The final definitive version of this manuscript was published in Australian Journal of Chemistry.
http://dx.doi.org/10.1071/CH12322
Role of H+ in polypyrrole and poly(3,4-ethylenedioxythiophene)
formation using FeCl3.6H2O in the room temperature ionic liquid,
C4mpyrTFSI
Graeme A. Snook,a Anand I. Bhatt,*b Muhammad E. Abdelhamid,a,c Adam S. Bestb
a
Commonwealth Scientific Industrial Research Organisation (CSIRO), Process Science and
Engineering, Box 312, Clayton South, Vic., 3169, Australia
b
Commonwealth Scientific Industrial Research Organisation (CSIRO), Energy Technology, Box 312,
Clayton 3169. Victoria. Australia.
c
RMIT University, School of Applied Sciences, Melbourne 3001. Victoria, Australia
Email: [email protected]
Abstract:
The polymerisation reaction of pyrrole and 3,4-ethylenedioxythiophene using the chemical oxidant
FeCl3.6H2O in the room temperature ionic liquid butyl-methylpyrrolidinium
bis(trifluoromethanesulfonyl)imide has been investigated using cyclic voltammetry, UV/vis and IR
spectroscopy. The voltammetric data for the Fe2+/3+ reaction is complicated by presence of H+
introduced upon dissolution of the iron salt via decomposition of the coordinated waters. The
voltammetric and chemical reaction studies show that H+ itself, introduced to solution as HTFSI, can
itself act as the chemical oxidant for the polymerisation reaction. Voltammetric data also implies that in
this system the Fe2+/3+ redox couple may not actually be involved in the polymerisation reaction and
that the H+ introduced upon dissolution of the FeCl3.6H2O may be the sole cause of the oxidation
reaction.
Keywords:
room temperature ionic liquid; conducting polymer; chemical polymerisation; polypyrrole; poly(3,4ethylenedioxythiophene); 1-butyl-methylpyrrolidinium bis(trifluoromethanesulfonyl)imide
1
Introduction:
The use of conducting polymers in electrochemical devices is being explored
for numerous applications including formation of large scale electrodes for
supercapacitor and battery devices [1-12]. Although the electrochemical formation
mechanism of polypyrrole and poly(3,4-ethylenedioxythiophene) (pedot) is well
known [13-16] widespread application of this methodology on an industrial scale may
be energy intensive.
An alternative synthesis methodology involves chemical oxidation of the
aromatic monomers to form conducting polymers and this methodology has been
extensively studied in the past [17-19]. However, the use of molecular solvents for
this purpose can result in industrial applications that may be environmentally
damaging.
Ionic liquids have been suggested as a replacement for molecular solvents and
synthesis of conducting polymers are being increasingly studied in ionic liquids [2028]. However, in many cases synthesis is still performed in molecular solvents and the
conducting polymer layers transferred into ILs for further studies.
REPLACE REF 20 with:
G.A. Snook, A.S. Best, J. Mater. Chem. 2009, 19, 4248-4254,
The use of IL as a reaction media for chemically preparing conducting
polymers in a one-pot synthetic methodology has only been investigated in recent
times [22-24]. However, there is a substantial knowledge gap in the mechanism of the
chemical polymerisation reaction in this media. In work by Pringle, et al. [22,23]
oxidants such as iron (III) perchlorate hydrate, iron (III) p-toluenesulfonate
hexahydrate (Fe-tosylate), silver (I) nitrate, gold (III) chloride and iron (III) chloride,
and hydrated ferric chloride, FeCl3.6H2O, were used. Yang, et al. [24] used
2
ammonium peroxydisulfate in an aqueous solution containing 0.2M aniline, 2M
H2SO4, 30% IL (v/v) to polymerise aniline and aniline derivatives. Often chemical
oxidants are used to synthesise the conducting polymers in ionic liquids without much
thought as to which chemical species is reacting and causing the oxidation process to
occur.
In this manuscript we present details of our studies on the mechanism of
polypyrrole and poly(3,4-ethylenedioxythiophene) (pedot) formation from the room
temperature ionic liquid, butyl-methylpyrrolidinium
bis(trifluoromethanesulfonyl)imide (C4mpyrTFSI) by the use of hydrated ferric
chloride (FeCl3.6H2O) as the oxidant. We discuss our voltammetric investigations of
the FeCl3.6H2O redox chemistry and present our findings on the presence of H+ in the
IL produced from dissolution of the FeCl3.6H2O into C4mpyrTFSI. We also discuss
our investigation into the role of H+ and present our finding that H+ can itself be used
as the chemical oxidant rather the FeCl3.6H2O complex, thus, simplifying the
reaction.
Results and Discussion:
1. Cyclic voltammetry study of FeCl3.6H2O in C4mpyrTFSI
Addition of FeCl3.6H2O to pyrrole and edot (3,4-ethylenedioxythiophene)
monomer solutions in C4mpyrTFSI results in chemical oxidation of the monomer to
form the corresponding polymers, polypyrrole or pedot. In order to ascertain further
information regarding the mechanism of this chemical reaction, electrochemistry was
employed as the investigative tool.
Cyclic voltammograms of 10 mM solution of FeCl3.6H2O dissolved in
C4mpyrTFSI recorded at a GC and Pt working electrodes are shown in Figure 1 with
3
selected parameters summarised in Table 1. This work was done under an Argon
atmosphere in an Argon filled glove box (O2 at less than 2 ppm). At a GC electrode
(Figure 1A), a reductive process, Epred, at -0.6 V and an oxidative process, Epox, at ca.
