Activation Barriers Provide Insight into the Mechanism of Self

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Activation Barriers Provide Insight to the Mechanism of Self-Discharge in Polypyrrole Henrik Olsson, Maria Stromme, Leif Nyholm, and Martin Sjödin J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/jp510690p • Publication Date (Web): 02 Dec 2014 Downloaded from http://pubs.acs.org on December 3, 2014

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The Journal of Physical Chemistry

Activation Barriers Provide Insight to the Mechanism of Self-Discharge in Polypyrrole Henrik Olsson,a Maria Strømme,a Leif Nyholm,b Martin Sjödin a* a

Nanotechnology and Functional Materials, The Ångström Laboratory, Uppsala University, Box 534, 751 21 Uppsala, Sweden

b

Department of Chemistry – Ångström, The Ångström Laboratory, Uppsala University, Box 538, 751 21 Uppsala, Sweden

ACS Paragon Plus Environment


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ABSTRACT

Conducting polymers are envisioned to play a significant role in the development of organic matter based electrical energy conversion- and storage systems. However, successful utilization of conducting polymers relies on a fundamental understanding of their inherent possibilities and limitations. In this report we have studied the temperature dependence of the self-discharge in polypyrrole and show that the rate of self-discharge is kinetically controlled by a polymer intrinsic endergonic electron transfer reaction forming a reactive intermediate. We further show that this intermediate is intimately linked to a process known as overoxidation. This process is general for most, if not all, p-doped conducting polymers irrespective of medium. The results herein are therefore expected to significantly impact the development of future energy storage systems with conducting polymer based components.

KEYWORDS Polypyrrole, self-discharge, activation-controlled faradaic reaction, stability, maleimide, degradation


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Introduction With the transition to intermittent renewable energy sources and to electrical and hybrid vehicles, triggered by environmental concerns related to the combustion of fossil fuels, as well as with the improved well-being in the developing world comes an increasing need for electrical energy storage. Batteries and capacitors constitute the two main technologies for reversible electrical energy storage and, of these two; batteries are likely to have the greatest impact on the development of environmentally benign energy conversion systems for large and intermediate scale applications due to their superior energy storage capacity. Current day secondary batteries are, however, far from environmentally benign as they are made from non-renewable resources of limited supply in addition to being energy demanding both in their rawmaterial extraction and in their refinement.1 The inorganic materials used in rechargeable batteries also add a large part to the final cost of the batteries.2 These issues are currently hampering the development of energy- and cost efficient energy conversion systems for a carbon dioxide neutral society and research on organic alternatives that can replace the inorganic materials has therefore gained impetus. Conducting polymers is one class of organic battery materials that has been proposed and tested as both cathode and anode materials for rechargeable energy storage.3-8 Such materials rely on a process, commonly referred to as doping, in which the polymer is oxidized (p-doping) or reduced (n-doping) with a concomitant uptake of negative and positive ions, respectively, from the electrolyte solution. The doping process in conducting polymers generally gives a capacitive charge-discharge response, in contrast to common battery materials that rely on redox processes with more or less well-defined potentials during charge and discharge. The charge capacities of conducting polymers thus depend on the level of doping that can be realized without reaching degrading potentials. For p-doped materials, including polyacetylene, polyparaphenylene, polypyrrole (PPy) and polythiophene overoxidation limits the achievable charge capacities and the use of conducting polymers for high capacity applications. Poor



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stability and self-discharge have also inhibited the use of conducting polymers in battery applications.912

Figure 1. The proposed reaction scheme for PPy self-discharge. An initial, and polymer intrinsic redox reaction leads to an unstable intermediate (Int+) that reacts with its surroundings to form maleimide and other degradation products while reducing the polymer. The reaction rate for producing the intermediate state increases with polymer potential by decreasing the potential barrier for the initial reaction. One strategy to overcome the limited charge storage capacity of known conducting polymers, which has recently gained considerable attention,13-18

is to incorporate high capacity redox groups in the

conducting polymer matrix, thereby increasing the specific capacity while retaining benefits of electronic conduction and low solubility. However, with this strategy, self-discharge and limited stability of the conducting polymer matrix is still a concern; even though the charge capacity is mainly carried by the high capacity redox group, these conducting polymer based problems will be conserved. In order to be able to address these issues on a synthetic strategic basis, a better and more fundamental knowledge of the self-discharge mechanism is desirable. We have therefore undertaken an investigation of the self-discharge behavior of PPy in a PPy/cellulose composite material, previously investigated as electrode material in both energy storage19-23 and ion-exchange24,25 applications. We have very recently shown that the self-discharge in this system, when used as positive electrode, is due to a faradaic reaction that is likely to be of polymer intrinsic origin.26,27 A model was proposed where an initial


