Download 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 2005
DOI 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 MSthiolate bonds can be
very covalent, and involve different orbital interactions. Spectroscopic features of these active sites (intense, lowenergy charge transfer transitions) reflect the high covalency of the M
Sthiolate bonds. The energy of the metal–thiolate 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 constants (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 environment 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.
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Key words: covalency; XAS; metal–thiolate bonds; spectroscopy; TD-DFT
Introduction
Metal–thiolate bonds are present for many classes of metalloprotein active sites and make major contributions to function. The
protein centers involved in electron transfer include the blue copper 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 residues coupling them into the protein matrix.1 The main heme
enzyme involved in O2 activation is P450, which has an axial thiolate–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 thiolateZn bond as in alcohol dehydrogenase.3 Other metalloenzymes
with bridging thiolate ligation include the hydrogenases, a binuclear (Fe–Fe or Ni–Fe) cluster involved in the reduction of dihydrogen, sulfite reductase, an FeIII porphyrin bridged to a Fe4S4
cluster involved in reduction of sulfite, and CO dehydrogenase,
which has a binuclear Ni cluster bridged to a Fe4S4 cluster (Acluster) 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 valence 3p orbitals, one of which is greatly stabilized in energy due
to the carbon–sulfur bond, and thus does not significantly contribute to the thiolate sulfur–metal bond (Fig. 2). The two remaining 3p orbitals, which are perpendicular to the S
C bond, dominate 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 inplane 3p orbital pseudo bonds to the metal ion (pseudo in the
sense that when the C–S–M angle is greater than 908 its electron
density is shifted off the S–M bond; C–S–M bond angles are generally in the range of 100–1208). The specific bonding interactions
depend on the metal ion, its 3dn configuration and its ZEff (effective 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 covalent, 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.).
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Figure 1. Biological centers having functionally relevant M-Sthiolate bonds.
bonding, where the increase in effective nuclear charge for the ferric 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 Å
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
dx2y2 orbital (i.e., oriented in the xy plane) which is highly covalent 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 -spin
LUMO, which closely reflects the total spin density). The 1s orbital 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 dx2y2 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–thiolate bond with 15% covalency. [In this model complex, the Cu
atom has a five-coordinate geometry, Cu–Sthiolate distance is
2.36 Å, and the Sthiolate p ! Cu and Sthiolate p ! Cu LMCT
transitions are at higher energy and much lower intensity—23,500
cm1 (" ¼ 430 M1 cm1) and 27,800 cm1 (" ¼ 360 M1 cm1),
respectively, than those in blue-copper sites. Both the XAS inten-
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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–Williams series for the stability constants.21,22
Spectroscopic Signatures of p and s Metal–Thiolate
Bond Covalency
Figure 2. Sulfur-based valence orbitals of methyl thiolate.
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 intensity 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 dx2y2 orbital to form bonds to donor ligands. This interaction derived from analysis
of the charge transfer (CT) absorption spectrum using low-temperature 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 cm1 (" 5000 M1
cm1), 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
cm1 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 cm1 band is the lowest energy; hence, CT
transition, while a higher energy weak feature is the Sthiolate
pseudo to CuII CT transition. The intense-/weak- CT absorption 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 dx2y2 orbital be oriented so that it interacts with the thiolate sulfur. This bond is
due to the short thiolate S–CuII bond length of 2.1 Å.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.
A combination of absorption, MCD, and resonance Raman (rR)
spectroscopies was used23 to distinguish the ligand field (LF) tran-
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
Figure 3. (A) Structure and the LUMO of the blue copper center;
(B) S K-edge XAS of plastocyanin (PCu) and the [Cu(tet
b)(SC6H4CO2) complex; (C) low-temperature absorption and (D)
MDC spectra of poplar plastocyanin.
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and Ni complexes, to less intense Sthiolate p ! MII dx2y2 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 cm1 due
to the Sthiolate p LMCT transition is more intense than the transition at 25,000 cm1 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
cm1 region. The Sthiolate p ! Fe dxy LMCT excitation appears as
the intense absorption (" 8000 M1 cm1) at 32,000 cm1, compared to the less intense (" 2200 M1 cm1) Sthiolate p ! Fe
dx2y2 CT excitation contributing to the absorption intensity at 28,000 cm1. The Sthiolate p ! Fe dxy LMCT transition is higher
in energy (by 2000 cm1) and lower intensity (by 3000 M1 cm1),
and the Sthiolate p ! Fe dx2y2 LMCT transition is also higher in
energy (by 4000 cm1) and lower in intensity (by 800 M1 cm1)
relative to the corresponding excitations in the Co complex.
