Journal of The Electrochemical Society, 159 (11) G151-G159 (2012) G1510013-4651/2012/159(11)/G151/9/$28.00 © The Electrochemical Society

The Electrochemical Oxidation of 6-Aminoquinoline:Computational and Voltammetric Study

Milica C. Stevic,a,z Gordana Ciric-Marjanovic,a Budimir Marjanovic,bLjubisa M. Ignjatovic,a and Dragan Manojlovicc

aFaculty of Physical Chemistry, University of Belgrade, 11158 Belgrade, SerbiabCentrohem, 22300 Stara Pazova, SerbiacFaculty of Chemistry, University of Belgrade, 11158 Belgrade, Serbia

The theoretical study of the 6-aminoquinoline (6-QNH2) electrochemical oxidation mechanism, based on the semi-empirical quantumchemical computations of the heat of formation, ionization energy, and spin density of reaction intermediates, taking into accountthe influence of pH and solvation effects, has been conducted. Two possible 6-QNH2 electro-oxidation pathways are investigated,namely, the initial single-electron oxidation of 6-QNH2 at lower electrode potentials, leading to the formation of cation radicals[6-QNH2]•+ in acidic solutions and neutral radicals [6-QNH]• in alkaline solutions, as well as the two-electron oxidation of6-QNH2 leading to the initial formation of nitrenium cations [6-QNH]+ at higher electrode potentials. The regioselectivity of6-QNH2 dimerization reactions, which follow both the single- and two-electron transfer reactions, is computationally studied. Cyclicvoltammetry experiments, conducted at a glassy carbon paste electrode in Britton-Robinson buffer/methanol media, are correlatedwith the computationally predicted 6-QNH2 electro-oxidation mechanism. Differential pulse voltammetric and adsorptive strippingdifferential pulse voltammetric analyses of 6-QNH2 oxidation have also been performed.© 2012 The Electrochemical Society. [DOI: 10.1149/2.004212jes] All rights reserved.

Manuscript submitted July 30, 2012; revised manuscript received August 21, 2012. Published September 21, 2012.

Electrochemical oxidations of heterocyclic aromatic compounds,such as pyrrole and thiophene, as well as the electro-oxidations ofcarbocyclic arylamines, such as aniline, aminonaphthalenes, etc.,have been the subject of numerous studies during the decades.1

Special attention has been paid to the electrochemical oxidativepolymerizations of heterocyclic aromatic compounds and carbocyclicarylamines, which lead to the formation of corresponding conductingpolymers (polypyrrole, polythiophene, polyaniline, etc.).2 Muchless attention, both from the experimental and theoretical pointsof view, has been focused on the electrochemical oxidation of hete-rocyclic arylamines, e.g. aminopyridines,3–8 aminoimidazoles,9

aminothiophenes,10–14 aminoindoles,15 aminoacridines,16

aminophenazines,17 aminopurines,18–20 aminothiazoles,21,22 andaminobenzothiazoles.23

Aminoquinolines are heterocyclic arylamines with considerableindustrial use, especially as intermediates in the pharmaceuticalindustry.24 The electrochemical oxidation of aminoquinolines hasscarcely been investigated.25–29 Based on various electroanalyticaltechniques, 2,2′-azoquinoline was proposed as the major product ofthe oxidation of 2-aminoquinoline at a stationary pyrolytic graphiteelectrode in methanol-phosphate buffer.25 No products were isolatedfrom the attempted electrochemical oxidation of 3-aminoquinoline(3-QNH2) in 75−80% sulfuric acid solution at a platinum electrode.26

Poly(5-aminoquinoline) films were prepared on gold electrodes byanodic polymerization of 5-aminoquinoline (5-QNH2) in acetonitrileusing cyclic voltammetry and potential step methods.27 Voltammetricdetermination of 3-QNH2, 5-QNH2, 6-QNH2, and 8-QNH2, separatelyor in their mixtures, by differential pulse voltammetry and adsorptivestripping differential pulse voltammetry on carbon paste electrode wasreported by Zima et al.28,29

Recently, we have successfully applied a combined computa-tional and experimental approach in order to obtain new insightsinto the mechanism of chemical oxidative polymerization of hetero-cyclic arylamines such as aminophenazines30 and aminoacridines,31

as well as in order to determine unambiguously the main structuralunits of the obtained oligomeric and polymeric oxidation products.We have recently used successfully a combined computational andexperimental approach in the case of electrochemical oxidation ofhydroxyquinolines.32,33 In the present communication we reporteda semi-empirical quantum chemical study of the electro-oxidationof 6-QNH2 (3, Scheme 1) in aqueous solution. To our knowledge,there has been no computational study of the electrochemical ox-

zE-mail: milica@ffh.bg.ac.rs

idation of aminoquinolines. We correlated a theoretically proposed6-QNH2 electro-oxidation mechanism with cyclovoltammetric (CV)behavior of 6-QNH2 at a glassy carbon paste (GCPE) working elec-trode, in a supporting electrolyte of mixtures of Britton-Robinsonbuffer/methanol. GCPE exhibits a very low background current, has awide potential range,34 combines the favorable electron-transfer kinet-ics of glassy carbon (GC) with the attractive advantages of compositepaste electrode materials,35 and is compatible with a high contentof methanol.36,37 Differential pulse voltammetry (DPV) was used forestablishing the optimum conditions for 6-QNH2 determination andfinding the limits of detection (LOD). Adsorptive stripping differentialpulse voltammetry (AdSDPV) has also been explored as a techniquefor sensitive quantitative analysis of the investigated substance.

