The first example of a nitrile hydratase model complex that reversibly binds nitriles

Article abstract of DOI:10.1021/ja012555f

Nitrile hydratase (NHase) is an iron-containing metalloenzyme that converts nitriles to amides. The mechanism by which this biochemical reaction occurs is unknown. One mechanism that has been proposed involves nucleophilic attack of an Fe-bound nitrile by

Full text of DOI:10.1021/ja012555f

Published on Web 08/29/2002
The First Example of a Nitrile Hydratase Model Complex that
Reversibly Binds Nitriles
Jason Shearer,† Henry L. Jackson,† Dirk Schweitzer,† Durrell K. Rittenberg,†
Tanya M. Leavy,‡ Werner Kaminsky,§ Robert C. Scarrow,*,‡ and Julie A. Kovacs*,†
Contribution from the Department of Chemistry, UniVersity of Washington, Box 351700,
Seattle, Washington 98195, and The Department of Chemistry, HaVerford College,
HaVerford, PennsylVania 19041
Received November 16, 2001
Abstract: Nitrile hydratase (NHase) is an iron-containing metalloenzyme that converts nitriles to amides.
The mechanism by which this biochemical reaction occurs is unknown. One mechanism that has been
proposed involves nucleophilic attack of an Fe-bound nitrile by water (or hydroxide). Reported herein is a
five-coordinate model compound ([FeIII(S2Me2N3(Et,Pr))]+) containing Fe(III) in an environment resembling
that of NHase, which reversibly binds a variety of nitriles, alcohols, amines, and thiocyanate. XAS shows
that five-coordinate [FeIII(S2Me2N3(Et,Pr))]+ reacts with both methanol and acetonitrile to afford a six-coordinate
solvent-bound complex. Competitive binding studies demonstrate that MeCN preferentially binds over ROH,
suggesting that nitriles would be capable of displacing the H2O coordinated to the iron site of NHase.
Thermodynamic parameters were determined for acetonitrile (∆H ) -6.2((0.2) kcal/mol, ∆S ) -29.4((0.8)
eu), benzonitrile (-4.2((0.6) kcal/mol, ∆S ) -18((3) eu), and pyridine (∆H ) -8((1) kcal/mol, ∆S )
-41((6) eu) binding to [FeIII(S2Me2N3(Et,Pr))]+ using variable-temperature electronic absorption spectroscopy.
Ligand exchange kinetics were examined for acetonitrile, iso-propylnitrile, benzonitrile, and 4-tert-butylpyridine
using 13C NMR line-broadening analysis, at a variety of temperatures. Activation parameters for ligand
exchange were determined to be ∆Hq ) 7.1((0.8) kcal/mol, ∆Sq ) -10((1) eu (acetonitrile), ∆Hq
)
5.4((0.6) kcal/mol, ∆Sq ) -17((2) eu (iso-propionitrile), ∆Hq ) 4.9((0.8) kcal/mol, ∆Sq ) -20((3) eu
(benzonitrile), and ∆Hq ) 4.7((1.4) kcal/mol ∆Sq ) -18((2) eu (4-tert-butylpyridine). The thermodynamic
parameters for pyridine binding to a related complex, [FeIII(S2Me2N3(Pr,Pr))]+ (∆H ) -5.9((0.8) kcal/mol,
∆S ) -24((3) eu), are also reported, as well as kinetic parameters for 4-tert-butylpyridine exchange
(∆Hq ) 3.1((0.8) kcal/mol, ∆Sq ) -25((3) eu). These data show for the first time that, when it is contained
in a ligand environment similar to that of NHase, Fe(III) is capable of forming a stable complex with nitriles.
Also, the rates of ligand exchange demonstrate that low-spin Fe(III) in this ligand environment is more
labile than expected. Furthermore, comparison of [FeIII(S2Me2N3(Et,Pr))]+ and [FeIII(S2Me2N3(Pr,Pr))]+
demonstrates how minor distortions induced by ligand constraints can dramatically alter the reactivity of a
metal complex.
Performing reactions under environmentally friendly condi-
tions has recently become a desirable goal in the search for new
catalysts.1 Enzymes are often viewed as ideal in this respect
because of their ability to perform chemical transformations
under mild conditions with high yields and stereospecificity.2
Nitrile hydratase (NHase) is an example of a metalloenzyme
that is used industrially, in the multikiloton production of
acrylamide from acrylonitrile.3 NHase is a mononuclear thiolate-
ligated non-heme iron (or noncorrinoid cobalt) metalloenzyme,
which converts nitriles to amides.4 Of the two metal-containing
forms, Co-NHase and Fe-NHase, the iron form has been the
most extensively studied.5 On the basis of EXAFS, EPR, and
two X-ray crystal structures, Fe-NHase has been shown to
(4) (a) Kobyashi, M. Nat. Biotechnol. 1998, 16, 733-736. (b) Shearer, J.;
Kovacs, J. A. Nitrile hydratase: An unusual Fe(III) ligated metalloenzyme.
Encyclopedia of Catalysis; John Wiley & Sons: New York, 2002, in press.
(5) (a) Huang, W.; Jia, J.; Cummings, J.; Nelson, M.; Schneider, G.; Lindqvist,
Y. Structure 1997, 5, 691-699. (b) Nagashima, S. N.; Naoshi, D.;
Tsujimura, M.; Takio, K.; Odaka, M.; Yohda, M.; Kamiya, N.; Endo, I.
Nat. Struct. Biol. 1998, 5, 347-351. (c) Nelson, M. J.; Jin, H.; Turner, M.
I., Jr.; Grove, G.; Scarrow, R. C.; Brennan, B. A.; Que, L., Jr. J. Am. Chem.
Soc. 1991, 113, 7072-7073. (d) Nagamune, T.; Honda, J.; Kobayashi, Y.;
Sasabe, H.; Endo, I.; Ambe, F. Hyperfine Interact. 1992, 71, 1271-1274.
(e) Brennan, B. A.; Cummings, J. G.; Chase, D. B.; Turner, M. I., Jr.;
Nelson, M. J. Biochemistry 1996, 35, 10068-10077 (the pH ) 9.0 form
of NHase has been shown to be inactive). (f) Jin, H.; Turner, I. M.; Nelson,
M. J.; Gurbiel, R. J.; Doan, P. E.; Hoffman, B. M. J. Am. Chem. Soc.
1993, 115, 5290-5291. (g) Payne, M. S.; Wu, S.; Fallon, R. D.; Tudor,
G.; Stieglitz, B.; Turner, I. M., Jr.; Nelson, M. J. Biochemistry 1997, 36,
5447-5454. (h) Nojiri, M.; Nakayama, H.; Odaka, M.; Yohda, M.; Takio,
K.; Endo, I. FEBS Lett. 2000, 465, 173-177. (i) Odaka, M.; Fujii, K.;
Hoshino, M.; Noguchi, T.; Tsujimura, M.; Nagashima, S.; Yohda, M.;
Nagamune, T.; Inoue, Y.; Endo, I. J. Am. Chem. Soc. 1997, 119, 3785-
3791. (j) Popescu, V.-C.; Mu¨nck, E.; Fox, B. G.; Sanakis, Y.; Cummings,
J. G.; Turner, I. M., Jr.; Nelson, M. J. Biochemistry, 2001, 40, 7984-
7991.
* To whom correspondence should be addressed. E-mail: kovacs@
chem.washington.edu.
† University of Washington.
‡ Haverford College.
§ Staff crystallographer.
(1) Clark, J. H. Pure Appl. Chem. 2001, 73, 103-111.
(2) Faber, K.; Patel, R. Curr. Opin. Biotechnol. 2000, 11, 517-519.
(3) Kobyashi, M.; Nagasawa, T.; Yamada, H. Tibtech 1992, 10, 402-408.
9
10.1021/ja012555f CCC: $22.00 © 2002 American Chemical Society
J. AM. CHEM. SOC. 2002, 124, 11417-11428
11417

