Two-Electron, Two-Proton Oxidation of Catechol: Kinetics and

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The Two Electron, Two Proton Oxidation of Catechol: Kinetics and Apparent Catalysis Qianqi Lin, Qian Li, Christopher Batchelor-McAuley, and Richard G Compton J. Phys. Chem. C, Just Accepted Manuscript • DOI: 10.1021/jp511414b • Publication Date (Web): 22 Dec 2014 Downloaded from http://pubs.acs.org on January 4, 2015

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The Journal of Physical Chemistry

The Two Electron, Two Proton Oxidation of Catechol: Kinetics and Apparent Catalysis Qianqi Lin, Qian Li, Christopher Batchelor-McAuley and Richard G. Compton* Department of Chemistry, Physical and Theoretical Chemistry Laboratory, Oxford University, South Parks Road, Oxford OX1 3QZ, United Kingdom

ACS Paragon Plus Environment

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ABSTRACT

The study of proton-coupled electron transfer reactions is of great current interest. In this work, the catechol redox process was studied voltammetrically in the pH range from 1.0 to 14.0 using a glassy carbon electrode. Analysis of the peak potentials and currents together with Tafel analysis allowed the inference of the likely transition states and electrode reaction mechanism. Modification of the glassy carbon electrode surface with sparse coverages of alumina particles was shown to lead to strong apparent catalysis of the catechol redox process at low pH. A possible mechanism for this is proposed.

KEYWORDS Proton-coupled electron transfer, scheme of squares, cyclic voltammetry, pH, electrode surface modification, alumina


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INTRODUCTION Catechol and catecholamine redox systems have been extensively investigated due to their importance in neurochemistry.1 To probe brain chemistry, diverse analytical methods have been developed to monitor neurotransmitters including spectrophotometry2, high-performance liquid chromatography3-4 and gas chromatography-mass spectrometry5. Catechol can be detected in vivo by electrochemical methods such as cyclic voltammetry based on the oxidation of catechol to o-benzoquinone at the electrode.6-8 The advantage of in vivo voltammetry is that it can trace many of the dynamic chemical changes associated with neurotransmitter release. Specifically it probes the timescale of a few milliseconds to seconds, giving information not accessible by any other technique.8 In fully buffered media the quinone/hydroquinone redox process involves changes in the protonation state of the molecule, resulting in the observation that potentiometric or amperometric equilibrium potentials vary with pH in a Nernstian manner.9 This behaviour can be demonstrated in potential-pH diagrams and used as a basis of an electrochemical pH sensor.10 This calibration-free pH sensor can effectively overcome the limitations of conventional glass electrodes relating to their fragility, alkali error and potential drift caused by dehydration of glass membrane.10-13 The pH dependence of catechol redox process can be ascribed to a two protontwo electron (2H+2e–) transfer, commonly known as a proton-coupled electron transfer (PCET) reaction. Quinone/hydroquinone based PCET agents play an important role in the metabolic oxidative phosphorylation of ADP to ATP.14 The PCET process can be understood using the ‘scheme of squares’ as originally introduced by Jacq15. Specifically the 2H+2e– nine-member ‘scheme of squares’ shown in


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Scheme 1 can be applied to interpret the electrochemical properties of the quinone/hydroquinone redox systems. This model is usually based on the assumptions that electron transfer is the rate determining step and that all protonations are at equilibrium. The reaction pathway depends heavily on the pKa values of the species within the model, as well as the pH of the local environment. The ‘scheme of squares’ was later further developed by Laviron16, assuming dimerization and disproportionation could be neglected, so that the system behaves as a simple reaction with two successive one-electron transfers, with apparent electrochemical rate constants and redox potentials. This approach was applied to the voltammetric analysis of p-benzoquinone reduction on a platinum electrode17 and the oxidation of substituted catechols on a carbon paste electrode18-19. Recently a simulation software DIGISIM (BASi, West Lafayette, IN, US) has been utilised to facilitate the ‘scheme of squares’ analysis on the reduction of flavin adenine dinucleotide20-21 and the reduction of anthraquinone-2,6-disulfonate and anthraquinone-2sulfonate22. The electrochemical detection of catechol has been previously examined on both metal17, 23-25

and carbon18-19,

26-28

electrodes. Of particular relevance to the work reported below are

studies of catechol redox processes on native and modified glassy carbon surfaces, with investigation of the heterogeneous electron transfer between solid electrodes and catechol.29-34 The present work explores for the first time the system of catechol oxidation on a glassy carbon electrode over the pH range from 1.0 to 14.0. The present paper additionally studies apparent catalysis using alumina modification of the electrode surface showing the role of adsorbed catechol on the latter, which assists the PCET between catechol and glassy carbon electrode. A possible application of this work would be to increase sensitivity of detection of catechol and catecholamines on microfluidic devices.35