Figure 4. UV-Vis absorption spectra of [MIIL(SC6F5)] in cyclohexane at room temperature.
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 (TDDFT)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 cm1, due to the Sthiolate
p ! Cu dx2y2 (-spin HOMO ! LUMO) CT transition. As
described earlier, this is also characteristic of the blue copper proteins 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 cm1,
which is due to the Sthiolate p ! Ni dx2y2 LMCT transition. This
transition is higher in energy (from 15,000 cm1 to 20,000 cm1)
and lower in intensity (from 12,000 M1 cm1 to 8000 M1
cm1) compared to the corresponding transition in the CuII complex (Figs. 4 and 5B). In the higher energy region (>25,000 cm1)
there are the Sthiolate p ! Ni dx2y2, benzyl based intraligand CT
transitions, the Sthiolate p ! Ni dz2, and the (pyrazolyl) p ! Ni
dx2y2 transitions.
In the absorption spectrum of [CoIIL(SC6F5)] (Figs. 4 and 5D),
The Sthiolate p ! Co dx2y2 charge transfer transition is now at
24,000 cm1 (" 3000 M1 cm1); this transition is higher in
energy and lower in intensity compared to the corresponding band
in the NiII complex (20,000 cm1 and " 8000 M1 cm1) as
observed experimentally. The interesting feature in the CoII complex is an additional intense (" 11,000 M1 cm1) Sthiolate p !
Co dxy charge transfer transition at 30,000 cm1. In CoII, the dxy
orbital is unoccupied and interacts with the Sthiolate p orbital, producing an intense thiolate-to-metal CT absorption band at 30,000
cm1. This reverses the pattern of more intense Sthiolate p and less
intense Sthiolate p ! MII dx2y2 CT transitions observed in the Cu
Figure 5. Simulated electronic spectra of [MIIL(SC6F5)] (TD-DFT)
calculations at the B3LYP/6-311þG(d) level. The calculated energies and intensities of electronic transitions were transformed into
simulated spectra as described in ref. 23 using Gaussian functions
with half-widths of 2500 cm1. The contributions from individual
electronic transitions are shown in gray.
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Metal–Thiolate bonds in Bioinorganic Chemistry
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.
The [MnIIL(SC6F5)] complex does not have absorption intensity in the 5000–25,000 cm1 region (Figs. 4 and 5F). This is
because the LF transitions in high spin Mn(d5) complexes are
both spin and Laporté forbidden.28 The absorption intensity at
35,000 cm1 is due to Sthiolate p ! Mn dx2y2 LMCT transition, and at 38,000 cm1 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 dx2y2 transitions shift to higher energy and
decrease in intensity, and (2) a new intense Sthiolate p ! M dxy
band 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 features arising due to the different metal–thiolate interactions in
these complexes, trends in the metal–thiolate bonding were evaluated using DFT.
Molecular Orbital Description of Metal–Thiolate s and p
Bonding 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
1419
extension of the charge decomposition analysis,31 indicates that
the metal–thiolate covalent interactions in these series are dominated by charge donation from the SC6F5 ligand to the MLþ fragment (Table 1). This charge donation has both - and -orbital
components, and the two highest occupied molecular orbitals of
the thiolate ligand, HOMO (S p) and HOMO-1 (S p), are principal 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 -spin
occupied 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 donation 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 contribute little to the covalent bonding relative the stronger MnII 4s,4p–
Sthiolate orbital interaction. The Mn 3d orbital character of the thiolate-based HOMO() and HOMO-3() (the thiolate character is
95 and 86%, respectively) is 3 and 5%, respectively, and reciprocally the Sthiolate contribution is 2% for the Mn dx2y2 (d) and 3%
for dxy (d) orbitals (Fig. 6).
In the FeII complex with one more valence electron (which
occupies the -spin dyzxz 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 Zeff
nuc for
Fe relative to Mn. The increased Zeff
nuc shifts the energies of metalbased 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 Zeff
nuc for Co is greater than that of Fe, the HOMO-LUMO gap again becomes smaller (3.2 eV) and the covalency of the M
Sthiolate bond is further increased (the Co 3d character 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 spectroscopic data (Figs. 4 and 5), which show an intense the Sthiolate
p ! Co dxy CT transition.