A R T I C L E S
Shearer et al.
Figure 1. (a) Active-site of NHase as determined by X-ray crystallography
(2.65 Å resolution). (b) Active site of NO-inactivated NHase as determined
by X-ray crystallography (1.7 Å resolution).
Figure 2. Spartan figure comparing the structures of [Fe(III)(S2Me2N3-
(Pr,Pr))]+ (1) and [Fe(III)(S2Me2N3(Et,Pr))]+ (2).
Scheme 1. Three Proposed Mechanisms for Nitrile Hydration by
NHase: (1) Involving the Metal Site Serving as a Basic Hydroxide
Source, (2) Involving the Metal Site Serving as a Nucleophilic
Hydroxide Source, or (3) Involving the Metal Ion Acting as a Lewis
Acid and Activating a Metal-Coordinated Nitrile
upon the introduction of nitriles to NHase.6 One could interpret
this to mean that nitriles coordinate to the metal ion; however,
it is also possible that the nitrile indirectly perturbs the metal
ion by slightly distorting its structure when it “wedges” itself
in the active-site pocket. There is also literature precedent for
mechanism 3.7,8 Ford and co-workers have shown8 that by
coordinating nitriles to a metal ion, nitrile hydrolysis rates are
enhanced by as much as 106 relative to that in the absence of
the metal ion. Metal ions in the +3 oxidation state were shown
to enhance rates more than those in the +2 oxidation state,
suggesting that the metal ion activates the nitrile by acting as a
Lewis acid (ie, via an inductive effect), as opposed to via
π-back-donation into the nitrile π* orbitals. Recently Mascharak
and co-workers described9 a five-coordinate Fe(III)-NHase
model complex which binds hydroxide, water, and pyridine;
this complex would not bind nitriles. On the basis of these
results, Mascharak proposed that mechanism 2 was more likely
to be operative in NHase than the other two.
contain a non redox-active Fe(III) ion ligated by two nitrogen
donors (derived from amides in the peptide backbone), three
cysteinate donors, and either a water or hydroxide (active form),
or an NO (inactive form), in a site trans to a cysteinate (Figure
1). The higher-resolution crystal structure5b of NO-inhibited Fe-
NHase revealed that two of the three cysteinates were oxygen-
ated to either a sulfinate or a sulfenate. It has yet to be proven
definitively whether both of these oxidized cysteinates are
required for activity.
Although the active-site structure of NHase has been defined,
a number of unanswered questions have yet to be adequately
addressed. Most importantly, the mechanism of nitrile hydrolysis
has yet to be determined. Three distinct mechanisms have been
proposed (Scheme 1).5a Two of these (mechanisms 1 and 2)
involve the metal ion serving as a hydroxide source for the
nucleophilic attack at an uncoordinated nitrile, with (mechanism
1), or without (mechanism 2) the involvement of an intervening
water molecule. The third (mechanism 3) involves the metal
ion acting as a Lewis acid in the activation of a coordinated
nitrile for nucleophilic attack by an uncoordinated water. The
last two mechanisms are differentiated from the first, in that
they involve metal-bound intermediates, and require that the
metal ion release product at reasonably fast rates. One might
predict that this would be problematic for “substitution inert”
low-spin Co(III) and low-spin Fe(III). We’ve shown, however,
that trans-thiolates enhance ligand dissociation rates signifi-
cantly, making low-spin Co(III) and Fe(III) reasonably labile.13,41
There is spectroscopic evidence that supports mechanism 3.6
Both the EPR signal and electronic absorption spectrum change
We previously reported a five-coordinate coordinatively
unsaturated NHase model compound [Fe(III)(S2Me2N3(Pr,Pr))]+
(1) containing Fe(III) ligated by cis-thiolates and imine nitrogens
(Figure 2).10 Munck and Nelson recently proposed, on the basis
of low-temperature Mossbauer and EPR experiments, that a five-
coordinate NHase intermediate is formed via the flash photolysis
of NO from the NO-inactivated form of Fe-NHase.11 The
parameters of these spectra compare well with those of 1.12
Compound 1 was shown to bind azide10a (a NHase inhibitor)
and NO10b (a NHase inactivator) trans to a thiolate. The azide-
bound complex reproduces the key magnetic and spectroscopic
1
features of NHase including spin state (S ) /2), EPR, low-
energy LMCT band in the electronic absorption spectrum (λmax
≈ 700 nm), and average metal-ligand bond lengths.10a How-
ever, it appeared that 1 would not bind either the substrate or
product of NHase. To increase the reactivity of [Fe(III)(S2Me2N3-
(Pr,Pr))]+ (1), and yet preserve the coordination environment
about the metal center, one methylene group was removed from
the ligand backbone to afford [Fe(III)(S2Me2N3(Et,Pr))](PF6)
(2).13 This caused the structure to distort, from approximately
(7) Kim, J. H.; Britten, J.; Chin, J. J. Am. Chem. Soc. 1993, 115, 3618-3622.
(8) Zanella, A. W.; Ford, P. C. Inorg. Chem. 1975, 14, 42-47.
(9) Noveron, J. C.; Olmstead, M. M.; Mascharak, P. K. J. Am. Chem. Soc.
2001, 123, 3247-3249.
(10) (a) Ellison, J. J.; Niestedt, A.; Shoner, S. C.; Barnhart, D.; Cowen, J. A.;
Kovacs, J. A. J. Am. Chem. Soc. 1998, 120, 5691-5700. (b) Schweitzer,
D.; Ellison, J. J.; Shoner, S. C.; Lovell, S.; Kovacs, J. A J. Am. Chem. Soc.
1998, 120, 10996-10997.
(11) Popescu, V.-C.; Munck, E.; Fox, B. G.; Sanakis, Y.; Cummings, J.; Turner,
I. M., Jr.; Nelson, M. J. Biochemistry 2001, 40, 7984-7991.
(12) Krebs, C.; Pereira, A. S.; Tavares, P.; Huynh, B. H.; Schweitzer, D.; Ellison,
J. J.; Kovacs, J. A. Manuscript in preparation.
(6) Sugiura, Y.; Kuwahara, J.; Nagasawa, T.; Yamada, H. J. Am. Chem. Soc.
1987, 109, 5848-5850.
9
11418 J. AM. CHEM. SOC. VOL. 124, NO. 38, 2002

Reactivity of Fe(III)−NHase Model Complex
A R T I C L E S
LMCT band centered near 800 nm to grow in. In this manner, the
reactivity of 1 and 2 with a variety of ligands was examined.
Thermodynamics of Nitrile and Pyridine Binding. Equilibrium
constants for neutral ligands binding to 1 and 2 were determined in
the manner previously described for anionic ligands.10a Methylene
chloride solutions of 1 and 2 were prepared and then diluted with a
methylene chloride solution containing the ligand (acetonitrile, benzo-
nitrile, iso-propionitrile, or pyridine), so as to afford a final concentration
of ligand-bound complex of 0.300 mM. Ligand concentrations were
selected so that between 20 and 80% of the ligand-bound complex
would result. Extinction coefficients for the nitrile complexes [Fe(III)-
(S2Me2N3(Et,Pr))(NCR)]+ (2-NCR) were determined at -90 °C using
a 4000-fold excess of nitrile, and were found to vary insignificantly
with varying R groups. All binding measurements for 2-NCMe and
2-NCPh were performed at 546, and 839 nm, respectivelysthe wave-
lengths which gave the most consistent results. Extinction coefficients
for [Fe(III)(S2Me2N3(Pr,Pr))(py)]+ (1-py) and [Fe(III)(S2Me2N3(Et,Pr))-
(py)]+ (2-py) were determined at -90 °C using a 4000-fold excess of
pyridine. Equilibrium constants were measured at 656 nm for 1-py and
at 554 nm for 2-py, over the temperature range -40° to -90 °C.
Thermodynamic parameters were then obtained from van’t Hoff plots.
NMR Exchange Kinetics. Kinetic parameters for nitrile exchange
from 2 were measured using NMR line-width analysis over the
temperature range of 190-300 K in CD2Cl2. The line-width of the R
carbon on the nitrile, with and without the metal complex, (∆νcplx and
∆νsolv), are related to the total transverse (spin-spin) relaxation rate,
T2P-1, by eq 1:
trigonal bipyramidal toward a more square pyrimidal structure,
and increased the complex’s reactivity toward a number of
ligands. For example, (Et,Pr)-ligated 2 was shown to have an
∼10-fold greater affinity for azide relative to 1.13 Complex 2
was also found to bind a variety of other ligands, including
nitriles. Herein, we describe the reactivity of 2 with a variety
of ligands, and examine the kinetics and thermodynamics of
nitrile binding.
Experimental Section
General Synthetic Methods. All manipulations were performed
under an atmosphere of dinitrogen using either glovebox or standard
Schlenk techniques. Chemical reagents purchased from commercial
vendors were of the highest purity available and used without further
purification. Solvents were dried using standard procedures and
rigorously degassed prior to use. Labeled Me13CN and CD2Cl2 were
both obtained from Cambridge Isotope Labs (Andover, MA). [Fe(III)-
(S2Me2N3(Pr,Pr))]+ (1), [Fe(III)(S2Me2N3(Et,Pr))]+ (2), and [Fe(III)(SMe2N5-
(tren)(MeCN))]2+ (3) were all prepared as previously described.10,13-14
Elemental analyses were performed by Atlantic Microlabs (Norcross,
GA).
General Instrumental Methods. Solution-state magnetic moments
were determined using the Evans’ method as modified for supercon-
ducting solenoids, on a Bruker AF300 FTNMR spectrometer.15 EPR
spectra were recorded on an X-band Bruker EPX CW EPR spectrometer
at 130 K in a CH2Cl2/toluene glass, unless otherwise stated. XAS data
refinements were performed using Igor Pro 3.15.16 NMR spectra were
recorded on a Bruker DPX 500 and DRX 750 FTNMR and are
referenced to an external standard of TMS. The temperature of the
spectrometer was determined using Van Geet’s method.17 Peaks were
fit to a Lorenzian curve using the NMR data analysis package NUTS.18
Electronic absorption spectra were recorded on a Hewlett-Packard
8452A diode array spectrophotometer. Variable-temperature electronic
absorption spectra were recorded in a custom-built low-temperature
copper-block fitted with a heater, controlled via a thermocouple, and
inserted into a stainless steel cryostat.
Preparation of [Fe(III)(S2Me2N3(Et,Pr))(NCS)] (2-NCS). Tetrakis-
methylammonium thiocyanate (410 mg, 3.10 mmol) was added to a
MeCN solution (25 mL) of 2 (800 mg, 1.55 mmol), and the solution
was stirred for 0.5 h. Diethyl ether (25 mL) was added to precipitate
out less soluble impurities. The solids were filtered through a bed of
Celite and the volatiles evaporated to dryness under reduced pressure.
The resulting red residue was extracted with THF, filtered, and
evaporated to dryness, affording 300 mg of 2-NCS as a red powder
(45% yield). Electronic absorption spectrum (CH2Cl2): λmax (ꢀM) 478
(1350), 833 (1020). EPR: g ) 2.183, 2.164, 2.004. IR (KBr pellet) ν
(cm-1): 2058 (thiocyanate). Anal. Calcd for FeC16H29N4S3: C, 44.75;
H, 6.81; N, 13.05. Found: C, 45.28; H, 6.93; N, 12.67.
Ligand Binding Studies. The reactivity of 1 and 2 with neutral
and anionic ligands was screened by X-band EPR at 150 K. All EPR
spectra were recorded in a toluene/methylene chloride glass (1:1) using
a 1000-fold excess of ligand relative to metal complex. Ligand binding
causes the distorted pseudorhombic EPR signal associated with 1 and
2 to become more rhombic and the g-spread to increase. Once ligands
were identified which bind to these metal centers at 150 K, they were
then screened for reactivity at higher temperatures (between -90 °C
and ambient temperature) using variable-temperature electronic absorp-
tion spectroscopy. Ligand binding to 1 and 2 causes a low-energy
-1
(T2PPM)-1 ) π(∆νcplx - ∆νsolv)PM
(1)
where PM is the partial molar fraction of nitrile bound to the metal
complex. (T2PPM)-1 can also be expressed as:
2
T2M-2 + (T2Mτm)-1 + ∆ωM
-1
-1
(T2PPM)-1 ) τm
+ T2O
(2)
2
(T2M-1 + τm-1)2 + ∆ωM
-1
where T2O is the transverse relaxation rate of the nitrile in the
outersphere, ∆ωM is the chemical shift difference between the
-1
coordinated and uncoordinated nitrile, τm is the resident time of the
-1
nitrile ligand on the metal complex, and T2M is the transverse
relaxation rate due to the coordinated nitrile which is determined by:
2
S(S + 1)
2πA
h
1
-1
T2M
)
(3)
(
)
-1
Tle-1 + τm
3
where S is the spin of the metal complex with the sixth ligand bound,
2πA/h is the hyperfine coupling constant between the unpaired electron
and the R-carbon (which was assigned as 10 MHz over the temperature
-1
range investigated) and T1e is the electronic longitudinal relaxation
-1
rate. T1e was determined as previously described by Yamamoto et
al.19 over the temperature range investigated. The temperature depen-
2
dence of ∆ωM is given by:
geµB 2πA
S(S + 1)
3kBT
∆ωM ) ω1
(4)
(
)
γ1
h
where ω1 is the resonance frequency of the carbon nucleus, ge is the
g-value for a free electron, µB is the Bohr magneton, γ1 is the
gyromagnetic ratio for carbon, kB is the Boltzmann’s constant, and T
is temperature. The exchange rate (kex) was then determined by:
(13) Schweitzer, D.; Shearer, J.; Rittenberg, D. K.; Shoner, S. C.; Ellison, J. J.;
Loloee, R.; Lovell, S.; Barnhart, D.; Kovacs, J. A. Inorg. Chem. 2002, 41,
3128-3136.
(14) Shearer, J.; Nehring, J.; Lovell, S.; Kaminsky, W.; Kovacs, J. A. Inorg.
Chem. 2001, 40, 5483-5484.
τm-1 ) kex
(5)
(15) Grant, D. H. J. Chem. Educ. 1995, 72, 39-40.
(16) Igor Pro, 3.15; Wavemetrics, Inc.: Lake Oswego, OR, 2000.
(17) Van Geet, A. L. Anal. Chem. 1968, 40, 2227-2229.
(18) NUTS; Acorn NMR, Inc.: Livermore, CA, 1997.
(19) Yamamoto, Y.; Nanai, N.; Chuˆjoˆ, R. Bull. Chem. Soc. Jpn. 1991, 64, 3199-
3201.
9
J. AM. CHEM. SOC. VOL. 124, NO. 38, 2002 11419