Computational Methods

A semi-empirical AM1 method38 (included in the molecular or-bital package MOPAC 97,39 a part of the Chem3D Pro 5.0 pack-age, CambridgeSoft Corporation) has been used to obtain the heatof formation (�Hf), and spin density of investigated species. Thismethod is accurate enough to have useful predictive power and it isfast enough to allow the processing of oligomers.32,33,40–45 The ion-ization energy (Ei) of 6-QNH2 in different acid-base forms, as well asEi of their oxidation products was computed by using RM1 method46

(improved/reparameterized version of AM1 included in the MOPAC2009, a part of the ChemBio3D Ultra 12.0, CambridgeSoft Corpo-ration). Solvation effects were taken into account using COSMO47

(the conductor-like screening model) to approximate the effect of sol-vent surrounding the analyzed molecules, ions and free radicals. Con-formational analysis of reaction intermediates was done. The stericenergy was minimized using MM2 molecular mechanics force-fieldmethod.48 Input files for the semi-empirical quantum-chemical com-putations of dimeric 6-QNH2 intermediates were the most stable con-formers of the investigated molecular structures. The geometrical op-timization was performed by the eigenvector following procedure.49

The restricted Hartree−Fock method has been used for the inves-tigated molecules and ions, while the unrestricted Hartree−Fockmethod has been used for free radical species.

Experimental

Reagents and chemicals.— 6-QNH2 (95%) and glassy carbonpowder, comprising spherical particles with diameters in the range2−12 μm, were supplied by Sigma−Aldrich (Poole, UK). Methanol

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G152 Journal of The Electrochemical Society, 159 (11) G151-G159 (2012)

Scheme 1. 6-Aminoquinoline acid-base equlibria and the initial electro-oxidation.

(HPLC grade) was purchased from J.T. Baker (Deventer, The Nether-lands). Paraffin oil (liquid petrolatum; white mineral oil; Nujol), ortho-phosphoric acid, boric acid, acetic acid, and sodium hydroxide, allof analytical grade purity, were obtained from Merck (Darmstadt,Germany). Deionized water, produced by an Ultra Clear Basic SG Wa-ter apparatus (SG Wasseraufbereitung GmbH, Germany), was used.

Apparatus.— CV and DPV measurements were carried out using acomputerized voltammetric analyzer CH Instruments (Austin, USA)driven by voltammetric software CHI (Version 4.03). A three elec-trode measurement system was used, comprising a saturated calomelelectrode (SCE) as a reference electrode, in respect to which all mea-sured potentials are given, a platinum wire as an auxiliary electrode,and a GCPE for CV and DPV, as a working electrode. The electrodeswere kept in triangular geometry. For pH measurements, a pH meter,model 744, equipped with combined glass pH electrode and tem-perature sensor, all from Metrohm (Herisau, Switzerland) was used.The UV-visible (UV-vis) spectra were recorded on a Cintra GB-10UV-visible spectrophotometer (GBC Scientific Equipment Pty Ltd.,Australia). Quartz cuvettes of 1.00 cm optical path length were used.

Solutions and electrode preparation.— The stock standard solu-tion (5 × 10−2 M) of the 6-QNH2 was prepared by dissolving accu-rately weighed amounts of 6-QNH2 in methanol. More dilute solutionswere prepared by serial dilution of the stock solution with methanol.All solutions were stored in a dark place. A spectrophotometric studydemonstrated that the stock standard solution was stable for at leastthree months.

The Britton−Robinson (BR) buffer solutions were prepared bymixing appropriate volumes of acidic and basic buffer components.The acidic buffer component comprises 0.04 M ortho-phosphoric acid,0.04 M boric acid, and 0.04 M acetic acid. The basic buffer componentis 0.2 M sodium hydroxide solution. Supporting electrolytes weremade by mixing the BR buffer of appropriate pH and methanol in theratio 80:20 (v/v).

The glassy carbon paste was prepared by hand-mixing 250 mgof glassy carbon powder with 90 μL of mineral oil in a mortar andpestle until a completely homogeneous stiff paste was obtained. The

resulting paste was pressed firmly into the 2 mm-internal diameter and20 mm long Teflon electrode body cylinder, to produce the GCPE.

Procedures.— UV-vis spectra were obtained for 3 × 10−5 Msolution of 6-QNH2 in BR buffer/methanol, containing 20% ofmethanol (v/v), pH range 2−12, in spectral range 200−500 nm. As ablank solution, supporting electrolyte, BR buffer/methanol, contain-ing 20% of methanol (v/v), of appropriate pH, was used.

