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A simple treatment of the fundamentals of solution-phase voltammetry is followed by six examples of porphyrinoids highlighting how structural changes of the redox species influence formal reduction potentials. Electronic communication among different molecular groups is demonstrated, and it is shown how the correct solvent and supporting electrolyte choice may result in observing 17 of a possible 18 redox processes in a cadmium triple-decker phthalocyanine. The difference between outer- and inner-sphere electron transfer processes is used to introduce adsorption effects of redox species on electrodes. This forms the bases of electrocatalysis, and a short theoretical introduction to electrocatalysis is presented. The reader is then eased into electrocatalysis concepts through a discussion of electrocatalyzed alcohol oxidation and sulfite sensors with adsorbed porphyrin polymers. Covalent binding or adsorption to the electrode or solution-phase electrocatalysts are demonstrated. Electrocatalysis in fuel production or energy storage systems concludes this chapter. CO2 reduction by an Fe0 porphyrin, an octaalkoxylated Co phthalocyanine, which aggregates less, and covalent– and metal–organic framework complexes of Co porphyrinoids, are considered. Hydrogen evolution by Ni porphyrins, oxygen reduction to either H2O or H2O2, and H2O oxidation to O2 utilizing carefully chosen porphyrinoids are other reactions that are discussed.

In this chapter, first, we hope to bring to readers some basic electrochemical fundamentals of solution-phase voltammetry and then show how these simple equations may be applied to characterize specifically porphyrinoids electrochemically. Each example was chosen to demonstrate a different electrochemical aspect. The focus then develops into changes brought about when porphyrinoids are adsorbed on the surface of an electrode first by nonspecific absorption and then by specific adsorption. When in solution, the analyte can diffuse to and from the electrode, and electron transfer between porphyrinoid and electrode is a relatively slow outer-sphere process. Provided the rate of electron transfer is still much faster than the rate of diffusion, the porphyrinoid continues to exhibit electrochemically reversible or quasi-reversible behavior. However, when specifically adsorbed, either by multiple secondary forces such as π–π stacking or by formal bond formation, diffusion is not anymore possible. Electron transfer takes place by an inner-sphere mechanism that is much faster than outer-sphere electron transfer. Measured electrochemical responses also differ. For example, for specifically adsorbed porphyrinoids, peak currents are linearly related to scan rates, while for dissolved porphyrinoids, peak currents are linearly related to the square root of scan rate.

Specific adsorption is a key aspect in electrocatalysis, and we then spontaneously digress into some simplified fundamental aspects of heterogeneous electrocatalysis. Electrocatalytic materials are researched in a variety of fields including fuel cells,1  chemical synthesis,2  energy storage,3  and sensors.4,5  The chapter concludes with examples in which porphyrinoids have been used for electrocatalysis.

The examples discussed here have never been intended as an exhaustive review of electrocatalysis by porphyrinoids. Rather, they were chosen to demonstrate how electrocatalytic activity changes due to porphyrinoid structural changes. That said, examples other than those discussed, have also been cited for the benefit of readers. We hope this work will assist many beginning and experienced researchers in the field of solution-phase electrochemistry and electrocatalysis.

Electrochemical properties of a compound are conveniently probed in solution by techniques such as cyclic voltammetry (CV), square wave voltammetry (SWV), and linear sweep voltammetry (LSV).6  When two closely overlapping redox processes of a compound follow each other, SWV can make it a little easier to resolve the potentials at which they occur. Following the procedure developed by Taube,7  for electrochemical reversible redox processes, the limit of potential resolution of two separate redox processes (differences in formal reduction potential, ) by both CV and SWV is in the vicinity of 50 mV. However, to accurately obtain them by CV using the Taube method is much more cumbersome when Ep values cannot be measured directly. The key factor is identifying peak potentials, Ep, of both peaks accurately. Experimental measurements and theoretical simulations show Ep measurements can still be done easily with SWV if ΔE°′ = 80 mV, but in CV, Ep identification struggles even at ΔE°′ = 100 mV. Only when both peaks can be measured directly with CV (at ΔE°′ = ca. 250 mV) will direct CV-measured E°′ values be accurate (within 1 mV of the correct potential).7  An LSV is a CV performed at a very slow scan rate, 1 or 2 mV s−1. It is the electrochemical equivalent of NMR integration giving the relative number of electrons that flow in each redox process in a multi-redox system.

Electrochemical reversible redox processes are characterized in CV by three equations.8,9 

ΔEp = EpaEpc = 59/n
Equation 1.1
ipa/ipc = 1
Equation 1.2
E°′ = (Epa + Epc)/2
Equation 1.3

In eqn (1.1)(1.3), n is the number of electrons transferred in the balanced equation of the redox process, ipa is the peak anodic current at the peak anodic potential (Epa), ipc is the peak cathodic current at the peak cathodic potential (Epc), and E°′ is the formal reduction potential of the reversible redox process. For eqn (1.2), ipa/ipc is meant to be the “reverse scan current” divided by the “forward scan current.” If this is not so, the ratio should be inverted to ipc/ipa = 1. Care must be taken when choosing a solvent and supporting electrolyte for a particular system because popular solvents like acetonitrile may complex, or at the very least associate, with a metal, and often-used supporting electrolytes like [nBu4N][PF6] may form ion pairs of the type (cation)+(PF6). For this reason, many CV/SWV/LSV measurements are now conducted in CH2Cl2 utilizing a super non-ion pair forming supporting electrolyte like [nBu4N][B(C6F5)4]. Working electrodes are chosen in such a way as to minimize any layer forming (coating) interactions with analytes.

Figure 1.1 shows the general structure of porphyrin (por), phthalocyanines (Pcs), and subphthalocyanines (subPcs). Electrochemically, most porphyrinoids exhibit two ring-based oxidations and four ring-based reductions, but all six redox processes are seldom observed in a solvent potential window.

Figure 1.1

Structure of some porphyrinoids: Porphyrin (top left), sapphyrin (an expanded corrole, top right), substituted phthalocyanines (Pcs, bottom left), and contracted phthalocyanines = subphthalocyanines (subPcs, bottom right). R =substituents to enhance desired properties (e.g., better solubility in organic solvents). M can be any of more than 70 elements for most porphyrinoids. SubPcs can only have boron coordinated in the cavity and are constructed from just three iminoisoindoline units.

Figure 1.1

Structure of some porphyrinoids: Porphyrin (top left), sapphyrin (an expanded corrole, top right), substituted phthalocyanines (Pcs, bottom left), and contracted phthalocyanines = subphthalocyanines (subPcs, bottom right). R =substituents to enhance desired properties (e.g., better solubility in organic solvents). M can be any of more than 70 elements for most porphyrinoids. SubPcs can only have boron coordinated in the cavity and are constructed from just three iminoisoindoline units.

Close modal

Scheme 1.1 shows the synthetic route toward metallated tetrabenzoporphyrins, TBPs, tetrabenzomonoazaporphyrins, TBMAPs, tetrabenzodiazaporphyrins, TBDAPs, and tetrabenzotriazaporphyrins, TBTAPs, 14.10  These porphyrinoid types have not been extensively studied.11,12  The metal-free derivatives showed in CH2Cl2/0.1 mol dm−3 [nBu4N][B(C6H5)4] two ring-based oxidations and two ring-based reductions. Some copper complexes also exhibit a CuII reduction wave in this potential region (see Figure 1.2). The relationship between E°′ and ΣχR where ΣχR is the sum of the Pauling scale group electronegativities of the apex atoms N and CH in 14 may also be found in Figure 1.2.10  ΣχR is calculated, for example for 2, as follows: ΣχR(2) = χN + 3 · {χC + χH} = 3.04 + 3(2.55 + 2.20) = 17.29. The larger this number, the more electron-withdrawing the macrocycle becomes, meaning that 4 with three aza groups and ΣχR = 13.87 is more electron-rich than 2 having only one aza group and ΣχR = 17.29.

Scheme 1.1

Synthesis of metal-free and copper-coordinated hybrid benzoporphyrins bearing monoaza, diaza, and triaza apexes (nitrogen meso atoms); R = hexyl. Reproduced from ref. 10 with permission from American Chemical Society, Copyright 2015.

Scheme 1.1

Synthesis of metal-free and copper-coordinated hybrid benzoporphyrins bearing monoaza, diaza, and triaza apexes (nitrogen meso atoms); R = hexyl. Reproduced from ref. 10 with permission from American Chemical Society, Copyright 2015.

Close modal
Figure 1.2

Left: CVs of ca. 0.5 mmol dm−3 solutions of Cu-TBP (1b, green), Cu-TBMAP (2b, blue), Cu-TBDAP (3b, red), and Cu-TBTAP (4b, black) in CH2Cl2/0.1 mol dm−3 of [nBu4N][B(C6H5)4] at 100 mV s−1. Fc* = Decamethylferrocene, the internal standard. The notation 0N to 3N denotes the number of meso nitrogen atoms in each compound. Right: Relationship between E° of redox waves 1–4 and ΣχR for metal-free (green lines) and copper complexes (blue lines) 14. Reproduced from ref. 10 with permission from American Chemical Society, Copyright 2015.

Figure 1.2

Left: CVs of ca. 0.5 mmol dm−3 solutions of Cu-TBP (1b, green), Cu-TBMAP (2b, blue), Cu-TBDAP (3b, red), and Cu-TBTAP (4b, black) in CH2Cl2/0.1 mol dm−3 of [nBu4N][B(C6H5)4] at 100 mV s−1. Fc* = Decamethylferrocene, the internal standard. The notation 0N to 3N denotes the number of meso nitrogen atoms in each compound. Right: Relationship between E° of redox waves 1–4 and ΣχR for metal-free (green lines) and copper complexes (blue lines) 14. Reproduced from ref. 10 with permission from American Chemical Society, Copyright 2015.

Close modal

For wave 3, linearity in the relationship between copper complex E°′ and ΣχR value breaks down. This is consistent with the source of this reduction in 1b (N = 0) and 2b (N = 1) not being the same as for 3b (N = 2) and 4b (N = 3). DFT (density functional theory) calculations showed that for 1b and 2b, the first reduction wave labeled 3 in Figure 1.2 is not a ring-based reduction as it is for 3b and 4b, but a CuII reduction.10 

The linear relationship between E°′ and ΣχR (see Figure 1.2) is indicative of good communication within the porphyrinoid macrocycle. Good electronic communication also extends to substituents on the macrocycle and axial ligands of coordinated metals. Figure 1.3 shows the structures of 5,15-diphenylporphyrin, 5; 5-ferrocenyl-10,20-diphenylporphyrin, 6; 5-ferrocenyl-15-phenylporphyrin, 7; and 5,15-diferrocenyl-15-phenylporphyrin, 8 obtained after scrambling of intermediate products of the MacDonald-type 2 + 2 condensation of 5-ferrocenyldipyrromethane, dipyrromethane and benzaldehyde in CH2Cl2 at room temperature in the presence of trifluoroacetic acid.13  Oxidation was with 2,3-dichloro-5,6-dicyanobenzoquinone, and the trifluoroacetic acid was quenched at the end of the reaction with triethylamine.

Figure 1.3

Left: Structures of porphyrins 5–8. Top right: CVs at scan rates of 100, 200, 300, 400, and 500 mV s−1, LSVs (2 mV s−1), and SWVs at 10 Hz of 8 in CH2Cl2 containing 0.1 mol dm−3 of [nBu4N][B(C6F5)4] at 25 °C. Fc* = Decamethylferrocene, the internal standard. Bottom right: Relationship between E°′ and Σχmeso for waves 1, 2, and 3 (the ferrocenyl waves) and wave 4 of complexes 58 as well as 5,15-diferrocenyl-10-20-dipentafluorophenylporphyrin, 9 (data from literature14 ). aWave 3b for porphyrin 8. bData point14  corresponds to wave 3b for porphyrin 9, tpp = 5,10,15,20-tetraphenylporphyrin. Reproduced from ref. 13 with permission from the Royal Society of Chemistry.

