Electrontransfer Processes

In electrochemical cells, electron transfer occurs within the electrode/solution interface, with electron removal (oxidation) at the anode and electron introduction (reduction) at the cathode. The current through the solution is carried by the ions of the electrolyte, and the voltage limits are those for electron removal from and electron insertion into the solvent/electrolyte {e.g., H2O/(H+Oaq)(ClO4aq); [Na1 (OH2)+](Cl-q)}

In the gas phase, electron removal from atoms is limited by their ionization potential (e.g., H-, 13.6 eV; K-, 4.3 eV; Na-, 5.1 eV; Cu-, 7.7 and 20.3 eV; Ag, 7.6 eV; Fe, 7.9,16.2, and 30.7 eV). However, in the solution phase, electron removal (oxidation) from the solvent may be facilitated by the presence of substrate atoms (rather than be from them).

For example, with water at pH 0 the process of Eq. (73) is shifted -4.82 V when hydrogen atoms are present,

A. Oxidative Electrochemistry: Metals, Metal Complexes, Lewis Bases

1. Metals

The transformation of metal electrode surfaces via ele-ctrooxidation to their metallo-oxides, solvated ions, and metal complexes is fundamental to most anodic electrochemical processes (batteries, electrorefining, anodic stripping analysis, and reference electrodes). Although this is traditionally represented as the removal of one (or more) valence electron from a metal atom at the electrode surface to give a metal ion [e.g., Ag (s) - e- ^ Ag+; E°, +0.80 V vs NHE], the gas-phase ionization potential [e.g., Ag ■ (g) - e- ^ Ag+ (g); IP, 7.6 eV] is far greater than the observed oxidation potential. The difference is attributed to the solvation energy for the metal ion [e.g., Ag+ + «H2O^ Ag+(aq); -AG(aq)«70-100 kcal mol-1]. However, such a sequential path would not obviate the 7.6-V energy barrier for the initial step and is in conflict with the observed thermodynamic reversibility for many metal/solvated-metal-ion redox couples.

All reactions, and particularly redox processes, occur via the easiest and lowest energy pathway that is available (mechanistically feasible) to the system. In the case of a metal electrode/electrolyte interface undergoing anodic transformations, the electrons can come from (a) surface metal atoms (energy limit; first ionization potential), (b) solvent molecules (energy limit; oxidation potential of solvent), (c) electrolyte anions (energy limit; oxidation potential of anions), and (d) base ligands (energy limit; oxidation potential of ligand). All metal electrodes are elec-trochemically transformed via path b, c, or d and never via path a. This general conclusion is illustrated for silver and copper electrodes in aqueous and acetonitrile (MeCN) solutions that contain inert electrolyte, chloride ion (Cl-), or bipyridine (bpy).

In aqueous solutions at pH 0, the silver electrode facilitates oxidation of water

The gas-phase ionization potential for a silver atom is 7.6 eV. In contrast, water is oxidized (gives up an electron)

at much lower potentials [Eq. (73); at pH 0, +2.72 V vs NHE]. At a silver electrode, water oxidation is facilitated via formation of an AgI-OH+ bond; the shift in oxidation potential from +2.72 to +0.80 V is a measure of the bondformation energy (-AGBF),

- AGbf = (+2.72 - 0.80) 23.1 = 44.4 kcal mol-1. (78) At pH 14 the anodic process is the oxidation of HO-, e-

which is facilitated via formation of an AgI-OH bond [-AGbf = (1.89 - 0.34) 23.1 = 35.8 kcal mol-1]

The data of Eqs. (77) and (79) can be combined to give a value for the solubility product (Ksp) for Ag-OH (s),

In the presence of chloride ion, metal electrodes facilitate its oxidation [Eq. (75); E°, +2.41 V vs NHE] via formation of metal-chlorine covalent bonds, e.g.,

