Electrochemical Investigation of Iron-Catalyzed Atom Transfer Radical Polymerization

Use of iron-based catalysts in atom transfer radical polymerization (ATRP) is very interesting because of the abundance of the metal and its biocompatibility. Although the mechanism of action is not well understood yet, iron halide salts are usually used as catalysts, often in the presence of nitrogen or phosphorous ligands (L). In this study, electrochemically mediated ATRP (eATRP) of methyl methacrylate (MMA) catalyzed by FeCl3, both in the absence and presence of additional ligands, was investigated in dimethylformamide. The electrochemical behavior of FeCl3 and FeCl3/L was deeply investigated showing the speciation of Fe(III) and Fe(II) and the role played by added ligands. It is shown that amine ligands form stable iron complexes, whereas phosphines act as reducing agents. eATRP of MMA catalyzed by FeCl3 was investigated in different conditions. In particular, the effects of temperature, catalyst concentration, catalyst-to-initiator ratio, halide ion excess and added ligands were investigated. In general, polymerization was moderately fast but difficult to control. Surprisingly, the best results were obtained with FeCl3 without any other ligand. Electrogenerated Fe(II) effectively activates the dormant chains but deactivation of the propagating radicals by Fe(III) species is less efficient, resulting in dispersity > 1.5, unless a high concentration of FeCl3 is used.


Introduction
Reversible deactivation radical polymerization (RDRP) techniques allow for the preparation of polymeric materials with precisely tailored architectures, low dispersity and well-preserved chain end functionality [1]. Among these methods, the most important and widely used techniques are atom transfer radical polymerization (ATRP [2][3][4]), reversible addition fragmentation chain transfer polymerization [5,6] and nitroxide-mediated polymerization [7,8]. The success of these methods relies on establishing a dynamic equilibrium between the propagating radical and a dormant species, which drastically reduces the rate of termination reactions, resulting in a remarkable increase in radical lifetime.
ATRP uses a transition metal complex, mainly Cu with a nitrogen-based ligand [3], to establish the equilibrium between a dormant polymer chain and the corresponding active radical. A metal complex at a low oxidation state (Mt z /L) reacts with an alkyl halide (RX) initiator or a halogen-capped dormant polymer chain (P n -X) to produce a propagating radical (P n • ) and the metal complex at a higher oxidation state (X-Mt z+1 /L) with the halide ion as an additional ligand (Scheme 1). This activation step occurs according to an inner-sphere electron transfer mechanism [9][10][11]. The radical propagates for a short period before it is quenched by the deactivator complex X-Mt z+1 /L to form a dormant species. The equilibrium constant, K ATRP = k act /k deact , defines the equilibrium concentration of Scheme 1. General mechanism of ATRP with a reduction loop for activator regeneration (X = Cl, Br).
Several advanced variants of ATRP have been developed to reduce catalyst loading and improve the industrial attractiveness of the process. In these methods, the process starts with an oxygen-stable metal catalyst at a high oxidation state, while the activator catalyst-the metal at a lower oxidation state-is (re)generated in situ through various approaches [12]. Examples of these methods include activators regenerated by electron transfer (ARGET) ATRP [13][14][15] and supplemental activator and reducing agent (SARA) ATRP (also known as SET LRP) [16][17][18], which use a homogeneous reducing agent and a zero-valent metal, respectively, and methods based on external stimuli via electrochemistry [19,20], photochemistry [21][22][23] and sonochemistry [24,25]. Metal-free photo-induced ATRP is also widely used [26,27].
Although Cu is the most popular metal in ATRP, polymerizations have successfully been carried out with catalysts based on a large variety of other metals, such as Re, Fe, Ru, Os, Rh, Co, Ni and Pd [28]. Of these, iron has attracted considerable interest because it presents several advantages over copper catalysts. Indeed, iron is an extremely abundant metal in the Earth's crust and biocompatible iron-based systems can potentially be used for biomedical applications [29,30]. Therefore, iron can be considered a very promising metal in view of industrial development of ATRP. Like Cu and other metal catalysts, the activity of Fe-based catalysts strongly depends on the nature and structure of the ligands [31]. Most Fe catalysts used in ATRP are based on iron halide salts, FeX 3 and/or FeX 2 , in the absence or presence of added ligands, mainly based on nitrogen and phosphorous [32].
Herein, we report a deeper investigation on the electrochemical behavior of the most common iron-based catalysts used in ATRP. The study also addresses the role of added ligands, such as chloride anions, amine-bis(phenolate), phosphines and amines (structures in Scheme 2), on eATRP of mMA. halide salts, with or without tris(4-methoxyphenyl)-phosphine as an additional ligand, respectively.
Herein, we report a deeper investigation on the electrochemical behavior of the most common iron-based catalysts used in ATRP. The study also addresses the role of added ligands, such as chloride anions, amine-bis(phenolate), phosphines and amines (structures in Scheme 2), on eATRP of MMA.  Figure 1 shows cyclic voltammetry (CV) of FeCl3 in DMF + 0.1 M Et4NBF4, recorded at different scan rates. A well-defined cathodic peak (B) around −0.5 V and a shoulder (A) around −0.2 V are observed in the cathodic scan, while two anodic peaks, labelled A' and B', are observed upon scan reversal. The voltametric behavior of the iron salt strongly depends on scan rate (v). The cathodic peak B and the shoulder shift to more negative potentials when v is increased while the normalized peak current of peak B, Ip(B)/v 1/2 , shows only a slight decrease. Conversely, major changes with v occur in the reverse scan.

Cyclic Voltammetry of FeCl3
As v increases, peak A′ increases at the expense of peak B′, which instead decreases. It is clear that more than one Fe(III) species is present in solution and, above all, the electrogenerated Fe(II) is distributed in different species with distinct oxidation potentials.

