1355
REDUCTION POTENTIALS OF COMPL,EX Ioxs
Reduction Potentials of Complex Ions. The Tris( pyridine-2-aldoxime)iron(III)-Tris( pyridine-2-aldoxime)iron(11) System' by George I. H. Hanania, Dennis H. Irvine, and Fahd R. Shurayh Department of Chemistry, American Universitg of Beirut, Beirut, Lebanon
(Receised October 24, 1967)
The reduction potential E for the tris(pyridi1ie-2-aldosime)iron(III)-trisipyrictine-2-aldo.;ime~iron(ll) coiiple has been measured potentiometrically in dilute aqueous solution within the ranges of pH 6.0-7.5, ionic strength 0.085-0.015 ,It', temperature 14.5-33.2'. The data on the variation of E with pH are consistent with the assumption of one protonation equilibrium involving the ironiII) complex. .It 25.0' and I = 0 the thermodynamic pH-independent reduction potential is E,' = 0.348 i 0.001 V with AH' = -20.7 i 0.5 kcal/mole and AS' = -42.6 i. 1.6 eu. The equilibrium between the above redox couple and the hexacyanoferrate(II1)hexacyanoferrate(I1) redox couple was investigated by a direct spectrophotometric method yielding K,' = 0.9 i 0.1 with AH' = -5.2 f 1.3 kcal/mole and AS' = -18 i 5 eu a t 25.0' and 1 = 0. The large uncertainty in these results reflects the limited precision with which the equilibrium constant could be measured. These thermodynamic quantities are also calculated, with higher precision, from the known data on the two redox couples involved. The tris(pyridine-2-aldoxime)iron(III-II) system is examined in relation to analogous iron complexes of neutral organic nitrogen ligands and also to inorganic anionic complexes of the heuacyaiioferrate type.
I n a preceding comparison of thermodynamic data on redox cell reactions in aqueous solutions2 it was noted that there are wide variations in reduction potential for a series of analogous iron complexes. Thus the tris(o-phenanthro1ine)- and tris(dipyridyl)iron(III)-(11) couples have reduction potentials of about 1.1 V, whereas for the tris(4,7-dihydroxyorthophenanthroline)iron(III-11) couple the value is -0.1 V when all six hydroxyl groups are ionized. This shows the profound influence of the electrostatic environment of the metal on its oxidation-reduction potential. Moreover, the data show that the effect appears primarily as a change in exothermicity and, to a lesser extent, as a change in entropy for the cell reaction. The tris(pyridine-2-aldoxime) complex of iron is similar to the dipyridyl and phenanthroline complexes in being a !ow-spin octahedral complex in which iron is bonded to six nitrogen atoms. It differs however in having three polar oxime -OH groups. The system can therefore be used for a detailed study of the effect of acidic side groups in a ligand on the reduction potential of the chelate. The converse effect, viz., the influence of the metal on the thermodynamics of ionization of the oxime side groups in this system, has already been e ~ a m i n e d . ~I n that case, pK of the oxime group was found to decrease from 10.22 in the free ligand to 7.13 in the complex, although the charges involved are the same in both cases. This 1000-fold increase in acid strength paralleled a decrease of 6 kcal/mole in the endothermicity of ionization. I n this paper. we report a potentiometric study of the reduction potential of the tris(pyridine-2-aldoxime)iron(II1-11) couple covering ranges of pH, ionic
strength, and temperature. Since this potential is close to that of the hexacyanoferrate(II.1-11) couple, the equilibrium between the two systems was also investigated by an independent spectrophotometric method. A preliminary account of the work has been rep~rted.~
Theory Pyridine-2-aldoxime (PH) forms well-defined ferrous and ferric complexes in aqueous solution. I n each case the complex undergoes three successive ionizations Fe(P3H3)2+ -+ -+ -+ Fe(P3)-
(1)
Fe(P,H3)3+ +
(2)
-A
-+- Fe(P3)0
It is known3that for the iron(I1) complex, in eq 1, pKl< 3, pK2 = 3.4, and pK, = 7.1, and that for the iron(II1) complex, in eq 2, pK1 and pKz < 3 while pKs = 3.5. Since measurements on these systems were largely confined to the range of optimum stability of both complexes, pH 6.0-7.5, the only prototropic equilibrium involved is the third oxime ionization of the iron(1T) complex
