Metal–thiolate bonds in bioinorganic chemistry

Metal–Thiolate Bonds in Bioinorganic Chemistry

EDWARD I. SOLOMON, SERGE I. GORELSKY, ABHISHEK DEY

Department of Chemistry, Stanford University, 333 Campus Drive, Stanford, California 94305

Received 15 November 2005; Accepted 20 December 2005DOI 10.1002/jcc.20451

Published online in Wiley InterScience (www.interscience.wiley.com).

Abstract: Metal–thiolate active sites play major roles in bioinorganic chemistry. The M��Sthiolate bonds can be

very covalent, and involve different orbital interactions. Spectroscopic features of these active sites (intense, low-

energy charge transfer transitions) reflect the high covalency of the M��Sthiolate bonds. The energy of the metal–thio-

late bond is fairly insensitive to its ionic/covalent and �/� nature as increasing M��S covalency reduces the charge

distribution, hence the ionic term, and these contributions can compensate. Thus, trends observed in stability con-

stants (i.e., the Irving–Williams series) mostly reflect the dominantly ionic contribution to bonding of the innocent

ligand being replaced by the thiolate. Due to high effective nuclear charges of the CuII and FeIII ions, the cupric–

and ferric–thiolate bonds are very covalent, with the former having strong � and the latter having more � character.

For the blue copper site, the high � covalency couples the metal ion into the protein for rapid directional long range

electron transfer. For rubredoxins, because the redox active molecular orbital is � in nature, electron transfer tends

to be more localized in the vicinity of the active site. Although the energy of hydrogen bonding of the protein envi-

ronment to the thiolate ligands tends to be fairly small, H-bonding can significantly affect the covalency of the

metal–thiolate bond and contribute to redox tuning by the protein environment.

q 2006 Wiley Periodicals, Inc. J Comput Chem 27: 1415–1428, 2006

Key words: covalency; XAS; metal–thiolate bonds; spectroscopy; TD-DFT

Introduction

Metal–thiolate bonds are present for many classes of metallopro-

tein active sites and make major contributions to function. The

protein centers involved in electron transfer include the blue cop-

per site (to be discussed in the next section), the mixed valent

binuclear CuA site, and the Fe(SR)4, (Fe2S2)SR4, (Fe4S4)(SR)4,

etc., iron sulfur sites all have thiolate–metal bonds of cysteine res-

idues coupling them into the protein matrix.1 The main heme

enzyme involved in O2 activation is P450, which has an axial thio-

late–Fe bond, while enzymes involved in superoxide reactivity

include superoxide reductase having a nonheme Fe–thiolate bond

and Ni superoxide dismutase having nickel thiolate bonds.2 There

are also classes of enzymes involved in Lewis acid catalysis that

have thiolate–Fe bonds including nitrile hydratase (FeIII also CoIII,

with some thiolates oxidized) and deformylase (FeII) or a thiolate-

Zn bond as in alcohol dehydrogenase.3 Other metalloenzymes

with bridging thiolate ligation include the hydrogenases, a binu-

clear (Fe–Fe or Ni–Fe) cluster involved in the reduction of dihy-

drogen, sulfite reductase, an FeIII porphyrin bridged to a Fe4S4cluster involved in reduction of sulfite, and CO dehydrogenase,

which has a binuclear Ni cluster bridged to a Fe4S4 cluster (A-

cluster) is involved in the assimilation of CO (Fig. 1).1,4

The general feature of all these metalloproteins and enzymes is

that they have unique spectral features (i.e., intense, low-energy

absorption bands and unusual spin Hamiltonian parameters) that

reflect highly covalent thiolate–metal bonding that can make

major contributions to reactivity. The thiolate ligand has three va-

lence 3p orbitals, one of which is greatly stabilized in energy due

to the carbon–sulfur � bond, and thus does not significantly con-

tribute to the thiolate sulfur–metal bond (Fig. 2). The two remain-

ing 3p orbitals, which are perpendicular to the S��C bond, domi-

nate the thiolate interaction with the metal center and split in

energy as the C–S–M angle decreases from 1808. The 3p orbital

out of the C–S–M plane is involved in � bonding, while the in-

plane 3p orbital pseudo � bonds to the metal ion (pseudo � in the

sense that when the C–S–M angle is greater than 908 its electrondensity is shifted off the S–M bond; C–S–M bond angles are gen-

erally in the range of 100–1208). The specific bonding interactions

depend on the metal ion, its 3dn configuration and its ZEff (effec-

tive nuclear charge).

In this review, we will first briefly consider the bonding in the

blue copper site, which has one d hole and a high ZEff; therefore, it

is a particularly covalent site. We then extend these studies over

the series of first row transition metal ions to probe trends in cova-

lent, ionic, � and � bonding. Finally, we will consider FeIII–Sthiolate

Correspondence to: E. I. Solomon; e-mail: [email protected]

;Contract/grant sponsor: NSF; contract/grant number: CHE-0446304.

Contract/grant sponsor: NIH; contract/grant number: GM-2FBM405 (to

E.I.S.).

q 2006 Wiley Periodicals, Inc.

bonding, where the increase in effective nuclear charge for the fer-

ric center again leads to highly covalent bonding and the high-spin

d5 configuration allows both � and � contributions to this bonding.

Here we focus on the fact that the protein can provide hydrogen

bonds to the thiolate, which can change the covalency of this bond

and affect reactivity.

Nature of the Thiolate–Cu Bond in the Blue

Copper Active Site

The details of this bonding description have been described,5 and

only a brief summary of key features is presented here. The crystal

structure of this site was first determined by Hans Freeman in

1978 for plastocyanin.6 It has a distorted tetrahedral geometry

with a short Sthiolate ligand in the x,y plane, which is determined

by this sulfur and two histidine nitrogen ligands all with strong

ligand–metal bonds. There is also a weak axial thioether S ligand

in some blue copper sites with a long bond to the copper (�2.8 A

in plastocyanin7). The Cu–S–C bond angle for the thiolate is 1108,and this dihedral plane is oriented approximately perpendicular to

the x,y plane. A DFT calculation of the ground state of the blue

copper site was first published in 19858 (Fig. 3A), and shows a

dx2�y2 orbital (i.e., oriented in the xy plane) which is highly cova-

lent and the covalency is highly anisotropic being delocalized into

the S 3p� orbital of the thiolate ligand. Modern DFT calculations

give an equivalent description but vary in the metal 3d-thiolate

covalency, defined here as the amount of thiolate character mixed

into the antibonding metal 3d-based molecular orbitals.

A direct experimental probe of the covalency of the thiolate S–

Cu bond is sulfur K-edge X-ray absorption spectroscopy (XAS).9

This involves a transition at 2469 eV (Fig. 3B) from the sulfur 1s

orbital into the half-occupied HOMO (i.e., SOMO or �-spinLUMO, which closely reflects the total spin density). The 1s or-

bital is localized on the sulfur atom and the s ! p transition is the

electric dipole allowed; thus, the intensity of the sulfur 1s ! �LUMO transition directly reflects the sulfur 3p character mixed

into the dx2�y2 orbital due to covalent bonding. The intensity of

this transition for the blue copper site in plastocyanin is very high:

2.5 times the intensity of a model complex with a normal CuII–thi-

olate bond with �15% covalency. [In this model complex, the Cu

atom has a five-coordinate geometry, Cu–Sthiolate distance is

2.36 A, and the Sthiolate p� ! Cu and Sthiolate p� ! Cu LMCT

transitions are at higher energy and much lower intensity—23,500

cm�1 (" ¼ 430 M�1 cm�1) and 27,800 cm�1 (" ¼ 360 M�1 cm�1),

respectively, than those in blue-copper sites. Both the XAS inten-

Figure 1. Biological centers having functionally relevant M-Sthiolate bonds.