Figure 1.3

Left: Structures of porphyrins 5–8. Top right: CVs at scan rates of 100, 200, 300, 400, and 500 mV s−1, LSVs (2 mV s−1), and SWVs at 10 Hz of 8 in CH2Cl2 containing 0.1 mol dm−3 of [nBu4N][B(C6F5)4] at 25 °C. Fc* = Decamethylferrocene, the internal standard. Bottom right: Relationship between E°′ and Σχmeso for waves 1, 2, and 3 (the ferrocenyl waves) and wave 4 of complexes 58 as well as 5,15-diferrocenyl-10-20-dipentafluorophenylporphyrin, 9 (data from literature14 ). aWave 3b for porphyrin 8. bData point14  corresponds to wave 3b for porphyrin 9, tpp = 5,10,15,20-tetraphenylporphyrin. Reproduced from ref. 13 with permission from the Royal Society of Chemistry.

Close modal

The CV, LSVs, and SWVs of 8 as representative examples are also shown in Figure 1.3. Key observations relate first to the LSVs that demonstrate each redox process involving the same number of electrons; here it is representative of a one-electron transfer process. Next, it is clearly easier to identify accurately peak potentials of wave 3 in the SWV than in the CV. Finally, the two ferrocenyl (= Fc) groups of 8 do not exhibit equivalent electrochemical behavior. When the first is oxidized to ferrocenium, the Fc group's Gordy scale group electronegativity14  is changed from χFc = 1.87 to χFc + = 2.82 (i.e., the strongly electron-donating Fc group became strongly electron-withdrawing). This caused the second Fc group to be oxidized at a potential 132 mV larger than the first Fc group. Figure 1.3, bottom right, shows that communication between different meso substituents in 58 is good enough that a linear dependency exists for all four redox waves between E°′ of waves 1, 2, 3 (the ferrocenyl redox processes) and 4 and ΣχR = χR1 + χR2 + χR3 + χR4. χR1 through χR4 can be deduced from the structures of 58, and χC6H5 = 2.21, χC6F5 = 3.01 (9 possess the C6F5 group), and χH = 2.20.14  Electronic communication between Fc substituents was also confirmed for 3,3′-diferrocenyl azaBODIPYs (difluoroboryl complex of azodipyrromethene),15  tetraferrocenylporphyrin (H2TFcP),16  corroles,17  and many others.18  The H2TFcP study also highlighted the advantage of using CH2Cl2 over acetonitrile, DMF, and THF as solvent (but see advantages under certain conditions of THF as solvent in Section 1.3.4).16  Nemykin wrote a good review on the magic of compounds obtained as a result of a marriage between metallocenes and porphyrinoids.19 

Communication to axial ligands is just as good as to ring substituents. Figure 1.4, top left, shows the structures and voltammograms of subphthalocyanines (subPcs) 1013.20  SubPcs 10 and 11 have hydrogen atoms on the annulated benzene rings, while 12 and 13 have fluorine atoms. Only one ring-based oxidation, wave A, was observed in the CH2Cl2 potential window, but E°′ for this wave for the fluoro complexes 12 and 13 was 395 mV more positive than E°′ of 10 and 11. E°′ of the first reduction wave, wave I, of 13 was offset with 544 mV from that of 11, and the second ring-based reduction of 12 occurred at a potential almost 600 mV more positive than that of 10.

Figure 1.4

Structures of subphthalocyanines 1013 and CVs associated with them obtained from 0.5 mM solutions of analytes in CH2Cl2 containing 0.1 mol dm−3 of [nBu4N][B(C6F5)4] at 25 °C at a scan rate of 100 mV s−1. Reproduced from ref. 20 with permission from American Chemical Society, Copyright 2020.

Figure 1.4

Structures of subphthalocyanines 1013 and CVs associated with them obtained from 0.5 mM solutions of analytes in CH2Cl2 containing 0.1 mol dm−3 of [nBu4N][B(C6F5)4] at 25 °C at a scan rate of 100 mV s−1. Reproduced from ref. 20 with permission from American Chemical Society, Copyright 2020.

Close modal

Most striking, the fluorinated complexes 12 and 13 exhibited a third ring base reduction, while compounds 10 and 11 did not show this redox process within the potential window of the solvent.

The axial ligand in 1013 is either OOC–CHCH–Fc (in 10 and 12) or OOC–CH2–CH2–Fc (in 11 or 13). The CHCH moiety allows communication between the subPc core of 10 and 11 and the Fc group, while the CH2CH2 atom chain in the axial ligand of 12 and 13 isolates the ferrocenyl group from the subPc core. This isolating effect of the CH2CH2 group manifested in the Fc group of 11 and 13 being oxidized at almost the same potential, −0.058 and 0.050 mV versus FcH/FcH+, respectively, despite 11 having 12 hydrogen atoms and 13 having 12 fluorine atoms bound to the subPc macrocycle.20 

The situation is rather different for 10 and 12 where the CHCH group separates the Fc group from the subPc cores. The formal reduction potentials of the Fc group of 10 and 12 (the compound with 12 F atoms) were 0.119 and 0.243 V versus FcH/FcH+, respectively. These four Fc E°′ values indicate that axial ligands and substituents on the macrocycle core, when conjugation is possible, influence the redox properties of each other. Communication between the Fc group of ferrocene-containing β-diketonato ligands and the phthalocyaninato macrocycle when coordinated with Zr and Hf Pcs was also observed and discussed.21 

It follows that a careful choice of ring substituents and axial ligands provides a route to fine adjustments of formal reduction potentials, E°′, of redox-active components in porphyrinoid complexes.

Sandwiched (double-decker) and triple-decker porphyrinoids are especially known for lanthanide and actinide complexes.22–25  Utilizing the example of a non-rare earth cadmium triple-decker phthalocyanine,26  it is shown how the careful choice of phthalocyanine (Pc) ring substituents can limit Pc aggregation (i.e., enhance solubility) and how solvent choice with a large working potential window may allow observation of more than just three or four ring-based porphyrinoid redox processes. The CVs of subPcs 10 and 11 exhibited three ring-based redox processes, and 12 and 13 exhibited four of the expected six redox processes in the solvent (CH2Cl2) potential window (see Figure 1.4). To observe all six expected porphyrinoid ring-based redox processes in a particular solvent is unusual. One such example in which it was observed is the cadmium phthalocyaninato triple-decker, 14, for which 17 of the possible 18 redox processes could be observed in THF as solvent (see Figure 1.5).26  The relative number of electrons flowing for each redox process is shown by the LSVs. Especially the SWV in CH2Cl2 of waves IV and C and the SWV of waves V1 and V2 in THF demonstrate better potential resolution than the CVs. Notable in Figure 1.5 are the differences in measured anodic potentials in THF and CH2Cl2. The free electron pairs on THF make it nucleophilic and capable of interacting or associating with oxidized species having a positive charge. This caused the six anodic waves of 14 to coalesce into two major waves, waves A and B, while in CH2Cl2 (which do not interact with positively charged species), the same redox processes were observable in three redox waves, waves A, B, and C. However, the interaction of THF with negatively charged reduced species is very weak or non-existent; hence, the cathodic CV fingerprint of 14 in THF and CH2Cl2 is similar. Because THF has a wider cathodic potential window than CH2Cl2, the redox process labeled VI of 14 could be detected in THF but not in CH2Cl2.26 

Figure 1.5

The structure of the triple-decker phthalocyanine 14 is shown top left; voltammograms of 14 are shown on the right. Blue: CVs of a 0.5 mM solution of 14 in CH2Cl2/0.1 mol dm−3 of [nBu4N][B(C6F5)4] at −40 °C scan rates are 100 (smallest currents), 200, 300, 400, and 500 mV s−1. Black: CV recorded in THF at 100 mV s−1. Purple: LSVs at 2 mV s−1; currents were multiplied by 2 for better display purposes. Red: SWVs at 10 Hz. Bottom left: Changes in the UV–vis spectrum of 14 as a solution in CH2Cl2 upon addition of aliquots of iodine. Reproduced from ref. 26 with permission from the Royal Society of Chemistry.

Figure 1.5

The structure of the triple-decker phthalocyanine 14 is shown top left; voltammograms of 14 are shown on the right. Blue: CVs of a 0.5 mM solution of 14 in CH2Cl2/0.1 mol dm−3 of [nBu4N][B(C6F5)4] at −40 °C scan rates are 100 (smallest currents), 200, 300, 400, and 500 mV s−1. Black: CV recorded in THF at 100 mV s−1. Purple: LSVs at 2 mV s−1; currents were multiplied by 2 for better display purposes. Red: SWVs at 10 Hz. Bottom left: Changes in the UV–vis spectrum of 14 as a solution in CH2Cl2 upon addition of aliquots of iodine. Reproduced from ref. 26 with permission from the Royal Society of Chemistry.

Close modal

One of the reasons 14 shows such well-defined CV shapes is the good solubility the eight non-peripheral C6H13 substituents impose onto the porphyrinoid. It was shown that solubility increases substantially (i.e., the tendency to aggregate decreases) in organic solvents as the alkyl substituent chain length (n in CnH2n + 1) increases. Listed below is how the onset aggregation concentration increases with increasing n where eight CnH2n + 1 groups are bound to the non-peripheral positions of ZnPc.27  {(n; [octa-subst. ZnPc]/μM) = (5; 9), (6; 11), (7; 32), (8; 27), (9; 42), (10; 152), (12; 400)}.

Lowering the tendency to aggregate is much enhanced from n = 10. This effect is less pronounced if the substituents are on the non-peripheral positions.

The above-described CVs have near-ideal shapes and obeyed eqn (1.1)(1.3) because they were involved in near-ideal, outer-sphere, heterogeneous electron transfer between electrode and porphyrinoid substrate.28  This means that porphyrinoid reactants, intermediates, and products do not adsorb (e.g., via aggregation) on the electrode surface. Heterogeneous electron transfer between electrode and porphyrinoid must occur by electron tunneling across at least a mono layer of solvent separating electrode from porphyrinoid.28,29  This process is demonstrated in Figure 1.6.

Figure 1.6

Left: Model of heterogeneous outer-sphere electron-transfer between substrate (here porphyrinoids) and electrode. Redox processes are relatively slow as the transferred electron must tunnel randomly through at least a solvent layer separating porphyrinoid from electrode (the red hosepipes demonstrate a random electron tunnel path to and from the electrode).28  The rate of electron-transfer slows down as the distance between electrode and substrate increase, but if diffusion rate ≪ electron transfer rate, then the redox process is electrochemically reversible and ΔEp = 59/n mV. Right: The redox potential associated with decamethylferrocene, Fc* versus FcH/FcH+ depends on the solvent and electrolyte used; see also Table 1.1. In the presence of ca. 0.7 mM FcH, the redox potential of ca. 0.5 mM solutions of Fc* in CH2Cl2/0.2 M [nBu4N][B(C6F5)4] is −610 mV (blue CV), in CH2Cl2/0.2 M [nBu4N][PF6], it is −550 mV (red CV), and in CH3CN/0.2 M [nBu4N][PF6] (black CV), it is 508 mV versus FcH/FcH+.

Figure 1.6

Left: Model of heterogeneous outer-sphere electron-transfer between substrate (here porphyrinoids) and electrode. Redox processes are relatively slow as the transferred electron must tunnel randomly through at least a solvent layer separating porphyrinoid from electrode (the red hosepipes demonstrate a random electron tunnel path to and from the electrode).28  The rate of electron-transfer slows down as the distance between electrode and substrate increase, but if diffusion rate ≪ electron transfer rate, then the redox process is electrochemically reversible and ΔEp = 59/n mV. Right: The redox potential associated with decamethylferrocene, Fc* versus FcH/FcH+ depends on the solvent and electrolyte used; see also Table 1.1. In the presence of ca. 0.7 mM FcH, the redox potential of ca. 0.5 mM solutions of Fc* in CH2Cl2/0.2 M [nBu4N][B(C6F5)4] is −610 mV (blue CV), in CH2Cl2/0.2 M [nBu4N][PF6], it is −550 mV (red CV), and in CH3CN/0.2 M [nBu4N][PF6] (black CV), it is 508 mV versus FcH/FcH+.

