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Voltammetry of Electroinactive Species UsingQuinone/Hydroquinone Redox: A Known Redox System Viewed ina New PerspectiveMohammad Rafiee, Davood Nematollahi*

Faculty of Chemistry, Bu-Ali Sina University, Hamadan 65174, Iran*e-mail: nemat@basu.ac.irReceived: January 22, 2007Accepted: March 27, 2007

AbstractElectrochemical behavior of p-benzoquinone (Q), hydroquinone (H2Q) and quinhydrone complex (QH) have beeninvestigated in aqueous unbuffered solutions. The results revealed that in unbuffered solution the half wave potentialof hydroquinone has a significant difference with quinone6s half wave potential. It shown that, added acid in anunbuffered solution of Q, give rise to a new reduction peak at a more positive potential than original reduction peakof Q. The half-peak potential of this new peak is dependent on acidity of added acid (pKa) and its height isproportional to the acid concentration. Also, added base in an unbuffered solution of H2Q, give rise to a newoxidation peak at a more negative potential than original oxidation peak of H2Q. The half-peak potential of this newpeak is dependent on basicity of added base (pKb) and its height is proportional to the base concentration. This papershows new perspectives for a known system in aqueous unbuffered solutions, by means of voltammetric responses thatcan be exploited for the electrochemical investigation of non electroactive species such as HPO2�

4 , HCO�3 or

CH3COOH.

Keywords: Quinhydrone, Unbuffered solution, Cyclic voltammetry, Non electroactive species

DOI: 10.1002/elan.200703864

1. Introduction

Since the introduction of quinhydrone electrode by Biil-mann [1], it has been used as the first and a powerful devicefor determination of pH. At constant temperatures andwell-buffered solutions its potential is a linear function ofpHwith a slope of 59 mV [2].Also poorly buffered solutionswere studied with this electrode and the pH values thusfound agreedwell with results obtained in differentways [3].Furthermore, because of its simplicity and sufficiently shortresponse time, this electrode was used in FIA and mini-aturized systems [4, 5]. Contrary to buffered solutions, inunbuffered solutions with pH values greater than 5.0 theelectrode gives erroneous values [6]. While p-benzoqui-none/hydroquinone system is perhaps to be the most wellstudied organic redox couple [2, 6 – 12], there are still someaspects of its redox chemistry that are not well known. Inparticular, while it is established that p-benzoquinone, Q,undergo a 2 e�, 2 Hþ reduction to give the hydroquinone,H2Q, in buffered aqueous solutions, little works have beendone on the electrochemistry of p-benzoquinone/hydro-quinone system in unbuffered aqueous solution [13 – 16].Therefore, in this work, electrochemical behavior of p-benzoquinone (Q), hydroquinone (H2Q) and quinhydronecomplex (QH) has been investigated in aqueous unbufferedsolutions. In otherwords, in thiswork, a known redox systemis viewed in a new perspective and shows that the potentialof the peaks are dependent on basicity or acidity of the

species that are present in solution. These results are usedfor electrochemical investigation of non electroactivespecies in aqueous unbuffered solutions.

2. Experimental

2.1. Chemicals and Solutions

p-Benzoquinone, hydroquinone and quinhydrone werereagent-grade materials and NaH2PO4, Na2HPO4 and KClwere of pro-analysis grade from E. Merck. These chemicalswere used without further purification. The stock solutionsof quinone, hydroquinone and quinhydrone prepared daily.Buffered solutions were prepared based on Kolthoff table[17] using NaH2PO4/Na2HPO4 solutions with total concen-tration of 0.15 M. Unbuffered solutions were adjusted withaddition of 0.01 M HCl or NaOH. The ionic strength of allexperimental solutions maintained at 0.15 M by adding KClsolution. Nitrogen gas with a purity of 99.999% was used toremove oxygen from solution during the experiments.

