Combining Nanosecond and Millisecond Time Scale Techniques

Article pubs.acs.org/JPCA

Combining Nanosecond and Millisecond Time Scale Techniques: Determination of Thermodynamic and Kinetic data of Primary Alkyl Amine Cation Radicals Hugo Cruz, Jose Luis Bourdelande, Iluminada Gallardo, and Gonzalo Guirado* Departament de Química, Univeristat Autònoma de Barcelona, 08193-Bellaterra, Barcelona, Spain ABSTRACT: Primary alkyl amines are one of the most commonly used and effective reagents in CO2 capture. Most of the amines used for CO2 capture are recycled, but a minor portion of the amines are degraded after one electron oxidation process, leading to highly toxic substances. The combination of the complementary information obtained from photoinduced electron transfer (flash photolysis) and heterogeneous electron transfer (electrochemistry) appears to be very attractive to fully characterize the electron transfer reaction mechanism of reactive species in general, as well as for determining important thermodynamic properties, such as standard potentials (E°) or pKa values. It is particularly difficult to determine these crucial data accurately in the cases of alkyl primary amines. Hence, in this manuscript we focus on the establishment of the several alkyl primary amines oxidation mechanism in organic aprotic solvents. In order to achieve this, this work combines information provided by flash photolysis (nanosecond), cyclic voltammetry (millisecond), and digital simulation (nanomile-second). Moreover, the accuracy of the E° values calculated using the nanosecond equilibrium method allows not only revising them, but also estimating new important thermodynamic data concerning the bond dissociation energies (BDEs) of ammonium cations (N+−H) and of the amine cation radicals (α-C−H), as well as their corresponding pKa values.

1. INTRODUCTION Primary alkyl amines are one of the most commonly used and effective reagents in CO2 capture, with that from natural gas being particularly well-known by the oil and gas industry.1 The most mature CO2 capture processes rely on the use of amine solvents to wash CO2 out of a gas mixture, such as flue gas. Most of the aliphatic amines used for CO2 capture are recycled, but a minor portion of the amines are either degraded or emitted into the air.2 The most important decomposition process takes place after the oxidation of the amines. The high oxygen concentration as well as other single electron oxidants (such as Fe 3+/ Fe 2+ or Cu+ to generate the oxide radical) are in the liquid holdup at the bottom of the absorber.3 This oxidative degradation process will not only increase the amine loss and amine waste, decreasing the capability of the system for CO2 capture, but will also lead to the formation of amine intermediates, which represent important negative risks on human health.4 As the chemistry of oxidative degradation is complex and not fully understood, fundamental research in this area would be highly desirable so as to design further applications of those systems. The main problem for fully establishing the amine oxidation mechanisms and characterizing the intermediates involved in this, is mainly related to standard potential values (E°). The high oxidation potential value of aliphatic alkyl amines gives rise to unstable intermediates after oxidation, probably due to the lack of the delocalization of the charge. © XXXX American Chemical Society

Aliphatic amine cation radicals can isomerize following an intramolecular hydrogen abstraction reaction under certain experimental conditions.5−23 It is known that extensive intramolecular hydrogen atom exchange takes place in gas phase after the initial formation of the amine cation radicals.5−16 The initial cation radical evolves to different distonic isomers, and the relative stability of those isomers have been established by computational thermochemistry studies, theoretical calculations, and mass spectrometry (Scheme 1).5−16 Electron paramagnetic resonance (EPR) studies of amine cation radicals at low temperature and low amine concentration in freon matrices also show intramolecular hydrogen atom abstraction of alkyl amine cation radicals to their distonic forms.17,18 At high concentrations of allylamines under the above-mentioned conditions, there is a notable change in the reaction mechanism. The amine cation radical changes, through either deprotonation or hydrogen atom abstraction reaction, from a neutral amine. In both cases, the radical (aminyl or αamino alkyl radical) is formed together with the ammoniumtype cation (R-NH3+) (Scheme 2).17,18 In this latest case (77 K and high amine concentrations), Knolle et al. proposed the preferential formation of the aminyl radical (Scheme 2-2a) Received: October 31, 2014 Revised: January 8, 2015

A

DOI: 10.1021/jp5109366 J. Phys. Chem. A XXXX, XXX, XXX−XXX

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The Journal of Physical Chemistry A Scheme 1

Scheme 2

Scheme 3

instead of α-amino alkyl radical (Scheme 2-2b, the most stable intermediate) through association of dimeric or small clusters of amine derivatives by hydrogen bonding. In a similar, more recent study, Belevskii et al. proposed the formation of the αamino alkyl radical (Scheme 2-2b) instead of the aminyl radical (Scheme 2-2a) in an ethylamine solution of spin trap irradiated between 77 and 293 K on the basis of EPR.18 Finally, it is important to note that few studies have been performed in solution. The use of N-halogenated amines under strong acidic

conditions may lead to the formation of cyclization products.19−25 The oxidation mechanism for secondary and tertiary amines has been widely studied in solution, although its electrochemical mechanism has been recently reported by some of us.26 It has been commonly accepted that the initially formed amine cation radical deprotonates, leading to the carbon centered α-radical. Later, this radical can evolve, either by oxidation to iminium cations or by dismutation to the related amines or enamines. Note that the cation radicals of secondary B