Voltammetric investigations were started by recording cyclicvoltammograms (CVs), using different scan rates, of the supportingelectrolyte at pH 2.00. The scan rates used for CV measurements inthis study were 5, 10, 20, 50, and 100 mV/s. After recording theseblank CVs, using 20 mL of the supporting electrolyte pH 2.00 in thevoltammetric cell, 160 μL of the 5 × 10−2 M 6-QNH2 stock standardsolution was added thus producing 4 × 10−4 M solution of 6-QNH2.

The solution mixing was achieved by a mechanical rotating stirrerand, after allowing the solution to equilibrate for 10 s, CV measure-ments at the above-mentioned scan rates, in the potential range from−0.4 to 1.5 V were carried out. The GCPE had to be cleaned afterevery measurement by wiping its surface on a filter paper wetted withdeionized water. After recording CVs for the 4 × 10−4 M 6-QNH2

solution at pH 2.00, the whole procedure was repeated using support-ing electrolytes with the following pH values: 5.05, 5.94, 7.06, 8.98,and 12.03.

Further CVs of the 1 × 10−4 M solution of 6-QNH2 were measuredin the potential range from −0.4 to 1.3 V for the supporting electrolytesBR buffer/methanol, containing 20% of methanol (v/v), pH range2−12, at scan rates 5, 10, 20, 50, and 100 mV/s. DPVs of the same so-lution were measured in the potential range 0.0−1.2 V, for the same pHrange 2−12, scan rate 20 mV/s, pulse amplitude 50 mV.

Continuous recording of fifty CVs was carried out for1 × 10−4 M 6-QNH2 solutions of pH 2.00, 5.05, 7.06, and 12.03 in BRbuffer/methanol supporting electrolyte, containing 20% of methanol(v/v). The scan rate used here was 50 mV/s and, as opposed to thepreviously described procedure, the GCPE surface was not cleanedafter each measuring cycle.

The 6-QNH2 calibration curves were determined recording differ-ential pulse voltammograms (DPVs) in triplicate in the concentrationrange 1 × 10−6−1 × 10−4 M in a supporting electrolyte pH 2.00,and evaluated by the least-squares linear regression method. LOD for6-QNH2 was obtained in the widely-accepted conventional way.50

Results and Discussion

Computational study.— The possible single-electron and two-electron transfer reaction pathways of the electrochemical oxidation of6-QNH2 have been computationally analyzed. The known acid−baseequilibria of 6-QNH2

51 (1 → 2, pKa1 = 1.63; 2 → 3, pKa2 = 5.59;Scheme 1) indicate the initial formation of trication radicals [H-6-QNH3]•3+ (4), dication radicals [6-QNH3]•2+ (5), and cation radicals[6-QNH2]•+ represented as a resonance hybrid (6), upon the single-electron oxidation of [H-6-QNH3]2+ (1), 6-QNH3

+ (2), and 6-QNH2

(3), respectively (Scheme 1). RM1/COSMO computations of ioniza-tion energy (Ei) of different acid−base forms of 6-QNH2 have shownthat their oxidizability increases in the following order: [H-6-QNH3]2+

(1, Ei = 10.18 eV) < 6-QNH3+ (2, Ei = 9.61 eV) < 6-QNH2 (3, Ei

= 8.15 eV), i.e. the oxidizability of 6-QNH2 increases with increasingpH. It should be noted that the major significance of the computation-ally determined values of Ei of investigated species is their relativeorder rather than their absolute values, which could be obtained moreaccurately by advanced DFT or ab initio methods. The initial forma-tion of [H-6-QNH3]•3+ species (4) is expected to occur prevalentlyat pH < 1.6 where more than 50% of 6-QNH2 molecules exist indiprotonated form [H-6-QNH3]2+. However, it is important to notethat their initial formation should not be neglected up to pH ∼ 3, since∼9% of 6-QNH2 molecules exist in diprotonated form [H-6-QNH3]2+

at pH 2.6. The [6-QNH3]•2+ species (5) could initially be formed insubstantial quantities only in the pH region 1.6 < pH < 5.6 wheremore than 50% of 6-QNH2 molecules exist in monoprotonated form

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Journal of The Electrochemical Society, 159 (11) G151-G159 (2012) G153

Scheme 2. 6-Aminoquinoline cation radical ([6-QNH2]+•) dimerization ac-companied by the two-electron oxidation of the formed dimer. Spin density ofthe main reactive sites of [6-QNH2]+• is also shown.

(6-QNH3+). The initial formation of [6-QNH2]+• species is expected

to prevail at pH > 5.6 where more than 50% of 6-QNH2 moleculesexist in nonprotonated form. Once formed at electrode surface, theradical species [H-6-QNH3]•3+ and [6-QNH3]•2+ as a much moreacidic species than parent [H-6-QNH3]2+ and 6-QNH3

+, respectively,undergo fast deprotonation to [6-QNH2]•+ in acidic aqueous solution.Taking into account the fact that the pKa of the aniline cation radi-cal is 7.05,52 it can be supposed that neutral radicals [6-QNH]• (10,Scheme 1), formed by deprotonation of [6-QNH2]•+, prevail in alka-line aqueous solutions, while [6-QNH2]•+ are the prevalent free radi-cal species in acidic solutions. It can be concluded that free radical re-combination reactions of reactive species [6-QNH2]•+ and [6-QNH]•

represent the main 6-QNH2 dimerization pathways which accompanythe initial single-electron oxidation in acidic and alkaline aqueoussolutions, respectively.

