Homogeneous Electron-Transfer Reaction between Electrochemically

Apr 30, 2013 - The homogeneous catalytic oxidation of propylamine (PrA), diethylamine (DEA), pyrrolidine (Pyr), and triethylamine (TEA) has been inves...
0 downloads 5 Views 1MB Size
Download PDF
Article pubs.acs.org/Organometallics

Homogeneous Electron-Transfer Reaction between Electrochemically Generated Ferrocenium Ions and AmineContaining Compounds Angel A. J. Torriero,*,† Muhammad J. A. Shiddiky,‡ Iko Burgar,§ and Alan M. Bond*,∥ †

Institute for Frontier Materials, Deakin University, Burwood, Victoria 3125, Australia Australian Institute for Bioengineering and Nanotechnology (AIBN), University of Queensland, Brisbane, Queensland 4072, Australia § CMSE Division, CSIRO, Clayton, Victoria 3800, Australia ∥ School of Chemistry, Monash University, Clayton, Victoria 3800, Australia ‡

S Supporting Information *

ABSTRACT: The homogeneous catalytic oxidation of propylamine (PrA), diethylamine (DEA), pyrrolidine (Pyr), and triethylamine (TEA) has been investigated for the first time in the presence of electrochemically generated ferrocenium ions as the catalyst. Mechanistic details for this electrocatalytic process have been investigated by cyclic voltammetry and other electrochemical techniques. A oneelectron oxidative process was observed for all amines. Deviation from this mechanism was only observed during controlled-potential bulk electrolysis of ferrocene in the presence of high concentrations of propylamine, where problems with electrode fouling and catalyst deactivation processes were encountered. The catalytic efficiency and the catalytic oxidation rate constants were estimated and found to follow the order TEA > Pyr > DEA > PrA. Interestingly, the catalytic reaction was not observed when ferrocene was replaced by decamethylferrocene. This observation was analyzed in terms of thermodynamic and structural effects, and a hypothesis is presented. The homogeneous catalytic oxidation reported opens a new avenue to achieve simple, low-cost, and efficient amine oxidation, which is potentially useful in several areas of chemistry.

■

INTRODUCTION The oxidation of amines plays an important role in biology and is crucial to several biochemical processes.1,2 Amines are also important as additives in plating baths, as reagents in the syntheses of dyes and polymeric materials, and for the manufacture of many medicinally important chemicals and hence are of considerable interest to the chemical and pharmaceutical industries.2 Electrochemical detection of aliphatic amines by direct oxidation at conventional electrodes, such as glassy carbon (GC) and platinum, is complicated due to surface fouling by oxidation products (most commonly with oxidation products of primary amines). The need to apply very positive potentials also presents a problem. The potential for oxidation is dependent on their structure, with primary amines (1.6−1.7 V vs NHE) being more difficult to oxidize than secondary or tertiary amines (1.0−1.3 V vs NHE).1 In order to decrease the overpotential for oxidation as well as minimize the surface fouling effect, increase the sensitivity, and enhance the reliability and reproducibility of the data, catalytic oxidation of aliphatic amines with a reversible redox couple acting as a mediator has been introduced.3,4 An ideal catalyst (or mediator) needs to have a standard reversible potential less positive than the oxidation potential of the substrate, exhibit © 2013 American Chemical Society

fast electron-transfer kinetics, and be stable in both the oxidized and reduced forms toward the species present in the reaction media.5−7 Effective mediators that meet these requirements are frequently based on ferrocene (Fc) and its derivatives, either homogeneously dispersed in the solution or immobilized in a monolayer or multilayer configuration onto the electrode surface.8−15 A well-defined reversible Fc ⇆ Fc+ + e− process is generally observed when both Fc and Fc+ are soluble in the media of interest. This feature has influenced the recommendation of ferrocene and some of its derivatives in the provision of a reference potential scale in voltammetric studies in ionic liquids and organic solvents containing added supporting electrolyte.16−20 A selective and sensitive electrochemical assay based on DNA-mediated charge transport coupled to a catalytic reaction between surface-bound ferrocene moieties and diethylamine present in solution has been recently developed by our group.7 However, quantitative analysis of such an electrocatalytic Special Issue: Ferrocene - Beauty and Function Received: March 19, 2013 Published: April 30, 2013 5731

dx.doi.org/10.1021/om4002318 | Organometallics 2013, 32, 5731−5739

Organometallics

Article

interaction between ferrocenium ions and amine-containing compounds has yet to be reported. In the present work, mechanistic details of the homogeneous catalytic oxidation of primary, secondary, and tertiary aliphatic amines in acetonitrile (0.1 M Bu4NPF6) have been investigated by voltammetric and related techniques when the electrochemically generated ferrocenium ion is used as the catalyst. The catalytic oxidation rate constant (kf), the catalytic efficiency, and the influences of the experimental conditions on the catalytic oxidation current are also discussed. The effect of substitution of all 10 hydrogen atoms present in the cyclopentadienyl ring of ferrocene with methyl groups on the catalytic efficiency has also been evaluated. The results presented in this work open a new avenue to achieve simple and efficient amine oxidation, which is potentially useful in many areas of chemistry.

