xianglin, J.; Zuohua, P.; Meicheng, S.; Youqi, T.; Depei, H.; Zihou,

T.; Jinqi, Z. K. Chin. Sci. Bull. 1983, 28, 1334.

(6) (a) Kurtz, D. M. Chem. Rev. 1990, 90, 585–606. (b) Mirica, L. M.;

Ottenwaelder, X.; Stack, T. D. P. Chem. Rev. 2004, 104, 1013–1046. (c) Lewis,

E. A.; Tolman, W. B. Chem. Rev. 2004, 104, 1047–1076. (d) Wu, A. J.;

Penner-Hahn, J. E.; Pecoraro, V. L. Chem. Rev. 2004, 104, 903–938.

3648 Inorganic Chemistry, Vol. 49, No. 8, 2010

Shook and Borovik

Figure 5. Molecular structure of a diiron-(μ-1,2-peroxo) complex con-

taining a bulky binucleating ligand.

Figure 4. Molecular structure of the Fe-O₂picket-fence porphyrin

complex.

adducts of nonheme complexes. For nonheme iron com-

plexes, only three diiron-(μ-1,2-peroxo) species have

been stable enough to be structurally characterized: two

of these are only isolable at temperatures below -20 ꢀC.⁹

Suzuki developed the other complex using a hindered

binucleating ligand containing several appended aryl

groups (Figure 5).^9dThe diiron-(μ-1,2-peroxo) complex

with this ligand is prepared at room temperature and

reversibly binds dioxgyen. An amazing aspect of this

system was that removal of the dioxygen from the

diiron-(μ-1,2-peroxo) complex was achieved in boiling

acetonitrile, conditions expected to lead to irreversible

oxidation to give oxo-bridged species. Two reasons, both

involving the aryl groups, were postulated to account for

the observed chemistry. First, the presence of two aryl

rings on each imidazole moiety caused a weakening of the

Fe-N bonds because of steric repulsions, which, in turn,

affected the iron centers’ ability to further reduce the

peroxo ligand. Second, the aryl rings created a sterically

constrained hydrophobic cavity around the metal centers,

which allowed entry of dioxygen but not other species,

especially other diiron complexes that would lead to

irreversible oxidation.

and copper complexes, in which dioxygen binding under

ambient conditions ultimately leads to species with brid-

ging cores (Figure 2). With this said, there have been some

notable examples demonstrating that steric bulk within

the secondary coordination sphere can produce impor-

tant findings. Arguably, the best-known example of this

approach is the ferrous “picket-fence” porphyrin com-

plex of Collman (Figure 4).⁷The steric bulk in this system

comes from four functionalized aryl groups attached at

the meso positions of the porphyrin ring. One regioisomer

has all of the bulky amide groups positioned on the same

side of the ring, producing a sterically constrained cavity

that is suitable for dioxygen binding. Efficient dioxygen

binding requires a five-coordinate iron(II) center, which

is accessible by binding of an imidazole ligand to the

vacant site on the unhindered side of the complex. Im-

portantly, the cavity is sufficiently confined to prevent the

formation of six-coordinate, bis(imidazole) species.

There is enough room to have smaller molecules, such

as dioxygen, enter the cavity and bind to the iron center:

there are several reports illustrating the reversible binding

of dioxygen and other small molecules under ambient

conditions. Over the years, Collman have used this system

to explore the fundamental aspects of dioxygen binding

and, more recently, elaborated on the design to investi-

gate other multimetal enzymatic systems, including cyto-

chrome c oxidase.⁸

A similar use of sterically bulky groups was needed to

0

control dioxygen binding in copper systems with Tp^R,R

ligands [Tp = hydrotris(3,5-R,R⁰-pyrazolyl)borates].

Early work with [Cu^ITp^Me,Me] systems gave some indica-

tions that a Cu-O₂adduct could be stabilized, but the

data were incomplete because of complications from

competing side reactions.¹⁰Kitajima realized that chan-

ging the R groups to Tp^tBu,iPrwould provide enough

steric limitation to isolate the first copper-superoxo

complex (Figure 6A).¹¹Moreover, a change to isopropyl

groups at the 3 position of the pyrazole rings led to

a relatively stable dimeric copper-peroxo complex,

with a μ-η²:η²-peroxodicopper(II) core (Figure 6B).¹²

Even though the isopropyl groups did not directly inter-

act with the μ-η²:η²-peroxodicopper(II) core, their pre-

sence was clearly important by permitting dimerization

but limiting other intermolecular oxidative processes.

Incorporating steric bulk into other ligand types

has also been instrumental in obtaining stable dioxygen

(7) (a) Jameson, G. B.; Molinaro, F. S.; Ibers, J. A.; Collman, J. P.;

Brauman, J. I.; Rose, E.; Suslick, K. S. J. Am. Chem. Soc. 1978, 100, 6769–

6770. (b) Collman, J. P.; Brauman, J. I.; Doxsee, K. M.; Halbert, T. R.;

Bunnenberg, E.; Linder, R. E.; LaMar, G. N.; Del Gaudio, J.; Lang, G.; Spartalian,

K. J. Am. Chem. Soc. 1980, 102, 4182–4192. (c) Collman, J. P.; Brauman, J. I.;

Iverson, B. L.; Sessler, J. L.; Morris, R. M.; Gibson, Q. H. J. Am. Chem. Soc.

1983, 105, 3052–3064. (d) Collman, J. P.; Zhang, Z.; Wong, K.; Brauman, J. I.

J. Am. Chem. Soc. 1994, 116, 6245–6251. (e) Momenteau, M.; Reed, C. A.

Chem. Rev. 1994, 94, 659–698.

(8) Collman, J. P.; Boulatov, R.; Sunderland, C. J.; Fu, L. Chem. Rev.

2004, 104, 561–588.

(9) (a) Dong, Y.; Yan, S.; Young, V. G., Jr.; Que, L., Jr. Angew. Chem.,

Int. Ed. Engl. 1996, 108, 618–620. (b) Kitajima, N.; Tamura, N.; Amagai, H.;

Fukui, H.; Moro-oka, Y.; Mizutani, Y.; Kitagawa, T.; Mathur, R.; Heerwegh, K.;

Reed, C. A.; Randall, C. R.; Que, L., Jr.; Tatsumi, K. J. Am. Chem. Soc. 1994,

116, 9071–9085. (c) Kim, K.; Lippard, S. J. J. Am. Chem. Soc. 1996, 118, 4914–

4915. (d) Ookkubo, T.; Sugimoto, H.; Nagayama, T.; Masuda, H.; Sato, T.;

Tanaka, Y.; Maeda, Y.; Okawa, H.; Hayashi, Y.; Uehara, A.; Suzuki, M. J. Am.

Chem. Soc. 1996, 118, 701–702.

(10) Thompson, J. S. J. Am. Chem. Soc. 1984, 106, 4057–4058.

(11) Fujisawa, K.; Tanaka, M.; Moro-oka, Y.; Kitajima, N. J. Am. Chem.

Soc. 1994, 116, 12079–12080.

(12) (a) Kitajima, N.; Fujisawa, K.; Moro-oka, Y. J. Am. Chem. Soc.

1989, 111, 8975–8976. (b) Kitajima, N.; Fujisawa, K.; Fujimoto, C.; Moro-oka,

Y.; Hashimoto, S.; Kitagawa, T.; Toriumi, K.; Tatsumi, K.; Nakamura, A. J. Am.

Chem. Soc. 1992, 114, 1277–1291.

Article

Inorganic Chemistry, Vol. 49, No. 8, 2010 3649

Figure 6. Molecular structures of copper-superoxo complex [CuTp^t-Bu,iPr

(η²-O₂)] (A) and μ-η²:η²-peroxodicopper(II) complex [Cu^IITp^iPr,iPr]₂-

(O₂) (B) showing the controlling effects of steric bulk.

-

Note that this η²:η²-peroxodicopper(II) complex was a

landmark discovery in synthetic bioinorganic chemistry

because it correctly predicted the active-site structure

of the oxygenated form of the respiratory protein hemo-

cyanin. Furthermore, this series of copper complexes

provides an excellent illustration of how incremental

adjustments within the secondary coordination sphere

can have a major impact on functional aspects of dioxy-

gen binding to transition-metal complexes.

More recently, bulkier ligand systems have been em-

ployed to create larger cavities around metal centers for

dioxygen binding. Notable examples are the calix[6]arene

ligands of Reinaud, which contain various tripodal li-

gands for the binding of a single metal ion. Application of

these systems toward dioxygen activation by copper

complexes has been reported.¹³

Figure 7. Active-site structures of the oxygenated forms of human

hemoglobin (A) and nematode hemoglobin (B) illustrating the different

H-bonding networks.

imidazolyl residue of histidine and the coordinated dioxy-

gen.^15,16Altering the position of the distal histidine in

these proteins has pronounced functional effects, an

example of which is the enhanced peroxidase activity

observed for myoglobins, whose active sites have been

engineered to eliminate bifurcated H-bond formation.¹⁷

Moreover, changes in the H-bonding network can influ-

ence the O₂affinity, as found when comparing the hemo-

globins from human and parasitic nematode Ascaris suum

(Figure 7B).¹⁸Nematodes normally exist in environments

having low dioxygen concentrations, especially when

compared to those experienced by humans. Notice that

the primary coordination spheres around the iron centers

are identical in the two hemoglobins. However, their

H-bonding networks differ, which appears to affect the

dioxygen binding, causing the nematode hemoglobin to

have 10⁴times greater affinity for O₂.

