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A Mechanistic Study of the Validity of Using Hydroxyl Radical Probes to Characterize Electrochemical Advanced Oxidation Processes Yin Jing, and Brian P. Chaplin Environ. Sci. Technol., Just Accepted Manuscript • DOI: 10.1021/acs.est.6b05513 • Publication Date (Web): 10 Jan 2017 Downloaded from http://pubs.acs.org on January 12, 2017

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Environmental Science & Technology

2 3 4 5

A Mechanistic Study of the Validity of Using Hydroxyl Radical

6

Probes to Characterize Electrochemical Advanced Oxidation

7

Processes

8 9

Yin Jing, Brian P. Chaplin*

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Department of Chemical Engineering, University of Illinois at Chicago, 810 South Clinton Street,

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Chicago, Illinois 60607, United States

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ACS Paragon Plus Environment

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ABSTRACT

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The detection of hydroxyl radicals (OH•) is typically accomplished by using reactive probe

15

molecules, but prior studies have not thoroughly investigated the suitability of these probes for

16

use in electrochemical advanced oxidation processes (EAOPs), due to the neglect of alternative

17

reaction mechanisms. In this study, we investigated the suitability of four OH• probes (coumarin,

18

p-chlorobenzoic acid, terephthalic acid, and p-benzoquinone) for use in EAOPs. Experimental

19

results indicated that both coumarin and p-chlorobenzoic acid are oxidized via direct electron

20

transfer reactions, while p-benzoquinone and terephthalic acid are not. Coumarin oxidation to

21

form the OH• adduct product 7-hydroxycoumarin was found at anodic potentials lower than that

22

necessary for OH• formation. Density functional theory (DFT) simulations found a

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thermodynamically favorable and non-OH• mediated pathway for 7-hydroxycoumarin formation,

24

which is activationless at anodic potentials > 2.10 V/SHE. DFT simulations also provided

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estimates of Eo values for a series of OH• probe compounds, which agreed with voltammetry

26

results. Results from this study indicate that terephthalic acid is the most appropriate OH• probe

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compound for the characterization of electrochemical and catalytic systems.

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KEYWORDS: hydroxyl radical, probe molecule, electrochemical systems, density functional

29

theory, Marcus theory, reaction mechanism


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Graphical Abstract

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1. Introduction

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The hydroxyl radical (OH•) is a highly reactive oxidant species, which is produced via

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autoxidation in biological systems, advanced oxidation processes (AOPs), and electrochemical

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advanced oxidation processes (EAOPs).1–20 Its ability to react with organic compounds, often

36

with diffusion limited rate constants, and its rapid decay to an innocuous end product of water,

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has led to an increased interest in using technologies that generate OH• for toxic pollutant

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degradation.17,21–25 The formation of OH• at the solid-electrolyte interface is key to the operation

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of photocatalytic and EAOP technologies, and it is an unwanted side product in several energy

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technologies (e.g., polymer electrolyte membrane fuel cells; secondary batteries; redox flow

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batteries).26–35 Therefore, quantifying OH• production and understanding formation mechanisms

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are goals across numerous scientific fields. The short lifetime of OH• (~10-9 s in biological cells,22,24,25,36,37 and 10-6 to 10-3 s in water38–40)

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and high reactivity complicate its analytical determination. Therefore, two indirect methods are

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primarily employed for their detection. The first method is using electron spin resonance (ESR)

46

spectroscopy in the presence of a spin trap, which is a compound that selectively reacts with a

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radical to form an adduct species that has a sufficient lifetime to allow ESR detection. The

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second method uses probe compounds that selectively react with OH• and often form adduct

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species that are detected analytically by conventional chromatographic techniques.

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The most common spin traps for OH• determination are substituted nitrones, where 5,5-

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Dimethyl-1-pyrroline N-oxide (DMPO) is commonly used due to its high selectivity for OH•.41–50

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However, spin traps can result in the false positive detection of OH•, which is documented in the

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biochemical and toxicology fields.51–56 False positive detection of OH• has been attributed to

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inverted spin trapping52–54,57 or the Forrester-Hepburn mechanism.51,55,56 Inverted spin trapping


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involves a direct electron transfer oxidation reaction of the spin trap (R) followed by nucleophilic

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attack by water (Reaction 1). !!!

R•!

!!"!

57

R

58

The Forrester-Hepburn mechanism is the reverse of this process, which involves nucleophilic

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R − OH

(1)

attack of the spin trap by water followed by direct electron transfer oxidation (Reaction 2). !!"!

R − OH !

!!!