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intrinsic redox reaction forms a highly reactive intermediate state (Int+) that reacts subsequently and kinetically non-limiting to form maleimide, protons as well as other degradation products (Fig. 1). In the present work we study the temperature dependence of the self-discharge reaction of oxidized PPy to reach the activation barriers involved in the initial redox reaction. From the activation energy, we estimate the redox potential of the intermediate and show that the faradaic self-discharge has the same origin as the reaction responsible for overoxidation commonly encountered in conducting polymers28-32 and is thus general for most, if not for all, p-doped conducting polymers. Experimental Synthesis of PPy/cellulose composites Cellulose from the Cladophora sp. algae was extracted as described elsewhere.33 Pyrrole was purchased from Sigma-Aldrich, while Iron (III) chloride hexahydrate (>99%), Tween 80 and sodium chloride (>99.5%) were purchased from VWR international. The chemicals were used as received. The cellulose was dispersed by sonicating 300 mg cellulose in 60 mL of water for 10 minutes. 1.5 mL pyrrole and one drop of Tween 80 was dissolved in 50 mL of 0.5 M HCl. An iron salt solution was prepared by dissolving 12.86 g FeCl3 in 100 mL 0.5 M HCl. After mixing the pyrrole solution with the cellulose dispersion, and stirring for 5 minute, the iron solution was added to the mixture. After 30 minutes, the polymer composite was washed with 5 L 0.5 M HCl and 1 L 0.1 M NaCl. The mixture was then sonicated for an additional 2 minutes before the polymer composite was dried in the funnel to form a black wet composite cake. The cake was pressed in a vise overnight and dried for at least 3 days at ambient conditions resulting in a black paper-like material. Electrochemical measurements The electrochemical experiments were carried out using an Autolab PGSTAT302N potentiostat (Metrohm-Autolab, the Netherlands) and in all measurements an Ag/AgCl reference electrode with 3M NaClaq was used as reference electrode and a Pt-wire served as counter electrode. The 1M NaClaq



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electrolyte used was thoroughly purged with N2 prior to measurements and was kept under argon atmosphere throughout the measurements. A constant electrolyte volume of 20 mL was also employed. All samples were pre-cycled with cyclic voltammetry (CV) between -0.7 and 0.35 V vs. Ag/AgCl for 5 cycles, employing a 5 mV/s scan rate. The three-electrode self-discharge measurement procedure was designed to correspond to the positive electrode in a pouch cell device. In a previous study,26,27 it was found that the positive electrode in a two-electrode device attained 0.53 V vs. Ag/AgCl when the cell was charged to 0.6 V. For this reason, 0.53 V vs. Ag/AgCl was chosen as the starting potential for all three-electrode self-discharge measurements in this paper. The composites were held at 0.53 V vs. Ag/AgCl for 600 s in order to ensure that the samples were completely charged. The applied potential was then turned off and the potential was logged continuously. The self-discharge was left to continue at least until the sample had reached 0.33 V vs. Ag/AgCl (i.e. a self-discharge of at least 200 mV). During both charge and self-discharge, the temperature of the electrolyte was held constant using a circulating temperature bath and a thermostated glass cell (Metrohm-Autolab, the Netherlands). Self-discharge experiments were performed at temperatures between 2 °C and 75 °C. Overoxidation measurement The overoxidation of PPy was studied using 11 µm diameter carbon fiber microelectrodes covered with a thin layer of PPy, at ambient temperatures. PPy was prepared by electrochemical polymerization, using 50 mM pyrrole in 0.5 M HCl. The polymerization was carried out by scanning the microelectrode between -0.3 and 0.9 V vs. Ag/AgCl for one scan. The scan rate employed during the polymerization was 500 mV/s, and the polymerization was kept to a minimum in order to ensure a thin layer deposited on the electrode. After polymerization, the microelectrode was rinsed quickly with water and placed in a cell with 1.0 M NaCl. The PPy was first cycled between -0.7 and 0.45 V vs. Ag/AgCl in order to ensure full electroactivity of the polymer. The polymer formed on the electrode yielded a capacity of approximately 1.5 mC/cm2 based on the oxidative sweep in the CV. After this, the polymer was


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overoxidized32 by executing a CV sweep to degrading potentials (> 1.5 V vs. Ag/AgCl). The scan rate during overoxidation was varied, and a newly polymerized electrode was used for each overoxidation. Results Self-discharge Fig. 2 shows the measured self-discharge at various temperatures. From these experiments it is clear that the self-discharge rate is higher at higher potentials and we have previously shown that this dependence agrees well with a faradaic self-discharge mechanism.26,27 There is also a clear increase in the self-discharge rate with temperature indicating an activated reaction.