Close modal

Redox potentials in organic media according to IUPAC should be referenced versus the FcH/FcH+ redox couple at 0.00 V.30  However, the potential of the FcH/FcH+ couple is dependent on the solvent used. If this redox process overlaps with redox processes of the analyte, decamethylferrocene, Fc*, is often used in its place. The Fc* redox potential relative to ferrocene in different solvents is well described elsewhere,31–34  but for convenience, some conversion factors are given in Table 1.1.

Solvent0.1 M electrolyteE°′(ref. electr.)/mVE°′(other ref. electr.)/mVE°′(other ref. electr.)/mVReference
CH2Cl2 [nBu4N][B(C6F5)4FcH0/ +; 0 Fc* 0/ +; −610 Ag/Ag+; −220tab1fna 26  
CH2Cl2 [nBu4N][PF6]tab1fnb FcH0/ +; 0 Fc* 0/ +; −550 SCE; −460/ − 475 26, 32, 34  
CH2Cl2 [nBu4N][ClO4FcH0/ +; 480 SCE; 0  34  
CH3CN [nBu4N][B(C6F5)4FcH0/ +; 0 Fc* 0/ +; −516  26  
CH3CN [nBu4N][PF6]tab1fnc FcH0/ +; 0 Fc* 0/ +; −508 SCE; −400/ − 382 26  
CH3CN [nBu4N][ClO4FcH0/ +; 380 SCE; 0  34  
CH3CN [nEt4N][PF6FcH0/ +; 380 SCE; 0  34  
THF [nBu4N][B(C6F5)4FcH0/ +; 0 Fc* 0/ +; −515 Ag/Ag+; −176tab1fna 26  
THF [nBu4N][PF6FcH0/ +; 547/560 SCE; 0 CoCp*20/ +; −1295 32, 34  
THF [nBu4N][ClO4FcH0/ +; 530 SCE; 0  34  
DMF [nBu4N][PF6FcH0/ +; 450/470 SCE; 0 CoCp*20/ +; −1402 32, 34  
DMF [nBu4N][ClO4FcH0/ +; 470 SCE; 0  34  
DMSO [nBu4N][PF6FcH0/ +; 435 SCE; 0 CoCp*20/ +; −1425 32  
DME [nBu4N][PF6FcH0/ +; 510 SCE; 0 CoCp*20/ +; −1260 32  
CH3NO2 [nBu4N][PF6FcH0/ +; 350 SCE; 0  34  
Solvent0.1 M electrolyteE°′(ref. electr.)/mVE°′(other ref. electr.)/mVE°′(other ref. electr.)/mVReference
CH2Cl2 [nBu4N][B(C6F5)4FcH0/ +; 0 Fc* 0/ +; −610 Ag/Ag+; −220tab1fna 26  
CH2Cl2 [nBu4N][PF6]tab1fnb FcH0/ +; 0 Fc* 0/ +; −550 SCE; −460/ − 475 26, 32, 34  
CH2Cl2 [nBu4N][ClO4FcH0/ +; 480 SCE; 0  34  
CH3CN [nBu4N][B(C6F5)4FcH0/ +; 0 Fc* 0/ +; −516  26  
CH3CN [nBu4N][PF6]tab1fnc FcH0/ +; 0 Fc* 0/ +; −508 SCE; −400/ − 382 26  
CH3CN [nBu4N][ClO4FcH0/ +; 380 SCE; 0  34  
CH3CN [nEt4N][PF6FcH0/ +; 380 SCE; 0  34  
THF [nBu4N][B(C6F5)4FcH0/ +; 0 Fc* 0/ +; −515 Ag/Ag+; −176tab1fna 26  
THF [nBu4N][PF6FcH0/ +; 547/560 SCE; 0 CoCp*20/ +; −1295 32, 34  
THF [nBu4N][ClO4FcH0/ +; 530 SCE; 0  34  
DMF [nBu4N][PF6FcH0/ +; 450/470 SCE; 0 CoCp*20/ +; −1402 32, 34  
DMF [nBu4N][ClO4FcH0/ +; 470 SCE; 0  34  
DMSO [nBu4N][PF6FcH0/ +; 435 SCE; 0 CoCp*20/ +; −1425 32  
DME [nBu4N][PF6FcH0/ +; 510 SCE; 0 CoCp*20/ +; −1260 32  
CH3NO2 [nBu4N][PF6FcH0/ +; 350 SCE; 0  34  
Aqueous reference electrodes
SolventElectrode; E°′/VConditionsE°′(other ref. electr.)/VE°′(other ref. electr.)/V
H2SHE; 0.000 V a(H + ) = 1 Mtab1fnd Ag/AgCl; 0.20tab1fne  
H2NHE; 0.000 V [H+] = 1 M Ag/AgCl; 0.23tab1fne SCE; 0.241 
H2RHE; 0.000 V E = 0 − 0.0591 pH   
Aqueous reference electrodes
SolventElectrode; E°′/VConditionsE°′(other ref. electr.)/VE°′(other ref. electr.)/V
H2SHE; 0.000 V a(H + ) = 1 Mtab1fnd Ag/AgCl; 0.20tab1fne  
H2NHE; 0.000 V [H+] = 1 M Ag/AgCl; 0.23tab1fne SCE; 0.241 
H2RHE; 0.000 V E = 0 − 0.0591 pH   
tab1fna

[Ag+] = [AgNO3] = 0.05 M in CH3CN.

tab1fnb

CoCp*20/ + in CH2Cl2/[nBu4N][PF6] has E°′ = −1497 mV versus SCE.32 

tab1fnc

CoCp*20/ + in CH3CN/[nBu4N][PF6] has E°′ = −1525 mV versus SCE.32 

tab1fnd

a = activity, a = 1 M at [H+] ca. 1.2 M HCl.

tab1fne

In saturated KCl(aq).

If, however, the porphyrinoid interacts weakly with the electrode, it may nonspecifically adsorb on the surface of the electrode.35,36  When a porphyrinoid in the oxidized (or reduced) form is nonspecifically (= weakly) absorbed on an electrode surface, the height of the cathodic peak, ipc, (or the anodic peak current, ipc) will increase faster than the height of peaks when the porphyrinoid is involved in normal ideal outer-sphere electron transfer processes because both absorbed and diffusing oxidized porphyrinoid contribute to the current.35  The effects of nonspecific adsorption of a compound on an electrode are demonstrated in Figure 1.7, left, for phthalocyanine 15,37  where the cathodic peak height of reduction half wave IIIads is larger than the anodic peak height (ipa of the anodic CV half wave III), as well as the cathodic peak heights of waves I and II. This implies that ipa/ipc is substantially less than 1. Since ipc would become larger more quickly with faster scan rates than ipa35  would, the current ratio ipa/ipc would deviate further from unity at faster scan rates. Aggregation probably plays a large role in observation of this effect and could be minimized or completely avoided if the substituents on the peripheral positions had C10 or longer chain lengths than the 2-ethylhexyl chains that were used here (see also Section 1.3.4). Other examples of this effect related to ferrocene-containing porphyrinoids may be found in the literature.38,39 

Figure 1.7

Left: The cathodic (reducing) portion of CV of 15 (with = 2-ethylhexyl) in THF/0.1 [nBu4N][B(C6F5)4] at 25 °C and scan rate 100 mV s−1. The triple-reduced form of 15 adsorbs nonspecifically (weakly) on the electrode; hence, ipc of wave IIIads are much larger than the other peak currents. Ads = Adsorbtive, pre = prewave (CV). Reproduced from ref. 37 with permission from the Royal Society of Chemistry. Right: A theoretical simulation of a species that adsorbs strongly on an electrode, generates wave 1 (ΔEp = 0 mV). If some substrate is still left in solution after adsorption, it will generate wave 2 (ΔEp = 59/n mV) during recording of a CV. Note for this figure, the X-axis does not run in the normal Cartesian direction to be consistent with the original publication. CVs plotted in this direction follow the American convention.6  Reproduced from ref. 36 with permission from American Chemical Society, Copyright 1967.

Figure 1.7

Left: The cathodic (reducing) portion of CV of 15 (with = 2-ethylhexyl) in THF/0.1 [nBu4N][B(C6F5)4] at 25 °C and scan rate 100 mV s−1. The triple-reduced form of 15 adsorbs nonspecifically (weakly) on the electrode; hence, ipc of wave IIIads are much larger than the other peak currents. Ads = Adsorbtive, pre = prewave (CV). Reproduced from ref. 37 with permission from the Royal Society of Chemistry. Right: A theoretical simulation of a species that adsorbs strongly on an electrode, generates wave 1 (ΔEp = 0 mV). If some substrate is still left in solution after adsorption, it will generate wave 2 (ΔEp = 59/n mV) during recording of a CV. Note for this figure, the X-axis does not run in the normal Cartesian direction to be consistent with the original publication. CVs plotted in this direction follow the American convention.6  Reproduced from ref. 36 with permission from American Chemical Society, Copyright 1967.

Close modal

However, if the porphyrinoid is strongly absorbed onto the electrode surface, it cannot continue to diffuse away from the electrode, eqn (1.1)(1.3) are not valid anymore, and ΔEp becomes zero (see Figure 1.7 right, wave 1). Strong adsorption is normally associated with specific adsorption (or in some cases, even bond formation) of a substrate onto an electrode forming a layer of adsorbate on it. Specific adsorption may be described by the Langmuir, Temkin, or other isotherms.40  Electron transfer in such a system is very fast as it proceeds via an inner-sphere mechanism.28  Different potentials are observed for specific adsorbed substrates and dissolved substrates. If some dissolved substrate is still left in the solvent after specific substrate adsorption on the electrode, this will be electrochemically observable as a near-ideal, outer-sphere, heterogeneous electron transfer process between electrode and porphyrinoid substrate [i.e., almost a “normal” CV portion as shown in wave 2 (see Figure 1.7)].35,36  This figure describes the following scenario.35  First, the porphyrinoid substrate is dissolved in the oxidized form, Odissolved. A portion of Odissolved is then reduced (wave 1), and the reduced species R simultaneously adsorbs strongly on the electrode forming a layer of this reduced product (Radsorbed). This wave 1 reduction occurs at potentials more positive than the diffusion-controlled reduction process at wave 2 because the free energy of adsorption of the reduced form, Radsorbed, makes simultaneous reduction and adsorption of the dissolved oxidized form, Odissolved, easier than to generate the reduced form in solution. Wave 1 associated with the half reaction Odissolved + e → Radsorbed is then followed by wave 2, which is associated with the half reaction Odissolved + e → Rdissolved. The peak current of wave 1 (i.e., of the specific adsorbed species) increases linearly with scan rate υ, while the peak current of wave 2 increases linearly with υ1/2. Finally, if the electrode covered with Radsorbed is then removed from the electrochemical cell, rinsed to remove all solvent containing substrate and electrolyte, and then immersed into a fresh solution of electrolyte only, and if the adsorbed layer is still intact (i.e., it was adsorbed strong enough not to be washed away), a new CV experiment would give only wave 1 in the voltammogram.

A key factor to observe is that electron transfer between electrode and specific adsorbed species follows an inner-sphere pathway that allows for much faster rates of electron transfer than outer-sphere processes. In homogeneous terms, this is called the kinetic advantage, and the ratio of inner-sphere to outer-sphere rate constants, kinner/kouter, has recently been shown41  to be as high as 104.81 for electron transfer reactions between a nickel(iii) peptide and IrIIICl63 − or 104.12 between a nickel(iii) peptide and Mo(CN)84 −. As shown in Figure 1.8, inner-sphere heterogeneous electron transfer between electrode and substrate occurs while the substrate is in firm, stable contact with the electrode (i.e., specifically adsorbed) either through a formal bond or via very strong secondary forces. This adsorption mode provides the shortest possible distance between electrode and substrate. The overlapping orbitals of substrate and electrode provide a “highway” for electron flow between electrode and substrate. In this scenario, no diffusion of the substrate is possible and ΔEp = 0.