-0.4 V are observed which are assigned to the Fe3+/Fe2+ reduction and corresponding
Fe2+/Fe3+ oxidation processes. The peak separations, defined as
Ep = Epox-Epred, are
in the range of 160 to 310 mV over the scan rate range of 20 - 200 mV s-1. This value
of
Ep is far larger thaen the 58 mV expected for a reversible process with a one
electron transfer. Later this will be shown not to be simply excessive ohmic (IR) drop.
Additionally, the ratio of the peak currents for the reduction and oxidation processes
do not equal unity as would be expected for a reversible process. The data obtained
does not show the expected response of the Fe2+/3+ processes, namely a reversible or
quasi reversible electron transfer process. Previously, Katayama et al. have studied
the voltammetric response of a range Fe salts in C4mpyrTFSI and found that the
FeCl4– (formed by dissolution of the anhydrous FeCl3 in the chloride based IL),
Fe(CN)63–, and FeBr4– complexes exhibit reversible 1 e– transfer redox
electrochemistry at a Pt electrode [29]. Additionally, they found that the FeII/IIITFSI
complex exhibits a quasi reversible or irreversible redox electrochemistry in
C4mpyrTFSI [29]. It should also be noted that Katayama's work showed that the
FeTFSI complex exhibits redox processes at E1/2 1.1 V [29]. In the case of the
FeCl3.6H2O, scanning into this voltage region does not show any processes that can
be attributed to an Fe-TFSI complex. There is a concentration dependence for the Emid
(mid point potential) of the FeCl3.6H2O complex, which is likely to be related to the
dissociation of the proton from the complex (as discussed in later sections).
Comparison of the peak potentials for the Fe2+/3+ processes in C4mpyrTFSI to those
obtained in aqueous media for the FeCl3.6H2O complex shows that in the IL a
4
significant reductive shift in peak potentials is observed (E1/2 in IL = ca. -440 at 10
mM cf. Emid in water = -160 mV vs. Ag/AgOTf reference electrode) and is attributed
to the speciation differences of the Fe complex within the IL.
Interestingly, if the FeCl3.6H2O concentration is reduced to 5 mM,
voltammograms obtained at a GC electrode show behaviour as would be expected for
a reversible 1 e- transfer process, namely,
Ep values of ca. 60 mV over the scan rate
range of 20 - 200 mV s-1 and the ratio of peak currents closer to unity thaen in the
case of the 10 mM solution (Figure 1B). However, closer analysis of the
voltammograms obtained under these conditions again show deviations from ideality.
For an electrochemically reversible reaction (Nernstian system) the separation
between the peak potential and the halfwave peak potential, Ep-Ep/2, is 55.9 mV at 22
°C [30]. Analysis of the voltammograms obtained under the apparent reversible case
show that the Ep-Ep/2 values do not follow this relationship. Clearly, the electrode
reactions for FeCl3.6H2O in C4mpyrTFSI is more complex thaen for the Fe2+/3+
systems studied by Katayama et al.
Based on the data presented below, the deviations from expected ideality for
the Fe2+/3+ process observed in the voltammograms is attributed to low levels of H+ in
solution. Consequently, the concentration dependent voltammetric profiles and
apparent quasi reversibility is consistent with presence of additional and related
electroactive processes occurring simultaneously to the Fe2+/3+ redox reaction.
Consequently, attempts to calculate the Fe diffusion coefficient in C4mpyrTFSI using
cyclic voltammetry were not attempted due to the uncertainties involved with this
process. In principle, a steady state method such as voltammetry at a rotating disk
electrode could be used to ascertain the diffusion coefficient. However, steady state
voltammograms recorded for the FeCl3.6H2O in C4mpyrTFSI were not able to
5
decouple the two overlapping reactions. A plot of peak current as a function of square
root of rotation rate (shown in the supporting information) shows a deviation from
linearity at high rotation rates. Thus, the Levich equation could not be used to
calculate the diffusion coefficient. Clearly the electrochemical response is complex
for this system.
At a Pt electrode the voltammograms recorded (Figure 1C) are far more
complex than in the case of the GC electrode. When the scan direction is negative
from an initial potential of +1.0 V to a switching potential of -1.3 V, in addition to the
Fe2+/3+ processes previously observed at the GC electrode, a number of extra
processes are also observed. These processes show a substantial increase in peak
currents when the Pt electrode is pre-oxidised by holding the potential at +1.5 V for
15 seconds prior to the scan. Based on the data obtained in the following sections
these additional processes observed in the voltammogram are assigned to H+
reductions and related processes in solution.
Interestingly, at a Pt electrode if the potential is scanned in the oxidative
direction from an initial potential of -1.5 V to a switching potential of -0.1 V after
holding at -1.5 V for 30 seconds no protic processes are observed in the
voltammogram (Figure 1D). Presumably, scanning using this regime consumes any
available protons at the electrode surface prior to commencement of the scan and thus
the response of only the Fe2+/3+ process can be observed in the voltammogram.
Additionally, once the H+ has been consumed from the electrode surface, then the
Fe2+/3+ couple shows more reversible behaviour, that is
Ep values of 80 mV and
peak current ratios close to unity. However, it should be noted that the process still
does not fulfil the criterion for electrochemical reversibility.
