Computational Electrochemistry Study of 16 Isoindole-4, 7-diones as

Article pubs.acs.org/JPCC

Computational Electrochemistry Study of 16 Isoindole-4,7-diones as Candidates for Organic Cathode Materials Christoffer Karlsson,* Erik Jam ̈ storp, Maria Strømme,* and Martin Sjödin* Division for Nanotechnology and Functional Materials, Department of Engineering Sciences, The Ångström Laboratory, Uppsala University, Box 534, SE-751 21 Uppsala, Sweden S Supporting Information *

ABSTRACT: Prediction of the redox behavior of electroactive molecules enables screening of a variety of compounds and can serve as a guideline in the search for organic molecules for use as cathode materials in, for example, Li ion batteries. In this study, we present a computational strategy, based on density functional theory, to calculate redox potentials and acid dissociation constants for a series of 16 isoindole4,7-dione (IID) derivatives. The calculations take all possible electron and proton transfers into account, and the results were found to correlate very well with electrochemical and spectroscopic measurements. The possibility of polymerizing the IID derivatives was also assessed computationally, as polymerization serves as a straightforward route to immobilize the active material. Three of the considered IIDs (5,6-dicyano-2-methyl-isoindole-4,7-dione, 5,6-dihydroxy-2-methyl-isoindole-4,7-dione, and 2-methyl-5-(trifluoromethyl)-isoindole-4,7-dione) are predicted to be particularly interesting for making polymers for organic cathodes because these are calculated to have high redox potentials and high specific capacities and to be readily polymerizable. The presented strategy is general and can be applied in the prediction of the electrochemical behavior of quinones as well as other systems involving proton and electron transfers.

Figure 1. Structure of isoindole-4,7-dione (1) with numbering of the ring atoms.

also been considered to be promising radiosensitizers.10 The bioreductive properties of the IIDs are highly influenced by pH,4 which renders an investigation thereof very important. For the last three decades, pyrrole and quinone compounds have also been investigated for use as electrode material in lithium ion batteries (LIBs).2,7,11−13 Commercially available LIBs currently use inorganic Li intercalation compounds, such as LiCoO2, as cathodes. The supply of such minerals is limited, and the production of inorganic LIBs is, thus, expected to become even more energy consuming than today, resulting in higher final product costs.14 Research has therefore started to aim at developing redox active organic molecules that can replace the inorganic materials used in LIBs today. Apart from being environmentally benign and inexpensive, organic-based electrode materials could have other advantages, such as higher specific capacities, reduced problems with swelling upon cycling,15 (since they are amorphous) and the possibility to produce bendable materials.16 Several classes of organic compounds have been proposed for this purpose in recent years, for example, nitroxides,17,18 disulfides,19,20 carboxylates,21,22 and carbonyl compounds.7,11−13,23,24 We propose 1, and derivatives thereof, as new attractive candidate materials for organic LIB cathodes. The corresponding dilithium salt has a theoretical capacity of 333 mAh/g. The molecule is, however, soluble in organic solvents, which limits

a bifunctional molecule and is both a pyrrole and a quinone derivative. Members of the IID group have previously been shown to exhibit pronounced antibacterial activity9 and have

Received: December 10, 2011 Revised: December 26, 2011 Published: December 30, 2011

1. INTRODUCTION Pyrrole, p-benzoquinone, and their derivatives are very important electroactive organic compounds frequently encountered in nature and in a number of industrial applications.1−7 In nature, the heterocyclic electron rich pyrrole moiety is found in, for example, chlorophyll, heme, tryptophan, and vitamin B121 and has been investigated due to its ability to form conducting polymers with uses in energy storage devices2 and sensors.3 Quinones are also common in nature and serve vital roles in many aspects of the biochemistry of the cell, such as in the electron transport chain.4 For this reason, they have been extensively used in drug molecules such as antibacterial and antitumor agents.5 The activity of both natural and synthetic quinones depends on their unique electrochemistry and their selective bioreductive properties.4,8 Ring fusion of pyrrole and p-benzoquinone results in isoindole-4,7-dione (IID, 1, Figure 1), which consequently is

© 2011 American Chemical Society

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dx.doi.org/10.1021/jp211851f | J. Phys. Chem. C 2012, 116, 3793−3801

