Mechanistic Study of the Conversion of Superoxide to Oxygen and Hydrogen Peroxide in Carbon Nanoparticles.

Hydrophilic carbon clusters (HCCs) are oxidized carbon nanoparticles with a high affinity for electrons. The electron accepting strength of HCCs, employing the efficient conversion of superoxide (O2(•-)) to molecular oxygen (O2) via single-electron oxidation, was monitored using cyclic voltammetry and electron paramagnetic resonance spectroscopy. We found that HCCs possess O2 reduction reaction (ORR) capabilities through a two-electron process with the formation of H2O2. By comparing results from aprotic solvents to those obtained from ORR activity in aqueous media, we propose a mechanism for the origin of the antioxidant and superoxide dismutase mimetic properties of poly(ethylene glycolated) hydrophilic carbon clusters (PEG-HCCs).

Introduction

Superoxide (O2•–) is a radical anion of enormous importance in biology and medicine. O2•– is generated as a primary byproduct of the partial reduction of O2 during aerobic metabolism and is a key component of the more broadly termed reactive oxygen species (ROS).[1−3] Other ROS include hydrogen peroxide (H2O2), singlet oxygen, organic peroxides, hydroperoxyl radical (HO2•), and hydroxyl radical (HO•), which are the products of subsequent proton-coupled reactions.[1,4] ROS are implicated in many diseases and can cause severe damage to cellular components and metabolism. To protect against ROS, living organisms have developed defensive metalloenzymes such as superoxide dismutase (SOD), peroxidase, and catalase.[5−7] In addition to enzymes, protection against ROS-induced oxidative stress is provided by antioxidants such as vitamin A, vitamin E, ascorbic acid, uric acid, glutathione, and hydroquinones.[8,9] Due to the existence of reactive redox coupling of metal ions, SOD has been widely investigated by electrochemical techniques, which has led to the development of various SOD-modified amperometric biosensors for the measurement of O2•–.[10,11]

In addition to the biological importance of the oxygen reduction processes, the cathodic oxygen reduction reaction (ORR) has been an important counterpart and half-reaction in fuel cells.[12−14] ORR can be a two-electron process, yielding H2O2, or a four-electron process with the reduction of O2 all the way to H2O. Synthetic metalloporphyrins are common molecular ORR catalysts inspired by the heme-based cofactors.[15] However, platinum still remains the first choice as an electrocatalyst for current fuel cells despite its high cost and low abundance. Various SOD-modified catalysts have been investigated for their ORR performances.[16] The existence of O2•– as an intermediate during the ORR has been shown in aprotic solvents.[17] In our continuing study of carbon-based nanomaterials for applications in both medicine[18−23] and as metal-free catalysts to convert chemical to electrical energy,[24−26] we explore here the intersection of the two fundamental carbon-based electrochemical pathways.

We recently developed water-soluble highly oxidized carbon nanomaterials, poly(ethylene glycolated) hydrophilic carbon clusters (PEG-HCCs), that act as highly active mimetics of superoxide dismutase (SOD), a selective antioxidant that dismutases two O2•– molecules to afford O2 and H2O2 according to Scheme .[27]

Scheme 1

Mechanism of Catalytic O2•– Dismutase by PEG-HCCs[27]

Another important feature of PEG-HCCs is their inertness to nitric oxide (NO•). Although PEG-HCCs have 1 equiv of stable π-conjugated carbon radical per highly oxidized carbon core, PEG-HCCs still do not activate NO• through either radical–radical coupling or single electron transfer processes. When the difference in the redox potentials of O2•– and NO• [E°(O2•–/O2) = −0.60 V vs SCE and E°(NO•/NO+) = 1.36 V vs SCE or E°(NO–/NO•) = −1.6 V vs SCE] is known,[28−30] it is clear that oxidative single electron transfer activation of O2•– by PEG-HCCs is thermodynamically more favorable than either oxidative or reductive activation of NO•.