A R T I C L E S
Shearer et al.
4-tert-Butylpyridine was used to determine pyridine exchange rates from
1 and 2. In this case there is negligible unpaired spin density on the
tert-butyl group carbons so that the exchange rate can be determined
by:
scattering theory. Our interest was primarily in the Fe-ligand bond
lengths and disorder, and these parameters were virtually the same in
the Fourier-filtered and -unfiltered fits. The least-squares error index
ꢀ2 is defined as:
T2PPM-1 ) π(∆νcplx - ∆νsolv)PM-1 - T2O ) kex
(6)
ꢀ2 ) [nidp(nidp - np)] × average{[(ydata - ycalc)/σdata]}2 (10)
The parameters ∆Hq and ∆Sq and Ea were then determined through
standard procedures.20
where nidp is the number of independent25 points in the data (ydata
)
Fourier-filtered k3ø), np is the number of refined parameters, and σdata
is the uncertainty attributed to the (Fourier-filtered) XAS spectrum,
which is linearly interpolated from (k‚Å, σdata‚Å3) ) (2.2, 0.19), (4.0,
0.26), (8.0, 0.26), (10.0, 0.42), (13.0, 0.88), (14.3, 0.56).33a Using criteria
established by the International Workshops on Standards and Criteria
in XAFS,25 any model that gives shell-specific ꢀs2 values within 1 of
the minimum achieved value should be considered a model consistent
with the EXAFS data, although models with lower ꢀs2 values may be
considered more likely. The reported parameter uncertainties were
determined in the recommended manner by fixing the parameter in
question at a series of values and refining all other parameters (that
were refined in the original fit) to determine the variation that causes
an increase in ꢀ2 by 1. The allowed variation range was generally
slightly asymmetric; the larger of the allowed positive or negative
variation from the best-fit value is reported as the ( value.
XAS Sample Preparation and Data Collection. Complex 2 was
dissolved in either CH2Cl2, MeCN, or MeOH (5-25 mg in ∼300 µL),
injected into an aluminum sample holder between two pieces of
translucent electrical tape (3M, no. 1205, Minneapolis, MN), and
quickly frozen in liquid nitrogen. Data were collected at the National
Synchrotron Light Source (Brookhaven National Laboratories, Upton,
NY) on beamline X-9B using a focused Si(111) double-crystal
monochromator, a low-angle nickel mirror for harmonic rejection, and
a 13-element Ge solid-state fluorescence detector (Canberra). A helium
Displex cryostat was used to keep the samples at a constant 77 K
throughout the data collection. X-ray energies were calibrated by
simultaneous measurement of the absorption spectrum of Fe foil (first
inflection point assigned to 7111.2 eV).21 The spectra were measured
in 5 eV increments in the preedge region (6960-7100 eV), 0.5 and
1.5 eV increments in the edge region (7100-7132 and 7132-7171
eV), and 3 eV increments in the EXAFS region (7171-7910 eV).
XAS Data Analysis. The XAS baseline was normalized using a
three-region cubic-spline function. The EXAFS region of the XAS was
plotted against the wavevector k:
Bond valence sums analysis (BVS) of the EXAFS data was
performed according to eq 11:
BVS )
sij
(11)
(12)
∑
where sij is the valence of an individual bond described by:
sij ) exp[(Ro - Rij)/0.37]
k ) π[8me(E - Eo)]1/2/h
(7)
where Eo ) 7125 eV. Simulations used the single-scatterer (SS) EXAFS
equation:
where Ro is the length of a valence unit and Rij is the experimentally
determined bond length. The values used for Ro are: 1.759 Å for Fe-
O, 1.831 Å for Fe-N, and 2.151 Å for Fe-S.33a
#shells
øcalc
)
ni fiki-1ri-2 exp(-2σi2ki2) sin(2kiri + Ri)
(8)
∑
X-ray Crystallography of 2-NCS. X-ray-quality crystals were
grown by slowly diffusing diethyl ether into a THF solution of 2-NCS
at -35 °C. A suitable single crystal was immersed in Paratone 8277
oil, mounted on a glass capillary, and immediately placed under a stream
of cold dinitrogen gas. X-ray data was collected at 130 K using a Nonius
Kappa CCD diffractometer. The structure was solved using direct
methods and refined using SHELEX 97.26 All hydrogen atoms were
placed in an idealized geometry and refined using a riding model.
Crystallographic data for 2-NCS are listed in Table 1.
i)1
where ni is the number of scatterers in the ith shell, ri is their average
distance from the metal center, σi is the variability (disorder) in that
shell, and ki is defined as:
ki ) π[8me(E - (7125 eV + ∆Ei))]1/2/h
(9)
where ∆Ei is the difference in energy between the nominal edge and
the true ionization energy of that shell. This parameter fixed at ∆Ei )
-0.5 eV based on previous studies of similar complexes with known
structures.22 The amplitude and phase functions, generated by FEFF
7.0123 simulations, were the same as we used previously.22 We
determined the baseline and edge height from simulations of the entire
XAS (including Gaussian simulations of preedge peaks and first-
coordination sphere fits to the EXAFS region) and then performed
subsequent fits only on the EXAFS region to investigate alternative
coordination models and the effects of including outer sphere scatterers.
For fits that we report in this paper, we used weighted Fourier-filtered
fits of k3ø (FT: 1.0-14.3 Å; BT: r′ ) 0.8-2.0 Å).33a We also
performed (but do not report) fits to unfiltered k2ø in which the outer-
sphere atoms were modeled as Fe-C shells at ca. 2.9, 3.4, and 4 Å;
only the distances from the first of these shells corresponded to distances
seen in crystal structures, while the longer distances probably were
artifacts due to modeling multiple-scattering EXAFS using single-
Results and Discussion
Reactivity of Five-Coordinate [Fe(III)(S2Me2N3(Pr,Pr))]+
(1) and [Fe(III)(S2Me2N3(Et,Pr))]+ (2) toward a Sixth Ligand.
Recent calculations by Richards and co-workers indicate that
the metal-amide bond in NHase has significant double-bond
character between the nitrogen and carbon, and is best repre-
sented as an imido-metal bond (Scheme 2).27 This is supported
by modeling studies done in our group which show that the
spectroscopic and electronic properties of NHase can be nicely
reproduced with six-coordinate metal complexes containing
imine nitrogens and cis-thiolates.28 [Fe(III)(S2Me2N3(Pr,Pr))]+
(24) Ankudinov, A. L.; Ravel, B.; Rehr, J. J.; Conradson, S. D. Phys. ReV. B.
1998, 58, 7565-7576.
(25) Bunker, G., Hasnain, S., Sayers, D. In X-ray Absorption Fine Structure;
Hasnain, S. S., Ed.; Ellis Horwood: New York, 1991; pp 751-770.
(26) (a) Otinowski, Z.; Minor, W. Methods Enzymol. 1996, 276, 307-326. (b)
Blessing, R. H. Acta Crystallogr. Sect. A 1995, 51, 33. (c) Altomare, A.;
Cascarano, G.; Giacovazzo, C.; Burla, M. C.; Polidori, G.; Camalli, M. J.
Appl. Crystallogr. 1994, 27, 435-442. (d) Sheldrick, G. M. SHELEX97
University of Gottingen.
(20) Moore, J. W.; Pearson, R. G. Kinetics and Mechanism: A Study of
Homogeneous Chemical Reactions, 3rd ed; John Wiley & Sons: New York,
1981; Chapter 5.
(21) Bearden, J. A.; Burr, A. F. CRC Handbook, 60th ed.; Weast, R. C., Astle,
M. J., Eds.; CRC Press: Boca Raton, Fl. 1979. p E-193.
(22) Scarrow, R. C.; Stickler, B. S.; Ellison, J. J.; Shoner, S. C.; Kovacs, J. A.;
Cummings, J. C.; Nelson, M. J. J. Am. Chem. Soc. 1998, 120, 9237-9245.
(23) Ankudinov, A. L.; Rehr, J. J. Phys. Rev. 1997, B56, R1712-R1715.
(27) Boone, A. J.; Cory, M. G.; Scott, M. J.; Zerner, M. C.; Richards, N. G. J.
Inorg. Chem. 2001, 40, 1837-1845.
9
11420 J. AM. CHEM. SOC. VOL. 124, NO. 38, 2002

