The Effect of Complex Formation upon the Redox Potentials of Metallic Ions Cyclic Voltammetry Experiments Jorge 0. Ibaiiezl Universidad Iberoamericana, Cerro de las Torres 395, 04200 Mexico, D.F., Mexico lgnacio Gonzalez Universidad Autonoma Metropolitana-lztapalapa, 09340 Mexico, D.F.. Mexico Marw A. Cardenas CINVESTAV-IPN, Ap. Postal 14-740, 07000 Mexico, D.F.. Mexico The modification of one or more of the properties of transition metal ions (e.g., solubility, optical density, redox potential, electric charge, stability) is frequently required in order to use them for specific applications, such as medicinal and personal hygiene products (I), food, cleaning, and photography (2), photoelectrocbemical cells (3-5), redox flow microcells (6, 7), electrochemical reaction initiators (8,9), biological fuel cells (lo), electron acceptors for hydrogen production ( l l ) , electrocatalysis (12,13),etc. Suchmodification of properties can he achieved in most cases by complex formation, where the transition metal ion reacts with an organic or inorganic electron donor ligand. In the following experiments, students taking analytical, inorganic, or electrochemistry courses will he able to prepare in situ soluble complexes of Fe(II1) with different ligands and to observe and estimate the change in the formal potential E0' (14) that the Fe3+ undergoes upon complexation. In addition, they will he able to form soluble complexes of two different metal ions, Fe(II1) and Co(I1) with the same ligand (ortho-phenanthroline), and likewise observe the effect pro-
' Author to whom correspondence should be addressed.
duced upon the formal potentials of the two ions by COGplexation with the same ligand. The variations of such potentials can be estimated by using the technique of cyclic voltammetry (CV), which has been widely described and used in this Journal (15-21). Even though CV is not generally used to determine exact Eo' values, it allows to perform a fast and reasonably good determination by doing a single scan in just one solution, whereas, if equilibrium techniques were used (such as potentiometry), several measurements with several solutions would be required for each metal ion or complex. Theory
The standard reduction potential of the redox pair of a metal ion (e.g., Fe2+/3+)is modified upon complexation due to the different changes in free energy of each half. The following sequence of equations (22) shows this effect : s a solvated metal ion in different oxidation states (Mm+, M(m-n)+)that reacts with a noncharged ligand (L): Mm+ + ne- + ~ i m - n ~ t
Mmt
+ pL =ML;+
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A G ~= -nFE.,
(1)
AG= ~ -RT In K ,
(2)
February 1988
173
where
and
are the formation constants of ML:+ and MLim-")+, respectively. If we add eq 3 to eq 1and substract eq 2, we obtain: MLF+ + ne-
+ ML:,-")+
+ ( p - q)L AG, = -nFE:q
+ RT In [KmIK,_,I
(4)
Now, if we divide AG, by -nF,
which shows that the standard reduction potential of the metal ion is modified hv a term that involves the natural log of the ratio of tee equilibrium constants. It is necessary to point out that E,, (, is a standard potential, whereas the formal potentials t i a t are obtained from the following experiments contain both the standard potential and the contributions from other components of the medium (14). However, a relationship similar to eq 5 is found for the formal notentials when such contributions remain essentiallv constant; in this case, the ratio of the formation constants is actually the ratio of the conditional formation constants (which depend upon the conditions of the medium), and the changes in En' will be basically the same as those of E n . In order to simplify the experiments described below, we chose to prepare metal complexes which exhibit Nernstian or nearly~e;nstian behavior, and which are stable under the chosen conditions. From the theory of CV (14), the peak potential E (for reversible systems) is related to the formal potential EE' in the forward potential sweep as follows:
where Do and D, are the diffusion coefficients of Mm+and M[m-"l+, respectively. If the concentration of the ligand L is high (and thus the reaction is not controlled by the diffusion of L) and the ML;+ and MLim-") species can be formed, we obtain:
Figure 1. Cyclic voltammograms at a platinum electrode in H20. Scan rate: 100 mV/s. Reference electrode: saturated calomel electrode W E ) . Curves: (A) 5 X MFe(SO& In 0.1 M NaCIO,. (B)5 X I O P M FeNH,(SO&in 0.1 M Na2EDTA. (C) 5 X lo-' M Fe(N03)3.9H20, 0.01 M c-phenanthroline In 0.1 M HNOs. Vertical scale: 5 pA/dlv. for CuNeS A and 6. 0.5 pAldlv. for curve C.
