Synthesis and Structural Characterization
As ligand and iron salt used for the synthesis of 6 are soluble in demineralized water, attempts were made to prepare a "water based" dinuclear iron complex. When dissolving K5LCOO and Fe(BF4)2·7 H2O in water a red solution is obtained. This solution was stirred for a few minutes and quickly filtered. Diffusion of acetone into the reaction mixture afforded red single crystals suitable for x-ray diffraction. The obtained neutral
complex7crystallizes in the triclinic space groupP¯1with one complex molecule, 5.5 H2O molecules and one acetone molecule in its unit cell. The structure of 7features two seven-coordinate iron sites bridged by a hydroxo ligand. A {N3O4} donor set can be assigned for each metal. Aside from the chelating ligand and the hydroxo bridge, water molecules coordinate to both iron sites. Interestingly, these two water molecules point towards the same side of the complex in a syn fashion. The Fe···Fe distance in 7 is 3.75 Å, which is fairly short compared to the corresponding diiron complexes in this work with entirely nitrogen based ligands. Figure 7.4 displays the molecular structure of the neutral complex 7.
Figure 7.4: Crystal structure of [LCOO{Fe(H2O)}2(µ–OH)], (CH3)2CO, 5.5 H2O 7. Thermal ellipsoids are set to 50 %. Not relevant hydrogen atoms as well as solvent molecules were omitted for clarity. Selected bond lengths [Å] and angles [◦] Fe1–N1 2.089(4), Fe1–N5 2.315(4), Fe1–N6 2.344(4), Fe1–O5 2.126(4), Fe1–O7 1.993(3), Fe1–O9 2.015(4), Fe1–O11 2.004(3), Fe2–N2 2.089(4), Fe2–N3 2.308(4), Fe2–N4 2.385(4), Fe2–O2 2.106(4), Fe2–O4 1.986(3), Fe2–O10 2.012(4), Fe2–O11 1.997(3), Fe2–O11–Fe1 139.65(19), Fe1–N1–N2–Fe2 -3.166(6).
Since the ligand is fivefold negatively charged and the hydroxo bridge is a monoanionic ligand a +iii charge results for each iron atom.
As shortly addressed before the two water ligands in 7coordinate to the metal centers in asyn fashion. This can be seen as excellent preorganisation of the substrate water for the catalysis. Generally two water oxidation mechanisms are described for dinuclear water oxidation catalysts in literature.[125] After the oxidation of the bimetallic complex, two metal oxo units preferably in syn orientation are formed.
Fe Fe
O O
H2O
Fe Fe
O O
Water Nucleophilic Attack (WNA)
Interaction of Two M-O Units (I2M)
Scheme 7.3:Discussed mechanisms of O–O bond formation in water oxidation catalysis.
For O–O bond formation a water molecule attacks as a nucleophile at one of the
metal-oxo units to yield a hydropermetal-oxo and a hydrmetal-oxo ligand. A second water molecule then liberates dioxygen and the initial two water molecules coordinating to the dinuclear core are regained. This mechanism is called water nucleophilic attack (WNA). A second option involves O–O bond formation via a direct interaction of two metal oxo units (I2M).[125]
However, for both reaction pathways 7 might be an appropriate precursor due to itssyn water coordination.
ESI– mass spectrometry from a mixed water/isopropanol solution reveals as observed for 6, a "naked" oxo-bridged diiron complex without any coligands (Appendix 11.5, Figure 11.20). The main signal in ESI+spectra can be attributed to the hydroxobridged complex without the water coligands together with a sodium cation.
Mössbauer Spectroscopy
To determine the spin state of the two iron sites in 7, Mössbauer measurements were conducted. In this respect diiron complexes differ greatly from nobel metal complexes as these feature mostly low spin configurations. Among the 3d metals however, due to smaller ligand field splittings the spin state of a metal in coordination complexes strongly depends on the coordination sphere and nature of the donor sites. Zerofield Mössbauer spectroscopy of a polycrystalline sample of 7 at 80 K suggests a single iron(iii) species which accounts for both iron atoms, as both feature the same coordination geometry. The isomer shift of δ= 0.57 mm s–1 and the rather small quadrupole splitting point towards a largely spherical electron distribution and thus a Fe(iii) high spin state.[204] Figure 7.5 illustrates the recorded spectrum of 7.
Figure 7.5: Zerofield Mössbauer spectrum of a polycrystalline sample of 7at 80 K. Isomer shift δand quadrupole splitting |∆EQ| in [mm s–1]: δ= 0.57, |∆EQ| = 0.57.
High spin complexes usually exhibit different electronic properties compared to low spin complexes which express themselves in profound changes in optical, vibrational, magnetic and structural properties.[4]
Electrochemistry
In analogy to 6, the redox potentials of 7were examined by means of cyclic voltammetry.
A polycrystalline sample of 7 was dissolved in an aqueous solution (0.1M NaClO4). As
working and counterelectrode a glassy carbon and platinum electrode were employed.
A saturated calomel electrode (SCE) was used as reference electrode. The respective measurements are shown in Figure 7.6.
Figure 7.6: Left: Cyclic voltammogram of 7in neutral H2O (0.1MNaClO4) vs SCE reference.
The inset shows the corresponding square wave diagram. Right: Large window view. Currents were inverted for a better comparison with CV’s according to the water oxidation formalism. The green trace describes the solvent background. The cyclic voltammogram of7is depicted in black.