Reactivity of Fe(III)−NHase Model Complex
A R T I C L E S
Table 1. Crystallographic Data for [Fe(III)S2Me2N3(Et,Pr)(SCN)]
(2-NCS)
Table 2. Ligand Binding to [Fe(III)(S2Me2N3(Et,Pr))]+ (2) Monitored
by Low-Temperature UV/Vis Spectroscopy and EPR
2-NCS
low-T electronic
spectrum, λmax
a
b
2-L
EPR
formula
molecular mass
crystal color, habit
T (K)
crystal system
space group
a (Å)
C16H29FeN4S3
429.48
plate, dark-red
130(2)
orthorhombic
Pbca (No. 61)
10.8846(5)
15.6740(5)
23.5810(10)
90
2
2.12, 2.07, 2.02
2.17, 2.14, 2.01
2.17, 2.13, 2.01
2.16, 2.13, 2.01
2.17, 2.13, 2.01
2.17, 2.15, 2.00
2.18, 2.15, 1.98
2.18, 2.16, 2.00
2.15, 2.10, 2.01
2.18, 2.15, 2.00
2.17, 2.14, 2.02
2.18, 2.16, 2.00
2.16, 2.12, 2.00
2.16, 2.00
-
2-MeCN
2-EtCN
2-tBuCN
2-PhCN
2-py
830 nm
830 nm
813 nm
845 nm
765 nm
760 nm
c806 nm
c877 nm
753 nm
742 nm
d
-
b (Å)
c (Å)
2-N3
2-SCN-
2-tBuNC
2-MeOH
R (deg)
â (deg)
90
90
γ (deg)
2-EtOH
V (Å3)
4023.0(3)
8
1.418
2-CH3C(O)NH2
2-CH3C(O)CH3
2-CH3SC(NH)CH3
Z
F
745 nm
d
calcd (g cm-3
)
µ(Mo KR) (mm-1
Θ range (deg)
)
1.067
2-26
a X-Band at 150 K; CH2Cl2/toluene. b In CH2Cl2 solution at -90 °C.
Only the low-energy band (above 700 nm) is reported. c Binds at room
temperature in CH2Cl2. d Insoluble at temperatures below -20 °C.
collected reflections
unique reflections
number of parameters
GOFa on F2
7342
3938
227
0.885
3.94 (7.85)
R1b (wR2c) (%)
2
2
a GOF ) {∑[w(Fo - Fc )2]/(M - N)}1/2 (M ) number of reflections,
N ) number of parameters refined). b R1 ) ∑||Fo| - |Fc||/∑|Fo| for
2
2
2
I > 2σ(I). c wR2 ) {∑[w(Fo - Fc )2]/∑[w(Fo )2]}1/2 using all data.
Scheme 2. Metal Amide Resonance Form Showing Relation to
Metal Imine
(1) was previously shown to reversibly bind azide; however, it
appeared to be unreactive toward neutral ligands such as MeCN
and alcohols, on the basis of ambient temperature electronic
absorption studies.10 Reexamination of the low-temperature EPR
spectrum of 1 in a noncoordinating solVent mixture (CH2Cl2/
toluene glass) indicated that this may not necessarily be the case.
When the EPR spectrum of 1 is recorded in a rigorously dried
CH2Cl2/toluene (1:1) glass (Figure S-19), a distorted pseudo-
rhombic EPR spectrum is observed (g ) 2.15, 2.07, 2.00), which
differs from that of 1 in a MeOH/EtOH glass (2.20, 2.15, 2.00).
In the noncoordinating solvent, a less rhombic signal, with a
narrower g-spread, is observed. Addition of MeOH or EtOH to
this CH2Cl2/toluene solution converted the g ) 2.15, 2.07, 2.00
signal to a more rhombic signal (2.20, 2.15, 2.00), identical to
that of 1 in MeOH/EtOH. This suggested that alcohols bind to
the metal center of 1 at 150 K. In this manner, it was determined
that in CH2Cl2/toluene solution, a wide variety of neutral ligands
(including nitriles), as well as SCN-, bind to 1 at extremely
low temperatures (e150 K).
Figure 3. X-band EPR spectrum of [Fe(III)(S2Me2N3(Et,Pr))]PF6 (2) in
CH2Cl2/toluene (1:1) glass, both in the absence and presence of MeCN.
with azide than 1 and that azide dissociates an order of
magnitude slower from the resulting complex 2-N3 versus 1-N3.
It was therefore not surprising that 2 was also found to bind a
wider variety of ligands at higher temperatures relative to 1.
Ligand binding was screened using EPR (Table 2), in a CH2Cl2/
toluene glass at 150 K. As was seen with 1, the distorted pseudo-
rhombic EPR spectrum of 2 converts to the more rhombic
spectrum of six-coordinate 2-L upon the addition of Lewis bases
(L) in noncoordinating solvents. This is illustrated in Figure 3
for L ) MeCN.
EPR screening of reactivity was performed in frozen glasses
at 150 K, whereas enzymes function at considerably higher
temperatures, and in solution. To examine the reactivity of 1
and 2 under more relevant conditions, reactivity was screened
in solutions of methylene chloride over the temperature range
-90-0 °C. This is where the major differences in reactivity
between 1 and 2 become apparent. For both complexes, as a
sixth ligand binds to the metal, a low-energy charge-transfer
band centered near 800 nm grows in (Figure 4). A similar
charge-transfer band is also observed in the active form of Fe-
NHase.5,6 Under these conditions (methylene chloride, -90 °C),
1 does not appear to bind any nitriles or alcohols, and the only
amine it will bind under these conditions is pyridine. With
anionic ligands such as azide or thiocyanate, on the other hand,
only 1 equiv of ligand is required to observe complete binding
to 1 (and 2) at -90 °C in CH2Cl2.
Recently,13 we showed that the removal of a methylene unit
from the ligand backbone of 1 to afford [Fe(III)(S2Me2N3-
(Et,Pr))]+ (2), results in a more distorted structure (Figure 2,
τ ) 0.52 (2) vs τ ) 0.76 (1)).29 This subtle distortion in struc-
ture altered the magnetic properties and increased the reactivity
of 2 relative to that of 1.13 It was determined that 2 reacts faster
(28) (a) Shoner, S. C.; Barnhart, D.; Kovacs, J. A. Inorg. Chem. 1995, 34, 4517-
4518. (b) Jackson, H. L.; Shoner, S. C.; Rittenberg, D.; Cowen, J. A.; Lovell,
S.; Barnhart D.; Kovacs, J. A. Inorg. Chem. 2001, 40, 1646-1653.
(29) τ is defined as (R-â)/60, where R ) largest angle, â ) second largest
angle (τ ) 1.0 for ideal trigonal bipyramidal; τ ) 0.0 for ideal square
pyramidal). Addison, A. W.; Rao, T. N.; Reedijk, J.; Van Rijn, J.;
Verschoor, G. C. J. Chem. Soc., Dalton Trans. 1984, 1349-1356.
9
J. AM. CHEM. SOC. VOL. 124, NO. 38, 2002 11421