To investigate both the antioxidant and the catalytic properties of HCCs, we sought to analyze the properties of the carbon core of PEG-HCCs, namely HCCs, to address their mechanism for the efficient conversion of O2•– to O2. HCCs are conjugated graphenic domains with highly oxidized edges.[18,19] These are then appended to PEG moieties to render them more soluble in commonly used phosphate buffered saline (PBS) and serum and to increase their blood circulation time in vivo.[20−23] Therefore, the objective of this work was to investigate (a) the oxidizing strength of HCCs by electrochemical analysis of their redox potentials, (b) the reactivity of HCCs with O2•– in homogeneous conditions (in aprotic solvents), and (c) the ORR activity of the HCC through its immobilization on an electrode acting as a heterogeneous catalyst (in aqueous media).

Results and Discussion

Preparation and Characterization of HCCs

Preparation of HCCs was reported earlier.[18−20] Oxidation of single-walled carbon nanotubes (SWCNTs) under harsh oxidation conditions, employing oleum and fuming nitric acid, resulted in extremely small, water-soluble, generously oxidized and nontubular hydrophilic carbon clusters (Scheme ). The nanotubes unzip to give planar graphene-like fragments that are devoid of any radial breathing modes in the Raman spectra. The average size of HCCs is ∼3 nm wide and ranges from 30 to 40 nm long, and they have been extensively characterized by atomic force microscopy (AFM) and transmission electron microscopy (TEM).[18−20] Additionally, harsh oxidation conditions facilitated functionalization of HCCs with various oxygen-containing functional groups.

Scheme 2

Synthesis of HCCs from SWCNTs as the Starting Material

To estimate the extent of the harsh oxidation step and to evaluate the oxygen content and functionality of the HCCs, we characterized HCCs by X-ray photoelectron spectroscopy (XPS). In our previous work, we determined the degree of oxidation of HCCs by identifying the C/O ratio in HCCs as 2/1 (Figure S1).[20] Here, we report the details of the high-resolution XPS spectra (C 1s and O 1s) of the HCCs. Figure shows the C 1s and O 1s deconvoluted peaks of HCCs (Figure S1).

Figure 1

High-resolution deconvoluted XPS of (a) C 1s and (b) O 1s spectra of HCCs.

From the deconvoluted C 1s peaks, it is apparent that the oxygen-containing groups are predominantly in the form of hydroxyl, ether, and epoxy (C–O and C–O–C, 286.6 eV) groups. There are also carbonyl (C=O, 287.5 eV) groups in the form of ketones or quinones and carboxyl (HO–C=O, 288.9 eV) groups. Note that the spectra were centered at 284.8 eV for the sp3- and sp2-hybridized carbons (C–C/C=C). Since the HCCs are heavily oxidized, the contribution from sp3-hybridized carbons are favored. The deconvoluted O 1s peak of HCCs further shows large amounts of carbonyl and carboxyl (C=O and –O–C=O, 531.5 eV) groups, and the hydroxyl and ether (C–O and C–O–C, 532.4 eV) groups. We have also identified the peak at 530.3 eV as the peroxyl group (C–O–O). Relative contributions of each functional group are summarized in Table S1. In general, from both the C 1s and O 1s peaks, HCCs show a presence of hydroxyl, ether, carbonyl, and carboxylic groups in large amounts as compared to those in sp3- and sp2-hybridized carbon (C–C/C=C) groups, indicating the large degree of oxidation of SWCNTs that results from this procedure. It is also noteworthy to mention that the common oxidation procedures for graphene oxide (GO) preparation were not reported to give such a high degree of disrupted π-conjugated networks during the oxidation of graphene.[31]