Figure 1.8

Model of heterogeneous inner-sphere electron transfer between substrate and electrode. The substrate is either bound to the electrode or kept firmly in place on it by many secondary forces; for porphyrinoids, this is often π–π stacking. Redox processes under these conditions are orders of magnitude faster than in outer-sphere redox processes because the transferred electrons have a “highway” of primary and/or secondary bond orbitals at its disposal to travel between electrode and substrate. This is shown by the dotted red line encapsulated in the “orbital barrel” between the electrode and substrate.28  CVs of such systems, because diffusion is not possible, show ΔEp = 0, while ΔEp = 59/n mV for outer-sphere redox processes. Peak currents obtained for outer-sphere processes increase linearly with the square root of scan rate υ1/2, while for inner-sphere redox processes, it increases linearly and much faster with υ. The CV shown is of decamethylferrocene Fc* exhibiting outer-sphere electron transfer in solution, while the wave to the right is of −S(CH2)4Fc (Fc = ferrocenyl) covalently bound to a gold electrode. Reproduced from ref. 42 with permission from American Chemical Society, Copyright 2016.

Figure 1.8

Model of heterogeneous inner-sphere electron transfer between substrate and electrode. The substrate is either bound to the electrode or kept firmly in place on it by many secondary forces; for porphyrinoids, this is often π–π stacking. Redox processes under these conditions are orders of magnitude faster than in outer-sphere redox processes because the transferred electrons have a “highway” of primary and/or secondary bond orbitals at its disposal to travel between electrode and substrate. This is shown by the dotted red line encapsulated in the “orbital barrel” between the electrode and substrate.28  CVs of such systems, because diffusion is not possible, show ΔEp = 0, while ΔEp = 59/n mV for outer-sphere redox processes. Peak currents obtained for outer-sphere processes increase linearly with the square root of scan rate υ1/2, while for inner-sphere redox processes, it increases linearly and much faster with υ. The CV shown is of decamethylferrocene Fc* exhibiting outer-sphere electron transfer in solution, while the wave to the right is of −S(CH2)4Fc (Fc = ferrocenyl) covalently bound to a gold electrode. Reproduced from ref. 42 with permission from American Chemical Society, Copyright 2016.

Close modal

Closer proximity of the redox center to the electrode surface has a profound influence on the rate of electron transfer. It was shown that the rate of electron transfer between ferrocenyl derivatives, Fc(CH2)nS–, bound to the surface of a gold electrode differs with n as follows:42,43 

(n; kET/s−1) = (1; 2 × 109), (2; 5 × 108), (3; 2 × 108), (4; 5 × 107), (5; 1.6 × 107), (6; 2.4 × 106), (8; 4.4 × 105), (10; 4 × 104), (11; 1.2 × 104), (12; 1.700 × 103), (16; 28).

The dissolved decamethylferrocene in the CV shown in Figure 1.8 highlights the linear dependence of peak currents against the square root of scan rate for solution-based CVs according to the Randles–Sevcik equation,8,9 eqn (1.4) and (1.5), where A is the electrode surface area, C is the bulk analyte concentration, F is the Faraday constant, D is the diffusion constant, υ is the scan rate, and n is the number of electrons transferred in the balanced redox reaction.

ip = 0.4463 nFAC(nFυD/RT)0.5
Equation 1.4
= (2.69 × 105)n3/2AD1/2υ1/2C at 25 °C
Equation 1.5

In contrast, for electrochemical reversible systems, anodic and cathodic waves of surface-bound analytes, in Figure 1.8, the self-assembled monolayer of Fc(CH2)4S– on gold, should be symmetrical, and ΔE = EpaEpc should be 0 mV.42  The relationship between peak currents and scan rate is given by eqn (1.6) where Asur is the electrode surface area, and Γ is the surface coverage.42,43 

ip = (n2F2υAsurΓ)/(4RT)
Equation 1.6

For dissolved species, the slope of an ipversus υ1/2 plot is used to obtain the diffusion constant of the species, while the slope of ipversus υ plots leads to the surface coverage Γ.

A final thought is that some species dissolved in a solvent may adsorb and pollute the electrode to the point where it becomes pacified or useless for electrochemical measurements. But this is not always the case. The above-described specific adsorption effects that lead to inner-sphere electron transfer are, in fact, the bases of electrocatalysis. The remainder of this chapter will focus on simplified fundamental thoughts required for treatment of data of electrocatalytic systems followed by some selected examples of this important heterogeneous process.

Electrocatalysis is just another form of heterogeneous catalysis. It catalyzes a reaction involving electron flow to and from an electrode. Electron flow takes place at the surface of the electrode, but as shown below, the catalyst does not need to be adsorbed on the electrode. It can also be in solution as long as it has access to the electrode via, for example, diffusion. Being involved in “catalytic” processes, the catalyst must also fit in with general principles of heterogeneous/homogeneous catalysis. The most important catalyst characteristics are as follows:

  • a. It must be stable.

  • b. It must be active.

  • c. It must be selective for the process for which it was chosen.

One would like to have the catalytic process operating as fast as possible at potentials as close as possible to the thermodynamic reversible electrode potential. A key equation to evaluate this, is the Butler–Volmer equation, eqn (1.7).44 

formula
Equation 1.7

In eqn (1.7), i is the total electrode current density of the electrode reaction in units of A m−2, the first term in brackets is the cathodic component, the second term is the anodic component, i0 is the exchange current density (A m−2), n = number of electrons transferred in the rate-determining step, αa is the dimensionless anodic transfer coefficient, η is the overpotential of the reaction in volt, F = Faraday constant, R = gas constant, and T = temperature in Kelvin. Note 1 − αc = αa where subscripts a and c denote anodic and cathodic, respectively. Also, η = EmeasuredE°eq and E°eq = E°red hrEox hr, hr = half reaction, red = reduction, and ox = oxidation. Provided there is no mass transport in the rate-determining step, the overpotential associated with any given current serves only as activation energy. In Figure 1.9(a), eqn (1.7) is used to obtain the green and black curves (when negative currents are found, log i is obtained by using the absolute value) as well as the purple line in Figure 1.9(c). Figure 1.9(a) shows the relationship between log i and overpotential. This particular plot is not symmetrical since αc was chosen to be 0.4, implying that αa = 1 − 0.4 = 0.6. If αa = αc = 0.5, then the plot is symmetrical. Figure 1.9(c) shows the relationship between i (rather than log i) and overpotential.

Figure 1.9

(a) Tafel plots for anodic and cathodic branches of the current–overpotential curve for O + 1 e⇌R with αc = 0.4, 1–αc = 0.6, T = 298 K and i0 = 1 × 10−5 A cm−2. Green and black lines are fitted to eqn (1.7). (b) Empiric Tafel plots fit eqn (1.10) using the correct α. (c) The effect of α on total current (eqn (1.7)) − overpotential plots; conditions as for (a). (d) Volcano plot for the oxidation of ethyl-2-mercaptoacetate in 0.1 M NaOH on an OPG (ordinary pyrolytic graphite) electrode modified with cobalt porphyrinoids. Data obtained from Tafel plots at E = −0.250 V versus SCE. Vit B12 = vitamin B12, CoOEHPc = cobalt(Co)-octaethylhexyloxy-phthalocyanine(Pc), CoOMePc = –octamethoxy–, CoTNPPc = –tetraneopentoxy–, CoTAPc = –tetraamino–, CoTsPc = –tetrasulfonato–, CoF6Pc = –hexadecafluoro–, CoT2APP = Co-tetra(-2-aminophenyl)porphyrin, CoTsPP = –tetra(p-sulfonatophenyl)porphyrin, CoF20PP = Co-tetra(pentafluorphenyl)porphyrin. Interestingly, “nature” is still beating us with the Co-corole, Vit B12! Reproduced from ref. 46 with permission from Springer Nature, Copyright 2007.

Figure 1.9

(a) Tafel plots for anodic and cathodic branches of the current–overpotential curve for O + 1 e⇌R with αc = 0.4, 1–αc = 0.6, T = 298 K and i0 = 1 × 10−5 A cm−2. Green and black lines are fitted to eqn (1.7). (b) Empiric Tafel plots fit eqn (1.10) using the correct α. (c) The effect of α on total current (eqn (1.7)) − overpotential plots; conditions as for (a). (d) Volcano plot for the oxidation of ethyl-2-mercaptoacetate in 0.1 M NaOH on an OPG (ordinary pyrolytic graphite) electrode modified with cobalt porphyrinoids. Data obtained from Tafel plots at E = −0.250 V versus SCE. Vit B12 = vitamin B12, CoOEHPc = cobalt(Co)-octaethylhexyloxy-phthalocyanine(Pc), CoOMePc = –octamethoxy–, CoTNPPc = –tetraneopentoxy–, CoTAPc = –tetraamino–, CoTsPc = –tetrasulfonato–, CoF6Pc = –hexadecafluoro–, CoT2APP = Co-tetra(-2-aminophenyl)porphyrin, CoTsPP = –tetra(p-sulfonatophenyl)porphyrin, CoF20PP = Co-tetra(pentafluorphenyl)porphyrin. Interestingly, “nature” is still beating us with the Co-corole, Vit B12! Reproduced from ref. 46 with permission from Springer Nature, Copyright 2007.

Close modal

If the overpotential, η, is large (η > 120/n mV), eqn (1.7) simplifies to contain the anodic term only. (Remember that 1 − αc = αa.)

formula
Equation 1.8a

or in logarithmic form

formula
Equation 1.8b

Eqn (1.8a) was used to obtain the green curve in Figure 1.9(c); eqn (1.8b) generated the linear green part in Figure 1.9(a). In accordance with eqn (1.8b), at η = 0, i0 may be found from the Y-intercept of a log i versus η plot [Figure 1.9(a)].

Y-intercept = log i0
Equation 1.9

Rearrangement of eqn (1.8b) to make η the subject of the formula gives

formula
Equation 1.10
= a + b log ia
Equation 1.11

Eqn (1.11) is the empirical Tafel equation44  that dates back to 1905 for anodic (oxidation) reactions. Comparison of eqn (1.10) and (1.11) gives

formula
Equation 1.12
formula
Equation 1.13

An empirical Tafel plot utilizing experimental data would give something like Figure 1.9(b). Note that the X and Y axes in Figure 1.9(a) must just be interchanged to fit Figure 1.9(b), and one should use the absolute values of the overpotentials. One can then calculate either n or α from eqn (1.12), and i0 may be found either from the X-axis intercept or (but less accurately) from the Y-intercept using eqn (1.13).

Similar equations may be derived for cathodic (reduction) processes.

From the Tafel diagram of eqn (1.10) (see Figure 1.9, right) and eqn (1.12), the Tafel slope, b, is only dependent on the variables α and n. The larger b is, the faster the overpotential η increases. Thus, for an electrochemical reaction to obtain a large current at low overpotentials, it should exhibit a small Tafel slope b that translates to a large αn. Figure 1.9(b) demonstrates this: α for the green line is 0.6, and this line has a smaller slope than the black line that has α = 0.4.

Volcano plots [see Figure 1.9(d)] are generated when electrocatalyst activity expressed as log[(i0 (A′ cm−2)] is graphed (normally) against heats of adsorption (ΔGH in eV or other suitable units) of a reacting molecule that must react with a series of different catalysts following the same electron transfer mechanism.45  The reaction conditions under which the reactions were performed and the rate-performing step for each reaction must be the same for all catalysts used.46  If ΔGH is not available from literature sources, other parameters may be used, such as redox potentials [used in Figure 1.9(d)]46  or number of d electrons if the catalyst metal differs.47  Bond strengths between the active site and adsorbed molecule are also used often.48  Any parameter related to the ability of the catalyst to form chemical bonds with reactants, reaction intermediates or products would suffice. Volcano plots are important in catalysis and electrocatalysis since they offer a guide to identify the most active (electro)catalyst for a particular reaction, and they give clues on what structural modifications could be considered to further enhance catalytic activity. The name “volcano” plots are self-explaining. The shape of the graph looks like a volcano mountain.