Hydrogen bonds formed within the secondary coordi-

nation sphere also influence dioxygen activation in a

number of iron-heme enzymes. For instance, the func-

tion of cytochrome P450 compounds, heme-containing

monooxygenases that are ubiquitous in nature, has

been linked with a protein-derived H-bonding network

within the secondary coordination sphere of the iron

center (vide infra).¹⁹In addition, recent structural studies

have implicated active-site H-bonding networks in the

function of nonheme halogenases.²⁰SyrB2, a nonheme

iron(II)/R-ketoglutarate (RKG)-dependent halogenase,

catalyzes the conversion of ^L-Thr to 4-Cl-L-Thr, a mod-

ification that is essential for the antifungal activity of

syringomycin E.²¹The proposed mechanism involves Cl^-,

H-Bonding Networks

Metalloproteins. Further control of the binding and

activation of dioxygen is possible through the use of H

bonds within the secondary coordination sphere. Exam-

ples of this approach using intermolecular H bonds have

already been discussed in connection with the host-guest

chemistry. Intramolecular H bonds are now recognized to

also be extremely important in controlling chemistry

occurring at metal centers. Thus, coupling of the effects

of site isolation and the H-bonding network, both within

the secondary coordination sphere, can be used to impact

metal-mediated transformations.

The best examples of intramolecular H-bonding

networks are from biomolecules. H bonds have bond

strengths ranging from 5 to 15 kcal/mol in the condensed

phase, which are generally much weaker than covalent

interactions found within the primary spheres.¹⁴Even

though H bonds are weak interactions, they are essential

for the binding of dioxygen in hemoglobins and myoglo-

bins. As depicted in Figure 7A with human hemoglobin,

dioxygen binds to the iron centers in these proteins via a

covalent Fe-O bond and a H bond between the distal

(16) (a) Shaanan, B. Nature 1982, 296, 683–684. (b) Condon, P. J.; Royer,

W. E., Jr. J. Biol. Chem. 1994, 269, 25259–25267. (c) Philips, S. E. V.;

Schoenborn, B. P. Nature 1981, 292, 81–82.

(17) Ozaki, S.; Roach, M. P.; Matsui, T.; Watanabe, Y. Acc. Chem. Res.

2001, 34, 818–825.

(18) Yang, J.; Kloek, a. P.; Goldberg, D. E.; Matthews, F. S. Proc. Natl.

Acad. Sci. U.S.A. 1995, 92, 4224–4228.

(19) (a) Martinis, S. A.; Atkins, W. M.; Stayton, P. S.; Sligar, S. G. J. Am.

Chem. Soc. 1989, 111, 9252–9253. (b) Gerber, N. C.; Sligar, S. G. J. Am. Chem.

Soc. 1992, 114, 8742–8743. (c) Schlichting, I.; Berendzen, J.; Chu, K.; Stock,

A. M.; Maves, S. A.; Benson, D. E.; Sweet, R. M; Ringe, D.; Pestko, G. A.; Sligar,

S. G. Science 2000, 287, 1615–1622.

(20) For recent reviews on halogenases, see:(a) Vaillancourt, F. H.; Yeh,

E.; Vosburg, D. A.; Garneau-Tsodikova, S.; Walsh, C. T. Chem. Rev. 2006,

106, 3364–3378. (b) Blasiak, L. C.; Drennan, C. L. Acc. Chem. Res. 2008, 42,

147–155.

(13) (a) Blanchard, S.; Le Clainche, L.; Rager, M.-N.; Chansou, B.;

Tuchagues, J. P.; Duprat, A.; Le Mest, Y.; Reinaud, O. Angew. Chem.,

ꢀ ꢁ

Int. Ed. 1998, 37, 2732–2735. (b) Rondelez, Y.; Seneque, O.; Rager, M.-N.;

Duprat, A. F.; Reinaud, O. Chem.;Eur. J 2000, 6, 4218–4226. (c) Izzet, G.;

ꢀ

Douziech, B.; Prange, T.; Tomas, A.; Jabin, I.; Le Mest, Y.; Reinaud, O. Proc.

Natl. Acad. Sci. U.S.A. 2005, 102, 6831–6836. (d) Izzet, G.; Zeitouny, J.;

Akdas-Killig, H.; Frapart, Y.; Menage, S.; Douziech, B.; Jabin, I.; Le Mest, Y.;

Renaud, O. J. Am. Chem. Soc. 2008, 130, 9514–9523.

(14) (a) Elmsley, J. Chem. Soc. Rev. 1980, 9, 91–124. (b) Etter, M. C. Acc.

Chem. Res. 1990, 23, 120–126.

(15) (a) Perutz, M. F.; Fermi, G.; Luisi, B.; Shaanan, B.; Liddington,

R. C. Acc. Chem. Res. 1987, 20, 309–321. (b) Springer, B. A.; Sligar, S. G.;

Olsen, J. S.; Philips, G. N., Jr. Chem. Rev. 1994, 94, 699–714.

(21) Vaillancourt, F. H.; Yin, J.; Walsh, C. T. Proc. Natl. Acad. Sci. U.S.

A. 2005, 102, 10111–10116.

3650 Inorganic Chemistry, Vol. 49, No. 8, 2010

Shook and Borovik

dipole contribution that aids in the stabilization of the

dioxygen adduct. The situation found with the picket-

fence porphyrin complexes highlights a design problem

with these systems for producing intramolecular H-bond-

ing networks. The meso-C atoms are the most convenient

positions to append H-bond donors/acceptors; however,

attachment at these positions places functional groups

too far from the metal center to engage in H bonds.

To circumvent this problem, porphyrins have been de-

signed to have rigid groups that “hang-over” the metal

centers and thus bring H-bond donors/acceptors within

critical distances to form intramolecular H bonds

(Figure 9). Reed’s urea/amide-appended porphyrins²⁸

and Chang’s use of Kemp’s acid-appended porphyrins

are two of the most prominent examples of this ap-

proach.²⁹Many of these models show a significant in-

crease in affinity for dioxygen binding over comparable

systems lacking H-bond donors. These results are con-

sistent with the idea that intramolecular H bonds help to

stabilize an Fe-O₂adduct. In addition, Nocera have

prepared “Hangman” porphyrins that use xanthane units

to position H-bond donors. In one example, they showed

that a small network of H bonds is present around a

Fe^III-OH unit, in both solution and the solid state

(Figure 9C).³⁰

Other Synthetic Systems. Nonheme systems have also

been developed to incorporate functional groups that

are capable of forming intramolecular H-bonding net-

works.^31,32One of the first examples pertinent to dioxy-

gen binding was reported by Kitajima, who isolated

two isomers of a monomeric manganese(III)-peroxo

complex that differ only by the absence or presence of

an intramolecular H bond (Figure 10A).³³Butler have

reported an intriguing [V^V(η²-O₂)(O)] complex with a single

intramolecular H bond (Figure 10B).³⁴By far the most

investigated nonheme systems are those containing tripodal

ligands. Masuda et al. were the first to show the applicability

of these systems using polypyridine tripods containing

Figure 8. Active-site structure of SyrB2 showing the hydrogen-bonding

network surrounding the Cl^-ion. The red spheres represent water

molecules.

RKG, and O₂and proceeds through a [(Cl)Fe^IVdO] inter-

mediate (Figure 8). XRD studies on the reduced form of

SyrB2 revealed that Cl^-is coordinated to the iron(II) center,

replacing a carboxylate ligand that is normally found in this

class of enzymes.²²The Fe^II-Cl distance is unusually long

at 2.44 A,²³in part because of two intramolecular H bonds

˚

involving Cl^-(Figure 8).²⁴The functional consequences of

these noncovalent interactions await further studies but

may assist in the selective transfer of a Cl atom to the

substrate.

Synthetic Heme Systems. Introducing similar hydrogen-

bonding networks into synthetic systems has been

challenging. Difficulties occur because of our inability

to develop systems that are sufficiently rigid to promote

intramolecular H-bond formation. The usual situations

are that H-bond formation involving metal complexes

occurs with other species present in the reaction mixture,

leading to intermolecular products that do not have the

chemical properties to promote the targeted function. To

overcome this drawback, synthetic complexes must in-

corporate rigid frameworks that are positioned proximal

to metal centers and contain H-bond donors or acceptors.

These types of systems often require multistep syntheses

that are difficult and time-intensive.

There have been some impressive examples of hydro-

gen-bonding networks within synthetic complexes.

The H-bonding networks in hemoglobins and myogolo-

bins have inspired investigations into the role of H bonds

in binding dioxygen by synthetic heme complexes,²⁵all

of which contain H-bond donors appended from the

meso positions of the porphyrin rings. Collman’s pick-

et-fence iron porphyrin is an example of such a system

(Figure 4):^26,27although placement of the amide groups is

not within H-bonding distances, there appears to be a

(28) Wuenschell, G. E.; Tetreau, C.; Lavalette, D.; Reed, C. A. J. Am.

Chem. Soc. 1992, 114, 3346–3355.

ꢀ

(29) Chang, C. K.; Liang, Y.; Aviles, G.; Peng, S.-M. J. Am. Chem. Soc.

1995, 117, 4191–4192.

(30) (a) Yeh, C.-Y.; Chang, C. J.; Nocera, D. G. J. Am. Chem. Soc. 2001,

123, 1513–1514. (b) Chang, L. L.; Chang, C. J.; Nocera, D. G. Org. Lett. 2003, 5,

2421–2424.

(31) (a) Mareques Rivas, J. C.; de Rosales, R. T. M.; Parsons, S. Dalton

Trans. 2003, 2156–2163. (b) Mareques Rivas, J. C.; Salvagni, E.; de Rosales, R.

T. M.; Parsons, S. Dalton Trans. 2003, 3339–3349. (c) Mareques Rivas, J. C.;

Prabaharan, R.; de Rosales, R. T. M. Chem. Commun. 2004, 76–77. (d) Mareques

Rivas, J. C.; Salvagni, E.; Parsons, S. Chem. Commun. 2004, 460–461. (e)

Mareques Rivas, J. C.; de Rosales, R. T. M.; Parsons, S. Chem. Commun. 2004,

610–611. (f) Mareques Rivas, J. C.; Prabaharan, R.; Parsons, S. Dalton Trans.

2004, 1648–1655. (g) Mareques Rivas, J. C.; Prabaharan, R.; de Rosales, R. T. M.;

Metteau, L.; Parsons, S. Dalton Trans. 2004, 2800–2807. (h) Mareques Rivas, J.