Figure 2. Self-discharge curves for PPy/cellulose composite samples in 1M NaClaq solution at different electrolyte temperatures. A faster discharge is seen at higher temperatures indicating a thermally activated reaction. In order to separate the temperature dependence and the potential dependence it is convenient to evaluate apparent rate constants for the self-discharge reaction at both different temperatures and different potentials. These apparent rate constants were determined from the self-discharge experiment by evaluating the rate of potential drop (dE/dt) at different potentials and converting the potential drop to a current (dQ/dt) through



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,

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(1)

where C is the capacitance value for the material evaluated as a mean from potential steps (charging) for the samples used in the study.34 A mean value was used in order to minimize the variation seen in the capacitance evaluation.

Figure 3. Cyclic voltammogram of the PPy/cellulose composite recorded at 5 mV/s in 1M NaClaq solution at 20 °C. The CV was integrated to obtain the capacity at different potentials. The colored regions indicate the time-scales (1/kapp at 20 °C) for the faradaic and polymer intrinsic mechanism found to be the dominant cause of self-discharge. Assuming a first order reaction with respect to charge an apparent rate constant, kapp, for the selfdischarge reaction was evaluated through

,

(2)

where Q is the charge per unit mass (C/g) of the material at the potential where dE/dt was evaluated. Q was determined from CV measurements (Fig. 3) by integration of the oxidative current response up to the various potentials (e.g. the self-discharge rate at 0.4 V vs. Ag/AgCl uses a Q from an integration of the positive current in an oxidative CV sweep up to 0.4 V vs. Ag/AgCl). The resulting apparent rate


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constants for the self-discharge at 20 °C are presented in Table 1 and in Fig. 4. At this temperature the potential dependence of the apparent rate constant was well described by

, 293 ∙ ∙

(3)

(solid line in Fig. 4), with the parameters A = 8·10-11 s-1 and b = 25 V-1. An exponential potential dependence of the observed apparent rate constant was found at all temperatures and the apparent rate constants at other temperatures are given in the supporting material.

Figure 4. Apparent rate constants (open circles) evaluated at different potentials from self-discharge experiments for PPy/cellulose composite samples in 1M NaClaq solution at 20 °C together with a fit to equation 3 (solid line) with A = 8*10-11 s-1 and b = 25 V-1 as evaluated from the fit.



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Table 1. Charge consumption rates, amount of charge and evaluated apparent rate constants at 20°C at different potentials during self-discharge in 1M NaClaq solution. E vs. Ag/AgCl (V) 0.32 0.34 0.36 0.38 0.40 0.42 0.44 0.46 0.48 0.50

dQ/dt (A/g) 4.26 · 10-5 7.62 · 10-5 1.22 · 10-4 2.06 · 10-4 3.56 · 10-4 6.00 · 10-4 9.77 · 10-4 1.59 · 10-3 2.65 · 10-3 4.59 · 10-3

Q (C/g) 170 174 177 181 184 188 192 196 200 204

kapp (s-1) 2.50 · 10-7 4.38 · 10-7 6.87 · 10-7 1.14 · 10-6 1.93 · 10-6 3.19 · 10-6 5.09 · 10-6 8.11 · 10-6 1.33 · 10-5 2.25 · 10-5

The temperature dependence of the apparent rate constants was evaluated as described above at specific potentials. Each temperature point thus corresponds to one self-discharge experiment from which ten potential points were evaluated. In total, seven different temperatures are included in this work which thus includes data from seven independent self-discharge experiments. The temperature dependence of the apparent rate constant was fitted to the Arrhenius equation,

∙

(4)

and the activation energy, Ea, was determined for each of the ten potentials. In this equation is a temperature independent constant and kB is the Boltzmann constant. The activation energies obtained from the fits showed a linear decrease with increasing potential, cf. solid circles in Fig. 5.


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Figure 5. Activation energies (solid squares) evaluated from the temperature dependence of apparent rate constants as well as evaluated redox potentials for the electron donor (intermediate) (open circles) as function of potential. Solid lines serve as guides for the eye. For endergonic electron transfer reactions a relatively small error is introduced if the activation energy is taken to be the driving force for the reaction in the Marcus treatment of electron transfer reactions (see discussion). The driving force for a one-electron transfer reaction is given by Eq. 5, where E0(A/A+) and E0(D/D+) are the standard potentials for the electron acceptor and the electron donor, respectively

!" # nF& # A/A) # D/D) +

(5)

