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Mixed-ligand complexes of iron and hydroxy-1,10-phenanthrolines
Talanro, Vol. 21. pp. 1007 to 1012 6 Pergamon Press Ltd 1980. Printed in Grcal Britain
MIXED-LIGAND COMPLEXES OF IRON AND HYDROXY-l,lO-PHENANTHROLINES DONALD P. POE and ALAN D. EPPEN Department of Chemistry, University of Minnesota, Duluth, MN 55812, U.S.A.
(Received 17 August
1979.
Accepted
7 November
1979)
Summary-The half-wave potentials of all of the possible tris-complexes formed between iron(I1) and I,lO-phenanthroline (A). 4-hydroxy-l,lO-phenanthroline (B). and 4,7-dihydroxy-l.lO-phenanthroline (C) have been measured at a rotated platinum electrode. Values on the hydrogen scale are 1.06 V for A,Fe(II), 0.78 for A,B. 0.58 for A,C, 0.60 for AB,, 0.39 for B,, 0.39 for ABC, 0.21 for AC2. 0.22 for B&Y, 0.06 for BC2. and -0.10 for C3 at pH 11.0, 25”, and ionic strength 0.2. Half-wave potentials of the parent binary complexes are constant from pH 9 to 13 and are equal to the conditional reduction potentials. The mixed complexes are stable from pH 10 to 12, and form a redox indicator system with a continuous range of accessible potentials from -0.10 to + 1.06 V vs. NHE.
The introduction of tris( 1, IO-phenanthroline)iron(II) (ferroin) as a reversible, high-potential redox indicator in 1931’ was followed by intense study of the iron l,lO-phenanthroline complexes as redox indicators. Several reviews2-4 are available which discuss their oxidation-reduction chemistry. It is notable that almost all the potentials reported fall between 0.84 V 3,4,6,7_tetramethyl- I, IO-phenanthroline for the 1.26 V for the 5-nitro-l,lOcomplex 5 and phenanthroline complex.6 Two exceptions are 0.71 V for the 3-carbethoxy-4-hydroxy-l,lO-phenanthroline compound’ and -0.1 V for the 4,7-dihydroxy-l,lOphenanthroline complex.s-10 The last named phenanthroline has several properties which are unique in the class of l,l@phenanthrolines. It forms a stable coloured complex with iron in highly alkaline solutions, and has been used as a spectrophotometric reagent for iron in such media.’ ’ In addition the iron(II1) tris-chelate is stable indefinitely in mildly alkaline solutions,‘* a property possessed by no other iron(III)-phenanthroline complex at any pH. The stability of the iron(N) derivative and the extremely low reduction potential of the iron(II&II) chelate couple have resulted in the use of the iron(I1) derivative as a reagent in the spectrophotometric determination of dissolved oxygen’ and as a redox indicator in the titration of sodium dithionite with potassium ferricyanide.” The development of additional iron-phenanthroline complexes with potentials in the intermediate range from -0.1 to +0.8 V seems highly desirable. Although a large number of redox indicators with potentials in this range are already available, such as the indophenols and indigosulphonates,13*14 these compounds are characterized by multielectron redox reactions with varying degrees of irreversibility. In addition the potentials of the latter compounds vary with pH, and in most cases the oxidized form is the highly coloured species. In contrast the ironphenanthrolines undergo rapidly reversible single-
electron reactions, the potentials are independent of pH, and the reduced forms are the highly coloured species. The low reduction potential of the tris(4,7-dihydroxy-IJO-phenanthroline)iron(III,II) couple is apparently a direct result of the powerful electron-donating effect of the hydroxy groups, which leads to stabilization of the iron(III) derivative. Thus one would look to other hydroxy-1,lSphenanthrolines in the search for low-potential indicators. To date at least 39, substituted l,lO-phenanthrolines incorporating the hydroxy group have been synthesized;” only a few of these have been studied in detail, and in general their synthesis is rather difficult. On the other hand, available ligands can be used to prepare mixed-ligand which have properties intermediate complexes between those of the respective binary complexes. This approach has been used by Taylor and Schilt,” who found that the reduction potentials of the mixedligand complexes of iron and various substituted l,lO-phenthrolines were approximately equal to the weighted averages of the reduction potentials of the corresponding binary complexes. In that study the range of potentials involved was 0X7-1.28 V. In this paper we report potentials for all the possible tris complexes formed by mixing iron(I1) with l,lO-phenanthroline, Chydroxy-l,lO-phenanthroline and 4,7-dihydroxy-l,lO-phenanthroline, covering a range of potentials from -0.10 to 1.06 V. EXPERIMENTAL Apparatus
The instrumentation used for obtaining current-potential curves at a rotated platinum electrode is described in the accompanying paper. ’ ’ Reagents
4-Hydroxy-I .IO-phenanthroline and 4.7-dihydroxy-1 ,lOphenanthroline were prepared according to Snyder and Freier’s with minor modifications.” 4-Hydroxy-l,lO-phenanthroline was recrystallized three times from water.