■

(E°P/Q , k°P/Q , αP/Q )

(Keq , k f , k b)

(1)

(2)

(E° A/B , k° A/B , αA/B)

(k1)

(3)

where E°A/B, k°A/B, αA/B, and k1 are the formal reversible potential, formal heterogeneous rate constant, transfer coefficient, and forward rate constant, respectively. This mechanism, also known as homogeneous redox catalysis, is concerned with systems in which the catalyst diffuses freely in the solution that contains the substrate. The key step involves the electron-transfer reaction between the substrate and the catalyst molecules, which serves as the electron sink in place of the electrode, the active form of the catalyst being generated by the electrode. In the case of an oxidation process, the reduced form P of the catalyst couple P/Q is introduced to the solution. In the potential region where the catalytic reaction happens, the mass transport for all diffusing species of interest follows the relationships

∂cQ ∂t

= DQ

∂ 2cQ ∂x 2

− k f cAcQ + k bc Bc P

(9)

t ≥ 0, x = 0 (electrode/solution interface): cP = exp[(F /RT )(E − E°P/Q )] cQ

(10)

(11)

0 ≤ t ≤ tλ , E = E i + vt

(12)

0 < t ≤ 2tλ , E = E i + 2v − vt

(13)

(14)

where a, F, and D are the electrode area, Faraday constant and diffusion coefficient of P, respectively. Under these conditions, setting the electrode potential at a value that achieves oxidation of P results in the formation of Q, which oxidizes A in the solution. An increase of the oxidation current of P as in comparison to its value in the absence of the substrate is then observed, accompanied by a loss of reversibility of the P ⇆ Q + e− process, hence providing the features associated with the catalytic process. The classical outcome entails detection of a sigmoidal-shaped catalytic steady-state response, with the plateau current being independent of scan rate and scan direction.22 The steadystate catalytic current observed in this situation is well described by the expression

(4)

∂c P ∂ 2c = DP 2P + k f cAcQ − k bc Bc P ∂t ∂x

c P = c*P , cA = c*A , cQ = 0, cA = 0

⎛ ∂c ⎞ I = FaDP⎜ P ⎟ ⎝ ∂x ⎠x = 0

and follow up reactions that are a result of transient formation of intermediate B B→C

(8)

where Ei is the initial potential, v is the scan rate, and tλ is the time required for a complete sweep of potential. The current, I, can be calculated according to eq 14, once the space- and timedependent concentrations are known

where E°P/Q (V) is the formal potential, k°P/Q (cm s−1) is the formal heterogeneous rate constant, and αP/Q is the transfer coefficient associated with reaction 1 and Keq, kf (M−1 s−1), and kb (M−1 s−1) are the equilibrium constant and forward and backward rate constants associated with reaction 2, respectively. In the present study, P = ferrocene, Q = ferrocenium, and A = the reduced form of triethylamine, diethylamine, propylamine, and pyrimidine. Implicitly involved in the EC′ mechanism is the seldom-discussed heterogeneous electron transfer6 A ⇆ B + e−

∂c B ∂ 2c = DP 2B + k f cAcQ − k bc Bc P − k1c B ∂t ∂x

where ci and Di (i = P, Q, A, B) are the concentration and diffusion coefficient of the species i, respectively, x is the distance from the electrode surface, c*P and c*A are the initial concentration of P and A, and kf, kb, and k1 are the rate constants defined in eqs 2 and 4. If the electrode potential E is scanned linearly:

and Q+A⇆P+B

(7)

∂cQ ∂c B ∂c P ∂c + = 0, A = 0, =0 ∂x ∂x ∂x ∂t

THEORY The classical electrochemical catalytic mechanism,5,21 often referred to as the EC′ mechanism, is traditionally described by the two reactions P ⇆ Q + e−

∂cA ∂ 2c = DA 2A − k f cAcQ + k bc Bc P ∂t ∂x

IL = Fac*P DPk f c*A

(15)

where the half-wave potential, E1/2, coincides with E° of the couple P/Q. From all the key parameters involved in the theoretical model outlined above, the influence of c*P, c*A, kf, and v on the observed voltammograms are of particular interest in this work. Simulated voltammograms for the mechanism described above were calculated at 294 K, using an electrode area of 0.72 mm2 and a DA value of 1.0 × 10−5 cm2 s−1 (DA being the estimated diffusion coefficient of triethylamine, diethylamine, propylamine, and pyrimidine, respectively). It is worth noting that DA was selected arbitrarily, as the values for the amines being studied are not available in acetonitrile (0.1 M Bu4NPF6). Therefore, a value previously reported in aqueous media23 was used in the simulations. However, IL (eq 15) does not depend

(5)

(6) 5732

dx.doi.org/10.1021/om4002318 | Organometallics 2013, 32, 5731−5739

Organometallics

Article

upon DA. Other parameters used for the simulations are provided in the relevant figure captions.