Results and Discussion
2.1. Cyclic Voltammetry of FeCl 3 Figure 1 shows cyclic voltammetry (CV) of FeCl 3 in DMF + 0.1 M Et 4 NBF 4 , recorded at different scan rates. A well-defined cathodic peak (B) around −0.5 V and a shoulder (A) around −0.2 V are observed in the cathodic scan, while two anodic peaks, labelled A and B , are observed upon scan reversal. The voltametric behavior of the iron salt strongly depends on scan rate (v). The cathodic peak B and the shoulder shift to more negative potentials when v is increased while the normalized peak current of peak B, I p (B)/v 1/2 , shows only a slight decrease. Conversely, major changes with v occur in the reverse scan. As v increases, peak A increases at the expense of peak B , which instead decreases. It is clear that more than one Fe(III) species is present in solution and, above all, the electrogenerated Fe(II) is distributed in different species with distinct oxidation potentials.

respectively.
Herein, we report a deeper investigation on the electrochemical behavior of the most common iron-based catalysts used in ATRP. The study also addresses the role of added ligands, such as chloride anions, amine-bis(phenolate), phosphines and amines (structures in Scheme 2), on eATRP of MMA.  Figure 1 shows cyclic voltammetry (CV) of FeCl3 in DMF + 0.1 M Et4NBF4, recorded at different scan rates. A well-defined cathodic peak (B) around −0.5 V and a shoulder (A) around −0.2 V are observed in the cathodic scan, while two anodic peaks, labelled A' and B', are observed upon scan reversal. The voltametric behavior of the iron salt strongly depends on scan rate (v). The cathodic peak B and the shoulder shift to more negative potentials when v is increased while the normalized peak current of peak B, Ip(B)/v 1/2 , shows only a slight decrease. Conversely, major changes with v occur in the reverse scan.

Cyclic Voltammetry of FeCl3
As v increases, peak A′ increases at the expense of peak B′, which instead decreases. It is clear that more than one Fe(III) species is present in solution and, above all, the electrogenerated Fe(II) is distributed in different species with distinct oxidation potentials.  The dependence of the relative intensities of peaks A and B on scan rate points out the existence of chemical reactions between the Fe(II) species involved in the oxidation processes at these peaks. Before making any hypothesis on the identity of these Fe(II) species, we attempted to identify the principal Fe(III) species present in solution. To this end, CV of a DMF solution of FeCl 3 was investigated in the presence of added chloride ions. As shown in Figure 2, both the shoulder (A) and the anodic peak A decreased and eventually disappeared as [Cl − ] was increased. When a large excess of Cl − was added, a single peak couple (B/B ) remained. This peak couple can be assigned to the FeCl 4 − /FeCl 4 2− redox couple. The separation between the anodic and cathodic peak potentials, ∆E p = E p,B -E p,B , is 71 mV at v = 0.01 V/s and increases with scan rate (at v = 1 V/s, ∆E p = 118 mV). Therefore, the redox process under the peak couple is a quasi-reversible electron transfer. The half-wave potential calculated as the mid-point between the cathodic and anodic peaks, E 1/2 = (E p,B + E p,B )/2 does not depend on v and the average value measured at different scan rates is −0.499 ± 0.003 V vs. Fc + /Fc. cluded that dissolving FeCl3 in DMF produces mainly FeCl4 − and FeCl2 + together with small amounts of FeCl 2+ . Kim and Park [57] investigated the structure of the adduct obtained by complexation of FeCl3 with DMF. They prepared the adduct by dissolving ferric chloride in DMF, followed by slow vacuum evaporation of excess DMF. After careful analysis of the solid with various techniques, they assigned the chemical formula [Fe III Cl2(DMF)1.2(H2O)2.7] + [Fe III Cl4(DMF)2.1] − to the adduct. Although the adduct was isolated only in the solid state, it likely arises from dichloro-and tetrachloro-iron(III) species in solution, which crystallize as an adduct when all excess solvent molecules are removed by evaporation. This study also confirmed that the preferred coordination number of iron(III) is 6. Throughout this paper, we will assume that all iron halide species are hexacoordinated, but for the sake of simplicity, we will omit coordinating solvent molecules. Interestingly, CV of FeCl2 showed almost the same pattern previously observed for FeCl3, except that the process now starts with the oxidation of Fe(II) to Fe(III) species. A broad anodic peak standing for two ill-resolved peaks is observed in the initial positivegoing scan, while two well-defined peaks are observed in the reverse cathodic scan (Figure 3). When FeCl2 is dissolved in DMF, possibly the principal species present in solution is FeCl2 with minor formation of FeCl4 2− so that the main anodic peak is A′. After oxidation, the electrogenerated iron(III) species will undergo a new speciation equilibrium with When FeCl 3 is dissolved in DMF, it undergoes speciation into FeCl 4 − , FeCl 2 + and arguably some other Fe(III) species. Dass and George [56] investigated the speciation of iron(III)-chloro complexes in DMF by ultraviolet absorption spectroscopy. They concluded that dissolving FeCl 3 in DMF produces mainly FeCl 4 − and FeCl 2 + together with small amounts of FeCl 2+ . Kim and Park [57] investigated the structure of the adduct obtained by complexation of FeCl 3 with DMF. They prepared the adduct by dissolving ferric chloride in DMF, followed by slow vacuum evaporation of excess DMF. After careful analysis of the solid with various techniques, they assigned the chemical formula [Fe III Cl 2 (DMF) 1.2 (H 2 O) 2.7 ] + [Fe III Cl 4 (DMF) 2.1 ] − to the adduct. Although the adduct was isolated only in the solid state, it likely arises from dichloro-and tetrachloro-iron(III) species in solution, which crystallize as an adduct when all excess solvent molecules are removed by evaporation. This study also confirmed that the preferred coordination number of iron(III) is 6. Throughout this paper, we will assume that all iron halide species are hexacoordinated, but for the sake of simplicity, we will omit coordinating solvent molecules.
Interestingly, CV of FeCl 2 showed almost the same pattern previously observed for FeCl 3 , except that the process now starts with the oxidation of Fe(II) to Fe(III) species. A broad anodic peak standing for two ill-resolved peaks is observed in the initial positivegoing scan, while two well-defined peaks are observed in the reverse cathodic scan ( Figure 3). When FeCl 2 is dissolved in DMF, possibly the principal species present in solution is FeCl 2 with minor formation of FeCl 4 2− so that the main anodic peak is A . After oxidation, the electrogenerated iron(III) species will undergo a new speciation equilibrium with significant presence of both FeCl 4 − and FeCl 2 + . Reduction of this mixture in the reverse scan shows both peaks A and B. When increasing amounts of Cl − were added, the B/B peak couple progressively increased while A/A decreased and disappeared in the presence of excess Cl − (Figure 3). The peak couple obtained from FeCl 2 with excess Cl − had the same characteristics of that obtained from FeCl 3 in the same conditions, indicating that the same redox couple is involved in both cases. significant presence of both FeCl4 − and FeCl2 + . Reduction of this mixture in the reverse scan shows both peaks A and B. When increasing amounts of Cl − were added, the B/B′ peak couple progressively increased while A/A′ decreased and disappeared in the presence of excess Cl − (Figure 3). The peak couple obtained from FeCl2 with excess Cl − had the same characteristics of that obtained from FeCl3 in the same conditions, indicating that the same redox couple is involved in both cases. The set of analyses so far discussed allows for unambiguous assignment of peak couple B/B' to FeCl4 − /FeCl4 2− . Considering the well-documented presence of FeCl2 + in DMF solutions of FeCl3 and the voltametric pattern of FeCl2, we assign peaks A/A′ to the FeCl2 − /FeCl2 redox couple.
Let us now discuss the unusual scan-rate dependence of peaks A′ and B′ in Figure 1. The B/B′ peak couple is partially reversible at low scan rates, which can be taken to be indicative of partial disappearance of FeCl4 2− during the scan. One would then expect to see more reversibility when the scan rate is increased and, hence, the overall electrode reaction time is reduced. Surprisingly, the opposite scenario is observed: the peak couple tends toward full irreversibility at higher scan rates, while peak A′ increases ( Figure 1). To rationalize this behavior, we have to consider a fast equilibrium between two iron(II) species: If this equilibrium is fast and well shifted to the right, FeCl4 2-will be converted to FeCl2 as soon as it is generated by electroreduction of FeCl4 − at peak B. At high scan rates, the voltametric pattern will be more akin to the equilibrium situation: small equilibrium FeCl4 2− concentration, small or hardly perceptible anodic peak B′. At low scan rates, the overall oxidation process at peak B′ continues occurring for a much longer time. When the small quantity of FeCl4 2− initially present at equilibrium is consumed, reaction (3) is shifted to the left to restore the equilibrium, but FeCl4 2− continues to be oxidized, producing, at the end, a signal much higher than predicted according to the equilibrium concentration of FeCl4 2− . In other words, the anodic peak B′ arises from a sequence of two reactions, as shown in Equation (4). Indeed, this sequence is at play also at high scan rates, but its effect Let us now discuss the unusual scan-rate dependence of peaks A and B in Figure 1. The B/B peak couple is partially reversible at low scan rates, which can be taken to be indicative of partial disappearance of FeCl 4 2− during the scan. One would then expect to see more reversibility when the scan rate is increased and, hence, the overall electrode reaction time is reduced. Surprisingly, the opposite scenario is observed: the peak couple tends toward full irreversibility at higher scan rates, while peak A increases ( Figure 1). To rationalize this behavior, we have to consider a fast equilibrium between two iron(II) species: If this equilibrium is fast and well shifted to the right, FeCl 4 2− will be converted to FeCl 2 as soon as it is generated by electroreduction of FeCl 4 − at peak B. At high scan rates, the voltametric pattern will be more akin to the equilibrium situation: small equilibrium FeCl 4 2− concentration, small or hardly perceptible anodic peak B . At low scan rates, the overall oxidation process at peak B continues occurring for a much longer time. When the small quantity of FeCl 4 2− initially present at equilibrium is consumed, reaction (3) is shifted to the left to restore the equilibrium, but FeCl 4 2− continues to be oxidized, producing, at the end, a signal much higher than predicted according to the equilibrium concentration of FeCl 4 2− . In other words, the anodic peak B arises from a sequence of two reactions, as shown in Equation (4). Indeed, this sequence is at play also at high scan rates, but its effect is negligible because the scan rate is so high that the experiment ends before FeCl 4 2− is regenerated to any appreciable extent by reaction (3). Iron(II) is present in solution either as FeCl 4 2− or FeCl 2 . At low scan rates, the kinetics of equilibrium (3) affects the overall oxidation process and peak B becomes more prominent than peak A . Conversely, when at high scan rates, the effect of the backward reaction in Equation (3) becomes negligible, the oxidation process is more representative of the equilibrium conditions and A becomes the only observable anodic peak.
Starting from either FeCl 3 or FeCl 2 , it is possible to prepare DMF solutions containing exclusively FeCl 4 − or FeCl 4 2− and determine the redox properties of FeCl 4 − /FeCl 4 2− . Independent experiments on FeCl 3 and FeCl 2 performed in DMF + 0.1 M Et 4 NBF 4 in the presence of a large excess of Cl − gave E 1/2 = −0.499 ± 0.003 V vs. Fc + /Fc and diffusion coefficients, D, of (8.0 ± 0.8) × 10 −6 cm 2 s −1 and (2.3 ± 0.1) × 10 −6 cm 2 s −1 for FeCl 4 − and FeCl 4 2− , respectively ( Figures S1 and S2). The diffusion coefficients were calculated from cyclic voltammetry according to the equation of Randles-Sevcik [58]. These data can be used to determine the standard potential of the redox couple according to Equation (5): where E • is the formal potential, R is the universal gas constant, F is the Faraday constant and Ox and Red stand for the oxidized and reduced forms of the redox couple. Approximately assuming E • = E • gave E • = −0.483 ± 0.003 V vs. Fc + /Fc. The standard electron transfer rate constant, k • , was determined by the method of Nicholson [59]. Assuming a value of 0.5 for the transfer coefficient, α, ∆E p values obtained from CVs of both FeCl 4 − and FeCl 4 2− were fit to the theoretical working curve of Nicholson ( Figure S3). Both data series showed excellent fits and produced an average k • value of (1.21 ± 0.36) × 10 −2 cm s −1 .

