Article pubs.acs.org/JPCC
Electrochemical Properties of Phenols and Quinones in Organic Solvents are Strongly Influenced by Hydrogen-Bonding with Water Malcolm E. Tessensohn, Hajime Hirao,* and Richard D. Webster* Division of Chemistry and Biological Chemistry, School of Physical and Mathematical Sciences, Nanyang Technological University, Singapore 637371 S Supporting Information *
ABSTRACT: The electrochemical behavior of several phenols, quinones and hydroquinone in acetonitrile (CH3CN) with varying amounts of water were investigated to understand the effect of hydrogen-bonding on their voltammetric responses. Karl Fischer coulometric titrations were performed to obtain an accurate reading of the water concentrations. The solvent/electrolyte mixture was carefully dried using 3 Å molecular sieves to obtain an initial water content that was close to the substrate concentration (∼1 × 10−3 M), and higher water contents were then achieved via the addition from microliter syringes. It was found that small changes in what is often considered “trace” amounts of water were sufficient to substantially change the potential and in some cases the appearance of the voltammetric waves observed during the oxidation of the phenols/hydroquinones and reduction of the quinones. Density functional theory calculations were performed on the reduced/oxidized species in the presence of varying numbers of water molecules to better understand the hydrogen-bonding interactions at the molecular level. The results highlight the importance of accurately knowing the trace water content of organic solvents when used for voltammetric experiments. concentration varies.1 The reaction in eq 1 represents an extreme situation where the equilibrium constant strongly favors the forward direction, such that all the reduced compounds are hydrogen-bonded. The potential dependence of reaction 1 can thus be described by the Nernst equation (eq 2) (using concentrations in place of activities, neglecting the charge on the molecules and assuming the reaction is kinetically favorable). Eobs is the observed voltammetric potential, E0f is the formal potential, R (8.3143 J K−1 mol−1) is the gas constant, T the temperature and F is the Faraday constant (96485 C mol−1).
1. INTRODUCTION Voltammetric experiments are frequently conducted in nonaqueous solvents in order to aid analyte solubility, increase the available voltage window and prevent reactions of the reduced or oxidized species with water.1,2 However, despite it being universally recognized that water is itself an ubiquitous impurity in organic solvents, its influence on the voltammetric behavior of the analytes within the nonaqueous medium is seldom quantitatively examined. Electrochemistry experiments in nonaqueous solvents are generally performed under three major conditions in relation to the water content of the solutions;3 (i) a very low water regime ( > 1
[Red(H 2O)y ] 2.303RT × log [Ox][H 2O]y nF
(1)
(2)
When there is an equal amount of [Ox] and [Red(H2O)y] at the electrode surface, the observed potential (Eobs) in eq 2 depends on log 1/[H2O]y and the voltammetric wave will shift linearly (more positively) for every 10-fold change in the water concentration. For example, a small change in water concentration from 1 × 10−3 M to 10 × 10−3 M will result in the same positive shift in potential as a change in water concentration from 1 to 10 M. Therefore, the logarithmic dependence of the voltammetric peak potential on the water Received: November 7, 2012 Revised: December 20, 2012
A
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chosen from a number of aprotic organic solvents suitable for electrochemistry because it is easily dried by storing over 3 Å molecular sieves, does not undergo strong hydrogen-bonding with water (unlike dimethyl sulfoxide and dimethylformamide), affords good solubility of the substrates, has a reasonably high dielectric constant (ε = 37) to enable dissolution of electrolytes, and because it is miscible with high concentrations of water (unlike dichloromethane). The experimental data are compared with results from density functional theory calculations on the starting and intermediate species produced during the oxidation and reduction processes in order to better understand the interactions with water at the molecular level.