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Scheme 1. The ‘scheme of squares’ for catechol oxidation. Each molecule is assigned abbreviation, Q, QH, QH2, S, SH, SH2, C, CH and H2C. ǂ Location of transition state for the pH range 7.0 < pH < 9.0 and 9.0 < pH < 12.8. The potentials relate to the heterogeneous electron transfers. The pKa values are associated with the proton transfers. *The literature values of pKa4~6 for catechol were measured using pulse radiolysis: pKa4 = 5,14, 36 pKa5 = 9.45 and pKa6 = 12.8.37 †pKa2 is not known for catechol but it is – 1.1 for duroquinone.38 pKa3 cannot be determined due to the instability of benzoquinone in concentrated acid.39 However, the behaviour of benzoquinone was reportedly examined by UV spectrophotometry and pKa3 was estimated to be – 6 or smaller.17 pKa1 is not known but is likely significantly more negative than pKa3. In this work we use pKa1 = – 10, pKa2 = – 1 and pKa3 = – 6 as assumptions in line with previous literature.17-19


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EXPERIMENTAL Chemical Reagents and Solutions. All chemicals were purchased from Sigma-Aldrich and were of analytical grade. They were used without any further purification. All solutions were prepared using deionised water of resistivity not less than 18.2 MΩ· cm at 298 K (Millipore, Billerica, MA, USA). Catechol solutions were prepared from catechol with the purity of 99.5+%. The structure of catechol (H2C) is shown in Scheme 1, along with the oxidation products, obenzoquinone (Q) and the semiquinone species (SH). Buffer solutions for different pH ranges were prepared using various recipes: 0.1 M hydrochloric acid for pH 1.0 ~ 2.5, 0.1 M citric acid/sodium citrate for pH 2.5 ~ 5.5, 0.1 M monosodium

phosphate/disodium

phosphate

for

pH

5.5

~

8.5,

0.1

M

sodium

carbonate/bicarbonate for pH 8.5 ~ 11.5, and either NaOH or KOH for pH 12.0 ~ 14.0. All solutions were made with 0.1 M KCl as the supporting electrolyte. pH buffers were freshly made, measured by using a Hannah pH213 pH meter, and saturated with nitrogen prior to addition of 0.5 mM catechol. Apparatus.

The voltammetric measurements were recorded using a µAutolabIII potentiostat

(Autolab, Netherlands) with a standard three electrode configuration. Either a glassy carbon macroelectrode (GC, diameter 3.0 mm, ALS distributed by BASi, Tokyo, Japan) or a carbon microelectrode (µ-C, radius calibrated as 4.98 µm, BASi, West Lafayette, IN, U.S., for work reported in Supporting Information) was used as the working electrode.

A platinum wire

(99.99% GoodFellow, Cambridge, U.K.) and a saturated calomel electrode (SCE, ALS distributed by BASi, Tokyo, Japan) were used as the counter and reference electrodes, respectively. The radius of the µ-C (rd = 4.98 µm) was calibrated by analysing the steady-state


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limiting current (Iss) for the one electron reduction of 1 mM hexaamineruthenium(III) chloride in aqueous solution with 0.1 M KCl as supporting electrolyte. A diffusion coefficient of 8.43 × 10-6 cm2· s-1 at 298 K was used40 to determine rd by the following equation:41 ‫ܫ‬௦௦ = 4݊‫ݎܥܦܨ‬ௗ

(1)

where n = 1 is the number of electrons transferred, F is the Faraday constant (96485 C· mol-1), D is the diffusion coefficient (cm2· s-1) and C is the bulk concentration of hexaamineruthenium(III) chloride (mol· cm-3). The GC working electrode was polished by using diamond sprays of decreasing particle sizes (3.0 ~ 0.1µm, Kemet, Kent, U.K.) for an unmodified surface. The µ-C working electrode was polished by using alumina slurry of decreasing particle sizes (1.0 ~ 0.05 µm, Buehler, IL, U.S.). The electrode preparation was completed with 1 minute sonication (FB15050, Fisher Scientific, 50/60 Hz, 80 W, Germany) in deionised water. To obtain a modified surface, the GC working electrode was prepolished by using alumina slurry of decreasing particle sizes. After cleaning by sonication, the electrode was then polished by using 1.0 µm alumina and rinsed with water. All working electrodes were dried by nitrogen before each measurement. All electrochemical experiments were performed in a three-necked flask with rubber stoppers within a Faraday cage. The temperature was controlled at (25 ± 0.2) oC by using a water bath. Experimental Procedure.

20 mL of each solution was transferred into a three-necked flask

and bubbled with nitrogen to remove the dissolved oxygen. Cyclic voltammetry was recorded at variable scan rates ranging from 16 to 400 mV· s–1, with potential windows adjusted at different pH.


