Anal. Chem. 1994,66, 2005-2009
Multistep Mechanism in the Electrochemical Oxidation-Reduction of Fe-Citrate Complexes Lorenza Battistinit and Jests Lbpez-Palacios' Department of Analyfcal Chemistry, Colegio Universitario de Burgos, Apdo. 23 1, E-09080 Burgos, Spain A study was made of the electrochemical behavior of the Fecitrate system in aqueous medium using cyclic voltammetry. Potential sweeps between 200 and -650 mV vs Ag/AgCl show the existence of two Fe(III)/Fe(II) redox systems. Tbe first, with a half-wave potential close to 0 V, proves to be a reversible systdm, while the second, with a half-wave potential around -0.4 V, shows the characteristics of an E& process. Competitive reactions between hydroxy and citrate complexes of Fe(II1) are considered. Both equilibrium measurements and kinetic evolution of the system are used in order to propose a scheme for the set of chemical and electrode rections studied. The different behavior observed in solutions of Fe(II1) or Fe(XI) is suggested as a method for the easy electrochemical differentiation of the oxidation state of Fe in solution.
The formation of complexes with polycarboxylic ligands is one of the most accepted methods used for the electrochemical determination of iron in an aqueous medium. In particular, the citrate ion has been used as a complexing agent of the Fe2+and Fe3+ions with excellent results from Lingane's polarographic studies' to the most recent workse2 Despite satisfactory analytical results, the complex redox behavior of the Fe-citrate system in aqueous media still presents many dark areas as, for example, the exact formulationof the species participating in the electrode reactions. The possibilitiesthat this system offers for the voltammetric speciation of iron in its I1 and I11 oxidation states have barely been considered.3 However, the very different thermodynamic stabilityconstants of the ferrous4trate and ferric-citrate c ~ m p l e x e ssuggest ~.~ sufficiently distinct electrochemical behavior which should, in theory, be observable using voltammetry. Speciation of iron in its oxidation states by e l e c t r ~ c h e m i c a l and ~ ~ ~non.~ electrochemicaP1Omethods continues to be an interesting subject. Consequently, the aim of this work is to analyze the mechanism of the reactions taking place during the electrochemical oxidation-reduction of iron in aqueous citrate t Current address: Deprtment of Analytical Chemistry, University ofcamerino, Camerino, Italy. (1) Lingane, J. J. 1.Am. Chem. Soc. 1946, 68, 2448. (2) Bond, A. M.;Pfund, B. V.;Newman, 0. M.G. Anal. Chim. Acta 1993,277 (I), 145. (3) Heyrovsky, J.; Zuman, P. Practical Polarography;Academic Press: London, New York 1968; pp 116-118. (4) SillCn, L. G. Stability Constants of Metal-Ion Complexes;Supplement No. 1, Part I, Special Publication No. 25; The Chemical Society: London, 1971. (5) Ha", R. E.; Shull, S. M.,Jr.; Grant, D. M.J. Am. Chem. Soc. 1954, 76, 2111. ( 6 ) Kennedy, C. D. Analyst 1990, 115, 1067. (7) Doherty. A. P.;Forster, R. J.; Smyth, M.R.;Vog, J. G. Anal. Chem. 1992, 64, 572. (8) Luque, de Castro, M.D. Talanta 1986,33, 45. (9) Faizullah, A. T.; Townshend, A. Anal. Chim. Acta 1985, 167, 225. (10) Cox, J. A.; AI-Shakshir, S. Anal. Lett. 1988, 21, 1757.
0003-2700/94/036&2005$04.50/0 0 1994 American Chemical Soclety
solutions. In the characterization of the different steps of the process, an uncomplicated well-known experimental technique such as staircase cyclic voltammetry with mercury electrode has been chosen, thereby allowing the comparison of the experimental results with a number of models developed some years ago. Special attention has been paid to the different behavior of Fe(II1) and Fe(I1) solutions, which could be used in the later development of simple procedures for the speciation of iron.