When a rotating disk electrode setup is available, the Koutecky–Levich equation [eqn (1.14a, b)] may be used to obtain insight into the mechanism of electron transfer.49  One may also evaluate half reaction rates occurring at an electrode that is hampered by a combination of sluggish kinetics and mass transport.

formula
Equation 1.14a
formula
Equation 1.14b

In eqn (1.14a), i is the measured current density (A cm−2), and iK is the kinetic current. The first term in eqn (1.14a, b) is the Levich current density, ilev, which is used to measure the rate of current-limiting chemical reactions. This means that one can also rewrite eqn (1.14a) as in eqn (1.14b). The value of n in eqn (1.14a) is the number of electrons transferred in the overall electrode reaction, F is the Faraday constant (96 485.3 C mol−1), D and Co* are the diffusion coefficient (cm2 s−1) and bulk concentration of the species that undergoes electrocatalysis, ν is the kinematic viscosity (cm2 s−1) of the solvent, and ω is the angular rotation speed (rad s−1) of the electrode.

The first step in an experiment is to collect CVs at different rotation rates at the rotating disk electrode while scanning from low to high potential (or from high to low). The measured current should start at zero, rise quickly, and then flatten off as overpotentials get large. The shape of the iE curve will be sigmoidal.50  This is demonstrated in Section 1.6.3 (see Figure 1.16).

A plot of 1/i versus 1/(ω)0.5 for a number of data points at the same potential should be linear, and the number of electrons flowing at the rate determining step, n, can be calculated from the slope.49  Data collected at a large enough overpotential (i.e., from the limiting region in a current potential plot) will pass through the origin in a Koutecky–Levich plot, which implies that there are no kinetic limitations in electron transfer. Only mass transport limits the current. When currents are measured at an overpotential so small that i is not constant with an increase in potential (i.e., i still rises with increasing E in a CV), a Koutecky–Levich plot gives a non-zero intercept of 1/iK, which indicates a kinetic limitation. This non-zero intercept implies that even at infinite rotation rate (i.e., mass transport is infinite), the rate of the electrode reaction is limited by slow kinetics at the electrode surface. The intercept term 1/iK is the reciprocal kinetic current in the absence of any mass transport limitations. By measuring the kinetic current at a variety of different overpotentials along the voltammogram, it is possible49,50  to determine the standard rate constant kf for the electrochemical half reaction.

With this background the four quantities important for electrocatalyst activity may be summarized as follows:

  • a. The exchange current density i0.

  • b. The Tafel slope b (mV dec−1) obtained from a plot of overpotential η against log i. The shape of this graph is also shown in Figure 1.9, and the linear portion satisfies the equation η = a + b log i. b may also be obtained from eqn (1.12) if α is known. Or once b has been determined experimentally, eqn (1.12) may be used to obtain α.

  • c. The current density i (in units of mA cm−2) at a given overpotential. Current density may be reported in any of three ways:

    i. Amp per geometric area, A cm−2geo,

    ii. Amp per surface area, A cm−2real, and

    iii. Amp per electrochemically active surface area, A cm−2act. Here, a roughness correction must be performed (see literature42 ).

    The last option of these three is the best to use because it gives the correct turnover frequency (TOF), but the first option is often used because it is easier to obtain.

  • d. The overpotential ηi required to reach a target current density

To illustrate some of the preceding concepts, this section starts with a discussion of nickel hydroxide nanoparticles (= Ni(OH)2 NPs) supported on polycrystalline boron-doped diamond electrodes (pBDD) for electrocatalysis of oxidation reactions.51 In situ generated Ni3 + in NiIIIOOH are the active catalytic species in alcohol oxidations because empty d-orbitals or unpaired d-electrons are available for bond formation with redox intermediates, adsorbed species, or electrodes. To deposit Ni(OH)2 on pBDD electrodes, nitrate ions from the nickel source, Ni(NO3)2, were electrochemically reduced at −1.1 V versus Ag/AgCl according to the reaction

NO3 + 7H2O + 8e → NH4+ + 10OH
Equation 1.15

The electrochemically generated OH ions reacted with Ni2 + to form poorly soluble Ni(OH)2 (Ksp = 5.48 × 10−16, 25 °C), which immediately adsorbed on the electrode. This method is easy to use, is cheap, and provides the capability of controlling Ni(OH)2 NPs size by choosing appropriate times of electrolysis. Ni(OH)2 NPs were characterized by XPS, AFM, and field-emission scanning electron microscopy. XPS measurements confirmed that the NPs were void of Ni metal, NiOOH, or NiO species; it consisted only of Ni(OH)2. AFM measurements showed Ni(OH)2 NPs aggregates had conelike structures, heights of 250 nm, and base diameters of ca. 200 μm; smaller NPs had a controlled height of 10–20 nm and base size of ca. 33 nm. The amount of deposited Ni(OH)2 could be controlled and was calculated51  as an effective surface concentration Γ from eqn (1.16) as ca. 20 ± 4, ca. 140 ± 20, or ca. 420 ± 70 nmol cm−2 (Qox = charge associated with oxidation), depending on electrolysis time. FE-SEM showed that adsorbed particles are distributed uniformly over the pBDD.

Γ = Qox/nFA
Equation 1.16

Figure 1.10 shows the CV of a pBDD electrode modified with ca. 20 nmol cm−2Ni(OH)2 in water containing 0.1 M KOH. Although the ratio ipc/ipa ≠ 1, the charge ratio Qred/Qox ≈ 1 (Qred = reduction charge and Qox = oxidation charge that flowed). The redox process that takes place is

NiII(OH)2(s) + OH(aq)NiIIIOOH + H2O + e
Equation 1.17
Figure 1.10

CVs at a scan rate of 5 mV s−1 for a Ni(OH)2-modified (Γ ∼ 20 n mol cm−2) pBDD electrode recorded in 0.1 M KOH only (left) and in 0.1 M KOH + 0.5 M ethanol (right). Inset: CV of bare pBDD in a 0.1 M KOH solution containing both 1 M ethanol and 1 M methanol. Reproduced from ref. 51 with permission from American Chemical Society, Copyright 2011.

Figure 1.10

CVs at a scan rate of 5 mV s−1 for a Ni(OH)2-modified (Γ ∼ 20 n mol cm−2) pBDD electrode recorded in 0.1 M KOH only (left) and in 0.1 M KOH + 0.5 M ethanol (right). Inset: CV of bare pBDD in a 0.1 M KOH solution containing both 1 M ethanol and 1 M methanol. Reproduced from ref. 51 with permission from American Chemical Society, Copyright 2011.

Close modal

This Ni(OH)2-modified pBDD electrode was then tested as electrocatalyst in oxidation of glucose, methanol, and ethanol. In the absence of glucose or alcohol, the parasite (background) current was found to be negligible in 0.1 M KOH (inset in Figure 1.10). When 0.5 M ethanol was added to the solution, the Ni(OH)2 was electrochemically oxidized at ca. 0.22 V versus Ag/AgCl [eqn (1.17)]. The oxidized product, NiIIIOOH, then oxidizes alcohols and regenerates Ni(OH)2. The next catalytic cycle continues when the electrocatalyst NiIIIOOH is electrochemically regenerated. At a potential of 0.3 V versus Ag/AgCl, NiIIIOOH generation is fast with a large current flow, which is the result of alcohol oxidation driving the catalytic cycle. This potential (0.3 V) is so low that OH is not oxidized simultaneously with alcohol oxidation. The activity of the Ni(OH)2 NPs is high. Keeping in mind that the most efficient electrodes were covered with only ca. 2 μg Ni(OH)2 cm−2, the normalized catalyst current for ethanol oxidation of 1010 A g−1 is impressive.51 

The authors51  found that the most effective catalyst particles had cone heights of ca. 25 nm. NPs with cone heights of 12 nm are less active than NPs with cone heights of 25 nm. Increasing the NPs cone heights beyond 25 nm again caused a lowering of catalyst efficiency.

This study51  also demonstrates the properties of good electrode material. The properties must be as follows:

  • a. The electrode surface must be such that strong Ni(OH)2–electrode surface material interaction exists. Here, no covalent bond between pBDD electrode and Ni(OH)2 is formed, yet the specific adsorption of Ni(OH)2 onto pBDD led to secondary forces that are strong enough to prevent loss of the NPs under the conditions used.

  • b. The NPs itself must exhibit chemical reversibility in the redox process with which it is involved over many cycles while staying adsorbed on the electrode.

  • c. The electrode itself must be robust and chemically inert under the conditions used (temperature, solvent type, electrolytes, and other species in the solution to avoid catalyst poisoning by any products).

  • d. The potential window in which the electrode operates must also be wide enough to prevent electrochemically induced processes to shorten the active life of the catalyst/support system.

  • e. The contribution of the parasitic (background) current from the electrode in the operating potential window should be negligible.

As described,51  pBDD electrodes are well suited as inert support electrode for electrocatalysis.

The above study made use of a catalyst, Ni(OH)2, that was not covalently anchored to the pBDD electrode. In general, this is dangerous in that the secondary forces that keep the specific adsorbed particles in place on the electrode surface may be too weak to prevent catalyst particles to fall off from the electrode partially or completely under the conditions of an electrocatalysis experiment. To make sure this does not happen, one would measure the current during electrocatalysis for an extended time at a fixed potential. In the ideal case, no current loss will be observed.

The importance of doing blank studies (i.e., checking whether the electrode itself does not have catalytic activity) cannot be overstressed. Kozub and Compton52  demonstrated how to prevent interpretation of results as an electrocatalytic effect by a potential catalyst adsorbed on the electrode surface when, in fact, the electrode itself behaved as the electrocatalyst. To do this, they observed that cobalt(II) phthalocyanine, CoPc, specifically adsorbed on a suitable electrode has been used as an electrochemical sensor to detect molecules or ions such as nitrite, glutathione, hydrazine, glucose, hydrogen peroxide, and many others (43 different species have been listed where CoPc was used for detection).52  They focused on nitrite detection. The popularity of CoPc as detecting substrate specifically adsorbed onto a suitable electrode stems from a volcano plot where different MPcs and MTSPcs (M = metal, TSPcs = tetrasulfonic acid phthalocyanines) were used to detect a specific species X.47  It was found that the effectivity of electro-oxidation of cysteine by these compounds was in the M order of Co > Fe > Mn > Ni > Cu. The detection principle to detect such species (e.g., X) operates as follows:52 

CoPc⇌CoPc+ + e
Equation 1.18
CoPc+ + X → CoPc + Xoxidized
Equation 1.19

Because the potential required for the mediated oxidation of compounds or ions X to be detected is less than the potential required for the direct electrochemical oxidation of these species according to reaction (1.20), X may be detected in a cheaper and more effective electrocatalytic way.52 

X → Xoxidized + e
Equation 1.20

It was shown that the electrochemical response of an edge plane pyrolytic graphite (EPPG) electrode exhibited within 98% of the same current responses during the oxidation of nitrite than a CoPc-modified EPPG electrode in water at pH = 7.2.52  This means that CoPc is not the active catalyst during nitrite detection. That the CoPc was (initially) adsorbed on the EPPG electrode stems from the fact that ipυ plots (υ = scan rate) were linear in DMSO and water for CoII/I processes [i.e., eqn (1.6) was obeyed]. In THF, the plot ip–υ1/2 was linear, which, according to eqn (1.5), is indicative of a diffusion-controlled dissolved species undergoing redox processes. However, nitrite oxidation does not benefit from the CoII/I redox reaction. It is the CoIII/II couple that is electroactive in the oxidation of nitrite to nitrate. Further research showed that CoIIPc is adsorbed on the electrode surface in microcrystallite form.52  These crystals are insoluble in water. However, when it is oxidized to CoIIIPc+, as in eqn (1.18), it undergoes solubilization with loss of electrode surface modification. Electrode surfaces were characterized with SEM and EDS. Thus, it was clearly shown in this project that if one does not know that the electrode itself can catalyze the desired oxidation of the compound to be detected, one can easily fall into the trap of making wrong conclusions by attributing catalytic activity to the envisaged catalyst while the electrode is actually the active catalyst. As the authors so elegantly demonstrated, the modified electrode should be checked for stability under the conditions of use (here utilizing the CoIII/II couple), not conditions that activate another redox process unimportant for the task (here utilizing the CoII/I couple).