The Journal of Physical Chemistry C

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2.1. Electrochemistry. Electrochemical experiments were performed in aqueous or organic electrolyte using an Autolab PGSTAT302N potentiostat (Ecochemie, The Netherlands). The organic electrolyte used was 0.1 M t-butylammonium hexafluorophosphate (TBAHFP) in acetonitrile (MeCN), and the supporting electrolyte in water was 1 M NaNO3. The aqueous electrolyte was buffered with phosphate, borate, and acetate ions (10 mM each), and the pH was tuned with HNO3 (aq). All potentials measured in aqueous solutions are reported versus the standard hydrogen electrode (SHE), and all potentials measured in MeCN are reported versus the ferrocene couple (Fc0/Fc+), unless otherwise specified. A polished and sonicated glassy carbon disk electrode was used as working electrode, and a Pt wire was used as counter electrode. In aqueous solutions, a Ag/AgCl (3 M NaCl) reference electrode was used (0.197 V vs SHE), and in MeCN solutions, a Ag0/Ag+ (0.01 M AgNO3, 0.1 M TBAHFP) reference electrode was used (−0.096 V vs Fc0/Fc+). Analyte solutions in electrolyte were thoroughly degassed with N2(g) and kept under N2(g) atmosphere throughout the measurements. UV/vis spectroelectrochemical measurements were made using a quartz cuvette as a thin layer cell and a Au grid as the working electrode. Peak currents were measured versus a baseline that was either linearly extrapolated from the previous data points or measured using linear sweep voltammetry (LSV) by sweeping the potential to a value before the peak and then keeping it constant. 2.2. Synthesis of 2-Methyl-isoindole-4,7-dione (2). 2 was synthesized according to a procedure adapted from a work by Schubert-Zsilavecz et al.32 p-Benzoquinone (2.0 mmol, 216 mg), sarcosine (4.0 mmol, 356 mg), and paraformaldehyde (10 mmol, 300 mg) were dissolved in toluene (20 mL). The reaction mixture was refluxed for 1 h using a Dean−Stark apparatus. The solvent was evaporated, and the crude residue was purified with flash chromatography (EtOAc/pentane gradient, 1:4 to 1:1). The product was obtained as yellow crystals; 48 mg (15% isolated yield). 1H NMR (CDCl3, 400 MHz) δH = 3.77 (s, 3H, CH3), 6.64 (s, 2H, H5, H6), 7.20 (s, 2H, H1, H3). 13C NMR (CDCl3, 100 MHz) δC = 37.2 (CH3), 121.7 (C3a, C7a), 124.6 (C1, C3), 139.7 (C5, C6), 182.0 (CO). NMR spectra are available in the Supporting Information. The synthesis is summarized in Scheme 1.

its use as cathode material in itself. A common strategy to solve this problem is to increase the molecular weight by polymerization.11,18,25 Isoindoles can be polymerized on the 1 and 3 positions,26,27 resulting in a conducting polymer with a polypyrrole backbone. Polymerizing 1 would produce polymers similar to that made by Häringer et al.,25 in which a polyaniline derivative fused with p-benzoquinone moieties was investigated. Polymerization not only serves to solve solubility problems but also, depending on the relative energetics of the quinone moiety and the resulting pyrrole polymer backbone, provides a conductive polymer network for charge transport that would decrease or eliminate the need for a conductive additive such as carbon black. Such a polymer would ideally have the welldefined redox potential of the quinone side groups as well as the capacitive charging of the conducting polymer. In the development of redox-active molecules, computations provide useful tools for predicting key properties, enabling screening of a multitude of compounds. One of the most interesting parameters for electrode materials is the redox potential, which together with the capacity determines the specific energy of the material. Other properties that are of interest for electrode materials as well as for antimicrobial agents and radiosensitizers include stability in different solvents, pKa values, and the propensity toward polymerization. The redox properties of simple quinone systems such as pbenzoquinone have been thoroughly studied both experimentally8,28 and computationally.29,30 Several computational models such as Hartree−Fock and density functional theory (DFT) have previously been used for this purpose. Quinones exhibit complex electrochemistry due to the existence of a radical semiquinone state and the possibility of proton transfers coupled to all redox reactions,8 which makes prediction of quinone electrochemistry difficult. For example, many studies do not consider all forms of oxidation and protonation that are available for the species in solution. Including all possible forms is often necessary to elucidate the complete picture of the reactions that can take place upon redox cycling. Experimental conditions can also significantly affect the observed electrochemical behavior. For example, depending on the choice of electrode material, adsorption onto the working electrode surface can occur,8 and the buffer capacity of a water solution has been found to change fundamentally the quinone electrochemistry.28 In this work, we present a computational study of the electrochemical reactions of a series of IIDs, performed according to a strategy that includes all possible proton and electron transfers. The results of the calculations were compared with electrochemical and spectroscopic measurements on 2-methyl-isoindole-4,7-dione (2) and with experimental data from a study conducted by Schubert-Zsilavecz et al.31 to verify the computational strategy.