The grey box highlights the range in which water oxidation might occur.
The cyclic voltammogram of 7exhibits as previously described for6two reduction events marking the Fe(ii)Fe(iii) couples. The two redox potentials E1/2from square wave voltam-metry lie at 0.13 V vs. SCE (iii,iii/iii,ii) and –0.34 V vs. SCE (iii,ii/ii,ii), respectively.
Zooming out of the cathodic reduction range, into a large window view (Figure 7.6 right) an enhanced current can be observed above 0.5 V vs. SCE. However, this oxidation event does not resemble the typical plateau shape of electrocatalytic events as observed for well-studied water oxidation catalysts. Generally, a more accurate prediction could be made with additional measurements at different scan rates for potentials above 0.5 V vs. SCE.
Simple electrocatalytic events exhibit an S-shaped curve that is independent of the scan rate and can be described by the following equation:[205]
i= ipl 1 + eRTF (E–E1/2)
(7.2)
i current
ipl current at the plateau E electrode potential
E1/2 half-wave potential of the steady-state catalytic wave
A plateau shaped current was not observed. No scan rate dependent measurements were performed as problems were encountered with the solubility of the complex in neutral water. Between each CV cycle vigorous stirring of the solution ensured a maximum sol-vation of the sample. However, still slight solid sample residues were observed during the measurement. The inhomogeneity of the solution can have basically two most probable reasons. One is that7is simply poorly soluble in neutral water, but stable over the course of the measurement. Another reason could be that over the course of the measurement iron oxide nanoparticles are produced, which leads to depletion. A variety of other events
that might contribute to inhomogeneity are certainly also possible.[206] A comprehensive experiment to rule out the formation of nanoparticles would involve a rinse test in which a CV with the same electrodes but a fresh solution without sample would be recorded. If metal oxide particles formed on the surface of the electrodes, enhanced currents would be observed. Another possibility could involve a blank test, in which simply a CV of the iron salt is recorded. However, a more thorough electrochemical analysis of 7 was beyond the scope of this work.
To probe whether 7 performs water oxidation under the well studied common water oxi-dation conditions in triflic acid (pH = 1) and with (NH4)2[CeIV(NO3)6](CAN) as oxidant the oxygen evolution was measured with a Clarke electrode. Apparently 7 is not very well soluble at low pH values. The oxidant was added anyway to a suspension of 7 in triflic acid. A color change from a red suspension to a yellow solution with white pre-cipitate was observed in the course of minutes. Additionally, no oxygen evolution was detected. Consequently, it can be assumed that these standard conditions are not suitable for the oxidation of 7. The complex decomposed. ESI mass spectra of the compound after oxidation did not contain any signals typical for the diiron complex.
It seems that the oxidizing strength of CeIV is too large for the oxidation of 7. The redox potential of CAN of 1.75 V vs. NHE (at pH = 0.9) is very high.[207] The observed oxidation event of 7 in neutral water features an onset of about 0.5 V vs. SCE which is significantly lower. Maybe the choice of a milder oxidant would be wise in this context.
Further well established oxidants for water oxidation catalysis are [RuIII(bpy)3]3+ (1.21 V vs. NHE), sodium peroxodisulfate (2 V vs NHE), Oxone (1.82 V vs. NHE) and sodium periodate (1.6 V vs. NHE).[207] The only oxidant of these reported ones, [RuIII(bpy)3]3+, which would be milder, requires similar to CeIV acidic conditions, in which7, as shortly addressed above, is only poorly soluble. Moreover, the modification of 7 at low pH values is unknown. In the future, a speciation analysis would be helpful. Also excluded should be the possibility of complex degradation in acidic media. In that respect it would be necessary to probe the complex stability of 7 within a longer time window in respective acids. However, again the poor solubility of 7 did not allow for a detailed analysis by spectroscopic monitoring.
pH-Titration followed by UV/Vis Spectroscopy
Since cyclic voltammtry revealed solubility problems of 7 in acidic and neutral aqueous solutions a pH-titration was performed followed by UV/Vis spectroscopy. For this mea-surement a definite amount of polycrystalline 7 was dissolved in a solution of 0.1M HCl and 0.1M KCl. This mixture was titrated with defined amounts of 0.1M NaOH and the pH was recorded. KCl was used to maintain a close to constant ionic strength during the measurement. For every pH value a UV/Vis spectrum was measured. The corresponding spectra are depicted in Figure 7.7.
For pH values below pH = 7 two absorption maxima at 243 and 290 nm can be observed for7. The intensity of these absorption bands increases with increasing pH value. Above pH = 7 the bands shift about 12 nm to larger wavelengths. A general observation is
more-Figure 7.7: pH dependent UV/Vis titration of 7. The black line represents the initial spectrum at low pH values, the red line depicts a spectrum under basic conditions.
over that7 dissolves well in basic solution, which also explains the increasing intensity.
This suggest that in future experiments water oxidation catalysis with 7 should be best performed in buffer solutions to maintain a slightly basic pH, which ensures full solubility of the complex. A suitable non coordinating buffer that fulfills this requirement has not been found yet. However, a buffer is inevitably important to perform these experiments as at lower pH values iron oxide nanoparticles might form, which can also take part in catalytic events and thus tamper with the entire experiment.