1416 Solomon, Gorelsky, and Dey l Vol. 27, No. 12 l Journal of Computational Chemistry

Journal of Computational Chemistry DOI 10.1002/jcc

sity analysis and the SCF X� calculations, calibrated to reproduce

the EPR spin Hamiltonian parameters, give �15% S 3p character

in the LUMO. This result and the error estimates in the S character

from the XAS edges are discussed in detail in ref. 10.] The inten-

sity of the blue copper sulfur K-pre-edge transition quantitates to

38 6 3% Sthiolate character in the ground-state wave function.5,10

The � nature of the ground state in Figure 3A is quite unusual,

as CuII normally utilizes its 1/2 occupied dx2�y2 orbital to form �bonds to donor ligands. This � interaction derived from analysis

of the charge transfer (CT) absorption spectrum using low-temper-

ature magnetic circular dichroism (LT MCD) spectroscopy. From

Figure 3C, the dominant transition in the absorption spectrum of

blue copper proteins is a band at �16,000 cm�1 (" �5000 M�1

cm�1), which is a characteristic unique spectroscopic feature

of the blue copper site. Correlation to the LT MCD spectrum

(Fig. 3D) shows that lower energy weak absorption features are the

most intense in the LT MCD spectrum, while the intense 16,000

cm�1 band has only limited LT MCD intensity. LT MCD intensity

(known as a C-Term) derives from spin-orbit coupling and reflects

the metal character (which dominates the spin-orbit interaction) in

the excited states. Thus, the low-energy weak absorption bands

that are intense in LT MCD can be assigned as d ! d transitions,

while the 16,000 cm�1 band is the lowest energy; hence, � CT

transition, while a higher energy weak feature is the Sthiolatepseudo � to CuII CT transition. The intense-�/weak-� CT absorp-

tion pattern is inverted from the normal behavior observed for CT

transitions in copper complexes. As absorption band intensity

reflects the overlap of the donor and accept orbitals involved in

the electronic transition, this requires that the dx2�y2 orbital be ori-

ented so that it � interacts with the thiolate sulfur. This � bond is

due to the short thiolate S–CuII bond length of 2.1 A.7 This short

Cu–S bond originates from the fact that the Cu–SMet is very weak,

and there are only three strong donor ligands (one SCys and two

NHis) oriented in the xy plane of the copper coordination sphere.

The high covalency of this bond activates specific superexchange

pathways through the protein for rapid directional electron transfer

as discussed in ref. 5.

Ionic and Covalent, r and p Contributions

to MII–Thiolate Bonds

A series of metal-varied model complexes [MIIL(SC6F5)] (where

L ¼ HB(3,5-iPr2pz)3� and MII ¼ Mn, Fe, Co, Ni, Cu, and Zn),

related to blue copper sites in proteins, has been synthesized and

crystallographically characterized.11,12 The metal atoms in these

complexes also have a distorted tetrahedral coordination sphere,

with one M–S bond, two equatorial M–N bonds, and an elongated

axial M–N bond. The spectroscopic features of [MIIL(SC6F5)] are

similar to those of the corresponding metal-substituted sites in the

proteins.13–20 The complexes show interesting systematic changes

in their metal–thiolate bond lengths, which follow the order: MnII

> FeII > CoII > NiII > CuII < ZnII. This parallels the Irving–Wil-

liams series for the stability constants.21,22

Spectroscopic Signatures of p and s Metal–Thiolate

Bond Covalency

A combination of absorption, MCD, and resonance Raman (rR)

spectroscopies was used23 to distinguish the ligand field (LF) tran-

Figure 2. Sulfur-based valence orbitals of methyl thiolate.

Figure 3. (A) Structure and the LUMO of the blue copper center;

(B) S K-edge XAS of plastocyanin (PCu) and the [Cu(tetb)(SC6H4CO2) complex; (C) low-temperature absorption and (D)

MDC spectra of poplar plastocyanin.

1417Metal–Thiolate bonds in Bioinorganic Chemistry


sitions from the charge transfer (CT) transitions. The latter provide

insights into covalent interactions at the metal center. Figure 4

compares the electronic absorption spectra of the [MIIL(SC6F5)]

complexes. The experimental electronic spectra of the series are

well reproduced by time-dependent density functional theory (TD-

DFT)24,25 at the B3LYP/6-311þG* level (Fig. 5) and confirm the

assignments of the bands.

In contrast to the ZnII complex, which has the d10 metal ion

configuration and does not show absorption in the visible region,

the absorption spectrum of the CuII complex (Figs. 4 and 5B) has

an intense absorption band at �15,000 cm�1, due to the Sthiolatep� ! Cu dx2�y2 (�-spin HOMO ! LUMO) CT transition. As

described earlier, this is also characteristic of the blue copper pro-

teins having the highly covalent Sthiolate p�–Cu interaction in the

ground state.5,26,27 The absorption spectrum of the NiII complex

(Figs. 4 and 5C) exhibits a dominant feature at �20,000 cm�1,

which is due to the Sthiolate p� ! Ni dx2�y2 LMCT transition. This

transition is higher in energy (from 15,000 cm�1 to 20,000 cm�1)

and lower in intensity (from 12,000 M�1 cm�1 to 8000 M�1

cm�1) compared to the corresponding transition in the CuII com-

plex (Figs. 4 and 5B). In the higher energy region (>25,000 cm�1)

there are the Sthiolate p� ! Ni dx2�y2, benzyl based intraligand CT

transitions, the Sthiolate p� ! Ni dz2, and the (pyrazolyl) p� ! Ni

dx2�y2 transitions.

In the absorption spectrum of [CoIIL(SC6F5)] (Figs. 4 and 5D),

The Sthiolate p� ! Co dx2�y2 charge transfer transition is now at

24,000 cm�1 (" �3000 M�1 cm�1); this transition is higher in

energy and lower in intensity compared to the corresponding band

in the NiII complex (20,000 cm�1 and " � 8000 M�1 cm�1) as

observed experimentally. The interesting feature in the CoII com-

plex is an additional intense (" � 11,000 M�1 cm�1) Sthiolate p� !Co dxy charge transfer transition at 30,000 cm�1. In CoII, the dxyorbital is unoccupied and interacts with the Sthiolate p� orbital, pro-

ducing an intense thiolate-to-metal CT absorption band at 30,000

cm�1. This reverses the pattern of more intense Sthiolate p� and less

intense Sthiolate p� ! MII dx2�y2 CT transitions observed in the Cu

and Ni complexes, to less intense Sthiolate p� ! MII dx2�y2 and

more intense Sthiolate p� ! MII dxy CT transitions. This behavior

explains the spectral differences in Figure 4, where for the Co

complex in solution the absorption intensity at �30,000 cm�1 due

to the Sthiolate p� LMCT transition is more intense than the transi-

tion at �25,000 cm�1 due to the Sthiolate p� LMCT transition.

The TD-DFT calculated spectrum for the FeII complex (Fig. 5E)

shows low intensity ligand field transitions in the 5000–20,000

cm�1 region. The Sthiolate p�! Fe dxy LMCT excitation appears as

the intense absorption (" � 8000 M�1 cm�1) at 32,000 cm�1, com-

pared to the less intense (" � 2200 M�1 cm�1) Sthiolate p� ! Fe

dx2�y2 CT excitation contributing to the absorption intensity at �28,000 cm�1. The Sthiolate p� ! Fe dxy LMCT transition is higher

in energy (by 2000 cm�1) and lower intensity (by 3000 M�1 cm�1),

and the Sthiolate p� ! Fe dx2�y2 LMCT transition is also higher in

energy (by 4000 cm�1) and lower in intensity (by 800 M�1 cm�1)

relative to the corresponding excitations in the Co complex.

Figure 4. UV-Vis absorption spectra of [MIIL(SC6F5)] in cyclohex-

ane at room temperature.

Figure 5. Simulated electronic spectra of [MIIL(SC6F5)] (TD-DFT)

calculations at the B3LYP/6-311þG(d) level. The calculated ener-

gies and intensities of electronic transitions were transformed into

simulated spectra as described in ref. 23 using Gaussian functions

with half-widths of 2500 cm�1. The contributions from individual

electronic transitions are shown in gray.



The [MnIIL(SC6F5)] complex does not have absorption inten-

sity in the 5000–25,000 cm�1 region (Figs. 4 and 5F). This is

because the LF transitions in high spin Mn(d5) complexes are

both spin and Laporte forbidden.28 The absorption intensity at

�35,000 cm�1 is due to Sthiolate p� ! Mn dx2�y2 LMCT transi-

tion, and at 38,000 cm�1 to the Sthiolate p� ! Mn dxy LMCT

transition.

Thus, in both the experimental and calculated spectra, the

interesting trends observed are: (1) in going from Cu ! Ni ! Co,

the Sthiolate p� ! M dx2�y2 transitions shift to higher energy and

decrease in intensity, and (2) a new intense Sthiolate p� ! M dxyband appears in going from Ni to Co, which shifts to higher

energy and decreases in intensity in going from Co to Fe to Mn.