6
If the voltammetry is repeated using the corresponding Fe(II) starting material
FeCl2.4H2O, then similar results are obtained. The voltammograms obtained at a
platinum electrode (Figure 1E) show markedly similar behaviour to the
voltammograms presented in Figure 1D. An oxidative process, Epox, at -0.5 V and an
reductive process, Epred, at ca. -0.6 V are observed which are assigned to the Fe2+/Fe3+
oxidation and corresponding Fe3+/Fe2+ reduction processes. Similar to the Fe3+
system where the potential is held negative prior to scanning, beginning with the Fe2+
starting material a peak separation of ca. 80 to 120 mV is observed over the scan rate
range of 20 to 200 mV s-1 (see supporting information for data table). However, peak
current ratios of less thaen unity are observed for this system and are in the range of
0.9 to 0.8 over the scan rate range. Since a pre-reduction step was not carried out,
some trace levels of H+ would be present in the system and thus interfere with the
Fe2+/Fe3+ oxidation process. This can be observed at fast scan rates where a shoulder
at ca. +0.35 V due to the H+ oxidation is observed in Figure 1E.
Overall, the voltammetry study of FeCl3.6H2O show that the Fe2+/3+ processes
are complicated by presence of H+ in solution. The voltammetric processes detected
are strongly dependent on concentration and electrode substrate used.
In order to explore the effects of Cl- ligands on the H+ generation mechanism,
preliminary studies conducted in our laboratories show that use of a non chloride
related compound, Fe(NO3)3.9H2O does not produce any bands due to H+ (see
supporting information for voltammograms). Clearly one of the roles of chloride in
this system is in the formation of complexes whereby waters are directly coordinated
to the Fe centre. This then enables the deprotonation to occur and thus liberates H+
into solution. Simply having waters of crystallisation dissolve into the IL does not
7
allow any deprotonation to occur, as also evidenced by the voltammetry where water
is deliberately added to the system as discussed in later sections.
2. UV/vis spectroscopic study of FeCl3.6H2O in C4mpyrTFSI
The voltammetric data obtained for FeCl3.6H2O is complicated by the
presence of additional electroactive processes. It is feasible that upon dissolution of
the FeCl3.6H2O, deprotonation of the Fe(H2O)63+ cation can occur which may liberate
H+ into solutions to form a iron-hydroxy species. It is well known that thymol blue
shows strong absorption bands in the UV/visible region and that the peak absorbance
at λmax is strongly dependent on the protonation level of this compound. Since thymol
blue is soluble in C4mpyrTFSI, this compound can be used as a proton indicator in
combination with UV/vis spectroscopy.
UV/vis spectra of solutions of 5 - 50 mM HTFSI dissolved in water are shown
in Figure 2(A) and C4mpyrTFSI in Figure 2(B) in the presence of thymol blue. In all
cases a strong absorption band at ca. 550 nm attributable to the protonated thymol
blue absorption band is observed. Comparing the ionic liquid solutions (0.2 µM) with
the aqueous solutions of thymol blue (0.1 mM), the transition is much more intense
(with an extinction coefficient roughly 500 times that in water). In order to confirm
that this band is due to a protonated thymol blue absorption, UV/vis spectra of thymol
blue in aqueous acidic media were recorded in concentration range of 5 mM to 50
mM. In all cases these spectra showed an absorption band with a λmax of 544 nm. The
deprotonated form of thymol blue (low acid concentration) had a band at λmax = 430
nm, that disappears upon increase in acid concentration. There is a small shift in the
λmax value which is attributed to solvent interactions going from aqueous to IL media,
from the spectra observed in Figure 2(B).
8
The UV-Vis spectrum of FeCl3.6H2O in the ionic liquid (in the absence of
thymol blue) is shown in Figure 3(A). Previously de Laat et al. have studied the
UV/vis spectra of a number of Fe3+ monomeric and oligomeric water-hydroxy
complexes in acidic media [31]. de Laat's work showed that the OH- or OH2 to Fe3+
ligand metal charge transfer (LMCT) transitions occur in the 190 - 330 nm range for
the monomeric [Fe(H2O)6]3+, [Fe(H2O)5(OH)]2+ and [Fe(H2O)4(OH)2]+ species and in
the range of 205 - 470 nm for the oligomeric [Fe2(OH)2(H2O)8]4+ species [31]. It can
be observed in Figure 3(A), that three major bands at 274, 312 and 363 nm exist.
Additionally, two weak transitions at 676 and 751 nm are also observed due to the d5
electronic transition which are similar in position to those observed for FeBr4- in
C4mpyrTFSI [29,32,33]. Comparison of the spectrum obtained to the data presented
by de Laat et al. suggests that the major bands are due to the LMCT of the OH- to Fe
transitions and could be [Fe(H2O)5(OH)]2+ or [Fe(H2O)4(OH)2]+ species, which are
deprotonation products of the hydrated Fe3+ ion [31]. Previously de Laat et al. have
experimentally observed these LMCT transitions (O → Fe) in aqueous medium at 210
and 245 nm (assigned to the p-dx2y2 and p-dz2/dx2y2 transitions) and at 335 nm (assigned
to the p-dxy/dxz/dyz transitions) [31]. Although the bands for the LMCT transitions in
the C4mpyrTFSI IL are shifted compared to the aqueous system similar results have
also previously been observed by Hussey and Seddon et al. where shifts from aqueous
media to IL media were observed for [MoCl6]3- in the UV/vis spectrum [34].
Utilising the thymol blue indicator, as shown in Figure 3(B) and 3(C), the
FeCl3.6H2O does in fact liberate protons, as evidenced by the increase in the band at
544 nm. Using a calibration curve derived from data obtained from Figure 2, the
intensity of this band suggests that approximately 0.8 protons are liberated per
FeCl3.6H2O dissolved in the ionic liquid. We assume that this is related to the
9
dissociation of 1 proton described by the following equation (where the 0.2 remaining
protons suggests that full dissociation from the dissolved Fe salt does not occur):
C mpyrTFSI
[Fe(H2O)6]3+ 4   
→ [Fe(H2O)5(OH)]2+ + H+
Equation 1
3. Cyclic voltammetric study of H+ in C4mpyrTFSI
In order to understand the complex electrochemical behaviour of the Fe2+/3+
process in C4mpyrTFSI when the Fe source is the FeCl3.6H2O and since the UV/vis
study demonstrates there is H+ in solution, voltammograms of HTFSI in C4mpyrTFSI
were recorded.