Cyclic Voltammetry of
The amine-bis(phenolate) complex Fe III L(Cl) in solution likely exhibits a trigonal bipyramidal five-coordinate geometry, with four coordination sites occupied by L and one by the chlorine anion or by a solvent molecule [60]. Cyclic voltammetry of Fe III L(Cl) shows two cathodic peaks labeled C and D and only one anodic peak (labeled C ) in the reverse scan ( Figure 4). This pattern did not change when the scan rate was increased. When instead excess Cl − was added to the solution, peaks C and C disappeared while a new anodic peak (labeled D ) appeared. The voltametric behavior of Fe III L(Cl) can be rationalized by considering the presence of a small amount of Fe III L + together with Fe III L(Cl), which is the principal Fe(III) species present in solution. When Cl − is added in large excess over Fe(III), Fe III L + is converted to Fe III L(Cl). Therefore, peaks C and D can be attributed to the reduction of Fe III L + and Fe III L(Cl), respectively. The assignment of the anodic peaks D and C is also straightforward. They are coupled with the cathodic peaks C and D, respectively. Thus, the observed voltametric pattern stands for the following one-electron redox reactions: Fe III L + arises from partial dissociation of Fe III L(Cl) in DMF, but a small amount of Cl − is enough to convert all Fe III L + in solution into Fe III L(Cl). Indeed, peak C almost disappears as soon as the [Cl − ] / [Fe III L(Cl)] reaches 0.5. Notably, however, the appearance of peak D and its development to a full peak, as well as the disappearance of peak C , require addition of a large excess of Cl − over Fe III L(Cl). Fe II L(Cl) − generated in reaction (7) rapidly and reversibly dissociates to Fe II L and Cl − (Equation (8)). Thus, when low [Cl − ] / [Fe III L(Cl)] is used, the oxidation peak of Fe II L remains well evident, even if there is no Fe III L + in solution (peak C is absent). A large excess of Cl − is required to fully suppress the dissociation reaction and set the conditions for one-electron reduction of Fe III L(Cl) without complications due to a chemical reaction following the electron transfer.
The redox properties of Fe III L(Cl) were further investigated to determine with scan rate-dependent ∆E p was observed ( Figure S4). The half-wave potential calculated as the average of the values measured at different scan rates is −0.853 ± 0.001 V vs. Fc + /Fc. Assuming similar diffusion coefficients for Fe III L(Cl) and Fe II L(Cl) − and neglecting the activity coefficient contribution in Equation (5), the measured E 1/2 can be approximately taken as E • . The standard heterogeneous rate constant of electron transfer was determined by the method of Nicholson [59], assuming α = 0.5. The best fit of the experimental data on the working curve ( Figure S5) gave k • = (1.10 ± 0.04) × 10 −3 cm s −1 . Fe III L + arises from partial dissociation of Fe III L(Cl) in DMF, but a small amount of Cl − is enough to convert all Fe III L + in solution into Fe III L(Cl). Indeed, peak C almost disappears as soon as the [Cl − ] / [Fe III L(Cl)] reaches 0.5. Notably, however, the appearance of peak D′ and its development to a full peak, as well as the disappearance of peak C', require addition of a large excess of Cl − over Fe III L(Cl). Fe II L(Cl) − generated in reaction (7) rapidly and reversibly dissociates to Fe II L and Cl − (Equation (8)). Thus, when low [Cl − ] / [Fe III L(Cl)] is used, the oxidation peak of Fe II L remains well evident, even if there is no Fe III L + in solution (peak C is absent). A large excess of Cl − is required to fully suppress the dissociation reaction and set the conditions for one-electron reduction of Fe III L(Cl) without complications due to a chemical reaction following the electron transfer.
The redox properties of Fe III L(Cl) were further investigated to determine E° and k° in DMF in the presence of Cl − with [Cl − ] / [Fe III L(Cl)] = 30. A quasi-reversible peak couple with scan rate-dependent ΔEp was observed ( Figure S4). The half-wave potential calculated as the average of the values measured at different scan rates is −0.853 ± 0.001 V vs. Fc + /Fc. Assuming similar diffusion coefficients for Fe III L(Cl) and Fe II L(Cl) − and neglecting the activity coefficient contribution in Equation (5), the measured E1/2 can be approximately taken as E°. The standard heterogeneous rate constant of electron transfer was determined by the method of Nicholson [59], assuming α = 0.5. The best fit of the experimental data on the working curve ( Figure S5) gave k° = (1.10 ± 0.04)  10 −3 cm s −1