Here !" # is the Gibbs free energy of the electron transfer reaction, n is the number of moles of electrons exchanged in the reaction and F is the Faraday constant. We can determine the standard potential for the reaction intermediate, in this case the electron donor, if we know the potential of the acceptor. However, since the acceptor is the doped state of PPy we cannot assign a unique reduction potential for the acceptor but rather define an acceptor Fermi level. The Fermi level of the conducting polymer is evaluated from the potential versus the reference electrode, in our case


11


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the Ag/AgCl electrode, and hence we can evaluate the standard potential of the intermediate against that reference at each measured potential through Eq. 6, where ∆G0 is the driving force for the electron transfer reaction as evaluated from the activation energy of the apparent rate constant and E is the measured potential at which the activation energy was evaluated. It should be noted that, in this treatment, we assume that electron transfer reactions from the intermediate to polymer states with energies above the Fermi level can be neglected

# ,/D)

-. / 01

4

+ ≈ 01 + .

(6)

In Eq. 6, Ea is in joules, while we will use eV throughout the rest of this paper. The evaluated redox potentials for the intermediate is shown in Fig. 5 (open circles) as function of potential at which the activation energy and redox potential for the intermediate was evaluated. Contrary to the activation energy, the redox potential for the intermediate is relatively constant, with a mean value of 0.89 V vs. Ag/AgCl, despite the complexity of the studied system. At this potential the apparent rate constant at 20 °C is found to be 0.43 s-1 using the potential dependence of the apparent rate constant given in Eq. 3 above. Overoxidation When potentials higher than about 0.5 V vs. Ag/AgCl are applied to PPy the polymer degrades in a process referred to as overoxidation. In order to reach kinetic information of this reaction cyclic voltammograms were recorded at different scan rates.


12

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Figure 6. Cyclic voltammetry of PPy covered carbon microelectrodes in 1M NaClaq solution at 1, 10 and 100 Vs-1 in black, blue and red respectively. The currents are normalized with respect to scan rate. The vertical line corresponds to the formal potential of the intermediate state during self-discharge as evaluated from the activation energy of self-discharge. Fig. 6 shows the normalized current response at 1, 10 and 100 V/s from PPy covered glassy carbon microelectrodes. A clear shift in peak position towards higher potentials is observed with increasing scan rate. This shift is found to be very well described by a kinetically limited redox reaction where a linear increase of the peak potential, Ep, with the logarithm of the scan rate,5, is expected (Fig. 7) as given by Eq. 7.

0 +

7 1901

:;

19015 7 0? is the rate constant connected to overoxidation, R is the

gas constant and 9 is the charge transfer coefficient. From the scan rate dependence of Ep, and the derived # value, a rate constant of 9 s-1 was found. Discussion From Fig. 4 an exponential increase of the apparent rate constant with potential is evident and in agreement with the previously inferred rate limiting faradaic self-discharge mechanism. However, the redox reaction is merely the initial step in the complex self-discharge reaction and the formed intermediate reacts further to form several degradation products, including maleimide and protons (Fig. 1). The apparent rate is thus a measure of the rate for the complete reaction and hence composed of several elementary reaction steps, the initial step being an electron transfer reaction. From the temperature dependence of the over-all rate constant for selfdischarge the activation energies was found to increase linearly, from 0.35 eV to 0.55 eV, with


14

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potential. The potential dependence of the activation energy suggests that the activated reaction is in fact the redox reaction indicated in Fig. 1, and the magnitude of the activation energy indicates that this redox reaction is endergonic as an exergonic reaction would give unrealistically high reorganization energy to account for the experimentally determined activation energies (see below). Based on the activation energy, the redox potential for the intermediate was evaluated to 0.89 V vs. Ag/AgCl. During the self-discharge, where the PPy Fermi level dropped from the initial 0.53 V vs. Ag/AgCl to 0.3 V vs. Ag/AgCl, the Gibbs free energy for the initial step in fig. 1 increased from +0.33 eV to +0.56 eV. The redox reaction thus becomes more and more energetically unfavorable, as the self-discharge proceeds, with increased activation energy as a direct result of the decreased potential. At no stage during self-discharge is the initial redox reaction energetically favorable. Therefore, the back-reaction (rate constant kb) must be faster than the forward reaction (rate constant kf) and we may safely introduce a steady state approximation for the concentration of the formed intermediate. The resulting expression for the observed rate constant (Eq. 8a) gives two limiting cases (Eq. 8b and 8c, respectively), one where the electron transfer back-reaction is much faster than the follow-up reaction (rate constant kd) and one where the opposite is true, i.e. the follow-up reaction is much faster than the electron transfer back reaction.