The voltammograms of HTFSI in the concentration range of 15 to 100 mM in
C4mpyrTFSI recorded at a Pt electrode are shown in Figure 4 and typical data
obtained is summarised in Table 2 with full data shown in the supporting information.
In order to record reproducible and stable voltammograms the Pt electrode was
preoxidised at +1.5 V for 15 seconds prior to scanning. This electrode pre-treatment
regime has previously been utilised by Stojek and Osteryoung et al. for studies of H+
in HClO4/water medium [35] and Compton et al. for H2 oxidation studies in ionic
liquids [36]. Although Compton et al. used an anodic oxidation potential of +2.0 V
for 30 seconds in a variety of ILs [36], we find that in the C4mpyrTFSI medium an
oxidation potential of +1.5 V for 15 seconds is sufficient to activate the Pt surface.
The voltammograms obtained at the activated Pt electrode show a reduction
process, I, at ca. -0.48 V and a related oxidation process, II, at ca. -0.38 V. An
additional, unrelated reductive process, III, is also observed at ca. 0.25 V. The
Ep
values for the related processes I and II are in the range of 65 to 120 mV over the scan
rate range of 20 to 200 mV s-1 for the 15 mM solution. Analysis of the peak current
10
ratios, show that at low scan rates values approaching unity are observed which
quickly decay as scan rate is increased (Table 2). Similar results are obtained for the
higher concentrations of HTFSI investigated (see supporting information).
Hardacre and Lagunas et al. have reported on a comprehensive study of H+
(introduced either as HTFSI or HCl) in the butyl-methylimidazolium TFSI
(C4mimTFSI) ionic liquid [37]. Based on this report and allowing for small variations
in redox behaviour due to solvation effects going from the C4mimTFSI to
C4mpyrTFSI ionic liquids the peak processes for HTFSI in C4mpyrTFSI can be
assigned.
Process I is attributed to the reduction of H+ to H2. Previous studies by
Silvester and Compton et al. on HTFSI reduction in C2mimTFSI and C4mimTFSI
hypothesised that the H+ reduction on Pt involves adsorption of H+ followed by
subsequent combination of the reduced H• to form H2 [37,38] or the reaction of the H•
with H+ and an other subsequent electron transfer to form the H2 as shown in the
following equations:
H+ + e– → Hadsorbed
Equation 2
2Hadsorbed → H2
Equation 3
Or Hadsorbed + H+ + e– → H2
Equation 4
For H+ reduction in C2mimTFSI and C4mimTFSI, the voltammetric processes
showed apparent quasi-reversibility which was attributed to the detection of two
overlapping processes [38]. In the C4mpyrTFSI medium the voltammetric profiles
show apparent reversibility, however upon closer inspection an additional process can
be detected as a shoulder at the beginnings of the reduction peak. Presumably, the
11
reduction of H+ in the C4mpyrTFSI medium follows a similar mechanism to that
proposed by Silvester and Compton et al., albeit with subtle differences most likely
related to kinetic differences between the H+ processes in the different ILs. Analysis
of the reduction peak currents as a function of square root of scan rate show that for
all the concentrations investigated, deviations from linearity occur at higher scan rates
(Figure 4(B)). Furthermore, for all the concentrations a non-zero intercept is observed
in the plot. This data is consistent with mechanism proposed by equations 2 to 4
where the combination of chemical reactions with electron transfer processes is
known to have influences in the electrochemical parameters. The related process II is
assigned to the corresponding oxidation to the electrogenerated H2 to form H+.
Turning to process III, previous studies have suggested that this is due to reduction of
a proton mediated anion oxidation product [37,38]. It is probable that a similar
process is occurring in the C4mpyrTFSI.
Looking more closely at the voltammograms of FeCl3.6H2O in C4mpyrTFSI
the processes observed can now be tentatively assigned. A voltammogram of
FeCl3.6H2O with overlaid data for HTFSI and FeCl3.6H2O (recorded in the positive
scan direction) are shown in Figure 4(C). As can be observed the FeCl3.6H2O
processes appear to be a mixture of the Fe2+/3+ couple and the H+/H2 processes.
Presumably, this is the cause of the shift in potentials observed for the ferric system.
It should be noted that the presence of H+ in the Fe solution voltammograms
could in principle arise from trace water in the IL. Voltammetric data for deliberately
added water in C4mpyrTFSI do show voltammetric processes attributable to H+.
These processes can only be detected after electrochemical decomposition of the
water at reductive potentials to generate protons as a reaction product. Additionally,
peak currents magnitudes approaching those observed in the FeCl3.6H2O data are
12
obtained only at significantly high concentrations of added water, much higher thaen
would be expected for trace impurities in the ppm range.
4. H+ generation mechanism in C4mpyrTFSI
Based on the voltammetry and UV/vis studies it is clear that the protons
detected in the IL solution are due to the disassociation of the FeCl3.6H2O complex
upon dissolution. Since the UV/vis data suggests that the [Fe(H2O)5(OH)]2+ species is
formed in solution and the voltammetry studies suggests that no Fe-TFSI complex is
formed. It has previously been shown for a La-TFSI complex that stronger bonding
between the metal centre and water O is observed when compared to the metal to
TFSI- bond thus showing that the TFSI- is a poor anion for metal species [39]. Based
on the data presented in this paper, this also appears to be the case for the Fe-TFSI
bond when water is present. Thus, it appears as if the [Fe(H2O)]3+ cation in the
FeCl3.6H2O complex does not undergo any ligand exchange with the TFSI- anions
upon dissolution, rather decomposition of the coordinated waters. Based on the
experimental data obtained the previously stated reaction in equation 1 is proposed for
the proton generation mechanism.