1007
1008
DONALD
P. POE and
complexes
and
electrolysis
EPPEN
RESULTS
Electrode ptetreatment At the beginning of each day, the rotated platinum electrode was soaked in sulphuric acid-dichromate solution for 5 min and rinsed thoroughly with distilled water. Before each run, the potential was kept at the foot of the oxygen evolution wave for 2 min, then at the foot of the hydrogen wave for 2 min, then at +0.3 V relative to the hydrogen wave for 2 min. according to a procedure suggested by Adams.” mixed
D.
general procedure are described in the accompanying paper. ’ 7 Cyclic voltammograms were obtained at pH 11.O and 2.5”, in ammonia-ammonium chloride buffer (ionic strength = 0.02) and dilute sodium hydroxide (ionic strength = 0.1) media, in presence and absence of I.0 x IO-‘M tris(4-hydroxy-I,lO-phenanthroline)iron(II).
yielding the anhydrous reagent (99.9% pure by potentiometric titration). 4,7-Dihydroxy-l,IO-phenanthroline was recrystallized twice from 6M hydrochloric acid, yielding a crystalline compound corresponding to the formula ClzHsN,02.0.31 HCI. Solutions of these two phenanthrolines were prepared by adding IOM sodium hydroxide dropwise to an aqueous suspension of the crystals until dissolution was complete. Solutions of I.lO-phenanthroline were prepared by adding the minimum required amount of concentrated sulphuric or hydrochloric acid dropwise to an aqueous suspension of the solid. Solutions of iron(H) were prepared by dissolving FeSO,. 7H,O (99.9% pure by titration with standard dichromate) in deaerated I M sulphuric acid. Ammonia buffers were prepared by adding SM ammonium chloride to 5M ammonia to give pH IO.
Preparation of the
ALAN
pro-
cedure
Solutions of the binary complexes were prepared, containing 5.0 x 10m5mole of iron and 5.0 x 10m4 mole of ligand in a total volume of 50 ml. Solutions of the mixed complexes contained 5.0 x 10e5 mole of iron(I1) and 2.5 x 10m4 mole of total ligand in 50 ml. The typical procedure is illustrated here for bis(l,lO-phenanthrolinet(4hydroxy-l,IO-phenanthroline)iron(II). IJO-Phenanthroline monohydrate (0.0335 g, 1.69 x 10e4 mole) was dissolved in a few ml of dilute hydrochloric acid and transferred to the electrolysis cell. Then 0.0162 g (0.83 x IO-” mole) of 4-hydroxy-l.lO-phenanthroline was added directly to the electrolysis cell, followed by IO ml of 5M ammonia buffer and enough water to give a total volume of 50 ml. The pH was adjusted to 11.0 by dropwise addition of IOM sodium hydroxide, the solution purged with nitrogen for 20 min. a final pH-adjustment made, and a current-potential curve obtained. Next 0.50 ml of O.lOOM FeSO, was added, the solution was stirred and left for 40 min to come to equilibrium, a final pH-adjustment was made, and a currentpotential curve for the mixed complexes was obtained. Potentials were scanned from 