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R

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Both methods form a radical adduct (R-OH) which is identical to that formed by OH• attack

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on the original spin trap. Cyclic voltammetry experiments have shown that spin trap compounds

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used for OH• detection have peak potentials (Ep) of 1.47 to 2.17 V/SHE on a Pt electrode in

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nonaqueous solvents (e.g., DMPO, Ep = 1.87 V/SHE),15,58 and thus they are not suitable for OH•

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detection during photocatalytic and EAOP studies. Studies have also shown that metal ions (e.g.,

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Cu(II), Fe(III), Cr(V)) can catalyze the nucleophilic addition step of the Forrester-Hepburn

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mechanism,46,59 which greatly decreases the redox potential necessary for the subsequent direct

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electron transfer oxidation step.53 Eberson showed that the DMPO-OH- product of nucleophilic

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addition was oxidized to the DMPO-OH adduct by compounds with low redox potentials (-0.26

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to 0.74 V/SHE).53 While these mechanisms are well reported in the biochemistry literature,51–55

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this knowledge has not been adequately transferred to catalytic and electrochemical fields, and

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thus

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photocatalytic,41,49,60–62

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studies.46,50,61,64–68

spin

traps

R − OH

have

(2)

been

used

electrochemical

to

quantify

OH•

oxidation,6,11,42,44,63

production and

during

numerous

Fenton-based

oxidation

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The alternative approach of using probe molecules to detect OH• production has also been

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widely used in photocatalytic,8,24,37,62,69–76 electrochemical,7,11–16,20,63,77–83 Fenton,5,23,84–86 and


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electrochemical Fenton studies.10,87,88 Several compounds have been used, including coumarin

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(COU),13,14,16,24,37,69,72,79–81 salicylic acid,4,5,11,20 benzoic acid (BA),84,85,89,90 terephthalic acid

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(TA),8,10,13,23,69–71,73–76,87,91 1,4-benzoquinone (p-BQ),12,14,15,92–94 p-chlorobenzoic acid (p-

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CBA),7,95–103 p-nitrosodimethylaniline (RNO),15,63,82,83 luminol,73 3-carboxyproxyl,62 1,4-

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dioxane,88 1-propanol,85 and methylene blue dye.86 A summary of commonly used OH• probes

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used in various processes are included in the Supporting Information (SI Section S-1), and

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oxidation potentials of select molecules are summarized in the SI (Table S-5). The most

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commonly used OH• probes are COU, TA, BA, p-CBA, and p-BQ. The use of probe molecules

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relies on either the detection of the adduct product, or the disappearance of the probe molecule.

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Both direct electron transfer and Forrester-Hepburn mechanisms are also possible with probe

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molecules, and thus could lead to the false positive detection of OH•. These mechanisms are

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largely missing from the catalytic and electrochemical literature.

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In principle an appropriate OH• probe should 1) react with OH• at diffusion-limited rates; 2)

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be resistant to direct electron transfer oxidation reactions at solid/electrolyte interfaces and with

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soluble redox species; 3) have sufficient solubility in aqueous solutions; and 4) yield a stable

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product that is detectable by conventional analytical techniques. In this study, we experimentally

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investigated the suitability of four common probes (i.e., COU, p-CBA, TA, and p-BQ) for OH•

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detection at a boron-doped diamond (BDD) electrode using electrochemical methods. Density

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functional theory simulations were conducted to interpret the experimental data and determine

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mechanisms for probe compound oxidation. The primary objectives of this study were to

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determine a stable, selective, and sensitive probe for the detection of OH• produced during

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EAOPs, and determine the oxidation mechanisms for their removal. Results are discussed in a

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broader context related to catalytic and photocatalytic systems.


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2. Materials and Method

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2.1 Reagents. Chemicals were reagent-grade and used as received (Sigma-Aldrich). Solution

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pH was adjusted with H3PO4 or NaOH. All solutions were made with deionized water obtained

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from a NANOPure water purification system (Barnstead) with resistivity greater than 18.2

104

MΩ·cm (25o C).

105

2.2 Electrochemical Characterization. Linear sweep voltammetry (LSV) and cyclic

106

voltammetry (CV) experiments were performed using a Gamry Reference 600 potentiostat

107

(Warminster, PA). A three-electrode setup was used, with BDD as the working electrode (0.35

108

cm2), a 10 cm long, 0.3 mm diameter Pt wire as counter, and a single-junction saturated Ag/AgCl

109

electrode as reference (Pine Research Instruments, Grove City, Pennsylvania). Experiments used

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a 50 mL, 1.0 M KH2PO4 background electrolyte (pH = 5.8) at 25 °C. Potentials were reported

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versus the standard hydrogen electrode (SHE) by subtracting 0.197 V from the potential

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measured versus the reference. Before each experiment, the BDD electrode was preconditioned

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in a blank electrolyte solution (100 mM KH2PO4) at an anodic current density of 20 mA cm-2 for

114

10 min, followed by a cathodic current density of 20 mA cm-2 for another 10 min. The

115

preconditioning step was used to remove adsorbed organic contaminants from the electrode

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surface. LSV simulations were performed in Matlab 2015b, and details can be found in the SI

117

Section S-3. 2.3 Bulk Oxidation Experiments. Experiments were performed using a rotating disk

118 119

electrode (RDE) setup at 5000 rpm using a Pine Research Instruments (Grove City, Pennsylvania)

120

rotator assembly (Model: AFMSRCE). Solution temperatures (15, 25, 35 and 45° C) were

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maintained in a divided and jacketed glass reactor (H-cell) using a re-circulating water bath

122

(Thermo Scientific Neslab RTE 7). A Nafion N115 membrane separator (Ion Power Inc., New


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Castle, Delaware) was used in the H-cell to isolate anodic/cathodic reactions. A schematic of the

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reactor setup is shown in the SI Figure S-1.