The [Fe(H2O)5(OH)]2+ species formed in C4mpyrTFSI solution is most likely
counter balanced with a mixture of Cl- and/or TFSI- anions via electrostatic
interactions. Presumably, the formation of this species is the cause for the shift in
redox potentials in the IL compared to the aqueous [Fe(H2O)6]2+/3+ system as noted
previously. This shift in potentials renders the iron oxidant incapable of significant
oxidation of the monomer and opens up the proton as the more likely candidate for
causing the polymerisation.
13
The other consideration is that the reaction to form, for example, polypyrrole
also involves the liberation of a proton. The simplified reaction for this process is:
C4H5N → -[C4H3Nx+]- + 2 H+ + (2-x)e–
Equation 5
We propose that once the oxidation is started, this continued generation of protons
means that the reaction self propagates and is sustained. Perhaps supporting this
assertion is the observed lack of reaction of aniline with HTFSI (even though the
oxidation potential lies between that of pyrrole and edot). Formation of polyaniline
involves the initial consumption of protons to form anilinium [40] and consequently
the reaction may be quenched in this case.
Previous studies by Hardacre and Lagunas et al. on HCl dissolved in TFSI
based ILs have shown a number of processes related to Cl– anions and the stable,
protonated HCl2– species [37]. In the voltammetry studies of FeCl3.6H2O in
C4mpyrTFSI, no evidence of these processes were observed. Thus the H+ liberated
upon dissolution most likely forms a H-TFSI species compared to the stable HCl2–
species.
5. H+ mediated polymerisation of pyrrole and edot in C4mpyrTFSI
It has previously been postulated that the chemical polymerisation of pyrrole
and edot can be accomplished by the reaction with Fe3+/2+ redox couple [22-24].
However in the case of FeCl3.6H2O as the source of the Fe couple this is not a
straightforward a hypothesis as proposed since both the voltammetry and UV/vis
studies show that upon dissolution of FeCl3.6H2O in C4mpyrTFSI, H+ are liberated
from the [Fe(H2O)6]3+. In principle, since the proton reduction process is more
14
cathodic than the Fe2+/3+ couple it is feasible that the H+ is the more likely reacting
species in the monomer polymerisation reaction. In order to test this hypothesis,
solutions of pyrrole and edot were reacted with deliberately added H+ in C4mpyrTFSI.
10, 50 and 100 mM solutions of HTFSI in C4mpyrTFSI were reacted with solutions of
pyrrole and edot and the polymerisation time observed and recorded. The
experimental conditions are summarised in Table 3.
Under all conditions tested as shown in Table 3, polymerisation of both
pyrrole and edot by HTFSI to the corresponding polypyrrole and pedot occurs. It is
important to note that these reactions were performed in the absence of oxygen in the
Argon filled glovebox. In order to obtain further information regarding the
morphology of the conducting polymers formed using the acid oxidation method,
scanning electron microscope (SEM) images of the conducting polymers were
collected. Samples were prepared by deposition of the polymers onto filter paper
following the procedure outline in the experimental section. Care was taken to ensure
that the polymers were free from IL residues. As a secondary confirmation that the
polymers were formed all SEM images were obtained without the use of conductive
coatings. Since the SEM imaging requires that the samples being imaged are
conductive, the absence of any conductive coatings provides secondary evidence of
the formation of the conductive polymer. There is a known reaction of acid with the
pyrrole monomer to form the polymer in the presence of oxygen [41]. In this work
[41], 2 mL of pyrrole dissolved in 10 mL of ethanol was added with stirring to a petri
dish containing 20 mL of 1.9M sulfuric acid. The solution was left standing for 10-15
h. The resulting films, however, were yellow in colour and insulating at 1.5x10-11
S/cm [41]. This is, obviously, not the same reaction occurring as seen in the ionic
liquid in the absence of oxygen.
15
Representative SEM images of polypyrrole and pedot are shown in Figure 5.
The SEM images show that in the case of the polypyrrole a relatively smooth
morphology is formed and the polymer coats the filter paper fibres evenly (Figure
5(A)). In the case of the pedot, the polymer again coats the filter paper fibres evenly,
although evidence for some nodular morphology superimposed onto the smooth films
is present (Figure 5(B)).
IR spectra of the polypyrrole and pedot synthesised by the acid route were
measured and are shown in Figure 6. The polypyrrole and pedot spectra show distinct
differences when compared to the spectra of the neat monomers. This result is to be
expected due to the polymerisation and these spectral changes has previously been
observed for both pyrrole and edot and their corresponding polymers [42,43]. Overall
the IR spectra are complicated due to the number of active bands observed. In the case
of the pedot, bands due to the C=C, C-C and C-S can be observed in the 1600-800 cm1
region and are comparable in position to those reported previously for pedot [42].