In [NiIIL(SC6F5)], the dxy orbital is now occupied. This cancels 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 decreases. Because the dxy orbital is now below the Sthiolate p
orbital (Fig. 6), the latter is destabilized by 1 eV. The greater
Zeff
nuc for Ni results in a further lowering of the energies of M 3dbased orbitals and favors more efficient covalent bonding with
the Sthiolate p orbital. As a result, the LUMO (dx2y2) has 8% S
contribution (Fig. 6).
In the CuII complex, the dz2 orbital is populated and the Cu
dx2y2 is the only unoccupied d orbital. As a result of the still
larger Zeff
nuc, the -HOMO-LUMO gap for [CuL(SC6F5)] is lowest
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29,30
Table 1. Extended Charge Decomposition Analysis
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 SC6F
5 Fragments, and
þ
the Net Charge Donation from the SC6F
5 Thiolate to the ML Fragment (B3LYP/TZVP Calculations).
Metal
%
%
%
%
-spin HOFO(SC6F5) a
-spin HOFO(SC6F5) a
-spin HOFO-1(SC6F5) b
-spin HOFO-1(SC6F5) b
Mn
Fe
Co
Ni
Cu
Zn
2.9
6.8
17.1
17.5
3.1
10.2
16.8
19.4
3.1
9.2
16.2
24.0
2.9
14.1
16.8
18.2
3.2
41.1
16.0
16.2
3.5
3.5
20.8
20.8
28.2
47.3
26.2
81.7
32.5
32.5
þ
þ
donation(SC6F
5 ! ML ) and polarization(ML )
þ
% -spin UFOs(ML )
% -spin UFOs(MLþ)
29.5
32.0
29.6
33.6
28.2
48.9
þ
donation(SC6F
5 / ML ) and polarization(SC6F5 )
% -spin UFOs(SC6F
5)
% -spin UFOs(SC6F
5)
c
PL(MLþ) - PL(SC6F
5)
þ
c
PL (ML ) - PL (SC6F5 )
donation(SC6F5 ! MLþ)d
donation(SC6F5 ! MLþ)d
donationþ(SC6F5 ! MLþ)
3.1
2.7
2.4
0.2
0.240
0.295
0.54
3.5
4.8
2.1
5.0
0.240
0.335
0.58
3.5
4.2
1.3
7.4
0.234
0.372
0.61
3.4
3.7
1.1
6.2
0.237
0.375
0.61
4.3
4.0
0.9
16.0
0.229
0.617
0.85
2.7
2.7
1.3
1.3
0.285
0.285
0.57
þ
The principal component of donation(SC6F
5 ! ML ): the HOMO(SC6F5 ) contribution to the unoccupied MOs of the complex.
þ
The main component of donation(SC6F
5 ! ML ): the HOMO-1(SC6F5 ) contribution to the unoccupied MOs of
the complex.
c
The difference (orbital %) in electronic polarization (PL) between the MLþ and SC6F
5 fragments for - and -spin
orbitals, respectively.
d
þ
The net charge donation from the SC6F
5 thiolate to the ML fragment for - and -spin orbitals, respectively. The
former is largely donation and the latter is both and donation.
a
b
(2.4 eV) maximizing the orbital interaction between the Sthiolate p
and Cu dx2y2 (Fig. 6). These two fragment orbitals mix to form
the bonding -spin HOMO [49% of HOMO(SC6F5) and 41% LUMO(Cu dx2y2 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 dx2y2 CT
transition. In addition, ECDA29 (Table 1) shows that the electronic
polarization of the MLþ fragment by the thiolate ligand is the largest for the CuII complex, 16 orbital%, in agreement with the smallest HOMO-LUMO gap of the CuLþ fragment in the series. This
polarization provides additional stabilization to the highly covalent CuII–thiolate bond.
In the ZnII complex, all the d orbitals of the metal ion are occupied, including the Zn dx2y2–Sthiolate p orbital. Thus, the contribution to covalency is cancelled, the covalent bonding interaction 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 interaction comes mostly from the ZnII 4s orbital.
32
Table 2. Metal and Sulfur NPA -Derived Charges and Metal 3d, 4s, and 4p Atomic Orbital
NPA Populations (-spin)a in the [MIIL(SC6F5)] (B3LYP/TZVP Calculations).