To obtain insight into the variation in the observed spectral fea-

tures arising due to the different metal–thiolate interactions in

these complexes, trends in the metal–thiolate bonding were eval-

uated using DFT.

Molecular Orbital Description of Metal–Thiolate s and pBonding and Ligand-to-Metal Charge Donation

Figure 6 shows MO diagrams for the [MIIL(SC6F5)] complexes,

with the metal 3d orbitals and the thiolate ligand orbitals labeled.

These MOs are relevant for chemical bonding between the metal

atom and the thiolate ligand and define the optical spectra. The

extended charge decomposition analysis (ECDA),29,30 which is an

extension of the charge decomposition analysis,31 indicates that

the metal–thiolate covalent interactions in these series are domi-

nated by charge donation from the SC6F5� ligand to the MLþ frag-

ment (Table 1). This charge donation has both �- and �-orbitalcomponents, and the two highest occupied molecular orbitals of

the thiolate ligand, HOMO (S p�) and HOMO-1 (S p�), are princi-pal donor orbitals for the M��S bond.

Due to strong overlap between metal 4s and Sthiolate p� fragment

orbitals, the � component of the M��S covalent bond is substantial

(Tables 1 and 2) and its strength remains approximately the same

over the series. This is reflected by the observation that the �-spinoccupied molecular orbitals of each [MIIL(SC6F5)] complex contain

26–32% of the unoccupied molecular orbitals of the MLþ fragment

(Table 1) or, alternatively, the natural population analysis (NPA)32

derived occupation of the metal 4s orbital is 0.14–0.23 electrons

(Table 2). Taking electronic polarization of the MLþ fragment into

account (Table 1),29 compositions indicate the net � charge dona-

tion of �0.5 e� from SC6F5� to the MLþ fragment.

In the MnII complex, all the �-spin 3d orbitals of the central

atom are unoccupied.

The MnII 3d–Sthiolate orbital interactions are weak and contrib-

ute little to the covalent bonding relative the stronger MnII 4s,4p–

Sthiolate orbital interaction. The Mn 3d orbital character of the thio-

late-based HOMO(�) and HOMO-3(�) (the thiolate character is

95 and 86%, respectively) is 3 and 5%, respectively, and recipro-

cally the Sthiolate contribution is 2% for the Mn dx2�y2 (d�) and 3%

for dxy (d�) orbitals (Fig. 6).In the FeII complex with one more valence electron (which

occupies the �-spin dyz�xz orbital) than the MnII complex, the �-HOMO-LUMO gap becomes smaller, 3.7 eV (in comparison to

the 4.4 eV gap in the MnII complex), reflecting the greater Zeffnuc forFe relative to Mn. The increased Zeffnuc shifts the energies of metal-

based orbitals closer to the occupied thiolate ligand orbitals and

increases the covalent mixing between the metal and the thiolate

(the Fe 3d characters in the Sthiolate p� and p� orbitals are 5 and

11%, respectively; Fig. 6).

In [CoIIL(SC6F5)], the next metal d orbital (dyzþxz) becomes

occupied. Because Zeffnuc for Co is greater than that of Fe, the �-HOMO-LUMO gap again becomes smaller (3.2 eV) and the cova-

lency of the M��Sthiolate bond is further increased (the Co 3d char-

acter in Sthiolate p� is 9%, Fig. 6). The important contribution to �covalency comes from HOMO-5, which is a bonding combination

of the Sthiolate p� orbital and Co dxy. This agrees with the spectro-

scopic data (Figs. 4 and 5), which show an intense the Sthiolatep�! Co dxy CT transition.

In [NiIIL(SC6F5)], the dxy orbital is now occupied. This can-

cels the Sthiolate p�–M dxy bonding component for the covalent

M��S bond, eliminates the Sthiolate p� ! M dxy CT band (Fig. 5),

and, as will be discussed later, the M��Sthiolate bond order de-

creases. Because the dxy orbital is now below the Sthiolate p�orbital (Fig. 6), the latter is destabilized by �1 eV. The greater

Zeffnuc for Ni results in a further lowering of the energies of M 3d-

based orbitals and favors more efficient covalent bonding with

the Sthiolate p� orbital. As a result, the LUMO (dx2�y2) has 8% S

contribution (Fig. 6).

In the CuII complex, the dz2 orbital is populated and the Cu

dx2�y2 is the only unoccupied d orbital. As a result of the still

larger Zeffnuc, the �-HOMO-LUMO gap for [CuL(SC6F5)] is lowest

Figure 6. �-Spin frontier molecular orbitals of the [MIIL(SC6F5)]

complexes (MOs with a0 and a@ symmetry are shown in gray and

black, respectively) from B3LYP/TZVP calculations. From the Ni

and Cu complexes, both S p� amd dxy orbitals are occupied and, as

a result, their mixing cannot contribute to the M-S covelency. Thus,

the M 3d character in the S p� orbital and sulfer character in the M

dxy orbital are not shown.



(2.4 eV) maximizing the orbital interaction between the Sthiolate p�and Cu dx2�y2 (Fig. 6). These two fragment orbitals mix to form

the bonding �-spin HOMO [49% of HOMO(SC6F5�) and 41% �-

LUMO(Cu dx2�y2 of MLþ)] and the antibonding �-LUMO [39%

of HOMO(SC6F5�) and 44% �-LUMO(MLþ)]. Alternatively, the

Sthiolate character of the �-LUMO is 27% and the Cu 3d character

of the �-HOMO is 21%. Thus, the �-HOMO/LUMO compositions

in the complex indicate the most covalent M��Sthiolate bond in the

series and largest net charge donation from the thiolate to the

metal (Table 1). This agrees with the spectroscopic data (Figs. 4

and 5), which show an intense the Sthiolate p� ! Cu dx2�y2 CT

transition. In addition, ECDA29 (Table 1) shows that the electronic

polarization of the MLþ fragment by the thiolate ligand is the larg-

est for the CuII complex, 16 orbital%, in agreement with the small-

est HOMO-LUMO gap of the CuLþ fragment in the series. This

polarization provides additional stabilization to the highly cova-

lent CuII–thiolate bond.

In the ZnII complex, all the d orbitals of the metal ion are occu-

pied, including the Zn dx2�y2–Sthiolate p� orbital. Thus, the � con-

tribution to covalency is cancelled, the covalent bonding interac-

tion between ZnII 3d orbitals and the thiolate orbitals becomes

very small, and the charge donation from the thiolate to the metal

is the lowest in the series (Table 1). The remaining covalent inter-

action comes mostly from the ZnII 4s orbital.

Table 1. Extended Charge Decomposition Analysis29,30 of the Metal–thiolate Bonding in the [MIIL(SC6F5)]

Complexes: Thiolate HOMO(�) and HOMO-1(�) Contributions to the Unoccupied MOs of [MIIL(SC6F5)],

Net Contributions (%) of Unoccupied Fragment Molecular Orbitals (UFOs) to the Occupied MO of

[MIIL(SC6F5)], the Difference in Electronic Polarization (PL) between the MLþ and SC6F5� Fragments, and

the Net Charge Donation from the SC6F5� Thiolate to the MLþ Fragment (B3LYP/TZVP Calculations).

Metal Mn Fe Co Ni Cu Zn

% �-spin HOFO(SC6F5�) �a 2.9 3.1 3.1 2.9 3.2 3.5

% �-spin HOFO(SC6F5�) �a 6.8 10.2 9.2 14.1 41.1 3.5

% �-spin HOFO-1(SC6F5�) �b 17.1 16.8 16.2 16.8 16.0 20.8

% �-spin HOFO-1(SC6F5�) �b 17.5 19.4 24.0 18.2 16.2 20.8

donation(SC6F5� ! MLþ) and polarization(MLþ)

% �-spin UFOs(MLþ) 29.5 29.6 28.2 28.2 26.2 32.5

% �-spin UFOs(MLþ) 32.0 33.6 48.9 47.3 81.7 32.5

donation(SC6F5� / MLþ) and polarization(SC6F5

�)

% �-spin UFOs(SC6F5�) 3.1 3.5 3.5 3.4 4.3 2.7

% �-spin UFOs(SC6F5�) 2.7 4.8 4.2 3.7 4.0 2.7

PL�(MLþ) - PL�(SC6F5

�)c 2.4 2.1 1.3 1.1 �0.9 1.3

PL�(MLþ) - PL�(SC6F5

�)c �0.2 5.0 7.4 6.2 16.0 1.3

donation�(SC6F5� ! MLþ)d 0.240 0.240 0.234 0.237 0.229 0.285

donation�(SC6F5� ! MLþ)d 0.295 0.335 0.372 0.375 0.617 0.285

donation�þ�(SC6F5

� ! MLþ) 0.54 0.58 0.61 0.61 0.85 0.57

aThe principal component of � donation(SC6F5� ! MLþ): the HOMO(SC6F5

�) contribution to the unoccupied MOs of the complex.bThe main component of � donation(SC6F5

� ! MLþ): the HOMO-1(SC6F5�) contribution to the unoccupied MOs of

the complex.cThe difference (orbital %) in electronic polarization (PL) between the MLþ and SC6F5

� fragments for �- and �-spin

orbitals, respectively.dThe net charge donation from the SC6F5

� thiolate to the MLþ fragment for �- and �-spin orbitals, respectively. The

former is largely � donation and the latter is both � and � donation.