Cyclic Voltammetry of FeCl3 in the Presence of Other Ligands
Phosphorous-and nitrogen-based ligands are often used as additional ligands in FeX3-catalyzed ATRP. FeX3 is often mixed with a ligand L with the assumption that a more active catalyst species FeX3/L is formed. We investigated the role of triphenylphosphine (TPP) and tris(2-pyridylmethyl)amine (TPMA), taken as examples of phosphorous and nitrogen ligands, respectively, on the redox behavior of FeCl3. Figure 5a shows CVs recorded for FeCl3 before and after addition of TPP. No change in the voltametric pattern of the iron chloride complex(es) was observed. A similar result was obtained by UV-vis analysis. The absorption spectrum recorded after addition of TPP was simply the superimposition of the separate spectra of FeCl3 and TPP ( Figure S6). The

Cyclic Voltammetry of FeCl 3 in the Presence of Other Ligands
Phosphorous-and nitrogen-based ligands are often used as additional ligands in FeX 3 -catalyzed ATRP. FeX 3 is often mixed with a ligand L with the assumption that a more active catalyst species FeX 3 /L is formed. We investigated the role of triphenylphosphine (TPP) and tris(2-pyridylmethyl)amine (TPMA), taken as examples of phosphorous and nitrogen ligands, respectively, on the redox behavior of FeCl 3 . Figure 5a shows CVs recorded for FeCl 3 before and after addition of TPP. No change in the voltametric pattern of the iron chloride complex(es) was observed. A similar result was obtained by UV-vis analysis. The absorption spectrum recorded after addition of TPP was simply the superimposition of the separate spectra of FeCl 3 and TPP ( Figure S6). The voltametric pattern of FeCl 2 was also unaffected by TPP ( Figure S7). These results clearly show that TPP does not form new complexes with FeCl 3 or FeCl 2 in the investigated reaction conditions. Walker and Poli [61] reported the synthesis of FeCl 3 -phosphine complexes of general formula FeCl 3 (PR 3 ) 2 (R = Me, Ph, cyclohexyl). The compounds were prepared at low temperature in nonpolar solvents, such as benzene and toluene, but their solutions were unstable at room temperature or higher. Phosphine complexes of FeX 2 were reported by Sawamoto and co-workers as FeX 2 (L 2 ) (X = Cl, L = PMePh 2 ; X = Br, L = PMePh 2 , PPh 3 , P(n-Bu) 3 ) [34]. They were prepared in toluene and used as ATRP catalysts in the same solvent at 80 • C. Although these complexes seem to be stable in these reaction conditions, their stability and chemical structure in polar solvents, such as DMF, are yet to be proved.
To check whether TPP can reduce Fe(III) to Fe(II), a mixture of FeCl 3 with a two-fold excess of TPP in DMF + 0.1 M Et 4 NBF 4 was monitored by linear sweep voltammetry (LSV) at a rotating disc electrode at room temperature. LSV recorded at t = 0 was identical with that observed before TPP addition and showed a well-defined wave for the reduction of Fe(III) to Fe(II) with a cathodic limiting current, |I L,c | = 56 µA. Since the system initially contained only Fe(III), the anodic limiting current, I L,a , was zero. The LSV recorded after 1 h showed a slight decrease in |I L,c | together with small I L,a . The reaction was left overnight and after 20 h, a significantly decreased Fe(III) concentration accompanied by the build-up of Fe(II) was observed, clearly indicating that FeCl 3 was slowly reduced by TPP. The ability of triphenylphosphines to act as reducing agents in iron-catalyzed ATRP has previously been suggested and, in a few cases, experimentally demonstrated [62,63]. The reaction between FeCl 3 and TPP is very slow at room temperature and probably will not affect the rate of polymerization in such conditions. However, many Fe-catalyzed polymerizations, especially in the case of mMA and styrene, are typically carried out at temperatures as high as 110 • C, where Fe(III) reduction by phosphines might be fast. of general formula FeCl3(PR3)2 (R = Me, Ph, cyclohexyl). The compounds were prepared at low temperature in nonpolar solvents, such as benzene and toluene, but their solutions were unstable at room temperature or higher. Phosphine complexes of FeX2 were reported by Sawamoto and co-workers as FeX2(L2) (X = Cl, L = PMePh2; X = Br, L = PMePh2, PPh3, P(n-Bu)3) [34]. They were prepared in toluene and used as ATRP catalysts in the same solvent at 80 °C. Although these complexes seem to be stable in these reaction conditions, their stability and chemical structure in polar solvents, such as DMF, are yet to be proved. To check whether TPP can reduce Fe(III) to Fe(II), a mixture of FeCl3 with a two-fold excess of TPP in DMF + 0.1 M Et4NBF4 was monitored by linear sweep voltammetry (LSV) at a rotating disc electrode at room temperature. LSV recorded at t = 0 was identical with that observed before TPP addition and showed a well-defined wave for the reduction of Fe(III) to Fe(II) with a cathodic limiting current, |IL,c| = 56 μA. Since the system initially contained only Fe(III), the anodic limiting current, IL,a, was zero. The LSV recorded after 1 h showed a slight decrease in |IL,c| together with small IL,a. The reaction was left overnight and after 20 h, a significantly decreased Fe(III) concentration accompanied by the buildup of Fe(II) was observed, clearly indicating that FeCl3 was slowly reduced by TPP. The ability of triphenylphosphines to act as reducing agents in iron-catalyzed ATRP has previously been suggested and, in a few cases, experimentally demonstrated [62,63]. The reaction between FeCl3 and TPP is very slow at room temperature and probably will not affect the rate of polymerization in such conditions. However, many Fe-catalyzed polymerizations, especially in the case of MMA and styrene, are typically carried out at temperatures as high as 110 °C, where Fe(III) reduction by phosphines might be fast.