A R T I C L E S
Shearer et al.
actiVe-site of NHase. However, this would require that the
hydroxide, shown (by ENDOR) to be coordinated to the iron
site,30 be protonated prior to substrate binding. NHase activity
has been shown to decrease5e,31 as the pH is raised above pH
) 7.3 consistent with this possibility. Mascharak recently
reported that the pKa of water bound to Fe(III) in a ligand
environment similar to 2 is around 6.8.9 Furthermore, he reports
an extremely large -∆H value for water binding to this
compound. When the results described herein are taken into
account, it is conceivable that Mascharak is dealing with a
hydroxide-bound, as opposed to water-bound species. However,
it is also conceivable that the amide ligands make his system
sufficiently electron-rich to raise the pKa of his Fe(III)-OH to
the point where the protonated form is favored. With 2,
hydroxide attacks the carbon of the imine bond in the ligand
backbone, rather than the metal, and water does not compete
adequately enough with the coordinating solvents required to
solubilize 2 (acetone, MeCN, DMF) to generate significant
quantities of water-bound 2-H2O, even at low temperatures. The
spectrum of 2 in a MeCN/H2O (75:25) is identical to that of 2
in MeCN. In noncoordinating solvents, such as CH2Cl2, the
concentration of water is too low to favor the six-coordinate,
water-bound form.
Figure 4. Temperature-dependent electronic absorption spectrum of
[Fe(III)(S2Me2N3(Et,Pr))]PF6 (2) + MeCN (5000 equiv) in CH2Cl2. As
the temperature is lowered, two bands corresponding to [Fe(III)(S2Me2N3-
(Et,Pr))(MeCN)]PF6 (2-MeCN) appear at 454 and 833 nm, and a shoulder
at 544 nm, corresponding to 2, disappears. The ratio of MeCN bound to
five-coordinate 2 changes from 55% at -60 °C, to 69% at -70 °C, 77%
at -75 °C, 83% at -80 °C, 90% at -85 °C, and 100% at -90 °C.
In contrast to 1, 2 binds a wide variety of “substrates”
including nitriles (benzonitrile, acetonitrile, iso-propionitrile) and
alcohols (methanol, ethanol, propanol), as well as pyridine,
acetamide, acetone, and thiocyanate in methylene chloride
solutions at -90 °C (Table 2). Nitrile binding is reversible, and
an excess of nitrile is required to observe complete binding.
For example, in a methylene chloride solution of 2, 5000 equiv
of MeCN are required to observe complete binding at -90 °C.
As the temperature is increased, the concentration of five-
coordinate 2 increases relative to that of six-coordinate 2-MeCN
(Figure 4). When 2 is dissolved in neat MeCN, 12% of MeCN-
bound 2-MeCN is detected at ambient temperature. In contrast,
<1% of MeCN-bound 1-MeCN is detected in neat MeCN,
even at temperatures as low as -46 °C. This illustrates how
dramatically the affinity for nitriles increases upon removal of
a methylene unit from (Pr,Pr)-ligated 1 and is yet another
indication of the inherent increase in reactivity of 2 compared
with that of 1.
(Et,Pr)-ligated 2 has a much lower affinity for alcohols
relative to nitriles. Complete binding of methanol to 2 is not
observed even in neat methanol just above its freezing point
(-98 °C). In methanolic solutions, nitriles preferentially bind
to 2, even when there is a large excess of MeOH present.
Competitive binding studies show that the equilibrium constant
for 2-MeOH formation is an order of magnitude smaller than
Keq for 2-MeCN formation. When 2 is dissolved in a solvent
consisting of an 11:1 MeOH/MeCN mixture, a broad absorption
band (width at half-height ) 205 nm) is observed, which is
red-shifted (789 nm) almost halfway between the λmax for
2-MeOH (753 nm) and λmax for 2-MeCN (830 nm). This
indicates that roughly a 1:1 ratio of 2-MeCN:2-MeOH forms
when an 11-fold excess of MeOH is present, indicating that
Keq(2-MeCN) is roughly an order of magnitude larger than
Keq(2-MeOH). Given that water is a poorer Lewis base than
MeOH, this result implies that Keq(H2O) would be more than
an order of magnitude smaller than Keq(MeCN). Proposed
mechanism 3 would require that a nitrile displace water in the
first step of nitrile hydrolysis by NHase. Our results suggest
that a nitrile would be capable of displacing water from the
Synthesis and Structure of [FeIII(S2Me2N3(Et,Pr))] (2-NCS).
As was observed with 1, anionic ligand binding is more favored
than neutral ligand binding to 2 in methylene chloride, or
-
acetonitrile. Only 1 equiv of N3 or SCN-, for example, is
required to observe complete binding to 2 at -90 °C. This
encouraged us to prepare a thiocyanate-bound derivative of 2
(2-NCS). Having a structurally characterized thiocyanate com-
plex would aid in the analysis of EXAFS data for nitrile-bound
2-MeCN (vide infra), by providing an example of a linearly
coordinated ligand that might be involved in multiple scattering.
Thiocyanate-ligated 2-NCS was synthesized by adding 2
equiv of thiocyanate to a MeCN solution of 2 at ambient
temperature. The νSCN stretch at 2058 cm-1 (vs 2107 cm-1 for
free thiocyanate) in the IR demonstrated that SCN- was
coordinated to the metal. Single crystals (red plates) of 2-SCN
were grown from THF/diethyl ether. The solution-state magnetic
moment (2.31 µB (in MeCN) and EPR parameters (g ) 2.183,
2.164, 2.004) indicate that 2-SCN is low-spin (S ) 1/2) at both
ambient temperature and 150 K, respectively. In CH2Cl2, 2-NCS
displays a LMCT band at 833(1020) nm in its electronic
absorption spectrum. This band is not observed with the five-
coordinate parent compound 2. If crystalline 2-NCS is dissolved
in methanol at ambient temperature, this LMCT band disappears,
and the spectrum reverts back to that associated with five-
coordinate 2, indicating that thiocyanate binds reversibly under
these conditions. To unambiguously confirm the structure of
2-NCS an X-ray crystal structure was determined.
The X-ray crystal structure shows that the Fe center of 2-NCS
is contained in a distorted octahedral coordination environment
with a nearly linear N-bound thiocyanate (Fe-N(4)-C(16) )
178.5°) coordinated trans to a thiolate (Figure 5). Substrates
bind in a similar fashion, that is, trans to a cysteinate sulfur,
in NHase. The CdN bond length of the bound thiocyanate
(1.172(4) Å) is slightly longer than that of free thiocyanate
(30) Doan, P. E.; Nelson, M. J.; Jin, H.; Hoffman, B. M. J. Am. Chem. Soc.
1996, 118, 7014-7015.
(31) Cramp, R. A.; Cowan, D. A. Biochim. Biophys. Acta 1999, 1431, 249-
260.
9
11422 J. AM. CHEM. SOC. VOL. 124, NO. 38, 2002

Reactivity of Fe(III)−NHase Model Complex
A R T I C L E S
Figure 5. ORTEP diagram of thiocyanate-bound [Fe(III)(S2Me2N3(Et,Pr)-
(NCS)] (2-NCS), showing 50% probability ellipsoids and atom-labeling
scheme. All H atoms, except for the N(2)-H proton, have been omitted
for clarity.
(1.15 Å). The average Fe-S bond length (2.214(9) Å) in 2-NCS
is ∼0.1 Å longer than that of 2 (2.117(1)). This increase in
bond length is the expected result of increasing the coordination
number and is similar to that observed upon coordination of
azide to 110 and 2.13 The average Fe-S bond length in these
model compounds compares well with the average EXAFS-
determined Fe-S distance (2.20 Å) in the active form of Fe-
NHase. The average Fe-N bond lengths (1.986 (3) Å) in 2-NCS
also compare well with the EXAFS-determined average Fe-N
bond lengths found in Fe-NHase (1.99 Å).
Fe K-Edge X-ray Absorption Spectroscopy (XAS) of
2, [FeIII(S2Me2N3(Et,Pr))(MeCN)](PF6) (2-MeCN), and
[FeIII(S2Me2N3(Et,Pr))(MeOH)](PF6) (2-MeOH). Compounds
2-MeCN and 2-MeOH would not yield isolable solids; there-
fore, XAS was used to determine the structures of solvent-bound
2. The X-ray absorption spectra of 2 were recorded in frozen
solutions of methylene chloride, acetonitrile, and methanol. The
preedge region of the spectra of 2 in all three solutions is
dominated by an absorbance at 7112 eV that is attributed to a
1s f 3d transition (Figure 6). Upon dissolution of 2 in the
coordinating solvents acetonitrile and methanol, the area under
this peak decreases by a factor of 2 when compared to that when
2 is dissolved in the noncoordinating solvent methylene chloride.
Que and co-workers have previously demonstrated that the size
of this transition is an indication of the degree of symmetry
about the metal center.32 In complexes with centrosymmetric
metal environments, such as square planar or octahedral, the
area of the 1s f 3d transition is less than that of the area for
noncentrosymmetric five-coordinate environments. The results
from analysis of the preedge indicates that the Fe(III) center of
2 dissolved in acetonitrile and methanol is more centrosymmetric
than that of 2 dissolved in methylene chloride, consistent with
the addition of a sixth ligand from solvent in acetonitrile and
methanol, but not in the methylene chloride solution.
Figure 6. 1s f 3d transition found in the XAS of 2 (a), 2-MeOH (b), and
2-MeCN (c). Relative heights for all spectra are identical. Data are depicted
as circles (ooo), fit of data is depicted as the solid line (s), and the
deconvolution of the preedge into an edge shape and two Gaussian peaks
is depicted as the dashed line (- - -) The refined peak positions and areas
(in units of eV × fractional edge height) are indicated.
solution compare with the average bond lengths previously
determined for crystalline five-coordinate 1 (1.99 and 2.15 Å)10a
and 2 (1.95 and 2.12 Å).13 The biggest discrepancy between
the solution and crystal structure distances is the average Fe-N
bond length, which is 0.05 Å shorter in the crystal structure.
This difference may be due to distortions caused by crystal
packing or even possibly due to systematic errors in the Fe-N
bond lengths caused by the symmetry-imposed disorder of the
Et and Pr ligand backbones within the structure of 2.13
Bond valence sum analysis (BVS) has previously proven
useful in interpreting the EXAFS results for metal complexes.33
Here the EXAFS determined bond lengths for the different
scatterers are compared with a reference metal-ligand distance
to estimate the valence of each bond. These are then summed,
giving a BVS. For a metal complex in given oxidation- and
spin states, the BVS for the complex should center around a
given value. If a ligand is added to the metal complex, the
average bond lengths about the metal center should increase;
however, the BVS should not dramatically change if there is
no change in spin- or oxidation state. Thus, BVS can provide
good estimates of the oxidation- and spin states of a metal
complex if only the bond lengths are known. If the oxidation-
and spin states for the complex are already known, then BVS
can provide an internal check to determine if the number and
Fourier transforms and simulations of the EXAFS for all three
compounds (Figure 7) show that contributions from the first
coordination sphere dominate the EXAFS. The parameters from
the fits shown in Figure 7 are given in Table 3. The coordination
model assumed complete chelation by the S2Me2N3(Et,Pr) ligand
and the presence of a sixth oxygen or nitrogen from solvent
2-MeOH or 2-MeCN, respectively, based on the XANES
analysis presented above. The calculated bond lengths of 2.00(2)
Å (Fe-N) and 2.12(2) Å (Fe-S) for five-coordinate 2 in CH2Cl2
(33) (a) Scarrow, R. C.; Brennan, B. A.; Cummings, J. G.; Jin, H.; Duong, D.
J.; Kindt, J. T.; Nelson, M. J. Biochemistry 1996, 35, 10078-10088. (b)
Brown, I. D.; Altermatt, D. Acta Crystallogr., Sect. B 1985, 41, 244-247.
(c) Thorp, H. H. Inorg. Chem. 1998, 37, 5690-5692. (d) Garner, C. D.
Physica B 1995, 209, 714-716. (e) Hati, S.; Datta, D. J. Chem. Soc., Dalton
Trans. 1995, 1177-1182. (f) Thorp, H. H. Inorg. Chem. 1992, 31, 1585.
(g) Liu, W.; Thorp, H. H. Inorg. Chem. 1993, 32, 4102-4105. (h) Palenik,
G. J. Inorg. Chem. 1997, 36, 4888-4890. (i) Palenik, G. J. Inorg. Chem.
1997, 36, 3394-3397. (j) Palenik, G. J. Inorg. Chem. 1997, 36, 122.
(32) (a) Randall, C. R.; Shu, L.; Chiou, Y.-M.; Kagen, K. S.; Ito, M.; Kitajima,
N.; Lachicotte, R. J.; Zang, Y.; Que, L., Jr. Inorg. Chem. 1985, 34, 1036-
1039. (b) Roe, A. L.; Schneider, D. J.; Mayer, R. J.; Pyrz, J. W.; Widom,
J.; Que, L., Jr. J. Am. Chem. Soc. 1984, 106, 1676-1681.
9
J. AM. CHEM. SOC. VOL. 124, NO. 38, 2002 11423