concentration means that the largest shifts in potential occur at low water concentrations. Since the natural variation in trace water content of the organic solvent typically occurs in mM levels (∼0.001−0.1 M), it would be expected that small variations in trace water would significantly shift the voltammetric waves for compounds that undergo strong hydrogen-bonding interactions. The predicted potential shift in eq 2 is analogous to what is observed for electron transfer reactions that are coupled to proton transfer steps in buffered aqueous solutions, where the reversible half-wave potential (Er1/2) varies linearly with the pH.6 Eq 2 can be used to predict that the potential will change by approximately 60 mV per logarithmic unit change in water concentration for a one-electron process at 298 K involving one water molecule hydrogen-bonding. However, in situations where the equilibrium in eq 1 does not strongly favor the product, or there are multiple species and hydrogen-bonding steps involved, the individual equilibrium constants need to be incorporated into the logarithmic term in eq 2.7,8 It has been known for many years that reduced quinones undergo strong hydrogen-bonding interactions with water,6−22 although only recently have the effects of trace (mM) levels of water been studied in detail.21,22 In organic solvents, quinones undergo two one-electron reduction processes; the first process at E1 produces the anion radical (semiquinone) and the second process at E2 produces the dianion. As water is added to the solvent, the E1 and E2 processes both shift to more positive potentials due to hydrogen-bonding, although the E2 shifts much more than E1 due to stronger interactions of H2O with the dianion (ΔE2 ≫ ΔE1). When large amounts of water are added to the solvents, E2 shifts so much that the two oneelectron processes eventually merge into one two-electron process.6−22 Concerted proton-coupled electron transfer (PCET) reactions have been reported to occur during the oxidation of phenols in the presence of water,23−31 although there is little information currently available regarding how the oxidation of phenols are affected by trace amounts of water in organic solvents. Phenolic compounds (such as antioxidants) are often present in biological systems in lipophilic (low water) environments; thus studying the effects of hydrogen-bonding with low levels of water is potentially important in understanding their biological functions. A recent study examined the phenol vitamin E in CH3CN and observed shifts in the voltammetric oxidation wave to more negative potentials as small amounts of water were added to the solvent.32 NMR spectra that were obtained during the addition of D2O to solutions of phenols displayed a shift in frequency of the hydroxyl proton resonance, which was interpreted as due to hydrogen-bonding interactions.32 However, the hydrogen− deuterium exchange reaction meant that the shifts became difficult to detect as the D2O concentration increased above a 10-fold excess.32 FTIR spectroscopic experiments are problematic because of the strong infrared absorption of the “free” H2O masking the hydrogen-bonding interactions of the hydroxyl groups (of phenols) or carbonyl groups (of quinones) with H2O. Voltammetric measurements have the advantage that relatively high concentrations of water in the solvent do not interfere with the current−potential measurements. In this study, we have examined the interactions of a number of substituted para-quinones, phenols and 1,4-benzohydroquinone in CH3CN with deliberately added water during voltammetric reduction and oxidation processes. CH3CN was
2. EXPERIMENTAL SECTION 2.1. Chemicals. 1,4-Benzoquinone (1), 2,5-di-tert-1,4benzoquinone (2), 2,6-di-tert-1,4-benzoquinone (3), 2,4,6trimethylphenol (4), 2,6-di-tert-butyl-4-methylphenol (5), 2,4,6-tri-tert-butylphenol (6), 1,4-benzohydroquinone (7), and ferrocene (Fc) were obtained from Sigma-Aldrich and used as received (Chart 1). The supporting electrolyte, tetra-nChart 1. Structures of the Quinones, Phenols, and Hydroquinone