Hence, the differential bond-formation energy [A(-AGbf)] [Ag-Cl bond energy, minus the energy required to break an Ag-Ag bond at the Ag (s) surface)] is given by the difference in oxidation potentials [Eqs. (75) and (82)],

Because the escape energy for an Ag- atom from Ag (s) is 68 kcal mol-1, a reasonable approximation for the breakage of a single bond is 22.6 kcal mol-1 [(1 /3) 68]. When combined with Eq. (83), this gives a reasonable value for -AGbf,

The literature value for the dissociative bond energy (AHDBE) ofAg-Cl (g) is 81.6 kcal mol-1, which is equivalent to an estimated -AGBF value of 73.8 kcal mol-1

[-AGBF = aHdbe - TaS = 81.6 - 7.8 (est) = 73.8kcal mol-1]. Thus, the proposition that metal electrode oxidations are solvent or ligand centered with potentials that reflect the metal-solvent/ligand bond-formation free energies (-AGBF) is supported by independent bond-energy data. The data of Eqs. (77) and (82) provide a measure of the solubility product for Agcl (s).

Another important example is the oxidation of Cl- at a mercury electrode [Hg2 (l)] to form calomel [mercurous chloride, Hg2Cl2 (s), Cl-HgII-HgII-Cl (s)].

The potential shift for the Cl-/Cl- couple from +2.41 [Eq. (75)] to +0.27 V in the presence of Hg2 (l) is a measure of the [Cl-HgHg] bond energy [-AGbf = (2.41 -0.27) 23.1 = 49.4 kcal mol-1].

Similar metal-facilitated oxidations of H2O and of Cl-occur for all metal electrodes. The respective potentials for the oxidation of each at a copper electrode are

Cu(s) + 2H2O ^ (H2O)CuI-OH+ (E°)pe0, +0.52V [-AGbf, 46kcal mol-1] (87)

Additional redox data for oxidations of H2O/HO- at Cu, Ag, and Au electrodes in aqueous and acetonitrile (MeCN) solutions are summarized in Table I. At pH 0 with an iron electrode the water oxidation of Eq. (73) is shifted by -3.12 V,

which indicates that the [H2O(H2O)+] species is stabilized by a strong [(H2O)4(H2O)+]FeII-OH2+ covalent bond (-AGbf, ~71 kcal mol-1).

2. Metal Complexes

In an analogous fashion, the removal of an electron (ox idation) from water via Eq. (73) is aided by the presence of transition-metal ions [e.g., CuI(OH2)+, FeII(OH2)2+, and CeIII(OH2)6+, each with one, two, and three M-OH+ covalent bonds, respectively],

(H2O)52+ FeIII


In none of these examples has the potential for removal of an electron approached the ionization potentials of the metals. Although traditional treatments attribute the potentials of Eqs. (77), (87), and (89)-(92) to the removal of electrons from the metals, coupled with large ionic solva-tion energies, this requires a pathway with the ionization potential as a kinetic barrier. Furthermore, the spontaneous reaction of iron with acidified water is driven by the formation of Fe-OH+ and H-H covalent bonds that facilitate hydrogen-atom transfer from water (rather than electron transfer from iron),

{Note: To ionize a gas-phase iron atom (Fe —Fe3+) requires 54.8 eV (1266 kcal mol-1); in turn, this species reacts upon dissolution into liquid water {Fe3+(g) + 7H2O(l) ^ [(H2O)5+]FeIII-OH + H+O, -A H « 1000 kcal mol-1 (1266 - 266)}; the net energy change often is ascribed as the solvation energy for Fe3+ (g) (heat of hydration)}.