Further evidence of the electron-deficient oxygen-containing functional groups was obtained by electrochemical analysis of the HCCs. Figure a illustrates the comparison of cyclic voltammograms (CVs) of GO and HCCs in aqueous media at pH 7.4. The electrochemistry of the electrode-immobilized carbon materials has been reported before, and particularly from the detailed studies on GO, the nature of electrochemically reducible oxygen-functional groups was shown to be dependent on the method of preparation.[32] Out of four reported GO preparation methods using the Staudenmeier,[33] Hoffman,[34] Hummers,[35] and Tour[31] procedures, only Staudenmeier’s procedure, which employs chlorate as an oxidant in the presence of fuming nitric acid (HNO3), produces GO with a positive reduction wave starting at −0.7 V and a peak potential at −1.1 V vs Ag/Ag+. All other methods, which use permanganate as the oxidant, exhibit reduction waves for GO starting at −1.0 V or below with the peak potentials ranging from ∼ −1.8 V to −2.0 V vs Ag/Ag+ (Figure a).[31,32] From the CV of HCCs (Figure a) in PBS, which was prepared using the procedure employing only oleum and fuming nitric acid, without an additional oxidant, HCCs have dominant electron-deficient domains that are far more extensive than those seen in GO. HCCs reveal reduction peaks starting at 0.0 V with the peak maxima at ∼ −0.8 V, which make HCCs stronger oxidants than GO by an average of almost 1.0 V. Oxygen-containing functional groups show different reduction potentials; specifically, only peroxide and aromatic quinone-type functional groups are reduced at potentials as high as those of HCCs (from 0.0 to −0.7 V vs Ag/Ag+).[36,37] Therefore, the strong oxidizing properties of HCCs could be due to the presence of both peroxyl and quinone moieties. The presence of peroxyl and quinone groups was also consistent with the XPS analyses shown above.

Figure 2

(a) CVs of 50 mM PBS at pH 7 on a bare glassy carbon (GC) working electrode coated with HCCs or GO recorded under a N2 atmosphere. (b) Comparison of CVs of GC/HCCs (solid line, recorded under a N2 atmosphere), the O2/O2•– couple (dot-dash line, O2-saturated buffer), and GC (dotted line, recorded under a N2 atmosphere). Scan rate: 100 mV/s.

To estimate the electron accepting strength of HCCs, we compared the reduction potentials of HCCs and O2 (Figure b) in aqueous buffer. Unlike in aprotic solvents,[27] in aqueous media, due to the disproportionation and proton-transfer processes, O2 exhibits an irreversible reduction wave potential. The reduction wave for O2 corresponds to the reversible one-electron wave followed by a second irreversible proton-coupled one-electron reduction wave according to eqs and 4:

As shown in Figure b, the O2 reduction wave potentials nicely overlap with the reduction wave potential of HCCs. Hence, the single-electron transfer oxidation of O2•– by HCCs is a thermodynamically favorable and exothermic process.

Reactivity of HCCs with Superoxide in Apolar Solvents

In addition to the electron-deficient oxygen-containing functional groups, an HCC also possesses one (on average) stable unpaired radical (Figure a) in DMSO as observed previously.[27] HCCs display an intense and symmetric electron paramagnetic resonance (EPR) signal with g = 2.0017 with an overall line width of 7 G. The sharp and structureless peak implies that π-delocalization across the extended π-network of HCCs is limited, and the radical character is more localized in nature. The stable radical character of HCCs is retained after appending the PEG groups to form water-soluble PEG-HCCs.[27] PEG-HCCs show remarkable selective antioxidant superoxide dismutase-like properties.[27] The origin of efficient catalytic quenching of O2•– was correlated to the stable radical species of the carbon core (HCCs) of PEG-HCCs (Figure S2). Consequently, HCCs showed the same behavior as PEG-HCCs toward the quenching of O2•–. The addition of KO2 to the solution of HCCs in dry DMSO resulted in an immediate loss of the EPR signal of HCCs. This implies that HCCs are capable of oxidizing O2•– with the formation of a reduced and near EPR-silent form of HCCs (Figure a). This is the first evidence that the stable intrinsic radical in HCC is capable of redox communication with O2•–.

Figure 3

(a) EPR spectra of 0.05 mg/mL of HCC in DMSO (blue) and after addition of KO2 (red). (b) CVs in DMSO containing 0.1 M [(n-Bu)4N]+ClO4– as a supporting electrolyte at 298 K at a glassy carbon working electrode with a platinum wire as a quasi-reference electrode. E(Fc/Fc+) = 0.44 vs SCE in DMSO. Scan rate: 100 mV/s. HCCs in DMSO under N2 (blue), after addition of KO2 (red), and after bubbling of the resulting solution with N2 for 5 min to remove any gaseous O2 from the solution (blue broken lines).