E E

@ABCD @AAB E F)EG H

(8a)

G

E E

E

IBBJKLMLNILKO E F)EG ≈ EF C ≈ C H

G

E E

MLOLALPQ R@B S E F)EG ≈ H

G

(8b)

H

EF EG EG

≈ T

(8c)


15


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In both cases the observed rate constant is expected to show an exponential increase with potential (as discussed below) and qualitatively account for the observed potential dependence seen in Fig. 4. In the first limiting case the fast back-reaction ensures that equilibrium between the reactant state and the intermediate state is established. In this pre-equilibrium case the potential dependence is governed by the ratio of the forward and backward electron transfer rate constants. Starting from the Nernst equation (here written for an oxidation), this ratio is given by Eq. 9,

# +

UV W Y

EF EH

UV W

:;

(9a)

0 :;

Z

&X?+

:; &UBC+ # +

[

\

44 /

[

\

4/

(9b) [

\

4

,

(9c)

where E0 is the redox potential for the intermediate. In this case the exponent b (from Eq. 3) equals 39 V-1, in clear contrast to the experimentally determined value of b = 25 V-1 at 293 K, making a pre-equilibrium limit unlikely. In the second, more likely, limiting case the over-all reaction rate constant is given solely by the forward electron transfer rate constant kf, which can be treated with the standard Marcus equation for electron transfer reactions (Eq. 10).

] ^

∆"0 +`2 4 `

(10)

where λ is the reorganization energy, ∆G0 is the electron transfer driving force and ^ is a preexponential factor that includes the electronic coupling element. Hutchison et al.35 have


16

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computed the inner reorganization energy for the self-exchange reaction in PPy oligomers of various chain length. They found that the reorganization energy decreased linearly with the square root of the number of monomer units, but they also state that this dependence is anticipated to level off as the polymer length exceeds the localization length of the polaron. At oligomer lengths of four and six monomers the inner reorganization energy was 0.281 eV and 0.228 eV, respectively. This polymer length corresponds to the determined polaron extension that Brédas et.al.36 reported based on combined theoretical and spectroscopic results and the corresponding reorganization energy is therefore used as an estimate of the inner reorganization energy in the present study and constitutes a lower boundary for the total reorganization energy. The solvent reorganization often gives the dominant contribution to the total electron transfer reorganization energy and in water solution it can reach about 1 eV for small donor-acceptor molecules. In the less polar PPy environment the outer sphere contribution is expected to be significantly lower and Hutchison et al.35 even suggested that the inner reorganization energy should dominate the total reorganization energy and ~1 eV is therefore taken as an upper limit for outer sphere reorganization. The total reorganization energy for the electron transfer reaction should thus be within 0.2 – 1.3 eV. For this range of reorganization energies the experimentally determined activation energies (Fig. 5) can be used to compute the driving force for the electron transfer reaction based on Eq. 10, where Ea is given by (∆G0+λ)2/4λ. Fig. 8 shows the determined driving forces required to account for the activation energies 0.35, 0.45 and 0.6 eV. Note that the reaction is endergonic for the entire interval of activation energies and reorganization energies, thus justifying the steady state treatment of the intermediate.


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Figure 8. Required driving force as function of reorganization energy to account for activation energies of 0.35 eV (black squares), 0.45 eV (grey circles) and 0.6 eV (red triangles) as obtained from the temperature dependence of the self-discharge rate. It is seen that the initial rate-limiting step is energonic for all reasonable reorganization energies. The inset shows an estimate of the error introduced by using activation energies as the driving force for the involved reaction. We use the activation energy to estimate the redox potential of the intermediate (Eq. 6) and the inset in Fig. 8 shows the error introduced in this assumption for a reasonable reorganization energy of 0.7 eV as function of driving force. Clearly the introduced error is, at all driving forces, less than 0.06 eV and at the mean driving force of 0.44 eV during self-discharge, based on the estimates in this work, the error is 0.02 eV. In Fig. 5 a small, but still evident, drift of about 0.04 V over the investigated potential region of the evaluated redox potential for the intermediate can be seen that is most likely related to the error in the approximation used. A lower reorganization energy, as suggested,35 would yield an even smaller error and hence the reported intermediate redox potential is determined to 0.89 V vs. Ag/AgCl with an accuracy of 0.06 eV. Having established that the reduction potential for the intermediate is approximately at 0.89 V vs. Ag/AgCl, an inspection of the redox processes occurring at those potentials through cyclic