Table 1 Aminea

ΔE (V)b

γ

IL/γId

kf (M−1 s−1)c

RESULTS AND DISCUSSION Ferrocene (Fc) undergoes a reversible, one-electron-oxidation process under voltammetric conditions in both acetonitrile and dichloromethane solutions containing 0.1 M Bu4NPF6 as the supporting electrolyte, to give the ferrocenium cation (Fc+) (eq 16, Figure 1, curve a)). The midpoint potentials (Em),

PrA DEA TEA Pyr

1.71 0.79 0.57 0.85

18 18 18 18

0.046 0.057 0.084 0.062

1.6 4.2 21.0 7.1

■

Fc ⇌ Fc+ + e−

a

PrA = propylamine, DEA = Diethylamine, TEA = triethylamine, and Pyr = pyrrolidine. bDifference between the reversible potential of the Fc0/+ process and the irreversible oxidation peak potential for the designated amine compounds. cValues obtained by comparison of digital simulations of cyclic voltammograms and experimental data over a wide range of scan rates, Fc concentrations, and amine concentrations.

(16)

calculated from the average of the oxidation (Epox) and reduction (Epred) peak potentials (Em = 1/2(Epox + Epred)), are 0.509 and 0.548 V, respectively, vs the DmFc0/+ redox couple, used to provide the potential reference scale. The ratio Ipc/Ipa = 0.99 (Ipc and Ipa being the peak currents for the cathodic and anodic peaks, respectively) confirms the chemical reversibility of the Fc0/+ system on the voltammetric time scale. The oxidation of Fc to the corresponding Fc+, in the presence of amine-containing compounds, was also studied. Figure 1 provides examples of catalytic voltammograms observed upon addition of 2.6 M diethylamine (DEA, Figure 1A, c) and 1.81 M propylamine (PrA, Figure 1B, c). These voltammograms show a clear loss of reversibility with an enhancement of the anodic peak current. Given that the irreversible peak potentials for the direct oxidation of DEA and PrA compounds at the GC electrode surface are 0.79 and 1.71 V more anodic with respect to oxidation of Fc (Figure 1A,B, curve b and Table 1), respectively, the increase in the anodic current on voltammograms c is attributed to the regeneration of ferrocene resulting from the reaction of Fc+ with DEA and PrA, respectively (eqs 17 and 18). Consequently, the current

detected, I (eq 19), is the sum of two contributions: (i) the diffusion-controlled current (Id) of Fc in the absence of amineI = Id + IL

containing compound, and (ii) the catalytic current (IL) as described by eq 15. The cyclic voltammogram for the direct oxidation of DEA on a GC electrode in CH3CN (0.1 M Bu4NPF6) depicted in Figure 1A, b shows an oxidation peak at around 1.30 V vs DmFc0/+ when the potential is scanned in the anodic direction. When the potential scan direction is reversed, no complementary reduction peak is observed over the scan rate range studied (0.02−5 V s−1). This behavior is typical for a fast irreversible chemical reaction coupled to the charge-transfer step. The electrochemical oxidation of secondary and tertiary aliphatic amines has been demonstrated to proceed via a oneelectron-oxidation process.24 The reaction pathway proposed by Mann et al. for tertiary amines is summarized by eqs 20−24 (R = aliphatic groups).25,26 Oxidation leads to a radical cation which deprotonates to give a radical (eqs 20 and 21).

CH3CH 2NHCH 2CH3 + Fc+ •+

⇌ CH3CH 2 NHCHCH3 + Fc

(19)

(17)

•+

R1R 2NCH 2CH 2R3 ⇌ R1R 2 NCH 2CH 2R3 + e− •+

CH3CH 2CH 2NH 2 + Fc+ ⇌ CH3CH 2CH 2 NH 2 + Fc

•+

•

R1R 2 NCH 2CH 2R3 ⇌ R1R 2NCHCH 2R3 + H+

(18)

(20) (21)

This radical can be involved in a disproportionation process to yield the starting amine and an enamine

Figure 1. Cyclic voltammograms obtained with a scan rate of 0.1 V s−1at a glassy-carbon electrode (1 mm diameter) for the oxidation of 1.0 mM Fc (A and B, a), 19.3 × 10−3 M DEA (A, b), 24.3 × 10−3 M PrA (B, b), and 1 mM Fc in the presence of 2.6 M DEA (A, c) and 1.81 M PrA (B, c) in 0.1 acetonitrile (0.1 M Bu4NPF6). T = 21 ± 1 °C. 5733

dx.doi.org/10.1021/om4002318 | Organometallics 2013, 32, 5731−5739

Organometallics

Article

Figure 2. Cyclic voltammograms obtained for 0.51 mM Fc in acetonitrile (0.1 M Bu4NPF6) at a glassy-carbon electrode (d = 1 mm) over the potential region where the Fc oxidation occurs (A) with the addition of 0.0, 0.12, 0.30, 0.58, 1.11, and 1.81 M PrA at a scan rate of 0.02 V s−1, (B) with the addition of (a) 3.63 M and (b) 4.68 M PrA at a scan rate of 0.02 V s−1, and (C) in the presence of 1.11 M PrA at scan rates of 0.1, 0.2, 0.3, 0.4, 0.5, 0.7, and 1.0 V s−1.