Metal
qNPA(M)
qNPA(S)
-spin 3d(M) populationa
-spin 4s(M) populationa
-spin 4p(M) populationa
a
Mn
Fe
Co
Ni
Cu
Zn
1.37
0.36
0.35
0.14
0.02
1.30
0.32
1.39
0.15
0.02
1.27
0.31
2.40
0.16
0.01
1.24
0.32
3.40
0.18
0.01
1.13
0.19
4.49
0.19
0.01
1.51
0.40
4.99
0.23
0.01
NPA populations of metal 3d, 4s, and 4p -spin orbitals are 4.92–4.99, 0.18–0.23, 0.01–0.03 electrons, respectively.
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Metal–Thiolate bonds in Bioinorganic Chemistry
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force constants implies that there are varying contributions to the
metal–thiolate bonding in this series, and a more detailed examination 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 complexes:
Eprep ¼ Eprep ðMLþ Þ þ Eprep ðSC6 F
5Þ
Over this series, Eprep(MLþ) and Eprep(SC6F5) was 5.3–10.3
kcal mol1 and 1.0–1.2 kcal mol1, respectively. Eint is the interaction 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 meaningful way:
Figure 7. (A) Experimental and calculated (at the B3LYP/6311þG* 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.
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 nonmonotonic variations in the M
S bond lengths mark important changes
in the metal–thiolate bonding, depending on the electronic configuration of the metal ion. The MLþ–SC6F5 binding energy was
calculated as the interaction energy between the metal–pyrazolyl
fragment and the thiolate ligand:
þ
ML þ
SC6 F
5
Eint ¼ Ecov þ Eionic :
Here, Ecov is the covalent or orbital interaction energy (including the exchange repulsion energy33–35) and Eionic is the electrostatic interaction energy. The latter is estimated as a sum of electrostatic 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–Kollman 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 mol1) is close to the NPA-derived value,
104.5 kcal mol1. 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 ¼
¼ ½M LðSC6 F5 Þ:
II
In the [MIIL(SC6F5)] series, it ranges (Fig. 7B) from 124.7 kcal
mol1 (Fe) to 132.7 kcal mol1 (Zn), and its variation does not
mirror the changes in the M
S bond lengths.23 M
S force constant is the lowest for the MnII complex (1.13 mDyne Å1),
increases from MnII to FeII to CoII, where it reaches a local maximum, then decreases for the NiII complex, and increases for the
CuII complex (the maximum value, 1.43 mDyne Å1), 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
X qNPA qNPA
a
b
ðin atomic unitsÞ
r
ab
a 2 ML b 2 SC F
X
6 5
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–Morokuma33,34,37 and Ziegler,38 which is calculated with undistorted
charge distributions corresponding to those in the isolated fragments. (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 decomposition is the noninteracting MLþ and SC6F5 fragments.)
Because the metal–pyrazolyl fragment and the thiolate ligand
carry opposite charges, the electrostatic interaction is attractive
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(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 sulfur charges are highest in the MnII and ZnII complexes and lowest
in the CuII complex. Thus, it can be expected that the ionic contribution to the M
S bond energy is lowest in the CuII complex and
highest in the MnII and ZnII complexes. Indeed, as will be discussed later in this section, the ionic bonding plays an important
role in determining the overall strength of the MLþ–SC6F5 interaction.
The analysis of covalent contributions to the chemical bonding
can be done using Mayer bond orders,39 BAB, and its - and -spin
23
orbital components, BAB
and BAB
.
BAB ¼
XX
ðPSÞba ðPSÞab þ ðPs SÞba ðPs SÞab ¼ BAB þ BAB ;
a2A b2B
BAB ¼ 2
XX
ðP SÞba ðP SÞab ;
BAB ¼ 2
XX
ðP SÞba ðP SÞab ;
a2A b2B
a2A b2B
where P and Ps are the density and spin-density matrices, respectively (P ¼ P þ P; Ps ¼ P P), P and P are - and -spin
electron 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–SC6F5
bond order) in the [MIIL(SC6F5)] complexes follows the Mn < Fe
< Co > Ni < Cu > Zn progression found in the M
S force constants (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 -bond
contributions, symmetry-adapted orbitals were used in the bondorder analysis:23
BAB ¼
X
BAB ð Þ;
where BAB (G) is the bond order contribution from orbitals with irreducible 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 resulting bond contributions to the Mayer bond orders from the -spin
MOs, BM
S (), remain similar for all complexes in this series.