Within an aprotic solvent (e.g., acetonitrile, MeCN) oxidation of metals and metal complexes also is ligand centered with the potential determined by the oxidation potential of the ligand and the metal-ligand covalent bondformation free energy (-AG BF). For example, the freebpy ligand in acetonitrile is oxidized near the solvent limit at a glassy-carbon electrode (GC) (Fig. 5),



but at a copper electrode the oxidation occurs at a negative potential,

TABLE II Oxidation Potentials for Ligands (l) and Their ml3 Complexes with Zn(II), Mn(II), Fe(II), and Co(II) in MeCN (0.1 M Tetraethylammonium Perchlorate)

Ligand (L or L-)a

Eb1/2, V vs NHEC

L/L+ • ( L-/ L•)

Zn11 L3

MnI! L3

Fen L3

Con L3





+ 1.84





+ 1.55

+ 1.30


PA -

+ 1.50

+ 1.54
















a Key: bpy, 2,2r-bipyridine; PA-, picolinate (2-carboxylate pyridine); acac-, acetylacetonate; 8-Q-, 8-quinolinate.

b E\/2 taken as (Ep a + Ep c)/2 for reversible couples of Mn11 L3 and Fe11 L3 complexes; as Ep a/2 + 0.03 V for L (or L—) and Zn11 L3; and as Ep c/2 — 0.03 V for the Co11 L3 complex that exhibits separated redox couples.

a Key: bpy, 2,2r-bipyridine; PA-, picolinate (2-carboxylate pyridine); acac-, acetylacetonate; 8-Q-, 8-quinolinate.

b E\/2 taken as (Ep a + Ep c)/2 for reversible couples of Mn11 L3 and Fe11 L3 complexes; as Ep a/2 + 0.03 V for L (or L—) and Zn11 L3; and as Ep c/2 — 0.03 V for the Co11 L3 complex that exhibits separated redox couples.

Table II summarizes the oxidation potentials for ligands (L) and their M11 L3 complexes with zinc(II), man-ganese(II), iron(II), and cobalt(II). The difference in the potentials for the free and complexed ligands is a measure of the metal(III)-ligand covalent-bond-formation energies (—AGbf); these are summarized in Table III. All of the data are consistent with ligand-centered redox processes.

FIGURE 5 Cyclic voltammograms of 3 mM solution in MeCN (0.1 M tetraethylammonium perchlorate [TEAP]): (a) bpy; (b) FeII(bpy)3+ ; (c) CoII(bpy)3+ ; (d) ZnII(bpy)2+ .Conditions: scan rate, 0.1 Vs—1 ; 250C; glassy-carbon working electrode (0.09 cm2); SCE vs NHE, +0.244 V.

FIGURE 5 Cyclic voltammograms of 3 mM solution in MeCN (0.1 M tetraethylammonium perchlorate [TEAP]): (a) bpy; (b) FeII(bpy)3+ ; (c) CoII(bpy)3+ ; (d) ZnII(bpy)2+ .Conditions: scan rate, 0.1 Vs—1 ; 250C; glassy-carbon working electrode (0.09 cm2); SCE vs NHE, +0.244 V.

Figure 5 illustrates that the oxidation of the FeII(bpy)3 complex is reversible and ligand centered,



Fem(bpy)3+ d5sp2

(Noteworthy are the three reversible one-electron reductions for this complex.) The electron that is removed from the FeII(bpy)3+ complex comes from the ligands to give bpy+-, which couples with one of the four unpaired electrons of the iron(II) center (d6sp) to give a third covalent bond [FeIII(bpy)3+, d5sp2; S = 5/2]. The difference in oxidation potentials for FeII(bpy)2+ and free bpy [Eq. (94)] is a measure of the FeIII-bpy+ bond energy [-AGbf = (2.32 - 1.30) 23.1 = 23.6 kcal mol-1]. The potential that would be required to remove an electron from the d6sp manifold of the iron(II) center of FeII(OH2)6+ or FeII(bpy)3+ would be greater than the first ionization potential of atomic iron (7.9 eV).