Scheme 1. Synthesis of 2-Methyl-isoindole-4,7-dione (2)

2.3. DFT Calculations. Gaussian 0933 was used to perform the DFT calculations. Molecular structures were optimized in the gas phase at the B3LYP level of theory using the 6311+G(d) basis set. This model is expected to give a good description of the electronic structure at reasonable computational times.34 Frequency calculations were performed on the optimized structures to obtain thermal correction terms. The PCM solvation model with the UAKS topological model was used to calculate the energies of molecules in solution (either water or MeCN), with geometries optimized in the gas phase. If atomic charges were desired, then natural bond orbital (NBO) analysis was also performed. The Gibbs free energy at

2. EXPERIMENTAL AND THEORETICAL METHODS Reagents were purchased from Sigma Aldrich and were used without further purification. Flash chromatography was performed using VWR Normasil 60 silica gel (40−63 μm, 60 Å). NMR spectra were recorded on a Varian INOVA (1H at 399.97 MHz, 13C at 100.58 MHz) spectrometer. Chemical shifts are reported using the residual chloroform solvent signal as an indirect reference to TMS (δH = 7.26, δC = 77.00). 13C NMR signals were assigned with HSQC and HMBC. UV/vis spectra were recorded on a UV-1650PC Shimadzu spectrophotometer using a quartz cuvette. 3794



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Figure 2. Electrochemical characterization of 2-methyl-isoindole-4,7-dione (2): DPV (50 mV/s) (a,b), SWV (50 mV/s) (c,d), and CV (100 mV/s) (e,f) in MeCN (a,c,e) and in water buffered at pH = 7.0 (b,d,f). Arrows denote the sweep direction. The displayed curves are all from the first cycle.

Here EH+ is the energy of bringing a proton into solution at a certain pH, which is given by basic thermodynamics and can be expressed as

298.15 K of each molecule was calculated as the sum of the energy in solution and the thermal correction of Gibbs free energy34 according to 0 0 + ΔG 0 ΔGtot = ΔGsol therm

E H+ = ΔG 0+′

(1)

H (aq)

ΔGk0 =

pi =

∑j gjNj

=

(7)

(8)

can be calculated from the energies obtained according to 0 = ΔG 0 − ΔG 0 ΔGred k−n k

(9)

and E0 = −

0 ΔGred

nF

(10)

Here n is the number of electrons transferred and F is the Faraday constant. The potential E0 can be referenced toward SHE, which has an absolute potential of 4.44 V at 298.15 K.38 The potentials of the one- and two-electron transfer reactions obtained in this fashion are pH-dependent (eq 6). To reference potentials in MeCN toward Fc0/Fc+, the absolute potential of that redox reaction was calculated with the same method (using the LANL2DZ basis set).

(3)

where Ei is the energy of that form and j is the range of all protonation forms available in that oxidation state. Furthermore, Ni is the number of molecules in state i, gi is the degeneracy of state i, kB is the Boltzmann constant, and T is the temperature. It is necessary to include the Gibbs free energy of the protons in Ei. Hence, for protonation level 0 (left-hand side of eq 2), the energy is

3. RESULTS 3.1. Electrochemical and Spectroscopic Characterization of 2. Differential pulse voltammetry (DPV), square wave voltammetry (SWV), and cyclic voltammetry (CV) was performed on 2 in organic and aqueous electrolytes (Figure 2) for comparison with computational data. The CV of 2 in MeCN at 100 mV/s (Figure 2e) shows two quasireversible pairs of redox peaks at −1.25 (peak 1) and

(4)

and for protonation level 1 (right-hand side of eq 2) E1 = EAH+

∑ Ej ·pj

Ak + ne− → Akn‐−n

gie−Ei / kBT

E0 = EA + E H+

(6)

The absolute standard potential of the reaction where A is reduced from oxidation state k to k-n,

(2)

∑j gje−Ej / kBT

+ pH ·59 meV

j

it is possible to calculate the population of one form, pi, using the Boltzmann distribution: giNi

H (aq)

where the solvation energy in the standard state is ΔGH0′+(aq) = 1107.8 kJ/mol.37 The total Gibbs free energy can then be calculated for each oxidation state, k, as