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2.4 Mass Transfer Rates Constants. The mass transfer rate constant (km) was determined

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by methods described by Donaghue and Chaplin.104 Briefly, the limiting current technique was

127

used to estimate the limiting current density and km to the electrode surface at a given rotation

128

speed. The km values at 25 °C are listed in the SI (Table S-6). The diffusion coefficient was

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calculated by the Stokes-Einstein equation, and the temperature dependent viscosity was

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calculated using the equation proposed by Al-Shemmeri.105

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2.5 Reaction Rate Constants and Curve Fitting. The reaction rate constants were

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determined by by simultaneous regression of duplicate experimental data sets, and the fitting

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results are reported with 95% confidence intervals. The specific, surface area normalized

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reaction rate constant (ks (m s-1)) was obtained by dividing the first order rate constant by the

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specific surface area (0.7 m-1).

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2.6 Analytical Methods. Liquid chromatography (LC) (Shimadzu, LC-20AD) used a SPD-

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M30A variable wavelength UV-Visible detector set at 254 nm and a reverse-phase C18 column

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(Phenomenex, 4.6 × 250 mm, 5 μm). Methanol:water (40:60, v/v for COU and 50:50 for p-BQ)

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and methanol:water with 0.1% (v/v) formic acid (60:40 v/v for TA) were used as mobile phases

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at a total flow rate of 1.0 mL·min-1. The injection volume was programmed at 5 µL (COU or p-

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BQ) or 10 µL (TA). LC with a fluorescent detector (TF-20A, Shimadzu) was used for 7-

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hydroxycoumarin (7-OH COU) (λex=332 nm and λem = 471 nm) and 2-hydroxyterephthalic acid

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(2-OH TA) (λex = 315 nm and λem = 435 nm) quantification. Analysis was performed in duplicate

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and evaluated using LabSolutions Lite Version 5.63 software.


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2.7 Quantum Mechanical Simulations. Density functional theory (DFT) simulations were

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performed using Gaussian 09 software.106 Unrestricted spin, all-electron calculations were

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performed using the 6-31G+(d) basis set for geometry optimization and frequency calculations

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and the 6-311G+(3df, 2p) basis set for energy calculations. Scale factors of 0.9806 and 0.989

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were used to correct for known systematic errors in frequency and thermal energies,

150

respectively.107 Frequency calculations allowed for adjustment of the Gibbs free energies of the

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reactants and products to standard state conditions (∆! ! ). The gradient corrected Becke, three-

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parameter, Lee−Yang−Parr (B3LYP) functional was used for exchange and correlation. Implicit

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water solvation was incorporated using the SMD model.108 The standard reduction potential of the direct electron transfer reaction (! ! ) was calculated

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according to Equation (3). !! = −

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∆! ! ! !"

! − !!"# (SHE)

(3)

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where ∆! ! ! is the free energy of the reduction reaction, n is the number of electrons transferred,

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! F is the Faraday constant, and !!!" (SHE) is a reference value for the absolute standard

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! reduction potential of the SHE (!!"# SHE = 4.28 eV).109–111 Values for ∆! ! ! were calculated

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from the ∆! ! values determined from the geometrically optimized reactant and product

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structures. Gibbs free energy of activation (Ea) versus electrode potential profiles for direct electron

162 163

transfer oxidation reactions were determined using Marcus theory according to Equation (4).112 !! =

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!! !

1−

!".!(!!! ! )

!

(4)

!!