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The Journal of Physical Chemistry A Scheme 4

and tertiary alkyl amines show the formation of primary and secondary amines, respectively, while for primary amines the formation of the iminium cation will be the only reaction pathway followed, since it is not possible to form the enamine intermediate (Scheme 3).26−30 Since the study of electrochemical behavior of primary amines is more difficult than for secondary and tertiary amines, fewer results have been published.24−26,31 The propensity of those amines and their cation radicals to be attached onto the electrode surface makes it tedious and difficult to obtain reliable data, since polishing must be mandatory after each measurement. A one-electron irreversible wave is even obtained at high scan rates (up to 1000 V·s−1),31 which also makes it impossible to determine the E° values or other parameters from the conventional electrochemical measurements (milli- and low microsecond range).32−34 Thus, for shorter lifetime intermediates (less than 0.02 ms), the use of faster techniques is required in order to have access to reliable thermodynamic and kinetic data. On the other hand, it has recently been demonstrated35−37 that the determination of meaningful thermodynamic standard potential values (E°) can be successfully addressed using the nanosecond equilibrium. This method allows determining the electron transfer equilibrium constant by transient absorption spectroscopy using nanosecond laser excitation (flash photolysis), which is an ideal technique for performing these experiments (Scheme 4). The appropriate setup of the photo-oxidant system, based on photoinduced electron transfer, allows the E° values for short lifetime intermediates (100 ns) to be accurately determined. The determination of E° values for organic molecules is fundamental to understand the innumerable chemical and biological electron transfer processes.38 In this sense, their thermodynamic determination makes it possible to precisely determine bond dissociation energies (BDEs) and pKa’s of short-lived intermediates. One of the most “simple” methods was developed by Bordwell in the late nineties. The BDEs of acidic bonds of nitrogenated bases and their corresponding conjugated acids (RNH3+) can be estimated for several systems, such as primary amines, in acetonitrile using eq 1 (Scheme 5).39

Scheme 5

However, most of the BDE determinations of the N+−H bonds, which are calculated following eq 1, do not use an accurate standard potential value. The main reason is because it is not always possible to use electrochemical techniques for determining those E° values. In the case of primary amines, cyclic voltammetry experiments revealed that the oxidation of those primary amines is irreversible even at high scan rates. Thus, the E° value was replaced by anodic peak potential (Epa) values, which obviously introduces an uncertainty into these values. The constant (C) is empirically determined (C = 59.5 kcal/mol in acetonitrile), and the redox potential values are referred to as the ferrocenium/ferrocene couple (−0.288 V vs SCE).40 BDE R ‐ NH3+ = 1.37pK R ‐ NH3+ + 23.1E R°‐ NH2•+ /R ‐ NH2 + C (1)

It is also possible to calculate the pKa of the alkyl amine cation radicals using a similar strategy. The relationship between the thermodynamic properties of the ions and radicals derived from the amine (R-NH2) can be conveniently schematized in the following thermochemical mnemonic (Scheme 6).41 Thus, like in the above-described example, in order to determine the pKa and the BDEs of the α-C−H on the basis of a thermodynamic cycle, a high accuracy is mandatory in the experimental determination of the E° values (eq 2). ° − F(E R°‐ CH2 ‐ NH •+ /R ‐ CH2 ‐ NH2 − E H° + /H•) ΔG K°a = ΔG BDE 2

(2) C


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The Journal of Physical Chemistry A

for 10−15 min before equilibrium measurements were taken. Initially, separate solutions were prepared with each of the equilibrium components in order to record the spectra of their corresponding cation radicals. Increasing concentrations of the second component were then added to one of the cuvettes from a dioxygen saturated stock solution via calibrated micropipettes. All equilibrium measurements were made at 20 °C (room temperature). In general, equilibrium measurements were made 70−120 ns after the laser pulse. At the end of each equilibrium experiment, the laser power was checked to make sure that it had not changed during the course of the experiment. Equilibrium spectra were fitted to equilibrium equations using Igor Pro (version 4.07; Wavemetrics, Inc.). The oxidation potential differences given in Table 2 are averages of 2−3 independent determinations.35−37 Kinetic Measurements. The cation radicals were produced following the energy diagram shown in Scheme 2. Experiments using oxygen-saturated solutions were prepared by bubbling a dry stream of gas through the solution. Quenching rate constants were measured on static samples at 20 °C. Quenchers (alkyl amines) were added directly to the sample by a microliter syringe as aliquots of a suitable solution of the quencher in acetonitrile (ACN) or as a neat compound. C. Cyclic Voltammetry Experiments. An electrochemical conical cell equipped with a methanol jacket, which allows fixing the temperature at 20 °C by means of a thermostat, was used for the setup of the three-electrode system. For cyclic voltammetry experiments, the working electrode was in all cases a glassy carbon disk with a diameter of 0.5 mm. It was polished after each measurement by using a 1 μm diamond paste. The counter electrode was a Pt disk with a diameter of 1 mm. All the potentials are reported versus an aqueous saturated calomel electrode (SCE) isolated from the working electrode compartment by a salt bridge. The cyclic voltammetry apparatus was composed of a home-built solid-state amplifier potentiostat with positive feedback iR drop compensation and a Tacussel GSTP 4 generator. The voltammograms were displayed on a Tektronix (2212) instrument. Solutions were prepared using either ACN or DMF as a solvent, and they were purged with argon before the measurements, and argon was allowed to flow under the solution during the measurements. The concentration of the amines was ∼10−3 M, while the supporting electrolyte concentration was 0.1 M of TBABF4. D. Methodology for Quantum Chemical Calculation. The quantum mechanical calculations were performed within the framework of the density-functional theory (DFT), which is believed to be free from flaws, such as the spin polarization and spin contamination problem,42 in the calculation of organic molecules. The spin-unrestricted formalism has been used for solving Kohn−Sham DFT equations. The exchange-correlation functional in the Kohn−Sham DFT equation used the Becke three-parameter hybrid method using Lee, Yang, Parr correlation functional (B3LYP),43 which is very reliable for obtaining equilibrium geometries of radical ions. The 6311G(d,p) basis set was used to take the cation radical into account, which is defined for the first row as 6-311G, and for the second row as the Maclean−Chandler basis with d and p polarization functions. In order to optimize the geometries of amine cation radicals, the Berny algorithm was used with redundant internal coordinates, which is a fast and reliable algorithm to calculate the local minima by using the combined methods of rational function optimization (RFO) and the

Scheme 6

In this sense, the combination of the complementary information obtained from photoinduced electron transfer (flash photolysis) and heterogeneous electron transfer (electrochemistry) appears to be very attractive to fully characterize the electron transfer reaction mechanism of reactive species in general, and particularly crucial in the cases of alkyl primary amines. The aim of this article is the establishment of the alkyl primary amines oxidation mechanism. To achieve this, this work combines information provided by flash photolysis (nanosecond), cyclic voltammetry (millisecond), and digital simulation (nanomile-second) to determine the oxidation mechanism of the alkyl primary amines. Moreover, the accuracy of the E° values calculated using the nanosecond equilibrium method not only allows them to be reviewed, but also allows one to estimate new important thermodynamic data concerning the BDEs of ammonium cations (N+−H) and of the amine cation radicals (α-C−H), as well as their corresponding pKa values.