Interestingly, a distinct broad band at 1721 cm-1 which has previously been attributed
to the doped state of pedot (1722 cm-1) is also observed. Comparison of the pedot IR
spectrum with those of HTFSI [44] also provides tentative evidence for the TFSI
anion present within the polymer due to the observation of bands which can be
attributed to the SNS vibrations of the TFSI anion. Tentative evidence for the cation
also being present in the polymer matrix is suggested by observation of bands in the
3000-2800 cm-1 region which can be attributed to the C-N vibrations. The data would
then suggest that the dopant in the pedot is C4mpyrTFSI. It should be noted that
spectra of samples of the polymer soaked in isopropanol for 12 hours and then
allowed to dry show no difference from those shown in Figure 6 which were recorded
after washing and drying the sample. Similar results were obtained for the polypyrrole
16
IR data where again significant differences between monomer and polymer are
observed. Bands which can be attributed to the C-H in plane and out of plane modes
of the pyrrole ring [42] can be observed as well as those possibly due to TFSI- [44].
Again, tentative evidence for the cation is again observed in the 3000-2800 cm-1
region from bands attributed to the C-N vibrations.
Overall the IR spectra show that the polymerisation reaction has occurred as
well as suggesting that for both the pedot and polypyrrole polymers that C4mpyrTFSI
may also be incorporated into the polymer structure.
Conclusions:
The voltammetric study of FeCl3.6H2O in C4mpyrTFSI shows that the
response obtained is complicated with both electrode and scan direction dependence.
The root cause of the complex voltammetry lies in the detection of H+ redox processes
which overlap the Fe redox processes. Voltammetric studies of HTFSI in the
C4mpyrTFSI ionic liquid show that the H+/H2 reaction shows apparent reversibility.
The H+/H2 redox mechanism follows that previously postulated for H+ in TFSI based
ionic liquids [37,38].
UV/vis spectroscopy studies of the FeCl3.6H2O in C4mpyrTFSI show that
upon dissolution H+ is produced from decomposition of the coordinated H2O. This
then is the cause of the complex voltammetry observed. Additionally, the UV/vis
study shows that the [Fe(H2O)5(OH)]2+ species is formed in solution, thus the iron
complex deprotonates and liberates H+ upon dissolution.
Studies of the polymerisation of pyrrole and edot show that H+ itself can be
used as the chemical oxidant for the polymerisation reaction. Interestingly, data
17
obtained implies that the Fe may not itself participate in the chemical polymerisation
when H+ is present in solution.
The observation of the proton mediated chemical polymerisation of pyrrole
and edot is an important find. Use of Fe, Ag or Au salts as oxidants for chemical
synthesis of conducting polymers from ILs results in additional steps required for
purification of the synthesised polymer and removal of not only the IL residues but
also the inorganic salts. However, by using the HTFSI acid in a TFSI based IL,
simpler purification processes for the synthesised polymers can be envisaged with the
use of less chemicals.
Acknowledgements:
The authors thank the Australian Defence Force Capability and Technology
Demonstrator program for funding this work. The authors also gratefully
acknowledge Mr Pon Kao, Dr John Ward and Mr Mark Greaves for provision of SEM
images.
Experimental:
1-3,-ethylenedioxothiophene (edot, Aldrich, 97 %), FeCl3.6H2O (Riedel-deHaan, 99 %), FeCl2.4H2O (BDH, 98%), bistrifluoromethanesulfonimide,
Fe(NO3)3.9H2O (Aldrich, 98%), HTFSI (Aldrich, 95 %), Butyl-methylpyrrolidinium
bistrifluoromethanesulfonimide, C4mpyrTFSI (Merck synthesis grade) and thymol
blue pH indicator (Aldrich, 95 %) were commercially available and used as received.
Pyrrole (Merck, 97 %) was distilled under inert atmosphere prior to use.
Cyclic voltammetry (CV) measurements were recorded on either a µ-AutoLab
III potentiostat or a AutoLab PGSTAT302N potentiostat, both operated by GPES
18
(ver. 4.9) software. All CV measurements were performed in a conventional threeelectrode cell comprising either a Pt (7.8 × 10-3 cm2) or glassy carbon (GC, 0.0707
cm2) working electrodes and a large surface area wound Pt wire counter electrode.
The reference electrode consisted of a silver wire immersed in a 0.01 M solution of
Ag(CF3SO3) (hereafter abbreviated to AgOTf) in C4mpyrTFSI, separated from the
bulk solution by a porous glass frit, as previously reported [45]. The IR-drop was
uncompensated. Unless otherwise stated, all peak potential data reported in this paper
are quoted versus the Ag/AgOTf couple. Voltammetric measurements were made in a
high purity Ar glovebox at ambient temperature (22 ± 2 °C). The Pt and GC working
electrodes were polished on 0.3 µm alumina and washed with distilled water,
sonicated and dried prior to use.
Conducting polymers chemically synthesised in C4mpyrTFSI using HTFSI as
an oxidant were deposited onto filter paper and washed thoroughly using isopropanol
to remove RTIL traces. The filter paper was then allowed to dry for 12 hours prior to
imaging. Scanning Electron Microscopy (SEM) images were recorded using a Philips
XL30FEG Scanning Electron Microscope. IR spectra were obtained on a PerkinElmer 400 FTIR/NIR spectrometer using a Golden Gate ATR attachment. All spectra
were obtained at a resolution of 4 cm-1 and background corrected. UV/visible spectra
were recorded on a Cary-Varian 50 Bio spectrophotometer using a 0.5 cm pathlength
cell. All spectra were run against an IL background and baseline subtracted. All ionic
liquid samples were prepared and sealed within the Ar glovebox.
References:
1.
G.A. Snook, P. Kao, A.S. Best, Journal of Power Sources, 2011, 196, 1-12.
19
2.
C. Peng, S. W. Zhang, D. Jewell and G. Z. Chen, Progress in Natural Science,
2008, 18, 777-788.
3.
G. A. Snook and G. Z. Chen, Journal of Electroanalytical Chemistry, 2008,
612, 140-146.