The regioselectivity of 6-QNH2 dimer formation through the re-combination of [6-QNH2]•+ (11, Scheme 2) and [6-QNH]• (10,Scheme 3) is analyzed by AM1/COSMO computational method. Thecomputationally determined spin density of both [6-QNH2]•+ and [6-QNH]• free radical species (Schemes 2 and 3) indicates C2, C4, C5,C7, and NNH2 as the main reactive sites for free radical coupling reac-tions. According to the Hammond postulate,53 the regioselectivity ofthese reactions is not controlled by the stability of final dimer products(13, Scheme 2, and 17, Scheme 3) but it is governed by the stabil-ity of intermediates (12, Scheme 2, and 16, Scheme 3) structurallyresembling the corresponding transition states. It is computationallyshown that C5–C5 coupling is the prevalent 6-QNH2 dimerizationroute in acidic solutions, i.e., in the case of [6-QNH2]+• recombina-tion reaction (Scheme 2, Table I), while NH–C5 coupled dimers areprevalently formed by the free radical recombination of [6-QNH]• in

Scheme 3. 6-Aminoquinoline radical ([6-QNH]•) dimerization accompaniedby the two-electron oxidation of formed dimer. Spin density of the mainreactive sites of [6-QNH]• is also shown.

alkaline solutions (Scheme 3, Table II). The two-electron oxidation ofthe formed C5–C5 (13, Ei = 8.02 eV) and NH–C5 (17, Ei = 8.11 eV)coupled 6-QNH2 dimers, which have slightly lower Ei in comparisonwith 6-QNH2 as indicated by RM1 computations, leads to the forma-tion of the corresponding quinonoid compounds in different acid-baseforms (14 and 15, Scheme 2, and 18 and 19, Scheme 3). It followsthat the initial single-electron oxidation of 6-QNH2 at relatively lowelectrode potentials in a broad pH range, accompanied by both thedimerization of the formed free radical species and the subsequentoxidation of prevalent 6-QNH2 dimers to corresponding quinonoidcompounds, represents an overall two-electron transfer process per6-QNH2 molecule.

The significantly lower oxidizability, i.e, the significantly higherEi of the cation radical species [H-6-QNH3]•3+ (4, Ei = 11.39 eV),[6-QNH3]•2+ (5, Ei = 10.72 eV), and [6-QNH2]•+ (the resonance hy-brid 6, Ei = 9.70 eV) in comparison with the parent molecular species[H-6-QNH3]2+ (1), 6-QNH3

+ (2), and 6-QNH2 (3), respectively(Scheme 1), indicates that the electrode potentials corresponding to thesingle-electron oxidations of cation radicals [H-6-QNH3]•3+,

Table I. The heat of formation, �Hf, of 6-QNH2 dimer dicationintermediates, formed by the radical recombination reaction of[6-QNH2]•+, computed by AM1/COSMO method.

Coupling �Hf Coupling �Hf Coupling �Hf

mode kcal mol−1 mode kcal mol−1 mode kcal mol−1

C2−C2 349.9 C4−C4 349.2 C5−C7 331.8C2−C4 349.1 C4−C5 332.9 C5−NH2 324.6C2−C5 331.1 C4−C7 345.9 C7−C7 348.7C2−C7 344.3 C4−NH2 336.8 C7−NH2 340.8C2−NH2 338.5 C5−C5 318.5 H2N−NH2 368.6

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G154 Journal of The Electrochemical Society, 159 (11) G151-G159 (2012)

Table II. The heat of formation, �Hf, of 6-QNH2 dimerintermediates, formed by the radical recombination reaction of[6-QNH]•, computed by AM1/COSMO method.

Coupling �Hf Coupling �Hf Coupling �Hf

mode kcal mol−1 mode kcal mol−1 mode kcal mol−1

C2−C2 167.3 C4−C4 166.8 C5−C7 141.3C2−C4 166.0 C4−C5 145.6 C5−NH 115.4C2−C5 146.8 C4−C7 160.4 C7−C7 157.8C2−C7 161.8 C4−NH 135.8 C7−NH 131.9C2−NH 139.2 C5−C5 126.4 HN−NH 119.5

[6-QNH3]•2+, and [6-QNH2]•+ are significantly higher in compari-son with the electrode potentials corresponding to the their formation.It is important to note that the ionization energy of [6-QNH]• (10, Ei