reaction would consume one electron per starting molecule.24−26 The electrochemical oxidation of primary amines has also been demonstrated to proceed via a one-electron-oxidation process.24 However, the radical cation formed as a product of the first electron transfer is more unstable than those for tertiary and secondary amines and is able to follow different pathways with respect to those described by eqs 20−24. Electrode fouling is also a possibility.24,27 Accordingly, Figure 1B (curve b) shows a voltammogram for oxidation of 24.3 × 10−3 M PrA in CH3CN (0.1 M Bu4NPF6) solution. This voltammogram exhibits an irreversible oxidation process at 2.22 V vs DmFc0/+ with a crossover loop after the potential scan direction is reversed, which may be associated with attachment of an organic functional group, most likely nitrogen, to the surface of the GC electrode. However, it should be noted that this interaction cannot be observed on the voltammetric time scale during the catalytic reaction and when the potential is scanned into the region where Fc oxidation occurs. Consistent with previous observations, the fouling effect was not observed for the tertiary and secondary amines used in this study. This difference is attributed to an increase in the steric hindrance of the radical cation formed after the electrochemical oxidative step.24 Figure 2A provides examples of cyclic voltammetric responses for 0.51 mM Fc obtained in the absence and presence of increasing concentrations of propylamine (see

or be involved in a second oxidation step to yield an iminium cation (eq 23).

It was also stated that the enamine and iminium ions should be in equilibrium (eq 24), as the enamine would be a stronger base than the starting saturated amine.26

It is important to note that both the enamine and iminium formed as products in eqs 22−24 can react with traces of water present in the organic solvent to yield an aldehyde as a final product.24 For example, in the case of a tertiary amine being oxidized, the overall reaction could provide a secondary amine, an aldehyde, and protons. This proton would protonate a starting amine or a secondary one (product from this reaction) to give an electrochemically inactive ammonium ion, and the 5734

dx.doi.org/10.1021/om4002318 | Organometallics 2013, 32, 5731−5739

Organometallics

Article

Figure 3. Cyclic voltammograms obtained for 1.05 mM (a−d), 5.08 mM (e−h), and 10.2 mM (i−l) Fc in acetonitrile (0.1 M Bu4NPF6) at a glassycarbon electrode (d = 1 mm) over the potential region where the Fc oxidation occurs with the addition of 0.77 M (a, e, i), 1.49 M (b, f, j), 3.49 M (c, g, k), and 5.35 M (d, h, l) DEA at a scan rate of 0.02 V s−1.

exhibits voltammograms obtained at a scan rate of 0.02 V s−1 upon increasing the ferrocene concentration from 1.05 to 10.2 mM (from top to bottom) and increasing the concentration of DEA from 0.77 to 5.35 M (from left to right). As a consequence, the excess factor, γ = c*A/c*Fc (c*A and c*Fc are the concentrations of DEA and ferrocene, respectively), is increasing from voltammogram i to d (Figure 3). When the γ values are large enough for the consumption of DEA to be negligible, the classical sigmoidal-shaped catalytic response is observed, with the cathodic and anodic traces almost superimposed and with IL following the relationship described in eq 15. Thus, a purely kinetically controlled condition is achieved when the homogeneous kinetics are rate limiting. Conversely, for small values of γ, the consumption of the DEA is faster than its diffusion from the bulk solution to the electrode, with its diffusion now being the rate-limiting step. Under this circumstance, a peak is detected when the potential is scanned in the positive direction and superimposed onto the sigmoidal-shaped component. This behavior is also found with other amines studied in this paper. The number of electrons exchanged (n or napp) during controlled-potential oxidation of 1.0 mM Fc in the absence (n) and presence (napp) of amine-containing compounds was determined in acetonitrile (0.1 M Bu4NPF6) by using a largesurface-area GC working electrode and applying a potential where only the catalytic oxidation takes place (an effective oxidation potential was found to be 0.6 V vs the DmFc0/+ redox couple, which is around 0.1 V more positive than the midpoint potential of the Fc0/+ process). Changes in the 1H and 13C NMR spectra accompanying exhaustive controlled-potential electrolysis were investigated. For the NMR measurements, acetonitrile was replaced by dichloromethane, which did not affect the catalytic behavior (the effective oxidation potential was found to be 0.65 V vs the DmFc0/+ redox couple in CH2Cl2). The current obtained at constant potential was recorded as a function of time until it decayed to about 2% of the initial value. Under the same conditions, controlledpotential electrolysis experiments were undertaken in both