To further test the single-electron transfer mechanism, we performed electrochemical detection of formed O2 as a second reaction product according to eq .The electrochemistry of O2 in aprotic solvents is well-documented and shows a reversible single-electron reduction process to form stable O2•–.[38] In the presence of a proton source, the O2/O2•– couple loses the reversibility upon the proton-transfer reaction pathways.[39] This behavior of O2•– was used to test the antioxidant capacities of various phenols and polyphenols.[40] A well-dispersed solution of HCCs in anhydrous DMSO with [(n-Bu)4N]+ClO4– as a supporting electrolyte under an inert atmosphere does not show any noticeable redox peaks associated with the dispersed HCCs (Figure b). This outcome indicates that the intrinsic stable radical of an HCC does not respond to the sweeping redox potential field in this aprotic solvent in contrast to that in the aqueous buffer (Figure ). However, addition of KO2 causes the appearance of a peak at E° = −1.25 V vs Fc/Fc+, which is the reversible reduction peak of dissolved O2 that was confirmed by separate measurements in the absence of HCCs. Interestingly, this peak disappears upon bubbling of the resulting solution with N2 for 5 min (Figure b), which confirms the conversion of O2•– into molecular O2. This is unlike that of the dissolved O2•–, which is persistent in solution even after being bubbled with N2 for 6 h (Figure S3). Therefore, this measurement further demonstrates that HCCs oxidize O2•– to O2 through a one-electron transfer reaction according to eq , and the product, molecular O2, can be observed electrochemically. In addition to O2•–, we also tested the water-soluble perylenedimine radical anion (PEG-PDI•–) as a one-electron reductant of HCCs (Figure S4). A one-electron transfer reaction between HCC• and PEG-PDI•– takes place only after sonication for 20 min, which could be due to the large difference in the diffusion rates between PEG-PDI•– and O2•–.

ORR Activity of HCCs

To further explore the properties of HCCs, particularly their interactions with O2, we performed electrocatalytic ORR activity tests. HCCs were immobilized on a GC working electrode, which served as an O2-electrode, and tested by CV and linear sweep voltammetry (LSV) in an O2-saturated 0.1 M NaHPO4/NaH2PO4 buffer solution at pH 7. CVs over a short potential range are shown in Figure a. ORR activity was demonstrated by comparing the CVs under an inert atmosphere and in the presence of O2 (Figure a). An Ar-saturated electrolyte exhibited weak CV peaks, but under the O2-saturated electrolyte, a prominent ORR peak with an onset potential at ∼0.65 V vs RHE was observed. The nature of the ORR peak and activity was evaluated from the estimation of the electron transfer number (n, number of electrons exchanged per oxygen molecule during the ORR) obtained using the Koutecky–Levich (K–L) equations (Figure b).[41] The ORR electron transfer number was determined to be n = 3.0 at a potential range from −0.10 to 0.20 V, which demonstrates almost equal contribution of two-electron and four-electron reduction pathways with the two-electron oxygen being reduced to H2O2 (according to eqs and 4) and the four-electron pathway with complete reduction to water. A mixture of of two- and four-electron pathways for ORR is common for carbon-based nanomaterials. The contribution of the glassy carbon substrate for the two-electron ORR is shown in Figure S5, revealing that the glassy carbon electrode contributes only a small catalytic effect on ORR performance. Though the HCCs are suggested for biomedical applications and not for use in fuel cells, the cycling stability shown in Figure S6 suggests that the catalytic sites are quite stable in the presence of produced aggressive radicals during the ORR.