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voltammetry on PPy covered electrodes indicates that the only candidate reaction is, in fact, the overoxidation reaction commonly seen in conducting polymers. We therefore suggest that the main process leading to self-discharge is the same as the one leading to overoxidation. The selfdischarge is likely to be limited by kf, as discussed above, and the same holds for the overoxidation reaction. From the exponential potential dependence of the rate constant displayed in Fig. 4 (cf. Eq. 3), a formal rate constant of 0.43 s-1 was received through extrapolation of the curve in Fig. 4 to 0.89 V vs. Ag/AgCl while the rate constant evaluated from the scan rate dependence of the overoxidation peak potential (cf. Eq. 7 and Fig. 7) was 9 s-1 at the same potential. Considering the uncertainty in the extrapolation procedure as well as in the evaluated formal potential the rate constants are, within experimental uncertainty limits, identical and give further support for the self-discharge and overoxidation originating from the same process. As overoxidation is a general feature, at least for the most common p-doped conducting polymers30, this mechanism is likely to give significant contributions to the self-discharge rate also in polyacetylene, polyparaphenylene and polythiophene based energy storage systems. However, when the conducting polymer is used as a conducting matrix for high capacity pendant groups, this mechanism of self-discharge can be mastered by proper energy matching of the pendant redox group and the conducting polymer since the specific capacity, in that case, does not rely on polymer doping. This is possible because the presented mechanism has an activation barrier that decrease linearly with operating potential and hence gives a rate constant that decrease exponentially with decreasing potential. As indicated in Fig. 3, the upper potential where PPy can operate depends on the timescale as given by the specific application demands. At potentials where PPy shows maximum conductivity, i.e. at modest doping levels,37 the time scale for


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intrinsic self-discharge (Fig. 3) is however expected to be large enough to render the contribution from this mechanism insignificant. Conclusions In this report we have studied the temperature dependence of the self-discharge rate in polypyrrole covered cellulose composites in water-based electrolyte. The self-discharge reaction was found to be kinetically controlled by an endergonic redox reaction forming a reactive intermediate. The redox potential for the intermediate was evaluated to 0.89 V vs. Ag/AgCl and we suggest that the process involved in self-discharge is identical to overoxidation that is common for known p-doped conducting polymers. Supporting Information Rate constants for self-discharge at all temperatures studied. This material is available free of charge via the Internet at http://pubs.acs.org Corresponding Author * +46 18 471 73 30; E-mail: [email protected] Acknowledgements The authors thank the Swedish Foundation for Strategic Research (SSF), the Swedish Science Council (VR), the Bo Rydin Foundation, the Swedish Energy Agency and the European Institute of Innovation and Technology under the KIC InnoEnergy NewMat for their financial support of this work. (1)

Armand, M.; Tarascon, J. M. Building Better Batteries Nature 2008, 451, 652-

657. (2) Gaines, L.; Cuenca, R. Costs of Lithium-Ion Batteries for Vehicles, Center for Transportation Research, Argonne National Laboratory, U.S. Department of Energy, 2000.