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where E is the electrode potential (V/SHE) and !! is the total reorganization energy for the

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forward reaction (oxidation reaction). The reorganization energy represents the energy needed to


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transform the geometric structure of the reactant and solvent to those of the product. Values for

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!! were calculated by subtracting ∆! ! of the reactant from that of a compound with an identical

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geometry of the product, but with the same charge as the reactant. Ionic effects of the supporting

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electrolyte on !! were not calculated, as previous work has shown the effect is minimal in polar

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solvents with high dielectric constants.113 Transition state structures were determined using the

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synchronous transit-guided quasi-newton (STQN) method114,115 using the QST3 scheme in the

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Gaussian 09 software. When this method did not converge, the pseudo reaction coordinate

174

method was used. Transition state structures were verified by confirming that frequency

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calculations yielded a single negative frequency corresponding to the normal mode along the

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reaction coordinate. Previous work indicates that errors in the calculation of the reaction energies

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by DFT methods may be up to ~ 16 kJ mol-1.116

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3. Results and Discussion

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3.1 Voltammetry Results. In order to determine if COU, TA, p-CBA, and p-BQ underwent

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direct electron transfer reactions at the BDD surface, LSV experiments were conducted with

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various concentrations of these compounds. Results from LSV experiments indicated that only

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COU (Figure 1a) and p-CBA (Figure 1c) exhibited an increased current with respect to the

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background electrolyte at approximately 1.7 and 2.0 V/SHE, respectively, while TA and p-BQ

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did not (SI Figure S-2). The absence of a current response for TA and p-BQ in this potential

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range suggested that these compounds did not undergo direct electron transfer reactions. By

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contrast, linear relationships between the peak currents and concentrations were observed for

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both COU and p-CBA (Figures 1b and 1d), suggesting they were oxidized by direct electron

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transfer reactions at the BDD surface. Additionally, LSV study of COU, p-CBA, and TA on a Pt

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electrode showed that oxidation peaks associated with direct electron transfer were found for


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both COU and p-CBA at 1.3 and 1.4 V/SHE, respectively (SI Figure S-3). Evidence for a direct

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electron transfer reaction for TA was not found (SI Figure S-3). Therefore, it is concluded that

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the findings of this study are not exclusive to BDD.

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Figures 2a and 2b show the background subtracted LSV scans of 5 mM COU at different

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scan rates and solution pH values. The reversal scans are not shown, since corresponding

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cathodic peaks were not observed in the CV scans (SI Figure S-4), indicating either a fast

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chemical reaction removed the direct electron transfer oxidation product or the product itself was

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not electrochemically active. The anodic peaks for COU oxidation shifted to less positive

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potentials with increasing solution pH, indicating OH- participated in the reaction.117 For

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example, at a scan rate of 300 mV s-1 the peak potentials for COU oxidation at pH 2.16, 5.80 and

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11.95 were 2.45, 2.38 and 2.12 V/SHE, respectively. The slope between the anodic peak

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potential and pH at a scan rate of 20 mV s-1 was -22.5 ± 3.40 mV per pH unit. Figure 2c shows

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that the peak current plotted against the square-root of scan rate follows a linear relationship,

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indicating a diffusion controlled process and therefore interaction with the electrode surface was

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not significantly affecting the electron transfer reaction.112

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The kinetic parameters governing the electrochemical reaction kinetics were determined by

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fitting a mathematical model to the experimental LSV data. The LSV model fits for COU (root-

207

mean-square error (rms) = 0.072) and p-CBA (rms = 0.058) are shown in Figures 2b and 1c. The

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simulation results provided E0 values for COU and p-CBA oxidation of 1.95 and 2.34 V/SHE,

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respectively. The oxidation kinetics of p-CBA were more facile than COU, as the first order

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heterogeneous kinetic rate constant ko for p-CBA oxidation (1.46 × 10-3 m s-1) was

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approximately 3 orders of magnitude higher than COU (6.00 × 10-6 m s-1). The values for


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transfer coefficient α were determined as 0.7 and 0.6 for COU and p-CBA oxidation,

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respectively, which agree with reported values.112

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Figure 1 LSV of (a) COU and (c) p-CBA (background subtracted) with different concentrations, and relationship between peak currents and concentrations with linear regression fit of the data for (b) COU and (d) p-CBA. In panel (c), hollow symbols and solid lines represent experimental and simulation data, respectively. Temperature: 25°C and electrolyte: 1.0 M KH2PO4 (pH 5.8).


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Figure 2 LSV scans of 5 mM COU (background subtracted) under different scan rates and solution pH values, (a) pH 2.16, (b) pH 5.80 and (c) pH 11.95 (d) Relationship between peak current and square-root of scan rate in different pH solutions. In panel (b), hollow symbols and solid lines represent experimental and simulation data, respectively. Temperature: 25°C and electrolyte: 1.0 M KH2PO4 (see each panel for pH).