Experimental
The metal ion complexes were prepared by dissolving the complexing agent in 30-40 mL of Hp0; then the metal ion salt was dissolved and finally the supporting (inert)electrolyte was added to the solution. The solution was stirred all throughout this preparation process. The final concentrations of the solutions were as folM FeNH1(SO&, 0.1 M NaCI04;5 X M FeNHd lows: 5 X (SO&, 0.1 M Nal EDTA; 5 X 10-3 M Fe (N03)3.9HzO20.01 M 0phenanthroline, 0.1 M HN03;5 X M Co (N03)~6HpO, 0.01M ophenanthroline, 0.1 M HNO, The applied potential was controlled with a CV-27 Voltammograph (Bioanalytical Systems), although simpler potentiastats would work as well; the signal (i vs. E ) was recorded with a Houston 100 X-Y recorder. The working electrode was a small stationary Pt disk encapsulated in a resin matrix. The potential was measured vs. a saturated calomel electrode, and the auxiliary (counter) electrode was a Pt wire. The solutions were deaerated with Np before each run. Results and Discussion As a rule of thumb, the formation of metal ion complexes stabilizes preferentiallv the ion with the highest oxidation state (22) due to the stronger bonds formed-and to the fact that the positive charges are better delocalized due to the ligand's polarizability. This observation is reflected in the decrease of the standard reduction potential of the redox metal ion pair upon complexation (see for example ref. 231, eq 5 shows that E, ,,, < E L Then, since, if K, > K ligands which make the K,/K,-, ratio larger, will make the reduction potentials to he displaced more negatively.
,-.,
where D' and D; are the diffusion coefficients of MLFf and MLP-")', respectively, and CLis the ligand's concentration In the solution (which is assumed to he larger than the total concentration of Mn+ and/or M(m-n)+).Now, if we consider that (DOID,) 3 (DJDJ and p = q , we obtain the much simpler equation: E
,,.,,I..-E,
=E&,,-Ez
= ln(K,-,lK,)
(8)
where K L and K k - , are conditional formation constants (which depend upon the conditions of the medium); this equation shows that the difference between the E, of Mm+ with and without the presence of L (if both Mm+ and M "-",- form complexes with L) permits the calculation of the ratio of the conditional formation constants. In order to observe these potential shifts, the preparation of Fe(II1) complexes with different ligands is described, as well as the preparation of a Fe(II1) and a Co(I1) complex with the same ligand. 174
Journal of Chemical Education
,
One Metal Ion, Different Ligands The results obtained with Fe(II1) (see Fig. 1)show that the presence of EDTA originates a negatiue displacement of the reduction peak of 600 mV; according to eq 8, this means that the conditional formation constant of Fe(II1)-EDTA is roughly 10'0 times larger than that of Fe(I1)-EDTA. On the contrary, the presence of o-phen originates a positiue displacement of the reduction peak of 370 mV (see Fig. 1); this means that the conditional formation constant of Fe(I1) ophen is roughly lo5 times larger than that of the Fe(II1) complex. Why does this happen? The explanation for this can he understood if we analyze the Fe ions spin states. Using the mnemonic suggested by Nance (24) we can remember that the Fez+ corresponds to a 3d6 configuration, whereas Fe3+ is a 3d5. Since the o-phen ligand has vacant orbitals (7, 251, it is obvious to think that if any of the
favors less the d ~ - T Lbond formation with the ligand than its counterpart, Co3+ (d6),which is in turn diamagnetic and has a low spin (t;,) basal state in the complex. Then, the metal ion c o m ~ l e xof o-phen with Co3+ is more stable than that formed with Co2-. This phenomenon is responsil,le for the negative displacement observed in the potential, as eq d correcily predicts.