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RESULTS AND DISCUSSION This work first investigates the electrochemical response of catechol in pH range from 1.0 to 14.0 with analysis of the peak potentials and mid-point potentials. It then proceeds to analyse the cyclic voltammetry using peak potentials, currents and Tafel analysis to find the likely reaction pathways and transition states for catechol oxidation at a glassy carbon electrode. Finally, and most importantly, the apparent electrocatalysis of the catechol oxidation using alumina modified glassy carbon electrodes is reported and discussed. Cyclic Voltammetry: Variation with pH. The electrochemical response of catechol was examined over the pH range from 1.0 to 14.0, as shown in Figure 1a and 1b. Cyclic voltammetry was run in an oxidative direction from – 0.4 to + 1.0 V (vs. SCE) and reversed to – 0.4 V using variable scan rates (ν) from 16 to 400 mV· s–1. As can be seen from Figure 1a and 1b, the voltammogram shifts to more negative potentials as the pH increases. The voltammetry is pH dependent since the removal of electrons from catechol induces the loss of protons. Such a deprotonation process is more facile at higher pH, resulting in a thermodynamically more feasible electron transfer occurring at more negative potential. This pH dependency will be analysed more quantitatively in the following discussion, leading to an inference of the likely reaction pathways and transition state for catechol oxidation at a glassy carbon electrode particularly in the light of Scheme 1. Figure 1a and 1b also shows that the reductive (back) peak is gradually diminished as the pH increases. The diminished reductive peak potentially indicates a mechanistic transition from a chemically-reversible redox process to a chemically-irreversible oxidation only. This likely occurs since after catechol is oxidised to o-benzoquinone, it is thought to undergo a nucleophilic attack by the hydroxyl ions through a 1,4-Michael addition reaction to


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give 1,2,4-trihydroxybenzene in relatively alkaline solutions.42-44 This hydroxylation consumes o-benzoquinone and hence the overall reaction becomes chemically irreversible.

25 20

(a)

Current / µA

15 10 5 0 -5

pH increases from 1.0 to 14.0

-10 -15 -20 -0.6 -0.4 -0.2

0.0

0.2

0.4

0.6

0.8

pH 1.0 pH 2.0 pH 2.9 pH 4.0 pH 5.0 pH 6.0 pH 7.0 pH 8.0 pH 9.0 pH 10.8 pH 13.0 pH 13.8 pH 14.0

1.0

Potential / V vs. SCE

Current / µA

10

(b)

5

0

pH increases from 1.0 to 14.0 -5 -0.6 -0.4 -0.2

0.0

0.2

0.4

0.6

0.8


1.0


0.7 0.6


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(c) slope1 = -0.059 V

0.5 0.4 0.3

slope ≈ -0.092 V

0.2 0.1 0.0 -0.1 -0.2

slope ≈ 0 V

slope2 = -0.040 V

pKa5 = 9.45 pKa6 = 12.8 0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15

pH


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Figure 1. Cyclic voltammograms of 0.5 mM catechol in pH buffered solutions supported with 0.1 M KCl recorded on an unmodified glassy carbon electrode over the pH range of 1.0 ~ 14.0 at a scan rate of (a) 144 mV· s–1 and (b) 16 mV· s–1, with voltammogram at each pH labelled. (c) A plot of oxidative peak potential (▲), reductive peak potential (▼) and mid-point potential (■) against pH in the pH range 1.0 ~ 14.0, measured on an unmodified glassy carbon electrode. The gradient of the best-fit line is – 0.059 V at pH 1.0 ~ 9.45 (slope1), which decreases to ca. – 0.040 V at pH 9.45 ~ 12.8 (slope2). Scan rate = 16 mV· s–1. A plot of the oxidative peak potential (Ep,ox, ▲), reductive peak potential (Ep,red, ▼), and mid-point potential (Emid, ■) against pH for the voltammograms at a scan rate of 16 mV· s–1 is depicted in Figure 1c. Such a scan rate is chosen to minimise any ohmic distortion. At pH values above 9.0, only Ep,ox is reported. The absence of Emid and Ep,red arises from the chemicallyirreversible behaviour mentioned above, leading to the absence of a back peak. The lines of best fit of Emid and Ep,ox are illustrated in Figure 1c, with gradient changes near pH 9.25 and pH 13.0. These breaks agree well with the previously reported first and second acid dissociation constants for catechol being 9.45 and 12.8 at 25 oC.37 Furthermore, the resulting gradients, within the three regions separated by the two acid dissociation constants, change from ca. 59 mV (Emid), to ca. 40 mV (Ep,ox), to ca. 0 mV (Ep,ox) as shown explicitly in the figure. For a simple electrochemically reversible voltammograms the midpoint potential obeys the following equation:45 ‫ܧ‬୫୧ୢ ~ܿ‫ݐݏ݊݋‬. −2.303

௠ோ் ௡ி

pH

(2)

where R is the gas constant (8.314 J· K-1· mol-1), T is the temperature (K) , F is the Faraday constant (96485 C· mol-1), m and n are the number of protons and electrons involved in the ௠

oxidation respectively. From Eqn. (2), it can be seen that Emid varies by the amount 0.059 ௡ V per pH unit. At pH values in the range 1.0 to 9.0, the experimental measurement resulted in a 59 mV/pH shift, consistent with a 2H+2e– transfer in the electrochemical oxidation of catechol. As


10

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shown in Scheme 1, the oxidation starts with H2C when the pH range is below 9.45 and finishes with Q corresponding to the observed slope of c.a. 59 mV/pH. When the pH is above 9.0, the oxidative peaks, Ep,ox, demonstrate possibly conflated chemical and electrochemical irreversible behaviour. For an electrochemically irreversible process,45 డா౦,౥౮ డ୮ୌ