= 9.12 eV) is considerably lower than that of all cation radical species([H-6-QNH3]•3+, [6-QNH3]•2+, and [6-QNH2]•+), thus indicating thatthe initial two-electron oxidation of 6-QNH2 occurs at significantlylower anodic potential in alkaline solutions than in acidic solutions.We propose that 6-QNH2, similarly to other arylamines such as ani-line which cannot give stable quinonoid compounds upon the two-electron oxidation, gives its nitrenium cation [6-QNH]+ (8, Scheme 1)as a major product of the initial two-electron oxidation at relativelyhigh electrode potentials in a broad pH range, these potentials beingmuch higher in acidic than in alkaline solutions. The formed nitre-nium cation is a highly reactive electrophilic species which can reactfurther with unoxidized 6-QNH2, forming 6-QNH2 dimers. The re-gioselectivity of the electrophilic aromatic substitution reaction of6-QNH2 with its nitrenium cation, governed by the stability of cationdimer intermediates (20, Scheme 4) rather than by the stability offinal dimeric products (17, Scheme 4), according to the Hammondpostulate,53 is analyzed by AM1/COSMO computational method. Itis computationally shown that the major product of the reaction of[6-QNH]+ with 6-QNH2 is the NH–C5 coupled dimer (Table III,Scheme 4), which further undergoes a two-electron oxidation to cor-responding quinonoid compounds in different acid−base forms (18and 19, Scheme 3). It follows that the initial two-electron oxidationof 6-QNH2 at relatively high electrode potentials in a broad pH range,accompanied by both the reaction of formed [6-QNH]+ with 6-QNH2

and the subsequent oxidation of prevalent NH–C5 coupled 6-QNH2

dimer to a corresponding quinonoid compound, represents an overalltwo-electron transfer process per 6-QNH2 molecule.

Experimental study.— UV-vis Spectrophotometry.— From the UV-vis spectra (Fig. 1) it can be seen that the investigated substance inthe pH region 2.0−4.0 (pKa1 + 0.35 ≤ pH ≤ pKa2 − 1.5) exhibits

Scheme 4. Electrophilic aromatic substitution reaction of 6-QNH2 with itsnitrenium cation [6-QNH]+.

Table III. The heat of formation, �Hf, of 6-QNH2 dimer cationintermediates, formed by the electrophilic aromatic substitutionreaction of 6-QNH2 with [6-QNH]+, computed by AM1/COSMOmethod.

Coupling �Hf Coupling �Hf Coupling �Hf

mode kcal mol−1 mode kcal mol−1 mode kcal mol−1

C2−C2 258.2 C4−C4 256.7 C5−C7 236.6C2−C4 257.0 C4−C5 237.3 C5−NH 208.9C2−C5 237.6 C4−C7 251.9 C7−C7 250.8C2−C7 254.1 C4−NH 228.0 C7−NH 223.2C2−NH 230.7 C5−C5 221.1 H2N−NH 236.5

an absorption maximum at 259 nm which can be attributed to the6-QNH3

+, while in the pH region 7.1−11.8 (pH ≥ pKa2 + 1.5) theabsorption maximum which can be attributed to the 6-QNH2 appearsat 241 nm. In the spectra at pH ∼5 and ∼6 absorption bands appear atboth mentioned wavelengths, indicating that both species, 6-QNH3

+

and 6-QNH2, are present in the pH range 5−6 (pH ∼ pKa2 ± 0.5).Our results are in accordance with the results of a previous spec-trophotometric study of 6-QNH2 acid−base equilibria.51 Taking intoaccount that pKa1 = 1.63,51 it can be calculated that ∼30% of 6-QNH2

molecules exist in diprotonated form [H-6-QNH3]2+ at pH 2, while∼4% of 6-QNH2 are diprotonated at pH 3. However, the distinctiveabsorption band which can be attributed to [H-6-QNH3]2+ specieswas not detected at pH 2 and 3 because it is most likely overlappedby absorption band of 6-QNH3

+ which is the dominant species in thatpH region.

Cyclic voltammetry.— In order to understand the 6-QNH2 oxida-tive electrochemical processes taking place at the surface of a GCPE,CV measurements were performed in pH range 2–12 (Fig. 2). It isobserved that the potentials of anodic peaks become more negativewith increasing pH of the supporting electrolyte. This CV behav-ior is in accordance with our computational prediction that 6-QNH2

oxidizability increases with increasing pH.In the potential range from 0.0 to 1.2 V vs. SCE (Fig. 2) two

poorly-separated anodic peaks appear in the potential range 0.9–1.2 Vat pH 2 and 3 (Fig. 2a, curves 1 and 2), the well defined anodic peakappears at about 0.9 V at pH 4 (Fig. 2a, curve 3), the anodic peakswith shoulders appear in the potential range 0.5–0.9 V at pH 5–9(Fig. 2a, curve 4; Fig. 2b, curves 5–8), while only one broad anodicpeak appears in the potential range 0.4–0.7 V at pH 10–12 (Fig. 2b,

Figure 1. UV-vis spectra of 3 × 10−5 6-QNH2 in BR buffer/methanol, con-taining 20% of methanol (v/v), pH: 1.98 (1); 3.03 (2); 4.01 (3); 4.94 (4); 6.05(5); 7.08 (6); 8.01 (7); 8.89 (8); 9.90 (9); 10.80 (10); 11.82 (11).