Because all -spin metal 3d-based orbitals are occupied and cannot contribute to covalent bonding with the thiolate donor, the bond component from -spin MOs, BM
S (), is close to zero
(Fig. 8A).
Thus, BM
S () and B M
S () provide a reference for changes
in -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–Sthiolate
p and M dx2y2–Sthiolate p orbital interactions, respectively. The
II
II
increasing Zeff
nuc (from Mn to Zn ) brings the energies of the dxy
and dx2y2 orbitals of the metal atom closer to the thiolate occupied
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 NPAderived net charge transfer from the SC6F5 ligand to the MLþ framents. (C) The ionic component of the MLþ–SC6F5 bonding
energy, Eionic, and the difference, Eint Eionic.
orbitals and makes the M 3d-thiolate covalency stronger, given that
the electron occupancy of the appropriate MOs allows the net contribution to be positive. This is the case for the component,
II
II
II
BM
S (), in going from Mn to Fe to Co , and for the component, BM
(),
when
proceeding
from
MnII to CuII. These
S
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 dx2y2 (M ¼ Cu, Ni, Co, and Fe) CT transition is due to reduced M 3d–thiolate covalency, supported by
the decrease in the pre-edge intensity in the S K-edge XAS data23
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Metal–Thiolate bonds in Bioinorganic Chemistry
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metal–thiolate interaction energy in the Zn complex is 2 kcal
mol1 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 contributions to bonding. In such a case, the general correlation of the
bond energy, length, and the force constant does not hold.
Irving–Williams Series
Figure 9. The relative formation energies (kcal mol1) 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.
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
dx2y2 charge transfer transition. This is consistent with its ground
state electronic structure description, indicating the strong interaction between Co dxy and the Sthiolate p orbitals. Similar to the
trend in the Sthiolate p ! M dx2y2 CT transition, the Sthiolate
p ! 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 determining 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 negative (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
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 complexes 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 mol1 (Mn) to 130.8 kcal mol1 (Cu), to
132.7 kcal mol1 (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 (calculated at the B3LYP/6-311þG(3df) level) decreases from 176.4
kcal mol1 (Mn) to 166.1 kcal mol1 (Cu) and then increases to
176.6 kcal mol1 (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 contribution 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
þ
Ef ¼ ½Eo ðMLþ SC6 F
5 Þ Eo ðMnL SC6 F5 Þ
½Eo ðMLþ F Þ Eo ðMnLþ 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 covalent 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,
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Solomon, Gorelsky, and Dey Vol. 27, No. 12 Journal of Computational Chemistry
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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 delocalization enabling facile long range ET, the weak covalency of the redox active orbital in the rubredoxin site limits the electronic coupling 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
Figure 10. Sulfur K-edge XAS spectra of the [FEIII/II(SPh)4]/2
rubredoxin models (solid/dashed lines). The inset shows the preedge 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.
H2O, etc.), the ionic term dominates and their binding energies are
affected by changes in covalency over the series. It is the competition 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 discussed 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 transition (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 configuration.) The fitted intensity gives a total hole covalency of
170% (summed over the five unoccupied Fe3d orbitals), which corresponds 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 covalency, which is consistent with the weak and strong charge
transfer transitions observed in the absorption spectrum for this
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 reactivity. 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 display 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 different proteins. This reduction in pre-edge intensity implies a reduction 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).
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Figure 12. Schematic diagram of the model complexes (left to right) FeP(SPh), FeP(SL1), and
FeP(SL2) (H-bonds shown as dashed lines).
The amount of charge donation from the ligands to the metal,
that is, covalency of the metal–ligand bond, can make a very significant contribution to the redox potential particularly in covalent
systems like rubredoxin. Increased covalency preferentially stabilizes 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 systematic 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 complexes FeP(SPh), FeP(SL1), and FeP(SL2) (where L1 ¼ 2-trifluoroacetamido 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 preedge 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 intensity 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) corresponding 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 interaction 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 toward 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 experimental 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 Å 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), similar to the effects found experimentally (Figure 13). The DE of Hbonding calculated for the ferric heme complex is about 5 kcal
mol1 for two H-bonds to the thiolate in the gas phase. This is in
Figure 13. S K-edge XAS of FeP(SPh) (in bold), FeP(SL1) (in
dots), and FeP(SL2) (in dashed).
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Table 3. DFT (BP86/TZP in ADF 2004) Calculated Bond Lengths and Covalencies for the FeP(X) Complexes
(Crystallographic Distance Are Given in Parenthesis).