Table 2. Metal and Sulfur NPA32-Derived Charges and Metal 3d, 4s, and 4p Atomic Orbital

NPA Populations (�-spin)a in the [MIIL(SC6F5)] (B3LYP/TZVP Calculations).

Metal Mn Fe Co Ni Cu Zn

qNPA(M) 1.37 1.30 1.27 1.24 1.13 1.51

qNPA(S) �0.36 �0.32 �0.31 �0.32 �0.19 �0.40

�-spin 3d(M) populationa 0.35 1.39 2.40 3.40 4.49 4.99

�-spin 4s(M) populationa 0.14 0.15 0.16 0.18 0.19 0.23

�-spin 4p(M) populationa 0.02 0.02 0.01 0.01 0.01 0.01

aNPA populations of metal 3d, 4s, and 4p �-spin orbitals are 4.92–4.99, 0.18–0.23, 0.01–0.03 electrons, respec-

tively.



Ionic and Covalent Contributions to the Metal–

Thiolate Bonds

The B3LYP calculations correlate well with the spectroscopic

data, and can be used evaluate the trends in M��S bond lengths,

energies, and force constants (Fig. 7), and correlate these to the

nature of bonding between the metal ion and the thiolate ligand.23

Overall, the B3LYP calculations reproduce experimentally

observed structural changes in the series (Fig. 7A). The nonmono-

tonic variations in the M��S bond lengths mark important changes

in the metal–thiolate bonding, depending on the electronic config-

uration of the metal ion. The MLþ–SC6F5� binding energy was

calculated as the interaction energy between the metal–pyrazolyl

fragment and the thiolate ligand:

MLþ þ SC6F�5 ¼ ½MIILðSC6F5Þ�:

In the [MIIL(SC6F5)] series, it ranges (Fig. 7B) from �124.7 kcal

mol�1 (Fe) to �132.7 kcal mol�1 (Zn), and its variation does not

mirror the changes in the M��S bond lengths.23 M��S force con-

stant is the lowest for the MnII complex (1.13 mDyne A�1),

increases from MnII to FeII to CoII, where it reaches a local maxi-

mum, then decreases for the NiII complex, and increases for the

CuII complex (the maximum value, 1.43 mDyne A�1), and then

decreases again for ZnII. The force constant (Fig. 7C) shows

greater variation (26%) than the bond energy (6%). Its variation is

more consistent with the changes in the M��S bond lengths. The

lack of correlation between the bond lengths and energies, and the

force constants implies that there are varying contributions to the

metal–thiolate bonding in this series, and a more detailed exami-

nation of the metal–thiolate binding energies and their ionic and

covalent components is warranted.

The MLþ–SC6F5� binding energy, Eo, can be partitioned into

several contributions. First, Eo is separated into two components

Eprep and Eint:

E0 ¼ Eint þ Eprep

Eprep is the preparation (deformation) energy23,29 necessary to

transform the MLþ and SC6F5� fragments from their equilibrium

geometries and electronic ground states to the those in the com-

plexes:

Eprep ¼ EprepðMLþÞ þ EprepðSC6F�5 Þ

Over this series, Eprep(MLþ) and Eprep(SC6F5�) was 5.3–10.3

kcal mol�1 and 1.0–1.2 kcal mol�1, respectively. Eint is the inter-

action energy between the MLþ and SC6F5� deformed fragments.

This interaction energy (Fig. 7B) can be further divided into two

major components that can be interpreted in a physically meaning-

ful way:

Eint ¼ Ecov þ Eionic:

Here, Ecov is the covalent or orbital interaction energy (includ-

ing the exchange repulsion energy33–35) and Eionic is the electro-

static interaction energy. The latter is estimated as a sum of elec-

trostatic interactions between atomic charges, qNPA, from the two

molecular fragments, MLþ and SC6F5�. The ESP-derived36

atomic charges could be also used to evaluate the electrostatic

contributions to bonding. For an example, the Merz–Singh–Koll-

man ESP-derived atomic charges of noninteracting ZnLþ and

SC6F5� fragments were similar to the NPA charges and, as a result,

the ionic interaction energy calculated from the ESP fragment

charges (�102.6 kcal mol�1) is close to the NPA-derived value,

�104.5 kcal mol�1. However, the ESP charges in the

[Zn(L)SC6F5] complex were different from both the NPA values

and charges in the noninteracting fragments. This underlines the

problem of finding ESP charges in large molecules using the

standard electrostatic potential fitting algorithms.

Eionic ¼X

a2ML

X

b2 SC6F5

qNPAa qNPAb

rabðin atomic unitsÞ

The charge distribution in this calculation corresponds to the

one in the complex, as opposed to the electrostatic interaction

energy from the energy decomposition analysis of Kitaura–Moro-

kuma33,34,37 and Ziegler,38 which is calculated with undistortedcharge distributions corresponding to those in the isolated frag-

ments. (If the undistorted charge distributions were used in the

analysis, the ionic contribution would show very little variation

with the metal because the reference state in the energy decompo-

sition is the noninteracting MLþ and SC6F5� fragments.)

Because the metal–pyrazolyl fragment and the thiolate ligand

carry opposite charges, the electrostatic interaction is attractive

Figure 7. (A) Experimental and calculated (at the B3LYP/6-

311þG* level) M-S bond lengths in the [MIIL(SC6F5)] complexes;

(B) calculated binding energies at the B3LYP/6-311þG(3df) level,

Eo, and interaction energies, Eint, between the MLþ and SC6F5�

fragments; and (C) calculated M-S force constants.



(Eionic < 0) in the [MIIL(SC6F5)] complexes and related to the

atomic charges of the metal and the sulfur of the thiolate ligand

(Table 2). It can be seen that the absolute values of metal and sul-

fur charges are highest in the MnII and ZnII complexes and lowest

in the CuII complex. Thus, it can be expected that the ionic contri-

bution to the M��S bond energy is lowest in the CuII complex and

highest in the MnII and ZnII complexes. Indeed, as will be dis-

cussed later in this section, the ionic bonding plays an important

role in determining the overall strength of the MLþ–SC6F5� inter-

action.

The analysis of covalent contributions to the chemical bonding

can be done using Mayer bond orders,39 BAB, and its �- and �-spinorbital components, BAB

� and BAB� .23

BAB ¼X

a2A

X

b2B

�ðPSÞbaðPSÞab þ ðPsSÞbaðPsSÞab� ¼ B�

AB þ B�AB;

B�AB ¼ 2

X

a2A

X

b2BðP�SÞbaðP�SÞab;

B�AB ¼ 2

X

a2A

X

b2BðP�SÞbaðP�SÞab;

where P and Ps are the density and spin-density matrices, respec-

tively (P ¼ P� þ P�; Ps ¼ P� � P�), P� and P� are �- and �-spinelectron density matrices, and S is the overlap matrix.

From the bond order analysis, the metal-sulfur bonding is a

dominant interaction between the MLþ fragment and the thiolate

ligand (Fig. 8A). The M��S bond order (and the total ML–SC6F5bond order) in the [MIIL(SC6F5)] complexes follows the Mn < Fe

< Co > Ni < Cu > Zn progression found in the M��S force con-

stants (Fig. 7C) and the charge transfer from the SC6F5� ligand to

the MLþ fragment (Table 1).

To understand this trend and to calculate the �- and �-bondcontributions, symmetry-adapted orbitals were used in the bond-

order analysis:23

BAB ¼X

�

BABð�Þ;

where BAB (G) is the bond order contribution from orbitals with ir-

reducible representation G.The �-spin LUMO of the MLþ fragment (this orbital has

�80% M 4s character) remains relatively unperturbed (in its

energy and composition) by the nature of the central atom. As a

consequence, the mixing between the �-spin LUMO of the MLþ

fragment and the HOMO-1 of the SC6F5� fragment and the result-

ing � bond contributions to the Mayer bond orders from the �-spinMOs, B�

M��S (�), remain similar for all complexes in this series.

