Download Heme Models - Bryn Mawr College

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Iron-Porphyrins as Models for Hemes
Transition metals play an critical roles in myriad biological processes. The metal ions and
their complexes serve as structural supports, electron transfer relays, oxygen transport and
storage depots, and catalytic centers.1-3 "Model" compounds have served to further our
understanding of many of these complex systems.
This is a series of experiments that involve the synthesis and characterization of iron
porphyrins. These 'heme' models are used to illustrate the effect of the axial ligands on spin
state4 and redox chemistry5 . The experimental results are used in conjunction with the
interpretation of (a) the stereochemical changes accompanying oxygen binding and the resulting
cooperativity exhibited between heme units in hemoglobin (i.e., the Perutz trigger mechanism)6
and (b) the effect of heme environment on Fe(III)/Fe(II) redox potential and how this pertains to
electron shuttling by cytochrome sequences5b,7. In addition, the synthesis of the
metalloporphyrin illustrates one method of macrocycle synthesis and the macrocyclic effect.
The sequence of experiments described below is carried out over several laboratory periods.
Each group will work with one of three different iron porphyrin complexes. Two different
techniques will be used during the course of this multi-week lab: cyclic voltammetry and
solution magnetic susceptibility measurement by the Evans’ method
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Synthesis of the Porphyrin-Ligand
You will synthesize one of three possible porphyrins:
1) tetraphenylporphyrin, TPP;
2) tetra(p-tolyl)porphyrin, TTP; or
3) tetra(p-chlorophenyl)porphyrin, TClPP,
using a modified version of the one-pot procedure developed by Adler et al.8 as depicted below
R
O
R
N
H
N
propionic acid
+
NH

HN
N
R
R
aldehyde
R
pyrrole
Procedure
**All manipulations including isolation of the product must be performed in a hood.**
A stir bar and 0.3L propionic acid are added to a 500-mL three-neck, round-bottom flask fitted
with a reflux condenser. Put the addition funnel in the center neck, the condensor on a side neck
and a septum on the third neck. The acid is brought to a reflux and 0.1 mol of the appropriate
substituted benzaldehyde is added.
Quantities of aldehydes:
benzaldehyde [for TPP synthesis] = 10.16 mL
p-tolualdehyde [for TTP synthesis] = 11.8 mL
4-chlorobenzaldehyde [for TClPP synthesis] = 14.06 g
To the refluxing solution, 10 mL (0.1 mol) freshly distilled pyrrole is cautiously added via a
dropping funnel. A recommended rate is about 1 drop/sec.
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Caution! The condensaton reaction of the pyrrole with the aldehyde is extremely exothermic,
and the rate of pyrrole addition must be continously monitored. The solution is refluxed for 30
min, cooled to ambient temperature, and chilled in an ice bath before vacuum filtering off the
purple, crystalline product. [Note: Fitrations can be slow. Try adding small amounts of the
reaction solution, rather than dumping the entire volume into the funnel. Also use of a coarse 4060 mL fritted funnel for filtering the product.] The product is washed with boiling water (about
500 mL) until free from the odor of propionic acid. Typical yields should be in the range of 1020%. The crude product is of sufficient purity for use in the metallation reaction.
WASTE. The reaction solution filtrate goes in the provided container in the hood. The
washings may be put down the drain.
Metallation of the Free-Base Porphyrin
The iron porphyrin is conveniently prepared from the free-base porphyrin by a scaled-down,
modified version of the procedure reported by Adler et al.10.
R
R
R
N
NH
R
N
FeCl2
Cl
N
HN
Fe
N
DMF
N
N
R
R
R
R
R = TPP:
R = TTP:
R = TClPP:
Cl
The progress of the reaction is conveniently monitored by checking the reaction medium for
fluorescence. Dilute solutions of the free-base (i.e., metal free) porphyrins fluoresce brightly red
when irradiated with long-wave UV light in a darkened room, whereas the metallated porphyrins
do not.
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Procedure
In a 250-mL three-neck, round-bottom flask equipped with a reflux condenser, a stirred
mixture of 1 g of free-base porphyrin in 100-mL dimethylflormamide, DMF, is brought to a
reflux. About 0.5 g of FeCl2 is added over a 5-min period, and the solution is allowed to reflux
for 10 min. Caution! The addition of FeCl2 to the reaction medium is accompanied with the
evolution of heat. The FeCl2 must be added in small increments. If an examination of the
reaction medium reveals fluorescence, additional FeCl2 is added and the reaction mixture
refluxed for an additional 10 min. The DMF solution is cooled to room temperature and an equal
volume of water added along with 1 mL of 6 M HCl. The mixture is chilled in an ice bath and
suction filtered. The dark brown-purple solid is washed with water until the filtrate is colorless
and aspirated to dryness. Typically, isolated yields are greater than 90%. WASTE DISPOSAL:
The initial filtrate should go in the metal containing waste container in the hood. The
subsequent water washings may go down the sink with plenty of water.