Figure S1, Supporting Information, for DEA behavior). The addition of PrA triggers a considerable increase of the catalytic current, with a concomitant variation of shape in the voltammograms. At a low concentration of PrA, the rate of regeneration of the catalyst is negligible in comparison to the diffusion rate. Thus, the Fc oxidation current remains totally diffusion-controlled. Nevertheless, at a given scan rate and when the IL/Id ratio increases as a consequence of increasing the concentration of PrA up to 3.63 M, the current reaches a limiting value. Simultaneously, the voltammograms change from peak- to sigmoidal-shaped curves, where the half-wave potential, E1/2, for the sigmoidal-shaped case coincides with the Em value for the Fc0/+ couple. With PrA concentrations higher than 4.0 M, as illustrated in Figure 2B, the process is again peak-shaped, and more so with a lower scan rate. This behavior may be a result of decay of the catalyst and electrode fouling, as revealed by exhaustive controlled-potential electrolysis (see below). The change from sigmoidal- to peak-shaped curve at high PrA concentrations was not observed for TEA, DEA, and Pyr at concentrations of up to 0.21, 5.35, and 0.35 M, respectively. Figure 2C shows the effect of scan rate on the catalytic process under conditions of excess substrate. As the scan rate decreases, the regeneration of ferrocene is kinetically favored and a sigmoidal-shaped curve is obtained. Em eventually becomes independent of the scan rate and equal to E1/2 when the sigmoidal shape is attained. Decreasing the scan rate further does not alter Em or IL. This steady-state regime achieved at low scan rates is a consequence of having the rate which the Fc disappears, due to oxidation at the GC electrode, being equal to the rate at which Fc is regenerated by PrA. This condition may arise when the catalysis is not too fast so that the concentration of the substrate throughout the reaction layer can be the same as its bulk concentration.28,29 Under this condition, IL is represented by the relationship in eq 15. In order to obtain further insights into the catalytic mechanism, the effect of variation of both ferrocene and DEA concentration on the voltammetry was studied. Figure 3 5735

dx.doi.org/10.1021/om4002318 | Organometallics 2013, 32, 5731−5739

Organometallics

Article

Figure 4. Cyclic voltammograms obtained before (A and B, a) and after (A and B, b) exhaustive controlled-potential electrolysis at 0.6 V vs DmFc0/+ for (A) 1.0 mM Fc in the presence of 2.0 mM TEA and (B) 1.0 mM Fc in the presence of 60 mM PrA in acetonitrile (0.1 M Bu4NPF6). T = 21 ± 1 °C.

Scheme 1. Proposed Mechanism for the Catalytic Oxidation of TEA

they are protonated at the end of the electrolysis. This is in agreement with the observations of Adenier et al. during their studies on the direct electrooxidation of amines.24 From the results of controlled-potential electrolysis with voltammetric monitoring and 13C and 1H NMR, it is possible to hypothesize that the electrochemical oxidation of ferrocene (eq 16) triggers the reaction sequence shown in Scheme 1. In this mechanism, TEA transfers one electron to Fc+ to regenerate Fc and to afford a radical cation that deprotonates to give a radical, which follows the reactions pathways given in eq 22−24 to produce triethylammonium and iminium ions as the final products. In the presence of water, and on the basis of the mechanism reported by Mann et al.,25,26 the expected final products would be triethylammonium, diethylammonium, and an aldehyde. Controlled-potential electrolysis of ferrocene in the presence of primary amines was also performed to determine the napp values and to examine the catalytic reaction between ferrocene and PrA on an experimental time scale larger than that which applies to cyclic voltammetry. Controlled-potential oxidation of 1.0 mM Fc in the presence of 1.0 mM PrA, under the same experimental conditions as the case of TEA and DEA, provides a napp value of 1.8 ± 0.1 electrons per molecule of Fc, which again indicates that one electron is also being transferred per molecule of PrA. This result is in agreement with the napp values observed previously during the direct electrooxidation of

acetonitrile (0.1 M Bu4NPF6) and dichloromethane (0.1 M Bu4NPF6) solutions in the absence of ferrocene and amines. The coulombs obtained were subtracted to the total values in order to obtain the ferrocene Faradaic component. As expected, the total coulombs gave an n value of 1.0 ± 0.1 electrons per molecule of Fc in the absence of amine. However, the Fc napp values were 1.9 ± 0.1 and 3.0 ± 0.1 electrons when 1.0 and 2.1 mM TEA or DEA was added to the Fc solution, which implies that one electron per molecule of TEA (or DEA) is transferred to Fc+ during the catalytic reaction. The 13C NMR spectrum of TEA in CH2Cl2 (0.1 M Bu4NPF6) shows resonances at δ 11.90 and 46.05, which progressively disappear during the course of exhaustive controlled-potential electrolysis and are replaced by resonances at δ 1.50, 11.60, 51.52, 56.15, and 116, which are in accord with those expected for the iminium ion. The 1H NMR spectrum in CH2Cl2 (0.1 M Bu4NPF6) after electrolysis shows new resonances at δ 1.99 and 2.75, which also suggests the presence of the iminium ion in the electrolyzed solution. Figure 4A shows cyclic voltammograms of a mixture of 1.0 mM Fc and 2.0 mM TEA in acetonitrile (0.1 M Bu4NPF6) obtained before (Figure 4A, a) and after (Figure 4A, b) exhaustive controlled-potential electrolysis. The fact that neither TEA (oxidation process at around 1.1 V vs DmFc0/+ in Figure 4A, a) nor a secondary amine, formed as a product of the catalytic reaction (in this case DEA), is detected by cyclic voltammetry in the fully electrolyzed solution indicates that 5736