To overcome the solubility issues of adsorbed porphyrinoid electrocatalysts as described in Section 1.5.2, one can either covalently bind the catalyst to the surface of the electrode or it can be structurally modified to be less soluble, for example, by increasing porphyrinoid aggregation properties. Covalent binding of the catalyst to the electrode utilizing electrochemical methods is demonstrated in Figure 1.11.53  An amino porphyrin was converted to a diazonium salt and then reduced electrochemically either at a platinum or glassy carbon (GC) or indium tin oxide (ITO) electrode utilizing five CV scan cycles between 0.5 and 0.05 V versus SCE. Cycles between 0.5 and 0.35 V were less efficient in generating the electrode–porphyrin bond. The electrode-grafted porphyrin could be modified chemically by inserting zinc into the cavity and then again removing it with 0.2 M HCl. Surprisingly though, the grafted porphyrin was not completely stable. Scanning between 0.00 and −1.80 V versus SCE for 100 cycles at 1 V s−1 of the metal-free porphyrin-modified glassy carbon electrode led to 37% signal intensity current loss. The Pt- and ITO-modified electrodes lost 19% signal strength. Zn2 +-porphyrin-modified GC electrodes were more stable than GC electrodes modified with the metal-free porphyrin because only 27% of the anodic and cathodic peak current values were lost after 100 CV cycles. However, when modified GC electrodes had Ni2 + coordinated to the center porphyrin cavity, 55% of the signal strength was lost. Slow scan rates of 0.1 V s−1 resulted in 100% loss of porphyrin redox features after 13 cycles.

Figure 1.11

(a) In situ generation of H4(PDTP-N2+) followed by electrochemical grafting on the electrode surface. (b) First five grafting scans at GC of 1 mM H4(PDTP-N2+) in 0.1 M [NEt4][BF4]/CH3CN (ACN) containing NaNO2 and CF3COOH (υ = 50 mV s−1, rt). (c) CVs of a H2(PDTP)-modified GC electrode in 0.1 M [NEt4][BF4]/CH3CN at υ = 50 mV s−1 and rt. Thirty cycles of cathodic and anodic scans of H2(PDTP)/GC films were collected. The black CV is the second scan of thirty and the red one is of the fifteenth. Reproduced from ref. 53 with permission from John Wiley & Sons, Copyright 2016 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim.

Figure 1.11

(a) In situ generation of H4(PDTP-N2+) followed by electrochemical grafting on the electrode surface. (b) First five grafting scans at GC of 1 mM H4(PDTP-N2+) in 0.1 M [NEt4][BF4]/CH3CN (ACN) containing NaNO2 and CF3COOH (υ = 50 mV s−1, rt). (c) CVs of a H2(PDTP)-modified GC electrode in 0.1 M [NEt4][BF4]/CH3CN at υ = 50 mV s−1 and rt. Thirty cycles of cathodic and anodic scans of H2(PDTP)/GC films were collected. The black CV is the second scan of thirty and the red one is of the fifteenth. Reproduced from ref. 53 with permission from John Wiley & Sons, Copyright 2016 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim.

Close modal

In an alternative approach, the porphyrin structure may be altered to be less soluble than simple CoPc. Shown in Figure 1.12 is an electrochemically induced polymerization reaction of magnesium porphyrin. By applying 0.65 V versus SCE to the electrode, porphyrin units were linked at the meso position in a potentiostatic polymerization.54  To prevent unwanted side reactions, polymerization was stopped when the deposition charge reached 10 mC cm−2. The species pMgP-I was deposited as a thin film on the surface of the electrode. Acid treatment of this film leaves the metal-free complex p2HP-I still intact on the electrode.55  CoII can then be inserted into this film on the electrode surface to give a pCoP-I modified electrode by treatment with a saturated CH3CN solution of cobalt acetate.

Figure 1.12

Electrochemical routes to obtain different polyporphyrin structures. Type I magnesium polymer films, pMgP-I, are obtained by applying potentials less than 0.65 V on the working electrode, and type II (pMgP-II) polymer films are obtained by applying potentials larger than 0.65 V versus SCE. pMgP-I films can be demetallated using acid on the electrode to give the metal-free polymer p2HP-1 films and then again remetalled with cobalt acetate to give a pCoP-I film on the electrode. Application of 20 CV cycles between −1.80 and 1.40 V to the pCoP-I film gave the desired insoluble pCoP-II film. Reproduced from ref. 54 with permission from the Centre National de la Recherche Scientifique (CNRS) and the Royal Society of Chemistry. The structure of polymer p2HP-I as well as the CV (solid line, scan rate 100 mV s−1 in 0.1 M [nBu4N][PF6]/CH3CN(= AN)) and specific conductance (dotted line) of pCoP-I films on a Pt disk electrode. Reproduced from ref. 55 with permission from Elsevier, Copyright 2016.

Figure 1.12

Electrochemical routes to obtain different polyporphyrin structures. Type I magnesium polymer films, pMgP-I, are obtained by applying potentials less than 0.65 V on the working electrode, and type II (pMgP-II) polymer films are obtained by applying potentials larger than 0.65 V versus SCE. pMgP-I films can be demetallated using acid on the electrode to give the metal-free polymer p2HP-1 films and then again remetalled with cobalt acetate to give a pCoP-I film on the electrode. Application of 20 CV cycles between −1.80 and 1.40 V to the pCoP-I film gave the desired insoluble pCoP-II film. Reproduced from ref. 54 with permission from the Centre National de la Recherche Scientifique (CNRS) and the Royal Society of Chemistry. The structure of polymer p2HP-I as well as the CV (solid line, scan rate 100 mV s−1 in 0.1 M [nBu4N][PF6]/CH3CN(= AN)) and specific conductance (dotted line) of pCoP-I films on a Pt disk electrode. Reproduced from ref. 55 with permission from Elsevier, Copyright 2016.

Close modal

This procedure was demonstrated on GC, Pt, and ITO electrodes. Oxidation of the thin pCoP-I film to a pCoP-II film having a ladder structure was achieved by cycling the electrode potential 20 times between −1.80 and 1.40 V versus SCE.54  Films of pCoP-II were characterized by XPS and SEM and, of course, electrochemically.

It was found that the (pCoP-II)-modified electrodes from the previous section could be used for oxidation and sensing of sulfite ions in aqueous solution (see Figure 1.13). The limit of quantification was 0.649 mmol dm−3, and the limit of detection was 0.195 mmol dm−3. This is about twice the sensitivity of glassy carbon electrodes modified with a copolymer of Co-tetrakis(para-aminophenyl)porphyrin and ortho-phenylene-diamine developed earlier by Arce and coworkers.56  Detection of sulfites is an important analytical experiment since sulfites are used to prolong the shelf life of many food stuffs, including red wine.

Figure 1.13

Top: The ladder polymer (pCoP-II)-modified GC electrode can be used for sulfite detection. The CVs shown were from a ca. 100 nm thick pCoP-II film at 100 mV s−1 scan rate of solutions of SO32 − in 0.1 M NaNO3/water. [SO32 −] = 0 (smallest current), 2, 4, 6, 8, and 10 mM (largest current). Plots of current intensity of oxidation peak O2(S)versus [SO32 −] (not shown) were found to be linear. Bottom left: CVs obtained at a scan rate of 100 mV s−1 in 0.1 M NaNO3/water on GC or on the (pCoP-II)/GC-modified electrode in the absence and presence of 10 mM of sulfite. Bottom right: A consecutive 4 cycle CV of the (pCoP-II)/GC-modified electrode in a 10 mM sulfite solution (NaNO3 0.1 M, υ = 0.1 V s1). Note that wave O1(S) is only present in the first cycle. Reproduced from ref. 54 with permission from the Centre National de la Recherche Scientifique (CNRS) and the Royal Society of Chemistry.

Figure 1.13

Top: The ladder polymer (pCoP-II)-modified GC electrode can be used for sulfite detection. The CVs shown were from a ca. 100 nm thick pCoP-II film at 100 mV s−1 scan rate of solutions of SO32 − in 0.1 M NaNO3/water. [SO32 −] = 0 (smallest current), 2, 4, 6, 8, and 10 mM (largest current). Plots of current intensity of oxidation peak O2(S)versus [SO32 −] (not shown) were found to be linear. Bottom left: CVs obtained at a scan rate of 100 mV s−1 in 0.1 M NaNO3/water on GC or on the (pCoP-II)/GC-modified electrode in the absence and presence of 10 mM of sulfite. Bottom right: A consecutive 4 cycle CV of the (pCoP-II)/GC-modified electrode in a 10 mM sulfite solution (NaNO3 0.1 M, υ = 0.1 V s1). Note that wave O1(S) is only present in the first cycle. Reproduced from ref. 54 with permission from the Centre National de la Recherche Scientifique (CNRS) and the Royal Society of Chemistry.

Close modal

Tests showed that a bare glassy carbon electrode in 0.1 M NaNO3/water was inactive in the absence or presence of 10 mM sulfite (see Figure 1.13). The (pCoP-II)-modified electrode showed in the absence of any sulfite only a redox wave associated with adsorbed CoII at ca. 0.240 V versus SCE, but in the presence of 10 mM SO32 −, sulfite oxidation was observed at 0.46 [wave O1(S)] and 0.79 V [wave O2(s)] at a pH of ca. 9.7.54 

Wave O1(S) is only present in the first of multi-sweep CVs (see Figure 1.13, bottom right). However, when the cell is allowed enough time between scans to homogenize, wave O1(S) is observed on every scan. Furthermore, wave O1(S) varies linearly with scan rate, while wave O2(S) varies linearly with (scan rate)1/2. Results were consistent with wave OI(S) representing oxidation of sulfite, which, prior to commencing the scan, diffused into the polymer film pores and was trapped there. Thus, oxidation of these sulfite ions is not dependent on diffusion; hence, a linear I/υ relationship was observed. It can only be observed in the first scan because sulfite diffusion into the polymer film is not fast enough to allow observation of wave O1(S) in successive uninterrupted CV scans. In contrast, wave O2(S) is considered the wave associated with electrocatalytic oxidation of dissolved sulfite that can diffuse freely to the surface of the (pCoP-II)-modified glassy carbon electrode; hence, its oxidation shows a linear dependence for the I/υ1/2 relationship.54  (pCoP-I)- and (pCoP-II)-modified Pt and glassy carbon electrodes were also found to be effective oxygen reduction electrocatalysts but will not be discussed here.57 

A discussion of electrocatalysis that does not include reduction of CO2(g), O2(g), and production of O2(g) and H2(g) from water (i.e., water splitting) would be incomplete because these processes are very important to fuel production and energy storage and a huge research effort has already gone into it. Important reactions in fuel production and consumption industries that would benefit from improved electrocatalysis include:

2H+ + 2e H2
Equation 1.21
CO2 + nH+ + ne CxHy + CaHbOcH + mH2O
Equation 1.22
2H2O O2 + 4H+ + 4e
Equation 1.23
CO2 + 2H+ + 2e CO + H2O
Equation 1.24

On the right-hand side of eqn (1.21)(1.24) are fuels, and on the left-hand side are the products of fuel consumption. Eqn (1.24) is included in the array of equations above since the Fischer–Tropsch reaction [eqn (1.25)] involves formation of paraffins (alkanes, including petrol and waxes), α-olefins (alkenes), and water.

3nCO + (6n + 1) H2 → CnH2n + 2 + CnH2n + (3n−1) H2O
Equation 1.25

CO may be generated from CO2 electrocatalytically and by reacting it with H2. Utilizing the sun as an electrical energy source (solar power) would further assist the energy requirements of an energy-hungry future world.

We begin the story with the reduction of CO2(g) to CO(g), but it can also be reduced to CH4(g), C2H4(g), CH3OH(g), or HCOOH(ℓ), to name only a few. CO is a feedstock material in the Fischer–Tropsch reaction and hydroformylation to produce liquid hydrocarbon fuels.