C.; Salvagni, E.; Parsons, S. Dalton Trans. 2004, 4185–4192. (i) Feng, G.;

Mareque-Rivas, J. C.; de Rosales, R. T. M.; Williams, N. H. J. Am. Chem. Soc.

2005, 127, 13470–13471. (j) Feng, G.; Mareque-Rivas, J. C.; Williams, N. H.

Chem. Commun. 2006, 1845–1847. (k) Mareque-Rivas, J. C.; Hinchley, S. L.;

Metteau, L.; Parsons, S. Dalton Trans. 2006, 2316–2322. (l) Metteau, L.;

Parsons, S.; Mareque-Rivas, J. C. Inorg. Chem. 2006, 45, 6601–6603. (m) Feng,

G.; Natale, D.; Prabaharan, R.; Mareque-Rivas, J. C.; Williams, N. H. Angew.

Chem., Int. Ed. 2006, 45, 7056–7059.

(22) Hegg, E. L.; Que, L., Jr. Eur. J. Biochem. 1997, 250, 625–629.

(23) Orpen, A. G.; Brammer, L.; Allen, F. H.; Kennard, O.; Watson, D.

G.; Taylor, R. In International Tables for Crystallography; Wilson, A. J. C.,

Ed.; Kluwer Academic Publishers: Dordrecht, The Netherlands, 1995; Vol. C,

pp 707-791.

(24) (a) Blasiak, L. C.; Vaillancourt, F. H.; Walsh, C.; Drennan, C. L.

Nature 2006, 440, 368–371. (b) Wong, C.; Fujimoir, D. G.; Walsh, C. T.;

Drennan, C. L. J. Am. Chem. Soc. 2009, 131, 4872–4879.

(25) Momenteau, M.; Reed, C. A. Chem. Rev. 1994, 94, 659-698 and

references cited therein.

(26) (a) Collman, J. P. Acc. Chem. Res. 1977, 10, 265–272. (b) Jameson,

G. B.; Drago, R. S. J. Am. Chem. Soc. 1985, 107, 3017–3022.

(27) Collman, J. P.; Zhang, X.; Wong, K.; Bauman, J. I. J. Am. Chem.

Soc. 1994, 116, 6245–6251.

(32) (a) Berreau, L. M.; Allred, R. A.; Makowski-Grzyska, M. M.; Arif,

A. M. Chem. Commun. 2000, 1423–1424. (b) Berreau, L. M.; Makowska-

Grzyska, M. M.; Arif, A. M. Inorg. Chem. 2001, 40, 2212–2213.

(33) Kitajima, N.; Komatsuzaki, H.; Hikichi, S.; Osawa, M.; Moro-oka,

Y. J. Am. Chem. Soc. 1994, 116, 11596–11597.

(34) Kinblin, C.; Bu, X.; Bulter, A. Inorg. Chem. 2002, 41, 158–160.

Article

Inorganic Chemistry, Vol. 49, No. 8, 2010 3651

Figure 9. Hydrogen-bonding porphyrin systems: urea-modified “picket-fence” porphyrin (only the urea group is shown for clarity) (A),²⁸system

developed by Chang (B),²⁹Fe^III-OH “Hangman” porphyrin complex (C).³⁰

Figure 10. Molecular structures of metal-peroxo complexes with intramolecular H bonds: [Mn^III(O₂)] (A);³³[V^V(O₂)(O)]^-(B);³⁴[Cu^II(OOH)]^þ(C).^35c

appended H-bond donors.³⁵They have done extensive

work in this field and have motivated many workers to

explore the effects of intramolecular H-bonding networks.

For instance, they developed a series of tris(carboxyamido-

pyridyl) tripods, in which the hindered amide groups serve

as H-bond donors. One of these ligands was used to isolate

the first example of a stable Cu^II-OOH complex, in which

two intramolecular H bonds are formed between the ligand

and proximal O atom of the hydroperoxide (Figure 10C).^35c

This complex was not generated from dioxygen, but rather it

was formed using a copper(II) precursor complex and

hydrogen peroxide. Nevertheless, it is an important con-

tribution to copper-peroxo chemistry and illustrates the

usefulness of intramolecular H-bonding networks.

Scheme 1

(35) (a) Harata, M.; Jitsukawa, K.; Masuda, H.; Einaga, H. Chem. Lett.

1995, 24, 61–62. (b) Ogo, S.; Wada, S.; Watanabe, Y.; Iwase, M.; Wada, A.;

Harata, M.; Jitsukawa, K.; Masuda, H.; Einage, H. Angew. Chem., Int. Ed. 1998,

37, 2101–2104. (c) Wada, A.; Harata, M.; Hasegawa, K.; Jitsukawa, K.; Masuda,

H.; Mukai, M.; Kitagawa, T.; Einaga, H. Angew. Chem., Int. Ed. 1998, 37, 798–

799. (d) Ogo, S.; Yamahara, R.; Roach, M.; Suenobu, T.; Aki, M.; Ogura, T.;

Kitagawa, T.; Masuda, H.; Fukuzumi, S.; Watanabe, Y. Inorg. Chem. 2002, 41,

5513–5520. (e) Yamaguchi, S.; Wada, A.; Funahashi, Y.; Nagatomo, S.; Kitagawa,

T.; Jitsukawa, K.; Masuda, H. Eur. J. Inorg. Chem. 2003, 4378–4386. (f) Wada, A.;

Ogo, S.; Nagatomo, S.; Kitagawa, T.; Watanabe, Y.; Jitsukawa, K.; Masuda, H.

Inorg. Chem. 2002, 41, 616–618.

We have explored dioxygen binding to manganese(II)

complexes with the hybrid tripodal ligands H₅bupa and

H₃bpaa (Schemes 1 and 2). Both ligands utilize the

carboxyamidopyridyl moieties introduced by Masuda

et al. and at least one other anionic donor from either a

3652 Inorganic Chemistry, Vol. 49, No. 8, 2010

Shook and Borovik

Scheme 2. Synthetic Route to the Hybrid Ligand H₃bpaa^a

^aConditions: (a) C₇H₉N, Et₃N, THF, 67 ꢀC, 76%; (b) C₆H₁₀, 20% Pd(OH)₂/C, EtOH, 78 ꢀC, 69%; (c) C₈H₇BrFNO, Et₃N, THF, 67 ꢀC, 92%.

II

˚

Figure 11. Synthetic details and molecular structure for [Mn Hbpaa]. Selected distances (A): Mn1-N1, 2.287(1), Mn1-N2, 2.157(1), Mn1-N3,

2.279(1), Mn1-N4, 2.266(2), Mn1-O2, 2.247(1), Mn1-O3, 2.047(1).

deprotonated amide or urea. This design takes advantage

of an anionic primary coordination sphere, which should

aid in the binding of dioxygen, and H-bonding networks

within the secondary coordination sphere to stabilize

M-O₂adducts directly from dioxygen. We have shown

that [Mn^IIH₂bupa]^-reacts with dioxygen to form a peroxo

adduct (Scheme 1), which further reacts with aldehydes to

afford ketones.³⁶These results were the first to show that

monomeric manganese-peroxo complexes can synthe-

tically be prepared from a manganese(II) complex and

dioxygen at room temperature.

H-atom source,³⁸a new species was observed, which

we propose is a manganese(III)-peroxo intermediate

([Mn^IIIH₂bpaa(O₂)]). Spectroscopic measurements sup-

port this assignment: after the addition of dioxygen to

[Mn^IIHbpaa], a new parallel-mode electron paramagnetic

resonance (EPR) signal appeared at a g value of 8.15 with

a six-line hyperfine splitting pattern with a = 57.7 G

(Figure 12A). Variable-temperature studies indicated

that this signal arises from a ground-state doublet with

a D value of -2.0(3) cm^-1and an E/D of 0.13. In addition,

the absorbance spectrum measured at room temperature

(Figure 12 B) has bands at λ_max= 590 nm (58 M^-1cm^-1

)

More recently, we have explored similar processes with

the neutral complex [Mn^IIHbpaa], which uses a new

ligand[Hbpaa]^2-, whose synthesisisoutlined inScheme2.

The solid-state structure of [Mn^IIHbpaa] determined by

XRD (Figure 11) reveals a six-coordinate complex in a

distorted octahedral geometry. Four of the coordination

sites are occupied by N donors from the [Hbpaa]^2-

ligand. The remaining two positions are occupied with

the O atoms from the appended carboxyamido groups,

which adopt conformations such that their carbonyl

moieties are positioned toward the metal center. The O

atoms bind to the manganese(II) center in different ways,

a result of the protonation states of the carboxyamide

groups. One is deprotonated, resulting in the Mn1-O3

and 460 nm (sh):³⁹taken together these spectroscopic

results are nearly the same as those observed for

[Mn^IIIH₃bupa(O₂)]^-and are consistent with a monomeric

manganese(III)-peroxo complex. Isotopic labeling

experiments also support the formulation of this

species as [Mn^IIIH₂bpaa(O₂)]. The generation of [Mn^III

-

H₂bpaa(O₂)] with ¹⁶O₂produces a new peak in the FTIR

spectrum at 891 cm^-1and a strong ion with a mass-

to-charge ratio of 633.2 (633.2) in the electrospray ioniza-

tion mass spectrum (ESI-MS): these features shift to

839 cm^-1and 637.1 (637.2) when samples are prepared

with ¹⁸O₂.

[Mn^IIIH₂bpaa(O₂)] has reactivity similar to those re-

ported for other metal-peroxo complexes. For instance,

treating [Mn^IIIH₂bpaa(O₂)] with cyclohexanecarboxalde-

hyde afforded cyclohexanone in 67% yield. We have also

found that cyclohexanone is formed when [Mn^IIHbpaa]

and the substrate are exposed to excess O₂at room tem-

perature. Presumably, formation of the manganese(III)-

peroxo complex precedesproduct formation, yetat present,

we have not been able to detect [Mn^IIIH₂bpaa(O₂)] under

these reaction conditions.

˚

bond distance of 2.047(1) A, whereas a significantly

˚

longer Mn1-O2 bond of 2.247(1) A is observed for a

protonated carboxyamide group.