2. MATERIALS AND METHODS A. Chemicals. Anhydrous acetonitrile (ACN) and N,Ndimethylformamide (DMF) stored in an inert atmosphere and molecular sieves were purchased from Across. The reference compound hexamethylbenzene (HMB) and substrates (alkyl primary amines from 1−12) of the highest available purity were purchased and recrystallized several times before use (solids) or fractionally distilled (liquids). Tetrabutylammonium tetrafluoroborate (TBABF4, Fluka, puriss.) was used without further purification. B. Laser Flash Photolysis Experiments. Nanosecond Laser Flash Photolysis experiments were performed using a LKS60 instrument from Applied Photophysics. Pulses of ca. 9 ns and energies of 5−7.5 mJ were provided by a Q-switched Nd:YAG laser (Spectron Laser Systems, UK). All experiments were carried out under oxygen and in acetonitrile in quartz cells. The transient spectra were obtained by recording the transient decays at different analysis wavelengths. Typically, the data from 4−8 laser pulses at 355 nm (absorbance 0.5) were averaged prior to computer processing. In addition, global analysis of the complete kinetic data set of the decays was carried out using GLint, which is a form of global analysis developed by Applied Photophysics Ltd. that uses the Marquardt−Levenberg algorithm and the fourth order Runge−Kutta numerical integration. Equilibrium Measurements. N-Methylquinolinium hexafluorophosphate (NMQ) solutions (absorbance −0.5 at 355 nm) containing 1 M toluene were prepared in anhydrous acetonitrile. Experiments were conducted in quartz cuvettes equipped with high-vacuum stopcocks carrying serum caps. Solvent-saturated dioxygen was bubbled through the solutions D


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Figure 1. Time-resolved absorption decay observed at 495 nm for excitation of NMQ hexafluorophosphate with toluene, as cosensitizer, in oxygenpurged ACN solutions: (a) HMB (24.93 mM) + dodecylamine, 8 (0.20 mM); (b) HMB (50.22 mM) + tert-butylamine, 12 (0.80 mM); (c) overlay of the two decays. The decay is mainly due to the cation radical of the HMB. Note that two fast components at early times (insets of panels a and b and to the left of the dashed line (panel c)), which are not included in the fitting. Those decays are due to reaction of NMQ radical with oxygen (the first one, less than 50 ns) and to the equilibrium process (the second one, less than 100 ns), thus the contribution is minimal on this time scale.

linear search process. The B3LYP analytical second derivatives of potential energy with respect to the nuclear Cartesian coordinates were used for calculating the vibration frequencies. In the case where imaginary frequencies occur in the calculation, the molecular structure will be modified to clear the imaginary frequencies. By doing this, all the amine cation radical can reach its energy minimum configuration. The solvent effect was studied by the polarization continuum model.44 Acetonitrile (ε = 36.64) has been employed as a solvent. All the electronic calculations were carried out by the Gaussian 03 package.45

a cosensitizer (note that the cosensitizer must have a higher oxidation potential than either HMB (reference compound) or amine). The cosensitizer is used in high concentration (1 M) to ensure that NMQ excited state reacts with PhCH3 instead of HMB or amine, leading to initial formation of the NMQ•/ PhCH3+• geminate radical/cation-radical pair (diffusion control electron-transfer reaction (∼1010 M−1 s−1)), which undergoes efficient separation (∼109 s−1).46−49 The “free” toluene cation radical is a powerful one-electron oxidant (2.26 V vs SCE) that effectively and irreversibly oxidizes the electron donors (HMB and/or amine) that are present at relatively low concentration (e.g., ∼1−20 mM) in solution with a bimolecular rate constant between ∼109−1010 M−1 s−1 (diffusion control electron transfer). Thus, in a pulsed laser experiment, each laser pulse produces the same concentration of PhCH3+• and, consequently, the same total concentration of HMB+• and/or amine+•. The reduced form of the photo-oxidant, NMQ•, can be rapidly scavenged with dioxygen leading to O2−•, which does not have interfering absorptions in the UV−vis region, where many organic cation radicals strongly absorb (>50 ns). Transient spectra are typically recorded between 60 and 120 ns after the laser pulse, which allows establishing the redox equilibration before significant decay of the cation radicals. Hence, primary amine cation radicals were generated using photoinduced electron transfer. The overall photo-oxidation system (Scheme 4) is based on direct irradiation of an anhydrous oxygen-purged acetonitrile sample at 355 nm. Each sample contains A+, electron acceptor (N-Methylquinolinium hexafluorophosphate in a suitable concentration, absorbance c.a. 0.5 at 355 nm), the cosensitizer (toluene, 1 M) and the reference (HMB) and substrate compound (∼10−2−10−3 M amine) for the generation of a mixture of HMB and amine

3. RESULTS AND DISCUSSION Free Cation Radical Amine Chemistry. Although described in detail elsewhere,46−49 the method is briefly outlined here for clarity. The sensitizer used in this study is NMQ, a positive charge electron acceptor (A+, Scheme 4). Toluene (C) acts as a cosensitizer. Irradiation of NMQ (A+) generates the corresponding NMQ excited state (A+*), which reacts with toluene (C) via exothermic electron-transfer (ket) to generate the geminate radical/cation radical pair (A•/C+•). The oxidation potential of the cosensitizer is chosen so that diffusional separation (ksep) competes efficiently with return electron transfer (k−et), resulting in a high yield of separated ion radicals in acetonitrile.47 The separated toluene cation radical (co-donor) reacts with the substrate to generate the corresponding cation radical (in this case either HMB or amine). NMQ is excited at a sufficiently long wavelength (343 nm) where there is no competitive absorption by the electron donors. The singlet excited state of NMQ is a powerful oneelectron oxidant (2.7 V vs SCE) allowing toluene to be used as E