dx.doi.org/10.1021/om4002318 | Organometallics 2013, 32, 5731−5739

Organometallics

Article

Figure 5. Comparison of experimental (black ) and simulated (red ○) cyclic voltammograms for the oxidation of 5.1 mM Fc in acetonitrile (0.1 M Bu4NPF6) at a glassy-carbon electrode (d = 1 mm) over the potential region where the Fc oxidation occurs with the addition of 3.49 M (a), 4.17 M (b), 5.35 M (c), and 6.12 M (d) DEA at a scan rate of 0.02 V s−1. The mechanism depicted in the Theory section was used for simulations. Simulation parameters: T = 294 K, electrode area = 0.72 mm2, charge-transfer coefficient (α) = 0.5, heterogeneous standard rate constant (k0) = 5 cm s−1, equilibrium constant (Keq) = 1 × 103. E°f = 0.5 V, kf = 4.2 s−1, Cdl = 0.40 μF, DFc = 2.3 × 10−5 cm2 s−1.

primary amines.24 Nevertheless, when the concentration of PrA is increased to 60 mM, a larger than expected increase in time required for exhaustive electrolysis was encountered. This could be associated with a slowing down of the heterogeneous electron transfer. As indicated above, the PrA radical formed after deprotonation could lead to binding of the organic moiety to the glassy-carbon surface. In addition, an napp value of 14.3 is obtained, which is far less than that expected for transfer of one electron per molecule of PrA. Figure 4B shows cyclic voltammograms of 1.0 mM Fc in acetonitrile (0.1 M Bu4NPF6) and in the presence of 60 mM PrA obtained before (Figure 4B, a) and after (Figure 4B, b) exhaustive controlled-potential electrolysis. The absence of the ferrocene redox process in the final solutions clearly indicates a mechanism involving deactivation of the catalyst during the controlled-potential electrolysis, which is a commonly encountered event in homogeneous catalysis.30 The fact that ferrocene deactivation and surface fouling are observed at concentrations ≥4 M on the cyclic voltammetric time scale, but at lower concentrations on the much larger bulk electrolysis time scale, implies that the reactions leading to electrode fouling are slow. On the basis of the bimolecular one-electron mechanism proposed in Scheme 1, the catalytic efficiency can be defined by the ratio IL/γId.30 In cases of significant overlap of the magnitudes of the catalytic and substrate currents, the value of IL was corrected from the background current obtained with

only substrate present (at the same concentration). The values of the catalytic efficiency obtained under these conditions are given in Table 1 together with the values of ΔE = Em(Fc0/+) − Ep, with Ep being the peak potential value for the irreversible oxidation of the amine-containing compounds. Even when the Ep values do not represent the thermodynamic value required for a correct correlation, a qualitative trend could be seen from data in this table, where the efficiency of the catalytic process diminishes as the irreversible peak potential for oxidation of the amines becomes more and more positive with respect to ferrocene. This may be attributed to an increase in the standard free energy of the reaction. On this basis, an identical situation is expected to arise when the reversible potential of the catalyst becomes less positive with respect to a particular amine compound. In this context, substitution of all 10 hydrogen atoms of the cyclopentadienyl ring of ferrocene with methyl groups lowers the reversible potential of the ferrocene process by 0.509 V in acetonitrile (0.1 M Bu4NPF6), which should reduce considerably the catalytic efficiency. Interestingly, no significant changes in the voltammetric behavior of 1 mM DmFc in acetonitrile (0.1 M Bu4NPF6) was observed on addition of 0.88 M DEA or 0.23 M TEA (Figure S2, Supporting Information). This result suggests that, in addition to providing an increase in the standard free energy of the reaction, methyl substitution on the cyclopentadienyl ring prevents specific 5737

dx.doi.org/10.1021/om4002318 | Organometallics 2013, 32, 5731−5739

Organometallics

Article

electrode and an Ag/AgCl (CH2Cl2, 0.1 M Bu4NPF6) single-junction quasi-reference electrode were used for studies in acetonitrile and dichloromethane, respectively. Nevertheless, all potential values are quoted relative to the decamethylferrocene/decamethylferrocenium (DmFc0/+) potential scale. A glassy-carbon (GC) working electrode with a diameter of 1.0 mm and an effective area of 7.2 mm2 was used for cyclic voltammetric experiments. The area or radii of the voltammetric working electrode was determined from cyclic voltammograms using the peak current derived as a function of the square root of scan rate from the oxidation of a 1.0 mM solution of ferrocene in CH3CN (0.1 M Bu4NPF6) (diffusion coefficient of 2.3 × 10−5 cm2 s−1) degassed with N2 and application of the Randles−Sevcik relationship.31 Prior to each voltammetric experiment, the working electrode was polished with 0.3 μm alumina (Buehler, Lake Bluff, IL) on a clean polishing cloth (Buehler), sequentially rinsed with distilled water and acetone, and then dried with lint-free tissue paper. Bulk electrolyses were performed inside a nitrogen-filled glovebox using a glassy-carbon tube working electrode and a platinum-gauze auxiliary electrode separated from the test solution by a fine-porosity glass frit. All voltammetric experiments were carried out at ambient temperature (21 ± 1 °C). The commercially available software package DigiElch was used to simulate the voltammetric responses. 1 H and 13C nuclear magnetic resonance (NMR) spectra were obtained at 21 ± 1 °C with a Varian 300 MHz VNMRS spectrometer in dichloromethane (0.1 M Bu4NPF6). No suppression of solvent or electrolyte resonances was required. Chemical shift values are reported in ppm and are referenced relative to TMS.