Treating [Mn^IIHbpaa] with excess O₂at room tempera-

ture showed no evidence of oxidation, which we ascribed

to the manganese center being coordinatively saturated.³⁷

However, when the same reaction was repeated in the

presence of 1,2-diphenylhydrazine (DPH), a known

(36) Shook, R. L.; Gunderson, W. A.; Greaves, J.; Ziller, J. W.; Hendrich,

M. P.; Borovik, A. S. J. Am. Chem. Soc. 2008, 130, 8888–8889.

(37) [Mn^IIHbpaa] also reacts directly with H₂O₂at room temperature in

DMSO to produce [Mn^IIIH₂bpaa(O₂)] in nearly quantitative yields.

(38) [Mn^IIIH₂bpaa(O₂)] was also produced with indene and fluorine as the

H-atom sources, whereas no reactions were observed in the presence of DHA

and xanthene.

(39) The absorbance spectrum for [Mn^IIIH₂bpaa(O₂)] was prepared by

treating [Mn^IIHbpaa] with H₂O₂at room temperature. Spectra obtained

with dioxygen were similar but contained a strong absorbance from

azobenzene.

Article

Inorganic Chemistry, Vol. 49, No. 8, 2010 3653

generated in the catalytic mechanism of iron-heme peroxi-

dase.^43,44A variety of spectroscopic and low-temperature

structural studies on compounds I and II of horsera-

dish peroxidases (HRPs) have indicated that the oxo of

the ferryl group is H-bonded to amino acid residues in the

active site (Figure 13B).⁴⁵This has led to the proposal that

peroxidase activity is partially regulated by this H-bonding

interaction.

Secondary Coordination Sphere, Synthetic Systems, and

Dioxygen Activation

Inspired by the findings from metalloproteins, we sought

to develop synthetic systems whereby the activation of

dioxygen is controlled, at least in part, by H-bonding net-

works in the secondary coordination sphere.⁴⁶Our objectives

were to understand how the H bonds affected the metal-

mediated activation process and the subsequent dioxygen-

derived species produced. To accomplish these goals, we

designed a series of tripodal ligands that when bonded to

transition-metal ions, created rigid H-bonding cavities

around vacant coordination sites. The rationale behind our

ligands has been reported^46d,eand will only briefly be dis-

cussed. The key component in our designs is the urea

moieties, which was a logical choice based on the vast

literature showing their strong tendency to H-bond in both

solution and the solid state. In most other systems, the two

NH moieties are actively involved in H-bonding to other

functional groups. Our systems use the urea units in a

different way because of the need to simultaneously bind a

metal ion and create an intramolecular H-bonding network.

The design principles are illustrated using the parent symme-

trical tripodal compound, tris[N⁰-tert-butylureayl-N-ethyle-

ne]amine (H₆buea; Figure 14), in which the RNH groups are

deprotonated to form a highly anionic metal-ion-binding

pocket and the remaining parts of the urea groups form the

scaffolding of a cavity. Be aware that the preferred confor-

mation of the urea groups place the R⁰NH groups within the

interior of the cavity, positioning them proximal to the metal

center and in the proper location to H-bond with external

ligands. In addition, using bulky R groups off the R⁰-N atoms

isolates the metal centers and limits unwanted metal-metal

interactions.

Figure 12. Parallel-mode EPR spectrum (black) and simulation (blue)

of [Mn^IIIH₂bpaa(O₂)] (2 mM in THF) recorded at 2.2 K (A) and

absorbance spectra of [Mn^IIH₂bpaa] (---) and [Mn^IIIH₂bpaa(O₂)] (;)

(10 mM in DMSO) measured at room temperature (B). EPR parameters:

microwave frequency and power, 9.26 GHz, 0.2 mW; modulation, 10 G.

EPR simulation parameters: S = 2, g = 2.0, D = -2 cm^-1, E/D = 0.13,

A = 160 MHz.

Dioxygen Activation: Generation of Metal-Oxo Inter-

mediates

The previous sections described approaches to stabilize

dioxygenadducts to metalcomplexes. Closely related to these

efforts are investigations into the fate of the M-O₂adduct, in

particular what is produced after breakage of the O-O bond.

This is an especially important question because of the

pivotal role oxidation that has in industrial and biological

processes. The consensus view is that O-O bond cleavage

produces high-valent metal species, many of which have

terminal oxo ligands that directly oxidize substrates. One

mechanism often invoked is that for the cytochrome P450

compounds, whereby reduction of dioxygen to a terminal

oxo-iron species (compound 1) proceeds through several

intermediates via the controlled flow of electrons and protons

into the active site. The secondary coordination sphere is

instrumental in the activation process, with site isolation and

H-bonding networks being needed to achieve the selective

oxidation of substrates. For instance, substrate binding in

P450, which initiates the catalytic cycles, often utilizes

H bonds formed between the substrate and amino acids

within the active site. H bonds involving the hydroxy

group of a nearby threonine (Thr-252) stabilize the Fe-O₂

moiety (Figure 13A).^40,41This H bond also appears to assist

in O-O bond cleavage because Thr-252 is essential for

the reduction of dioxygen to peroxide during the initial

stages of dioxygen activation. Moreover, a H-bonding

network, created with amino acid residues and structural

water molecules near the iron center, provides a pathway for

protons to traverse through the active site during turnover.

A similar role has been proposed for H bonds in the active

sites of peroxidases and catalases, where H-bonding to

Fe-O₂is believed to polarize the peroxy group, helping

facilitate heterolytic O-O bond cleavage.⁴²Moreover, H

bonds are formed to the high-valent Fe^IVdO units that are

Our first efforts led to the preparation of monomeric

iron(III) and manganese(III) complexes with terminal oxo

ligands. These types of complexes had been problematic to

isolate because of the strong thermodynamic tendency of

(43) Mukai, M.; Nagano, S.; Tanaka, M.; Ishimori, K.; Morishima, I.;

Ogura, T.; Watanabe, Y.; Kitagawa, T. J. Am. Chem. Soc. 1997, 119,

1758-1766 and references cited therein.

€ €

(44) (a) Fulop, V.; Phizackerley, R. P.; Soltis, S. M.; Clifton, I. J.;

Wakatuski, S.; Erman, J.; Hajdu, J.; Edwards, S. L. Structure 1994, 2,

201–208. (b) Bonagura, C. A.; Bhaskar, B.; Shimizu, H.; Li, H.; Sundaramoorthy,

M.; McRee, D. E.; Goodin, D. B.; Poulos, T. L. Biochemistry 2003, 42, 5600–

5608.

€

(45) Berglund, G. I.; Carlsson, G. H.; Smith, A. T.; Szoke, H.; Henriksen,

A.; Hajdu, J. Nature 2002, 417, 463–466.

(46) Selected references:(a) Hammes, B. S.; Young, V. G.; Borovik, A. S.

Angew. Chem., Int. Ed. 1999, 38, 666–669. (b) MacBeth, C. E.; Golombek, A. P.;

Young, V. G., Jr.; Yang, C.; Kuczera, K.; Hendrich, M. P.; Borovik, A. S. O2

Science 2000, 289, 938–941. (c) Gupta, R.; Borovik, A. S. J. Am. Chem. Soc.

2003, 125, 13234–13242. (d) MacBeth, C. E.; Gupta, R.; Mitchell-Koch, K. R.;

Young, V. G.; Lushington, G. H.; Thompson, W. H.; Hendrich, M. P.; Borovik, A.

S. J. Am. Chem. Soc. 2004, 126, 2556–2567. (e) Borovik, A. S. Acc. Chem. Res.

2005, 38, 54–61. (f) Shook, R. L.; Borovik, A. S. Chem. Commun. 2008, 6095–

6107.

(40) (a) Martinis, S. A.; Atkins, W. M.; Stayton, P. S.; Sligar, S. G. J. Am.

Chem. Soc. 1989, 111, 9252–9253. (b) Gerber, N. C.; Sligar, S. G. J. Am. Chem.

Soc. 1992, 114, 8742–8743. (c) Sono, M.; Roach, M. P.; Coulter, E. D.; Dawson,

J. H. Chem. Rev. 1996, 96, 2841–2887.

(41) (a) Schlichting, I.; Berendzen, J.; Chu, K.; Stock, A. M.; Maves, S. A.;

Benson, D. E.; Sweet, R. M; Ringe, D.; Pestko, G. A.; Sligar, S. G. Science

2000, 287, 1615–1622. (b) Nagano, S.; Poulos, T. L. J. Biol. Chem. 2005, 280,

31659–31663.

(42) Poulos, T. L. Adv. Inorg. Biochem. 1988, 7, 1–36.

3654 Inorganic Chemistry, Vol. 49, No. 8, 2010

Shook and Borovik

oxo ligands are low-spin, a finding that agrees with most

bonding schemes. All theories have stipulated that multiple

bonds are needed between the oxo ligand and a metal center

to prepare stable monomeric oxo-metal complexes.⁴⁷

π

bonds arise between filled p orbitals on the oxo ligand and

vacant d orbitals on the metal ion. Calculations predict that

low-spin oxo-metal complexes having C_4vsymmetry are

isolable for metal ions with less than three d electrons because

the initial electrons are housed in a nonbonding orbital

(Figure 16A).⁴⁸More electron-rich systems populate anti-

bonding π orbitals, leading to reactive species. Mayer and

Thorn showed that C₃-symmetricoxo-metalcomplexes with

up to four d electrons are also accessible, again provided that

the complexes are low-spin (Figure 16B).^48,49

Figure 13. Active sites of cytochrome P450 (A) and compound 1 of

horseradish peroxidase (PDB, 1HCH) highlighting the intramolecular

hydrogen-bonding networks surrounding the iron centers.