225

3.2 Bulk Oxidation Experiments. Results from LSV scans indicated that COU reacted via a

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direct electron transfer reaction at potentials > 1.7 V/SHE (Figure 1a), and therefore bulk

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oxidation of 1 mM COU was performed at anodic potentials of 1.7 and 2.0 V/SHE to determine

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the Ea for COU oxidation. The solution was also monitored for the presence of the 7-OH COU

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product, which is generally considered to be an adduct compound that is formed by OH• addition

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(Figure 3). Pseudo first-order reaction rate constants were fit to the concentration versus time

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profiles measured during COU oxidation, and results are summarized in the SI (Table S-7). The


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calculated km for COU was nearly one order of magnitude higher than ks values determined for

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COU oxidation at both potentials, and therefore the oxidation of COU was assumed to be

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kinetically limited. The Arrhenius plots yielded Ea values of 45.2 ± 2.75 (SI Figure S-5) and 16.8

235

± 3.63 kJ mol-1 (Figure 3d) at 1.7 and 2.0 V/SHE, respectively. Significant concentrations of the

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fluorescent 7-OH COU product were detected during the bulk oxidation of COU (Figure 3). At

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the conclusion of the oxidation experiments 2.12 × 10-5 and 4.36 × 10-4 mM of 7-OH COU was

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observed at 1.7 and 2.0 V/SHE, respectively, corresponding to average percent yields of 0.03%

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and 0.22%, respectively. However, OH• production should be negligible at the anodic potentials

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of the experiments (1.7 and 2.0 V/SHE), due to the high redox potential for OH• formation (Eo =

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2.80 V/SHE118). Negligible OH• production is also supported by the low background current (i.e.,

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0.29 - 0.31 mA) observed in LSV scans of the blank electrolyte (Figure 1a). Therefore, it was

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assumed that the production of 7-OH COU at anodic potentials of both 1.7 and 2.0 V/SHE was

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by an alternative reaction mechanism.

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Sweep voltammetry results indicate that TA and p-BQ were resistant to direct electron

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transfer oxidation, and thus the oxidation of both compounds at a potential greater than that

247

necessary for OH• formation was attributed solely to OH• oxidation. An anodic potential of 3.0

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V/SHE, which was sufficient to produce OH• at the BDD electrode,119 was chosen for the

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oxidation of both compounds. Results from the bulk oxidation of TA and p-BQ at various

250

temperatures are shown in the SI (Figure S-6 and S-7). The values of ks obtained for the

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concentration profiles of both compounds are summarized in the SI (Table S-7). A significant

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concentration (4.63 × 10-4 mM and average percent yield of 0.11%) of the fluorescent 2-OH TA

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product from the bulk oxidation of TA was detected at the conclusion of the experiment (SI;

254

Figure S-6). Meanwhile, there were not any products detected by either the UV-vis or fluorescent


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detectors from p-BQ oxidation, indicating that ring cleavage of p-BQ occurred, which is

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consistent with previous studies.93,120,121 The ks values for both TA and p-BQ oxidation were

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comparable to the km values (SI Table S-7), indicating the bulk oxidation of both TA and p-BQ at

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3.0 V/SHE were mass transfer controlled. Mass transfer controlled oxidation of p-BQ and TA

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was attributed to the fast reaction with OH• (kOH•, p-BQ = 1.2×109 M-1s-1 21 and kOH•, TA = 4.4×109 M-1s-1

260

91

261

therefore the disappearance of p-BQ during anodic oxidation experiments was used to assess

262

OH• formation. However, p-BQ is easily reduced to hydroquinone (E0 = 0.28 V/SHE122), and

263

therefore p-BQ is only a useful OH• probe in a divided cell, where the anode and cathode are

264

separated by a selective membrane (e.g., Nafion membrane). Therefore, p-BQ is also not a useful

265

probe in photocatalytic systems where the catalyst possesses both oxidative and reductive active

266

sites.

•

). There is not a signature product from p-BQ oxidation that is indicative of OH formation, and

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Figure 3. Bulk oxidation of 1 mM COU at 2.0 V/SHE and a temperature of (a) 15 ºC, (b) 25 ºC, and (c) 35 ºC, respectively. (d) Arrhenius plot for the oxidation of COU at 2.0 V/SHE. and


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represent reactant (COU) and product (7-OH COU), respectively. Experiments were performed in duplicate, and the 1st order reaction rate model was fitted to both data sets simultaneously (dashed line). Electrolyte: 100 mM KH2PO4 (pH 4.5).

273

3.3 DFT Modeling. DFT modeling was used to determine Eo and λf values for the direct

274

electron transfer reaction (reaction 1) of each of the OH• probe compounds investigated

275

experimentally (COU, p-BQ, TA, p-CBA). These values were used in Equation (4) to determine

276

Ea profiles as a function of electrode potential. Some additional common OH• probe compounds

277

were also investigated theoretically, including BA and RNO. Results are shown in Figure 4, and

278

Eo and λf values are reported in the SI (Table S-8). The Eo values of 0.84, 2.04, 2.48, 2.55, 2.76,

279

and 3.21 V/SHE were determined for RNO, COU, p-CBA, BA, TA, and p-BQ, respectively.