In contrast to TPP, the nitrogen ligand TPMA reacted rapidly and quantitatively with FeCl3. Addition of 1 equiv. of TPMA to a solution of FeCl3 caused full disappearance of the original signal, which was replaced by a reversible peak couple at a less negative potential ( Figure 6). Further addition of TPMA did not affect the voltametric response. It is obvious that a new complex, in which both Fe(III) and Fe(II) are stable, is formed. TPMA complexes of both Fe(III) and Fe(II) are known in the literature [64,65]. They have a distorted octahedral geometry arising from tetra-coordination by TPMA plus coordination of chloride ions, FeCl2(TPMA) + and FeCl2(TPMA). Therefore, reactions (9) and (10) occur when TPMA is added to a DMF solution of FeCl3. The standard reduction potential of FeCl2(TPMA) + (Equation (11)) can be estimated as E° ≈ E1/2 = (Epa + Epc)/2, where Epa and Epc are the anodic and cathodic peak potentials and the measured value is E1/2 = −0.211 V vs. Fc + /Fc. In contrast to TPP, the nitrogen ligand TPMA reacted rapidly and quantitatively with FeCl 3 . Addition of 1 equiv. of TPMA to a solution of FeCl 3 caused full disappearance of the original signal, which was replaced by a reversible peak couple at a less negative potential ( Figure 6). Further addition of TPMA did not affect the voltametric response. It is obvious that a new complex, in which both Fe(III) and Fe(II) are stable, is formed. TPMA complexes of both Fe(III) and Fe(II) are known in the literature [64,65]. They have a distorted octahedral geometry arising from tetra-coordination by TPMA plus coordination of chloride ions, FeCl 2 (TPMA) + and FeCl 2 (TPMA). Therefore, reactions (9) and (10) occur when TPMA is added to a DMF solution of FeCl 3 . The standard reduction potential of FeCl 2 (TPMA) + (Equation (11)) can be estimated as E • ≈ E 1/2 = (E pa + E pc )/2, where E pa and E pc are the anodic and cathodic peak potentials and the measured value is E 1/2 = −0.211 V vs. Fc + /Fc.
We examined the effect of Cl − excess on the stability of FeCl 2 (TPMA) + . Addition of a two-fold excess of Cl − over FeCl 3 did not affect the CV pattern of the iron-TPMA complex ( Figure S8). A decrease in the cathodic peak of FeCl 2 (TPMA) + together with the appearance of the cathodic peak of FeCl 4 − was observed only when a large excess of Cl − was added. This indicates that Fe(III) has a much higher affinity for TPMA than Cl − , but reaction (9) shifts to the left as the concentration of added Cl − is increased. Notably, when in the presence of excess Cl − , the principal iron species present in solution became FeCl 4 − , a reversible peak couple for the FeCl 4 − / FeCl 4 2− couple was not observed ( Figure S8). This means that the stability constant of FeCl 2 (TPMA) is higher than that of FeCl 2 (TPMA) + . Once FeCl 4 − is reduced to FeCl 4 2− , the latter rapidly reacts with TPMA to form FeCl 2 (TPMA), so that the oxidation peak of FeCl 4 2− is missing while the anodic peak due to FeCl 2 (TPMA) oxidation is always present ( Figure S8). We examined the effect of Cl − excess on the stability of FeCl2(TPMA) + . Addition of a two-fold excess of Cl − over FeCl3 did not affect the CV pattern of the iron-TPMA complex ( Figure S8). A decrease in the cathodic peak of FeCl2(TPMA) + together with the appearance of the cathodic peak of FeCl4 − was observed only when a large excess of Cl − was added. This indicates that Fe(III) has a much higher affinity for TPMA than Cl − , but reaction (9) shifts to the left as the concentration of added Cl − is increased. Notably, when in the presence of excess Cl − , the principal iron species present in solution became FeCl4 − , a reversible peak couple for the FeCl4 − / FeCl4 2− couple was not observed ( Figure S8). This means that the stability constant of FeCl2(TPMA) is higher than that of FeCl2(TPMA) + . Once FeCl4 − is reduced to FeCl4 2− , the latter rapidly reacts with TPMA to form FeCl2(TPMA), so that the oxidation peak of FeCl4 2− is missing while the anodic peak due to FeCl2(TPMA) oxidation is always present ( Figure S8).
In summary, the Fe(III) species formed in solution and their standard reduction potentials are summarized in Table 1. When no ligand is added or TPP is used as a ligand, the principal Fe(III) complex in solution is FeCl4 − with a characteristic E° of −0.483 V vs. Fc + /Fc. 2-pyridylamino-N,N-bis(2-methylene-4,6-dichlorophenolate) (L) forms a stable complex, Fe III L(Cl), with a cathodic shift of E° of Fe(III)/Fe(II) by 0.37 V. TPMA also forms a stable complex with the iron salt, but the standard potential of the Fe(III)/Fe(II) couple shifts anodically by 0.272 V. In ATRP, a metal at a low oxidation state (Fe(II) in this case) reacts with an alkyl halide initiator, RX or a halogen-capped dormant species, Pn-X, to generate the active radical species. The rate of this activation reaction is strongly dependent on E° of the metal complex. Therefore, the expected order of activity of the Fe complexes examined here is Fe III L(Cl) > Fe III Cl4 − > Fe III Cl2(TPMA) + . However, in addition to the ability to activate RX, an efficient ATRP catalyst must also be a good radical deactivator; hence, predicting the overall efficacy of an iron catalyst, on the basis of E° alone, is not straightforward. In summary, the Fe(III) species formed in solution and their standard reduction potentials are summarized in Table 1. When no ligand is added or TPP is used as a ligand, the principal Fe(III) complex in solution is FeCl 4 − with a characteristic E • of −0.483 V vs. Fc + /Fc. 2-pyridylamino-N,N-bis(2-methylene-4,6-dichlorophenolate) (L) forms a stable complex, Fe III L(Cl), with a cathodic shift of E • of Fe(III)/Fe(II) by 0.37 V. TPMA also forms a stable complex with the iron salt, but the standard potential of the Fe(III)/Fe(II) couple shifts anodically by 0.272 V. In ATRP, a metal at a low oxidation state (Fe(II) in this case) reacts with an alkyl halide initiator, RX or a halogen-capped dormant species, P n -X, to generate the active radical species. The rate of this activation reaction is strongly dependent on E • of the metal complex. Therefore, the expected order of activity of the Fe complexes examined here is Fe III L(Cl) > Fe III Cl 4 − > Fe III Cl 2 (TPMA) + . However, in addition to the ability to activate RX, an efficient ATRP catalyst must also be a good radical deactivator; hence, predicting the overall efficacy of an iron catalyst, on the basis of E • alone, is not straightforward.