2.2. Electrodes and Electrochemical Instruments

Cyclic voltammetry was performed using an Autolab modelPGSTAT 20 potentiostat/galvanostat. The working elec-trode used in the voltammetry experiment was a glassy

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Electroanalysis 19, 2007, No. 13, 1382 – 1386 I 2007 WILEY-VCH Verlag GmbH&Co. KGaA, Weinheim

carbon disk (1.8 mm diameter), and Pt wire was used ascounter electrode. The working electrode potentials weremeasured versus SCE (all electrodes from AZAR elec-trode). pH was measured using a Metrohm pH meter 744with a combined glass electrode. RDE voltammetricmeasurements were made with a Metrohm analyticalrotator system equipped with an 8.0 mm electrode and1.8 mmglassy carbondisk, aPt counter electrode, and aSCEreference electrode. The experiments were performed atelectrode rotation rates of 500 – 3000 rpm. The temperaturewas maintained at 25� 0.1 8C by using a water-jacketedelectrochemical cell connected to a circulating water bath.

3. Results and Discussion

3.1. Electrochemical Study in Buffered Solutions

Figure 1 shows the cyclic voltammograms recorded for1 mM p-benzoquinone (Q), hydroquinone (H2Q) andquinhydrone (QH) in a pH 7.0 buffered solution. For allcompounds, the CVs show only one anodic peak (A0) in thepositive-going scan and its cathodic counterpart peak (C0) inthe negative-going scan. The anodic peak is due to theoxidation of H2Q to Q and the cathodic peak is due to thereduction of the produced Q to H2Q [2, 6 – 11]. The halfwave potentials (E1/2) for all cases shifted to the negativepotentials by increasing pH, with the slope of 59 mV/pH.This is expected because of the participation of two protons/two electrons in the oxidation/reduction reactions.

3.2. Electrochemical Studies in Unbuffered Solutions

Figure 2 shows cyclic voltammograms recorded for 1 mMQ,H2Q and QH in unbuffered aqueous solutions containing0.15 M KCl. The CV of H2Q shows an anodic peak (A1) in

the positive-going scan and its cathodic counterpart peak(C1) in the negative-going scan withE1/2¼ 0.27 V. The CVofQ shows a cathodic peak (C2) in the negative-going scan andits anodic counterpart peak (A2) in the positive-going scanwith E1/2¼�0.14 V. The CV of QH in both negative andpositive going scans shows two anodic (A1 and A2) andrelated cathodic peaks (C1 and C2). The difference betweenoxidation potential peaks,EA1 –EA2, (and also their counter-parts (EC1 –EC2)) is about 0.41 V.

In order to describe of this behavior, the followingequation is proposed for oxidation/reduction of H2Q/Q:

QþH2OÐþ2e� Generation of peak C2

�2e�Generation of peakA2HQ� þOH� (1)

According to the proposed equation, water acts as a protondonor in the reduction of Q (generation of C2 peak) (Eq. 1).But, when the hydronium ion (H3O

þ) is present at thesurface of the electrode, it acts as a stronger proton donor toproduce peak C1 (Eq. 2).

QþH3OþÐ

þ2e� Generation of peak C1

�2e� Generation of peakA1HQ� þH2O (2)

On the other hand, water acts as an acceptor of the proton,which is generated in the oxidation of HQ� (peak A1) [18].And when the hydroxide ion is present at the surface of theelectrode, it acts as a stronger proton acceptor to produceA2

peak (Eq. 1).According to the proposed equations, the difference in

anodic potential peaks (DEpA) is related to difference inproton acceptor properties (basicity) of H2O and OH�.Also, we think that DEpC¼EC1�EC2 is related to thedifference in acidity of hydronium ion (H3O

þ) and H2O inreduction of Q. These behaviors can be confirmed mathe-matically using the Equations 3 – 5.

Fig. 1. Cyclic voltammograms of 1 mM (a) Q, (b) H2Q and (c)HQ in buffered solutions of pH 7.0 at a glassy carbon electrode.Scan rate¼ 100 mV s�1 and t¼ 25� 1 8C.

Fig. 2. Cyclic voltammograms of 1 mM (a) Q, (b) H2Q and (c)QH in aqueous unbuffered KCl 0.15 M solution at a glassy carbonelectrode. Scan rate¼ 100 mV s�1 and t¼ 25� 18.

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Electroanalysis 19, 2007, No. 13, 1382 – 1386 www.electroanalysis.wiley-vch.de I 2007 WILEY-VCH Verlag GmbH&Co. KGaA, Weinheim

EA1¼E8’þ 0.0592/2 log[Q][H3Oþ]/[HQ�] (3)

EA2¼E8’þ 0.0592/2 log[Q]/[HQ�][OH�] (4)

DEpA¼EA1�EA2¼ 0.0592/2 (log[H3Oþ][OH�])

¼ 0.0592/2 (log Kw)¼ 0.41 V (5)

where, Kw is the autoprotolysis constant of water. As shownby Equation 5, an excellent agreement is observed for theexperimental and calculated values of DEpA.