Figure 4

(a) CVs of HCCs under 1 atm O2 or Ar. All scans were collected at 100 mV s–1 using a GC working electrode with an area of 0.196 cm2. (b) LSV curves of RDEVs of HCCs in an O2-saturated 0.1 M NaHPO4/NaH2PO4 buffer solution at pH 7 with different rotating speeds ranging from 400 to 1600 rpm. Inset: Koutecky–Levich plots of HCCs where n = 3.0. (c) RRDEV of HCC in a 0.1 M NaHPO4/NaH2PO4 buffer solution at pH 7 with a rotation speed of 1600 rpm. The disk was scanned from 0.8 to −0.2 V, while the ring electrode was held at 1.4 V. (d) The number of electrons transferred and the H2O2 yield of HCCs during the ORR calculated by ring currents.

Rotating disk electrode voltammetry (RDEV) and rotating ring-disk electrode voltammetry (RRDEV) were used to further investigate the mechanism of the O2 reduction, specifically to estimate the potential range of the ORR and determine the yield of H2O2 during the ORR. Figure c displays the RRDEV of the O2-saturated solution at 1600 rpm. The disk current goes through a gradual change when it is scanned from 0.8 to −0.2 V (consistent with the voltammograms shown in Figures a and 4b), revealing the ORR activity. At the same time, it is possible to observe the symmetric (with respect to the disk current, Figure c) changes in ring current as a function of disk potential. The change in ring current, as a catalytic response, is coupled to the formation of H2O2 as a product at the disk.

Based on the disk and ring currents in Figure , the electron transfer number (n) was 3.5 and the yield of H2O2 produced was ∼25% according to eq (see the Experimental Section for the details of the calculations). Therefore, in addition to the efficient one-electron oxidation of O2•– to O2 by HCCs in aprotic media, HCCs catalyze O2 reduction to H2O2 by a two-electron process in the presence of a proton source. This work focuses on the oxidizing properties of HCCs and the formation of H2O2 in a 25% yield through a two-electron process. Another four-electron transfer process takes place, generating H2O in a 75% yield.[14] We propose a mechanistic pathway, as shown in Scheme , based on the H2O2-producing behavior of HCCs. Hydrolysis (proton-transfer steps) of the reaction intermediates [HCC/O2•–]− and [HCC/•OOH]− is faster than the two successive electron transfer steps, which shows that the electron transfer steps are the rate-determining steps of the ORR by HCCs to form H2O2. We are not certain that this is a biologically relevant scheme for HCCs, but it is seen in electrochemical ORR processes.

Scheme 3

ORR Mechanism by HCC

These observations suggest that the SOD mimetic activity of PEG-HCCs is linked to the properties of HCCs as oxidants. After the initial electron transfer step from O2•– to the HCCs, HCC– forms as a reactive intermediate. In the presence of protons, HCC– further proceeds through a fast reaction with another equiv of O2•–, giving rise to O2 and H2O2 as the final products. The overall SOD-type activity of HCCs is summarized in eq , precisely as was predicted by our earlier EPR and stop-flow studies.[27]

Conclusions

In summary, the reduction potential of HCCs has been measured to show that HCCs are stronger oxidants than GO. Harsh fuming acid oxidation conditions give rise to increased electron withdrawing functionalities on HCCs. Both π-conjugated carbon radical centers and electron withdrawing functionalities make HCCs an excellent oxidant to oxidize O2•– to O2 in aprotic solvent by a single-electron transfer mechanism. To assess the importance of subsequent proton-coupled steps, HCCs were immobilized on a GC working electrode, and the electrocatalytic reduction of O2 to H2O2 in a neutral aqueous environment was demonstrated. The ORR electrocatalysis by HCC was indicated to proceed with an electron transfer number of 3.5. RRDEV of HCCs shows the production of H2O2 in a ∼25% yield, indicating that proton-coupled hydrolysis of electron transfer intermediates is much faster than that in the electron transfer steps. Thus, electron transfer steps are the rate-determining steps.