The properties associated with [Fe^IIIH₃buea(O)]^2-and

[Mn^IIIH₃buea(O)]^2-clearly do not follow the general prin-

ciples used to describe other oxo-metal complexes. We

have probed our systems to understand the fundamental

properties associated with the M-O unit. A combination

of spectroscopic and theoretical studies supports the con-

cept that the intramolecular H-bonding network within the

complexes has a significant effect on the overall stability

of the complexes. Theory showed that the intramolecular H

bonds in a truncated version of [Fe^IIIH₃buea(O)]^2-contrib-

uted 25 kcal/mol to the stability of the complex.⁵⁰This

stabilization was somewhat offset by a decrease in the Fe-O

bond energy by 19 kcal/mol compared to that found in a

computer-generated control system that does not contain

intramolecular H bonds (Figure 17). The weakening of the

Fe-O interaction [Fe^IIIH₃buea(O)]^2-was credited exclu-

sively to a reduction in π bonds between the iron center

and oxo ligand to the extent that there was essentially no

multiple-bond character present. Thus, these findings show

that in some cases H bonds can replace π bonds to oxo-

metal complexes, a result that can have profound implica-

tions on the development of new oxidation catalysts

and understanding of functions of metalloproteins. We

have argued that it may be possible to tune the properties

and reactivity of oxo-metal species via modulation in the

H-bonding networks within the secondary coordination

sphere.^46d-fFor example, adjustments in the metal-oxo

bond energy, as seen in [Fe^IIIH₃buea(O)]^2-, could lead to

new means of controlling the reactivity and selectivity of

oxidation reagents. Moreover, it may be possible to develop

structure-function correlations based on the type of

H-bonding networks surrounding a metal-oxo unit.

Figure 14. Design criteria for complexes with the H-bonding ligand

.

[H₃buea]^3-

these metal ions to form M-μ-O-M species. Moreover,

there are longstanding theoretical predictions that metal

centers with four or more d electrons would not be able to

form terminal oxo species, especially if they were high-spin

complexes (vide infra). Despite these concerns, we were

successful in stabilizing and isolating [M^IIIH₃buea(O)]^2-

directly from the activation of dioxygen (Scheme 3). Isotopic

labeling studies confirmed that dioxygen was the source of

the oxo ligand; a ν(Fe-O) peak at 671 cm^-1was observed for

[Fe^IIIH₃buea(¹⁶O)]^2-

,

which shifted to 645 cm^-1in

[Fe^IIIH₃buea(¹⁸O)]^2-. Similarly, [Mn^IIIH₃buea(¹⁶O)]^2-had

an ν(Mn-O) of 700 cm^-1, which moved to 672 cm^-1in the

¹⁸O isotopomer.

The molecular structures for the two [M^IIIH₃buea(O)]^2-

were determined in the solid state by XRD methods to reveal

trigonal-bipyramidal geometries around the metal centers

(Figure 15). The oxo ligand binds in an axial position

and resides within the H-bonding cavity formed by the

[H₃buea]^3-ligand. Distances between heavy atoms are often

used to gauge whether H bonds are present in the solid state.

Using this criterion, there is strong metrical evidence

for intramolecular H bonds, with each complex having

(47) (a) Nugent, W. A.; Mayer, J. M. Metal-Ligand Multiple Bonds: The

Chemistry of Transition Metal Complexes Containing Oxo, Nitrido, Imido,

Alkylidene, or Alkylidyne Ligands, 1st ed.; Wiley: New York, 1988. (b) Griffith,

W. P. Coord. Chem. Rev. 1970, 5, 459-517. (c) Gulliver, D. J.; Levason, W.

Chem. Soc. Rev. 1982, 46, 1–127.

(48) Mayer, J. M.; Thorn, D. L.; Tulip, T. H. J. Am. Chem. Soc. 1985, 107,

7454-7462 and references cited therein.

(49) Selected examples of other oxo complexes with at least four d

electrons:⁴⁸(a) Wilkinson, G.; Hay-Motherwell, R.; Hussian-Bates, B.;

Hursthouse, M. B. Polyhedron 1993, 12, 2009–2012. (b) Cheng, W.-C.; Yu,

W.-Y.; Cheung, K.-K.; Che, C.-M. J. Chem. Soc., Dalton Trans. 1994, 57–62. (c)

Welch, T. W.; Ciftan, S. A.; White, P. S.; Thorp, H. H. Inorg. Chem. 1997, 36,

4812–4821. (d) Spaltenstein, E.; Conry, R. R.; Critchlow, S. C.; Mayer, J. M. J.

Am. Chem. Soc. 1989, 111, 8741–8742. (e) Rohde, J.-U.; In, J.-H.; Lim, M. H.;

Brennessel, W. W.; Bukowski, M. R.; Stubna, A.; M€unck, E.; Wam, W.; Que, L.,

Jr. Science 2003, 299, 1037–1039. (f) England, J.; Martinho, M.; Farquhar, E. R.;

Frisch, J. R.; Bominaar, E. L.; Munck, E.; Que, L., Jr. Angew. Chem., Int. Ed.

2009, 48, 3622–3626.

0

˚

O1 R N distances of less than 2.9 A. Further support came

3 3 3

from solid-state Fourier transform infrared (FTIR) studies

that showed a shift to lower energy and broadening of the

ν(N-H) peaks compared to H₆buea, features that are indica-

tive of strong H bonds. Notice that the three urea arms of

[H₃buea]^3-are symmetrically distributed around the M^III-O

units in the solid state, which we proposed also occurs in

solution. This premise is supported by the axial EPR spectrum

obtained for [Fe^IIIH₃buea(O)]^2-(g_x,y= 6.0, g_z= 2.0; E/D =

0), data that are consistent with a complex having C₃sym-

metry. Additional magnetic measurements on both complexes

indicated that [Fe^IIIH₃buea(O)]^2-and [Mn^IIIH₃buea(O)]^2-

have spin ground states of S = ⁵/₂and S = 2, respectively.

At first glance, the high-spin character of these complexes

was surprising because most known complexes with terminal

(50) Dey, A.; Hocking, R. K.; Larsen, P. L.; Borovik, A. S.; Hedman, B.;

Hodgson, K. O.; Solomon, E. I. J. Am. Chem. Soc. 2006, 128, 9825–9833.

Article

Inorganic Chemistry, Vol. 49, No. 8, 2010 3655

Scheme 3. Synthetic Procedure for the Isolation of [M^IIIH₃buea(O)]^2-(M^III= Fe, Mn)^a

^aConditions: (a) 4KH, DMA, Ar, rt; (b) M(OAc)₂, DMA, Ar, rt; (c) ¹/₂O₂, DMA, rt.

Figure 15. Molecular structures of [Fe^IIIH₃buea(O)]^2-(A) and

[Mn^IIIH₃buea(O)]^2-(B) determined by XRD methods.

Modulating Intramolecular H-Bonding Networks. The

possibility that H-bonding networks affect the function of

M-O₂and M-O complexes has been proposed to occur

in metalloproteins, where active-site structures and pro-

tein function are known.⁵¹A series of heme proteins

nicely illustrate this possibility. Myoglobin, heme oxyge-

nase, and horseradish peroxidase (HRP) all have the

same square-pyramidal geometries within their primary

coordination spheres and form M-O₂adducts. However,

myoglobin reversibly binds dioxygen, HRP cleaves the

O-O bond to destroy peroxide, and heme oxygenase also

breaks the O-O bond but to oxidize the porphyrin ring

during heme catabolism. These differences in function

could be related to the varied H-bonding networks sur-

rounding the iron centers within their respective active

sites.

Figure 16. Qualitative orbital diagrams for metal-oxo complexes with

C_4v(left) and C₃symmetry (right) (adapted from Mayer and Thorn).⁴⁸

We have developed systems to examine the effects of

varied H-bonding networks on dioxygen activation in

synthetic complexes.⁵²A straightforward design concept

was employed for the metal-containing precursors:

a series of four cobalt(II) complexes were prepared

with trigonal-monopyramidal coordination geometries

but tunable H-bonding networks within their secon-

dary coordination spheres. The investigation required

four ligands, each with a trianionic metal-binding pocket

Figure 17. Bonding decomposition scheme for the Fe^III-O complex

illustrating the effects of the intramolecular H-bonding network.

but differing in the number of H-bond donors. Building on

the work that we had done on the symmetrical urea-based

ligand [H₃buea]^3-and its amidate counterpart [0]^3-, we

prepared the two hybrid ligands [H₂2]^3-and [H1]^3-

(Figure 18). Their corresponding cobalt(II) complexes were

synthesized and characterized to show similar molecular

and electronic structures. Yet, differences in the reactivity

with dioxygen were observed throughout the series of

complexes.

(51) Selected examples:(a) Holmes, M. A.; Stenkamp, R. E. J. Mol. Biol.

€

1991, 220, 723–737. (b) F€ulop, V.; Phizackerley, R. P.; Soltis, S. M.; Clifton, I. J.;

Wakatuski, S.; Erman, J.; Hajdu, J.; Edwards, S. L. Structure 1994, 2, 201–208.

(c) Mukai, M.; Nagano, S.; Tanaka, M.; Ishimori, K.; Morishima, I.; Ogura, T.;

Watanabe, Y.; Kitagawa, T. J. Am. Chem. Soc. 1997, 119, 1758-1766 and

references cited therein. (d) Dunietz, B. D.; Beachy, M. D.; Cao, Y.; Whittington,

D. A.; Lippard, S. J.; Friesner, R. A. J. Am. Chem. Soc. 2000, 122, 2828–2839.

(e) Gherman, B. F.; Dunietz, B. D.; Whittington, D. A.; Lippard, S. J.; Friesner, R.

A. J. Am. Chem. Soc. 2001, 123, 3836–3837. (f) Du Bois, J.; Mizoguchi, T. J.;

Lippard, S. J. Coord. Chem. Rev. 2000, 200-202, 443–485. (g) Berglund, G. I.;

Treating [Co^IIH₃buea]^-and [Co^IIH₂2]^-with dioxygen

(0.5 equiv) under the same experimental conditions

(Figure 18) gave the Co^III-OH complexes in yields of

greater than 70%. Isotopic labeling studies and spectro-

scopic measurements confirmed the identity of these

€

Carlsson, G. H.; Smith, A. T.; Szoke, H.; Henriksen, A.; Hajdu, J. Nature 2002,

417, 463–468. (h) Tomchick, D. R.; Phan, P.; Cymborowski, M.; Minor, W.;

Holman, T. R. Biochemistry 2001, 40, 7509–7517.