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The Journal of Physical Chemistry A cation radicals. The HMB cation radical (HMB+•) absorbs between 400 and 500 nm, and it is stable enough to be followed by nanosecond transient absorption spectroscopy. Despite the alkyl primary amines extinction coefficient being zero at the UV−visible wavelengths range, the cation radicals evolution can be followed by monitoring the absorption of the HMB+•, λmax = 495 nm, and at different delay times (sufficient to reach the equilibrium and after reaching it). At the shorter times after the laser pulse (less than 120 ns), both cation radicals are in equilibrium with their corresponding neutral compounds. Appropriate analysis of the recorded spectra at the equilibrium time allows a meaningful standard potential of the primary alkyl amines to be accurately determined, as previously published.37 After this equilibrium time, it is reasonable to assume that both components evolve to their corresponding oxidation products.35−37 As a result, a complex transient kinetic trace is recorded at 495 nm (Figure 1). A contribution of the N-methylquinolium radical (λmax = 540 nm) within 40 ns in oxygen purged solution followed by a plateau zone due to the electron transfer equilibrium between the HMB cation radical and the amine (c.a 100 ns).37Moreover, a later longer-lived intermediate, decaying on a microseconds time scale, can be also distinguished. This longer-lived intermediate is assigned to the HMB cation radical reactivity, which is confirmed by the transient absorption spectra obtained.50 A plot of the pseudo-first-order rate for this decay as a function of the amine concentration demonstrates the base reactivity of the amine. A typical experiment is illustrated in Figure 2.

transfer equilibrium becomes shorter as the concentration of amine is increasing. This effect is in agreement with the previous observation, and confirms the existence of a protontransfer from the amine cation radical (Scheme 7). Scheme 7

Table 1. Standard Potentials (E° in V vs SCE), Bond Dissociation Energies (BDE’s), pKa’s, Structural and Solvation Effects on the Second-Order Rate Constants (k2) for the Reaction of Hexamethylbenzene Cation Radical (HMB+•) with Alkyl Primary Amines, 1−12, as Bases in ACN at 20 °C amine

λmax,a nm (HMB+•)

n-propyl (1) n-butyl (2) n-pentyl (3) n-hexyl (4) n-heptyl (5) n-octyl (6) n-decyl (7) n-dodecyl (8)

495

isopropyl (9) isobutyl (10) sec-butyl (11) tert-butyl (12)

495

k2 (M−1 s−1)b

E° (V) (amine+• /amine)c

Unbranched Acyclic Amines 6.0 × 106 1.517(±0.003)

BDE (kcal· mol‑1)

pKa+•d

93.1e

8.50

495 495

5.5 × 106 4.5 × 106

1.509 (±0.003) 1.500 (±0.002)

94.0e 95.5e

9.30 10.6

495 495

4.2 × 106 3.0 × 106

1.489 (±0.003) 1.496 (±0.004)

97.5f 96.3f

12.2 11.2

495 495 495

2.1 × 106 2.2 × 106 2.1 × 106

1.501 (±0.004) 1.503 (±0.005) 1.500 (±0.005)

95.5f 95.1f 95.6

10.5 10.2 10.6

Branched Acyclic Amines 7.5 × 105 1.538 (±0.005)

88.5g

4.78

495

3.0 × 105

1.515 (±0.004)

93.1f

8.54

495

7.5 × 105

1.517 (±0.003)

92.7f

8.21

495

2.5 × 105

1.526 (±0.003)

93.7h

12.0h

a Absorption maximum of HMB+• in the visible region. bRate constant for the decay of the HMB+•. cStandard potential value measured using the nanosecond equilibrium method from ref 37. dCalculated from Scheme 6 (eq 2). eData taken from ref 52. fEstimated from Figure 6. g Data taken from ref 53. hDifferent reactivity is showing from tertbutyl amine since no α-H is present in the molecule refs 51 and 53.

Figure 2. Plot of the pseudo-first-order rate constant for the decay of HMB cation radical (HMB+•) formed using the photo-oxidation system depicted in Scheme 4 in dried oxygen-purged acetonitrile solutions (see text), as a function of the concentration of added amines (n-butylamine (2) (solid line), n-dodecylamine (8) (dotted line), and tert-butylamine (12) (dashed line)).

The second-order rate constants (k2) were measured for several amines (Table 1) for the proton transfer reaction between the neutral amine with HMB+•. The k2 values for shorter alkyl amines decrease with the increasing electrondonating ability of the alkyl chain. Moreover, the rate constant decreases by a factor of 2 for heptyl (5), octyl (6), decyl (7), and dodecylamines (8), presumably due to steric hindrance around the amino group because of the fold ring-shape conformation adopted by those amines. As expected for the amine cation radical proton-transfer mechanism, the rate constants are very sensitive to the steric nature of the alkyl

Hence, absolute rate constants for the second order of primary amine cation radicals as a function of the concentration of added amines are estimated from the full fit of the decay of the HMB cation radical (kobs). It is well-known that the rates of proton transfer from HMB to a series of substituted pyridines are strongly related to the base strength. Since the pKa value of HMB cation radical49−53 and the pKa of primary ammonium cations is ca. 18 in ACN,50 one can expect an exothermic and reasonable fast proton transfer from the HMB to neutral amines. It is noticeable that the time range where the electron F


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The Journal of Physical Chemistry A substituents in the nitrogen. As shown in Table 1, changing the alkyl group, R, from methyl to isopropyl (9−10) or tert-butyl (12) also causes the rate constant to decrease ca. 1 order of magnitude. The conformation of those alkyl amine cation radicals will be deeply explored in the Theoretical Modeling section. It is assumed that the rate constant value, k2, should be similar to the reaction of the neutral amine with their cation radicals (k‑H+, Scheme 8 and Table 2), since, as in the case of