Because all �-spin metal 3d-based orbitals are occupied and can-

not contribute to covalent bonding with the thiolate donor, the �bond component from �-spin MOs, B�

M��S (�), is close to zero

(Fig. 8A).

Thus, B�M��S (�) and B�

M��S (�) provide a reference for changesin �-spin MO � and � contributions to the bond orders (Fig. 8A).

These changes result from an additional small � bond and larger �bond contributions to the covalent bonding due to M dxy–Sthiolatep� and M dx2�y2–Sthiolate p� orbital interactions, respectively. The

increasing Zeffnuc (from MnII to ZnII) brings the energies of the dxyand dx2�y2 orbitals of the metal atom closer to the thiolate occupied

orbitals and makes the M 3d-thiolate covalency stronger, given that

the electron occupancy of the appropriate MOs allows the net con-

tribution to be positive. This is the case for the � component,

B�M��S (�), in going from MnII to FeII to CoII, and for the � compo-

nent, B�M��S (�), when proceeding from MnII to CuII. These

changes in �- and �-components of the covalent bonding between

the metal and the thiolate produce the observed net bond order

trends with a local maximum at CoII and a global maximum at CuII.

Following this analysis of the covalent bonding in [MIIL(SC6F5)],

it is possible to explain their observed spectral features (Figs. 4–5).

The shift to higher energy and the decrease in intensity of the

strong Sthiolate p�! M dx2�y2 (M ¼ Cu, Ni, Co, and Fe) CT transi-

tion is due to reduced M 3d–thiolate covalency, supported by

the decrease in the pre-edge intensity in the S K-edge XAS data23

Figure 8. (A) The M-Sthiolate bond order (open square) and its �-

and �-components (circles and diamonds, respeceively) and the

bond order between the MLþ and SC6F5� fragments (solid squares)

in [MIIL(SC6F5)] from B3LYP/TZVP calculations. (B) The NPA-

derived net charge transfer from the SC6F5� ligand to the MLþ fra-

ments. (C) The ionic component of the MLþ–SC6F5� bonding

energy, Eionic, and the difference, Eint � Eionic.



and in MO calculations. An important feature is observed in the

Co complex, the Sthiolate p� ! Co dxy charge transfer transition is

now present and more intense relative to the Sthiolate p� ! Co

dx2�y2 charge transfer transition. This is consistent with its ground

state electronic structure description, indicating the strong interac-

tion between Co dxy and the Sthiolate p� orbitals. Similar to the

trend in the Sthiolate p� ! M dx2�y2 CT transition, the Sthiolatep� ! M dxy CT transition shifts to higher energy and decreases in

intensity in going from CoII to FeII to MnII.

If the metal–thiolate bonding in [MIIL(SC6F5)] were limited to

only covalent interactions, the metal–thiolate bond orders would

directly correlate with the metal–thiolate interaction energies and,

would show maxima of the metal–thiolate bond energies at CoII

and CuII (note the correlation between the M��S bond orders and

Eint – Eionic; Fig. 8). However, although the CoII complex shows

the second largest metal–thiolate binding energy (Fig. 7B), the

absolute maximum is observed for the ZnII complex, not for the

CuII complex. This indicates that the ionic contribution, along

with the covalent component, plays an important role in determin-

ing the overall strength of the metal–thiolate interaction (Fig. 8C):

�44% ionic and 56% covalent in MnII, becoming more and more

covalent (with maxima at CoII for �-type bonding and at CuII for

� type), and back to �50% ionic and 50% covalent in ZnII.

The higher effective nuclear charge of the metal atom in going

from Mn to Cu favors covalent metal–thiolate bonding but it also

causes a decrease in metal ionic radii. The latter would cause a

gradual increase in the ionic component of the MLþ–SC6F5� bond

energy in going from MnII to ZnII if the M��S covalency were

unperturbed. However, this effect is opposed by the changes in the

M��S covalency, which result in charge donation from the thiolate

ligand to the MLþ fragment (Table 1) and cause the metal charge

to become less positive and the sulfur charge to become less nega-

tive (Table 1). As a result, the ionic component of the M��S bond

energy decreases in going from MnII to CuII (Fig. 8C).

The metal–thiolate force constant is very similar in

[CuL(SC6F5)] and [ZnL(SC6F5)]. The Zn complex shows a longer

M��S bond distance than the Cu complex; however, the calculated

metal–thiolate interaction energy in the Zn complex is 2 kcal

mol�1 stronger than in for the Cu complex with the largest M��S

covalency. For the Zn complex, the M 3d–thiolate covalency is

lost (all 3d orbitals are occupied) and the remaining M 4s–thiolate

covalency contributes �50% to the metal–thiolate interaction

energy. This large ionic contribution compensates for the lost

M��S covalency and results in the strongest metal–thiolate bond

in this series (Fig. 7B). The lack of a correlation between bond

strength and length (Fig. 7) for the Cu vs. Zn complex reflects the

differences in distance dependence of the ionic vs. covalent contri-

butions to bonding. In such a case, the general correlation of the

bond energy, length, and the force constant does not hold.

Irving–Williams Series

Interestingly, the metal–thiolate binding energies in [MIIL(SC6F5)]

do not show the trend expected for the Irving–Williams series,

which indicate that, if the successive stability constants of com-

plexes of divalent ions of the first transition series are plotted

against the atomic number of the element, there is a monotonic

increase to a maximum at Cu irrespective of the nature of the

ligand.21,22 The calculated MLþ–SC6F5� binding energy varies

from �126.8 kcal mol�1 (Mn) to �130.8 kcal mol�1 (Cu), to

�132.7 kcal mol�1 (Zn). As we discussed above, this is a result of

the compensating effect in the metal–thiolate complexes where

the covalent contribution to bonding is comparable in magnitude

with the ionic contribution to bonding. To reconcile this fact

with the experimentally observed Irving–Williams series, we

have also analyzed the metal–ligand bonding in a series of

[MII(HB(pz)3)(F)] complexes where the fluoride models the fourth

ligand being replaced by the thiolate.23 Our analysis indicates that

the metal–ligand bonding in these complexes is dominated by the

ionic contribution and the metal–fluoride binding energy (calcu-

lated at the B3LYP/6-311þG(3df) level) decreases from �176.4

kcal mol�1 (Mn) to �166.1 kcal mol�1 (Cu) and then increases to

�176.6 kcal mol�1 (Zn). This variation in the binding energy is a

result of the increasing metal–ligand covalency in going from Mn

to Cu, which reduces the ionic interaction. However, because of

the dominant ionic contribution to bonding in the [MIIL(F)] series,

the compensating effect of the increasing covalent bonding contri-

bution is not as large as in the [MIIL(SC6F5)] series. As a result,

the difference between the relative formation energies of the

metal–thiolate complexes and the metal–fluoride complexes

�Ef ¼ ½EoðMLþ � SC6F�5 Þ � EoðMnLþ � SC6F

�5 Þ�

� ½EoðMLþ � F�Þ � EoðMnLþ � F�Þ�;

shows the Mn � Fe > Co > Ni > Cu < Zn progression (Fig. 8),

which is the Irving–Williams series. Moreover, stability constants

of the metal–thiolate complexes calculated from the binding

energy differences are consistent with the experimentally observed

quantities.21,22

Thus, the ‘‘softer’’ thiolate ligand can have comparable cova-

lent and ionic contributions to bonding and these compensate to

produce little change in binding energy over the series of metal

ions (open squares in Fig. 9). For the ‘‘harder’’ ligands (F�, OH�,

Figure 9. The relative formation energies (kcal mol�1) of the

metal–thiolate complexes, [MIIL(SC6F5)], and the metal–fluoride

complexes, [ML(F)], and their difference, DEf, calculated at the

B3LYP-6-311þG(3df) level.



H2O, etc.), the ionic term dominates and their binding energies are

affected by changes in covalency over the series. It is the competi-

tion between these behaviors that produces the Irving–Williams

series (solid circles in Fig. 9) in stability constants.