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(3) Nyholm, L.; Nyström, G.; Mihranyan, A.; Strømme, M. Toward Flexible Polymer and Paper-Based Energy Storage Devices Adv. Mater. 2011, 23, 3751-3769. (4) Naegele, D.; Bittihn, R. Electrically Conductive Polymers as Rechargeable Battery Electrodes Solid State Ionics 1988, 28–30, Part 2, 983-989. (5) Osaka, T.; Ogano, S.; Naoi, K.; Oyama, N. Electrochemical Polymerization of Electroactive Polyaniline in Nonaqueous Solution and Its Application in Rechargeable Lithium Batteries J. Electrochem. Soc. 1989, 136, 306-309. (6) Qiu, W.; Zhou, R.; Yang, L.; Liu, Q. Lithium-Ion Rechargeable Battery with Petroleum Coke Anode and Polyaniline Cathode Solid State Ionics 1996, 86–88, Part 2, 903906. (7) Yang, L.; Qiu, W.; Liu, Q. Polyaniline Cathode Material for Lithium Batteries Solid State Ionics 1996, 86–88, Part 2, 819-824. (8) Novák, P.; Müller, K.; Santhanam, K. S. V.; Haas, O. Electrochemically Active Polymers for Rechargeable Batteries Chem. Rev. 1997, 97, 207-282. (9) Shinozaki, K.; Kabumoto, A.; Sato, H.; Watanabe, K.; Umemura, H.; Tanemura, S. Mechanism of Self-Discharge in Conductive Polymer Electrodes Synth. Met. 1990, 38, 135141. (10) Passiniemi, P.; Österholm, J. E. Critical Aspects of Organic Polymer Batteries Synth. Met. 1987, 18, 637-644. (11) Novák, P.; Rasch, B.; Vielstich, W. Overoxidation of Polypyrrole in Propylene Carbonate: An in Situ Ftir Study J. Electrochem. Soc. 1991, 138, 3300-3304. (12) Novák, P.; Inganäs, O. Self‐Discharge Rate of the Polypyrrole‐Polyethylene Oxide Composite Electrode J. Electrochem. Soc. 1988, 135, 2485-2490. (13) Karlsson, C.; Huang, H.; Strømme, M.; Gogoll, A.; Sjödin, M. Polymer–Pendant Interactions in Poly(Pyrrol-3-Ylhydroquinone): A Solution for the Use of Conducting Polymers at Stable Conditions J. Phys. Chem. C 2013, 117, 23558-23567. (14) Milczarek, G.; Inganäs, O. Renewable Cathode Materials from Biopolymer/Conjugated Polymer Interpenetrating Networks Science 2012, 335, 1468-1471. (15) Rosciano, F.; Salamone, M. M.; Ruffo, R.; Sassi, M.; Beverina, L. Crosslinked Electroactive Polymers Containing Naphthalene-Bisimide Redox Centers for Energy Storage J. Electrochem. Soc. 2013, 160, A1094-A1098. (16) Xu, L.; Yang, F.; Su, C.; Ji, L.; Zhang, C. Synthesis and Properties of Novel Tempo-Contained Polypyrrole Derivatives as the Cathode Material of Organic Radical Battery Electrochim. Acta 2014, 130, 148-155. (17) Rose, T. L.; Kon, A. B.; Wang, F. High Charge Storage Conductive Polymer Abstr. Pap. Am. Chem. Soc. 1995, 209, 311-312. (18) Karlsson, C.; Huang, H.; Strømme, M.; Gogoll, A.; Sjödin, M. Probing Polymer– Pendant Interactions in the Conducting Redox Polymer Poly(Pyrrol-3-Ylhydroquinone) J. Phys. Chem. C 2014, 118, 23499-23508. (19) Olsson, H.; Carlsson, D.; Nyström, G.; Sjödin, M.; Nyholm, L.; Strømme, M. Influence of the Cellulose Substrate on the Electrochemical Properties of Paper-Based Polypyrrole Electrode Materials J. Mater. Sci. 2012, 47, 5317-5325. (20) Razaq, A.; Nyholm, L.; Sjödin, M.; Strømme, M.; Mihranyan, A. Paper-Based Energy-Storage Devices Comprising Carbon Fiber-Reinforced Polypyrrole-Cladophora Nanocellulose Composite Electrodes Adv. Energy Mater. 2012, 2, 445-454.


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(21) Nyström, G.; Razaq, A.; Strømme, M.; Nyholm, L.; Mihranyan, A. Ultrafast AllPolymer Paper-Based Batteries Nano Lett. 2009, 9, 3635-3639. (22) Nyström, G.; Mihranyan, A.; Razaq, A.; Lindström, T.; Nyholm, L.; Strømme, M. A Nanocellulose Polypyrrole Composite Based on Microfibrillated Cellulose from Wood J. Phys. Chem. B 2010, 114, 4178-4182. (23) Olsson, H.; Nyström, G.; Strømme, M.; Sjödin, M.; Nyholm, L. Cycling Stability and Self-Protective Properties of a Paper-Based Polypyrrole Energy Storage Device Electrochem. Commun. 2011, 13, 869-871. (24) Gelin, K.; Mihranyan, A.; Razaq, A.; Nyholm, L.; Strømme, M. Potential Controlled Anion Absorption in a Novel High Surface Area Composite of Cladophora Cellulose and Polypyrrole Electrochim. Acta 2009, 54, 3394-3401. (25) Ferraz, N.; Carlsson, D. O.; Hong, J.; Larsson, R.; Fellström, B.; Nyholm, L.; Strømme, M.; Mihranyan, A. Haemocompatibility and Ion Exchange Capability of Nanocellulose Polypyrrole Membranes Intended for Blood Purification J. R. Soc. Interface 2012, 9, 1943-1955. (26) Olsson, H.; Jämstorp Berg, E.; Strømme, M.; Sjödin, M. Self-Discharge in Positively Charged Polypyrrole–Cellulose Composite Electrodes Electrochem. Commun. 2015, 50, 43-46. (27) Olsson, H.; Sjödin, M.; Berg, E. J.; Strømme, M.; Nyholm, L. Self-Discharge Reactions in Energy Storage Devices Based on Polypyrrole-Cellulose Composite Electrodes Green 2014, 4, 27-39. (28) Christensen, P. A.; Hamnett, A. In Situ Spectroscopic Investigations of the Growth, Electrochemical Cycling and Overoxidation of Polypyrrole in Aqueous Solution Electrochim. Acta 1991, 36, 1263-1286. (29) Lewis, T. W.; Wallace, G. G.; Kim, C. Y.; Kim, D. Y. Studies of the Overoxidation of Polypyrrole Synth. Met. 1997, 84, 403-404. (30) Pud, A. A. Stability and Degradation of Conducting Polymers in Electrochemical Systems Synth. Met. 1994, 66, 1-18. (31) Park, D. S.; Shim, Y. B.; Park, S. M. Degradation of Electrochemically Prepared Polypyrrole in Aqueous Sulfuric Acid J. Electrochem. Soc. 1993, 140, 609-614. (32) Beck, F.; Braun, P.; Oberst, M. Organic Electrochemistry in the Solid StateOveroxidation of Polypyrrole Ber. Bunsen-Ges. Phys. Chem 1987, 91, 967-974. (33) Mihranyan, A.; Llagostera, A. P.; Karmhag, R.; Strømme, M.; Ek, R. Moisture Sorption by Cellulose Powders of Varying Crystallinity Int. J. Pharm. 2004, 269, 433-442. (34) Nyström, G.; Strømme, M.; Sjödin, M.; Nyholm, L. Rapid Potential Step Charging of Paper-Based Polypyrrole Energy Storage Devices Electrochim. Acta 2012, 70, 9197. (35) Hutchison, G. R.; Ratner, M. A.; Marks, T. J. Hopping Transport in Conductive Heterocyclic Oligomers: Reorganization Energies and Substituent Effects J. Am. Chem. Soc. 2005, 127, 2339-2350. (36) Brédas, J.-L.; Scott, J. C.; Yakushi, K.; Street, G. B. Polarons and Bipolarons in Polypyrrole: Evolution of the Band Structure and Optical Spectrum Upon Doping Phys. Rev. B 1984, 30, 1023-1025. (37) Zotti, G. Doping-Level Dependence of Conductivity in Polypyrroles and Polythiophenes Synth. Met. 1998, 97, 267-272.