3.3. The Effects of pH in Unbuffered Solutions

For more investigations, electrochemical behavior of H2Qand Q has been studied in various unbuffered pHs. Cyclicvoltammograms of 10.0 mM solution of H2Q in aqueousunbuffered solutions, at various pHs have been shown inFigure 3. In lower pHs (pH< 7.0) cyclic voltammogramsshow one anodic (A1) and the corresponding cathodic peak(C1), which corresponds to the transformation of H2Q (orHQ�) to Q and vice-versa, according to Equation 2. Theelectrochemical oxidation ofH2Q in higher pHswas studiedin some details. It is shown that, proportional to the increaseof pH, and in parallel with the decrease in height of cathodicC1 peak, a new anodic (A2) and its cathodic counterpartpeak (C2) appears and the height of them increases (Fig. 3,curves b and c). These new anodic (A2) and cathodic (C2)peaks are corresponding to the transformation of HQ� to Qand vice-versa according to Equation 1.

In this direction, cyclic voltammograms of 10.0 mMsolution of Q in aqueous unbuffered solutions at variousconcentrations of HCl have been shown in Figure 4. Asshownby curve a, in higher pHs (pH� 7.0) cyclic voltammo-gram shows one cathodic (C2) and the corresponding anodicpeak (A2), which corresponds to the transformation of Q to

HQ� and vice-versa according to Equation 1. The electro-chemical reduction of Q in lower pHs shows that, propor-tional to the increase in hydronium ion concentration(decrease of pH), and in parallel with the decrease in heightof anodic peak A2, new anodic and cathodic peaks (A1 andC1) appear and the height of them increase (Fig. 4, curvesb – d and inset). These peaks (C1 and A1) are correspondingto the transformation of Q to HQ� and vice-versa accordingto Equation 2.

In this direction, cyclic voltammograms of various con-centrations of Q in the presence of 0.5 mM HCl aqueoussolution are shown in Figure 5. As shown by curves a and b,when the concentration of Q is 1 mM or less, cyclicvoltammograms show only one cathodic (C1) and its anodic

Fig. 3. Cyclic voltammograms of 10.0 mM H2Q in aqueousunbuffered solution at various pHs, at a glassy carbon electrode.a) pH 7.0, b) pH 8.0 and c) pH 9.0. Supporting electrolyte: 0.15 MKCl. Scan rate¼ 50 mV s�1 and t¼ 25� 1 8C.

Fig. 4. Cyclic voltammograms of 10.0 mM Q in presence ofvarious concentration of HCl. a) in the absence of HCl, b) CHCl¼0.01, c) CHCl¼ 0.10, and d) CHCl¼ 1.0 mM at a glassy carbonelectrode. Supporting electrolyte: 0.15 M KCl. Scan rate¼ 50 mVs�1 and t¼ 25� 1 8C.

Fig. 5. Cyclic voltammograms of a) 0.5, b) 1.0, c) 2.0, and d)3.0 mM Q in presence of 0.5 mM aqueous HCl at a glassy carbonelectrode. Supporting electrolyte: 0.15 M KCl. Scan rate¼ 50 mVs�1 and t¼ 25� 1 8C.

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Electroanalysis 19, 2007, No. 13, 1382 – 1386 www.electroanalysis.wiley-vch.de I 2007 WILEY-VCH Verlag GmbH&Co. KGaA, Weinheim

counterpart peak (A1). But when the concentration of Q ismore than 1 mM, new anodic and cathodic peaks A2 and C2

appear in more negative potential and the height of themincrease with increasing the concentration of Q (Fig. 5,curves c and d).