interaction between the permethylated ferrocenium cation and the amine. The second-order rate constants, kf, for homogeneous electron-transfer reactions between the ferrocenium cation and various amine derivatives were estimated by comparison of digital simulations of the cyclic voltammograms and experimental data (Figure 5 and Figures S3 and S4 in the Supporting Information). Simulations based on the mechanism outlined in the Theory section (see also Scheme 1) give excellent agreement with experimental observations made over a wide range of catalyst concentrations (0.2−10 mM) and amine concentrations (0.02−0.15, 3.49−6.12, 0.30−1.81, and 0.10− 0.35 M for TEA, DEA, PrA, and Pyr, respectively), when the rate constants summarized in Table 1 are employed. Determinations of the rate constants were also performed at different scan rates (from 0.02 to 0.7 V s−1) to minimize the interference of natural convection in the electrochemical response, which is normally observed at scan rates ≤0.05 V s−1. Significantly, kf values decrease in the order TEA > Pyr > DEA > PrA, which is in agreement with the catalytic efficiency trend observed in Table 1.

■

CONCLUSIONS

The present study demonstrates that the electrochemically generated ferrocenium cation has the ability to oxidize propylamine, diethylamine, triethylamine, and pyrrolidine in acetonitrile containing 0.1 M Bu4NPF6 as the supporting electrolyte, thus serving as a suitable catalyst for the homogeneous oxidation of aliphatic amines in aprotic media. Mechanistic details that apply for this reaction have been investigated by cyclic voltammetry and bulk electrolysis techniques. The catalytic reaction between ferrocenium and relatively low concentrations of secondary and tertiary amines follow a oneelectron-oxidation mechanism with almost quantitative recovery of the catalyst at the end of the experiment. Propylamine behaves similarly to the other amines on the voltammetric time scale, but differences appear when techniques that require longer reaction times are employed. To this end, controlled-potential bulk electrolysis of ferrocene in the presence of high concentrations of PrA provides evidence of both ferrocene deactivation and surface fouling. The catalytic efficiency and the second-order rate constants for the homogeneous electron-transfer reaction between ferrocenium and the different amine derivatives were estimated and found to follow the order TEA > Pyr > DEA > PrA. Most notably, the catalytic reaction was not observed when ferrocene was substituted by decamethylferrocene, which suggests that, in addition to the increase in the standard free energy of the reaction, the methyl substitution on the cyclopentadienyl ring may prevent interaction between the permethylated ferrocenium cation and the amine molecules. The results presented in this work should open up a new avenue to achieve simple, low-cost, and efficient amine oxidation, which could be useful in several areas of chemistry.

■

■

ASSOCIATED CONTENT

S Supporting Information *

Figures giving voltammograms of ferrocene in the absence and presence of DEA, voltammograms of decamethylferrocene in the absence and presence of DEA and TEA, and simulations of voltammograms. This material is available free of charge via the Internet at http://pubs.acs.org.

■

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected] (A.A.J.T.); alan.bond@ monash.edu (A.M.B.). Fax: +61 3 9244 6868. Tel: +61 3 9244 6897. Notes

The authors declare no competing financial interest

■

ACKNOWLEDGMENTS We gratefully acknowledge Deakin University, Monash University, and the Australian Research Council for financial support. We thank Professor Leone Spiccia for his insightful discussions.

■

EXPERIMENTAL SECTION

REFERENCES

(1) Steckhan, E. In Organic Electrochemistry; Lund, H., Hammerich, O., Eds.; Marcel Dekker: New York, 2001. (2) Lawrence, S. A. Amines: Synthesis, Properties and Applications; Cambridge University Press: Cambridge, U.K., 2004. (3) Largeron, M.; Neudorffer, A.; Benoit, M.; Fleury, M.-B. Proc. Electrochem. Soc. 2003, 2003−12, 185. (4) Regino, M. C. S.; Brajter-Toth, A. Electroanalysis 1999, 11, 374. (5) Saveant, J. M.; Vianello, E. Electrochim. Acta 1965, 10, 905. (6) Feldberg Stephen, W.; Campbell Jennifer, F. Anal. Chem. 2009, 81, 8797. (7) Shiddiky, M. J. A.; Torriero, A. A. J.; Zeng, Z.; Spiccia, L.; Bond, A. M. J. Am. Chem. Soc. 2010, 132, 10053. (8) Ojani, R.; Raoof, J. B.; Alinezhad, A. Electroanalysis 2002, 14, 1197.