Fe(P3H)0 = Fe(P3)-
+ H+
(3)
The Redox Reaction. The parent cell reaction for the (1) Work supported by a n Arts and Sciences research grant from the Americah University of Beirut. (2) G. I. H. Hanania, D. H. Irvine, 1%'. A. Eaton, a n d P. George, J. Phys. Chem., 71, 2022 (1967). (3) G. I. H. Hanania and D. H. Irvine, J . Chem. Soc., 2746 (1962). (4) G. I. H. HananiR, D. H. Irvine, M . S. Michaelides, and F. R. Shurayh, "Proceedings of the 9th International Conference on Coordination Chemistry," w. Schneider, Ed., Verlag Helvetica Chimica Acta, Basel, Switaerland, 1966, p 224.
Volume 72, Number 4 April 1968
1356
G. 1. H. HANANIA, D. H. IRVINE, AND F. R.
tris(pyridine-2-aldoxime)iron(III-II) redox couple us. standard hydrogen electrode is Fe(P3)0
+ ‘/zHz
=
Fe(P3)-
+ H+
(4)
where the charges of 0 and -1 are the result of three oxime ionizations in the oxidant and reductant ions, respectively. The reductant ion is also involved in the prototropic equilibrium of eq 3. All other ionization and ion association equilibria are assumed to make negligible contributions to the measured free energy change within the experimental range of conditions. On this basis, the measured reduction potential E (defined at given pH, ionic strength I,and temperature T,for equal total molar concentrations of oxidant and reductant) will vary with pH and I in accordance with the relation
E
=
E O i
T + RIn (1 + (h/K,)) -t F
Fe(P3)-
SHURAYH
+ Fe(CN)63- = Fe(P3)0 + Fe(CN)e4-
(9) where both systems are in their fully ionized forms. The thermodynamic pH-independent equilibrium constant KO, is defined in terms of activities, and the measured equilibrium constant K is defined in terms of the total molar concentrations of all species for each reactant and product in eq 9. K varies with pH, ionic strength, and temperature. At constant temperature, within the range of our experiments (pH 6.0-7.5, and concentrations about M ) , the hexacyanoferrate system is not appreciably influenced by protonation or ion association e q ~ i l i b r i aand , ~ the only variation of K arises as a result of the ionization in eq 3. Thus, at finite ionic strength, one can calculate from the measured equilibrium constant K the corresponding value for the pH-independent constant Ki which refers to the equilibrium in eq 9. The relation is simply
K,
=
K(1
+ (h/K3))
(10)
The dependence of Ki on ionic strength is readily shown to be given by where Eoi is the thermodynamic pH-independent reduction potential of the couple in eq 4 relative to s.h.e., K3 the appropriate value of the ionization constant for eq 3, h the hydrogen ion activity in the solution computed on the assumption that the measured pH = -log h, and yo and yn are the mean molar activity coefficients of free oxidant and reductant ions, respectively. At given finite ionic strength, it is convenient to calculate from the measured reduction potential E the pH-independent value Ei which is the reduction potential for the parent couple in eq 4. Thus defining
Ei
=
Eoi
T + RIn F
(YO/YR)
Ei
=
E
-
F
(In 1
+ (h/Kd)
(7)
The ionic strength variation of Ei follows readily from eq 6 using an extended Debye-Huckel relation for activity coefficients with the mean distance of closest approach of ions taken to be about 7 d
Ei = E’;
+ cI1”/(l + 2.3I”’)
(8)
Here the parameter c = 2.303ART/F and is seen to include the Debye-Huckel term A which is defined as A = 1.S25 x 106 ( p O / ~ 3 T 3 ) ”po* being , the density and a the dielectric constant of the solvent at temperature T’K. For water at 25.0°, A = 0.510, and since R T / F = 59.17 mV, it follows that c = 30.2 mV. Reaction between T~is(pyridine-d-aldoxi17ze)il*on(II) and Hexacyanofel.rate(III), This reaction, an electron exchange between two redox couples, is represented by the equilibrium The Journal of Physical Chemistry
log KO,
+ 6A11”/(1 + 2.91”’)
(11)