(52) Lucas, R. L.; Muhkerjee, J.; Zart, M. K.; Sorrell, T. N.; Powell, D.

R.; Borovik, A. S. J. Am. Chem. Soc. 2006, 128, 15476–15489.

3656 Inorganic Chemistry, Vol. 49, No. 8, 2010

Shook and Borovik

Figure 18. Summary of the dioxygen reactivity for a series of cobalt(II) complexes with varied H-bonding networks.

complexes. [Co^IIH1]^-, a complex containing only one

H-bond donor, showed no appreciable reactivity with

dioxygen under these conditions. Excess dioxygen was

needed to initiate a reaction, with the first observable

species proposed to be a cobalt(II)-superoxo species with

an S = ¹/₂spin ground state. This complex converted to

an intermediate that had spectroscopic properties con-

sistent with a Co^III-OH complex, but it was too unstable

to isolate in pure form. The final complex in the series,

[Co^II0]^-contained no H-bond donors and did not react

with dioxygen under any conditions. Note that we have

ruled out steric constraints of the cavity as the reason for

this lack of dioxygen reactivity. We have numerous

examples of complexes with [0]^3-being able to form

there is a clear relationship between the stability of the

Co^III-OH complexes and the number of intramolecular

H-bond donors. [Co^IIIH₃buea(OH)]^-, with three H-bond

donors, was by far the most stable complex in solution,

leading to its crystallization. The other complexes reacted

to such an extent in solution that crystallization was not

possible. Two reasons could account for the observed

stabilities: (1) increasing the number of H bonds to the

Co^III-OH unit would lower its nucleophilicity, rendering it

less reactive; (2) the additional H bonds with an intra-

molecular network would rigidify the cavity structure,

limiting further reactivity with external species.

The structure-function relationships observed in our

cobalt(II) complexes support the idea that the activity

of a metalloprotein is linked to the H-bonding net-

works within secondary coordination spheres. This rela-

tionship is especially pertinent in systems involving

O₂binding and activation. It appears likely that other

correlations will be established as more data emerge

from both biological and synthetic systems. For example,

there are now specific examples of engineered metallo-

proteins, in which changes in H-bonding networks

within active sites affect the physical properties and

reacitivity.⁵³Moreover, other synthetic systems have

been reported that show strong links between H bonds

and physical properties, such as redox potentials and

pK_avalues.^31,35

five-coordinate complexes, including [Co^II0(CN)]^2-

,

which was isolated and structurally characterized.

The trigonal-monopyramidal cobalt(II) complexes

have different levels of reactivity toward dioxygen despite

their nearly identical primary coordination spheres. We

have suggested that this difference is caused in large part

by the intramolecular H-bonding networks within the

secondary coordination sphere that assist in the binding

of O₂, leading to activation of O₂and formation of

the Co^III-OH complexes. The data showed a distinct

correlation between the number of H bonds and O₂bind-

ing/activation. The connection is apparent when the reac-

tivities of [Co^IIH₂2]^-and [Co^IIH1]^-are compared, both of

which are capable of forming intramolecular H bonds.

[Co^IIH₂2]^-, with two H-bond donors, reacts at lower

O₂concentrations compared with [Co^IIH1]^-, which only

has the possibility of forming one intramolecular H bond.

An additional H-bond donor in [Co^IIH₂2]^-would assist

in dioxygen binding, the essential first step in metal-

mediated O₂activation. Moreover, [Co^II0]^-does not bind

dioxygen and is the only complex in the series that is

incapable of forming intramolecular H bonds. Finally,

Relationship between the Basicity of the Oxo Ligand and

Chemical Reactivity. The isolation of [M^IIIH₃buea(O)]^2-

has led to the formation of related complexes, namely, the

monomeric M^III/II-OH complexes [M^IIIH₃buea(OH)]^-

and [M^IIH₃buea(OH)]^2-, which are stable and have

been completely characterized, including their molecular

(53) (a) Miller, A. F. Acc. Chem. Res. 2008, 41, 501–510. (b) Jackson, T. A.;

Brunold, T. C. Acc. Chem. Res. 2004, 37, 461–470.

Article

Inorganic Chemistry, Vol. 49, No. 8, 2010 3657

structures by XRD. Both iron and manganese hydroxo

complexes have trigonal-bipyramidal coordination geome-

tries comparable to those of the parent metal-oxo com-

plexes with intramolecular H-bonding networks surroun-

ding the M-OH units. We have also been able to detect

[Mn^IVH₃buea(O)]^-, a rare example of a manganese(IV)-

oxo complex. Similar to the M^III-O complexes, the

M^III/II-OH and Mn^IVdO complexes are all high-spin. A

more complete description of the preparative details

and properties of these complexes has been reviewed

recently.^46e,f

We have used these complexes to investigate C-H

cleavage by metal-oxo complexes.^46c,54This type of re-

action has a wide variety of chemical and biology applica-

tions but is one of the most difficult chemical transforma-

tions, mostly because of the high dissociation energies of

C-H bonds.⁵⁵In many substrates, the homolytic bond

dissociation energies are greater than 90 kcal/mol, mak-

ing them thermodynamically difficult to cleave. Like

many researchers, we have turned to biology to gain

insights into how to accomplish this transformation.

Numerous metalloproteins have evolved that achieve this

reaction, and most studies suggest that metal-oxo spe-

cies, derived from the activation of dioxygen, are the

active species involved in C-H bond cleavage. The

oxidation strength of metal-oxo complexes is usually

expressed in terms of redox potentials, with compounds

having higher redox potentials being stronger oxidants.

Our metal-oxo complexes are also derived from dioxy-

gen and are able to cleave C-H bonds; however, their

redox potentials are unusually low for complexes that oxi-

dize substrates. For example, the [Mn^IV/IIIH₃buea(O)]^-/2-

and [Mn^V/IVH₃buea(O)]^/-redox couples are at -1.0

and -0.076 V vs Cp₂Fe^þ/Cp₂Fe. Moreover, we were

unable to observe the [Mn^III/IIH₃buea(O)]^2-/3-process,

suggesting that it is lower than -2.0 V vs Cp₂Fe^þ/Cp₂Fe.

However, [Mn^IIIH₃buea(O)]^2-reacts with substrates

Figure 19. Thermodynamic cycles used to evaluate BDE_OHfor

[Mn^IIH₃buea(OH)]^2-and [Mn^IIIH₃buea(OH)]^-.

atom and is dependent on the solvent and reference used

to measure the redox potentials.⁵⁷

BDE_OH¼ 23:06E₁₌₂þ 1:37pK_aþ C

ð1Þ

The analysis of BDE_OHwas readily accomplished because

the needed Mn-O(H) complexes had already been pre-

pared and characterized as discussed above. From these

studies, we found that the oxo ligands in the Mn^IV/III-O

complexes were basic, with pK_avalues of 28.3 and ∼15

for [Mn^IIIH₃buea(OH)]^-and [Mn^IVH₃buea(OH)]. Taken

together with the measured redox potentials, we obtained

a BDE_OHof 89 kcal/mol for [Mn^IIIH₃buea(OH)]^-and

77 kcal/mol for [Mn^IIH₃buea(OH)]^2-(Figure 19).⁵⁴It

is gratifying to note that these predicted values were in close

agreement with what was observed in our reactivity studies

between substrates having cleavable C-H bonds and the

Mn^IV/III-O complexes.^46c,54

One of the benefits of these thermodynamic analyses

was that they showed the importance that the basicity of

the oxo ligand has on C-H bond cleavage. It is clearly

evident from our work that the strongly basic character of

the oxo groups helped drive the observed reactivity with

external substrates. Furthermore, the large difference in

basicity for the two manganese-oxo complexes predicted

that they should react differently with substrates with

differing acidities. The reactivity with 2,4,6-tri-tert-butyl-

phenol illustrated this point: [Mn^IVH₃buea(O)]^-homo-

lytically cleaved the O-H bond, producing the phenoxyl

radical and [Mn^IIIH₃buea(OH)]^-, whereas the more basic

[Mn^IIIH₃buea(O)]^2-deprotonated the phenol,⁵⁸produ-

cing the phenolate anion and [Mn^IIIH₃buea(OH)]^-.

The formation of this series of Mn-O(H) complexes

also aided in a comparative kinetic analysis of the C-H

bond cleavage from 9,10-dihydroanthracene (DHA). For

containing

a

C-H bond of ∼80 kcal/mol, and

[Mn^IVH₃buea(O)]^-cleaves C-H bonds of ∼90 kcal/mol.

The findings with manganese-oxo complexes sug-

gested that redox potentials alone were not sufficient to

fully explain the observed reactivity. To gain a more

detailed understanding of our systems, we adapted the

approach pioneered by Mayer to use thermodynamic

cycles to evaluate the ability of metal-oxo complexes to

homolytically cleave C-H bonds.⁵⁶The basis of this

approach is to experimentally evaluate the homolytic

O-H bond dissociation energies (BDE_OH) of M-OH

complexes that are produced after C-H bond homolysis:

for this transformation to be thermodynamically achiev-

able, the energy required to cleave the C-H bond must be

comparable to that produced in forming a O-H bond.

BDE_OHcan be obtained from the appropriate experi-

mentally determined one-electron redox potentials and

pK_avalues (Figure 19), using eq 1. The constant C is

included in the analysis to account for solvation of the H

[Mn^IIIH₃buea(O)]^2-at 20 ꢀC, a corrected second-

^III-O 12,59

order rate constant (k^Mn

was determined, which is more than 1 order of magnitude

)

of 0.48(4) M^-1s^-1

^IV-O

larger than the k^Mn

of 0.026(2) M^-1s^-1found for

[Mn^IVH₃buea(O)]^-. This result is counter to what would

have been predicted based on the BDE_OHvalues for the

Mn-OH analogues (Figure 19), in which BDE_OH

of [Mn^IIIH₃buea(OH)]^-is greater than that for

[Mn^IIH₃buea(OH)]^2-by 12 kcal/mol. Insightsintothisun-

expected finding came from the measured kinetic isotope

(54) Parsell, T. H.; Yang, M.-Y.; Borovik, A. S. J. Am. Chem. Soc. 2009,

131, 2762–2763.