Nature of Fe–Thiolate Bonds

S K-edge XAS has been used to study the bonding in a series of

divalent first row transition metal tetra-thiolates, and the results

obtained are consistent with the general trend of bonding dis-

cussed above. Although CuII–Sthiolate exhibits a low energy intense

pre-edge feature, FeII–Sthiolate has its pre-edge transition at a high

energy due to its lower ZEff, and this overlaps the rising edge tran-

sition (Fig. 10, doted lines).

The S K-edge XAS of the ferric thiolate complex [FeIII(SPh)4]�

(Fig. 10) has well-resolved low-energy intense pre-edges at �2470

eV.41 Using an effective S4 site symmetry, its pre-edge was fit

using four peaks corresponding one-electron transitions to the low

symmetry (Td ! S4) split e(�)5A þ 5B (in red in Fig. 10 inset) and

t2(�)5B þ 5E (in blue in Fig. 10 inset) d6 excited states. (The peak

splittings were fixed using an energy diagram determined from

other spectroscopic techniques.41 The splittings were reduced by

20% to account for the fact that the final state has a reduced d6 con-

figuration.) The fitted intensity gives a total hole covalency of

170% (summed over the five unoccupied Fe3d orbitals), which cor-

responds to the contributions from the thiolate S � and pseudo-�from the four thiolates (�43% per Fe��S bond). Spectroscopically

calibrated SCF-X�-SW calculations gave 140% total covalency in

reasonable agreement with the experiment.40 The maximum �covalency was estimated to be 30% of the total observed cova-

lency, which is consistent with the weak � and strong � charge

transfer transitions observed in the absorption spectrum for this

complex.40 The redox active orbital in dz2, from DFT calculation,

is metal based and has very little (6%) S character (as indicated by

the S K-edge data and a weak � CT band in this complex).42 In

contrast to the blue copper protein, which has extensive � delocali-

zation enabling facile long range ET, the weak covalency of the re-

dox active orbital in the rubredoxin site limits the electronic cou-

pling of the metal into the protein and the ET process is localized

in the vicinity of the active site as discussed in ref. 42.

Effects of the Protein Environment

on FeIII–Thiolate Bonding

The reduction potential of the iron–tetrathiolate site in rubredoxin

is �1 V more positive than structurally similar model complexes.

Hence, it is important to understand the contributions to this shift

in redox potential of a protein active site that is key to its reactiv-

ity. Experimental and computational results suggest that H-bonds,

protein dielectric effects, solvent accessibility, and surrounding

peptide dipoles can make significant contributions to the redox

potentials.43 S K-edge XAS has proved to be a powerful probe of

H-bonding to sulfur ligands as this method directly probes the

covalency of the Fe��S bond.

The S K-edge XAS spectra of rubredoxins from three different

organisms, Clostridium pasteurianum (Cp) , Pyrococcus furiosus(Pf), and a mutant having half the sequence from each of the

above two proteins (Cp|Pf), are shown in Figure 11, along with

data for the model complex [Fe(S2-o-xyl)2]�.41 The proteins dis-

play an intense pre-edge feature at approximately the same energy

as the model complex. However, the intensity of this transition in

the proteins is lower than that of the model complex, and this

decrease in intensity varies somewhat in magnitude for the differ-

ent proteins. This reduction in pre-edge intensity implies a reduc-

tion in Fe��Sthiolate bond covalency, which was quantified by fits

to the experimental spectra. The total Fe3d hole covalency for the

four FeIII–Sthiolate bonds is 125–130% (�32% S3p per bond) in the

proteins compared to 170% in the model. The protein active sites

have 6 backbone N��H---Scys H-bonds which, from the pre-edge

intensity, reduces the covalency of these sites relative to the model

complexes.41

Figure 11. S K-edge XAS spectra of oxidized rubredoxin proteins

Cp (dashed-dot) Pf (gray-dashed), Cp/Pf (dotted), and the model

[FeII(S2-o-xyl)2]� (bold).

Figure 10. Sulfur K-edge XAS spectra of the [FEIII/II(SPh)4]�/2�

rubredoxin models (solid/dashed lines). The inset shows the pre-

edge region of the ferric complex and includes a representative

four-peak fit performed based on the d-orbital splitting diagram of

Gebbhard et al.40 The lower energy peaks represent � contributions

while the higher energy peaks represent the � contributions the the

pre-edge intensity.



The amount of charge donation from the ligands to the metal,

that is, covalency of the metal–ligand bond, can make a very sig-

nificant contribution to the redox potential particularly in covalent

systems like rubredoxin. Increased covalency preferentially stabil-

izes the higher oxidation state reducing the reduction potential. In

the presence of H-bonds to the sulfur, the charge donation of the

ligand to the metal should decrease and the redox potential

becomes more positive. Although the S K-edge XAS data on the

Rubredoxin proteins clearly demonstrate the effect of H-bonding

in reducing the covalency in the protein active site, a more system-

atic evaluation of this effect was required.

H-bonding in a Series of FeIII–Sthiolate Heme

Model Complexes

The effect of H-bonding on an Fe–Sthiolate bond was quantitatively

estimated in series of high-spin FeIII pophyrin (P) model com-

plexes FeP(SPh), FeP(SL1), and FeP(SL2) (where L1 ¼ 2-trifluor-

oacetamido benzene thiol and L2 ¼ 2,5-bis-trifluoroacetamido

benzene thiol; Fig. 12), where the number of H-bonds was

increased along the series (0 ! 1 ! 2) and a systematic variation

in Fe��S bond lengths and reduction potentials was observed.44

The S K-edge XAS of the three model complexes [FeP(SPh),

in bold], [FeP(SL1), 1H-bond in dots], and [FeP(SL2), 2 H-Bonds

in dashed], are given in Figure 13.45 The data show that there is a

decrease in pre-edge intensity along the series, and that the pre-

edge peak maxima progressively shift to higher energy. (The

energy shift was related to shifting of charge density away from

the sulfur due to electron-withdrawing effect of the substituent on

the phenyl ring. However, simulation showed that this-I effect of

the substituent did not affect the Fe��S bond.) The energy and in-

tensity of the pre-edge features are quantitatively estimated from

fits to the experimental spectra and their second derivatives. The

normalized intensity of the thiolate-based transitions is related to

the total percent ligand character in the Fe3d antibonding manifold.

This decreases from 1.30 for FeP(SPh) to 0.80 for FeP(SL2) corre-

sponding to a decrease of Fe��S bond covalency from 49 to 31%. The

t2-e orbital splitting in these complexes is not large enough to allow

experimental resolution of the � and � contributions to bonding.

DFT calculations were performed on the high-spin (S ¼ 5/2)

ground states of these heme complexes. The calculated geometries

are in general agreement with the crystal structures and reproduce

the Fe��S bond elongation on H-bonding as observed crystallo-

graphically (Table 3).44 The calculations can be correlated to the

experimentally observed changes in pre-edge intensity. The MO

diagram for the FeP(SPh) complex (Fig. 14) shows the dyz orbital

has a � type interaction with the Sthiolate 3p orbital in the plane of

the aromatic ring, and the dz2 orbital has a pseudo-�-type interac-

tion with the thiolate orbital out of the plane of the ring. In the

crystal structure and the optimized geometry of the complexes, the

thiolate binds such that the N��H bonds are oriented directly to-

ward the in-plane S 3p � orbital. The sum of Sthiolate 3p character

in these �-spin unoccupied Fe 3d orbitals (Table 3) decreases from

42 to 30% from FeP(SPh) to FeP(SL2), paralleling the experimen-

tal results in Table 3. The calculations also indicate that the

decrease in the thiolate contribution is solely in the � type orbital,

consistent with the orientation of the H-bonds in these complexes.

The energy of the H-bonding interaction was evaluated for

both the aryl thiolate and a simplified alkyl thiolate (�SMe) with

the same donor (H2O to model the H-bonding interaction from the

ligand) and the results were the same.45 DFT calculations for

FeP(SMe) and FeP(SMe) þ 2H2O show that the Fe��S bond

length increases on H-bonding by 0.02 A in the oxidized form

(Table 4 and Fig. 15). The covalency of the Fe��S bond decreases

from 30 to 20% in the � orbital due to H-bonding (Table 4), simi-

lar to the effects found experimentally (Figure 13). The DE of H-

bonding calculated for the ferric heme complex is about �5 kcal

mol�1 for two H-bonds to the thiolate in the gas phase. This is in

Figure 12. Schematic diagram of the model complexes (left to right) FeP(SPh), FeP(SL1), and

FeP(SL2) (H-bonds shown as dashed lines).