௠ோ்

~ − 2.303 ൫௡ᇲ ାఉ

(3)

೙ᇲ శభ ൯ி

௠

so that theoretically Ep,ox shifts by an amount of 0.059 ൫௡ᇲ ାఉ

೙ᇲ శభ ൯

V per pH unit, where n’ is the

number of electrons transferred before the rate-determining step and βn’+1 is the transfer coefficient of the rate-determining step. In the pH region 9.0 < pH < 12.8, the experimental Ep,ox shift is 40 mV/pH, consistent with m and (n’ + βn’+1) being 1 and 1.49 ± 0.03 respectively. The value for (n’ + βn’+1) is close to the ones extracted from Tafel plot (see Figure 2) with an average value of 1.40 ± 0.09. The values for m and (n’ + βn’+1) imply a 1H+2e– transfer, a monodeprotonation process accompanied with the loss of two electrons where the second electron transfer is the rate-determining step. Accordingly, the oxidation starts with CH as indicated in Scheme 1.


11


-12

-12

10

-13

8

-14

(n' + βn'+1)= 1.48

-16 -0.08

-0.06

-0.04

E

4

-0.02

0.00

2

-15

Current / µA

6

-14

8

-15

Current / µA

-13

(b)

ln Iox

(a)

ln Iox

10

0

6

-16 -0.10

-0.08

-0.06

E

-0.04

-0.02

4 2

-0.2

0.0

0.2

0.4

-2 -0.4

0.6

-0.2


10

-13

0.6

-12

ln Iox

-13

-16 -0.12

-0.10

-0.08

E

-0.06

-0.04

4 2

Current / µA

(n' + βn'+1) = 1.41

-15

6

0.4

(d)

8

-14

8

0.2

ln Iox

(c)

0.0


-12

10

(n' + βn'+1) = 1.45

0

-2 -0.4

Current / µA

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

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-14 -15

6

(n' + βn'+1) = 1.27

-16 -0.14

-0.12

4

-0.10

E

-0.08

-0.06

2 0

0 -2 -0.4

-0.2

0.0

0.2

0.4

0.6

-2 -0.4

-0.2


0.0

0.2

0.4

0.6


Figure 2. Cyclic voltammograms of 0.5 mM catechol in buffer solutions supported with 0.1 M KCl recorded on an unmodified glassy carbon electrode. (a) pH 9.5, (b) pH 10.0, (c) pH 10.5 and (d) pH 10.8. Scan rate = 16 mV· s–1. The forward scan region highlighted in orange was selected as the Tafel analysis region. Insets are the corresponding Tafel plots, giving the values of transfer coefficient, (n’ + βn’+1) = 1.40 ± 0.09. The voltammetric data, both Tafel and peak potential against pH plot, are all consistent with a transition state (ǂ) located as shown in the ‘scheme of squares’ (Scheme 1) for the pH range between 9.0 and 12.8. The threshold pH values reflect the two pKa values (pKa5 and pKa6) as shown in Scheme 1 and again in Figure 1c.


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We next consider the pH range between 7.0 and 9.0. Considering the peak potential against pH plot in Figure 1c, it can be seen that the oxidative peak, Ep,ox, varies with pH according to ∂Ep,ox / ∂pH ≈ – 0.092 V whilst for the reductive peak, ∂Ep,red / ∂pH ≈ 0 V. The oxidative Tafel slope in this region was consistent with a value of (n’ + βn’+1) as 1.28. These data are again consistent with the transition state for the reaction being located as shown in Scheme 1, corresponding to the second electron transfer. Last we consider the pH range between 1.0 and 7.0 where there is a large separation (over 200 mV) between the oxidative and reductive peak potentials with both of these peak potentials apparently showing a Nernstian regime of 59 mV per decade. Tafel analysis showed an apparent transfer coefficient of (n’ + βn’+1) ≈ 1 suggesting electrochemically reversible behaviour. It is clear that a change of mechanism has occurred relating to the behaviour at higher pH and consideration of the pKa values given in the ‘scheme of squares’ (Scheme 1) suggests that this is triggered by the protonated form, SH, of the semiquinone radical becoming stabilised relative to S (pKa4 = 5). The voltammetric behaviour hints at a behaviour where the H2C / SH potential occurs at more positive values than the SH / Q potential – a phenomena commonly referred to as ‘potential inversion’46-48. Thus both the oxidative and reductive voltammetric peaks behave in an apparently electrochemically reversible fashion but with a wider potential separation than expected for a simple two electron transfer. Given these observations, the oxidative mechanism (for pH below 5.0) is likely to be: H2C – e– – H+ ⇄

SH

(4)

followed by disproportionation 2 SH →

H2C + Q

(5)


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Conversely for reduction we have Q + e– + H+

⇄

2 SH →

SH

(6)

H2C + Q

(5)

These quantitative inferences were supported by DIGISIM modelling as reported in the Supporting Information. Apparent Electrocatalysis Using Alumina Modified Glassy Carbon Electrodes.