Reagents and Chemicals. Diethylamine (DEA; 99.5%, SigmaAldrich), triethylamine (TEA; 99.5%, Sigma), propylamine (PrA; 99.0%, Aldrich), pyrrolidine (Pyr; 99.0%, Fluka), decamethylferrocene (DmFc; 97%, Aldrich), ferrocene (Fc; 98%, Aldrich), and tetrabutylammonium hexafluorophosphate (Bu4NPF6; ≥99.0%, Fluka), were used as received from the manufacturer. HPLC grade acetonitrile (CH3CN; Merck), and dichloromethane (CH2Cl2; Merck) were distilled over CaH2 and stored inside a homemade nitrogen-filled glovebox. Instrumentation and Procedures. Electrochemical experiments were carried out in a standard three-electrode arrangement with a BAS Model 100B electrochemical workstation (Bioanalytical Systems, West Lafayette, IN). A platinum wire was used as an auxiliary electrode. An Ag/Ag+ (0.01 M AgNO3, 0.1 M Bu4NPF6) double-junction reference 5738

dx.doi.org/10.1021/om4002318 | Organometallics 2013, 32, 5731−5739

Organometallics

Article

(9) Gao, Z.-N.; Zhang, J.; Liu, W.-Y. J. Electroanal. Chem. Interfacial Electrochem. 2005, 580, 9. (10) Gao, Z.-N.; Ma, J.-F.; Liu, W.-Y. Appl. Organomet. Chem. 2005, 19, 1149. (11) Zhang, S.; Fu, Y.; Sun, C. Electroanalysis 2003, 15, 739. (12) Lertanantawong, B.; O’Mullane, A. P.; Zhang, J.; Surareungchai, W.; Somasundrum, M.; Bond, A. M. Anal. Chem. 2008, 80, 6515. (13) Rusling, J. F.; Wang, B.; Yun, S.-e.. In Bioelectrochemistry: Fundamentals, Experimental Techniques and Applications; Bartlett, P. N., Ed.; Wiley: Chichester, U.K., 2008. (14) Kaneko, Y.; Kadokawa, J.-I.; Tagaya, H.; Chiba, K. Recent Res. Dev. Pure Appl. Chem. 1998, 2, 563. (15) Ferrocenes: Homogeneous Catalysis, Organic Synthesis, Materials Science; Togni, A., Hayashi, T., Eds.; VCH: Weinheim, Germany, 2006. (16) Torriero, A. A. J.; Sunarso, J.; Forsyth, M.; Pozo-Gonzalo, C. Phys. Chem. Chem. Phys. 2013, 15, 2547. (17) Torriero, A. A. J.; Sunarso, J.; Howlett, P. C. Electrochim. Acta 2012, 82, 60. (18) Shiddiky, M. J. A.; Torriero, A. A. J.; Reyna-Gonzalez, J. M.; Bond, A. M. Anal. Chem. 2010, 82, 1680. (19) Shiddiky, M. J. A.; Torriero, A. A. J.; Zhao, C.; Burgar, I.; Kennedy, G.; Bond, A. M. J. Am. Chem. Soc. 2009, 131, 7976. (20) Zeng, Z.; Torriero, A. A. J.; Belousoff, M. J.; Bond, A. M.; Spiccia, L. Chem. Eur. J. 2009, 15, 10988. (21) Bard, A. J.; Faulkner, L. R. Electrochemical methods: fundamentals and applications, 2nd ed.; Wiley: New York, 2001. (22) Saveant, J.-M. Chem. Rev. 2008, 108, 2348. (23) Campbell, A. N.; Lam, S. Y. Can. J. Chem. 1973, 51, 4005. (24) Adenier, A.; Chehimi, M. M.; Gallardo, I.; Pinson, J.; Vila, N. Langmuir 2004, 20, 8243. (25) Mann, C. K.; Barnes, K. K. Electrochemical reactions in nonaqueous systems; Marcel Dekker: New York, 1970. (26) Portis, L. C.; Bhat, V. V.; Mann, C. K. J. Org. Chem. 1970, 35, 2175. (27) Deinhammer, R. S.; Ho, M.; Anderegg, J. W.; Porter, M. D. Langmuir 1994, 10, 1306. (28) Savéant, J. M.; Su, K. B. J. Electroanal. Chem. 1984, 171, 341. (29) Andrieux, C. P.; Hapiot, P.; Saveant, J. M. Chem. Rev. 1990, 90, 723. (30) Saveant, J.-M. Elements of molecular and biomolecular electrochemistry: an electrochemical approach to electron transfer chemistry; Wiley: Hoboken, NJ, 2006. (31) Torriero, A. A. J.; Siriwardana, A. I.; Bond, A. M.; Burgar, I. M.; Dunlop, N. F.; Deacon, G. B.; MacFarlane, D. R. J. Phys. Chem. B 2009, 113, 11222.

5739

dx.doi.org/10.1021/om4002318 | Organometallics 2013, 32, 5731−5739