butylammonium hexafluorophosphate (n-Bu4NPF6), was prepared by a literature procedure and stored under vacuum.33 1/ 16 in. rods with 3 Å pore size molecular sieves (CAS: 308080− 99−1) were obtained from Fluka. Analytical grade CH3CN was acquired from Merck and dried over 3 Å molecular sieves before use. 2.2. Measurement of Water Content. Karl Fischer titrations were performed with a Mettler Toledo DL32 coulometer using (Riedel-deHaën) HYDRANAL-Coulomat AG and HYDRANAL-Coulomat CG as the anolyte and catholyte, respectively. The coulometer was allowed to stabilize until a steady drift value close to 0 μg min−1 of water was achieved. Constant humidity (30%) measurements were carried out in a (122 cm × 61 cm ×61 cm) humidity control box using a dry nitrogen gas purge system from Coy Laboratory Products Inc. 5 Plastic disposable syringes were used to inject approximately 1 mL aliquots of the electrochemical solutions, via a silicon/Teflon septum; the individual measurements were found to conclude within 1 min, thereby indicating that the drift from atmospheric water was insignificant. 2.3. Voltammetry. Cyclic voltammetry and square-wave voltammetry experiments were conducted with a computercontrolled Eco Chemie Autolab PGSTAT302N potentiostat in a three-electrode cell where a 1 mm diameter planar glassy carbon (GC) disk (Cypress Systems) working electrode was used together with a platinum wire (Metrohm) auxiliary electrode. A platinum wire (Metrohm) reference electrode was used in the presence of 3 Å molecular sieves for the experiments under ultradry conditions whereas a silver wire (Cypress Systems) miniature reference electrode connected to the test solution via a salt bridge containing 0.5 M n-Bu4NPF6 in CH3CN was employed in all other experiments. The internal filling solution of the liquid-junction reference electrode was B
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functional36 within the Gaussian 09 suite of programs.37 The final geometries of the compounds were then verified as true minima via frequency calculations at the B3LYP/6-311+G(2df,p) level where no imaginary frequencies were detected. Zero-point energies were scaled by a factor of 0.9889.38 The methods for calculating adiabatic electron affinities (AEAs) of the quinones and adiabatic ionization energies (AIEs) of the phenols were similar to those used in the literature.39−41
freshly prepared before each experiment to reduce the amount of water that it contributed to the analyte solution.5 All voltammetric experiments were performed in a Faraday cage at room temperature and under an argon atmosphere. 2.4. Preparation of Dried Solutions. CH3CN (predried with 3 Å molecular sieves for at least 48 h) was used to dissolve the respective electroactive species and supporting electrolyte (dried under vacuum at 413 K for 6 h) to give a solution containing 2.0 × 10−3 M of the compound and 0.2 M nBu4NPF6 in CH3CN, before quantitatively transferring into a 25 mL vacuum syringe (SGE Analytical Science) containing 3 Å molecular sieves (dried under vacuum at 413 K for 6 h). The syringe was subsequently stored under a nitrogen atmosphere for a period of 48 h and shielded from light to prevent unwanted photochemical reactions. 2.5. Procedure for Voltammetric Experiments with Addition of Water. A dummy cell was used to obtain the initial moisture content of the experimental solution owing to the fluctuating natural humidity levels. Ten milliliters of the dried solution was accurately measured and injected into both the experiment and dummy cells that were previously heated at 373 K for an hour and cooled under an argon atmosphere, before deoxygenating with argon gas for at least 2 min. Consecutive cyclic voltammetry (CV) and square-wave voltammetry (SWV) potential scans were performed prior to extracting a 1 mL aliquot from the dummy solution for Karl Fischer titration to obtain an estimate of the water content present in the experiment solution. This was followed by careful additions of accurately controlled microliter volumes of purified water into the experiment solution and further deoxygenation by purging with argon gas. For the quinones, their experiments were concluded upon the merger of the two redox processes or separation of the aqueous/organic solution into two immiscible layers, whichever occurred earlier. The experiments for all other compounds were terminated upon addition of 1 mL of water. The water content at the end of each experiment was subsequently calculated based on the total amount of water added. 2.6. Procedure for Ultradry Experiments. In situ voltammetric experiments in the presence of 3 Å molecular sieves were performed in order to obtain the voltammetric behavior of the compounds under ultradry conditions. The molecular sieves were heated at 433 K for 6 h prior to cooling and transferring into an electrochemical cell, which itself was heated at 373 K for an hour and cooled under an argon atmosphere. The dried solution was subsequently injected into the electrochemical cell, deoxygenated by purging with argon gas, and subjected to cyclic voltammetric and square-wave voltammetric analyses. After completion of each voltammetric analysis, approximately 1 mL of the solution was withdrawn for a Karl Fischer titration. Careful attention was paid to ensure that the moisture content was less than the concentration of the electroactive species before the results were accepted as obtained under ultradry conditions. 2.7. Computational Methods. The molecular geometries of all compounds, inclusive of the intermediate species in their respective redox mechanisms, were drawn using GaussView 5, while their structures were optimized using the 6-311+G(2df,p) Pople-style Gaussian basis set34,35 (triple ζ quality with one set of diffuse s- and p-functions on heavy atoms, two additional dtype polarization functions and an additional f-type polarization function on heavy atoms, and one set of p-type polarization functions on hydrogen) and the B3LYP hybrid density