Next

the effect of surface modification of the glassy carbon (GC) with alumina was studied in the light of literature observations33, 49 and suggestions of possible ‘catalysis’. To obtain an alumina modified suface, the GC was prepolished by using alumina slurry of decreasing particle size (1.0, 0.3 and 0.05 µm). After cleaning by sonication, the electrode was polished by using 1.0 µm alumina, rinsed with water and dried under nitrogen before each measurement. Figure 3 shows the cyclic voltammetry of 0.5 mM catechol in pH 1.0 solution measured on an unmodified GC (dashed) in comparison with that on a modified GC (line) at a scan rate of 144 mV· s–1. It can be seen that ∆Ep is reduced from ca. 300 mV to ca. 40 mV, whereas the Emid stays at ca. 0.50 V vs. SCE. Thus the overall thermodynamics of the CH2 to Q process are unchanged, but the reaction mechanism has clearly been altered significantly, possibly usefully in the context of pH sensing.


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50 40 30

Current/ µA

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


20 10 0 -10 -20 -30 -0.4

-0.2

0.0

0.2

0.4

0.6

0.8

1.0


Figure 3. Comparison of cyclic voltammograms measured on an unmodified (dashed) and alumina modified (line) glassy carbon electrode, of 0.5 mM catechol in pH 1.0 buffered solution supported with 0.1 M KCl and saturated with nitrogen. Scan rate = 144 mV· s–1. Similar observations were found at higher pH, as depicted in Figure 4a, 4b and 4c. The electrochemical response of catechol over the pH range from 1.0 to 14.0 was recorded by using an alumina modified GC, as shown in Figure 4a and 4b. The corresponding Ep,ox (▲), Ep,red (▼) and Emid (■) were plotted against pH in Figure 4c. The small ∆Ep seen at pH values below pH 9.0 is all suggestive of an altered mechanism on the modified GC (Figure 4c). As noted earlier, above pH 9.0 the system will be insensitive to further redox processes after the chemicallyirreversible step, and hence the reductive back peak is missing from the voltammogram. The inset of Figure 4c depicts Emid below pH 9.0 and Ep,ox above pH 9.0 measured on the modified GC (red symbols) overlaid with the corresponding measurements on the unmodified GC (green symbols). This comparison further supports the finding that catechol oxidation is unchanged thermodynamically when studied at an alumina modified GC.


15

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Page 16 of 24


pH increases from 1.0 to 14.0 -0.2

0.0

0.2

0.4

0.6

0.8

1.0

Potential / V vs. SCE 15

(b)


Current / µA

10

5

0

pH increases from 1.0 to 14.0

-5

-10 -0.6

-0.4

-0.2

0.0

0.2

0.4

0.6

0.8

1.0

0.7 0.6 0.5 0.4

(c) slope1 = -0.059 V

0.6

E



1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Current / µA


0.2

0.4

0.0

0.3 -0.2

0.2

0

2

4

6

8

10

12

14

pH

0.1 0.0 -0.1

pKa5 = 9.45

slope2 = -0.040 V pKa6 = 12.8

-0.2

0 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15

pH

Figure 4. Cyclic voltammograms of 0.5 mM catechol in pH buffered solutions supported with 0.1 M KCl recorded on an alumina modified glassy carbon electrode over the pH range 1.0 ~ 14.0 at a scan rate of (a) 144 mV· s–1 and (b) 16 mV· s–1, with voltammogram at each pH labelled. (c) A plot of oxidative peak potential (▲), reductive peak potential (▼) and mid-point potential (■, red symbols) against pH in the pH range 1.0 ~ 14.0, measured on an alumina modified glassy carbon electrode. Inset depicts the comparison with the data measured on an unmodified glassy carbon electrode (green symbols). Scan rate = 16 mV· s–1. A ‘transfer experiment’ was conducted to further examine the adsorption of catechol to the alumina on the electrode surface. 2g of 1.0 µm alumina was stirred in 50 mL of 2 mM catechol pH 3.0 solution supported with 0.1 M KCl and saturated with nitrogen for 3 hr, allowing


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catechol to be fully adsorbed by the alumina. The milky mixture was then left still for 3 min, which allows the alumina to settle under gravity. This catechol-alumina slurry was transferred to a polishing pad for GC surface modification. After rinsing with water and drying with nitrogen, the electrode was submerged into a pH 3.0 solution supported with 0.1 M KCl and saturated with nitrogen. Cyclic voltammetry was run from – 0.3 V to 0.9 V and reversed to – 0.3 V for two consecutive scans at a scan rate of 144 mV· s–1, as shown in Figure 5 (black line: first, red line: second scan). For comparison, a ‘blank’ voltammogram was recorded in the same solution on a clean alumina modified GC (Figure 5, black dashed: first, red dashed: second scan). It is apparent that the oxidation at ca. 0.42 V and the reduction at ca. 0.35 V are the results of catechol redox process, which demonstrates the strong adsorption of catechol to the alumina. Moreover, Ep,ox and Ep,red in both scans remain unshifted, implying that catechol is substantially retained on the electrode surface even after electrolysis. The slightly reduced Ip,ox and Ip,red in the second scan is due to the background current drop (see dashed lines). 8 6