EXPERI MENTAL SECT1ON Instruments and Reagents. In the majority of the voltammetric experiments, an EG&G Princeton applied research system was used, consisting of a potentiostat Model 263 and an electrode stand Model 303A static mercury drop electrode in the handing mercury drop mode (HMDE) with a nominal area of 0.01 cm2. A platinum wire auxiliary electrode and an Ag/AgCl reference electrode wereused. Software of the same make, Model 270 (version 4.0), was used for the acquisition and treatment of the data on a COMPAQ 3/25 personal computer with coprocessor. To obtain the kinetic parameters and compare the experimental data to predefined models, software E G h G PARC Model 271 was used, based on the COOL algorithm."-'3 Repeated cyclic voltammograms on the same drop of mercury were obtained using an AMEL 55 1potentiostat, with a Tacussel GSTP signal generator and a Philips 8133 XY recorder. In this case, the mercury electrode was a Metrohm EA 290 HMDE. All the chemicals used were of analytical grade without further purification. The mercury was tridistilled beforehand. The solutions were prepared in the usual way with distilled water purified in a Barnstead Nanopure system. When it was possible, sodium citrate was used as a supporting electrolyte; in the experiments in which relatively low citrate concentrationswere required, NaNO3 was used with the same purpose. In the control of pH, when it was necessary, buffer solutions based on citric acid/citrate ions were used, avoiding in this way the introduction of strange species in the system. The solutions were deoxygenated by bubbling N2 (99.997%) at least 15 min before each experiment. Except when indicated, all experiments were done at a temperature of 25 OC. The remaining experimental conditions are indicated in the corresponding sections. (1 1) O'Dca, J. J.; Osteryoung, J.; Lane, T. 1.Phys. Chem. 1986.90, 2761. (12) Chin, K. Y.;Prasad, S.;ODca, J. J.; Osteryoung, J. AMI. Chfm.Acta 1992, 264, 197. (13) OD-, J. J.;Osteryoung, J.;Osteryoung,R.A.~..hys. Chem. 1983,87,3911.
AnalyticalChemisby, Vol. 66, No. 13, ,tu@ 1, 1994 2005
I
do0
-500 -400
-300 -200
-100
0
100
200
E (mVvs Ag/AgcI) Flguo 2. Repeated cyclic vo#a"ogam on HMDE of Fe(II1) 10-4 mold& h citrate lo-* mol bnJ at pH = 5. Fkst cyck k, noted 1. Scan rate. 250 mV 8-l.
1
-0.5
-0.4
-0.3
-0.2
, -0.1
0
0.1
E (mV vs Ag/AgCI) Figure 1. (A) Staircase cyclic voltammogram of Fe(II1) 2 X lo4 mol dm4 in citrate 2 X md dm4 at pH = 6, obtained on HMX wlth a scan rate of 250 mV s-l. Starting potential, 50 mV; electrode area, 0.01 cm2. (B) Analogous voltammogram for a solution of Fe(I1) 2 X lo4 mol dm4 In &rate 1 t 2 mol dm4. The same experimental conditions were used.
RESULTS AND DISCUSSION As is well-known, cyclic voltammograms of aqueous solutions of Fe(II1) salts in the absence of complexing agents show several poorly defined waves that can only be reproduced with difficulty. These signals can easily be attributed to electrode reactions of the different hydroxocomplexes of the Fe(III)/Fe(II) couple. The addition of citrate ion to an aqueous solution of Fe(II1) causes a drastic change in the voltammograms, which now show (Figure 1A) in a wide pH range two cathodic peaks, c1 and CZ,and one anodic peak, a], all well-defined and reproducible. The first and most basic difference is that analogous solutions of Fe(I1) allow one to obtain voltammograms where only one anodic/cathodic couple of peaks can be observed (al/cl in Figure 1B), which coincides exactly with what appears at more positive potentials in Fe(111) solutions.