UV/vis spectra were calculated by performing a timedependent (TD-SCF) energy calculation for the six lowest energy levels on the structures optimized in gas phase using the same settings as described above. Spin density plots were made using either gOpenMol35 or GaussView,36 color mapping an isosurface of the total electron density with the α minus β spin density. A unique Gibbs free energy was calculated for every oxidation state by pooling all available protonation states, as described below. The term “protonation level” will be used to refer to the extra number of hydrogen atoms in the molecule with respect to the fully oxidized form Q. For a species A in a certain oxidation state, going from protonation level zero to one, A + H+ → AH+

= ΔG 0+

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−1.60 V (peak 2) versus Fc0/Fc+, corresponding to the conversion between the quinone (Q) and the hydroquinone (HQ) species via the semiquinone (SQ). The two peaks contain an equal amount of charge (Qred,1/Qred,2 = 1.0, Qox,1/ Qox,2 = 1.0) and the reduction and oxidation peaks in each pair have approximately the same charge (Qox,1/Qred,1 = 1.1, Qox,2/ Qred,2 = 1.0). The peak splits (voltage difference between oxidation and reduction peaks) are larger than expected for a reversible system (ΔEp,1 = 99 mV, ΔEp,2 = 107 mV). Spectroelectrochemical measurements were performed (recording UV/vis spectra during potentiostatic oxidation and reduction, Figure S1 of the Supporting Information) and also indicated that neither of the two redox reactions is reversible. In buffered water solution, there is only one pair of redox peaks (Figure 2b,d,f), which at pH = 7.0 is located at −0.20 V versus SHE. The peaks correspond to a two-electron conversion between the quinone and the hydroquinone species.8 The reaction is nonreversible in water because the oxidation peak is much smaller than the reduction peak and there is a large peak split (Qox/Qred = 0.72 and ΔEp = 121 mV at 50 mV/s). Even at high scan rates of 100 V/s the reoxidation peak was significantly smaller than the reduction peak, indicating fast following reactions. At all scan rates employed (0.01 to 100 V/s) the logarithm of the peak currents, in both organic and aqueous solutions, increased linearly with the logarithm of scan rate with a slope of 0.5, indicating diffusion-limited reactions for all peaks observed (Figure S2 of the Supporting Information). 3.2. UV/vis Spectroscopy of 2. In a neutral aqueous solution, 2 has three absorption peaks in the UV/vis range located at 220, 247, and 390 nm, as shown in Figure 3. These

equation40 and fitting the absorbance at 390 nm gives a pKa value of −0.2 for this protonation (inset graph in Figure 3). 3.3. DFT Calculations. Gibbs free energies were calculated for the series of 16 IID derivatives (1−16, Table 1) and are displayed in Table S1 of the Supporting Information. The obtained energies were used to calculate the redox potentials of the one- and two-electron processes (Table 1). Calculated and experimental redox potentials for 2 are shown versus pH in the Pourbaix diagram in the lower part of Figure 4. Calculated Pourbaix diagrams for the entire IID series 1−16 are available in the Supporting Information. pKa values of the HQ and SQ species of all compounds were calculated as the pH where pA = pAH+, (eqs 2 and 3, Table 1). Corresponding calculations were also performed on compounds with known pKa values, for verification. The potentials needed for oxidation of Q to Q•+ were also calculated in both water and MeCN (Table 1). Calculated potentials in MeCN were referenced toward Fc0/ Fc+, which was calculated to have an absolute redox potential of 5.108 V. The spin densities of the oxidized radical form Q•+ for all compounds in the series were calculated. The calculated spin density distribution for 2 is shown in Figure 5, and maps of all compounds are available in the Supporting Information. The major part of the spin density is located on positions 1 and 3 for all molecules in the series except for 8, 11, 13, and 16.

4. DISCUSSION 4.1. DFT Calculations. Quinones can be reduced and protonated in multiple steps and in a number of combinations (Scheme 2). To calculate redox potentials, all different levels of oxidation and protonation that are available to the species of interest were taken into account. There are three possible oxidation states, namely, the fully oxidized (quinone) form, Q, the fully reduced (hydroquinone) form, HQ, and an intermediate radical (semiquinone) form, SQ. For each oxidation state, a number of protonation forms can exist. SQ, for example, can be protonated on none (Q•‑), one (QH•), or both (QH2•+) of the oxygens, depending on the pH (middle column of Scheme 2). Furthermore, some protonation forms of symmetrical IID derivatives (i.e., where R5 = R6, cf. Table 1) have degenerate energy levels; for example, QH• can have the OH group on either position 4 or 7, which are equivalent and have the exact same energy. For asymmetrical IID derivatives, the two positions are nonequivalent and must be treated separately in the calculations. More protonation levels can exist in theory, for example, both Q and QH2 could, in principle, be protonated. However, calculations show that the acid dissociation constants for those levels are very low and will hence not affect the calculation of redox potentials in aqueous solutions. It should be noticed that 6 and 16 have two quinone moieties and have a total number of five oxidation states each. It should also be noticed that compound 11 is substituted with a hydroxyl group, which in the Q form can be deprotonated at high pH values (calculated pKa = 10.2), forming an enolate.41 Therefore at pH > 10.2 it exhibits the quite rare feature of a 3H+, 2e− coupled redox reaction. 4.2. Electrochemical Behavior of 2. Calculated redox potentials for compound 2 are shown versus pH in the Pourbaix diagram in the bottom part of Figure 4. The experimentally observed data points that are overlaid in the Figure are DPV reduction peak potentials. The reason for choosing this parameter, rather than, for example, E0′, is that