Figure 13. S K-edge XAS of FeP(SPh) (in bold), FeP(SL1) (in

dots), and FeP(SL2) (in dashed).



good agreement with previous estimates of the H-bond energy of

sulfur donors.46

Energy of H-Bonding and Decomposition

This net change in energy (�5 kcal mol�1) is small considering

the large change in Fe��S bonding interaction (decrease in � cova-

lency by 33%; Table 3) involved in the process. The decrease of

ligand–metal bond covalency should be reflected in the energy of

metal–ligand bond more than in the total energy of the system.

This is indicated in the energy decomposition given in Scheme 1,

which shows that the bond energy of the H-bonded thiolate ligand

and the FeIII-heme fragment is about �129 kcal mol�1, while that

of the free thiolate is �151 kcal mol�1, 22 kcal/mol higher than

the H-bonded ligand. Thus there is, in fact, a 22 kcal mol�1

decrease in BE of the Fe��S bond between the H-bonded and the

non-H-bonded complexes that corresponds to the dramatic

decrease in the covalent interaction observed from the S K pre-

edge intensity (Fig. 13). However, the H-bonding energy of the

thiolate with H2O is �27 kcal mol�1, which compensates for this

difference in bonding energy and further stabilizes the system by

the �5 kcal mol�1.

Similar DFT calculations performed on the one-electron

reduced ferrous–heme complexes, show that there is an 0.05 A

increase in Fe��S bond length (Fig. 15) on H-bonding and the DEof H-bonding is about �12–15 kcal mol�1 for both alkyl and aryl

thiolates. Thus, the H-bonding stabilization is �6 kcal mol�1

greater (Table 4) for the reduced relative to the oxidized complex

for both alkyl and aryl thiolates after taking solvation into account

(this stabilization is 10 kcal mol�1 in the gas phase). This energy

difference contributes to the observed redox potential difference

for these complexes. This H-bond energy difference predicts a

260 mV higher redox potential for the H-bonded couple. The

experimentally observed difference in redox potential between

the FeP(SPh) and FeP(SL2) couple is 330 mV. [The calculated

FeII/III redox potentials for the FeP(SMe) and the FeP

(SMe)þ(2H2O) complexes are about �520 and �360 mV, respec-

tively, in CH2Cl2. The measured redox potentials in the same

solvent for FeP(SPh) and FeP(SL2) are �680 and �350 mV,

respectively.]

In summary, the covalency of a Fe��Sthiolate bond is about 42–

49% without H-bonding and 32% with it, which is similar to that

observed for the Cu��Sthiolate bond in plastocyanin (38% with a

single N��H----Scys H-bond). The covalency of this bond is signif-

icantly reduced in the active sites of rubredoxins due to N��H----S

H-bonding from the backbone. This effect is quantitatively esti-

mated in a series of high spin heme Fe–Sthiolate complexes where

there is about a 30% decrease (49 to 32% in absolute numbers) of

Fe–Sthiolate covalency for two H-bonds. DFT calculations repro-

duce these experimental trends and further show that the energy of

Table 3. DFT (BP86/TZP in ADF 2004) Calculated Bond Lengths and Covalencies for the FeP(X) Complexes

(Crystallographic Distance Are Given in Parenthesis).

Distance (A)

Covalency

(% S3p) Total covalency (% S3p)

Fe–S Fe–N N–S � � (% S3p)

FeP(SPh) 2.30 (2.30) 2.11 (2.06) N/A 24 18 42

FeP(SL1) 2.33 (2.33) 2.10 (2.05) 2.98 (2.93) 17 18 35

FeP(SL2) 2.38 (2.36) 2.10 (2.05) 2.98 (2.96) 12 18 30

Figure 14. DFT-calculated MO diagram (BP86/TZ2P) of the

FeP(SPh) complex. The � LUMO orbitals are pictured. The dz2 or-

bital has a pseudo-� (p�) interaction with the out-of-plane thiolate

donor orbital, and the dyz orbital has a � interaction with the in-

plane thiolate donor orbital. The inset shows the reference coordi-

nate system.



H-bonding is small (�5 kcal mol�1 in the gas phase) compared to

the large change in bonding observed in the S K-edge data. A

bond energy decomposition scheme shows that the overall stabili-

zation is a result of weakening in Fe–thiolate bond due to H-bond-

ing coupled to the energy of H-bonding to the free ligand. The H-

bonding stabilization increases by 10 kcal mol�1 (6 kcal mol�1 in

solvent) in the FeII form due to higher charge density on the thio-

late (less covalent FeII–Sthiolate bond), which increases the redox

potential as observed experimentally.

Concluding Comments

Metal–thiolate bonds play a dominate role in bioinorganic chemis-

try. These can be highly covalent and � or � (i.e., pseudo �; Fig.2) in nature depending on the effective nuclear charge of the metal

ion and its dn configuration. The energy of the metal–thiolate bond

is fairly insensitive to its ionic/covalent and �/� nature as increas-

ing covalency reduces the charge separation; hence, the ionic term

and these contributions can compensate. Thus, trends observed in

Scheme 1. Bonding energy decomposition scheme for H-bonding interaction. (Energies in kcal mol�1

and calculated in the gas phase at the BP86/TZ2P level.)

Table 4. DFT-Calculated FePSX Energies, MPA Populations, and Hirschfield Charges.

Covalency Charge

DEf (kcal mol�1)

(solvent)Fe–S (A) � � Fe S H2O

FeP(SPh) 2.30 24 18 0.34 �0.07 �0.10 �4.0 (1.0)

FeP(SPh) þ 2H2O 2.33 12 18 0.36 �0.03

FeP(SMe) 2.30 30 21 0.32 �0.06 �0.14 �5.1 (�1.5)

FeP(SMe) þ 2H2O 2.32 19 21 0.34 �0.03

FeP(SPh)(red) 2.34 10 9 0.24 �0.16 �12.2 (�7.2)

FeP(SPh) þ 2H2O (red) 2.39 4 7

FeP(SMe)(red) 2.32 13 14 0.24 �0.17 �0.26 �15.6 (�7.5)

FeP(SMe) þ H2O (red) 2.36 5 13 0.25 �0.09

Energies obtained using the PCMmethod and CH2Cl2 as a solvent (G03 BP86/6-311þG*) are reported in parenthesis.

Figure 15. Optimized geometries and relevant bond lengths of the oxidized (in black) and reduced (in

gray) complexes with (a) and without (b) H-bonding used to evaluate the energy of H-bonding.



stability constants (i.e., the Irving–Williams series) then mostly

reflect the dominantly ionic contribution to bonding of the inno-

cent ligand being replaced by the thiolate. Due to their high effec-

tive nuclear charges of CuII and FeIII, cupric– and ferric–thiolate

bonds are very covalent, with the former having strong � and the

latter having more � character. For the blue copper site, the high �covalency couples the metal ion into the protein for rapid direc-

tional long range electron transfer. For rubredoxins, because the

redox active molecular orbital is � in nature, electron transfer

tends to be more localized in the vicinity of the active site.

Although the energy of hydrogen bonding of the protein environ-

ment to the thiolate ligands tends to be fairly small, H-bonding

can greatly reduce the covalency of the thiolate–metal bond, which

destabilizes the oxidized more than the reduced state and can sig-

nificantly contribute to redox tuning by the protein environment.

Acknowledgments

S.I.G. is grateful to NSERC (Ottawa) for a postdoctoral fellow-

ship. E.I.S. thanks his students and collaborators as cited in the

references for their contributions to this research.

References

1. Rao, P.; Holm, R. H. Chem Rev 2004, 104, 527.

2. (a) Yeh, A. P.; Hu, Y.; Jenney, F. E., Jr.; Adams, M. W. W.; Rees, D.C.

Biochemistry 2000, 39, 2499; (b) Wuerges, J.; Lee, J.-W.; Yim, Y.-I.;

Yim, H.-S.; Kang, S.-O.; Carugo, K. D. Proc Natl Acad Sci 2004, 101,

8569; (c) Kovacs, J. A. Chem Rev 2004, 104, 825.