4.
G. A. Snook, G. Z. Chen, D. J. Fray, M. Hughes and M. Shaffer, Journal of
Electroanalytical Chemistry, 2004, 568, 135-142.
5.
G. A. Snook, C. Peng, D. J. Fray and G. Z. Chen, Electrochemistry
Communications, 2007, 9, 83-88.
6.
M. Q. Wu, G. A. Snook, V. Gupta, M. Shaffer, D. J. Fray and G. Z. Chen,
Journal of Materials Chemistry, 2005, 15, 2297-2303.
7.
C. Arbizzani, M. Mastragostino and F. Soavi, Journal of Power Sources,
2001, 100, 164-170.
8.
L. Z. Fan and J. Maier, Electrochemistry Communications, 2006, 8, 937-940.
9.
S. A. Hashmi, R. J. Latham, R. G. Linford and W. S. Schlindwein, Polym.
Int., 1998, 47, 28-33.
10.
S. A. Hashmi and H. M. Upadhyaya, Solid State Ionics, 2002, 152, 883-889.
11.
M. Mastragostino, C. Arbizzani and F. Soavi, Solid State Ionics, 2002, 148,
493-498.
12.
M. Mastragostino, F. Soavi and C. Arbizzani, Advances in Li-ion Batteries,
2002.
13.
A. F. Diaz, K. K. Kanazawa and G. P. Gardini, J. Chem. Soc.-Chem.
Commun., 1979, 635-636.
14.
M. Grzeszczuk and G. Bator, Pol. J. Chem., 2002, 76, 1143-1150.
15.
K. K. Kanazawa, A. F. Diaz, R. H. Geiss, W. D. Gill, J. F. Kwak, J. A. Logan,
J. F. Rabolt and G. B. Street, J. Chem. Soc.-Chem. Commun., 1979, 854-855.
20
16.
P. A. Topart and M. A. M. Noel, Anal. Chem., 1994, 66, 2926-2934.
17.
K. Lota, V. Khomenko and E. Frackowiak, Journal of Physics and Chemistry
of Solids, 2004, 65, 295-301.
18.
R. McNeill, D. E. Weiss, J. H. Wardlaw and R. Siudak, Aus. J. Chem., 1963,
16, 1056.
19.
J. Stejskal and R. G. Gilbert, Pure Appl. Chem., 2002, 74, 857-867.
20.
M. S. Cho, J. D. Nam, H. R. Choi, J. C. Koo and Y. Lee, in Asian Pacific
Conference on Fracture and Strength, eds. Y. J. Kim and H. D. Bae, Trans
Tech Publications Ltd, Jeju Isl, South Korea, 2004, pp. 641-645.
21.
S. Inagi and T. Fuchigami, Synthetic Metals, 2008, 158, 782-784.
22.
J. M. Pringle, O. Ngamna, J. Chen, G. G. Wallace, M. Forsyth and D. R.
MacFarlane, Synthetic Metals, 2006, 156, 979-983.
23.
J. M. Pringle, O. Ngamna, C. Lynam, G. G. Wallace, M. Forsyth and D. R.
MacFarlane, Macromolecules, 2007, 40, 2702-2711.
24.
Y. F. Yang and S. L. Mu, Electrochimica Acta, 2008, 54, 506-512.
25.
F. Endres, Abstr. Pap. Am. Chem. Soc., 2004, 228, U71-U71.
26.
Y. K. Koo, B. H. Kim, D. H. Park and J. Joo, Molecular Crystals and Liquid
Crystals, 2004, 425, 333-338.
27.
P. Kubisa, J. Polym. Sci. Pol. Chem., 2005, 43, 4675-4683.
28.
J. H. Mazurkiewicz, P. C. Innis, G. G. Wallace, D. R. MacFarlane and M.
Forsyth, Synthetic Metals, 2003, 135, 31-32.
29.
M. Yamagata, N. Tachikawa, Y. Katayama, T. Miura, Electrochim. Acta,
2007, 52, 3317
30.
A. J. Bard, L. R. Faulkner, Electrochemical methods: fundamentals and
applications. Wiley, New York, 2001
21
31.
L. Lopes, J. de Laat, B. Legube, Inorg. Chem., 2002, 41, 2505
32.
I-W. Sun, J. R. Sanders, C. L. Hussey, J. Electrochem. Soc., 1989, 136, 1415
33.
P. Day, C.K. Jørgensen, J. Chem. Soc., 1964, 6226
34.
T. B. Scheffler, C. L. Hussey, K. R. Seddon, C. M. Kear, P. D. Armitage,
Inorg. Chem., 1983, 22, 2099
35.
A. Jaworski, M. Donten, Z. Stojek, J. G. Osteryoung, Anal. Chem., 1999, 71,
243
36.
D. S. Silvester, K. R. Ward, L. Aldous, C. Hardacre, R. G. Compton, J.
Electroanal. Chem., 2008, 618, 53
37.
L. Aldous, D. S. Silvester, W. R. Pitner, R. G. Compton, M. C. Lagunas, C.
Hardacre, J. Phys. Chem. C, 2007, 111, 8496
38.
D. S. Silvester, L. Aldous, C. Hardacre, R. G. Compton, J. Phys. Chem. B,
2007, 111, 5000
39.
A. I. Bhatt, I. May, V. A. Volkovich, D. Collison, M. Helliwell, I. B. Polovov,
R. G. Lewin, Inorg. Chem., 2005, 44, 4934
40.