Table 3. Kinetic and Thermodynamics of Cation Radicals

Scheme 8

Table 2. Electrochemical Data for 1−12 in DMF at 20 °C amine

Epa (V)a


1.35d (1.49)f 1.36g 1.34d (1.45)f 1.36g 1.34e

isopropyl (9) isobutyl (10) sec-butyl (11) tert-butyl (12)

1.40

1.35e 1.34e 1.35e

ΔEpa (mV)a

n

αb

k‑H+ (M‑1 s‑1)c

mechanismd

Unbranched Acyclic Amines 170 1 0.3 6.5 × 106

EC2

110 156

1 1

0.4 0.3

5.5 × 106 4.5 × 106

EC2 EC2

110 152

1 1

0.4 0.3

4.0 × 106 3.5 × 106

EC2 EC2

125 128 115

1 1 1

0.4 0.4 0.4

2.0 × 106 2.5 × 106 2.0 × 106

EC2 EC2 EC2

Branched Acyclic Amines 80 1 0.6 8.0 × 105

EC2

1.38

85

1

0.6

3.5 × 105

EC2

1.38

88

1

0.5

7.5 × 105

EC2

1.40f,g

95

1

0.5

2.0 × 105

EC2

amine

pKaN+‑Ha


18.20a

isopropyl (9) isobutyl (10) sec-butyl (11)

BDE (kcal· mol‑1)b,c

ΔG°d (kcal· mol−1)

Unbranched Acyclic Amines 93.1b −13.49

ΔG‡e (kcal· mol−1)

ΔG°‡f (kcal· mol−1)

6.30

11.99

18.26a 18.30a

94.0b 97.5b

−12.47 −10.78

6.35 6.47

11.65 11.11

18.34a 18.35

97.5 96.3

−8.53 −9.94

6.51 6.71

10.26 11.03

18.35 18.35 18.35

95.5 95.1 95.6

−10.87 −11.33 −10.75

6.91 6.89 6.91

11.62 11.77 11.57

7.51

15.15

17.96a

Branched acyclic amines 88.5 −18.35

18.20c

93.1c

−13.45

6.71

12.41

17.92c

92.7

−13.51

7.51

13.30

a

Data taken from ref 55. bEstimated values from Figure 6. cCalculated values using thermodynamic cycle (Scheme 5). dΔGo = 2.3 RT (pKa+• − pKaN+‑H) eCalculated values using eq 3. fCalculated values using eq 4.

mentary character of the techniques and the covering of nearly ten magnitudes of time scale, both techniques have been rarely used together for elucidating reaction mechanisms.54 Electrochemical Behavior of Alkyl Primary Amines. The electrochemical study of primary alkyl amines is a real challenge. The one electron oxidation of those amines presumably generates radicals, which are rapidly attached to the electrode surface. It has been demonstrated by some of us26,55 that in this oxidative grafting mechanism, the α-amino alkyl radical is in equilibrium with the aminyl radical. The aminyl radical is the key intermediate in the attachment reaction, since the amino group binds to the electrode surface through a nitrogen bond (Scheme 9). The combined use of

At 0.1 V·s−1 bExperimentally determined (ΔEpa= 47.7 mV/α) Obtained using DigisimR and fixing the experimental parameters (E° values from Table 1 and α). dAn electrochemical reaction (E) followed by a Chemical one (C) eThis work. fData from refs 24 and 25. gData from ref 26. a c

Scheme 9

HMB cation radical, the pKa’s of the amine cation radicals of primary amines will not be higher than 14,51,54 (Table 3), and one can expect an exothermic and reasonable fast proton transfer from the amine cation radicals to neutral amines (Scheme 8). This assumption will be later demonstrated in terms of cyclic voltammetry, since in the electrochemical experiments the alkyl amine will be the only reactant present in solution. Photoinduced electron-transfer (Flash photolysis) and electrochemistry techniques provide complementary information for the study of reaction mechanisms involving electrontransfer. While flash photolysis allows characterizing short-lived intermediates by UV−visible transient spectroscopy and the determination of the kinetics starting in the low-microsecond range, electrochemical techniques are the most commonly used techniques at shorter time scales for the characterization of those intermediates formed upon heterogeneous electron transfer reaction and for the determination of the mechanism coupled to this first electron transfer. Despite the comple-

cyclic voltammetry, X-ray photoelectron spectroscopy (XPS), and infrared reflection absorption spectroscopy (IRRAS) allowed the attachment of the organic group to be confirmed. 1. Cyclic Voltammetry at Low Scan Rates. The electrochemical investigation of these amines is much more difficult, due to their strong propensity to attach on the electrode surface,26−31 so careful polishing of the electrode between every measurement is mandatory. Alkyl branched or unbranched acyclic amines (1−12) exhibit a broad irreversible oxidation peak (Epa) at ca. 1.4 V vs SCE in ACN or DMF solutions using a glassy carbon electrode as a working electrode at slow scan rates (v = 0.1 V·s−1). Typical cyclic voltammograms are shown in Figure 3, and their corresponding results are summarized in Table 2. The peak height indicates the exchange of nearly one G