Figure 3. UV/vis absorption spectra of 2-methyl-isoindole-4,7-dione (2) in buffered aqueous solution at different pH values: −0.13, 0.00, 0.12, 0.34, 0.48, 0.69, 1.15, and 1.64. Arrows indicate changes in the spectrum at decreasing pH. The inset shows the absorbance at 390 nm versus pH and a fit to the Hendersson−Hasselbalch equation.

transitions correspond to the π−π*, π−π*, and n-π* transitions of the quinone moiety, respectively.39 Upon lowering the pH below ∼2, an absorption peak at ∼300 nm appears (Figure 3) in conjunction with a decreased peak intensity for the 247 and 390 nm transitions. In addition, a weak band at 500 nm appears. The absorption spectrum that emerges at low pH values belongs to the protonated form of 2, QH+. Employing the Beer−Lambert law on the Hendersson−Hasselbalch 3796



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Table 1. Calculated Redox Potentials and Acid Dissociation Constants for the Series of IID Derivatives Studied as Well As Comparison with Experimental Values and Calculations for p-Benzoquinone, Pyrrole, and N-Methyl-pyrrole

a

Substituents on positions 2, 5, and 6 (Figure 1). bCalculated redox potentials (V vs SHE) for the one- (SQ/HQ and Q/SQ) and two-electron (Q/ HQ) processes in H2O at pH = 0 (E0) and 7 (E0′). Bold figures for E0 and E0′ indicate whether the one- or two-electron transfer, respectively, is favored. Sorted E0′ values are plotted in Figure 6. cExperimental redox potentials (V vs SHE) at pH = 7 reported by Schubert−Zsilavecz et al.31 The E0′ value for 2 determined by DPV (50 mV/s) in this work is −0.205 V. dComputational model utilized in this work accounts well for radical pKa values, whereas the model overestimates the QH2 pKa value (Table S2 of the Supporting Information). eCalculated Q/Q•+ oxidation potentials in H2O (V vs SHE) and MeCN (V vs Fc0/Fc+). f6 and 16 have two sets of redox potentials and pKa values because they have an additional quinone group. g11 has an additional calculated pKa at 10.2 for the deprotonation of Q. hAt pH = 7. (The potential is pH-dependent due to deprotonation of Q•+; see page S17 of the Supporting Information.) iReported by Shim et al.28

Figure 4. Pourbaix diagram with calculated and experimental redox potentials of 2-methyl-isoindole-4,7-dione (2). Bottom graph shows calculated redox potentials of the SQ/HQ, Q/SQ, and Q/HQ reactions versus pH. Thick lines show the thermodynamically favored transitions. Experimental potentials measured with DPV (50 mV/s, reduction peak) are overlaid as black filled circles. The top graph shows calculated populations of the different protonation forms of the SQ state at corresponding pH values.

Figure 5. Calculated total electron density isosurface for 2-methylisoindole-4,7-dione (2) in the oxidized Q•+ form, color mapped by spin density, from low (blue) to high (red) spin density.

A linear decrease of ∼59 mV/pH is indeed also observed experimentally in the pH range 3−8 (black filled circles in Figure 4). The one-electron transfer reactions, however, show a more complex behavior because different SQ species are dominating at different pH values. The top part of Figure 4 shows the calculated Boltzmann populations of the three protonation forms in the radical oxidation state, SQ. These calculations predict two pKa values at 5.7 and 11.7, respectively. Between these values, QH• is the dominating SQ species, and both one-

the redox reaction is nonreversible in water and the oxidation peak does, therefore, not give reliable data. The calculated redox potential of the Q/HQ transition versus pH is a straight line with a slope of −59 mV/pH (black solid line in Figure 4), corresponding to an equal number of protons and electrons being transferred: Q + 2H+ + 2e− ⇄ Q H2