10-4 M causes the disappearanceof c1 besides a badly defined c2. A ratio 3 X 10-4 mol dm-3 of Fe(III)/10-2 mol dm-3 of citrate allows one to observe on the voltammogram both c1 and cz fairly well defined. (ii)Procedure Followed in Preparing Solution. Solutions prepared by dissolvinga ferric salt directly in a citratesolution show a small peak, cz, and a noticeably higher one, c1. The opposite effect is observed when the ferric salt is dissolved in water and then mixed with thecitratesolution in theelectrolytic cell. (iii) Time Elapsed between Preparation of Solution and Measurement. Voltammogramssuch as that shown in Figure 1A evolve in time, and an increase in peak c1 is observed at the same time as a decrease in peak CZ. These three factors together would suggest the coexistence in the solution of two reducible species whase concentrations vary in an inverse way, a hypothesis which is confirmed by the evolution of the peak heights when repeated potential sweeps are made on the same drop of mercury (Figure 2); as can be seen, the increase of the one peak always occurs at the expense of the other. Furthermore, the total absence of an anodic peak corresponding to the cathodic peak cz suggests an irreversible chemical reaction after the electron transfer, which causes the rapid disappearance of the product of this reduction. Using these observations, an initial scheme describing the complete process can be proposed: 0x1
ne= Ri
(1)
abw The relative height of thecathodic peaksobserved in Figure 1A is shown to be independent of the pH, but varies with the followingfactors: (i) the ratio of concentrations C F ~ I I I ) / C & ~ ~ ; (ii) the way in which the solutions were prepared; and (iii) the time which elapses between the preparation of the solutions where 0x1 and 0 x 2 are the species giving the peaks c1 and CZ, and obtaining the voltammograms. respectively, and R1 and RZ are the corresponding reduced ( i ) C ~ , ( ~ ~ ~ / C cRatio. i r r a rFor r a given concentrationof Feforms. Each of the stages which make up the complete process (111), the peakcl became higher with increasingconcentrations is analyzed below. of citrate; the peak cz being small in these conditions. High Electrochemical Step 0 x 1 + R1. From the voltammetric values of the ratio CFe(III)/C&trate caused the opposite effect. data, theOxl/Rl system appearsas perfectlyreversible.Cyclic Thus, a solution of Fe(II1) 1 V M in citrate 0.2 M gives a potential sweep between 0.05 and -0.25 V carried out at voltammogram in which only C I can be observed, while a different concentrations of Fe(II1) and of citrate show the characteristics of reversible systems; the peak potentials and solution of the same concentration of iron with citrate also
2006
AmiyticaIChemb~,Vol. 06, No. 13, Ju& 1, 1994
Considering these facts altogether, a scheme of reaction for this stage of the process can be established, depending on the pH of the medium:
-u,4‘
;
2
4
5
PH
Flgure 3. Shift of the potential peaks wtth pH in solutions 3 X lo4 mol dm3 of Fe(II1) and 2 X lo-* mol dm3 of citrate. (+) Potential of the anodic peak al; (0)potential of the cathodic peak cl; (X) potential of the cathodic peak cp. Table 1. Equlllbrlum Constants for Cltrlc Acld Dlrroclatlone and Iron-Cltrate ComplexeS.
equilibrium
constant value
H3L * H+ + HzLH2L- * H+ + HL” HL-* H+ L3Fe3+ HL2- * FeHL+ Fe3+ + L s + FeL Fez(HIL)Z + 2H+ 2Fe3+ + 2Ls Fez+ + HL2- + FeHL Fez+ L3- ==FeL-
Ka1 = 1.096 X Ka2 = 4.169 X Ka3 = 2.089 X 10-6 KI = 2.00 X lo6 K2 = 2.51 X 10” K3 1.58 X 10” K4 4.41 X 10’ K5 = 2.51 X 104
+
+
+
a Ls = citrate ion. All the constants are taken from ref 4 except K4, which was obtained from ref 5.
the relationship ipox/ip’edremain unchanged when different scan rates between 0.02 and 1 V s-l were used. When the COOL algorithm is applied to the experimentally obtained curves, one can obtain a perfect correspondence with the reversible model for a number of electrons transfered, n,equal to 1; in each case fits to the model with correlation coefficients greater than 0.998 and values of x2around lo-” were obtained. However, both the anodic peak potential, Epl, and the cathodic peak potential, E,1, depend on the pH. Figure 3 demonstrates this dependence graphically, showing two straight segments for each peak. It is possible to conclude from the slopes of these lines that at pH values below approximately 4.0-4.5 the reduction of 0 x 1takes place with the previous taking of two hydrogen ions for each transfered electron (overall reaction: Ox1 + 2H+ + e- Rl), while at a higher pH only one proton participates for each electron (Oxl H+ eRl’).