Figure 2. Cyclic voltammogramsat a platinum electrode in HZO.Scan rate: (A) loo mV/s, (0) 50 rnV/s. Reference electrode: SCE. Solution: 5 X 10P M ColNO31r6H20. 0.01 M whenanthroline in 0.1 M HNOs. Curves: (A) the Co (ill Ill) e p h i p;ak, and ( 0 )the lull spectrum, showing that the Co(il/lll)when redox couple is me only couple observed within the water's potential window. Vertical scale: (A) 0.5 pA/div., (8) 1 pA/div.
Conclusions Cyclic voltammetry can be used to demonstrate the variations in the reduction potential that a metal ion undergoes upon complexation with different ligands that stabilize one oxidation state more than another, as well as the different variations that one ligand may cause upon the potentials of two ions of different metals. It can also be used, in the absence of important solution effects upon the stabilities of the complexes, to estimate the value of the conditional formation constants of complexes by using the method described above. Some important factors that may be analyzed in order to obtain more information about a given system characteristics, have been previously discussed in this Journal: scan rates, reversihility, and pH (20), detection of intermediate steps (15),factors affecting i-E curve shapes (18),electron transfer rate constants, transfer coefficients, return potenrials, and coupled chemical reactinns (17, 18).dissolvid oxygen, potential limits and surface effects (19). For simplicity, we have omitted a discussion of the i n t l u e n c ~of
the
lirnnd's
concentration; such discussion can he found in anzytical textbooks (30,31). oxidation states of Fe can form d,,~-~i,..d bonds, this comdex will he much more stable than that formed with the other oxidation state, since there will be a back donation of electronic charee from the metal ion to the ligand. Since ophen originates a large separation of the d orhitals in the metal ion (high ligand field) (23, 26), its complex with Fe" hay a low spin configuration '2,. which favors the d-n bond formation with the o-phen vacant orbitals. On the other hand. the formation of the FetllI) o-ohen c o m ~ l e xis less favorkd since (a) Fe3+has poorer r-donor propeities due to its hieher ~ositivecharee (7) . . and (b) the Fe in this metal ion comaex das a t$ configuration @), which favors less the formation of bonds with the ligand, (26). The agreement of the results obtained for the K,/K,-, ratio for the EDTA case, with the data reported by Laitinen and Harsis (27) shows that CV can be used to estimate quantitatively one formation constant if the other one is known. However, the same comparison in the o-phen case shows a disagreement with the reported data (27); such disagreement has recentlv been exnlained bv the effect of the DH upon the stabilities of thesecomplex& (28).In summary,ihe Fe3+ is stabilized bv the lieands studied in the followinnorder: EDTA > H ~ O ~> - ~ h ; n .
-
One ligand, DifferentMetal Ions The results obtained with the ligand o-phen (Figs. 1C and 2) show that this ligand forms complexes with both ions, Co2+ and Fe3+, modifying drastically their formal reduction potentials (the Co3+ eCo2+has a reduction potential of about +1.60 V vs. SCE (29),and the Fe3+ e- Fez+was discussed above). However, why does the potential for the Co2+"+ couple decrease upon complexation, contrary to the Fe2+/3+ case just discussed? The o-phen originates a lowspin basal state (t;,e,) in Co2+ (d7 configuration), which
+
-
+
-
Acknowledgement
The authors wish to thank Ralph S. Becker, Chausoo Choi, and Alberta Rojas for helpful discussions. This paper is dedicated to Margarita Watty (of the Universidad Iheroamericana) on the occasion of her 60th birthday.
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