For more details, electrochemical reduction of Q (C¼10.0 mM) has been studied in the presence of HCl (C¼1.0 mM) using rotating disk electrode (RDE). Fig. 6 I showsthese voltammograms at various rotation rates (w). Such asFigure 4, voltammograms show two cathodic plateaus C1

and C2 which are corresponding to reduction of Q to H2Qaccording to Equation 2 and Equation 1, respectively.Figure 6 II shows normalized voltammograms of Figure 6 I.The normalized voltammograms show the same currentsand potentials in all rotation rates, can be considered as acriterion for this fact that all currents are diffusion-controlled [19]. According to Equation 2 and by takinginto account the concentration of Q is excess, the current ofC1 is related to the concentration and the diffusioncoefficient of hydronium ion, whereas the current of C2 isrelated to the concentration and the diffusion coefficient ofQ (Eq. 1). So, in order to reconfirm Equations 1 and 2, weanalyzed voltammograms of Figure 6I based on Levichequation. The results indicate that the diffusion coefficientratio (DC1/DC2) is about 10, in agreement with diffusioncoefficient ratio of hydronium ion to Q (DHyd/DQ) [2, 20].

3.4. Electrochemical Study of H2Q in the Presence ofHPO2�

4

Cyclic voltammograms of 10.0 mM of H2Q in the presenceof various concentrations of HPO2�

4 in aqueous unbufferedsolutions are shown in Figure 7. As shown by curve a, in theabsence of HPO2�

4 , cyclic voltammogram shows one anodic(A1) and two cathodic peaks (C1 and C2). Also, as shown bycurves b – e, in the presence of HPO2�

4 , the height of C1

decreases and a new anodic and cathodic peak (A3 and C3)appears. The height of these new peaks is proportional toconcentration of HPO2�

4 and increase with increasing of it.For more details, electrochemical behavior of H2Q has

been investigated in the presence of HPO2�4 in aqueous

unbuffered solutions at pH 10.0. Figure 8, shows the cyclicand differential pulse voltammograms of H2Q (0.5 mM) inthe presence of HPO2�

4 (0.2 mM,) at pH 10.0. The cyclicvoltammogram in both negative and positive going scansshows three anodic (A1, A2 and A3) and their relatedcathodic peaks (C1, C2 and C3). Also, differential pulsevoltammogram in above conditions shows three welldefined anodic peaks A1, A2 and A3. The anodic peaks A1,A3 and A2 are related to participation of OH�, HPO2�

4 andH2O as proton acceptors in oxidation reaction of H2Q,

Fig. 6. I) Typical voltammograms of 10.0 mM Q in aqueous unbuffered solution in the presence of 1.0 mM HCl at a glassy carbonelectrode and at various rotation rates. II) Normalized voltammograms of I). Rotation rates from a) to d) are: 500, 1000, 2000, and3000 rpm, respectively. Supporting electrolyte: 0.15 M KCl. t¼ 25� 1 8C.

Fig. 7. Cyclic voltammograms of 10.0 mM H2Q in the presenceof a) 0.0, b) 1.0, c) 2.0, d) 3.0, e) 5.0 mM HPO2�

4 . f) as d) and g) ase) with different switching potential, at a glassy carbon electrode.Supporting electrolyte: 0.15 M KCl. Scan rate¼ 50 mV s�1 and t¼25� 18.

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Electroanalysis 19, 2007, No. 13, 1382 – 1386 www.electroanalysis.wiley-vch.de I 2007 WILEY-VCH Verlag GmbH&Co. KGaA, Weinheim

respectively, and the cathodic peaksC1,C3 andC2 are relatedto participation of H3O

þ, H2PO�4 and H2O as proton donors

in reduction of Q, respectively.The difference between potential of anodic peaks (EA1 –

EA3) is 0.21 V. In order to describe the new anodic andcathodic A3 and C3 peaks, the following equation isproposed for oxidation/reduction of H2Q/Q in the presenceof HPO2�

4 :

QþH2PO�4Ð

þ2e� Generation of peak C3

�2e�Generation of peakA3HQ� þHPO2�

4 (6)

According to the proposed equation, the difference be-tween potential of anodic peaks (DEpA¼EA3�EA1) isrelated to difference in basicity of H2O and HPO2�

4 . Thisbehavior can be confirmed mathematically using theEquations 7 – 9.