Current / µA

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


4 2 0 -2 -4 -0.4

-0.2

0.0

0.2

0.4

0.6

0.8

1.0


Figure 5. Comparison of cyclic voltammograms measured on a catechol-alumina modified glassy carbon electrode (line) and on a clean alumina modified glassy carbon electrode (dashed) of pH 3.0 buffered solution supported with 0.1 M KCl and saturated with nigtrogen. Black line


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and dashed correspond to the first scan, while red ones to the second scan. Scan rate = 144 mV· s–1. It is interesting to speculate on the causes of the apparent catalysis of the H2C / Q process by adsorbed alumina. The changed behaviour of the adsorbed catechol is possibly due to the alumina facilitating the two electron process: H2C – 2e– – 2H+

⇄

Q

(7)

on the surface of the solid. The presence of proton accepting oxide groups on the surface of the alumina may allow the process to proceed rapidly possibly even without the formation of a semiquinone species if the two electron transfer is concerted with proton transfer to the solid, as shown in Scheme 2. The scheme shows the possibility of two electrons and two protons being transferred to and from catechol adsorbed on the alumina. Such a reaction needs to occur at or close to the triple boundary formed by the electrode, the solution and the alumina particles. To ensure transfer of electrons into the electrode since the alumina is not conductive, the behaviour of the adsorbed catechol on the alumina and its altered voltammetry compared to solution phase catechol at the same electrode suggests that the facile two electron oxidation of the adsorbed species mediates the electron transfer to and from the solution phase species: C (ads) – 2e–

⇄

Q (ads) + C (soln)

⇄

Q (ads) C (ads) + Q (soln)

(8) (9)

so that the forward (oxidative) and back (reductive) reactions are both catalyzed.


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Scheme 2. Schematic diagram of the three-phase boundary between alumina oxide, catechol and glassy carbon electrode surface. Note the catechol is adsorbed on the alumina, not on the electrode, and reactions would be initiated at the triple boundary formed between the alumina particles, the solution and the electrode.


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CONCLUSIONS This work has investigated the 2e– 2H+ oxidation of catechol from two aspects. Fisrt, the electrochemical redox of catechol was studied voltammetrically in the pH range from 1.0 to 14.0 using a glassy carbon electrode. Through the study of potentials and currents together with Tafel analysis, the likely transition state and redox mechanism have been established in the ‘scheme of squares’ model. Second, modification of the glassy carbon electrode surface with sparse coverages of alumina particles was reported. Strong apparent catalysis of catechol redox process has been demonstrated at low pH, which was possibly due to the presence of proton accepting oxide groups on the surface of the alumina facilitating the proton-coupled electron transfer at the triple boundary formed by the electrode, the solution and the solid.

ASSOCIATED CONTENT Supporting Information Further information relating to the DIGISIM modelling of catechol redox mechanism can be found in the Supporting Information. This material is available free of charge via the Internet at http://pubs.acs.org.

AUTHOR INFORMATION Corresponding Author *Email: [email protected]. Phone: +44 (0) 1865 275957. Fax: +44 (0) 1865 275410