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Figure 2. Cyclic voltammograms of 1 × 10−4 M 6-QNH2, obtained at a clean GCPE surface, in BR buffer/methanol, containing 20% of methanol (v/v), pH(a): 2.00 (1); 3.08 (2); 4.06 (3); 5.05 (4); (b): 5.94 (5); 7.06 (6); 8.03 (7); 8.98 (8); 10.02 (9); 11.01 (10); 12.03 (11); scan rate 20 mV/s.

curves 9–11). Combining semi-empirical calculations and CV results,and taking into account the acid-base properties of 6-QNH2 and itsinitial electro-oxidation products, it can be proposed that the peakobserved at lower anodic potentials in the pH range 2–3 (Fig. 2a,curves 1 and 2), as well as the anodic peak observed about 0.9 V atpH 4 (Fig. 2a, curve 3) appears as a consequence of the initial single-electron oxidation of the 6-QNH3

+ (Reaction 1), accompanied withthe recombination of [6-QNH2]•+ and the oxidation of formed C5–C5coupled dimers (Scheme 2).

6-QNH3+ − e− → [6-QNH2]•+ + H+ [1]

The peak which appears at higher anodic potentials in the pH range2–3 (Fig. 2a, curves 1 and 2) can be attributed to the initial singleelectron oxidation of [H-6-QNH3]2+ (1) (Reaction 2), accompaniedby the recombination of [6-QNH2]•+ and the oxidation of formedC5–C5 coupled dimers (Scheme 2).

[H-6-QNH3]2+ − e− → [6-QNH2]•+ + 2H+ [2]

The anodic peaks with shoulders observed in the pH region 5–9(Fig. 2a, curve 4; Fig. 1b, curves 5–8) appear most probably as a con-sequence of the combined initial single-electron oxidations of boththe 6-QNH3

+ (Reaction 1) and 6-QNH2 (Reactions 3 and 4), accom-

panied by the recombination of [6-QNH2]•+ (Scheme 2), generated inacidic solutions (Reactions 1 and 3), and [6-QNH]• (Scheme 3), gen-erated in alkaline solutions (Reaction 4), and subsequently followedby the oxidation of formed C5–C5 (Scheme 2) and NH–C5 dimers(Scheme 3). The contribution of the initial single-electron oxidationsof 6-QNH2 (Reactions 3 and 4) to the peaks observed in the pH region5–9 (Fig. 2) increases with the increase of pH.

6-QNH2 − e− → [6-QNH2]•+ [3]

6-QNH2 − e− → [6-QNH]• + H+ [4]

The broad peaks which appear at pH ≥ 10 in the potential range0.4–0.7 V (Fig. 2b) are most likely due to the single-electron oxidationof the 6-QNH2 (Reaction 4), accompanied by the recombination of[6-QNH]• (Scheme 3), and subsequently followed by the oxidation offormed NH–C5 dimers (Scheme 3).

Our semi-empirical quantum chemical computations indicate thatthe peaks corresponding to the initial two-electron oxidations of6-QNH2 (6-QNH3

+, [H-6-QNH3]2+), accompanied by the reactionof generated [6-QNH]+ with 6-QNH2 (Scheme 4) and followed bythe oxidation of formed NH–C5 dimers (Scheme 3), are expected tobe observed at potentials higher than those corresponding to the ini-tial single-electron oxidations. Indeed, it can be seen that for pH≤ 6, besides the already observed anodic peaks in the potential

Figure 3. Cyclic voltammograms of 4 × 10−4 M 6-QNH2 obtained for a wider anodic potential limit of 1.5 V vs. SCE, at a clean GCPE surface, in BRbuffer/methanol, containing 20% of methanol (v/v), scan rate 20 mV/s, pH (a): 2.00 (1), 5.05 (2), 5.94 (3); arrows indicate peaks that appear at E > 1.2 V;(b): 7.06 (1); 8.98 (2); 12.03 (3).

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Figure 4. Continuous cyclic voltammograms of 1 × 10−4 M 6-QNH2 on GCPE in BR buffer/methanol, containing 20% of methanol (v/v), at different pH: 2,00(a); 5.05 (b); 7.06 (c); 12.03 (d); recorded after different numbers of cycle(s): one (1), two (2), three (3), ten (4), and fifty (5); scan rate 50 mV/s. The anodicpre-peaks (PP) are indicated. The enlarged areas with pre-peaks are shown in the insets.

range <1.2 V which correspond to the initial single-electron oxi-dations, the anodic peaks appear also in the potential range >1.2 V(Fig. 3a, the peaks marked with arrows), while for pH > 6 these peaksare not observed (Fig. 3b). It should be noted here that the investi-gation of these peaks in alkaline solutions, when the appearance ofthese peaks is expected at lower potentials than in the case of acidicsolutions, is limited due to the supporting electrolyte’s oxidation.

Continuous CVs at four pH values are also presented (Fig. 4).The decrease of the heights of the anodic peaks with the increase inthe number of cycles and the appearance of pre-peaks (PP) in theanodic region are observed in acidic as well as in alkaline solutions.This can be explained as the consequence of the complex 6-QNH2

electro-oxidation process which leads to 6-QNH2 dimerization at thebeginning of the oxidation, and subsequently to the oligomerization/polymerization. The pre-peaks correspond most probably to theelectro-oxidation of dimers and higher oligomers of 6-QNH2 whichare more oxidizable than 6-QNH2, as computed by RM1 method. Theformation of non-conducting oligo/poly(6-QNH2) at the electrode sur-face may cause an increase in its passivation with the increase in thecycles number.