A R T I C L E S
Shearer et al.
Figure 7. Unfiltered k3 EXAFS (left, circles), FT over k ) 1.0 to 14.3 Å-1 (right) and first sphere FF using backtransform limits of r′ ) 0.8-2.0 Å (left).
For FT and FF spectra, the thick line is the data, and the thin line shows the fit using parameters from Table 3; the FT magnitude of the difference spectrum
is shown by a line of medium thickness. Data are shown for 2 in CH2Cl2 (a), 2-MeOH (b), and 2-MeCN (c).
Table 3. EXAFS Data Refinements for 2, 2-MeOH, and 2-MeCN
in 2 is supported by the BVS analysis as well as by both
magnetic and EPR data.
2 (in CH Cl2)
2-MeOH
2-MeCN
2
The Fe-N and Fe-S distances determined by EXAFS
analysis lengthen by 0.02-0.06 Å in the six-coordinate 2-MeOH
and 2-MeCN compared with those in five-coordinate 2. Similar
bond distance increases are observed upon comparison of the
crystal structures of the azide adducts and the corresponding
five-coordinate 1 and 2.10,13 The lengthening of bonds is required
in order to keep the BVS within the range expected upon
addition of a sixth ligand.
NN+O
N+O (Å)
N+O (Å)d
3 Na
3 N + 1 Ob
2.033 ( 0.016
0.093 ( 0.025
2 S
2.179 ( 0.021
0.093 ( 0.025
4.07 ( 0.14
2.0
4 N
r
σ
2.000 ( 0.018c
0.075 ( 0.030
2 S
2.017 ( 0.016
0.108 ( 0.019
2 S
2.179 ( 0.010
0.067 ( 0.012
4.28 ( 0.09
0.7
NS
rS (Å)
σS (Å)
BVS
ꢀ2
2.117 ( 0.015
0.070 ( 0.019
4.09 ( 0.13e
2.2
a The number of N and S scatterers was held fixed at the values shown.
b For 2-MeOH, an oxygen scatterer was included with its r and σ constrained
to be the same as for the nitrogen shell. c Uncertainty range is for ꢀ2 no
more than 1.0 greater than the minimum value reported. d σ2, the Debye-
Waller disorder parameter, was refined. The best-fit values and uncertainty
ranges were converted to σ values because the uncertainty ranges were more
symmetrical when expressed in σ rather than σ2. e Uncertainty range
determined by mapping ꢀ2 in a series of refinments in which BVS was
fixed to a value (4.00, for instance) and rN+O and the two σ parameters
were refined, while rS was constrained to be consistent with the fixed BVS.
As noted above, the number of N and S scatterers in the
refinements of Table 3 were fixed at values expected on the
basis of the coordination numbers suggested by XANES. In all
three cases, the quality of fit improves (ꢀ2 is decreased) when
the number of nitrogen scatterers in the model is increased by
one (i.e., from 3 to 4, or from 4 to 5). This would suggest six-
coordination for 2 in CH2Cl2 and seven-coordination for
2-MeOH and 2-MeCN. Despite the better fit to the data, these
simulations yield unprecedentedly high BVS (4.6 to 4.8) and
also high refined disorder factors (σN > 0.10 Å) for the Fe-N
shell. The seven-coordinate models for the methanol and
acetonitrile adducts seem particularly unlikely. Thus, we
discount the suggestion of higher coordination numbers from
the EXAFS fitting and attribute the better fits to slight
inaccuracies in the amplitude and phase functions, an inac-
distance of the different scatterers are correct. When the bond
lengths determined for 2 (Table 3) are analyzed by this method,
the BVS is 4.1 ( 0.1. For comparison, the crystal structure bond
lengths of 1 and 2 give BVS values of 4.0 and 4.4, respectively.
All three of these values fall within the range of 3.7-4.5
previously determined for low-spin Fe3+ complexes (average
value ) 4.1).33a Thus, the existence of a low-spin ferric center
9
11424 J. AM. CHEM. SOC. VOL. 124, NO. 38, 2002