+
+
-
-
From simple mass and charge balances and taking into account the thermodynamic stability constants contained in Table 1, it can be obtained that Fe2(K1L)z2-, where L3- is the citrate ion, is the most stable ferric-citrate complex in solutions with enough high relative concentration of citrate. Taking into account that the voltammetric waves correspond to a one-electron system, a previous dissociationof the binuclear complex can be supposed.
(a) pH < 4.0
I/ZF~,(H-,L);-
(b) pH > 4.5
1/2Fe,(H-,L);-+
2H+ FeHL’ e-
H+
FeHL
(2)
c
FeL
FeL-
(3)
This involves a CE (chemical-electrochemical) mechanism. But, as has been stated before, the system shows voltammetric behavior characteristic of a reversible couple. That is to say, the preceding dissociation of the binuclear form proceeds in such a rapid way that it does not influence the electron-transfer step. In accordance with the analysis made by Savbnt and Vianello1”16 and Nicholson and ShainI7for this case of coupled chemical reactions, our system shows, under the experimental conditions mentioned here, a behavior typical of a diffusioncontrolled process. Therefore, relatively high values must be assumed for the kinetic parameter, A, and/or for the constant of the chemical reaction. The existence of the steps described below makes it difficult to obtain more detailed information on the chemical reaction in this phase of the process.
-
Chemical Step0x2 0x1. As stated above, the procedure by which the solutions to be electrolyzed are prepared seems to be an important factor in the relative value of the two cathodic peak currents, ih,/ih2. If a ferric salt is dissolved in an aqueous medium containing citrate ion, one of the complexing reactions which in Table 1 correspond to the equilibrium constants K1, K2, or K3 will immediately take place. But if the iron salt is simply dissolved in water, it is well-known’*that the resulting aquo-ion, F e ( H ~ 0 ) 6 ~which +, is only stable in extremely acid mediums, evolves rapidly toward one or more of the highly stable hydroxylate forms, Fe(OH)n(3-n)+.The addition of citrate ion to a solution which already contains hydroxo complexes of Fe( 111) causes the displacement of the OH- ion by the citrate ion, a process which takes place through a multistep mechanism19with some slow steps. The result is that complexes of Fe(II1)xitrate may coexist for a time with mixed complexes of the FeL(OH),” type. This explains the appearance of the second peak, c2, in the cathodic sweep, notably intense when the solutions have been prepared in the way described here and atributable to the reduction of the hydroxylate form. As has been pointed out before, the system evolves chemically in time, it being possible to measure the disappearance of this hydroxylate form by the decrease of peak c2. Taking into account the acidity of the working solutions, pH 4, it could be supposed as a first approximation that the transformation of the hydroxocitrate-Fe(II1) complex is due to a protonization reaction that can be formulated as (14) SavCant, J. M.; Vianello, E. Electrochim. Acta 1%3, 8, 905. (15) Savhnt, J. M.; Vianello, E. Electrochim. Acta 1967. 12, 629. (16) Savhnt, J. M.; Viancllo, E. Electrochim. Acta 1967, 12, 1545. (17) Nicholson, R. S.; Shain, I. Anal. Chem. 1964, 36(4), 706. (18) Kotrly, S.; Sucha, L. Handbook of Chemical Equilibria in Analytical Chemistry; Ellis Honvood: Chichcster, 1985; Chapter 13. (19) Mentasti, E.;Baiccchi, C . J. C w r d . Chem. 1980, IO, 229.