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REFERENCES (1) Anden, N. E.; Leander, S. Effects of 4-Aminopyridine on the Turnover of Mono-Amines in the Central Nervous-System of the Rat. J. Neural. Transm. 1979, 44, 1-12. (2) Figueiredo, E. C.; Tarley, C. R. T.; Kubota, L. T.; Rath, S.; Arruda, M. A. Z. On-Line Molecularly Imprinted Solid Phase Extraction for the Selective Spectrophotometric Determination of Catechol. Michrochem. J. 2007, 85, 290-296. (3) Lee, B. L.; Ong, H. Y.; Shi, C. Y.; Ong, C. N. Simultaneous Determination of Hydroquinone, Catechol and Phenol in Urine Using High-Performance Liquid Chromatography with Fluorimetric Detection. J. Chromatogr. B: Biomed. Sci. App. 1993, 619, 259-266. (4) Cui, H.; He, C.; Zhao, G. Determination of Polyphenols by High-Performance Liquid Chromatography with Inhibited Chemiluminescence Detection. J. Chromatogr. A 1999, 855, 171-179. (5) Lourenço, E. L. B.; Ferreira, A.; Pinto, E.; Yonamine, M.; Farsky, S. H. P. On-Fiber Derivatization of Spme Extracts of Phenol, Hydroquinone and Catechol with Gc-Ms Detection. Chroma 2006, 63, 175-179. (6) Boulton, A. A.; Baker, G. B.; Adams, R. N., Voltammetric Methods on Brain Systems; Humana Press: Totowa NJ, 1995. (7) Ewing, A. G.; Strein, T. G.; Lau, Y. Y. Analytical Chemistry in Microenvironments: Single Nerve Cells. Accounts Chem. Res. 1992, 25, 440-447. (8) Adams, R. N. In Vivo Electrochemical Measurements in the Cns. Prog. Neurobiol. 1990, 35, 297-311. (9) Kolthoff, I. M.; Lingane, J. J., Polarography, 2nd ed.; Interscience: New York, 1952; Vol. I, Chapters XIV and XV; II, Chapter XL. (10) Lu, M.; Compton, R. G. Voltammetric pH Sensor Based on an Edge Plane Pyrolytic Graphite Electrode. Analyst 2014, 139, 2397-2403. (11) Kristensen, H. B.; Salomon, A.; Kokholm, G. International pH Scales and Certification of Ph. Anal. Chem. 1991, 63, 885A-891A. (12) Skoog, D. A.; West, D. M.; Holler, F. J.; Crouch, S. R., Fundamentals of Electroanalytical Chemistry, 8th ed., 2004. (13) Xiong, L.; Batchelor-McAuley, C.; Compton, R. G. Calibrationless pH Sensors Based on Nitrosophenyl and Ferrocenyl Co-Modified Screen Printed Electrodes. Sensor Actuat. B: Chem. 2011, 159, 251-255. (14) Trumpower, B. L., Function of Quinones in Energy Conserving Systems; Academic Press: New York, 1982. (15) Jacq, J. Schema Carre: Etablissement Et Discussion De L'equation Generale De La Courbe Intensite-Potentiel En Regime Stationnaire Et Diffusion Convective. J. Electroanal. Chem. Interfacial Electrochem. 1971, 29, 149-180. (16) Laviron, E. Electrochemical Reactions with Protonations at Equilibrium: Part Viii. The 2e-, 2H+ Reaction (Nine-Member Square Scheme) for a Surface or for a Heterogeneous Reaction in the Absence of Disproportionation and Dimerization Reactions. J. Electroanal. Chem. Interfacial Electrochem. 1983, 146, 15-36. (17) Laviron, E. Electrochemical Reactions with Protonations at Equilibrium: Part X. The Kinetics of the p-Benzoquinone/Hydroquinone Couple on a Platinum Electrode. J. Electroanal. Chem. Interfacial Electrochem. 1984, 164, 213-227.


21


1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60

Page 22 of 24

(18) Deakin, M. R.; Wightman, R. M. The Kinetics of Some Substituted Catechol/o-Quinone Couples at a Carbon Paste Electrode. J. Electroanal. Chem. Interfacial Electrochem. 1986, 206, 167-177. (19) Deakin, M. R.; Kovach, P. M.; Stutts, K. J.; Wightman, R. M. Heterogeneous Mechanisms of the Oxidation of Catechols and Ascorbic Acid at Carbon Electrodes. Anal. Chem. 1986, 58, 1474-1480. (20) Smith, E. T.; Davis, C. A.; Barber, M. J. Voltammetric Simulations of Multiple Electron Transfer/Proton Transfer Coupled Reactions: Flavin Adenine Dinucleotide as a Model System. Anal. Biochem. 2003, 323, 114-121. (21) Cable, M.; Smith, E. T. Identifying the n=2 Reaction Mechanism of Fad through Voltammetric Simulations. Anal. Chim. Acta 2005, 537, 299-306. (22) Batchelor-McAuley, C.; Li, Q.; Dapin, S. M.; Compton, R. G. Voltammetric Characterization of DNA Intercalators across the Full Ph Range: Anthraquinone-2,6-Disulfonate and Anthraquinone-2-Sulfonate. J. Phys. Chem. B 2010, 114, 4094-4100. (23) Mu, S. Catechol Sensor Using Poly(Aniline-Co-O-Aminophenol) as an Electron Transfer Mediator. Biosens. Bioelecton. 2006, 21, 1237-1243. (24) Ghanem, M. A. Electrocatalytic Activity and Simultaneous Determination of Catechol and Hydroquinone at Mesoporous Platinum Electrode. Electrochem. Commun. 2007, 9, 2501-2506. (25) Su, L.; Mao, L. Gold Nanoparticle/Alkanedithiol Conductive Films Self-Assembled onto Gold Electrode: Electrochemistry and Electroanalytical Application for Voltammetric Determination of Trace Amount of Catechol. Talanta 2006, 70, 68-74. (26) Nasr, B.; Abdellatif, G.; Cañizares, P.; Sáez, C.; Lobato, J.; Rodrigo, M. A. Electrochemical Oxidation of Hydroquinone, Resorcinol, and Catechol on Boron-Doped Diamond Anodes. Environ. Sci. Technol. 2005, 39, 7234-7239. (27) Wang, S. G.; Li, Y. Q.; Zhao, X. J.; Wang, J. H.; Han, J. J.; Wang, T. Electrochemical Detection of Catechol at Integrated Carbon Nanotubes Electrodes. Diam. Relat. Mater. 2007, 16, 248-252. (28) Mersal, G. A. Electrochemical Sensor for Voltammetric Determination of Catechol Based on Screen Printed Graphite Electrode. Int. J. Eletrochem. Sc. 2009, 4, 1167-1177. (29) Cabaniss, G. E.; Diamantis, A. A.; Murphy, W. R.; Linton, R. W.; Meyer, T. J. Electrocatalysis of Proton-Coupled Electron-Transfer Reactions at Glassy Carbon Electrodes. J. Am.Chem. Soc. 1985, 107, 1845-1853. (30) Poon, M.; McCreery, R. L. In Situ Laser Activation of Glassy Carbon Electrodes. Anal. Chem. 1986, 58, 2745-2750. (31) DuVall, S. H.; McCreery, R. L. Control of Catechol and Hydroquinone Electron-Transfer Kinetics on Native and Modified Glassy Carbon Electrodes. Anal. Chem. 1999, 71, 4594-4602. (32) Allred, C. D.; McCreery, R. L. Adsorption of Catechols on Fractured Glassy Carbon Electrode Surfaces. Anal. Chem. 1992, 64, 444-448. (33) Zak, J.; Kuwana, T. Electrooxidative Catalysis Using Dispersed Alumina on Glassy Carbon Surfaces. J. Am.Chem. Soc. 1982, 104, 5514-5515. (34) DuVall, S. H.; McCreery, R. L. Self-Catalysis by Catechols and Quinones During Heterogeneous Electron Transfer at Carbon Electrodes. J. Am.Chem. Soc. 2000, 122, 6759-6764. (35) Wang, J.; Pumera, M.; Chatrathi, M. P.; Escarpa, A.; Konrad, R.; Griebel, A.; Dörner, W.; Löwe, H. Towards Disposable Lab-on-a-Chip: Poly(Methylmethacrylate) Microchip Electrophoresis Device with Electrochemical Detection. ELECTROPHORESIS 2002, 23, 596601.