Reactivity of Fe(III)−NHase Model Complex
A R T I C L E S
curately determined E0, or incomplete removal of outer sphere
EXAFS contributions by the Fourier filtering.
differences. Azide anions are very effectively solvated by
methanol, whereas MeCN is not effectively solvated by meth-
ylene chloride. Thus, azide is in an inherently more ordered
state in methanol than is MeCN in methylene chloride,
contributing to a smaller than expected ∆S.
It was previously demonstrated that multiple scattering (MS)
pathways were important in fitting the EXAFS of an Fe-NHase
model containing NO because of the nearly linear Fe-N-O
bond angle (172°). We therefore expected MS pathways to be
crucial in fitting the EXAFS data for 2-MeCN, because of the
expected near linear iron-nitrile bond. To test this expectation,
we used FEFF 7.0223 to simulate single- and multiple-scattering
EXAFS for a structure based on the coordinates of 2-N3, but
with the azide replaced by a linear NCCH3 ligand. From these
simulations we isolated contributions to the EXAFS due only
to the Fe-NCCH3 unit. Results of these simulations are shown
in the Supporting Information (Figure S-20). When only a linear
Fe-NCCH3 unit was considered, peaks in the k3ø FT EXAFS
spectrum occurred at r′ ) 2.8 and 4.2 Å, and these peaks were
about 3 times larger than the magnitude of the k3ø FT EXAFS
for 2-MeCN at the same values of r′. However, as the Fe-
N-C angle was varied from 180° to 164° (a variation found in
acetonitrile complexes in the Cambridge Structure Database),34
the magnitude of the r′ ) 4.2 Å peak was cut in half.
Furthermore, when the coordinates of the S2Me2N3(Pr,Pr) ligand
were added into the simulation, scattering from the C atoms of
the ligand caused destructive interference with the r′ ) 2.8 Å
peak due to the acetonitrile ligand, so that r′ ) 2.8 Å became
a minimum in the FT magnitude. Finally, when the FEFF 7.02
program was asked to simulate vibrational disorder, the
magnitude in both the first coordination sphere (r′ ) 1-2 Å)
and outer sphere (r′ > 2 Å) regions of the FT EXAFS were
reduced so that the magnitudes in the simulations were similar
to those of the data of 2-MeCN. Thus, our expectation of
observing multiple-scattering effects from a single acetonitrile
ligand was not realized, and may not be possible due to
competing single- and multiple-scattering pathways due to the
ligand S2Me2N3(Et,Pr).
Thermodynamics of Ligand Binding. Neutral ligands were
found to bind reversibly to 1 and 2 and induce dramatic changes
to the electronic absorption spectra. Thus, thermodynamic
parameters for ligand binding could easily be obtained spec-
trophotometrically. Equilibrium constants (Keq) for reactions
converting 1 or 2 to 1-L or 2-L were determined (Tables S-8-
S-11) at several different temperatures according to standard
procedures.10,13,35-36 Van’t Hoff plots (Figures S-8-S-11) were
then used to extract the enthalpy and entropy parameters for
the reaction.
Thermodynamic parameters for the binding of neutral ligands
to 1 and 2 were determined as previously described for azide,
except methylene chloride was used in place of MeOH. This
was done to ensure that the added ligand would not have to
compete with a coordinating solvent. Analysis of the temperature
dependence of Keq’s for MeCN binding to 2 using a van’t Hoff
plot showed that MeCN binds with an enthalpy similar to that
of azide (∆H ) -7.5 ( 0.6 kcal/mol). The entropy for nitrile
binding to 2 (∆S ) -30 ( 6 eu) was determined, however, to
be substantially more negative than for azide binding (∆S )
-16.6 ( 2.5 eu). Solvation effects could account for these
To determine if there is any selectivity in nitrile binding to
2, we examined the reactivity of 2 with a variety of nitriles
under the conditions necessary to measure thermodynamic
parameters. Unfortunately, nitriles such as iso-propionitrile and
tert-butyl nitrile were found to precipitate out of solution before
complete binding was observed at -90 °C. Aside from
acetonitrile, the only nitrile that did not precipitate under these
conditions was benzonitrile. The ∆H for benzonitrile binding
to 2 was determined to be -4.2((0.8) kcal/mol. This smaller
enthalpy change relative to MeCN can be rationalized from two
different perspectives. First, the electron-withdrawing properties
of the phenyl ring results in an inherently less Lewis basic nitrile
with PhCN when compared with MeCN. This would result in
a weaker Fe-N bond if the nitrile were mainly acting as a
σ-donor, as opposed to a π-acceptor. Second, given that the Ph
group is substantially larger than a methyl group, it is likely
that steric interactions between PhCN and the ligand backbone
of 2 would also weaken the Fe-NC bond relative to MeCN.
The ∆S associated with benzonitrile binding to 2 was determined
to be -18((4) eu. This is substantially lower than the entropy
change associated with MeCN binding. Such a result can be
rationalized if, prior to its coordination to the metal, the PhCN
molecules were involved in, for example, some type of
π-stacking interaction. This would create a more ordered
unbound state relative to MeCN.
To compare the affinities of (Et,Pr)-ligated 2 versus that of
(Pr,Pr)-ligated 1 for neutral ligands, we needed to identify a
ligand which would bind to both 1 and 2 at temperatures above
the freezing point of CH2Cl2. The only ligand that fit this
criterion was pyridine. Pyridine binding to 1 requires a large
excess of ligand, and the reaction is highly reversible. Complete
binding is only observed when >5000 equiv of pyridine are
added to a methylene chloride solution of 1 at -90 °C. As the
temperature is increased, the electronic absorption peaks due
to five-coordinate 1 grow at the expense of those due to 1-py.
Above -20 °C there is no evidence of pyridine-bound 1-py in
the spectrum. Pyridine binding to 2 is also reversible, and 6500
equiv of pyridine are required to observe complete binding at
-90 °C. The thermodynamic parameters for pyridine binding
to 1 and 2 (∆H(1-py) ) -6.3(1) kcal/mol, ∆S(1-py) ) -29.0(3)
eu, ∆H(2-py) ) -8.2(6) kcal/mol, ∆S(2-py) ) -41(3) eu) were
determined spectrophotometrically and are compared in Tables
S-10 and S-11. In contrast to what one might expect, at any
given temperature the equilibrium constant for py binding to 2
is smaller than that of MeCN binding. Most likely the greater
steric bulk of the phenyl ring is responsible for the unusually
low affinity of 2 for pyridine. This would be consistent with
the PhCN results. The smaller barrier to ligand exchange with
the bulkier PhCN and py ligands is also consistent with this
explanation (vide infra).
Kinetics of Ligand Exchange. Dynamic NMR analysis was
used to determined the rates of nitrile ligand exchange from 2
and the rates of pyridine exchange from both 1 and 2.37 Nitriles
were not observed to bind to 1 at temperatures accessible in
(34) Allen, F. H.; Davies, J. E.; Galloy, J. J.; Johnson, O.; Kennard, O.; Macrae,
C. F.; Mitchell, E. M.; Mitchell, G. F.; Smith, J. M.; Watson, D. G. J.
Chem. Inf. Comput. Sci. 1991, 31, 187-204.
(35) Guidry, R. M.; Drago, R. S. J. Am. Chem. Soc. 1973, 95, 6645.
(36) Epley, T. D.; Drago, R. S. J. Am. Chem. Soc. 1969, 91, 2883.
9
J. AM. CHEM. SOC. VOL. 124, NO. 38, 2002 11425

A R T I C L E S
Shearer et al.
Table 5. Activation Parameters and 275 K kex for 2-RCN,
2-py4-tBu, and 1-py4-tBu
kex275 (sec-1
)
∆H (kcal/mol)
∆S (eu)
E (kcal/mol)
a
q
q
2-MeCN
2-PhCN
2-iPrCN
2-py4-tBu
1-py4-tBu
8.50(2) × 104
3.27(2) × 104
3.27(2) × 104
3.56(3) × 104
1.19(3) × 105
7.1(8)
4.9(8)
5.4(6)
5(1)
-10(1)
-20(3)
-17(2)
-18(2)
-25.8(3)
7.3(8)
5.3(9)
6.0(7)
6(1)
Figure 8. Atom-labeling scheme of the (Et,Pr) backbone in 2.
3.1(8)
3.7(9)
Table 4. Proton Chemical Shift Assignments of the Et,Pr
Backbone of 2 as Derived through Proton Longitudinal Relaxation
Rates at 303 K (500 MHz)
105 sec-1 at 300 K to 1.71(3) × 102 sec-1 at ∼180 K (Table
S-1). When this analysis is applied to 1, no observable line-
broadening is observed in the 13C NMR (over the temperature
range of 190-300 K) that can be attributed to chemical
exchange, indicating that MeCN does not bind to 1 over this
temperature range.
The ambient temperature exchange rate, kex, for MeCN (Table
5) is considerably faster than one would expect for a low-spin
Fe(III) complex. For example, the low-spin thiolate-ligated
Fe(III) complex [Fe(III)(SMe2N5(tren)(MeCN))]2+ (3) (Figure 10)
also undergoes nitrile exchange at room temperature with bulk
atom
label
chemical shift
(ppm)
distance from Fe
(Å)a
distance from Fe
(Å)b
T
1
(msec)
Ha
Hb
Hc
Hd
He
Hf
Hg
Hh
Hi
22.92
-10.61
19.78
3.01
3.54
3.19
3.14
3.12
3.22
3.89
4.10
3.82
3.09
2.79
3.42
3.19
3.09
3.08
3.15
3.97
4.18
3.72
2.85
2.338
6.266
3.324
3.053
2.875
3.480
9.917
15.243
10.900
2.741
26.51
-22.05
5.65
-9.72
5.32
-11.40
-31.81
Hj
solvent; however, the rate of MeCN exchange from 3 (kex
)
a Calculated from T1 values. b Calculated from crystal structure.
1.25(1) × 101 sec-1 at 293 K) is nearly 4 orders of magnitude
slower than that found for 2-MeCN.40 This decrease in rate could
be explained on the basis of differences in molecular charge
(+2 with 3 vs +1 with 2-MeCN); however, it is unlikely that
charge differences could account for such disparity in rates.
Previously, we demonstrated that the trans-thiolate of the low-
spin Co(III) complex [Co(III)(S2Me2N3(Pr,Pr)]+ increases anionic
ligand dissociation rates by 4 orders of magnitude relative to
the NMR experiment. It was previously shown that eq 2 (see
Experimental Section) could be used to determine the rate of
nitrile exchange with bulk solvent for a series of Mn2+
complexes.38 Thus, it was reasoned that this equation could be
applied to our system.
The electronic transverse relaxation rate (T1e-1) was deter-
mined as previously described.19 Briefly, by measuring trans-
verse relaxation rates for protons (T1-1) in the backbone of 2
and plotting this as a function of r-6 (Figure S-1) one can extract
[Co(III)(NH3)5(H2O)]3+ 41
Complex 3 lacks a thiolate trans to
.
the MeCN binding site, whereas 2-MeCN contains this structural
feature. Thus one might conclude that the trans thiolate aids
the displacement of the nitrile from the metal center of 2-MeCN.
NHase also possesses a cysteinate trans to its substrate-binding
site, which might assist product displacement from the NHase
metal center. This would explain how low-spin metal centers
can be utilized as catalysts in metalloenzymes involved in
substrate binding and product release. As shown here and
elsewhere,41 a trans-cysteinate is capable of turning a relatively
inert metal center into a labile metal center.
-1
the T1e value from the slope using the following equation:
T1-1 ) (2/5)S(S + 1)γH gH â2T1e r-6
(13)
2
2
Proton chemical shift assignments (Figure 8) and proton-iron
distances were determined via T1 relaxation times, and NOE
difference spectroscopy (Table 4, Figure S-2). The resulting
distances were consistent with those expected on the basis of
the X-ray crystal structure of 2.13 Using this analysis for 2, T1e
-1
To determine how electronic effects influence nitrile exchange
rates from 2-RCN, we compared MeCN exchange rates with
those of benzonitrile (PhCN). Exchange rates, kex, for PhCN
are comparable to those of MeCN (Table S-2). Activation
parameters differ significantly, however (Table 5). As illustrated
in the Eyring plot of Figure 9 and in Table 5, ∆Hq for MeCN
exchange is roughly 2.3 kcal/mol higher than for PhCN
exchange.
was determined to obey the following relationship as a function
of temperature:
T1e-1 ) [2.33 × 10-1 exp(0.008147T)]1011 sec-1 (14)
over the temperature range of 190-313 K (Figure S-1). This
value falls within the expected T1e-1 value for a low-spin Fe(III)
complex.39
This smaller barrier to exchange for PhCN could be a
consequence of the electron-withdrawing properties of the
phenyl group which decrease the Lewis basicity of the nitrile,
thus weakening the Fe-N bond. Another possibility is that
unfavorable steric interactions between the bulkier phenyl group
and the ligand backbone of 2 favor the displacement of PhCN
from 2-NCPh. To determine how sterics might influence nitrile
exchange rates from 2-RCN, the rates of iso-propionitrile
(iPrCN) exchange were measured (Table S-3) and compared
The rates for nitrile exchange (kex) from 2-RCN were
calculated using eqs (1-5) (see Experimental Section). Using
this analysis, it was determined that kex for MeCN exchange
from 2-MeCN varies significantly with temperature. As the
temperature is lowered to near the freezing point of a CH2Cl2/
MeCN mixture, the exchange rate decreases from 2.51(1) ×
(37) (a) Kaplan, J. I.; Fraenkel, G. NMR of Chemically Exchanging Systems;
Academic Press Inc.: New York, 1980. (b) Sandstro¨m, J. Dynamic NMR
Spectroscopy; Academic Press Inc.: New York, 1982.
(38) Inada, Y.; Sugata, T.; Ozutsumi, K.; Shigenobu, F. Inorg. Chem. 1998,
37, 1886-1891.
(39) Walker, F. A. Spectroscopic Methods in Bioinorganic Chemistry; Solomon,
E. I., Hodgson, K. O., Eds; ACS Symposium Series 692; American
Chemical Society: Washington, DC, 1998; p 37.
(40) Shearer, J. Models For Metalloenzymes Containing Sulfur-Metal Bonds.
Doctoral Thesis, University of Washington, Seattle, WA, 2001.
(41) Shearer, J.; Kung, I. Y.; Lovell, S.; Kaminsky, W.; Kovacs, J. A. J. Am.
Chem. Soc. 2001, 123, 463-468.
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11426 J. AM. CHEM. SOC. VOL. 124, NO. 38, 2002