(55) Stone, K. L.; Borovik, A. S. Curr. Opin. Chem. Biol. 2009, 13, 114–

118.

(56) (a) Mayer, J. M. In Biomimetic Oxidations Catalyzed by Transition

Metal Complexes; Meunier, B., Ed.; Imperial College Press: London, 2000; pp

1-43. (b) Mayer, J. M. Annu. Rev. Phys. Chem. 2004, 55, 363–390.

(57) (a) Bordwell, F. G.; Cheng, J.-P.; Ji, G.-Z.; Satish, A. V.; Zhang, X. J.

Am. Chem. Soc. 1991, 113, 9790–9795. (b) Parker, V. D.; Handoo, K. L.; Roness,

F.; Tilset, M. J. Am. Chem. Soc. 1991, 113, 7493–7498.

(58) 2,4,6-Tri-tert-phenol: BDE_OH= 85 kcal/mol, pK_a= 18. Bordwell,

F. G.; Zhang, X.-M. J. Org. Chem. 1995, 8, 529–535.

(59) The second-order rate constants were normalized for the four

reactive C-H bonds per substrate molecule.

3658 Inorganic Chemistry, Vol. 49, No. 8, 2010

Shook and Borovik

Figure 20. Proposed mechanism for the reactions of [Mn^IVH₃buea(O)]^-and [Mn^IIIH₃buea(O)]^2-with DHA.

effects of 2.6 and 6.8 found for [Mn^IIIH₃buea(O)]^2-

and [Mn^IVH₃buea(O)]^-, respectively. These values indicate

that both complexes have primary kinetic isotope effects,

yet the large disparity in their values suggested the

possibility that the complexes might react via dif-

ferent mechanistic paths. This premise was supported

by Eyring analysis, which found a 35 eu difference

in the entropy of activation between the two complexes

(ΔS^q= -14(6) eu for [Mn^IIIH₃buea(O)]^2-and ΔS^q=

-49(4) eu for [Mn^IVH₃buea(O)]^-), which had a dominant

effect on the rate constants for C-H bond cleavage.

We proposed a mechanistic explanation for the above

rate data. Our kinetic finding, coupled with the large

difference of ∼15 between pK_avalues of DHA and

[Mn^IVH₃buea(OH)], suggested that these reagents

react via a proton-coupled electron transfer (PCET)

path (Figure 20A).⁶⁰This mechanism involves the coming

together of the substrate and Mn^IV-O complex to form a

molecules, which agree with the more positive ΔS^qvalue

determined for this reaction (Figure 20B).

The mechanistic findings for these two manganese-

oxo complexes highlight the contributions of the pK_a(i.e.,

oxo ligand basicity) on chemical reactivity. Our finding

that the rate of C-H homolysis is affected by oxo basicity

shows that there is another tunable parameter (other than

the redox potential) that can be used in the development

of efficient and durable oxidation catalysts. Furthermore,

this concept may help resolve a longstanding issue in

biochemistry. The evolutionary benefit of coupling O₂

activation with C-H bond cleavage is well recognized

and has been cited as a major factor for the success of

aerobic life. However, dioxygen-derived, metal-contain-

ing species, especially high-valent metal-oxo species,

exist long enough to do irreversible oxidative damage to

a protein active site. One means to prevent such detri-

mental processes is to couple the redox potentials of

the metal-oxo species with the basicity of the oxo ligand.

The benefit of this effect is shown graphically (Figure 21)

for the homolytic cleavage of the C-H bond in methane

(BDE_CH= 104 kcal/mol). At low pK_avalues, the

redox potentials needed to perform this reaction are too

high to sustain function. The metal-oxo center would

certainly attack other species within the active sites, such

as the side chains of amino acids. As the pK_avalue

increases, there is a concomitant lowering of the redox

potential of the metal-oxo species to values that ulti-

mately are compatible with a protein. It is the interplay

between these two fundamental properties that may

lead to the efficient oxidation of substrates within an

oxygenase.

highly ordered transition state. For [Mn^IIIH₃buea(O)]^2-

,

the basicity of the oxo ligand influenced the reaction to

such an extent that a two-step mechanism occurs, in

which proton transfer precedes electron transfer. More-

over, the similar pK_avalues for [Mn^IIIH₃buea(OH)]^-

and DHA (ΔpK_aless than 2) and the primary kinetic

isotope effect suggested rate-limiting proton transfer.

This mechanistic proposal would also produce a charge-

delocalized transition state, leading to less ordered solvent

(60) There have been several discussions in the literature on the proper

nomenclature for this process: two of the most common are H-atom transfer

and PCET. We chose to use the latter in this Forum Article. For more

detailed discussions, see ref 55b and: Huynh, M. H. V.; Meyer, T. J. Chem.

Rev. 2007, 107, 5004–5064.

Article

Inorganic Chemistry, Vol. 49, No. 8, 2010 3659

would be beneficial to the development of new or improved

function.

Experimental Section

General Methods. All reagents were purchased from com-

mercial sources and used as received, unless otherwise noted.

Solvents were sparged with argon and dried over columns

containing Q-5 and molecular sieves. Potassium hydride

(KH), as a 30% dispersion in mineral oil, was filtered with a

medium-porosity glass frit and washed five times each with

pentane and Et₂O. The solid KH was dried under vacuum and

stored under an inert atmosphere. 1,2-Diphenylhydrazine

(DPH) was recrystallized from Et₂O, dried under vacuum, and

stored under an inert atmosphere. ¹⁸O₂(99 atom % ¹⁸O) was

purchased from ICON Isotopes (Summit, NJ). Elemental ana-

lysis was accomplished at Robertson Microlit Laboratories

(Madison, NJ). The syntheses of H₃bpaa and its intermediates

were carried out under a dinitrogen atmosphere. The syntheses

of metal complexes were conducted in a Vacuum Atmospheres,

Co., drybox under an argon atmosphere. N-[6-(Bromomethyl)-

2-pyridyl]pivalamide⁶⁴was prepared according to literature

methods with minor variations.

Figure 21. Proposed structure of compound I in cytochrome P450 (A)

and the relationship between the redox potential and pK_afor a metal-oxo

species in the cleavage of a C-H bond in methane with BDE_C-H= 104

kcal/mol (B).

Green was the first to suggest the importance of the

basicity of the oxo ligand in oxygenases.⁶¹They have been

exploring the intermediates formed during turnover in

P450 compounds and proposed that the redox poten-

tials of the metal-oxo intermediates are much lower

than previously thought, thus limiting oxidative damage.

In his model, the lower redox potential in compound I,

the competent oxidizing species in P450 compounds

(Figure 21), is compensated for by the basicity of the

oxo ligand in compound II, the one-electron-reduced

form of compound I. Green further proposed that this

type of chemistry is possible because of thiolate ligation to

the iron center, which is supported by resonance Raman

studies. Our work described above supports this idea,

as do two recent reports on synthetic manganese-oxo⁶²

and iron-oxo⁶³systems.

Preparative Methods. N-(4-Fluorophenyl)-2-bromoacetamide.

A solution of 4-fluoroaniline (5.00 mL, 52.1 mmol) and Et₃N (8.5

mL, 61 mmol) in 50 mL of CH₂Cl₂was cooled to 0 ꢀC with an ice

water bath. Bromoacetyl bromide (4.60 mL, 53.0 mmol) was

diluted with CH₂Cl₂(∼50 mL) and added dropwise. The solution

was brought to room temperature and stirred overnight. The

solution was concentrated under reduced pressure to a solid,

dissolved in CHCl₃, and washed with 2 M HCl (3 ꢀ 50 mL),

H₂O (1 ꢀ 50 mL), and brine (1 ꢀ 50 mL). The organic layer was

dried over anhydrous sodium sulfate, filtered, and concentrated to

a solid under reduced pressure. The solid was triturated with

CH₂Cl₂and dried under vacuum to afford 8.06 g (67%). Mp:

132-134 ꢀC. ¹H NMR (500 MHz, CDCl₃): δ 8.14 (1H, s, NH),

7.50 (2H, m, ArH), 7.06 (2H, m, ArH), 4.03 (2H, s, CH₂). ¹³C

NMR (500 MHz, CDCl₃): δ 164.7, 159.2, 157.3, 135.0, 121.0,

115.6, 115.4, 30.3. ¹⁹F NMR (400 MHz, CDCl₃): δ 118.8.

Summary

It has been nearly 100 years since Werner first proposed the

significance of the secondary coordination sphere in transi-

tion-metal complexes. Unbeknownst to Werner were many

of the central components that define and control the

chemistry associated with this sphere. Some advances have

been made in determining the regulatory features for the

secondary coordination sphere, as highlighted in this Forum

Article. We know that secondary coordination spheres are

essential for the function of nearly all metalloproteins and

that similar characteristics can be designed into synthetic

metal complexes. Its importance is now obvious in the

chemistry associated with dioxygen binding and activation,

as illustrated with respiratory proteins, oxygenases, and

synthetic complexes. Yet, with all of the progress that has

been made, our utilization of the secondary coordination

sphere still lags far behind that of the primary sphere, in large

part because of our inability to manipulate the noncovalent

interactions that are needed within the secondary sphere.

Examples from our group and others showed that the design

of systems with rigid scaffolds allows the positioning of

functional groups to promote the formation of intramolecu-

lar H bonds while at the same time isolating metal centers in a

manner similar to what occurs within a protein active site.