0.6 or 0.7 V us. SCE to that for the hydrogen-evolution wave and back again. Scans initiated at more positive potentials showed almost complete suppression of anodic currents. The ionic strength of the solutions was approximately 0.2. Reversibility studies The reversibility of the oxidation of tris(Chydroxy-l,lOphenanthroline)iron(II) was studied bv cyclic voltammetry at a stationary pla&um electrode. -Thk equipment anh
AND
In this discussion the following symbols will be used: A = l,lO-phenanthroline, B = 4-hydroxy-l,lOphenanthroline, C = 4,7-dihydroxy-l,lO-phenanthroline, a, b, c = stoichiometric coefficients for A, B, C respectively. In the formation of a stable mixed-ligand complex A,B&,Fe(II), the sum a + b + c is equal to 3, where a, b, and c can take integral values from 0 to 3. The formation of A,Fe(II) is normally considered complete over the pH-range 2-9. In the presence of sodium dithionite, or in the absence of oxygen, formation of this complex as determined from its absorbance at 510 nm is quantitative up to pH 1I, with only a slight decrease in absorbance occurring at pH 12 and 13. Previous investigations have shown that the optimum conditions for formation of B,Fe(II) and &Fe(II) are from pH 11 to 2M sodium hydroxide” and from pH 9 to concentrated sodium hydroxide solution,’ ’ respectively. Spectrophotometric measurements of solutions of the mixed-ligand complexes indicated that maximum stability of the complexes is obtained from pH 10 to 12. Therefore pH 11 was selected as the optimum pH for study of the mixed complexes. This pH represents the upper limit of stability of tris( 1,lO-phenanthroline)iron(II) in the presence of excess of phenanthroline. In order to ensure complete formation of ‘the mixed complexes in solutions containing l,lO-phenanthroline, a 5: 1 molar ratio of ligands to iron was used. Half-wave potentials of the binary complexes
The half-wave potential of the A,Fe(III,II) couple is 1.06 V (all potentials will be quoted us. NHE) from pH 0 to 11 at ionic strength = 1, and is equal to the conditional reduction potential, as demonstrated in the accompanying paper.” The reduction potential of the B,Fe(III,II) couple was recently reported as 0.39 V, with no supporting data.” The E ,,,-values obtained in this work at pH 9, 11, 13 are all 0.39 V. A potentiometric measurement of the reduction potential was not attempted because
Table 1. Cyclic voltammetry of tris(4-hydroxy-l,lO-phenanthroline)iron(II) Electrolyte NH,, NH&I NH,, NH&I NaOH NaOH
DISCUSSION
Ionic strength
Scan-rate, mVfsec
AE,, V*
E ,,2, V vs. SCEt
0.02 0.02 0.1 0.1
5 100 5 50
0.060 0.080 0.067 0.080
0.135 0.130 0.151 0.140
* AE, = difference between anodic and cathodic peak potentials. was calculated as the mid-point between the peak potentials. anodic peak current, i, = cathodic peak current.