3. RESULTS AND DISCUSSION 3.1. Effect of Water on the Reduction of Quinones. Quinones may be reduced in two successive one-electron transfers to form first the radical anion (Q•−), followed by the dianion (Q2−) at more negative potentials (Scheme 1). Typical Scheme 1. Successive Reduction Mechanism of Quinones in Aprotic Organic Solvents
quinone voltammetric responses in very dry CH3CN are shown in Figure 1a for compounds 1 − 3 where the two reduction processes are clearly evident.6−22 Compounds 1 and 2 show additional reduction and oxidation processes (with much smaller current values) before and after the major waves which are possibly associated with surface based interactions. Increasing the water content during the experiments resulted in the first (Epred(1)) and second (Epred(2)) reduction peaks and half-wave potentials (Er1/2(1) and Er1/2(2)) for 1, 2, and 3 shifting toward less negative potentials (see Figures S1−S3 in the Supporting Information section). The potential difference between Epred(1) and Epred(2) (ΔEpred = |Epred(1) − Epred(2)|) was measured using square-wave voltammetry, as shown in Figure 1b. Since the shift in Epred(2) was much greater than Epred(1) when water was added to the solvent, the two one-electron processes eventually merged into one-two electron process at high water concentrations. An increase in the cathodic peak current, ipred, was detected upon the combination of the two reduction processes, but it was less than a 2-fold increase for doubling the number of electrons transferred and some additional processes became observable at the higher water concentrations. Depicted in Figure 2 are plots of ΔEpred as determined by square-wave voltammetry against the concentration of water for the different quinones. It can be observed that there is a gradual decrease in the potential between the reduction processes as the water concentration is increased, and that the largest relative shifts occur at the lowest water concentrations, which is expected based on the ΔEpred-values depending on the logarithmic function of the H2O concentration. The shifts in potential with respect to changing water concentrations were quite similar for 1, 2, and 3. This indicates that the bulky tertbutyl substituents in 2 and 3 did not appear to greatly hinder hydrogen-bonding interactions of the quinones. From the results in Figure 2, it is clear that when voltammetric experiments are performed on quinones (and possibly other carbonyl containing compounds) in organic solvents in order to investigate the interactions with hydrogenbonding additives (such as amines and alcohols),42−45 it is C
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with the substrate; thus different ΔEpred-values are obtained at different supporting electrolyte concentrations.22 Use of a pseudo Pt wire reference electrode at the lowest water concentrations ( Q−H2O and arises from the increased charge density at the oxygen atoms with successive reduction. A more detailed explanation of the hydrogenbonding interactions at the molecular level based on DFT calculations is given in section 3.5. 3.2. Effect of Water on the Oxidation of Phenols. Phenols typically undergo a −2e−/−H+ oxidation process (Scheme 3) in aprotic organic solvents, via an ECE mechanism
step occurs rapidly. The CVs of several phenols in CH3CN containing varying concentrations of H2O are shown in Figure 3, where chemically irreversible oxidation processes (with no reverse reduction peak at a scan rate of 100 mV s−1) were detected. This observed chemical irreversibility arises from the highly reactive PhO+ cation, which undergoes hydration or selfreactions (e.g., dimerization, dealkylation) depending on the amount of water present and the nature of the substituents in the ortho and para positions.3,49−53 It was not possible to directly compare the absolute peak potentials of the voltammetric waves recorded in ultradry CH3CN (red dashed lines in Figure 3) with the higher water concentrations, due to the different reference electrodes employed (see Experimental Section). Nevertheless, in ultradry CH3CN, the voltammograms of 4 and 5, as shown in Figure 3, parts a and b, displayed additional oxidation processes at potentials 200−300 mV more positive than the main oxidation process; the electrochemical behavior of 6, however, showed no observable difference (Figure 3c). It is interesting to note that changes in the appearance of the voltammetric waves of 4 and 5 occurred at water levels >1