AnalyricalChemistry, Vol. 66, No. 13, July 1, 1994
2007
FeL(OH),"
k
+ (n - 1)H'
.+
1/~Fe2(H-,L):-
+ nH20
(4)
Y -5(1
A
M
8
where k is the rate constant of transformation of FeL(OH),* in Fe2(H-1L)z2-. A voltammetric followup of the kinetics of reaction 4 was done. The decrease with time of the peak current,,i was measured in solutions with pH from 4 to 6 , with the same concentrations of Fe(II1) and citrate as those indicated in the previous paragraph (Fe, 3 X lo4 mol dm-3; citrate, 2 X 10-2 mol. dm-3) and varying the temperature between 30 and 70 OC. The scan rate of potential was in all cases 250 mV s-l, and the time of each experiment was about 120 min, starting at the moment in which the solutions of Fe(II1) and citrate were mixed in the cell. Under these conditions, values of kob, the observed rate constant, were obtained, which follow the rate equation
Flgure 4. Anheniw plot for the values of the fkstorder rate constant,
where kob = k[H+Ib,and b is thereaction order with to respect the hydrogen ions. No significatively different values of kob were obtained in solutions with pH 4, 5 , or 6 ; this involves b = 0, that is to say, no hydrogen ions take part in the reaction. The values of k obtained for the different temperatures give good linear Arrhenius plots (Figure4), which prove thevalidity of the first-order rate equation assumed. From the slope of these lines, the medium value of 9 1.2 kJ mol-' was obtained for the activation energy, Ea. This relatively high value explains the slowness of the transformation of the hydroxo complex in the dimer F ~ ~ ( H - I L ) ~ ~ - . Once the value of n is known, eq 4 can be rewritten as
is undoubtedly the most interesting stage in the process, not only because of the nature of the electrochemical-chemical step itself but also because we are dealing with a differentiating voltammetric signal that can be observed in the reduction of the Fe(III)-citratesystem but not in that of the Fe(II)-citrate. Using solutions of different concentrations of Fe(II1) ranging from 1od mol dm-3 to 6 X 1Pmol dm-3 and a constant concentration of citrate of 2 X mol dm-3, cyclic voltammograms from -210 to -600 mV at different sweep rates, u, were obtained. The absence of an anodic peak precluded the use of criteria based on the ratio of peak currents to characterize the system. The shift of the peak potential with log u follows the equation
k
FeL(OH)-- l/2Fe2(H-,L);-
+ H,O
(6)
If that is so, the 0 x 2 corresponds to the form FeL(0H)-, and eq 6 represents the overall reaction in the complicated process of formation of the binuclear Fe2(H-lL)z2- ion from the hydroxylate mononuclear complexes. Similar reactions involving the formation of oxo and/or hydroxo bridges have been described for iron(II1) in analogous systems.20*21 EC Step 0x2 ;r R2 R1. When the species FeL(0H)is reduced, one may expect the formation of FeL(OH)2-, where the iron is in oxidation state 11. But if the stability of the hydroxo complexes of Fe(II1) is perfectly admissible given the extremely acid character of the iron at this oxidation level, one cannot expect identical behavior for the Fe(II), which has a far smaller capacity to accept pairs of electrons. Consequently, it is very reasonable that a rapid protonation of the reduced form occurs after the electron transfer
-
FeLOH2- + H+
k'
FeL- + H,O
(7)
k'being the rate constant of the reaction. This involves an EC mechanism which characterizes the cathodic peak c2; it ~~
~~~~~~
A
E
i
A
a
-7
8
Q 8 -9 -1 2,95
3,OO
3,05
3,lO
3,15
3,20
3,25
Analytical Chemism, Vol. 66, No. 13, ~ u l y1, 1994
3,30
1000/T ( K ' ) k, obtained at temperatures between 30 and 70 OC. (0)pH 4; (0) pH 5; (A)pH 6. (Other expertmental detalls in the text.)