These principles should be applicable to most chemical

systems using transition-metal ions and undoubtedly

Bis[N-[6-pivalamido-2-(pyridylmethyl)benzyl]]amine. A solu-

tion of N-[6-(bromomethyl)-2-pyridyl]pivalamide (6.11 g, 22.5

mmol), benzylamine (1.2 mL, 11 mmol), and Et₃N (3.5 mL,

25 mmol) in 100 mL of THF was refluxed for 63 h. The solution

was cooled to room temperature, filtered, and concentrated

under reduced pressure to a yellow oil. The product was purified

by silica gel column chromatography using 2:3 hexanes/EtOAc

1

to yield 4.06 g (76%). Mp: 97-106 ꢀC. H NMR (500 MHz,

CDCl₃): δ 8.10 (2H, d, J = 8.3 Hz, NHArH), 7.98 (2H, s, NH),

7.68 (2H, t, J = 7.9 Hz, NHArH), 7.39 (2H, d, J = 7.5 Hz,

NHArH), 7.32 (4H, m, ArH), 7.24 (1H, m, ArH), 3.68 (4H, s,

NHArCH₂), 3.67 (2H, s, NCH₂Ar), 1.32 (18H, s, C(CH₃)₃). ¹³

C

NMR (500 MHz, DMSO): δ 177.5, 158.4, 152.0, 139.1, 139.0,

129.1, 128.8, 127.5, 117.8, 112.9, 59.3, 58.2, 39.8, 27.4.

Bis[N-(6-pivalamido-2-pyridylmethyl)]amine. To a solution of

bis[N-[6-pivalamido-2-(pyridylmethyl)]benzyl]amine (5.54 g,

11.4 mmol) and cyclohexene (33.0 mL, 326 mmol) in 80 mL of

EtOH was added 20% Pd/C (0.284 g). The suspension was

refluxed overnight, cooled to room temperature, and filtered

through a pad of Celite. The filtrate was concentrated under

reduced pressure to a yellow solid, washed with Et₂O, and then

dried under vacuum to afford 2.54 g (69%). Mp: 127-128 ꢀC.

¹H NMR (500 MHz, CDCl₃): δ 8.13 (2H, d, J = 8.3 Hz, ArH),

7.99 (2H, s, NHAr), 7.87 (2H, t, J = 7.9 Hz, ArH), 7.04 (2H, d,

J = 7.5 Hz, ArH), 3.86 (4H, s, ArCH₂), 1.33 (18H, s, C(CH₃)₃).

¹³C NMR (500 MHz, CDCl₃): δ 177.1, 157.8, 151.3, 138.8,

118.2, 112.1, 54.3, 39.9, 27.6.

(61) (a) Green, M. T.; Dawson, J. H.; Gray, H. B. Science 2004, 304, 1653–

165. (b) Green, M. T. Curr. Opin. Chem. Biol. 2009, 13, 84–88.

(62) Lansky, D. E.; Goldberg, D. P. Inorg. Chem. 2006, 45, 5119–5125.

(63) Sastri, C. V.; Lee, J.; Oh, K.; Lee, Y. J.; Lee, J.; Jackson, T. A.; Ray,

K.; Hirao, H.; Shin, W.; Halfen, J. A.; Kim, J.; Que, L., Jr.; Shaik, S.; Nam,

W. Proc. Natl. Acad. Sci. U.S.A. 2007, 104, 19181–19186.

(64) Harata, M.; Hasegawa, K.; Jitsukawa, K.; Masuda, H.; Einaga, H.

Bull. Chem. Soc., Jpn. 1998, 71, 1031–1038.

3660 Inorganic Chemistry, Vol. 49, No. 8, 2010

Shook and Borovik

N-[Bis(6-pivalamido-2-pyridylmethyl)](N⁰-4-fluorophenylcar-

bamoylmethyl)amine (H₃bpaa). A mixture of bis[N-(6-pivalami-

do-2-pyridylmethyl)]amine (3.82 g, 9.61 mmol), N-(4-fluoro-

phenyl)-2-bromoacetamide (2.23 g, 9.61 mmol), and Et₃N

(1.4 mL, 10.0 mmol) in 90 mL of THF was refluxed overnight.

The reaction mixture was allowed to cool to room temperature,

95 μmol) in 3 mL of anhydrous DMSO was treated with H₂O₂

(6 μL, 160 μmol, 71.2% aqueous) and stirred for 5 min.

MS (ES^-). Mass calcd for C₃₀H₃₅FN₆O₅Mn [M - H]: m/z

633.2. Found: m/z 633.2. λ_max[DMSO, nm (ε, M^-1cm^-1)]: 590,

460 (sh).

Reaction of [Mn^IIIHbpaa(O₂)] with Cyclohexanecarboxalde-

hyde. [Mn^IIIHbpaa(O₂)] was prepared by the method above and

allowed to stir for 10 min. A total of 1 equiv of cyclohexane-

carboxaldehyde was added via a syringe and allowed to stir

an additional 3 h. After treatment with H₂O (3 mL), the solution

was washed with Et₂O (3 ꢀ 3 mL). The Et₂O washings

were combined in a 10 mL volumetric flask and diluted. The

solution was analyzed by gas chromatography (GC), and the

amount of cyclohexanone was determined from calibration curves.

Physical Methods. Electronic absorbance spectra were re-

corded with a Cary 50 spectrophotometer using a 1.00 mm

quartz cuvette. FTIR spectra were collected on a Varian 800

and the Et₃N HBr that precipitated from the reaction was

3

removed by filtration. The filtrate was concentrated under re-

duced pressure to a yellow-orange solid, dissolved in CH₂Cl₂, and

washed with H₂O (3 ꢀ 100 mL) and brine (1 ꢀ 100 mL). The

organic layer was dried over anhydrous sodium sulfate, filtered,

and concentrated under reduced pressure to a solid. The solid was

triturated with Et₂O and dried under vacuum to afford 4.83 g

(92%). Mp: 64-67 ꢀC. ¹H NMR (500 MHz, CDCl₃): δ 10.12 (1H,

s, NHArF), 8.15 (2H, d, J = 8.3 Hz, ArH), 7.88 (2H, s, NHAr),

7.66 (2H, t, J = 7.9 Hz, ArH), 7.51 (2H, m, HArF), 7.01 (2H, m,

HArF), 6.97 (2H, d, J = 7.4 Hz, ArH), 3.82 (4H, s, ArCH₂), 3.51

(2H, s, CH₂CO), 1.27, (18H, s, C(CH₃)₃). ¹³C NMR (500 MHz,

CDCl₃): δ 117.1, 169.7, 158.4, 156.0, 151.6, 139.0, 134.1, 121.8,

119.1, 115.8, 112.8, 65.9, 60.0, 58.8, 39.8, 27.5, 15.3. HRMS (ES^þ).

Exact mass calcd for C₃₀H₃₇N₆O₃FNa [M þ Na]: m/z 571.2809.

Found: m/z 571.2797.

1

Scimitar series FTIR spectrometer. H and ¹³C NMR spectra

were recorded on a Bruker DRX500 spectrometer. EPR spectra

were collected using a Bruker EMX spectrometer equipped with

an ER041XG microwave bridge, an Oxford Instrument liquid-

helium quartz cryostat, and a dual-mode cavity (ER4116DM)

or a Bruker ESP300 spectrometer equipped with an Oxford

ESR910 cryostat and a bimodal cavity (Bruker ER4116DM).

MS of the manganese complexes was done on a Waters LCT

Premier mass spectrometer operated in negative-ion ESI mode.

Identification of the organic products was done on a Thermo

Trace MS^þGC-MS operated in ESI mode. Quantitative GC

analysis was done on a Hewlett-Packard 6890 series gas chro-

matograph equipped with a HP 7683 series injector. A calibra-

tion plot was established using standard procedures for the

quantitative determination of organic products.

N-[Bis(6-pivalamido-2-pyridylmethyl)](N⁰-4-fluorophenylcar-

bamoylmethyl)aminatomanganate(II) ([Mn^IIHbpaa]). A solu-

tion of H₃bpaa (0.147 g, 0.268 mmol) in 5 mL of anhydrous

DMA was treated with solid KH (0.023 g, 0.573 mmol) under an

argon atmosphere. The mixture was stirred until H₂evolution

ceased. Mn(OAc)₂(0.048 g, 0.272 mmol) was added and allowed

to stir an additional 45 min. The mixture was filtered to remove

KOAc (0.046 g, 88%) to afford 0.159 g (99%) of the pale-yellow

product. Mass calcd for C₃₀H₃₄FN₆O₃Mn [M - H]: m/z 600.2.

Found: m/z 600.2. EPR (^ mode, X band, 4 K) g values: 21.0, 5.6,

3.0, 1.66 1.31.

Preparation of [Mn^IIIH₂bpaa(O₂)] via O₂and DPH. A solu-

tion of [MnHbpaa] (0.114 g, 0.190 mmol) and DPH (0.018 mg,

98 mol) in 4 mL of anhydrous DMA was treated with excess dry

O₂and stirred for 10 min. MS (ES^-). Mass calcd for

C₃₀H₃₅FN₆O₃Mn¹⁶O₂[M - H]: m/z 633.2. Found: m/z 633.2.

Mass calcd for C₃₀H₃₅FN₆O₃Mn¹⁸O₂[M - H]: m/z 637.2.

Found: m/z 637.1. After treatment with H₂O (3 mL), the

solution was washed with Et₂O (3 ꢀ 3 mL). Concentration

of the Et₂O solution gave 16 mg (89%) of an orange solid.

Analysis by ¹H NMR spectroscopy indicated that the solid was

azobenzene. Via H₂O₂: A solution of [MnHbpaa] (57 mg,

Acknowledgment. Acknowledgment is made to the NIH

(Grant GM 050781) and the NSF (Grant 0738252) for financial

support. We are thankful for the numerous co-workers

who have contributed to the work presented here from our

laboratory, including C. MacBeth, R. Gupta, T. Parsell, M.-Y.

Yang, R. Lucas, M. Zart, M. Hendrich, W. Gunderson, and

E. Solomon.

Supporting Information Available:

CIF for the X-ray

experiment on [Mn^IIHbpaa]. This material is available free of

charge via the Internet at http://pubs.acs.org.

Article Doi

Role of the secondary coordination sphere in metal-mediated dioxygen activation

DOI: 10.1021/ic901550k

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Article abstract of DOI:10.1021/ic901550k

Full text of DOI:10.1021/ic901550k

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