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parameters of the electron transfer processes surveyed, as well as those of any homogeneous chemical reactions coupled to them, it is necessary to use numerical approximation methods.62−65 Cyclic voltammogram simulations of electron transfer reactions coupled to homogeneous chemical processes is a commonly used strategy for fitting kinetic and thermodynamic parameters in the establishment of the reaction mechanisms. In the case of alkyl primary amines, the transition between the irreversible behavior (cyclic voltammetry at low scan rates) and the hypothetical reversible behavior at high scan rates can only be obtained by digital simulation techniques, since the use of ultramicroelectrodes and ultrafast potentiostat are not possible due to the attachment of the reactants and intermediates in their surfaces. However, we can use the data obtained by laser flash photolysis in nanoseconds (E° of primary alkyl amines) to determine the kinetic constant value (k) of the chemical reaction coupled to the electron-transfer for the digital simulations of the cyclic voltammograms in the microsecond time scale. The simulation of the cyclic voltammograms at low scan rates using the data obtained photochemically allows checking those theoretical voltammograms with the experimental ones in a first approach. The fact that the standard potentials are highly accurate values makes it possible to fit the rate of the chemical reaction coupled to the electron-transfer (Table 2). Finally, when the simulation is performed at high scan rates, no effect of the coupled chemical reaction is observed on the voltammetric response, since the time scale of the experimental technique used for the digital simulation is much faster than the lifetime of the electrogenerated species. A comparison between the rate constant values obtained for k2 (Table 1) and for k‑H+ (Table 2) shows that they are almost identical, which verifies that the rate of both processes (Schemes 7 and 8) are the same. Theoretical Modeling. To rationalize the observed behaviors, a series of molecular modeling calculations were performed to determine both the geometry and the stability of the alkyl amine cation radicals. The cation radical geometries were determined by a full optimization of the conformation using the B3LYP/6-311G(d,p) density functional basis set. The solvent effect (ACN, ε = 36.64) was studied by the polarization continuum model.44 In the first step, the energy of the cation radicals in the gas phase was calculated. The result of n-amines (n = 1−8) in gas phase shows that the energy of n-amines decreased with the increase of carbon atom number, which means that the ring-shape structure, like hexylamine, is more stable than the open structure, like propylamine. The nbutylamine (2) and n-pentylamine (3) have open-ring structure trends to reach the energy minimum (Figure 4, Table 4). When the solvent effects were considered using the polarization continuous model, the same tendency was found. The solvent effect can be studied by the difference between the minimum energies in gas and solvent phases. Figure 4 shows that the solvation energy of the amines decreases for propylamine (1), butylamine (2), pentylamine (3), hexylamine (4), decylamine (5), and dodecylamine (6), in which the carbon atoms in the amines increase consequently. This trend was interrupted by heptylamine (7) and octylamine (8). The results suggest that, depending on the conformation adopted for the cation radical in solution, the oxidation potential value is different. The fact that longer amines have the same oxidation potential leads one to think that they should be folded in the

Figure 3. Cyclic Voltammetry (at 0.1 V·s−1 ; scan range: 0.00/1.60/ 0.00 V) of an alkyl primary amine argon-purged DMF + 0.1 M TBABF4 at 20 °C solution: (a) n-dodecylamine (8), 7.5 mM, and (b) isopropylamine (9), 9.4 mM. Working electrode: glassy carbon disk of 0.5 mm diameter; counter electrode: platinum disk of 1 mm diameter; reference electrode: saturated calomel electrode (SCE).

electron per molecule by comparison with ferrocene under the same experimental conditions.56 The interest of those experiments at slow scan rate is to show the electrochemical oxidation mechanism, and to characterize the first reaction coupled to the electron transfer by the study of the peak potential with the scan rate. The fact that the amines and amine cation radicals have a strong tendency to modify the electrode surface, and that the monoelectronic oxidation waves are irreversible up to 100 V·s−1 makes this study particularly unfavorable, so no direct mechanistic information can be obtained. As reversibility cannot be reached, it is not possible to determine the other parameters exclusively based on electrochemical techniques, and analysis of the peak width should be done. It is possible to determine the transfer coefficient (α) from the width of the peak (Table 2), being between 0.3 and 0.6. Thus, for primary alkyl amines, the cation radicals are much more unstable than those for the other amines, their lifetime being at least 0.2 ms,26 but the values of α, always close to 0.5, do not indicate an electron transfer concerted with the follow-up reaction.57−61 Despite the fact that the amine cation radicals are not able to be detected by electrochemical experiments, flash photolysis experiments confirmed their existence, which is also in favor of sequential mechanism (electron transfer (E) + chemical second order reaction (C2)). It is remarkable that, in the case of tertbutylamine,28 a different mechanism should be operating, since the α-C−H bond cannot be broken. In addition, this is the only primary alkyl amine that is not attached to the electrode surface.26 2. Digital Simulations at Low and High Scan Rates. In order to determine the kinetic and/or thermodynamic H


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the amine cation radicals to be completely characterized after the monoelectronic oxidation of the corresponding amine in ACN.37 The high precision of the method (experimental error less than 6 mV) makes it possible to determine thermodynamically meaningful data, providing some insight into the origin of the different E0, which had never been done until our previous work. For short chain alkyl primary amines (Table 1, 1−4), the average differences are c.a. 0.19 kcal/mol for each CH2 unit. This variation of the standard potential is rationalized in terms of inductive effect, since the stabilization of the amine cation radical by solvation of the amino group should be similar in all these cases. On the other hand, the standard potentials of 5 and 6 in ACN are indistinguishable within experimental error, while for longer unbranched alkyl primary amines (5−8), the E0 rises slightly ca. 0.16 kcal/mol. We suggest that these results could be the consequence of strain of the alkyl chain, which obviously causes an inhibition of solvation around the amino group by a solvent exclusion mechanism (Figure 5). The comparison of the E0 values between n-butylamine, 2, and tert-butyl amine, 12, in CH3CN and CH2Cl2 shows that the cation radicals of both amines are stabilized at 3.6 and 2.8 kcal/mol, respectively. The most plausible explanation is related to differences in the solvation environment around the NH2 group. Substrates 9−12 would also present steric solvation inhibition relative to 1 or 2, respectively, due to the branched alkyl group. It seems that each CH3 branched to the α carbon blocks the accessibility of the solvent to the polar group, which implies a 0.21 kcal/mol loss of stabilization solvation energy per methyl group bonded to the α carbon (Cα−(CH3)n, where n = 1,2 or 3)). The thermodynamic data obtained by laser flash photolysis support the electrochemical characteristics and mechanism obtained on the same cation radicals by electrochemistry. The electron transfer reaction is followed by a proton transfer reaction, the rate of this chemical reaction being strongly dependent on the amine cation radical conformation. From Nanosecond to Microsecond Time Scale: Determination of Microscopic Dissociation Constants. Deprotonation from the amine cation radicals to neutral seems to be an exothermic and fast reaction, as can be deduced from the previous flash photolysis determinations. To investigate the overall reactivity of amine cation radicals, it is particularly interesting to combine the results of flash photolysis, cyclic voltammetry-digital simulations experiments, and theoretical calculations. Indeed, the first one, flash photolysis, allows the short-lived intermediates to be characterized down to 100 ns and thus, the characterization of their initial chemical reactivity. While the second one, electrochemistry, permits the reactivity at a longer time range to be characterized by the determination of the mechanism of the reaction following the first electronic transfer.57−61 Besides, these results confirm that in solution, and in our amine concentration solutions, the initially formed amine cation radical follows a second-order reaction (deprotonation reaction) instead of an isomerization reaction (intramolecular hydrogen atom exchange, first order reaction). These results are in good agreement with previously published results, where the initially amine cation radical evolves to different distonic isomers at low temperatures, low concentration, and Freon matrices, where the amine molecules are isolated and the deprotonation reaction can be avoided.17,18 When the concentration of amine and the temperature are increased in liquid solutions,18,26 or even in Freon matrices,18 the initially generated amine cation radical deprotonates, yielding the corresponding radical and an ammonium type cation. Under