To obtain a good analogue for PEG-HCCs with very high SOD activity, we need to have its redox potential located between those of the two half reactions: O2•– oxidation to O2 and O2•– reduction to H2O2 (−0.16 and 0.94 V, respectively, relative to NHE in water). In this study, we demonstrated that HCCs meet redox requirements for O2•– oxidation to O2 in DMSO and aqueous buffer. To fully meet the expectation for both of the half reactions, synthetic small molecule analogues of HCCs and PEG-HCCs are currently being constructed and tested. The reactivity of the HCC intrinsic radical toward O2•– in DMSO observed in this study (Figure ) is encouraging and suggests its relevance in this process. Transient kinetic measurements by rapid-freeze quench EPR for the direct reaction between PEG-HCCs and O2•– in aqueous buffer are also under careful study to assess if this intrinsic radical indeed participates in the actual chemical steps of the catalytic mechanism.

Experimental Section

Materials

All chemicals were purchased from Sigma-Aldrich and used without further purification unless otherwise stated. SWCNTs (batch HPR 187.4) were obtained from the HiPco Laboratory at Rice University

Electrochemistry

CVs were obtained with a CHI1202 ElectroChemical Analyzer (CH Instruments) of 10 mL samples of electrolyte solutions (0.1 M [(n-Bu)4N]ClO4 solution in DMSO or PBS buffer, pH 7.4) using a three-electrode cell. A GC electrode served as the working electrode, platinum wires served as the counter electrode, and Ag/AgCl served as the reference electrode. A platinum wire was used as the pseudoreference electrode in DMSO; ferrocene (Fc) was used as the internal potential standard, and all potentials were referred to the Fc/Fc+ couple. CVs were recorded at a scan rate of 100 mV s–1.

RDE and RRDE experiments were conducted in an electrochemical cell (AutoLab PGSTST302) using a rotator (Pine Instrument, AFMSRCE) connected to an electrochemical analyzer (CH Instruments, 600D) with an Ag/AgCl reference electrode and a Pt wire counter electrode. An HCC ink solution was prepared by dispersing 4 mg of HCCs into 1 mL of 4/1 DCM/EtOH solvent, and 8 μL of the catalyst ink solution was loaded onto a GC electrode (5 mm in diameter). Constant bubbling by a stream of O2 in the cell solution was maintained throughout the measurements to ensure continuous O2 saturation. Measurements were carried out at pH 7 (0.1 M K2HPO4/KH2PO4 buffer). For RRDE experiments, the electrode rotation speed was 1600 rpm (scan rate: 0.05 V/s; platinum data collected from anodic sweeps), while the ring electrode potential was held at 1.1 V vs the reversible hydrogen electrode (RHE).

The O2 reduction current increases with increasing rotation rates following K–L eq :where JK is the potential dependent kinetic current and JL is the Levich current. JL is expressed as 0.62nF[O2](DO2)2/3ω1/2ν –1/6, where n is the number of electrons transferred to the substrate, F is the Faraday constant, [O2] is the concentration of O2 in an air-saturated buffer (0.22 mM in the case of pH 7) at 25 °C, DO2 is the diffusion coefficient of O2 (1.8 × 10–5 cm2 s–1 at pH 7) at 25 °C, ω is the angular velocity of the disc, and ν is the kinematic viscosity of the solution (0.009 cm2 s–1) at 25 °C. When eq is rewritten as eq and solved for JL, it gives eq :RRDE measurements were carried out to determine the H2O2 yield (%) and n, which were calculated by eqs and 11:where id and ir are the disk and ring currents, respectively. N is the ring current collection efficiency which was determined to be 25% by the reduction of 10 mM K3[Fe(CN)6] in 0.1 M KNO3.

XPS

XPS spectra were obtained on a PHI Quantera SXM scanning X-ray microprobe system using a 100 μm X-ray beam with a takeoff angle of 45° and a pass energy of 140 eV for the survey and 26 eV for the high-resolution elemental analysis.

Detection of Radicals by EPR Spectroscopy

EPR spectra of HCCs (0.05 mg/mL in DMSO) in a capillary tube at ambient temperature were recorded using the following parameters: center field 3320 G, sweep width 50 G, microwave frequency 9.3 GHz, microwave power 1 mW, modulation frequency 100 kHz, and modulation amplitude 1.0 G. The same sample was remeasured after adding a small amount of KO2 (1 mg in powder form).