eATRP Mediated by Amine-bis(phenolate) iron(III) Chloride, Fe III L(Cl)
The voltametric behavior of Fe III L(Cl) in 50 vol% mMA in DMF + 0.1 M Et 4 NBF 4 is slightly different from that in pure DMF. The pre-peak C is absent, whereas peak D shows partial reversibility ( Figure S9). It appears that dissociation of Fe III L(Cl) to give Fe III L + and Cl − is disfavored in the monomer/solvent mixture. Further, Fe II L(Cl) − is more stable in mMA/DMF than in pure DMF. These changes are likely due to modifications of medium polarity, which decreases when mMA with a dielectric constant, ε, of 6.53 at 25 • C [66] is added to DMF (ε = 38.25 at 20 • C [67]). The mixture has lower ability to solvate ions than pure DMF and, therefore, dissociation of Fe III L(Cl) and Fe II L(Cl) − becomes less favored in the mixture. eATRP of mMA mediated by Fe III L(Cl) was carried out in 50/50 (V/V) DMF/MMA mixtures with 5 mM catalyst and 15 mM initiator. As initiators, ethyl α-chlorophenyl acetate, ECPA, and ethyl α-bromophenyl acetate, EBPA, which are among the most active alkyl halides used in ATRP, were chosen. Nonetheless, CV tests revealed that electrogenerated Fe II L(Cl) − reacts very slowly with these initiators ( Figure S9). Since it was not easy to estimate E 1/2 of the catalyst in the typical reaction conditions, the applied potential, E app , was chosen with reference to the cathodic peak potential of the mediator (peak D). The reaction time was set at 4 h, but in some cases, polymerization was stopped earlier (2-3 h) because of drastic increase in viscosity.
The results of the electrochemically mediated polymerizations are collected in Table 2, whereas examples of reaction kinetics and trends of M n and dispersity are shown in Figure 7. Using ECPA at 70 • C and E app = E p , 42.3% conversion was achieved after 4 h, but the reaction was not controlled. The dispersity of the polymer was very high and its molecular weight exceeded the theoretical value by about one order of magnitude ( Table 2, entry 1). This first experiment indicated that activation was effective, but deactivation was inefficient. The experiment was then repeated with a 10-fold excess of Cl − with respect to the iron mediator to increase the concentration of the deactivator Fe III L(Cl). Again, the reaction was fast without any control (entry 2). Last, the temperature was lowered to 50 • C. The reaction was slower, but the dispersity remained very high and M n was 20-fold higher than the theoretical value (entry 3). Two more experiments were performed with this mediator by changing the initiator to EBPA. Activation of this initiator by ATRP mechanism will produce Fe III L(Br), which might be a better deactivator than Fe III L(Cl). Disappointingly, the results were very similar to those obtained with ECPA ( Table 2, entries 3-4 and Figure 7).

Entry
Initiator  catalyst concentration, which inevitably will involve costly polymer purification st eATRP with this type of mediator was here investigated to explore the possibility of ploying a low-catalyst load under mild conditions. The failure of control in low-cata load eATRP can be attributed to lack of deactivation. Schroeder and Buback [69] measu deactivation rate constants, kdeact, in iron-mediated ATRP. For an amine-bis(pheno iron(III) complex, they measured at 60 °C a kdeact value of about 10 4 M −1 s −1 , which is 2 or of magnitude lower than kdeact of copper complexes. This value increases by one ord magnitude if the temperature is raised to 120 °C. While deactivation rate in bulk mono at 120 °C with high catalyst load is high enough to ensure control over molecular we distribution, it is inefficient in eATRP. Indeed, due to lower T and lower catalyst load it can be estimated that deactivation rate in the eATRP experiments is at least 2 orde magnitude lower than in the reaction conditions employed by Shaver and co-worker