(4)
[Red][H 2O]y 2.303RT × log nF [Ox(H 2O)y ]
(5) 54
Although H2O is able to act as a weak base in CH3CN, the addition of water to solutions of CH3CN does not cause the deprotonation of the phenols. This was confirmed by UV−vis experiments that only showed the presence of the phenols and not the deprotonated phenols (phenolate anions) when water was added to the CH3CN. The phenolates, which can be easily prepared by the addition of a strong organic soluble base such as Bu4NOH, have substantially different UV−vis spectra than the phenols.32 Furthermore, phenolate anions are much easier to oxidize than their corresponding phenols by at least 1 V.3,50 Therefore, the potential shifts that are observed in this work are due to the weaker hydrogen-bonding interactions.
(where E and C represent electron- and proton-transfers, respectively), to produce phenoxonium cations (PhO+).3,48,49 In some systems the initial electron and proton transfer reactions have been proposed to occur in a concerted step.23,24,29 Even though the overall oxidation involves the transfer of two electrons, only one anodic peak is detected because the second oxidation process occurs at less positive potentials than the first and the intermediate deprotonation E
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Figure 3. Cyclic voltammetric data obtained at 22 ± 2 °C with 0.2 M n-Bu4NPF6 as the supporting electrolyte at a scan rate of 100 mV s−1 with a 1 mm diameter planar GC electrode: (a) 2.0 × 10−3 M 4 (b) 2.0 × 10−3 M 5, and (c) 2.0 × 10−3 M 6 at different concentrations of H2O. The plots in red were obtained in ultradry CH3CN using a platinum reference electrode.
Scheme 4. Electrochemical −2e−/−2H+ Oxidation Mechanism of Hydroquinones in Aprotic Organic Solvents
It can also be observed in Figure 3 that the variations of Epox in 4 are much more significant than those of the other phenols, which is likely because the ortho positions to the hydroxyl group are most critical in affecting the degree of hydrogenbonding. Since compound 4 contains only methyl groups adjacent to the hydroxyl group, it provides lesser steric hindrance to water molecules undergoing hydrogen-bonding compared to 5 and 6. The slight difference between 5 and 6 suggests that steric bulk at the para position also contributes to the shifting of Epox with changing water levels. The addition of water to the solutions lowered the bulk concentration of the phenols, bringing about a decrease in the peak anodic current, ipox, at very high water concentrations. Intermolecular and intramolecular hydrogen-bonding interactions have been reported for phenolic compounds in the presence of amines or other potential hydrogen-bonding additives.55−59 Although amines would be expected to undergo stronger hydrogen-bonding interactions than H2O, the difference in the voltammetric behavior between the amine hydrogen-bonded and non-hydrogen-bonded compounds would be clearer if the water could be completely removed from solution (i.e., removing the potential shifts caused by hydrogen-bonding with trace water). 3.3. Effect of Water on the Oxidation of 1,4Benzohydroquinone. Compound 7 contains two hydroxyl groups but, unlike the phenols, forms a chemically long-lived product after oxidation that can be reduced back to the starting material. Upon a −2e−/−2H+ oxidation via an ECEC mechanism, 7 gives 1 as the final product (Scheme 4). The oxidation of 7 in unbuffered organic solvents is not completely chemically reversible since one anodic peak is detected on the forward oxidation scan while a smaller cathodic peak is seen on the reverse. Only a single anodic peak is observed in the voltammograms because QH2•+ is more acidic than the starting material and the kinetics of the first deprotonation reaction occur very rapidly to form QH•. QH• is immediately further oxidized to QH+ because the oxidation potential for the QH•/ QH+ couple occurs at a less positive potential than the oxidation of QH2/QH2•+, which then deprotonates to form Q (1). CVs on compound 7 give rise to two oxidation peaks at low water concentrations, which merge into one process as the water content increases (Figure 4). Therefore, the presence of
Figure 4. Cyclic voltammograms of approximately 2.0 × 10−3 M 7 in CH3CN at various concentrations of water obtained at 22 ± 2 °C with 0.2 M n-Bu4NPF6 as the supporting electrolyte at a potential scan rate of 100 mV s−1 with a 1 mm diameter planar GC electrode.