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The fact that the predicted change in the slope of E versus pH is also observed in experiments verifies the calculated pKa value of 5.7 for the QH•/QH2•+ couple. It is noteworthy that this acid dissociation constant could be determined using only DFT calculations and voltammetric experiments, even though the semiquinone radical is not stable under these experimental conditions. Methods such as pulse radiolysis are usually required to measure such pKa values.8 Without the computational prediction of this behavior, the observed change in slope and the deviation from the Q/HQ line might have passed unnoticed. At very high pH values, calculations indicate that Q•− will be the dominant radical species and that the one-electron processes will become favored at pH >14.6. 4.3. Redox Potentials of the IID Series. Calculated redox potentials for the studied IID derivatives are shown in Table 1 and also sorted by potential in Figure 6. In general, the redox

Scheme 2. Oxidation and Protonation States of Isoindole4,7-dione (1)a

a

Each column corresponds to an oxidation state; from left to right: HQ, SQ, and Q. In each oxidation state, several protonation levels are possible.

electron transfers will be coupled to a one-proton transfer according to Q + H+ + e− ⇄ QH•

(12)

and QH• + H+ + e− ⇄ QH2

(13)

Figure 6. Calculated redox potentials at pH = 7 for the studied series of IIDs. Squares mark two-electron transfers and circles mark oneelectron transfers. Crosses show experimental values.31 Colors indicate the different groups of substituents (R5): B = fused benzo group; R = alkyl or H; X = halogen; Q = extra quinone group. The compounds in the latter group have additional two-electron transfers at low potentials.

These redox transitions are, however, not experimentally accessible individually because E0′ for the reaction in eq 12 is below that of the reaction in eq 13 in that pH region; therefore, the radical QH• is unstable with respect to disproportionation. Below pH = 5.7, QH2•+ is the dominating SQ species, and the first one-electron reduction will be coupled to a two-proton transfer, whereas the second reduction does not involve any proton transfer, according to Q + 2H+ + e− ⇄ QH•+ 2

potential increases with the electron-withdrawing effect of the substituents. The Hammett parameter σp can be used to describe the electron withdrawing effect of the substituents on positions 5 and 642 and calculated redox potentials correlate well with the sum of σp for R5 and R6 (r2 = 0.91, Figure 7). In all solution-phase DFT calculations performed in this work, the PCM solvation model has been used. PCM is an implicit solvation model; that is, it does not contain explicit solvent molecules but instead represents interaction with the solvent by placing the molecule in a cavity of the solution with a certain dielectric constant. Implicit solvation models account well for long-range effects of solvation but are less useful for predicting specific interactions with solvent molecules. Contributions from such interactions, most notably hydrogen bonding, can affect the outcome of the calculations, especially in water solution. For instance, the model tends to overestimate the energy for the deprotonated hydroquinone (QH−), where significant hydrogen bonding is expected to stabilize the negative charge on the oxygen, thereby overestimating the dissociation constant for the corresponding deprotonation of QH2 (Table S2 of the Supporting Information). Explicit solvation models, on the other hand, suffer from other problems, that is, difficulties in finding global conformational

(14)

and − QH•+ 2 + e ⇄ QH2

(15)

Hence, the slope of E versus pH changes to −118 mV/pH for the first reduction (red dotted line in Figure 4), whereas the second reduction becomes pH-independent (golden dashed line in Figure 4). The two E versus pH lines cross at pH = 2.8, and below that value the one-electron processes become favored instead of the two-electron process. It should be pointed out that pH values where the lines in the Pourbaix diagram cross (2.8 and 14.6) do not correspond directly to pKa values of the involved species. Below pH 2.8, the observed reduction peak potential does indeed deviate from the redox potential calculated for the two-electron transfer, and a stronger pH dependence is observed, as expected for the Q/SQ transition (compare black filled circles with red dotted line and black solid line in Figure 4). No additional redox peaks are present; that is, the Q/HQ transition is never observed, suggesting that the radical cation QH2•+ is unstable in water. 3798