eATRP Mediated by FeCl3
eATRP of MMA mediated by iron halide complexes, both in the absence and pres of a phosphorous ligand, has already been reported [44][45][46]. These studies used FeBr3 FeCl3.6H2O with tris(2,4,6-methoxyphenyl)phosphine as an additional ligand FeCl3.6H2O without added ligand and, in most cases, the best results were obtained w quite high concentrations of iron salt were used. For comparison with FeBr3 FeCl3.6H2O, which was the principal catalyst in their study, Wang et al. [46] report single eATRP experiment with FeCl3 [46]. Using ca 31 mM FeCl3 in N-methylpyrrolid at 95 °C, they achieved 43% monomer conversion after ca 6.5 h, obtaining a polymer dispersity Ð = 1.34. Overall, the amine-bis(phenolate) iron(III) complex proved to be inefficient as a mediator of controlled radical polymerization of mMA in DMF under electrochemical generation of Fe(II) activator species. This result is at variance with previous reports by Shaver and co-workers on well-controlled polymerizations of mMA and styrene mediated by various amine-bis(phenolate) iron(III) complexes in the presence of radical initiators, such as azobis(isobutyronitrile) [37,68]. There are, however, important differences between the experimental conditions. All previous polymerizations were performed in bulk or in mixtures of monomer with nonpolar solvents. Further, the typical concentration of the iron(III) complex was nearly 100 mM. Last, the reactions were carried out at temperatures as high as 120 • C. Obviously these conditions are not desirable, especially the use of high catalyst concentration, which inevitably will involve costly polymer purification steps. eATRP with this type of mediator was here investigated to explore the possibility of employing a low-catalyst load under mild conditions. The failure of control in low-catalystload eATRP can be attributed to lack of deactivation. Schroeder and Buback [69] measured deactivation rate constants, k deact , in iron-mediated ATRP. For an amine-bis(phenolate) iron(III) complex, they measured at 60 • C a k deact value of about 10 4 M −1 s −1 , which is 2 orders of magnitude lower than k deact of copper complexes. This value increases by one order of magnitude if the temperature is raised to 120 • C. While deactivation rate in bulk monomer at 120 • C with high catalyst load is high enough to ensure control over molecular weight distribution, it is inefficient in eATRP. Indeed, due to lower T and lower catalyst loading, it can be estimated that deactivation rate in the eATRP experiments is at least 2 orders of magnitude lower than in the reaction conditions employed by Shaver and co-workers.

eATRP Mediated by FeCl 3
eATRP of mMA mediated by iron halide complexes, both in the absence and presence of a phosphorous ligand, has already been reported [44][45][46]. These studies used FeBr 3 and FeCl 3 .6H 2 O with tris(2,4,6-methoxyphenyl)phosphine as an additional ligand or FeCl 3 .6H 2 O without added ligand and, in most cases, the best results were obtained when quite high concentrations of iron salt were used. For comparison with FeBr 3 and FeCl 3 .6H 2 O, which was the principal catalyst in their study, Wang et al. [46] reported a single eATRP experiment with FeCl 3 [46]. Using ca 31 mM FeCl 3 in N-methylpyrrolidone at 95 • C, they achieved 43% monomer conversion after ca 6.5 h, obtaining a polymer with dispersity Ð = 1.34.
Based on these previous studies, a relatively high concentration of FeCl 3 was first considered. Before eATRP experiments, the effect of [FeCl 3 ] and monomer on the voltametric behavior of Fe(III) was examined in conditions of eATRP. Addition of mMA to DMF did not change the general CV pattern of FeCl 3 , but peak A became more pronounced while peak B gained more reversibility. The equilibrium distributions of both Fe(III) and Fe(II) species are influenced by medium polarity as well as initial FeCl 3 concentration and temperature. The effect of [FeCl 3 ] is evidenced in Figure 8, which shows a remarkable enhancement of peak B as [FeCl 3 ] is increased. In DMF/MMA, it is possible to estimate E 1/2 of the Based on these previous studies, a relatively high concentration of FeCl3 was first considered. Before eATRP experiments, the effect of [FeCl3] and monomer on the voltametric behavior of Fe(III) was examined in conditions of eATRP. Addition of MMA to DMF did not change the general CV pattern of FeCl3, but peak A′ became more pronounced while peak B gained more reversibility. The equilibrium distributions of both Fe(III) and Fe(II) species are influenced by medium polarity as well as initial FeCl3 concentration and temperature. The effect of [FeCl3] is evidenced in Figure 8, which shows a remarkable enhancement of peak B′ as [FeCl3] is increased. In DMF/MMA, it is possible to estimate E1/2 of the FeCl4 − / FeCl4 2− redox couple as (Ep,B + Ep,B′)/2, especially at high [FeCl3] and low scan rates. The estimated value was −0.49 V vs. Fc + /Fc. The results of eATRPs for MMA performed at different conditions are summarized in Table 3. The applied potential, Eapp, is reported as Eapp − E1/2, with reference to the value of E1/2 measured from CV before electrolysis. A first set of eATRPs at 70 °C was carried out The results of eATRPs for mMA performed at different conditions are summarized in Table 3. The applied potential, E app , is reported as E app − E 1/2 , with reference to the value of E 1/2 measured from CV before electrolysis. A first set of eATRPs at 70 • C was carried out with 23.5 mM FeCl 3 at different E app values. All polymerizations were well controlled, obeying first-order kinetic rate laws and producing polymers with narrow molecular weight distribution, though the experimental molecular weights did not perfectly match the theoretical values ( Figure 9). M n,GPC was always greater than M n,th , indicating initiation efficiency < 1. The rate of polymerization was higher at lower E app values, but polymer dispersity was better at higher (less negative) applied potentials (Table 3, entries 1-3). Lowering the temperature to 55 • C at E app = E 1/2 − 0.06 V had no significant effect on reaction rate and polymer properties ( Table 3, entry 4). Therefore, most of the other experiments were conducted at 55 • C.  6 Determined by GPC. 7 Theoretical molecular weight. 8 In the presence of a 2-fold excess of TPP with respect to FeCl 3 . 9 In the presence of a 2-fold excess of TPMA with respect to FeCl 3 .
Molecules 2022, 27, x FOR PEER REVIEW 13 o reaction rate. As desired, the conversion increased more than twice while the disper remained unchanged (Table 3, entry 9).  Two more experiments were performed to test the role of added ligands. The eAT of entry 7 (Table 3) was repeated with a two-fold excess of TPP over FeCl3 (Table 3, en 10). Conversion after 5 h increased from 38.5% to 40.6%, while the dispersity remai unchanged. We previously showed that TPP acts as a weak reducing agent rather tha ligand able to form a new complex with FeCl3 in DMF. The slight rate enhancement m be attributed to the contribution Fe(II) regeneration via reduction by TPP in the homo neous phase. The same experiment was again repeated in the presence of TPMA in p of TPP (Table 3, entry 11). No polymerization was observed after 5 h, clearly indicat that FeCl2(TPMA) is not able to activate the initiator.
To sum up, ligand-free FeCl3 is a cheap ATRP catalyst able to achieve good polym ization control under appropriate conditions. The efficacy of the catalyst depends   (Table 3, entries 5 and 7).
In an attempt to improve polymerization control, the eATRP experiment of entry 7 in Table 3 was repeated at E app = E 1/2 . The dispersity was little affected (slight increase from 1.50 to 1.53) but the reaction rate decreased considerably, with conversion after 5 h dropping from 38.5% to 23.7%. The temperature was then raised to 70 • C to increase the reaction rate. As desired, the conversion increased more than twice while the dispersity remained unchanged (Table 3, entry 9).
Two more experiments were performed to test the role of added ligands. The eATRP of entry 7 (Table 3) was repeated with a two-fold excess of TPP over FeCl 3 (Table 3, entry 10). Conversion after 5 h increased from 38.5% to 40.6%, while the dispersity remained unchanged. We previously showed that TPP acts as a weak reducing agent rather than a ligand able to form a new complex with FeCl 3 in DMF. The slight rate enhancement may be attributed to the contribution Fe(II) regeneration via reduction by TPP in the homogeneous phase. The same experiment was again repeated in the presence of TPMA in place of TPP (Table 3, entry 11). No polymerization was observed after 5 h, clearly indicating that FeCl 2 (TPMA) is not able to activate the initiator.
To sum up, ligand-free FeCl 3 is a cheap ATRP catalyst able to achieve good polymerization control under appropriate conditions. The efficacy of the catalyst depends on many parameters, including concentrations of catalyst and initiator, type of solvent, temperature and applied potential. Further work systematically addressing the effects of these parameters and their combinations is necessary to define optimal controlled-polymerization conditions.