3. (a) Nagashima, S.; Nakasako, M.; Dohmae, N.; Tsujimura, M.; Takio, K.;

Odaka, M.; Yohda, M.; Kamiya, N.; Endo, I. Nat Struct Biol 1998, 5, 347;

(b) Huang, W.; Jia, J.; Cummings, J.; Nelson, M.; Schneider, G.;

Lindqvist, Y. Structure 1997, 5, 691; (c) Nguyen, K. T.; Hu, X.;

Colton, C.; Chakrabarti, R.; Zhu, M. X.; Pei, D. Biochemistry 2003,

42, 9952; (d) Moras, D.; Olsen, K. W.; Sabesan, M. N.; Buehner, M.;

Ford, G. C.; Rossmann, M. G. J Biol Chem 1977, 250, 9137.

4. Halcrow, M. A.; Christou, G. Chem Rev 1994, 94, 2421.

5. Solomon, E. I.; Szilagyi, R. K.; George, S. D.; Basumallick, L.

Chem Rev 2004, 104, 419.

6. Colman, P. M.; Freeman, H. C.; Guss, J. M.; Murata, M.; Norris, V. A.;

Ramshaw, J. A. M.; Venkatappa, M. P. Nature 1978, 272, 319.

7. Guss, J. M.; Bartunik, H. D.; Freeman, H. C. Acta Crystallogr Sect

B Struct Sci 1992, 48, 790.

8. Penfield, K. W.; Gewirth, A. A.; Solomon, E. I. J Am Chem Soc

1985, 107, 4519.

9. Solomon, E. I.; Hedman, B.; Hodgson, K. O.; Dey, A.; Szilagyi,

R. K. Coord Chem Rev 2005, 249, 97.

10. Shadle, S. E.; Penner–Hahn, J. E.; Schugar, H. J.; Hedman, B.;

Hodgson, K. O.; Solomon, E. I. J Am Chem Soc 1993, 115, 767.

11. Kitajima, N.; Fujisawa, K.; Tanaka, M.; Moro–Oka, Y. J Am Chem

Soc 1992, 114, 9232.

12. Matsunaga, Y.; Fujisawa, K.; Ibi, N.; Miyashita, Y.; Okamoto, K.

Inorg Chem 2005, 44, 325.

13. McMillin, D. R.; Rosenberg, R. C.; Gray, H. B. Proc Natl Acad Sci

USA 1974, 71, 4760.

14. Tennent, D. L.; McMillin, D. R. J Am Chem So. 1979, 101,

2307.

15. Nar, H.; Huber, R.; Messerschmidt, A.; Filippou, A. C.; Barth, M.;

Jaquinod, M.; van de Kamp, M.; Canters, G. W. Eur J Biochem

1992, 205, 1123.

16. Di Bilio, A. J.; Chang, T. K.; Malmstrom, B. G.; Gray, H. B.; Goran

Karlsson, B.; Nordling, M.; Pascher, T.; Lundberg, L. G. Inorg Chim

Acta 1992, 198–200, 145.

17. Bonander, N.; Vanngard, T.; Tsai, L.-C.; Langer, V.; Nar, H.; Sjolin,

L. Proteins Struct Funct Genet 1997, 27, 385.

18. Moratal, J. M.; Romero, A.; Salgado, J.; Perales–Alarcon, A.; Jime-

nez, H. R. Eur J Biochem 1995, 228, 653.

19. Funk, T.; Kennepohl, P.; Di Bilio, A. J.; Wehbi, W. A.; Young, A. T.;

Friedrich, S.; Arenholz, E.; Gray, H. B.; Cramer, S. P. J Am Chem Soc

2004, 126, 5859.

20. De Kerpel, J. O. A.; Pierloot, K.; Ryde, U. J Phys Chem B 1999,

103, 8375.

21. Irving, H.; Williams, R. J. P. Nature 1948, 162, 746.

22. Irving, H.; Williams, R. J. P. J Chem Soc 1953, 3192.

23. Gorelsky, S. I.; Basumallick, L.; Vura–Weis, J.; Sarangi, R.; Hed-

man, B.; Hodgson, K. O.; Fujisawa, K.; Solomon, E. I. Inorg Chem

2005, 44, 4947.

24. Casida, M. E. In Recent Advances in Density Functional Methods;

Chong, D. P., Ed.; World Scientific: Singapore, 1995, p. 155.

25. Stratmann, R. E.; Scuseria, G. E.; Frisch, M. J. J Chem Phys 1998,

109, 8218.

26. Gewirth, A. A.; Solomon, E. I. J Am Chem Soc 1988, 110, 3811.

27. Gewirth, A. A.; Cohen, S. L.; Schugar, H. J.; Solomon, E. I. Inorg

Chem 1987, 26, 1133.

28. Lever, A. B. P. Inorganic Electronic Spectroscopy; Elsevier: Amster-

dam, 1984; 2nd ed.

29. Gorelsky, S. I.; Ghosh, S.; Solomon, E. I. J Am Chem Soc 2006,

128, 278.

30. Gorelsky, S. I. AOMix: Program for Molecular Orbital Analysis;

York University: Toronto, Canada (http://www.sg-chem.net).

31. Dapprich, S.; Frenking, G. J Phys Chem 1995, 99, 9352.

32. Reed, A. E.; Curtiss, L. A.; Weinhold, F. Chem Rev 1988, 88,

899.

33. Morokuma, K. J. Chem Phys 1971, 55, 1236.

34. Kitaura, K.; Morokuma, K. Int J Quantum Chem 1976, 10, 325.

35. Chen, W.; Gordon, M. S. J. Phys Chem 1996, 100, 14316.

36. Besler, B. H.; Merz, K. M., Jr.; Kollman, P. A. J Comput Chem

1990, 11, 431; Francl, M. M.; Chirlian, L. E. Rev Comput Chem

2000, 14, 1.

37. Umeyama, H.; Morokuma, K. J Am Chem Soc 1977, 99, 1316.

38. Ziegler, T.; Rauk, A. Theoret Chim Acta 1977, 46, 1.

39. Mayer, I. Chem Phys Lett 1983, 97, 270.

40. Gebhard, M. S.; Deaton, J. C.; Koch, S. A.; Millar, M.; Solomon, E. I.

J Am Chem Soc 1990, 112, 2217.

41. Rose, K.; Shadle, S. E.; Eidsness, M. K.; Kurtz, D. M., Jr.; Scott,

R. A.; Hedman, B.; Hodgson, K. O.; Solomon, E. I. J Am Chem

Soc 1998, 120, 10743.

42. Kennepohl, P.; Solomon, E. I. Inorg Chem 2003; 42; 696.

43. Torres, R. A.; Lovell, T.; Noodleman, L.; Case, D. A. J Am Chem

Soc 2003, 125, 1923.

44. Okamura, T.; Takamizawa, S.; Ueyama, N.; Nakamura, A. Inorg

Chem 1998, 37, 18.

45. Dey, A.; Okamura, T.-A.; Ueyama, N.; Hedman, B.; Hodgson,

K. O.; Solomon, E. I. J Am Chem Soc, 2005, 127, 12046.

46. Meot–Ner, M. Chem Rev 2005, 105, 213.



https://www.researchgate.net/publication/220023836_Recent_Advances_in_Density_Functional_Methods_Part_I_Vol_1?el=1_x_8&enrichId=rgreq-cbedfb1b-31c1-4644-b401-55245464ee7a&enrichSource=Y292ZXJQYWdlOzY5NzY2ODc7QVM6MTU1NzY3NTQ2MDYwODAwQDE0MTQxNDkyODU2ODk=

https://www.researchgate.net/publication/220023836_Recent_Advances_in_Density_Functional_Methods_Part_I_Vol_1?el=1_x_8&enrichId=rgreq-cbedfb1b-31c1-4644-b401-55245464ee7a&enrichSource=Y292ZXJQYWdlOzY5NzY2ODc7QVM6MTU1NzY3NTQ2MDYwODAwQDE0MTQxNDkyODU2ODk=

https://www.researchgate.net/publication/287783610_Inorganic_Electronic_Spectroscopy?el=1_x_8&enrichId=rgreq-cbedfb1b-31c1-4644-b401-55245464ee7a&enrichSource=Y292ZXJQYWdlOzY5NzY2ODc7QVM6MTU1NzY3NTQ2MDYwODAwQDE0MTQxNDkyODU2ODk=

https://www.researchgate.net/publication/287783610_Inorganic_Electronic_Spectroscopy?el=1_x_8&enrichId=rgreq-cbedfb1b-31c1-4644-b401-55245464ee7a&enrichSource=Y292ZXJQYWdlOzY5NzY2ODc7QVM6MTU1NzY3NTQ2MDYwODAwQDE0MTQxNDkyODU2ODk=

Metal–thiolate bonds in bioinorganic chemistry

Documents