EA3¼E8’þ 0.0592/2 log([Q][H2PO�4 ]/[HQ�][HPO2�

4 ]) (7)

EA1¼E8’þ 0.0592/2 log([Q][H3Oþ]/[HQ�]) (8)

EA3�EA1¼ 0.0592/2 {log([H2PO�4 ]/[HPO2�

4 ])� log[H3Oþ]}

¼�0.0592/2 (log Ka)¼ 0.21 V (9)

where,Ka is the acid dissociation constant ofH2PO�4 (6.34�

10�8). As shown by Equation 9, an excellent agreement isobserved between the experimental (DEpA¼ 0.21 V.) andcalculated values. Also, all discussed peaks (A1, A2, A3, C1,C2 and C3) are diffusion peaks and their height changedlinearly with square root of scan rate (v1/2). Since, thecurrents of peaks A3 and C3 are proportional to concen-tration of HPO2�

4 in one hand, and to the potential of peaksA3 and C3 are related to basicity and acidity of HPO2�

4 and

H2PO�4 , respectively, on the other hand, in spite of the fact

that, HPO2�4 is not an electro-active species, we proposed

that, the voltammograms Figure 7, curves f and g, namedcyclic voltammograms of HPO2�

4 . Also we obtained similarresults for oxidation of H2Q in presence of HCO�

3 andreduction of Q in presence of weak acids such asCH3COOH.

4. Conclusions

The electrode processes for cathodic reduction of Q oranodic oxidation of H2Q in various unbuffered solutions arepresented inEquations 1, 2 and 6 and confirmedmathemati-cally using theEquations 3 – 5 and 7 – 9.Also, it is shown, thepotential of oxidation and reduction peaks dependent onbasicity or acidity of the species that are present in solution.The results of this work can be used for electrochemicalinvestigation of non electroactive species such as H2PO�

4 ,HPO2�

4 (see Fig. 7, inset), HCO�3 , CH3COOH and estima-

tion of acidity (or basicity) constant of unknown compounds(see Eq. 9) as well as determination of acid or base6sconcentration (see Fig. 4, inset)

5. References

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664.[3] H. Kahlert, J. R. Pcrksen, I. Isildak, M. Andac, M. Yolcu, J.

Behnert, F. Scholz, Electroanalysis 2005, 17, 1085.[4] A. Jang, J. H. Lee, P. R. Bhadri, S. A. Kumar, W. Timmons,

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[5] H. P. Cady, J. D. Ingle, J. Phys. Chem. 1936, 40, 837.[6] R. J. Best, J. Phys. Chem. 1930, 34, 1815.[7] M. Olive, J. Lammert, R. M. Livingston, J. Am. Chem. Soc.1932, 54, 910.

[8] A. Hemingway, Ind. Eng. Chem. Anal. Ed. 1935, 7, 203.[9] R. Rosenthal, A. E. Lorch, L. P. Hammett, J. Am. Chem. Soc.1937, 59, 1795.

[10] R. G. Bates, J. Electroanal. Chem. 1961, 2, 93.[11] C. Aquino-Binag, P. J. Pigram, R. N. Lamb, P. W. Alexander,

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1986, 51, 4760.[13] O. H. Muller, J. Am. Chem. Soc. 1940, 62, 2434.[14] K. Takamura, Y. Hayakawa, J. EIectroanal. Chem. 1971, 31,

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Quantitative Chemical Analysis, 4th ed., Macmillan, London1971.

[18] K. Izutsu, Electrochemistry in Nonaqueous Solutions, Wiley-VCH, Weinheim 2002, p. 103.

[19] A. J. Bard, L. R. Faulker, Electrochemical Methods, 2nd ed.,Wiley, New York 2001, p. 339.

[20] S. Guo, S. W. Feldberg, A. M. Bond, D. L. Callahan, P. J. S.Richardt, A. G. Wedd, J. Phys. Chem. B 2005, 109, 20641.

Fig. 8. Cyclic and differential pulse voltammograms of 0.5 mMH2Q in presence of 0.1 mM NaOH and 0.2 mM HPO2�

4 at a glassycarbon electrode. Supporting electrolyte: 0.15 M KCl. Scan rate¼50 mV s�1 and t¼ 25� 1 8C. Pulse amplitude: 100 mV, pulse width:50 ms and pulse interval: 100 ms.

1386 M. Rafiee, D. Nematollahi

Electroanalysis 19, 2007, No. 13, 1382 – 1386 www.electroanalysis.wiley-vch.de I 2007 WILEY-VCH Verlag GmbH&Co. KGaA, Weinheim