Covalency
(% S3p)
Distance (Å)
FeP(SPh)
FeP(SL1)
FeP(SL2)
Total covalency (% S3p)
Fe–S
Fe–N
N–S
(% S3p)
2.30 (2.30)
2.33 (2.33)
2.38 (2.36)
2.11 (2.06)
2.10 (2.05)
2.10 (2.05)
N/A
2.98 (2.93)
2.98 (2.96)
24
17
12
18
18
18
42
35
30
Energy of H-Bonding and Decomposition
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 reproduce these experimental trends and further show that the energy of
This net change in energy (5 kcal mol1) is small considering
the large change in Fe
S bonding interaction (decrease in covalency 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 mol1, while that
of the free thiolate is 151 kcal mol1, 22 kcal/mol higher than
the H-bonded ligand. Thus there is, in fact, a 22 kcal mol1
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 preedge intensity (Fig. 13). However, the H-bonding energy of the
thiolate with H2O is 27 kcal mol1, which compensates for this
difference in bonding energy and further stabilizes the system by
the 5 kcal mol1.
Similar DFT calculations performed on the one-electron
reduced ferrous–heme complexes, show that there is an 0.05 Å
increase in Fe
S bond length (Fig. 15) on H-bonding and the DE
of H-bonding is about 12–15 kcal mol1 for both alkyl and aryl
thiolates. Thus, the H-bonding stabilization is 6 kcal mol1
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 mol1 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, respectively, 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 significantly reduced in the active sites of rubredoxins due to N
H----S
H-bonding from the backbone. This effect is quantitatively esti-
Figure 14. DFT-calculated MO diagram (BP86/TZ2P) of the
FeP(SPh) complex. The LUMO orbitals are pictured. The dz2 orbital has a pseudo- (p) interaction with the out-of-plane thiolate
donor orbital, and the dyz orbital has a interaction with the inplane thiolate donor orbital. The inset shows the reference coordinate system.
good agreement with previous estimates of the H-bond energy of
sulfur donors.46
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Metal–Thiolate bonds in Bioinorganic Chemistry
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Table 4. DFT-Calculated FePSX Energies, MPA Populations, and Hirschfield Charges.
Covalency
FeP(SPh)
FeP(SPh) þ 2H2O
FeP(SMe)
FeP(SMe) þ 2H2O
FeP(SPh)(red)
FeP(SPh) þ 2H2O (red)
FeP(SMe)(red)
FeP(SMe) þ H2O (red)
Charge
Fe–S (Å)
Fe
S
H2O
DEf (kcal mol1)
(solvent)
2.30
2.33
2.30
2.32
2.34
2.39
2.32
2.36
24
12
30
19
10
4
13
5
18
18
21
21
9
7
14
13
0.34
0.36
0.32
0.34
0.24
0.07
0.03
0.06
0.03
0.16
0.10
4.0 (1.0)
0.14
5.1 (1.5)
0.24
0.25
0.17
0.09
0.26
12.2 (7.2)
15.6 (7.5)
Energies obtained using the PCM method 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.
Scheme 1. Bonding energy decomposition scheme for H-bonding interaction. (Energies in kcal mol1
and calculated in the gas phase at the BP86/TZ2P level.)
H-bonding is small (5 kcal mol1 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 stabilization is a result of weakening in Fe–thiolate bond due to H-bonding coupled to the energy of H-bonding to the free ligand. The Hbonding stabilization increases by 10 kcal mol1 (6 kcal mol1 in
solvent) in the FeII form due to higher charge density on the thiolate (less covalent FeII–Sthiolate bond), which increases the redox
potential as observed experimentally.
Concluding Comments
Metal–thiolate bonds play a dominate role in bioinorganic chemistry. 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 increasing covalency reduces the charge separation; hence, the ionic term
and these contributions can compensate. Thus, trends observed in
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Solomon, Gorelsky, and Dey Vol. 27, No. 12 Journal of Computational Chemistry
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stability constants (i.e., the Irving–Williams series) then mostly
reflect the dominantly ionic contribution to bonding of the innocent ligand being replaced by the thiolate. Due to their high effective 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 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 environment 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 significantly contribute to redox tuning by the protein environment.
Acknowledgments
S.I.G. is grateful to NSERC (Ottawa) for a postdoctoral fellowship. 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.
l
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.; Jimenez, 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.; Hedman, 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: Amsterdam, 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.
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DOI 10.1002/jcc