B. Reductive Electrochemistry

1. Free Electron (e-); Ultimate Base and One-Electron Reductant in Any Matrix

The free electron interacts with all atoms and molecules that have finite electron affinities to produce anions and, thus, is unstable in all but the most inert liquids. Electrochemistry attests to this general axiom and provides a convenient means to evaluate the energetics for the addition of an electron to solvent molecules and to species within a solution, e.g., e- + H2O(aq)--[H-] + HO-(aq)



E-1.86 V vs NHE (98) e" + Me2S(O)-^H2^S(O-)Me + [H-] (99)

TABLE III Apparent Metal-Ligand Covalent-Bond-Formation Free Energies (—Agbf) for Several Manganese, Iron, and Cobalt Complexes



A. Manganese

(8Q)2MnnI-8Q (acac)2MnIn-acac (PA)2MnnI-PA [(bpy)2MnnI-bpy]3+

B. Iron





C. Cobalt

(8Q)2CoIII-8Q (acac)2CoIII-acac (PA)2Com -PA [(bpy)2CoIn-bpy]3+

23 31

16 21 35

>46b a ( AGbf) — [£l/2[ZnL-/Zni2(L )] E1/2(ML-/M(-L)L2)] x 23.1 kcal mol-1.

b (-AGbf) — [Ep,a (ZnL/ZnL+0 - Ep,a(ML/M-L+)] x 23.1 kcal mol-1 ; L — (bpy)3 or (Ph3PO)4.

Hence, reductive electrochemistry converts electrons (e-) via the solution matrix at the interface to atoms and anions. The solution outside the inner double layer never is exposed to an electron. Some examples of such innerdouble-layer electron transfer include

The electrochemical reduction of permanganic acid [HOMnVII(O)3], which is traditionally represented as a metal-centered electron transfer to change Mn7+ to Mn6+, is another example of a ligand-centered process,

(O)3MnVII-OH + H+O + e- MnVI(O)3 + 2 H2O -AGbf = 28 kcalmol-1 (pH1). E0/, +1.45 V (106)

Comparison of this with the reduction of free hy-droxyl radical (HO-) [Eq. (104)] provides a measure of the HO-MnVII(O)3 bond energy [-AGbf = (2.66 -1.45)23.1 = 28 kcal mol-1]. The other strong oxidants [(HO)2CrVI(O)5 and HOCeIV(OH2)3+] that are used for aqueous redox titrations are reduced by a similar path,

(O)2CrVOCrVI(O)2OH + 2H2O

An important point in these electron-transfer reductions is that the primary electron acceptor is the hydronium ion (H+O), which is transformed to a hydrogen atom (H-) that reacts with HO- (either free or bound via a covalent bond to the metal center). Thus, in the reactions of Eqs. (103), (104), and (106)-(108), the oxidant in each is the hydro-nium ion (H3+O) and the reduction potential is determined by the H-OH bond energy (-AG BF) of the product H2O, minus the metal-OH bond energy [Eqs. (106)-(108)].

Under alkaline conditions MnVIIO4- is reduced via direct electron addition to one of the bound oxygen atoms,

The extent of the stabilization of the oxygen atom in MnVIIO4- is indicated by the reduction potential for a free •O' atom,

In summary, the electron-transfer reactions for metals and metal complexes are ligand centered (or solvent centered). In each case the potential for the oxidation of free ligand is decreased in the presence of metal or reduced-metal complex by an amount that is proportional to the metal-ligand bond energy (-AGBF).

2. Electro-Induced Hydrogénation

Reductive electron transfer in an electrochemical cell occurs by insertion of an electron from the electrode (cathode) into the solution matrix within the double layer of the electrode/solution interface {e.g., H2O/(H+O)(ClO-); [Na1 (OH2 )+](ClO- )}.

The reductive processes of Eq. (111) may be facilitated by the presence of substrates to stabilize the H-atom product. For example, in pH 0 water, the reduction process is shifted by +4.82 V when hydroxyl radicals (HO) are present,

(-AGBF, 111 kcal mol and by +2.10 V at a platinum electrode,

H+O + Pt^Pt-H + H2O E°, 0.00 V (-AGbf, 48 kcal mol-1)

In the presence of benzoquinone (Q), the shift is +2.80 V, O—H

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