The influence of the pH of supporting electrolyte on the potentialof the first anodic CV peak (Epa) is depicted in Fig. 5. The interceptthat can be detected at pH 5.41 corresponds well to pKa2 (6-QNH3

+/6-QNH2). The linear parts of the Epa− pH dependence in the particular

Figure 5. Influence of the supporting electrolyte pH on the CV first anodicpeak potential; scan rate 5 mV/s; 6-QNH2 concentration 1 × 10−4 M. Whitesquare represents pKa2 value of 6-QNH2. The equations of the linear parts aregiven.

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Figure 6. Cyclic voltammograms of 1 × 10−4 M 6-QNH2, obtained at aclean GCPE surface, in BR buffer/methanol, containing 20% of methanol(v/v), pH 2.00, at different scan rates: 5 mV/s (1); 20 mV/s (2); 50 mV/s (3);100 mV/s (4); 200 mV/s (5). The dependence of the first anodic CV peakcurrent logarithm, log Ip, on the scan rate logarithm, log v, is given in theinset.

pH regions are expressed by the following equations:

pH < 5.41 : Epa(V) = 1.28 − 0.08 × pH

pH > 5.41 : Epa(V) = 1.08 − 0.05 × pH

In the cathodic area of the investigated pH range 2−12 one peakappears (Fig. 1). Taking into account that the 6-QNH2 oxidation onGCPE represents an overall two-electron transfer process, accordingto our computational predictions, the following facts are apparent: (I)the anodic and the cathodic peak potential differences are much biggerthan 28 mV, expected for the reversible two-electron transfer process(Fig. 2); (II) the anodic and cathodic peak current ratio (Ipa/Ipc) aftera certain number of cycles is much higher than unity; (III) the peakpotential is a function of the scan rate (Fig. 6). It can be unambiguouslyconcluded, in accordance with computational predictions, that theelectrode process is irreversible in the investigated pH range due to thefast chemical reactions which follow the electron transfer processesat electrode.

The rate of the electrochemical oxidation of 6-QNH2 is controlledby the diffusion of the electroactive species (as the slowest step inthe electrochemical process). The diffusion nature of the 6-QNH2

electro-oxidation process is confirmed by the linear dependence ofthe logarithm of the anodic CV peak current on the logarithm of thescan rate, log Ip = f (log v), (Fig. 6, inset), with r = 0.9966 and slope0.49 (which is close to 0.50, the theoretical value for the slope of adiffusion controlled process54).

Determination of 6-QNH2 using DPV.— The influence of pH onthe DP voltammetric behavior of 6-QNH2 was investigated. The de-pendence of the DPV currents, I, as a function of pH in the potentialrange from 0.0 to 1.2 V vs. SCE is depicted in the Figure 7. Inaccordance with computational predictions and CV measurements,the potentials of the anodic peaks shift toward more negative val-ues with increasing pH. Similarly to CV findings, the anodic DPVpeaks with shoulders observed at pH 2−3 can be assigned to the ini-tial single-electron oxidations of the [H-6-QNH3]2+ and 6-QNH3

+,the anodic DPV peak at pH 4 is due to the initial single-electronoxidation of the 6-QNH3

+, the anodic DPV peaks with shouldersobserved at pH 5−9 can be attributed to the initial single-electronoxidations of the 6-QNH3

+ and 6-QNH2, while the broad DPV peaks

Table IV. Parameters of the calibration plots for DP voltammetricdetermination of 6-QNH2 in BR buffer/methanol, containing 20%of methanol (v/v), pH 2.00, on GCPE.

A ± S(A) B ± S(B) LODmol L−1 mA/mol L−1 nA r mol L−1

(1−10) × 10−5 22 ± 1 −64 ± 48 0.9956 −(1−10) × 10−6 31 ± 1 −24 ± 3 0.9996 4.3 × 10−7

which appear at pH ≥ 10 are due to the single-electron oxidation ofthe 6-QNH2.

Calibration curves for the quantitative determination of 6-QNH2

on GCPE were obtained at pH 2, for the anodic peak at about 0.95 V(marked with an arrow) corresponding to the single-electron oxidationof the 6-QNH3

+, in the concentration range 1 × 10−6−1 × 10−4 M.The statistical parameters − slope (A) and its standard error S(A),intercept (B) and its standard error S(B), correlation coefficient (r),and LOD are summarized in Table IV. For the concentration range(1−10) × 10−6 M calculated LOD for 6-QNH2 is 4.3 × 10−7 M.

In order to improve the sensitivity of 6-QNH2 determination wehave conducted a preliminary investigation using AdSDPV. The an-alyte concentration in the investigated solution was 1 × 10−6 M.The best deposition potential, i.e. the potential with the highest DPVpeak current, has been found by recording DPVs after 1 min depo-sition time at the following potentials: 300, 200, 100, and 0 mV.55

It was found that the highest DPV peak current was obtained by ap-plying a 1 min deposition at 0.0 mV, so this potential was selectedas the optimum deposition potential, Edep. Afterwards, at Edep, differ-ent deposition times, tdep: 0, 5, 10, 20, 30, 60, 150, 300, and 600 s,were applied. After each deposition time the DPV were recorded in0.0−1.2 V range.