Instrumentation
Electrochemical measurements were carried out in a 5-neck glass cell under an Ar atmosphere; an Autolab PGSTAT 30 or 30N potentiostat/galvanostat (EcoChemie, Utrecht, The Netherlands) run by a PC with GPES or NOVA software (EcoChemie) was used. The working electrode, counter electrode and reference electrode used during voltametric investigations were a 3 mM diameter GC disk (Tokai GC-20), a Pt ring and Ag/AgI/0.1 M n-Bu 4 NI in DMF, respectively. Before each experiment, the GC disk was cleaned by polishing with a 0.25 µm diamond paste, followed by ultrasonic rinsing in ethanol for 5 min. The reference electrode was always calibrated with ferrocene (Fc), which was added at the end of each experiment as an internal standard and all potentials are reported versus the fer-rocenium/ferrocene (Fc + /Fc) redox couple. Conversion of these potentials to the saturated calomel electrode scale can be achieved by using E o (Fc + /Fc) = 0.476 V vs. SCE [70].
eATRP experiments were carried out in a two-compartment cell equipped with a Pt mesh (Alfa Aesar, 99.9% metals basis) working electrode, a graphite rod counter electrode and the same reference electrode used in cyclic voltammetry. Before each experiment, the Pt mesh was electrochemically activated in 0.5 M H 2 SO 4 by cycling the potential from −0.7 V to 1 V vs. Hg/Hg 2 SO 4 at a scan rate of 0.2 V s −1 (60 cycles). The counter electrode was separated from the working solution by a glass frit filled with the same electrolyte solution used in the working electrode compartment and a methylcellulose gel saturated with Et 4 NBF 4 .
Gel permeation chromatography (GPC) was used to determine the number average molecular weight (M n ) and dispersity (Ð) of polymers prepared by eATRP. The GPC instrument was Agilent 1260 Infinity, equipped with a refractive index (RI) detector and two PLgel Mixed-D columns (300 mM, 5 µm) connected in series. The column compartment and RI detector were thermostated at 70 • C and 50 • C, respectively. The eluent was DMF containing 10 mM LiBr, at a flow rate of 1 mL/min. Before injection, the samples were filtered through alumina over a PTFE membrane of 200 nm pore to remove any particulate material and the iron catalyst. The column system was calibrated with 12 linear poly(methyl methacrylate) standards (M n = 540-2,210,000 Da). Monomer conversion was determined by 1 H-NMR spectroscopy with a 200 MHz Bruker Avance instrument, using CDCl 3 as a solvent.
UV-Vis spectra were recorded with an Agilent Cary 5000 spectrophotometer by using 10 mM optical path length quartz cuvettes.

eATRP of Methyl Methacrylate
A thermostated 5-neck electrochemical cell, flushed with an inert gas, was loaded with DMF/MMA (50:50, V/V) + 0.1 M Et 4 NBF 4 and the desired amount of iron catalyst. After recording a CV of the catalyst, the initiator RX was injected, and a CV was recorded. Polymerization was then started by applying the selected applied potential (E app ) and samples were withdrawn periodically to measure monomer conversion and M n and Ð of the polymer.

Conclusions
In summary, we have shown that FeCl 3 dissolves in DMF to yield FeCl 4 − and FeCl 2 + . Reduction of these species gives prevalently FeCl 2 . A large excess of Cl − is required to convert FeCl 3 and FeCl 2 to FeCl 4 − and FeCl 4 2− , respectively. TPMA forms stable complexes with both Fe(III) and Fe(II). 2-Pyridylamino-N,N-bis(2-methylene-4,6-dichlorophenolate) (L) also gives a stable iron(III) complex, Fe III L(Cl), which provides Fe II L(Cl) − , upon oneelectron reduction. In contrast, triphenylphosphine does not form a stable complex with FeCl 3 . It acts as a reducing agent. eATRP of mMA catalyzed by Fe III L(Cl) was fast and uncontrolled. FeCl 3 proved to be a much better catalyst than the amine-bis(phenolate) complex. FeCl 3 -mediated polymerization was well controlled, provided that a large amount of the iron salt was employed. No polymerization was observed when TPMA was used as an additional ligand. It appears that the best catalyst system is the ligand-free iron salt. A weakness of this system is in the deactivation step, characterized by a low rate constant.  Table S1: 1 H NMR spectral data of 2-pyridylamino-N,N-bis(2-methylene-4,6 -dichlorophenol; Table S2: Elemental analysis of chloro(2-pyridylamino-N,N-bis(2-methylene-4,6 -dichlorophenolate))iron(III). Ref. [60] is cited in Supplementary Materials.