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species should have high spin density on positions 1 and 3, and (3) there should be sufficient accessibility at those positions to allow dimerization, that is, low steric hindrance. Calculated oxidation potentials of Q (Q/Q•+) for the IID series are given in Table 1, both in water and in MeCN. Whereas the oxidation potentials in water are too high to be achieved, all oxidation potentials calculated in MeCN are within the accessible potential window (ranging from 0.8 to 2.0 V vs Fc0/Fc+). All compounds 1−16 thus meet the first criterion for polymerization. The spin density should be high on positions 1 and 3 to meet the second criterion for polymerization of the IID monomers. The calculations indicate that almost all spin density is located on those positions for all molecules in the series except for 8, 11, 13, and 16. 11 (hydroxy-substituted, which in the oxidized radical state is deprotonated at pH = 7) has the highest spin density on position C6. A radical at that position is necessary to avoid charge separation and is therefore favored. There is both an ortho- and a para-benzoquinone motif present in this state. Some spin density is also located on the oxygens attached to C5 and C7, which are positioned β to each other and which are in resonance through the radical on C6 (in α position). A similar pattern is found in 13 and 16, which have carbon−oxygen groups in β positions as well. The spin density in these two compounds is also located on the oxygens and on the α carbon. Interestingly, this is not the case for 12, which is 13 demethylated on C6. The extra stabilization by substitution of that position is apparently needed to favor the radical on C6 in that case. In the case of 8, which is Cl-substituted, the spin density is distributed over all atoms in the molecule. It is unclear why this is the case for 8 but not for any other compound in the series, for example, 9 (F-substituted). Compounds 1−7, 9, 10, 12, 14, and 15 meet the first two criteria for polymerization stated above, and if the steric hindrance is sufficiently low, then they should polymerize upon oxidation. Pyrrole can also be polymerized via protonation.45 By analogy to oxidative polymerization, the important parameters for polymerizability are the pKa value and the position of protonation. The pKa for the protonation of 2 in the Q form was determined to −0.2 by UV/vis spectroscopy (Figure 3, showing a peak at 283 nm for QH+). As for the position of protonation, DFT calculations of 2 predict that the most stable form of QH+ will be protonated on C1 and that only that form can account for the UV/vis absorption peak observed at 283 nm (Table S3 of the Supporting Information). That is also in agreement with the protonation of pyrrole, for which carbon rather than nitrogen is protonated.46 Hence, it should be possible to polymerize 2 and other IID derivatives via protonation by a strong acid. 4.5. Use of IIDs as Organic Cathode Material. Because of the electron-donating effect of the fused pyrrole ring, the redox potentials of the quinone moiety in the IID derivatives are lower than those of other quinone derivatives (cf. pbenzoquinone in Table 1). A comparison between the Mulliken or NBO atomic charges of 1 and the corresponding atoms in pbenzoquinone and pyrrole (Table S4 of the Supporting Information) reveals the electron-donating effect of the pyrrole ring. The electron density is clearly shifted from the pyrrole to the quinone part (the sum of NBO charges on the quinone ring is −0.36), making reduction of the quinone moiety more difficult. In a polyisoindole polymer of 2, the oxidation of the main chain would therefore occur at higher potentials than the

Figure 7. Calculated redox potentials at pH = 7 for compounds 2−14 versus the sum of the Hammett parameter σp for R5 and R6.

minima and underestimation of shielding effects of polar solvents. They are also considerably more time-consuming than implicit models.43 To validate the model, four of the compounds studied by Schubert-Zsilavecz et al.31 have been included in the present work (2, 10, 13, and 15). These redox potentials are CV reduction peak potentials measured in a water/ethanol or water/N,N-dimethylformamide mixture. Calculated and experimental redox potentials at pH = 7 are compared in Table 1 and in Figure 6. The redox potentials calculated here are very close to the experimental values,31 with a mean absolute deviation (MAD) of 15.4 mV. pKa values of phenols (corresponding to the HQ form, Table S2 of the Supporting Information) calculated for method validation were greatly overestimated (as discussed above, MAD = 8.2), whereas agreement for radical cations (corresponding to the SQ form) was better (MAD = 1.1). Because only the calculations of pKa values for the SQ species influence the comparison with experiments discussed in Section 4.2, the results are based on reliable pKa data. Considering this, and the results on 2, the PCM solvation model was deemed to be sufficient for this study. Calculated redox potential values are very dependent both on the level of theory used (e.g., the basis set) and on the Gibbs free energy of a proton in aqueous solution, ΔGH0 +(aq) . The latter parameter has not been precisely determined, and estimations in the literature vary widely.37 Poor choice of either of these parameters can result in a bias of calculated redox potentials. The fact that the redox potentials calculated here fit so well with the experimental data confirms that the method used in this work is a good choice for the IID family of compounds. Furthermore, the method used to treat the various forms of oxidation and protonation is applicable to all types of quinones as well as to other systems involving proton and electron transfers. 4.4. Propensity toward Polymerization. Isoindoles have previously been polymerized to polypyrrole-type polymers via chemical oxidation or electrochemical oxidation.26,27 For successful polymerization of a pyrrole derivative, there are, in principle, three requirements that have to be met: (1) the oxidation potential must fall within the potential window of the solvent used, enabling the formation of the oxidized radical cation Q•+, (2) the unpaired electron of Q•+ should be situated on carbons where dimerization occurs;44 that is, the radical 3799