Figure 4. Structural and Solvation Effects of Unbranched Acyclic Amines (1−8).

Table 4. Total Amine Energies (Hartree) at the 6-311G (d,p) Levela total energies (Hartree)

amine n-propyl (1) n-butyl (2) n-pentyl (3) n-hexyl (4) n-heptyl (5) n-octyl (6) n-decyl (7) n-dodecyl (8) sec-butyl (11) tert-butyl (12)

gas phase (B3LYP) −174.2288

acetonitrile (B3LYP)

ΔEnergies (GasSolvent)

Unbranched Acyclic Amines −174.3297 0.1000

ΔEnergies (GasSolvent) (kcal· mol‑1) 62.8

−213.5554

−213.6540

0.0986

61.9

−252.8839

−252.9801

0.0963

60.4

−292.19889

−292.2956

0.0967

60.7

−331.54929

−331.6373

0.0880

55.2

−370.8741

−370.9618

0.0877

55.0

−449.4964

−449.6069

0.1105

69.3

−528.1465

−528.2566

0.1101

69.1

Branched Acyclic Amines −213.6560 0.0948

60.0

−213.5612 −213.5626

−213.6560

0.0939

59.5

a

All the optimization calculations and frequency analyses in gas phase were performed at the B3LYP/6-311G (d,p) level. The solvent effect was performed at the B3LYP/6-311G (d,p) level, and the PCM model was used to study this effect. The minima structures are ensured by the frequency analysis (positive frequencies).

same way, so that the solvation energy around the NH2 polar group in a polar solvent should be the same (Figure 5).53,66,67 Hence, careful investigation of the oxidative behavior of alkyl primary amines indicates that the first intermediate is a cation radical that deprotonates to give a radical following a sequential reaction pathway. Thus, based on considerations of stability of radicals in solution, the reaction path involves the formation of a cation radical, in agreement with the electrochemical data. After that a deprotonation reaction coupled to the electrontransfer is observed according to the photochemical, electrochemical, and simulation data. Finally, taking into account previous studies,26 an isomerization reaction of the α-amino alkyl radical can take place (Scheme 10). Inductive vs Solvation Effects in Primary Alkyl Amines on the Standard Potential Values. The Eo values were obtained using the nanosecond equilibrium method allowing I


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Figure 5. Representation of optimized structures of amine cation radicals in solution: (a) n-butylamine (2), (b) n-hexylamine (4), (c) ndodecylamine (8), (d) iso-butylamine (10), (e) sec-butylamine (11), and (f) tert-butylamine (12).

(3), and isopropylamines (9) have been reported for the α-C− H bonds.52,53From our data of E° and reported BDEs values (Table 1), it is possible to demonstrate that there is a linear relationship between the standard potential of each primary alkyl amine branched or unbranched and the BDE (Figure 6). That good correlation permits us to estimate the BDE value of an unknown primary alkyl amine, such us n-dodecylamine (8), by interpolating the E° value. Once, the E° and the BDE values are obtained, it is possible to get pKa of the unknown substrates by applying eq 2. Note that the easier oxidizable one leads to the more acidic cation radical (Figure 7), which obviously will present the lowest BDE value. From the determination of the acid dissociation constant of primary monoamine cation radicals in ACN, it is possible to determine, by applying the above-mentioned thermodynamic cycle (Scheme 6, eq 2), the bond dissociation enthalpies of the whole amine group within an accuracy of 2 kcal/mol.

Scheme 10

our experimental conditions, the α-amino alkyl radical will be the first amino radical intermediate formed, Although according to previous works published in the literature,24−29 it can be in equilibrium with its tautomer (Scheme 10). Despite the fact that the determination of BDEs of amines has been the subject of intensive experimental and theoretical studies, only the values of n-propyl (1), n-butyl (2), n-pentyl

Figure 6. Correlation between E°’s and BDE’s of primary alkyl amines. J


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Figure 7. Correlation between E°’s and cation radical pKa’s of alkyl primary amines.