G.A. Snook, T.L. Greaves, A.S. Best, J. Mater. Chem. 2011, 21, 7622-7629
41. M. Salmon, K. K. Kanazawa, A. F. Diaz and M. Krounbi, Journal of Polymer
Science Part C-Polymer Letters, 1982, 20, 187
42.
H. Kato, O. Nishikawa, T. Matsui, S. Honma, H. Kokado, J. Phys. Chem.,
1991, 95, 6011
43.
S. V. Selvaganesh, J. Mathiyarasu, K. L. N. Phani, V. Yegnaraman, Nanoscale
Res. Lett., 2007, 2, 546
44.
I. Rey, P. Johansson, J. Lindgren, J. C. Lassègues, J. Grondin, L. Servant, J.
Phys. Chem. A, 1998, 102, 3249
22
45.
G. A. Snook, A. S. Best, A. G. Pandolfo, A. F. Hollenkamp, Electrochem.
Commun. 2006, 8, 1405
23
Tables:
Table 1: Cyclic voltammetry parameters obtained for FeCl3.6H2O in C4mpyrTFSI at
a GC electrode at 22 °C
Concentration
10 mM
5 mM
Scan
rate /
mV s-1
Epred /
mV
Epox /
mV
Emid /
mV
Ep /
mV
Jpox/Jpred
Epred /
mV
Epox /
mV
Emid /
mV
Ep /
mV
Jpox/Jpred
20
40
60
80
100
120
140
160
180
200
-624
-623
-624
-630
-632
-631
-549
-546
-543
-540
-536
-532
-587
-585
-584
-585
-584
-582
75
77
81
90
96
99
0.78
0.72
0.65
0.70
0.69
0.62
-552
-568
-579
-587
-586
-594
-597
-587
-590
-598
-394
-364
-345
-337
-337
-323
-326
-282
-300
-289
-473
-466
-462
-462
-462
-459
-462
-435
-445
-444
158
204
234
250
249
271
271
305
290
309
0.36
0.39
0.47
0.45
0.41
0.54
0.39
0.31
0.39
0.52
24
Table 2: Cyclic voltammetry parameters obtained for HTFSI in C4mpyrTFSI at a Pt
electrode at 22 °C
Concentration
15 mM
Scan rate
/ mV s-1
EpI / mV
EpII / mV
Emid /
mV
Ep /
mV
Jpred/Jpox
20
40
60
80
100
120
140
160
180
200
-466
-464
-475
-477
-479
-482
-483
-485
-483
-473
-400
-392
-383
-377
-375
-371
-365
-363
-351
-355
-433
-428
-429
-427
-427
-427
-424
-424
-417
-414
66
72
92
100
104
111
118
122
132
118
1.03
0.88
0.77
0.78
0.70
0.67
0.68
0.65
0.63
0.64
25
Table 3: Experimental conditions for synthesis of polypyrrole and pedot from
C4mpyrTFSI using HTFSI as chemical oxidant. All experiments performed in the
absence of air in an argon filled glovebox at 22 °C.
†
Monomer
HTFSI
concentration /
concentration /
mMol L-1
mMol L-1
pyrrole
10, 50 and 100
edot
10, 50 and 100
Temperature Polymerisation
/ °C
time†
10, 50 and 100
22
10 - 60 seconds
10, 50 and 100
22
10 -30 seconds
Note all reactions were allowed to proceed for 12 hours to ensure that reaction was
complete prior to SEM imaging and IR spectroscopic analysis
26
Figures:
Figure 1: Cyclic voltammograms of FeCl3.6H2O dissolved in C4mpyrTFSI recorded
at 22 °C under inert atmosphere. (A) 10 mM FeCl3.6H2O at a GC electrode, ν = 20 100 mV s-1 in 20 mV s-1 intervals; (B) 5 mM FeCl3.6H2O at a GC electrode, ν = 20 100 mV s-1 in 20 mV s-1 intervals; (C) 10, 20, 40 and 60 mM FeCl3.6H2O at a Pt
electrode, ν = 20 mV s-1; (D) 15 mM FeCl3.6H2O at a Pt electrode, ν = 20, 50, 100
and 200 mV s-1; (E) 10 mM FeCl2.4H2O at a Pt electrode, ν = 20 - 200 mV s-1 in 20
mV s-1 intervals.
27
Figure 2: UV/vis spectra of (A) 0.1 mM thymol blue in HCl/water at a concentration
of 5, 10, 25, 50 mM of acid and (B) 0.2 µM thymol blue with HTFSI dissolved in
C4mpyrTFSI at a concentration of 5, 10, 25 and 50 mM.
28
Figure 3: UV/vis spectrum of (A) FeCl3.6H2O dissolved in C4mpyrTFSI measured at
22 °C (no thymol blue); (B) 0.4 µM thymol blue with 5, 10, 20 and 50 mM
FeCl3.6H2O showing a zoomed image around the protonated thymol blue peak (at ca.
544 nm) region.
(A)
(B)
29
Figure 4: Cyclic voltammograms of HTFSI in C4mpyrTFSI recorded at 22 °C under
inert atmosphere. (A) 15, 20, 40, 60, 80 and 100 mM HTFSI at a pre-treated Pt
electrode. (B) Plot of process I reduction peak currents as a function of square root of
scan rate (ν1/2) for 15 mM and 20 mM data sets and (C) overlay of CVs obtained for
15 mM HTFSI and 15 mM FeCl3.6H2O recorded in the forward and backwards
direction.
30
Figure 5: Scanning Electron Microscope images of (A) polypyrrole and (B) pedot
deposited onto filter paper.
31
Figure 6: ATR-FTIR spectra of polypyrrole and polyedot synthesised by the reaction
of monomer with HTFSI in C4mpyrTFSI.
32