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Figure 1. The proposed reaction scheme for PPy self-discharge. An initial, and polymer intrinsic redox reaction leads to an unstable intermediate (Int+) that reacts with its surroundings to form maleimide and other degradation products while reducing the polymer. The reaction rate for producing the intermediate state increases with polymer potential by decreasing the potential barrier for the initial reaction. 94x49mm (150 x 150 DPI)


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Figure 2. Self-discharge curves for PPy/cellulose composite samples in 1M NaClaq solution at different electrolyte temperatures. A faster discharge is seen at higher temperatures indicating a thermally activated reaction. 84x64mm (300 x 300 DPI)



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Figure 3. Cyclic voltammogram of the PPy/cellulose composite recorded at 5 mV/s in 1M NaClaq solution at 20 °C. The CV was integrated to obtain the capacity at different potentials. The colored regions indicate the time-scales (1/kapp at 20 °C) for the faradaic and polymer intrinsic mechanism found to be the dominant cause of self-discharge. 84x58mm (300 x 300 DPI)


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Figure 4. Apparent rate constants (open circles) evaluated at different potentials from self-discharge experiments for PPy/cellulose composite samples in 1M NaClaq solution at 20 °C together with a fit to equation 3 (solid line) with A = 8*10-11 s-1 and b = 25 V-1 as evaluated from the fit. 58x41mm (600 x 600 DPI)



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Figure 5. Activation energies (solid squares) evaluated from the temperature dependence of apparent rate constants as well as evaluated redox potentials for the electron donor (intermediate) (open circles) as function of potential. Solid lines serve as guides for the eye. 58x41mm (600 x 600 DPI)


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Figure 6. Cyclic voltammetry of PPy covered carbon microelectrodes in 1M NaClaq solution at 1, 10 and 100 Vs-1 in black, blue and red respectively. The currents are normalized with respect to scan rate. The vertical line corresponds to the formal potential of the intermediate state during self-discharge as evaluated from the activation energy of self-discharge. 84x64mm (300 x 300 DPI)



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Figure 7. Overoxidation peak potentials as function of scan rate (solid squares). The peak potentials were obtained by cyclic voltammetry on PPy covered carbon microelectrodes in 1M NaClaq solution. The solid line shows a fit to equation 7 using E0 = 0.89 V vs. Ag/AgCl and an exchange rate constant of 9 s-1 and α = 0.33 as evaluated from the fit. 58x41mm (600 x 600 DPI)


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Figure 8. Required driving force as function of reorganization energy to account for activation energies of 0.35 eV (black squares), 0.45 eV (grey circles) and 0.6 eV (red triangles) as obtained from the temperature dependence of the self-discharge rate. It is seen that the initial rate-limiting step is energonic for all reasonable reorganization energies. The inset shows an estimate of the error introduced by using activation energies as the driving force for the involved reaction. 84x58mm (300 x 300 DPI)



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Activation Barriers Provide Insight into the Mechanism of Self

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