280

These results were consistent with values determined by model fits of the LSV scans for COU

281

and p-CBA, which yielded Eo values of 2.34 V/SHE (Figure 1c) and 1.95 V/SHE (Figure 2b),

282

respectively. A comparison between experimental and DFT determined Ea values for COU also

283

showed close agreement (Figure 4). Experimentally measured Ea values for COU were 45.2 ±

284

2.75 kJ mol-1 at 1.7 V/SHE and 16.8 ± 3.63 kJ mol-1 at 2.0 V/SHE, which were within -1.2 kJ

285

mol-1 and 12.1 kJ mol-1 of DFT determined values, respectively (Figure 4). The close agreement

286

between experimental and theoretical Ea values at 1.7 V/SHE shows the robustness of the

287

theoretical method. The difference between experimental and theoretical Ea values at 2.0 V/SHE

288

is largely a result of unactivated processes influencing the experimental Ea measurement. For

289

example, temperature effects on molecular diffusion, adsorption, or the structure of the double

290

layer can cause errors in accurately measuring Ea values in this range.123

291

Figure 4 also contains the Ea versus potential profile for OH• formation at an oxidized BDD

292

surface, taken from Chaplin et al.124 Results indicated that the direct electron transfer reaction for

293

RNO and COU proceeded at potentials much less than that for OH• formation, and thus neither


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RNO nor COU are useful OH• probes. The Ea versus potential profiles for p-CBA and BA in

295

Figure 4 are intersected by the profile for OH•, indicating direct electron transfer reactions were

296

likely dominant at potentials < 2.3 V/SHE, and that both direct electron transfer and OH•

297

oxidation likely occurred at potentials > 2.3 V/SHE for these compounds. The results from LSV

298

experiments showed the direct electron transfer oxidation of p-CBA occurred at potentials > 2.0

299

V/SHE (Figure 1c). By contrast, direct electron transfer reactions for TA and p-BQ were

300

predicted to occur after OH• formation. At anodic potentials > 2.3 V/SHE the water oxidation

301

reaction dominated the measured current and therefore LSV scans were not able to provide

302

evidence for direct electron transfer reactions involving p-BQ and TA. In addition, these

303

compounds have fast reaction rate constants with OH• (i.e., kOH•, BQ = 1.2×109 M-1 s-1; kOH•, COU =

304

2.0×109 M-1 s-1),21,91 and at high anodic potentials reactions became mass transfer limited (as

305

noted in SI Table S-7). Past studies have concluded that under mass transfer limited conditions

306

the indirect oxidation with OH• was the dominant oxidation mechanism, because the substrates

307

were depleted via reaction with OH• in the diffusion boundary layer and did not reach the

308

electrode surface.104 The results above suggest that p-BQ and TA are resistant to direct electron

309

transfer reactions and therefore should be appropriate OH• probes, while other probe compounds

310

investigated are not, and their use in electrochemical and heterogeneous catalytic systems should

311

be revisited.

312

An additional requirement for the OH• probes is that they are nonreactive by the Forrester-

313

Hepburn mechanism. Reactivity by this mechanism was assessed by determining Ea and ∆! ! !

314

values for the nucleophilic addition of OH- to each of the compounds, followed by calculation of

315

Eo and λf values for the subsequent direct electron transfer reaction (Table 1). Transition state

316

structures for the nucleophilic reactions are provided in the SI (Figures S-8). Simulations were


17


Page 18 of 41

317

conducted for OH- attack on the C7 and C2 atoms for COU and TA, respectively, as these were

318

the fluorescent adduct probes used to assess OH• production in this and other

319

studies.8,10,13,14,16,23,24,37,69–76,79,81,87,91,125 Simulations determined that nucleophilic addition of OH-

320

to COU yielded Ea = 130 kJ mol-1 and ∆! ! ! = 118 kJ mol-1, and addition of OH- to TA yielded

321

Ea = 123 kJ mol-1 and ∆! ! ! = 111 kJ mol-1 (Table 1). The relatively high Ea values and positive

322

∆! ! ! values suggested that these reactions were unlikely to be significant at room temperature.

323

Experimental results supported these theoretical calculations, as removal of COU and TA were

324

not observed under alkaline conditions (pH = 12.4, T = 25 oC). However, catalytic nucleophilic

325

addition has been reported to occur by metal ions in solution (e.g., Fe, Cu, Cr).46,59 Therefore,

326

this reaction may be catalyzed by Fenton type reactions and therefore warrants further study. The

327

Eo values for the corresponding direct electron transfer reactions are -0.23 and -0.05 V/SHE for

328

COU and TA (Table 1), respectively; indicating that these reactions would proceed at anodic

329

potentials provided the nucleophilic addition reactions occurred.

330


18

Page 19 of 41


331 332 333 334

Figure 4. DFT determined Ea versus electrode potential profiles. Profile for OH• taken from Chaplin et al., 2013.124 Squares represent experimental values determined for COU oxidation in RDE experiments.