Figure 8. Plot of the variation of the intrinsic barrier, ΔG°‡, of the deprotonation of the amine cation radicals with the pKa and BDEs.

required work for the formation of the precursor (wR) and successor (wP) complexes from the reactants at infinite distance is negligible since there is no electrostatic charge interactions, it is also possible to determine the ΔrG°‡ value (eq 4)

A Bronsted plot for unbranched amines shows a small slope that is usually associated with a small extent of the proton transfer in the transition state.68 The data fits reasonably well to a linear relationship when nonsterically encumbered bases have been used.69 The fitting of the log k with the driving force experimental data follows eq 3 (thermodynamic formulation of the transition state theory): k −H +

⎛ −Δ G‡ ⎞ r ⎟⎟ = Z exp⎜⎜ ⎝ RT ⎠

2 Δr G ° ⎞ 4Δr G°‡ ⎛ ⎜ ⎟ Δr G = ⎜1 + ⎟ 4 ⎝ 4Δr G°‡ ⎠ ‡

(4)

Note that ΔrG° is the driving force of the process, so it can be calculated between pKa(cation radicals) and pKa(neutral amines) (Tables 1 and 3). The ΔrG°‡ values are a measure of the acidities of the intrinsic amine cation radicals; at first sight the intrinsic barriers do not follow the “standard” acid behavior since the activation barrier is not related with the pKa of the acid, in the sense that

(3)

where k is the deprotonation rate constant, Z the collision frequency 3 × 1011 M−1 s−1, and ΔrG‡ is the activation energy of the free-barrier. Using the resulting quadratic Marcus-type activation-driving force relationship, and considering that the K


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The Journal of Physical Chemistry A the lower pKa does not correspond to the smaller intrinsic barrier. The current cation radicals do not fall in this category, as it is shown in Figure 8. The less acidic cation radical (1+•) is the fastest to deprotonate, while the most acidic is the slowest (12+•), which is the contrary to what is expected. Moreover, the large intrinsic barrier energy values (11.53 kcal/mol (0.5 eV)) do not increase with the pKa. This intrinsic barrier is practically the same for all the amines, and so is the homolytic dissociation energy (BDE(AH•+/A+ + H•) = 92.2 kcal/mol (4.0 eV)) while the pKa varies from 4 to 12 in the series. Thus, in this case the intrinsic kinetic acidity does not behave parallel the thermodynamic acidity. The same observations appear to be applying to another series of cation radicals, such as polymethylbenzenes and NADH analogues.50,59 Taking into account the model developed by Savéant for the dynamics of cation radical proton transfer in the study of NADH analogues, it appears that a concerted electron-H atom transfer is the mechanism operating in this case for the proton transfer reaction. Finally, it is remarkable that the homolytic bond dissociation energies of H−N bonds in the conjugated acids of nitrogen bases can also be conveniently addressed by our experimental values using the thermodynamic cycle depicted in Scheme 5. However, for tert-butylamine (12) a discrepancy of 21 kcal/mol in its BDE value is reported, in other words a difference from 93.3 to 114.3 kcal/mol, depending on the experimental method used (Table 5). The first determination of the experimental

Figure 9. Structural effects of unbranched acyclic primary amines on the relative acidities (ΔpKa = pKa+• − pKaN+‑H ) with the standard oxidation potentials (E°) in ACN vs SCE.

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CONCLUSIONS The physical chemistry, oxidation mechanism, and cation radicals of alkyl primary amines have been widely investigated combining photochemical and electrochemical studies. Those studies allowed: (i) the identification of the alkyl amine cation radicals for primary amines, and for the first time (ii) the establishment of the conformation of those cation radicals and their relationship with the E° and (iii) the establishment of the kinetic constant values of the chemical reaction linked to the first electron transfer. Moreover, the high accuracy of the experimental values in the study of representative alkyl primary amines BDEs has provide upward revisions for α-C−H and N− H BDEs as well as an improved understanding of their structure/property relationship. The rate of amine cation radical proton transfer is satisfactorily fitted with quadratic Marcus-type equation. The intrinsic kinetic acidity is not correlated with the pKa of the cation radicals, which seems to indicate, like previous reported studies, that the proton transfer appears to be a concerted electron-H atom transfer. The overall methodology developed allows accessing useful and indispensable information for understanding the insights of fundamental chemical and biochemical electron transfer processes that would undoubtedly help in the design of new amine systems, especially in the field of CO2 storage and reversible capture.

Table 5. BDE’s Values of sec- and tert-Butylamine (11 and 12) BDEa sec-butylamine (11) tert-butylamine (12) a

113.6 114.4

107.1 106.0

BDEb

BDEc

BDEd

116.6 115.6

93.3

112.4 112.9

Data from ref 39. bData from ref 70. cData from ref 52. dThis work.

values of BDE tert-butylamine using photoacustic calorimetry55 seems to be underestimated in comparison with the previously reported ones by Bordwell39 using irreversible oxidation potentials instead of standard potentials and the pKa values previously published in the literature for the application of the method described in Scheme 3. In this sense, the development of high-level theoretical calculations to re-examine the N+−H BDEs is a growing area, and the solvent effects were introduced by the polarized continuous model, as we previously mentioned. Depending on the level of theory, the estimated values can oscillate from 101 to 116 kcal/mol, again showing a large discrepancy. In this sense, it seems worthwhile to revise N+−H BDEs since the accurate determinations of E° lead to upward revisions of those energies. The application of our experimental data to the thermodynamic cycle depicted in Scheme 5 allows to unequivocally establish the experimental value for the N+−H BDEs in the tert-butyl amine 112.9 kcal/ mol. This methodology is also employed to determine the N+− H BDEs of those amines in which the pKa was previously determined (Table 3, Figure 9). It is noticeable, for unbranched alkyl amines, that the standard potential of the amines is linearly related with the ΔpKa. Thus, the pKa in ACN is estimated, and consequently the N+−H BDEs is also determined for the rest of the amines (Table 4).

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AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS Financial Support by the Universitat Autònoma de Barcelona (APOSTA Research Program), MICINN the Spanish Ministry of Science and Innovation (project CTQ2009-07469), and MINECO the Spanish Ministry of Economy and Competitiveness (Project CTQ2012-30853) is gratefully acknowledged. The authors are pleased to acknowledge Professor J. P. Dinnocenzo, Department of Chemistry, University of Rochester (NY) for his help and technical advice, and Dr. C. Wang, previously at the Department of Chemistry & Biochemistry (University of Missouri-St. Louis), for theoretical calculations. L


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