22

Page 23 of 24

1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 25 26 27 28 29 30 31 32 33 34 35 36 37 38 39 40 41 42 43 44 45 46 47 48 49 50 51 52 53 54 55 56 57 58 59 60


(36) Steenken, S.; O'Neill, P. Oxidative Demethoxylation of Methoxylated Phenols and Hydroxybenzoic Acids by the Hydroxyl Radical. An in Situ Electron Spin Resonance, Conductometric Pulse Radiolysis and Product Analysis Study. J. Phys. Chem. 1977, 81, 505508. (37) Steenken, S.; Neta, P. Electron Transfer Rates and Equilibriums between Substituted Phenoxide Ions and Phenoxyl Radicals. J. Phys. Chem. 1979, 83, 1134-1137. (38) Land, E. J.; Porter, G. Acidic Semiquinones. P. Chem. Soc. 1960, 84-85. (39) Handa, T. On the Color-Reactions of the Condensed Polycyclic Aromatic Hydrocarbons and Their Related Quinones in the Solutions of the Concentrated Sulfuric Acid. B. Chem. Soc. Jpn. 1955, 28, 483-489. (40) Wang, Y.; Limon-Petersen, J. G.; Compton, R. G. Measurement of the Diffusion Coefficients of [Ru(NH3)6]3+ and [Ru(NH3)6]2+ in Aqueous Solution Using Microelectrode Double Potential Step Chronoamperometry. J. Electroanal. Chem. 2011, 652, 13-17. (41) Saito, Y. A Theoretical Study on the Diffusion Current at the Stationary Electrodes of Circular and Narrow Band Types. Rev. Polarography, 15, 178-187. (42) Papouchado, L.; Petrie, G.; Adams, R. N. Anodic Oxidation Pathways of Phenolic Compounds: Part I. Anodic Hydroxylation Reactions. J. Electroanal. Chem. Interfacial Electrochem. 1972, 38, 389-395. (43) Kiani, A.; Raoof, J.-B.; Nematollahi, D.; Ojani, R. Electrochemical Study of Catechol in the Presence of Dibuthylamine and Diethylamine in Aqueous Media: Part 1. Electrochemical Investigation. Electroanal. 2005, 17, 1755-1760. (44) Ojani, R.; Raoof, J. B.; Hosseinzadeh, R.; Alinezhad, A. Electrochemical Oxidation of Catechol in the Presence of Dimethyl Chloromalonate and Its Digital Simulation. Asia J. Chem. 2008, 20, 5863-5872. (45) Compton, R. G.; Banks, C. E., Understanding Voltammetry. 2nd ed.; Imperial College Press: 2011. (46) Evans, D. H.; Hu, K. Inverted Potentials in Two-Electron Processes in Organic Electrochemistry. J. Chem. Soc., Faraday Trans. 1996, 92, 3983-3990. (47) Evans, D. H. The Kinetic Burden of Potential Inversion in Two-Electron Electrochemical Reactions. Acta. Chem. Scand. 1998, 52, 194-197. (48) Batchelor-McAuley, C.; Compton, R. G. Voltammetry of Multi-Electron Electrode Processes of Organic Species. J. Electroanal. Chem. 2012, 669, 73-81. (49) Zak, J.; Kuwana, T. Chemically Modified Electrodes and Electrocatalysis. J. Electroanal. Chem. Interfacial Electrochem. 1983, 150, 645-664.


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Table of Contents Image


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Two-Electron, Two-Proton Oxidation of Catechol: Kinetics and

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