E = -29.33 log u - 205.49
(R2 = 0.997)
(8)
the slope of which agrees exactly with that characteristic of the first-order chemical reactions in the EC model. Again, we applied the COOL algorithm to compare the experimental data to the E& model (reversibleelectrochemical step followed by an irreversible first-order chemical reaction) first defined in the works of Nicholson and Shainl7 and Nadjo and Sav6ant.22 Two parameters were obtained from the fit: the half-wave potential of the system, E1p, and log (k's,),wherest is thestep timeusedin thestaircase potential sweep (0.02 s in our case). From this last parameter k'was calculated, giving an average value near 5 X lo3 dm3 mol-' s-l. A series of these results, corresponding to solutions of 2 X 10-4 mol dm-3 of Fe(III), is summarized in Table 2. The participation of hydrogen ions in the followupchemical step explains the dependence of the peak potential of the wave c2 on pH, as can be seen in Figure 3. The slope of the fitted line would allow one to obtain a value of 1.34 protons for each electron interchanged, but this value should not be considered an average because the existence of the kinetic mechanism prevents one from making simple deductions in this respect. However, it is reasonable to believe that in a basic, neutral,
~
(20) Nisida, Y . J. Chem. Soc. Dolton Trans. 1985, 2375. (21) Spiro, T.G . Inorg. Chem. 1987, 26, 2063.
2000
-6-
(22) Nadjo, L.;Savtant, J. M.Electrochim. Actu 1973, 48, 113.
Table 2. FII of Expohenla1 Data to E& M0d.l Udng COOL
'/2Fe2(H-jL)22-
Algorlth"
quality of fit scan rate (mV s-I)
no. of exppoints
R
50 100 150 300 500
781 391 261 131 79
0.999 0.998 0.999 0.998 0.998
0
rate constant of chemical step, k'(dm3 mol-' s-I)
X2
5.42 X 4.34 X 2.37 X 1.42X 4.81 X
l@I' l@I7 l@I7 1@l6 1@16
The parameter x2 is here defined as x2 = [ 1/(!MU -N - i0b)2. R is the usual correlation coefficient.
+
'/2Fe2(H-1L)2-
P-
(9)
while in a more acidic medium, two hydrogen ions will be interchanged
FeLOH- + e- * FeLOH2-
-
+ZH*
FeHL + H,O
e= FeLfast
+e-
I
(1 1)
H+
* FeLOH2-
or in a very acidic solution,
or slightly acid medium the chemical step would proceed with the interchange of one proton
+ e- * FeLOHZ-+H* FeL- + H 2 0
FeL
8bW
~
1
-HP
(10)
Global Scheme of the Process. Once the different stages which make up the complete process of oxidation-reduction have been analyzed, the initial scheme (1) can be explained in detail in the following way:
e2H+ = FeHL' =FeHL
Slow
FeLOH'
FeLOH-
H+
FeLOH-
5081 5011 5046 508 1 5012
2)]&'olo
1
-H@
+e-
-
I
(12)
fast 2H+
FeLOH2-
From these schemes, one can understand the different electrochemical behavior of the iron in its I1 and I11 oxidation states. The nonexistence in aqueous solution of species Fe(11)-hydroxwitrate complexes makes it impossibleto observe the cathodic peak c2 in systems in which iron is only present in this oxidation state. This enabtes one to differentiate Fe(11) from Fe(1II) in solutions using simple voltammetric experiments. Based on this behavior, we are currently developing experimental procedures for iron speciation in its two oxidation states, using both the- voltammetric peaks designated c1 and c2 throughout this study. Other voltammetric techniques such as differential pulse normal (DPN) linear voltammetry are proving to give a very suitable sensibility for analytical purposes.
ACKNOWLEDGMENT This work was supported in part by Grant ALI-0441/91 from the Spanish CICYT (Comisi6n Interministerial de Ciencia y Tecnologfa). L.B. thanks theuniversityof Camerino for a postgraduate fellowship. Received for revkw October 29, 1993. Accepted March 29, 1994.. Abtract published in Advance ACS Abstracts, May
IS, 1994,
AmWbl chemlsby, Vd. 66, No. 13, Ju& 1, 1994
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