Often a misconception in electrocatalysis is that the electrocatalyst must be deposited on the surface of an electrode. This conception is incorrect. Electron transfer takes place at the electrode, and as long as a catalyst has access to the surface of the electrode, even from a homogeneous solution through diffusion, electrocatalysis can still take place. Officer58  demonstrated this recently with the elegant electroreduction of CO2 to CO catalyzed by a dissolved Fe(0) tetraphenylporphyrin/ionic liquid co-catalytic system in the presence of trifluoroethanol (TFE). [Fe0TPP]2− is generated from FeIIITPPCl by an in situ electrochemical reduction. To place the energy saving this brought about in perspective, one should recognize that simple CO2 reduction generating the first intermediate CO2 according to the reaction CO2 + e⇌CO2˙ takes place at −1.97 V in DMF versus NHE. Addition of protons to the reaction leads to less energy-consuming reductions.58,59  The electrode potential under acidic conditions falls with ca. 0.4 V for the reaction CO2 + 2 e + 2 H+⇌CO + H2O. By performing the reaction in the ionic liquid 1-butyl-3-methylimidazolium tetrafluoroborate, [BMIM][BF4], a further energy saving was achieved. [BMIM][BF4] does not ease the reduction potential of CO2. That is the job of the weak acid trifluoroethanol (TFE). As shown in Figure 1.14, top right, addition of [BMIM][BF4] to the reaction medium led to a decrease of 0.23 V (from −1.63 to −1.40 V versus NHE) in the reduction potential of the [FeITPP]/[Fe0TPP]2− couple, probably through an electrostatic interaction between [BMIM]+ cations, the anions [FeITPP] and [Fe0TPP]2−. The CV wave at ca. −0.98 to −1.06 V versus NHE corresponds to the [FeIITPP]/[FeITPP]2− couple (best seen in the inset), and the wave at −0.14 to −0.23 V is associated with the [FeIIITPP]/[FeIITPP]2− couple. Use of 0.3 M [BMIM][BF4] in DMF containing 1 M TFE led to a lowering of the onset potential of CO2 reduction where CO could first be detected from −1.51 V to −1.36 V versus NHE. Gas evolution was measured by gas chromatography. This corresponds to a decrease of the overpotential η from 0.82 to 0.67 V relative to the equilibrium CO2/CO potential of E° = 0.69 V in DMF.58 

Figure 1.14

Shown clockwise from top left is the structure of FeIIITPPCl, CVs at scan rate of 50 mV s−1 of 0.5 mM FeIIITPPCl; 0.5 mM FeIIITPPCl + 1 M TFE; 0.5 mM FeIIITPPCl + 0.3 M [BMIM]BF4; 0.5 mM FeIIITPPCl + 1 M TFE + 0.3 M [BMIM]BF4 in DMF + 0.1 M TBAPF6 at glassy carbon electrode (a) under Ar (inset shows an expansion of the active area); (b) under CO2 and, finally, Tafel plots of CO formation for 0.5 mM FeIIITPPCl + 1 M TFE (black) and 0.5 mM FeIIITPPCl + 1 M TFE + 0.3 M [BMIM]BF4 (red). Reproduced from ref. 58 with permission from John Wiley and Sons, Copyright 2016 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim.

Figure 1.14

Shown clockwise from top left is the structure of FeIIITPPCl, CVs at scan rate of 50 mV s−1 of 0.5 mM FeIIITPPCl; 0.5 mM FeIIITPPCl + 1 M TFE; 0.5 mM FeIIITPPCl + 0.3 M [BMIM]BF4; 0.5 mM FeIIITPPCl + 1 M TFE + 0.3 M [BMIM]BF4 in DMF + 0.1 M TBAPF6 at glassy carbon electrode (a) under Ar (inset shows an expansion of the active area); (b) under CO2 and, finally, Tafel plots of CO formation for 0.5 mM FeIIITPPCl + 1 M TFE (black) and 0.5 mM FeIIITPPCl + 1 M TFE + 0.3 M [BMIM]BF4 (red). Reproduced from ref. 58 with permission from John Wiley and Sons, Copyright 2016 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim.

Close modal

The CVs on the right were recorded under argon (top) and CO2 (bottom). Current densities were increased from 1.14 mA cm−2 when a potential of −1.75 V versus NHE was applied to systems in the absence of [BMIM][BF4] but increased to 3.26 mA cm−2 at an applied potential of −1.56 V for the system containing [BMIM][BF4]; both contained TFE. The electrode was stable over 4 h, and turnover frequencies (TOFs) were calculated as 47 (at −1.75 V, no ionic liquid) or 190 s−1 (at −1.56 V in the presence of ionic liquid). Turnover numbers were found to be 6.77 × 105 or 27.4 × 105, respectively, after 4 h.58 

Tafel plots (Figure 1.14) had slopes of 197 mV dec−1 for the system containing 0.3 M [BMIM][BF4] and 286 mV dec−1 for the system without any [BMIM][BF4]. The smaller slope indicated that the rate of electron transfer in the presence of [BMIM][BF4] is much faster (Section 1.4) and highlighted the important contribution of ionic liquids to this reaction.58 

Regarding hydrogen evolution reactions (= HERs), Wu and coworkers60  demonstrated that β-hydrogenation of porphyrin followed by NiII insertion to give nickel(II) porphyrin, N-1, chlorin, N-2, porpholactone, N-3, and β-hydroporpholactone, N-4 (see Figure 1.15) increased TOF values for H2(g) evolution from CH3CN/CF3CO2H from 265 for N-1 to 6287 s−1 for N-2 (i.e., a 24-fold increase in activity) and from 342 for Ni-3 to 1737 s−1 for N-4 (a 5-fold increase in activity). DFT calculations showed that after initial 2 e reduction to generate the pre-catalytic species [Ni-1]2− and [Ni-2]2−, the β-hydrogenated species [Ni-2]2− has a more electron-rich Ni center than [Ni-1]2−. Interaction of CF3CO2H with [Ni-2]2− then forms the active catalyst, which is a nickel hydride species, while in [Ni-1]2−, a protonated pyrrole species is generated. The nickel hydride catalyst generated from [Ni-2]2− then reacts much faster (k2 ≈ 3.03 M−1 s−1) with an additional CF3CO2H species than the pyrrole-protonated catalyst from [Ni-1]2− (k2 ≈ 0.0917 M−1 s−1) to generate H2(g) with TOF values that are 24 times larger for Ni-2 than Ni-1.

Figure 1.15

Synthesis of Ni-1 to Ni-4 complexes modified at the β-pyrrole positions, R = C6F5. Bottom left: Successive cyclic voltammograms at 25 °C and 100 mV s−1 of 1.0 mM Ni-2 in CH3CN containing 0.10 M [(nBu4)N][PF6] at increasing concentrations of TFA utilizing a 3 mm glassy-carbon working electrode. Bottom right: TOF values for Ni-2 flattens at ca. 6290 s−1 at [TFA] > 450 mM. In this region, TOF(Ni-2) ≈ 24TOF(Ni-1). Reproduced from ref. 60 with permission from the Royal Society of Chemistry.

Figure 1.15

Synthesis of Ni-1 to Ni-4 complexes modified at the β-pyrrole positions, R = C6F5. Bottom left: Successive cyclic voltammograms at 25 °C and 100 mV s−1 of 1.0 mM Ni-2 in CH3CN containing 0.10 M [(nBu4)N][PF6] at increasing concentrations of TFA utilizing a 3 mm glassy-carbon working electrode. Bottom right: TOF values for Ni-2 flattens at ca. 6290 s−1 at [TFA] > 450 mM. In this region, TOF(Ni-2) ≈ 24TOF(Ni-1). Reproduced from ref. 60 with permission from the Royal Society of Chemistry.

Close modal

The potential at which H2(g) forms for the Ni-2 catalyst range from ca. −1.75 V at [TFA] = 100 mM to ca. −2.2 V versus Fc/Fc+ (see Figure 1.15, bottom left), which is slightly better than what was found61  for Pd and Cu meso-substituted tetraferrocenylporphyrin (− 2.0 until −2.4 V versus Fc/Fc+), but TOFmax values were still very good (ca. 6000). Despite the high TOF values, by utilizing the standard potential estimates of the H+/H2 couple in organic media,62,63  overpotentials are in the region of 0.8–1.2 V. This is very high. Good catalysts should have high turnover at low (< 0.5 V) overpotentials. It follows that the above catalysts will have a low power:chemical conversion ratio and need further refinement for industrial applications.

As far as reduction of O2(g) is concerned, the work of Nemykin and Kadish comes to mind.64  These authors used a series of three different meso-ferrocenylated cobalt porphyrins, (Fc)n(p-CH3Ph)4−nPorCo with n = 0 − 4, adsorbed on an EPPG rotating disk as well as a rotating ring-disk electrode in DMF to reduce oxygen electrocatalytically. The CoII/III couple provides the electrons to reduce O2 to H2O2 selectively under acidic conditions in a two-electron process (see Figure 1.16). Hydrogen peroxide is a promising candidate as a sustainable energy carrier because it has a high energy density and emits no CO2. It is, therefore, an environmentally friendly oxidant. Fuel cells running on hydrogen peroxide produce only water and oxygen as waste products after power generation, and over the years, cell performance has been improved significantly.65  Interestingly, in the Nemykin study,64  the electron-donating properties of the Fc groups lowered the potential at which O2 reduction took place but did not participate in the electron transfer steps to O2. Because of its bulky size, the Fc group did prevent dimerization of porphyrin intermediates; thus, selective oxidation of O2 to only H2O2 was possible. In the absence of reactants that prevent dimerization of intermediates, four-electron oxidation to H2O often occurs selectively66  or a mixture of two- and four-electron oxidation takes place.64,67 

Figure 1.16

(a) Fc4PorCo reduces O2 selectively to H2O2. (b) Sigmoidal current–voltage curves at a rotating disk electrode coated with Fc(p-CH3Ph)3PorCo in 1.0 M HClO4 saturated with air at the indicated rotation rates (ω). Potential scan rate = 50 mV s−1. (c) Koutecky–Levich plot for the catalyzed O2 reduction by Fc(p-CH3Ph)3PorCo. Reproduced from ref. 64 with permission from American Chemical Society, Copyright 2014. (d) Unsubstituted CoPc adsorbs in thicker films on CCG (chemically converted graphene) than octa(octyloxy)-substituted CoPc. (e) In sequence, bright field TEM image showing the morphology of the CoPc-A-modified CCG electrode and EDS mapping of a homogeneous distribution of C and Co over the composite material. Reproduced from ref. 73 with permission from American Chemical Society, Copyright 2019.

Figure 1.16

(a) Fc4PorCo reduces O2 selectively to H2O2. (b) Sigmoidal current–voltage curves at a rotating disk electrode coated with Fc(p-CH3Ph)3PorCo in 1.0 M HClO4 saturated with air at the indicated rotation rates (ω). Potential scan rate = 50 mV s−1. (c) Koutecky–Levich plot for the catalyzed O2 reduction by Fc(p-CH3Ph)3PorCo. Reproduced from ref. 64 with permission from American Chemical Society, Copyright 2014. (d) Unsubstituted CoPc adsorbs in thicker films on CCG (chemically converted graphene) than octa(octyloxy)-substituted CoPc. (e) In sequence, bright field TEM image showing the morphology of the CoPc-A-modified CCG electrode and EDS mapping of a homogeneous distribution of C and Co over the composite material. Reproduced from ref. 73 with permission from American Chemical Society, Copyright 2019.

Close modal

Use of a rotating disk electrode gave sigmoidal CVs as shown in Figure 1.16(b). By plotting reciprocal i (written as j in this publication) against 1/(ω)0.5 [see Figure 1.16(c)], the number of electrons flowing to O2 at the electrode surface could be calculated using eqn (1.14a) as 2.1 (i.e., ≈ 2) even though the CoIII/II catalyst is involved in a one-electron redox process. This is consistent with H2O2 forming 95% selectively. The rotating ring-disk electrode gave slightly lower selectivity: 88% of products were H2O2.