335 336

Table 1. Theoretical parameters determined for the reaction of probe molecules by the ForresterHepburn mechanism. Reaction COU + OH- à COU-OHCOU-OH- à COU-OH + eTA + OH- à TA-OH-

-

TA-OH à TA-OH + e

p-BQ + OH- à p-BQ-OHp-BQ-OH- à p-BQ-OH + ep-CBA + OH- à p-HBA + Cl-

∆! !!

Ea

Eo

λf

(kJ mol-1) 118

(kJ mol-1) 130

(V/SHE) --

(kJ mol-1) --

--

--

-0.23

13.4

111

123

--

--

--

--

-0.05

11.9

14.9

42.1

--

--

--

--

1.22

14.9

-146

125

--

--

337 338

The nucleophilic addition of OH- to p-BQ was investigated at both the C1 and C2 positions.

339

Due to the symmetry of the molecule, these were the only two unique positions for OH- attack.

340

The ∆! ! ! for nucleophilic addition of OH- at the C1 and C2 positions of p-BQ were 14.9 and 289

341

kJ mol-1, respectively (Table 1). These results indicated that attack at the C1 position was

342

thermodynamically favorable and the calculated Ea of 42.1 kJ mol-1 indicated that the reaction

343

was feasible at room temperature (Table 1).126 The corresponding direct electron transfer

344

reaction has a calculated Eo = 1.22 V/SHE, and thus this reaction would readily occur at anodic

345

potentials used in EAOPs and during photocatalysis.

346

The nucleophilic addition of OH- to p-CBA was studied at the C4 position, and resulted in a

347

substitution reaction that formed p-hydroxybenzoic acid (p-HBA), with Cl- as the leaving group

348

(Table 1). The calculated Ea and ∆! ! ! values for the nucleophilic substitution reaction were 125


19


Page 20 of 41

349

and -146 kJ mol-1, respectively (Table 1), indicating the reaction was thermodynamically

350

favorable but likely not significant at room temperature. Since p-HBA is a stable product and p-

351

CBA is usually used as a OH• probe based on its disappearance, the direct electron transfer

352

reaction was not studied. These results indicated that p-BQ has the potential to react via the

353

Forrester-Hepburn mechanism, and therefore reactions should not be conducted in alkaline

354

conditions. Results for other compounds are consistent with studies of nucleophilic attack on

355

aromatic derivatives, which indicated that high temperatures are needed for significant reaction

356

rates to proceed.127–129

357

Experimental results showed that the anodic oxidation of COU produced the 7-OH COU

358

product at anodic potentials lower than that necessary to produce OH• (i.e., 1.7 and 2.0 V/SHE),

359

which suggested that this product was not exclusively formed by reaction with OH•. Therefore,

360

DFT simulations were conducted to determine a thermodynamically favorable mechanism for its

361

formation. The overall anodic reaction of COU to form 7-OH COU is given by Reaction (5). 362

+ OH- + H2O

→

+ H3O+ + 2e-

363

(5)

364

This reaction involves 2e- oxidation, which may occur by either direct electron transfer or

365

indirect oxidation with OH•. Results in Figure 4 show that the direct electron transfer oxidation

366

was activationless at potentials > 2.1 V/SHE, which was likely the initial step responsible for

367

conversion to 7-OH COU via a non-OH• radical mechanism. A possible reaction mechanism that

368

does not involve OH• oxidation is given in Figure 5, where ∆! ! ! values were calculated at an

369

electrode potential of E = 2.044 V/SHE (Ea = 2.44 kJ mol-1 (Figure 4)). The mechanism involves

370

three steps, one single electron transfer electrochemical (E) step, followed by one chemical (C)

371

step and another single electron transfer (E) step, and the overall ∆! ! ! = -390 kJ mol-1 at E =


20

Page 21 of 41


372

2.044 V/SHE. The mechanism was defined as an ECE mechanism. At an electrode potential of

373

2.044 V/SHE, the first electrochemical step (E1) had ∆! ! ! = 0 kJ mol-1, and values for Ea versus

374

electrode potential are given in Figure 4. All subsequent steps were exothermic. The first

375

chemical step (C1) involved nucleophilic attack by OH-, which was activationless and ∆! ! ! = -

376

152 kJ mol-1. The second electrochemical step (E2) involved a coupled H+ and e- transfer to form

377

7-OH COU. This reaction had a calculated Eo = -0.43 V/SHE, and thus was activationless at an

378

electrode potential of 2.044 V/SHE (∆! ! ! = -239 kJ mol-1). The participation of OH- in the

379

reaction mechanism was consistent with LSV experimental results that showed OH- dependence

380

on the measured peak potential for COU (Figure 2). The stepwise e- and H+ transfer reactions

381

were also investigated, which yielded Eo = 0.81 V/SHE for the e- transfer step and the H+ transfer

382

to H2O was activationless.