where the mean distance of closest approack in the Debye-Huckel expression is taken at nearly 9 A in this system. The term A was defined above. The mean enthalpy change for the reaction in eq 9 is obtained from the temperature variation of the calculated Ki values at low ionic strength, and as a first approximation this is taken to be equal to the thermodynamic quantity AH’. The above experimentally determined values of KO, and AH’ can now be compared with the derived values which are calculated from the thermodynamic quantities for the two redox couples involved. Thus AG”(eq 9) = AGO (hexacyanoferrate)
it follows from eq 5 that
R T --
log Ki
- AGO (eq 4)
(12)
- E’i(eq
(13)
from which it follows at 25.0” that 0.05916 log K o i = E’(hexacyan0ferrate)
4)
where the reduction potentials are in volts.
Experimental Section Reagents and Materials. Pyridine-2-aldoxime was purchased from Light & Co., Colnbrook, England, and recrystallized to constant melting point, 113”. All other chemicals were of AnalaR grade. The purity of the salt Fe(I\”4)2(S0J2 6Hz0 as determined by titration against Ki\lnOr was 99.9%, and that of Fe(n”4) (S04)~. 12Hz0 titrated against SnCL was 98.6Yc; in both cases allowance was made for this in the weighing 9
( 5 ) W, A. Eaton, P. George, and G. I. H. Hanania, J . Phys. Chem., 71, 2016 (1987).
1357
REDUCTION POTELVTIALS OF COMPLEX Tom of samples. K3Fe(CN)6 as used in the equilibrium experiments was taken from a freshly opened bottle without further purification. Buffer solutions were prepared from appropriate mixtures of 0.05 M S a O H and ?;aH$04, adjusted to the required ionic strength by dilution and/or addition of NaCl (vide infra). The iron(I1) complex of pyridine-2-aldoxime formed immediately when a 0.1 M solution of the latter was added to a solution of ferrous ammonium sulfate containing an equivalent amount of HC104, the ligand to metal ratio being 280. The stock solution so made, M , was diluted in the appropriate usually about buffer to give the required concentration. These solutions were fairly stable. For the iron(II1) complex, ligand was added to a solution of ferric ammonium sulfate containing 3 equiv of HClOJequiv of Fe, and 4 min was allowed for complete formation. S o buffer was added in order to give the maximum ligand to metal ratio of 480. Nevertheless, these solutions were not so stable. Consequently, fresh solutions of both complexes were made up for each experiment, as were all other solutions except buffers. Deionized glass redistilled water was used in making all the solutions. Measurement of Reduction Potential. Reduction potentials were measured in a thermostated two-compartment cell with an agar-saturated KC1 bridge. The reference half-cell was a selected commercial saturated calomel electrode constantly equilibrated in saturated KC1, while the other contained 14.0 ml of an equimolar mixture of the ferrous and ferric complexes, usually of concentration 2.50 X M each. One gold and one rhodium metal electrode dipped into this solution, and purified nitrogen was bubbled into the mixture to effect mixing. Potentials were read to 0.1 mV on a Radiometer P H l I 4 potentiometer. pH was then measured using the same solution and apparatus with a thermostated electrode assembly. In both cases, temperature was controlled to .t0.05” or better. NBS standards were used in calibrating the pH scale.6 Potentials were read at l-min intervals for about 10 min. A rise of 1-2 mV was usually observed in the first few minutes followed by a slow decay. Linear extrapolation from this region to zero time (involving no more than 2 mV) yielded emf values which were taken to be the redox potentials corresponding to the moment of mixing. The gold and rhodium electrodes gave concordant results and were used in preference to platinum which had a slower response and gave readings which were higher but not consistent. The reversibility of the system was tested using molar ratios of oxidant to reductant varying from 1: 4 to 4 : 1 at constant total ionic strength; the measured emf’s yielded a reduction potential constant within 0.5 mV. The effect of concentration was also tested over the range 1.0 X to 5.0 X M and gave reduction potentials constant within 0.5 mV.