Identification of the CoIII/II couple in the CVs came from spectro-electrochemical measurements. In a nutshell, the wavelength of maximum absorption of the Soret band of the studied complexes, which lies in the UV/vis wavelength region 400 < λmax < 450 nm, shifted by as much as 21 nm to longer wavelengths when CoII was oxidized to CoIII. In contrast, it stayed at the same wavelength when the Fc group(s) were oxidized. When the ferrocenyl groups were oxidized, the Soret band absorbance maximum fell by ca. 13%, but when ring-centered oxidation took place, the absorbance maximum of the Soret band fell more than 65%. This is an example of the use of spectro-electrochemistry to identify which electrochemical oxidation or reduction wave belongs to which part of a molecule. One could also have used DFT calculations to obtain the same information.10 

Although the Fc group did not participate in the electron transfer process of Figure 1.16, free small ferrocene derivatives can be involved with two-electron O2 reduction to H2O2 or four-electron reduction to H2O. Thus, as reactant partners, a cobalt chlorin porphyrinoid, CoII(Ch), catalyst with either dimethylferrocene (Me2Fc) or Br2Fc as reductants reduced O2 in acidic media to H2O2 with TON numbers exceeding 30 000 at low overpotentials.65  Cobalt corrole porphyrinoids also generate H2O2 in a two-electron acidic reduction of O2,68  while both catalytic two-electron69  and catalytic four-electron70  reduction of O2 by free small ferrocene derivatives in the presence of non-porphyrinoid copper catalysts generate H2O2 or H2O selectively.

In terms of water splitting reactions, energy requirements get a bit more demanding. Water is the biggest natural source available to generate H2, H+, and O2. O2 forms during a four-electron oxidation of water [eqn (1.26)].

2H2O ⇌ O2 + 4H+ + 4e
Equation 1.26

However, the oxygen evolution reaction (OER) above requires much energy since it occurs at potentials close to 2 V versus RHE. Not much success with porphyrinoids as electrocatalysts have been achieved in the OER reaction; hence, only two of them are discussed briefly. Cobalt phthalocyanine (CoPc, see Figure 1.16) and the perfluorinated derivative CoPcF16 have been shown71  to electrocatalyze reaction 1.26 when adsorbed onto a hematide (a photoanode material) film previously evaporated onto fluorine-doped tin oxide glass (FTO). The photoanode material allows photocatalysis to work in tandem with electrocatalysis during the OER. CVs showed that the onset potential of OERs for CoPcF16- and CoPc-modified FTO electrodes was ca. −1.67 and −1.78 V versus RHE, respectively. Irradiation of a CoPcF16/hematite/FTO-modified electrode with 5 s interval chopped light from an AM1.5G solar illumination source lowered the OER onset potential to ca. 0.8 V versus RHE, and the current density increased to ca. 1 mA cm−2 at 1.23 V versus RHE. This current density is more than twice as high as that of a known cobalt–phosphate/hematite water oxidation photocatalyst system under the same conditions.71  The CoPcF16/hematite/FTO-modified electrode was stable for 60 000 s at 1.7 V versus RHE under illumination; Faradaic efficiency was close to 100%. Oxygen was measured with an in-line micro gas chromatograph. In another report,72  a CoPcF16/FTO-modified electrode operating at 2 V versus RHE (OER onset potential = 1.7 V) for 8 h (TON ≈ 1 × 105) at pH = 7 produced about 4 μmol O2 h−1 cm−2 at a turnover frequency (TOF) of 3.6 s−1. This is an efficient electrocatalytic process with current knowledge, but it is not spectacular. It means that a 1 cm2 electrode produces over 1 h less than 0.1 mL O2 gas.

One of the problems with heterogeneous catalysis, which includes electrocatalysis by material adsorbed on the surface of an electrode, is particle size. If the particles are large, there is a large portion of the material in the interior of these particles to which reacting molecules never gain access. It means that the measured activity of an electrocatalyst is much less than the potential (theoretical) activity simply because large portions of a catalyst are hidden (or “dead”) in the interior of large particles. Compounds that easily aggregate, as many phthalocyanine porphyrinoids do, are especially prone to this limitation. To overcome it, CoPc and octa(octyloxy)CoPc (the latter is labeled CoPc-A in Figure 1.16) were prepared and adsorbed onto CCG (chemically converted graphene) electrodes. Then electrocatalyst activity for these two materials was compared in the reduction of CO2 to CO.73  Characterization of the surfaces was with UV–vis, Raman, XPS, and TEM coupled with energy-dispersive X-ray spectroscopy (TEM–EDS). Bright field TEM images showed the morphologies of a bare CCG electrode, and both electrodes modified with layers of phthalocyanine adsorbed on it were the same [see Figure 1.16 bottom, inset (e), left]. EDS mapping of the CoPc-A-modified CCG electrode showed a homogeneous distribution of Co, C [see Figure 1.16, inset (e), middle and right], and N over the composite material.

Integration of the CV CoII/I oxidation allowed estimation of the amount of electrochemical active CoPc or CoPc-A molecules on the CCG electrode surface as 1.4 × 10−9 and 9.0 × 10−10 mol cm−2, respectively. This led to a larger observed total current density (e.g., ca. 9 mA cm−2 if the applied potential = −0.8 V versus RHE) for the CoPc-A-modified CCG electrode compared with the CoPc-CCG-modified electrode (ca. 6 mA cm−2 under identical conditions) during electrocatalytic CO2 reduction. In addition, the faradaic efficiency for CO formation for the CCG/CoPc-A-modified electrode is ca. 77% at −0.59 V, while that of the CCG/CoPc-modified electrode is only 63% at the same potential.73  These results are consistent with an interpretation that the long-chain octyloxy substituents of CoPc-A prevented Pc aggregation much better than aggregation is prevented by simple CoPc.

The authors73  developed the concept of effective turnover frequency (eTOF) as a measure of the minimum activity of an aggregated catalyst. The lower aggregation of CoPc-A on the graphene sheets resulted in an effective catalytic activity per single CoPc-A molecule of ∼5 s−1 at −0.59 V compared to an eTOF of ca. 2 s−1 at −0.59 V for the CCG/CoPc-modified electrode. In addition, the CoPc-A-modified CCG electrode exhibited long-term stable CO production with eTOF increasing to ∼6 s−1 and eTON ∼6.7 × 105 for CO evolution over 30 h of electrolysis. The compound minimizing aggregation, octa(octyloxy)CoPc = CoPc-A, was thus [(5 to 6)/2] = more than 2.5 times more active as the compound that aggregated strongly, CoPc. The authors also showed that CoPc-A was more active than almost all other state-of-the-art CO2 reduction electrocatalyst systems.73  The idea of using long-chain substituents to prevent or minimize aggregation is, however, not new. Cook27  determined the onset aggregation concentration of non-peripherally substituted phthalocyanines as a function of substituent chain length, and others used this information to overcome aggregation in spectroscopic and electrochemical studies.10,26,37 

The last support system discussed is the covalent organic framework (COF) system where CoPc was covalently bound into a cavity-forming structure (see Figure 1.17).74,75  Films of these COFs were adsorbed onto an HOPG electrode (HOPG = highly oriented pyrolytic graphite76 ), characterized thoroughly and then tested in the reduction reaction of CO2 to CO. The most active CO2 reduction film was found to have R1 = R2 = F. For this complex, CO2 was reduced at a low (550 mV) overpotential with 87% faradaic efficiency. When a potential of −0.67 V versus RHE was applied, the current density achieved was 65 mA mg−1 COF.74  Derivatives with four rather than three phenyl rings in the linking unit, linking CoPc functionalities together and being void of F substituents, were slightly less active.75 

Figure 1.17

Top. Structure of COF films on HOPG for CO2 reduction. Reproduced from ref. 74 with permission from American Chemical Society, Copyright 2018. Bottom. Design strategies for porphyrinoid MOF electrocatalysts. Reproduced from ref. 78 with permission from John Wiley & Sons, Copyright 2018 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim.

Figure 1.17

Top. Structure of COF films on HOPG for CO2 reduction. Reproduced from ref. 74 with permission from American Chemical Society, Copyright 2018. Bottom. Design strategies for porphyrinoid MOF electrocatalysts. Reproduced from ref. 78 with permission from John Wiley & Sons, Copyright 2018 WILEY-VCH Verlag GmbH & Co. KGaA, Weinheim.

Close modal

In contrast to COFs, MOFs (metal–organic frameworks) are not held together by covalent bonds. Building blocks are held together by coordination bonds that are not as strong as covalent bonds. However, this does not mean MOFs are excluded from electrocatalysis. For example, a redox-active iron porphyrin was built into the robust Zr-based PCN-223-Fe MOF, [Zr6O4(OH)4(Fe(III)-(TCPP)3] with TCPP = meso(tetracarboxyphenyl)porphyrin)] and films of it adsorbed onto an FTO substrate.77  This working electrode was then used for the O2 reduction reaction (the ORR), and it was found that the MOF stayed intact and was capable of reducing O2 94% selectively to water in a four-electron process. Only 6% of products were found to be H2O2, which was generated in a two-electron reduction of O2 (compare Section 1.6.3). The current density was ca. 7.5 mA at 0.8 V versus NHE. A review concerning design strategies for MOF electrocatalysts is available.78 

This work started with some fundamentals of electrochemistry and showed how it is applied in simple solution-phase electrochemical experiments. These were all outer-sphere processes that have little or no interaction between the electrode surface and the redox species. Such redox processes are fully described by formal reduction potentials, E°′, standard rate constants, k°, charge transfer coefficients, α, and double layers next to an electrode surface as defined by standard Butler–Volmer treatments.79  Questions that these studies answer tend to focus on the influence structural and electronic changes the redox species has on its electrochemical properties, as well as the influence of solvent and electrolyte interactions with redox species. The chapter then progressed into the realm of inner-sphere redox processes. The chapter discussed some theoretical background and presented case studies. Here, the bonding or adsorption of reactants, intermediates, and products to the electrode surface enhances the rate of electrode reactions by several orders of magnitude and is the gateway to electrocatalysis. Inner-sphere redox processes are characterized by overpotentials, η, Tafel slopes (which is an indicator of rates of electron transfer in the rate-determining mechanistic step) and exchange current densities at equilibrium (also an indicator of charge transfer kinetics) as defined in the Tafel equation.79  Several examples of the application of these concepts were provided.

While all of these aspects remain important in a modern era, leading future studies need to strive to bring additional information to the party. These involve more information of how electrode interactions with adsorbed species enhance electrode reactions and what structural properties of the electrode as well as the adsorbed particle will enhance catalytic reactions. One needs to understand why certain changes will have no influence on electrocatalysis, what types of changes will actually be detrimental, and why other changes will be very beneficial to electrocatalysis. The emphasis is on the “why”. It means graduating from a black box approach, which simply reports “if you do ‘this,’ ‘this’ happens” to understanding “why” it happens. The new leading modern publication should report “why” it happens.

Toward realization of this goal, more use of heterogeneous surface characterization techniques applied to electrode surfaces and adsorbed particles becomes increasingly important. Spectroscopic techniques in a multidisciplinary approach to enhance insight and understanding already play a big role. The chapter tried to show how XPS, TEM and TEM-EDS, AFM, SEM, FE-SEM, and DFT calculations brought new insights and enhanced the quality of results. This is not an exhaustive list of possible techniques to utilize. Additionally, electrochemical scanning tunneling microscopy, SEM, scanning ion conductance microscopy, and PXRD come to mind; the list is almost endless. Understanding how adsorbed nanoparticles may diffuse, aggregate, or sinter during electrocatalysis and how it influences the electrocatalytic reaction is another important aspect of research. The properties of electrodes are another topic that needs to be researched thoroughly. It is not any more cutting-edge science to report that on platinum you get “this,” on glassy carbon you get “that,” and on single-walled carbon nanotubes or graphene, you get “another” result. One needs to develop knowledge on why these results are expected. One also needs to understand the effect of composite electrode material, grain boundaries in such electrodes, and the effect defects on the surface of single crystals have on electrode performance.79  This is only possible if several multidisciplinary labs collaborate on projects. This is also where the bigger funding opportunities are.

The days of a single researcher doing it all as described above are gone. We hope that this chapter has guided, challenged, and inspired new research avenues and collaborations for the benefit of all interested in electrochemistry and electrocatalysis.

The authors acknowledge the Central Research Fund of the University of the Free State, Bloemfontein, South Africa for financial support.

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