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focus on IIDs for use as organic cathode materials by exploring the potential of the most promising compounds from this study and to test these as cathode materials in LIBs.

quinone oxidation. At the potential where the quinone moiety is redox-active, the polypyrrole backbone would then be in the reduced state and hence nonconductive. An advantage of this actuality is that polypyrrole would serve as protection against overoxidation.25 The low potential of the quinone group also means that the cell voltage of a LIB will be low. In the organic electrolyte used here, the Fc0/Fc+ redox potential is 3.68 V versus Li0/Li+.47 Hence, the electron transfers of 2 in MeCN would give a cell voltage of ∼2.0 V in an LIB. This value is lower than inorganic Li intercalation compounds, which typically give a cell voltage of 3 to 4 V.48 However, because of the much higher theoretical capacity of 2, its theoretical specific energy is nevertheless comparable to that of the inorganic cathode materials. The calculated redox potentials of most compounds in the studied IID series are similar to those of 2. Their performance in a LIB would hence also be comparable to inorganic intercalation materials. According to the calculations, however, a few compounds are oxidized at higher potentials than 2, namely 4, 6, 7, and 16 (8 and 9 also have marginally higher redox potentials) due to the electron-withdrawing effect of their substituents. These molecules are more promising for use as LIB cathode material for two reasons. First, the cell potential is higher, giving a larger specific energy. Second, in the corresponding polymer, the quinone redox reaction occurs at a potential where the polypyrrole backbone is oxidized and hence conductive. 6 and 16 also have additional reductions at lower potentials, giving a protection against overreduction, alternatively a possibility for increased capacity. As discussed above, however, calculations indicate that 16 will not be polymerizable, leaving 4, 6, and 7 as the best candidates for organic LIB cathode materials.

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ASSOCIATED CONTENT

* Supporting Information S

Complete ref 33, spectroelectrochemical measurements on 2, scan rate dependence of peak currents of 2, NMR spectra of 2, calculated Gibbs free energies of all IID derivatives, calculated pKa values of phenols, calculated UV/vis transitions of QH+, calculated atomic charges on 1, and molecular structures, spin density plots, and calculated Pourbaix diagrams of all IID derivatives. This material is available free of charge via the Internet at http://pubs.acs.org.

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]; maria.stromme@ angstrom.uu.se; [email protected].

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ACKNOWLEDGMENTS Anna Lundstedt is gratefully acknowledged for valuable discussions and suggestions. We would also like to thank Tomas Edvinsson and Prof. Adolf Gogoll for useful comments and assistance. The Swedish Research Council (VR), the European Institute of Innovation and Technology, under the KIC InnoEnergy Electric Energy Storage project and the KIC InnoEnergy NewMat project, as well as the Nordic Innovation Centre (contract number 10014) are acknowledged for their financial support.

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5. CONCLUSIONS IIDs consist of a benzoquinone moiety with a fused pyrrole ring and exhibit properties characteristic of both compounds. The quinone moiety can be reduced in a two-electron reaction, allowing them to function as Li high-capacity cathode materials in organic LIB batteries. A computational strategy for prediction of the redox behavior of IIDs, in particular, and quinones, in general, is presented. This strategy takes into account all possible electron and proton transfers and was used to predict the redox behavior of a series of 16 IIDs in aqueous solution. pKa values of the HQ and SQ species were also calculated using the same method. One compound in the series (2) was synthesized and characterized electrochemically for method validation. Experimental observations (both from electrochemical measurements and UV/vis spectroscopy in this and previous work31) correlate very well with the results of the calculations (MAD = 15.4 mV). The pyrrole moiety in IIDs offers the possibility to polymerize the compounds, which is a convenient route to lower the solubility and to provide a conducting network, which would be interesting in a LIB application. The oxidation potential to form the oxidized radical species Q•+, as well as the spin density distribution of that radical, was used to assess the polymerizability of the compounds. On the basis of the calculations performed, three of the studied compounds (4, 6, and 7) are suggested to be the best candidates for making polymers for LIB cathodes. The strategy presented in this work opens up for prediction of the electrochemical behavior of a wide range of systems involving proton and electron transfers and can serve as a guide in the search for redox active molecules. Our work will continue to

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Computational Electrochemistry Study of 16 Isoindole-4, 7-diones as

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