Magnetic Susceptibility Measurements
The number and nature of axial ligands are the primary determinants of spin state in
ferriporphyrins4. Thus, as the axial ligand field strength increases, the spin state adopted by an
Fe(III)(P), where P is any porphyrin, progresses from intermediate-spin (S = 3/2), to high-spin (S
= 5/2), and finally to low-spin (S = 1/2),. This is represented in Figure 1 with simple ligand-field
splitting diagrams. An Fe(III)(P) with very weak-field ligands invariably adopts a spin state that
is not pure S = 3/2 but rather quantum mechanically admixed S = 3/5, 5/2.
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Figure 1. The variation of d-orbital energies with increasing axial field strength.
dx2 -y2
dx2 -y2
dx2 -y2 , dz2
dz2
dz2
dxy , dyz, dxz,
dxy , dyz, dxz,
increasing axial ligand field strength
The effective magnetic moments of the iron porphyrins can be determined for the solid state
by the Faraday Balance method and in solution using Evans' method. For the solution
susceptibility (Evans’ Method), you will do :
1) one measurement of your group’s Fe(P)Cl in CHCl3 and
2) one measurement using in a solution of 0.1 M imidazole in CHCl3 .
The presence of the imidazole in (2) causes in situ formation of low-spin bis-imidazole
complexes, [Fe(P)(Im) 2]Cl and provides another environment of the Fe atom found in biological
hemes.
(Note: in the case (2) the reference solution for the second solution must contain imidazole in the
same concentration). An Fe-porphyrin concentration of about 0.01 M works well for these
Evans’ experiments and you will need to calculate the appropriate amounts of Fe-P compound
needed for the experiments prior to coming to lab.
The results from each group will be shared with the rest of the class so that the magnetic
behavior of each iron-porphyrin in the series TPP, TTP and TClPP is examined and comparisons
of the magnetic state as a function of added imidazole and without imidazole can be made.
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Magnetic moments for Fe(P)Cl by both the Guoy and Evans methods are typically in the range
5.3-6.0 µB, corresponding to a high-spin state (S = 5/2) whereas the values obtained for
[Fe(P)(Im)2]Cl are 2.4-2.7 µB corresponding to a low-spin state (S = 1/2).
Samples will need to be corrected for diamagnetism using these values:
-700 x 10-6 cgs for TPP
-753 x 10-6 cgs for TTP
-760 x 10-6 cgs for TC1PP.
Electrochemical Characterization by Cyclic Voltammetry
The electrochemistry of iron porphyrins has been well characterized5, and the potentials for
various redox processes correlate with many aspects of iron-prophyrins: metal spin state,
coordination of axial ligands, the solvent system employed for the study, out-of-porphyrin-plane
displacement of the iron, counterion, and basicity of the porphyrin ring5c. In this lab,
experiments will focus only on the effect of porphyrin basicity and axial ligation on the Fe(III/II)
and Fe(II/I) redox potentials.
Cyclic voltammetry is used to obtain the half-wave reduction potentials for the Fe(III/II) and
Fe(II/I) couples in the absence and presence of imidazole. The results from each group will be
shared so that each porphyrin in the series is examined and comparisons of the E1/2’s can be
made. 0.1 M tetraethylammonium perchlorate, TEAP, in dimethylformamide, DMF, performs
well as a supporting electrolyte-solvent system combination. Complexities associated with the
displacement of anions from the coordination sphere are minimized in coordinating solvents such
as DMF (5c).
Procedure
A three-electrode cell supplied by Bioanalytical Systems, Inc., and composed of a Pt disk
electrode, Pt wire auxiliary electrode and a Ag/AgCl reference electrode will be used.
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A DMF solution 0.1 M in TEAP is prepared in the electrochemical cell and deaerated for 10
min with a stream of nitrogen. A voltammogram is obtained using a scan range of +0.3 to -1.4 V
vs. Ag/AgCl and a scan rate of 50 mV/s. A small amount of the Fe-prophyrin to be analyzed is
added to the electrolyte solution in the cell, deaerated for 5 minutes, then scanned over the same
potential range +0.3 to -1.4 V. Then 100 mg of imidazole is put into the DMF solution, which is
again deaerated. A second voltammogram is obtained. Finally, ferrocene is added to the solution
and the wave for the ferrocene-ferricinium couple is recorded for calibration of the cell.
Results
The data to be collected from each group is as follows:
1. E1/2 values for one of the three iron porphyrins , in the absence and in the presence of
imidazole.
2. The observed chemical shift and calculated magnetic susceptibility for one of the three
iron porphyrins , in the absence and in the presence of imidazole.
Analysis of the half-wave reduction potentials should reveal a number of trends that you should
attempt to explain using concepts of electronic effects first introduced in organic chemistry
courses.
1. J. Chem. Ed. 1985, vol. 62, p 917.
2. the Bioinorganic chapter of any modern Inorganic text
3. J. Chem. Ed. 1989, vol. 66, p 854.
4. Scheidt and Reed Chem. Rev. 1981, vol. 81, p 543.
5. “Electrochemical and Spectrochemical Studies of Biological Redox Components”; currently
in my office.
6. Acc. Chem. Res. 1987, vol. 20, p 309.
7. Acc. Chem. Res. 1972, vol. 5, p 234.
8. Adler, J. Org. Chem. 1967 vol. 32, p 476.
9. Adler, J. Amer. Chem. Soc. 1975 vol. 97, p 5107.
10. Adler, J. Inorg. Nucl. Chem. 1970 vol. 32, p 2443.