For every experiment, at given ionic strength, pH, and temperature, two or three independent measurements of reduction potential were made. Reproducibility varied between 0.3 and 1.0 rnV depending on conditions. Determination of Equilibrium Constant. The equilibrium between the tris (pyridine-2-aldoxime)iron (I11-1I) and the hexacyanoferrate(II1-11) systems was investigated spectrophotometrically. It was found convenient to prepare the various equilibrium mixtures by adding the tris(pyridine-2-aldoxime)iron(II) complex to a solution of hexacyanoferrate(II1) in buffer under the required conditions. The reverse approach, starting with the iron(II1) complex and hexacyanoferrate(11), proved unsatisfactory possibly owing to the instability of the iron(II1) complex. Measurements were made a t 515 mp where the iron(I1) complex has a strong absorption maximum while the iron(II1) complex has weak absorption (at pH 7.0, E 10,000 and 2450 M-’ cm-‘, respectively) ; the Corresponding absorption by hexacyanoferrate(I1) is negligible, and for hexacyanoferrate(II1) e is 16 M-l cm-l. For every experiment, three of four mixtures were prepared containing 1.00 X M iron(I1) complex and concentrations of hexacyanoferrate(II1) varying M . Absorbancies from 1.00 X 10-4 to 3.00 X at 515 mp were read on a Unicam SP ,500 spectrophotometer with a thermostated cell compartment. It was found that the absorbancy inereased perceptibly with time (averaging about 0.005 unit/min) ; readings were therefore taken at rapid intervals up to about 3 min, and the value corresponding to the moment of mixing was obtained by extrapolation. The effect was smaller at lower ferricyanide concentrations and at lower ionic strengths. As a further check, the absorbancy of the solution containing only iron(I1) complex was measured independently, and that for the iron(II1) complex was obtained from a series of independent measurements on solutions containing progressively increasing molar ratios of ligand to iron. This latter value is pH independent, whereas the corresponding value for the iron(I1) complex varied with pH in accordance with the ionization in eq 3. The measured equilibrium constant K is defined in terms of the total molar concentrations of the species in eq 9. It can readily be shown that if AI is the absorbancy of a given equilibrium mixture, A z that of the iron(I1) complex, and Aa that of the iron(II1) complex, all under the same conditions, then
K = c(A2 - W / { ( A 1 [f(A2
- A3) x - A3) - c(Az - AI)]} (14)
where c is the initial molar concentration of the iron(I1) complex, and f is the initial ferricyanide concentration. (6) R. G . Bates, “Determination of pH,” John Wiley and Sons, Inc., Kew York, N. Y., 1964,p 123.