Reactivity of Fe(III)−NHase Model Complex
A R T I C L E S
Variable-temperature pyridine binding kinetics were per-
formed using 4-tert-butyl pyridine (4-tBuPy) as the substrate.
The tert-butyl derivative of pyridine provides a nice spectro-
scopic handle, because there is negligible unpaired spin-density
located on the tert-butyl group. As a result, there should be no
observable line broadening of these NMR peaks due to Fermi
contact interaction. Thus, all of the line broadening should be
directly related to dynamic and outer-sphere relaxation pro-
cesses. Hence, eq 6 (see Experimental Section) can be used
t
directly to determine BuPy exchange kinetics. To further
eliminate error in the calculation of kex the tertiary carbon on
the tert-butyl group was monitored, as opposed to the methyl
group carbons. This eliminated contributions due to other
dynamic processes, such as rotation of the methyl groups, that
may be affected by temperature.
The kex for 4-tBuPy exchange from 1 was measured over the
temperature range of 190-300 K (Table S-5), and was
determined to vary from 6.56 × 103 to 2.31 × 105 s-1, which
is comparable to nitrile exchange from 2. The activation
parameters for 4-tBuPy exchange from 1 do noticeably differ,
however, from the activation parameters for nitrile exchange
from 2 (Table 5). For example, ∆Hq for 4-tBuPy exchange from
1 (3.1(8) kcal/mol) is considerably lower than ∆Hq for MeCN
exchange from 2 (7.1(8) kcal/mol). Rates for 4-tBuPy exchange
from 2 were determined over the temperature range 225-300
K (Table S-4). With this compound, considerably slower
exchange rates at temperatures below 200 K prevented deter-
mination of exchange rates for 4-tBuPy exchange from 2 at
lower temperatures. Rates varied from 1.74 × 103 to 5.79 ×
104 s-1, which is about an order of magnitude slower than those
observed for 1. This translates into a ∆Hq for 4-tBuPy exchange
from 2-py4-tBu of 5(1) kcal/mol, which is 1.9 kcal/mol higher
than ∆Hq (1-py4-tBu). This finding is consistent with the higher
barrier to azide exchange from 2-N3 versus 1-N3 and appears
to suggest that the transition-state enthalpies for 1 and 2 are
similar, with the difference in reaction enthalpies being derived
from the relative stabilities of the products. The ∆Hq for 4-tBuPy
exchange from 2-py4-tBu is comparable with that of PhCN
exchange from 2-PhCN, and both are considerably smaller than
that of MeCN exchange from 2-MeCN (Table 5). This provides
additional evidence to suggest that steric repulsions are in part
responsible for the rapid exchange of ligands which introduce
bulk directed toward the metal of the complexes described
herein.
Figure 9. Eyring plot for the exchange of MeCN from 2-MeCN.
Figure 10. Structure of [Fe(III)(SMe2N5(tren)(MeCN)]2+ (3).14
with those of MeCN and PhCN. The A value for a Ph group is
larger than the A value for an iPr group (APh ) 3.0 kcal/mol vs
AiPr ) 2.21 kcal/mol),42 suggesting that if sterics do play a role
in nitrile exchange, one would expect ∆Hq(iPrCN) to be smaller
than ∆Hq(MeCN), but larger than ∆Hq(PhCN). However, if
sterics do not play a significant role, then ∆Hq(iPrCN) should
be approximately the same as ∆Hq(MeCN). As was found with
PhCN, the ambient temperature kex rates for 2-iPrCN are
i
comparable to those of 2-MeCN, while the ∆Hq for PrCN
exchange from 2-iPrCN is significantly lower than that of
2-MeCN (Table 5). This suggests that steric factors do play a
i
role in promoting the displacement of both PrCN and PhCN
from 2-RCN when compared to MeCN. As was determined
with MeCN, iPrCN and PhCN do not appear to bind to 1 under
the conditions used to determine kex rates.
The rates of azide binding to 1 and 2 were previously
compared.13 It was determined that azide binds ∼3 times faster
to 2 relative to 1. This was attributed to the more distorted,
less stable geometry and more accessible metal ion of 2 relative
to that of 1. Furthermore, azide was found to dissociate from
2-N3 ∼1 order of magnitude slower than from 1-N3. This was
attributed to the more strained geometry of 1-N3 compared to
that of 2-N3. To obtain further evidence for this structural
difference having an influence on reaction kinetics, we attempted
to determine the kex for the exchange of neutral ligands from
both 1 and 2. Unfortunately 1 was shown to be unreactive
toward nitriles and alcohols under the conditions examined.
However, both 1 and 2 readily bind pyridine.
Summary and Conclusions
This work demonstrates for the first time that nitriles can
reversibly bind to iron when it is in a ligand environment
resembling that of NHase. The resulting nitrile-bound complex
displays spectroscopic (electronic-absorption, EPR) and struc-
tural (spin-state, average Fe-S and Fe-N bond lengths)
properties similar to that of Fe-NHase. Such work lends validity
to the proposed mechanism in which nitrile CtN bond
activation occurs via a coordinated Fe-NCR intermediate. Also,
we have demonstrated that ligand exchange in a system
resembling that of NHase is rapid, showing that a low-spin first-
row transition metal can be a competent catalyst in an enzyme
reaction that involves substrate displacement from the metal
center. Faster than expected ligand exchange rates with these
systems are attributed to the labilizing effects of a trans-thiolate.
When this structural feature is removed and replaced with a
(42) A is defined as the difference in energy between a functional group being
in the axial vs equatorial position of a cyclohexane ring, and is a measure
of steric bulk. (a) Booth, H.; Everett, J. R. J. Chem. Soc., Chem. Commun.
1976, 278. (b) Booth, H.; Everett, J. R. J. Chem. Soc., Perkin Trans. 2
1980, 255. (c) Hirsch, J. A. Top. Stereochem. 1967, 1, 199.
9
J. AM. CHEM. SOC. VOL. 124, NO. 38, 2002 11427

A R T I C L E S
Shearer et al.
trans nitrogen, rapid ligand exchange is no longer observed.
NHase also possesses a thiolate trans to the substrate-binding
site. It thus appears that nature utilizes the labilizing effect of
a trans-thiolate to improve product off-rates from the metal
center of NHase.
Supporting Information Available: Variable-temperature
NMR kinetics results for 2-MeCN, 2-iPrCN, 2-PhCN, 1-py,
2-py (Tables S-1-S-5), temperature dependence of T1e-1 for 2
(Figure S-1), the 1H NMR spectrum of 2 (Figures S-2),
equilibrium constants (Tables S-6-S-9), van’t Hoff plots
(Figures S-8-S-11) for ligand binding to 1 and 2, variable-
temperature electronic absorption spectra (Figures S-3-S-7),
and EPR spectra for “substrate”-bound 2 (Figures S-12-S-19),
and crystallographic data for 2-NCS (Tables S-10-S-14) (PDF).
This material is available free of charge via the Internet at
http://pubs.acs.org.
Acknowledgment. We thank Steven Schultz and Vincent
Pons for experimental assistance with the collection of low-
temperature NMR spectra. This work was supported by the NIH
(Grant No. GM 45881). Research carried out (in part) at the
National Synchrotron Light Source, Brookhaven National
Laboratory, was supported by the U.S. Department of Energy,
Division of Materials Sciences and Division of Chemical
Sciences.
JA012555F
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11428 J. AM. CHEM. SOC. VOL. 124, NO. 38, 2002