The resulting peak heights, Ip, are presented as a function of tdep inFigure 8, where it can be seen that 50 s is the optimal deposition pe-riod at 0.0 V. Under these conditions about a 50% higher peak currentwas produced than without any deposition period. Therefore, we con-cluded that AdSDPV is more sensitive than DPV for the quantitativedetermination of 6-QNH2.

Figure 7. DP voltammograms of 1 × 10−4 M 6-QNH2 on GCPE, in BRbuffer/methanol, containing 20% of methanol (v/v), pH: 2.00 (1); 3.08 (2);4.06 (3); 5.05 (4); 5.94 (5); 7.06 (6); 8.03 (7); 8.98 (8); 10.02 (9); 11.01 (10);12.03 (11); scan rate 20 mV/s, pulse amplitude 50 mV. The arrow indicates thepeak used for obtaining calibration plots.

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G158 Journal of The Electrochemical Society, 159 (11) G151-G159 (2012)

Figure 8. The influence of the deposition time on DP peak current, Ip, in 1× 10−6 M 6-QNH2 in BR buffer/methanol, containing 20% of methanol (v/v),pH 2.00, on GCPE; scan rate 20 mV/s, pulse amplitude 50 mV, depositionpotential 0.0 V.

Conclusions

CV and DPV measurements, in excellent agreement with RM1semi-empirical quantum chemical computations, have clearly shownthat the oxidizability of 6-QNH2 increases with increasing pH of theelectrolyte solution. Based on the acid-base properties of 6-QNH2

and the computationally determined ionization energy of its differ-ent acid-base forms, in combination with CV and DPV data in thepotential range 0.0–1.2 V vs. SCE, it was concluded that both the ini-tial single-electron oxidations of diprotonated ([H-6-QNH3]2+) andmonoprotonated (6-QNH3

+) form of 6-QNH2 occur at pH ≤ 3, closeto pKa1. It was also concluded that the initial single-electron oxida-tion of 6-QNH3

+ takes place at pH 4 close to (pKa1 + pKa2)/2, boththe initial single-electron oxidations of 6-QNH3

+ and non-protonated6-QNH2 occur in slightly acidic, neutral, and slightly alkaline solu-tions at pH ≥ 5 close to pKa2, while the initial single-electron oxidationof nonprotonated 6-QNH2 occurs in highly alkaline solutions at pH≥ 10. It was computationally determined using the AM1 methodthat the free radical recombination reaction of [6-QNH2]•+, generatedin acidic solutions, leads to the prevalent formation of C5–C5 cou-pled 6-QNH2 dimers, while the free radical recombination reaction of[6-QNH]•, generated in alkaline solutions, leads to the prevalent for-mation of NH–C5 coupled 6-QNH2 dimers. The initial single-electronoxidations of [H-6-QNH3]2+, 6-QNH3

+, and/or 6-QNH2 in acidic andalkaline solutions, accompanied by the recombination of generatedfree radical species [6-QNH2]•+ or [6-QNH]• and followed by two-electron oxidation of prevalently formed C5–C5 or NH–C5 coupled6-QNH2 dimers, represent an overall two-electron transfer process per6-QNH2 molecule. This is the first possible 6-QNH2 electro-oxidationpath, which occurs at relatively low anodic potentials (Epa ≤ 1.2 Vvs. SCE). It was confirmed that it is a diffusion controlled process.The second possible 6-QNH2 electro-oxidation path, which representsalso an overall two-electron transfer process per 6-QNH2 molecule,comprises the initial two-electron oxidation of 6-QNH2 in a broadpH range at higher potentials in comparison with the correspondingsingle-electron oxidation. In accordance with semi-empirical quantumchemical predictions, the anodic CV peak which can be attributedto the initial two-electron oxidation was observed at higher anodicpotentials (Epa ≥ 1.2 V) in acidic solutions. This peak was not ob-served in alkaline solutions because of the supporting electrolyte’soxidation. The AM1 computations indicate that this electro-oxidationprocess is followed by the reaction of the formed [6-QNH]+ nitre-nium cation species with 6-QNH2 and the subsequent oxidation of

prevalent NH–C5 coupled 6-QNH2 dimer. Continuous CV experi-ments have clearly shown that the electro-oxidation of 6-QNH2 in abroad pH range is an irreversible process ultimately leading to the for-mation of non-conducting 6-QNH2 oligomers and/or polymer filmsat electrode surface. Because of the rapid oxidative oligomerization/polymerization of 6-QNH2 the isolation of its dimeric oxidation prod-ucts and unambiguous determination of their molecular structure stillpresents a challenge. From DPV analysis in supporting electrolyte,BR buffer/methanol, containing 20% of methanol (v/v) at pH 2.00,the 6-QNH2 limit of detection of 4.3 × 10−7 M was calculated. Itis shown that AdSDPV can be used for more sensitive quantitativedetermination.

Acknowledgment

This study was supported by the Ministry of Education and Scienceof the Republic of Serbia (projects OI172030 and OI172043).

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