383

The results above provided an activationless and exothermic pathway for the formation of 7-

384

OH COU from COU oxidation at an applied anodic potential of E > 2.1 V/SHE, indicating that

385

COU was not a useful OH• probe. Since the initial direct electron transfer reaction was rate

386

limiting, the Ea versus E profile given in Figure 4 should accurately describe the energy barrier

387

for 7-OH COU formation as a function of potential for the non-OH• pathway.


21


Page 22 of 41

+

150 100 50 ΔG (kJ mol-1)

0 -50 -100

+ OH- + H2O

+ OH- + H2O + e-

ΔGE1 = 0 (E = 2.04 V/SHE) ΔGC1 = -102

+ H2O + e- ΔGr = -341

-150 -200

ΔGE2 = -239 (E = 2.04 V/SHE)

-250 -300 -350 -400

E1: COU-H à COU-H+ + e- (Eo = 2.04 V/SHE)

+ H3O+ + 2e-

C1: COU-H+ + OH- à COU-H-OH E2: COU-H-OH + H2O à COU-OH + H3O+ + e- (Eo = -0.43 V/SHE)

388 389

Figure 5. Free energy diagram of COU oxidation by ECE mechanism.

390

3.4 Broader Significance and Environmental Impacts. Although this study was focused

391

on selecting appropriate OH• probes for EAOPs, our findings have broader impacts that extend to

392

other processes involving electron transfer, such as photocatalysis and energy storage/conversion

393

technologies. In photocatalysis, the formation of OH• occurs through a direct electron transfer

394

reaction from H2O to the UV light excited holes in the valence band of the photocatalyst. This

395

same type of direct electron transfer reaction mechanism can occur for OH• probe compounds,

396

and thus may be falsely interpreted as an indication of OH• formation. Often the direct electron

397

transfer reaction mechanism was assessed in photocatalysis by adding a OH• scavenging

398

compound and assessing its effect on a given substrate.8,24,37,69,70,72–76,98,99 This method does not

399

provide definitive proof for or against a direct electron transfer mechanism, since the OH•

400

scavenger can block direct electron transfer reaction sites and thus slow the reaction of the


22

Page 23 of 41


401

substrate. In energy storage/conversion technologies, OH• can be formed from H2O2, which is a

402

product of the oxygen reduction reaction at the cathode surface. The formation of OH• is also

403

possible during the charging process in aqueous electrolytes. Therefore, the use of appropriate

404

probe compounds to characterize OH• formation during the charging/discharging process at the

405

electrodes or during fuel cell operation is necessary to prevent false positive detection of OH•.

406

A detailed literature review of OH• probes used in various processes is provided in the SI

407

Section S-1 (Table S-1 to S-4), and electrochemical oxidation potentials of select molecules are

408

summarized in the SI (Table S-5). Table S-5 indicates that OH• probes, such as salicylic acid,

409

luminol, 1-propanol and methylene blue, are susceptible to direct electron transfer reaction, since

410

their reported oxidation potentials are lower than 2.3 V/SHE (Figure 4), however, 1,4-dioxane,

411

with the oxidation potential at 2.4 V, may be an appropriate OH• probe in electrochemical and

412

catalytic systems, but requires further investigation.

413

Many prior studies have used EAOPs and photocatalysis for the oxidation of contaminants in

414

water, and therefore the correct identification and quantification of OH• is necessary in order to

415

understand and optimize these processes for environmental remediation. Careful attention should

416

be given when selecting the appropriate probes to detect the formation of OH• in these processes.

417

Spin trap compounds, RNO, COU, and p-CBA can react by direct electron transfer and

418

Forrester-Hepburn mechanisms, and therefore are not appropriate probes. The results of this

419

study indicated that TA is the most selective OH• probe investigated, since it has a high redox

420

potential (Eo = 2.76 V/SHE), is resistant to the Forrester-Hepburn mechanism, and is not easily

421

reducible.

422

ASSOCIATED CONTENT


23


Page 24 of 41

423

Supporting Information

424

Literature reviews of OH• probes used in various systems, Electrochemical oxidation potentials

425

of selective compounds, Details of LSV simulation, Schematics of System Setup, LSV of COU,

426

p-CBA and TA on Pt Electrode, Diffusion coefficients and mass transfer rates for COU, TA and

427

p-BQ, LSV of TA and p-BQ, CV of COU, Reaction rate constants for bulk oxidation of COU,

428

TA, and p-BQ, COU oxidation under various temperatures at 1.7 V vs SHE, TA and p-BQ

429

oxidation under various temperatures at 3.0 V vs SHE, Theoretical E0 and λf values determined

430

by DFT methods and used to determine Ea versus potential profiles, and Transition state

431

structure.

432

AUTHOR INFORMATION

433

Corresponding Author

434

*Phone: 312-996-0288; fax: 312-996-0808; e-mail: [email protected].

435

Acknowledgements

436

Funding for this work was provided by the National Science Foundation (CBET-1159764 and

437

CBET-1453081 (CAREER)).

438

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