Volume 72, Number
4
April 1968
G. I. H. HANANIA, D. H. IRVINE, AND F. R. SHURAYH
1358 Equation 14 was used in calculating K . In theory, the only variation in K at constant temperature and ionic strength arises from the ionization in eq 3, and on this basis one calculates the pH-independent equilibrium constant Ki using eq 10. K was determined from the measured absorbancies corresponding to between 40 and 75% reaction; beyond this a tailing was observed, K values tending to rise with increasing ferricyanide concentration. This, together with the uncertainty in the extrapolation of absorbancies to zero time, leads to uncertainty in K estimated at8 no less than
~
f.570. Ionic Strength. For every experiment, the total molar ionic strength was computed from the contributions of the phosphate buffer, acid, and salts in the solution. In redox experiments, buffer ions contributed about SOY, of the total. Since stock solutions of the iron salts mere made in the presence of HC104, these solutions invariably caused a lowering of the pH of phosphate buffers, the effect being more pronounced at higher pH and at lower ionic strengths. Calculation of buffer contribution had to take this effect into account and was therefore based on the measured pH of the final reaction mixture. Other contributing ions include NH4+, SO?-, C104-, and Fe(Pa)-. Independent conductivity measurements also showed that under our experimental conditions enough KCl diffuses from the salt bridge to contribute about 0.001 M . Thus in a typical redox experiment: I = 0.42 (buffer) 0.007 (other ions) 0.001 (bridge) = 0.050 M . In the equilibrium experiments, where measurements were spectrophotometric, the solutions were more dilute than in redox (potentiometric) experiments. Buffer contributions were accordingly higher, over 90% of total, and pH effects were smaller. Salt ions include NH4+, S042-, c104-,and Fe(P,)-, as well as I5.
Mean Activity Coefficient of Polyelectrolytes.
VIII.
Osmotic and Activity
Coefficients of Polystyrenesulfonates of Various Gegenions' by Norio Jse and Tsuneo Okubo Department of Polymer Chemistry, Kyoto Unioersity, Kyoto, J a p a n Accepted and Transmitted by The Faraday Society
( J u l y 66, 1967)
The osmotic and mean activity coefficients of polystyrenesulfonates of various gegeiiioiis were investigated by means of the isopiestic vapor-pressure measurements. The order of the magitude of the activity coefficient was H + > Li+ > Na+ > K+, Ca2+ > Ba2+,and N+(n-CdHg) > N+(n-CsH,) > N+(CzH6)4 > N+(CH3)4> N+(CHs)&H2C6H6 > "2. This relative order was accounted for in terms of the structural effects of the ions on water. It was inferred that the polystyreiiesulfonate ion acted as a fairly strong structure former because of the benzene ring. The structure-formingtendency of the polysulfonate ion was suggesbed to be enhanced with increasing concentration, in contrast with the structural influences of simple ions which are so far considered concentration independent. As is well recognized, the mean activity coefficient of electrolytes is a most basic and most important quantity for understanding the thermodynamic properties of the solutions. (In the present work, the mean activity coefficient is discussed, which should not be confused with the single-ion activity coefficient.) For low molecular weight electrolytes, systematic study, experimental and theoretical, on this quantity has been conducted thoroughly, as far as dilute aqueous solutions of 1-1 type electrolytes are concerned. I n the field of high molecular weight electrolytes, however, there is very little information concerning the activity coefficient, the only published data being those measured in aqueous media in this Thus, it is
strongly hoped to extend the experimental work to other polyelectrolyte systems. I n the present paper, we report the results obtained for various salts of a polystyrenesulfonic acid (PSt) in the binary aqueous solutions. The main purpose of this work is to study (1) Part VII: H. Matsui, N . Ise, and T. Okubo, J. Phys. Chem., submitted for publication. (2) N . Ise and T. Okubo, ibid., 65, 4102 (1965). (3) N. Ise and T. Okubo, ibid., 70, 1930 (1966). (4) N. Ise and T. Okubo, ibid., 70, 2400 (1966). (6) N . Ise and T. Okubo, ibid., 71, 1287 (1967). (6) N. Ise and T. Okubo, ibdd., 7 1 , 1886 (1967). (7) T. Okubo, N. Ise, and F. Matsui, J . Am. Chem. floc., 89, 3687 (1967).
Volume 76, Number 4
A p d 1968