trace water is clearly altering the voltammetric response, although the exact chemical reason for this phenomenon is currently uncertain. It is also noted that 7 displayed a smaller ipox compared to the current values obtained from the quinones and phenols. This is largely due to the drying procedures employed where storing of the solution in a vacuum syringe F
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point energy plus zero-point correction) of the compounds and their respective oxidized or reduced states. For the phenols, the AIE of PhOH was calculated from the difference in the ground state energies of PhOH and PhOH•+, while that of the radical species PhO• was obtained via the difference in energies of PhO• and PhO+. AEAs of quinones followed a similar trend where the computation for Q was obtained via the difference in ground state energies of Q and Q•− whereas that of Q•− was arrived through the difference in energies of Q•− and Q2−. Diagrams showing the energies, geometries and Cartesian coordinates of the calculated hydrogen-bonded structures and computed AEA and AIE values are provided in the Supporting Information section. Geometries arising from primary (substrate−water) and secondary (water−water) hydrogen-bonding interactions were also examined especially when numerous hydrogen-bonding sites were available. Quinones and phenols may undergo many secondary hydrogen-bonding interactions in addition to the primary hydrogen-bonding interactions with the oxygen atoms and hydroxyl groups, respectively, thereby potentially increasing the number of possible structural arrangements that may exist. Although careful attention was paid to ensure that most of these geometries were explored, not all conformers may have been determined because of the numerous conformational possibilities. The electron affinity of a molecule describes its propensity to accept an electron and, therefore, allows for the investigation of the effects of hydrogen-bonding on the reduction of the quinones. However, when dealing with how the medium (bulk solvent/electrolyte) affects the hydrogen-bonding between the water molecules and the substrate (quinone, phenol, hydroquinone) a number of factors can be considered, including; (i) interactions between the solvent and substrate (solvation), (ii) interactions between the solvent and water which are important for solvents such as dimethyl sulfoxide and dimethylformamide which interact very strongly with water,22 and (iii) interactions between the substrate and the supporting electrolyte, which have been observed during the reduction of quinones.21,22,46 For the compounds under study, it would appear that the solvation effects play relatively minor roles, since the addition of small amounts of water in the presence of a huge excess of solvent are still sufficient to shift the voltammetric waves; thus, the trends for the AIE and AEA values are unlikely to differ greatly from those which include solvation effects. Nevertheless, gas-phase and solvated geometry optimizations were performed on compound 1 to assess the differences in the calculated AEA values. The results obtained from the calculations (Table 1) indicate that the solvent affects the absolute AEA values but not the relative values for varying degrees of hydrogen-bonding. In both cases, it was found that the AEA values increased with increasing degrees of hydrogen-bonding. This supports the
containing molecular sieves was necessary to ensure minimum moisture content. The use of molecular sieves, however, also encourages 7 to undergo surface based adsorption interactions (with the sieves), consequently decreasing the bulk concentration and ipox value. It is clear that the Epox of compound 7 shifts to less positive potentials with increasing moisture content (Figure 4), similar to that observed for the phenols (Figure 3). It is interesting that both 4 and 7, which contain no bulky substituents, show two anodic processes at low water concentrations (