at pH 11
i&3 0.83 0.74 0.92 1.19
1009
Mixed-ligand complexes of iron and hydroxy-l,IO-phenanthrolines
1
a
0.40
,
I
I
0.00
-0.40
I
-0.80
E va SCE, volts Fig. 1. Linear-scan vohammetry of BC,Fe(II) system in pH 11 ammonia buffer at a rotated wire electrode: 1.00 x IO-‘M total iron added, scan-rate 4 mV/sec. a stable iron(W) tris-chelate is not formed under these
conditions. In order to relate El,,-values to the reduction potential, cyclic voltammetric studies at a stationary platinum electrode were performed. A summary of the results appears in Table 1. Background currents due to surface reactions of the platinum electrode were approximately equal to the faradaic currents for oxidation and reduction of the complex, making measurements of peak potentials and currents subject to considerable uncertainty. In the ammonia buffer system the peak separation indicates that the electrode reaction is reversible at a scan rate of 5 mV/sec. The calculated E,,, is therefore equal to the conditional reduction potential, 0.38 V, in good agreement with E,,r = 0.39 V determined by linear scan voltammetry at the rotated platinum electrode. The reduction potential of the C,Fe(III,II) couple has been reported three times previously as -0.13,
platinum
- 0.11, and - 0.06 V.s-’ ’ The ionic strength was not given in any of these reports. The E,,,-values found at ionic strength = 0.16 in this work were -0.10 V from pH 9 to 13. Potentiometric titration of a solution of sodium dithionite and C,Fe(II) at pH 10.6, 25” and ionic strength = 1.0 gave E” = -0.11 V. Because of the good agreement with potentiometric and cyclic voltammetry results we conclude that the E,,,-values found are equal to the conditional reduction potentials of the binary complexes, and that the reduction potentials do not vary over the pH range 9-13. Presumably the E,,, -values reported below for the mixed complexes are also equal to the conditional reduction potentials. Half-wave potentials of the mixed complexes
Current-potential curves for solutions of BC,Fe(II) and for ABCFe(I1) are shown in Figs. 1 and 2 as
I
E vs
SCE. volts
Fig. 2. Linear-scan voitammetry of ABCFe(II) system in pH I1 ammonia buffer at a rotated platinum wire electrode: 1.00 x IO-‘&f total iron added, scan-rate 4 mV/sec.
1010
DONALD P. POE and ALAN D. EPPEN Table 2. Half-wave
. Molar A
ratio L:Fe B C
E ,,*, Vvs. NHE calculated observed
3 2
0 I
0 0
2
1 0 0
0.67 0.6 1
0
0 2 3
I
I
I
0.45
I 0 0 0
0 2 1 0
2 I 2 3
0.29 0.23 0.06
I
potentials
0.84
I .06 0.78
of the mixed-ligand
AE,,?,
V*‘
complexes
at pH
I1
Species observedtg
- 0.06
1.0 A, 1.0 AIB: 0.5 AB,: 0.0 AJ
0.58 0.60 0.39 0.39
- 0.09 -0.01
1.0 A&: 0.2 AC,: 0.2 C,(III) 1.0 AB2: 0.7 A,B: 0.1 B,
0.21 0.22 0.06 -0.10
- 0.08 -0.01 0.00
I.0 8, -0.06
1.0 ABC: 0.7 (A,C + AB,): 0.6 (AC2 + B,C): 0.2 A2B: 0.1 C,(W) 1.0 AC,: 2.6 A$: 0.1 AC2(III) 1.0 B,C: 0.5 B,: 0.3 BC, 1.0 BCZ: 0.8 B,C: 0.2 C,: 0.3 CJII) I .o c,: 0.2 C,(U)
*AE I:Z = Eobrcrved - Eca,cu,n,cd. t Relative magnitude of the diffusion current for each observed species is given, normalized so that the value for the desired species equals I .O. $ All species indicated are the iron(I1) derivatives. except for those followed by a romal numeral III, used to indicate the iron(II1) derivatives. $ It is likely that some A3Fe(Il) was present, but could not be observed (see text).
examples. The curves for each mixed-Egand system showed two or more separate waves. The assignments in Table 2 were made by matching observed halfwave potentials with weighted averages calculated from the E,,,-values of the binary complexes by means of the equation : E i.ca1c =
a(1.06) + b(0.39)
+ c( -0.10)
3
In making the assignments the magnitudes of the diffusion currents and the probable equilibrium concentrations of individual species were also considered. In all except the AC2 system the largest diffusion current corresponded to the complex in which the ligand-toligand ratio matched the ratio in which the ligands were added to the system. Appearance of the waves at the predicted potentials in several different systems lends validity to the assignments. Some significant differences between the observed and calculated half-wave potentials, listed as AE,,, in Table 2, are apparent. In every case the observed value is equal to or less than the calculated value, in contrast to the findings of Taylor and Schilt.16 They determined the reduction potentials of I2 mixedligand iron-phenanthroline complexes, incorporating the methyl, chloro. phenyl. sulpho and unsubstituted derivatives of 1.10-phenanthroline. and found that the reduction potentials agreed with the weighted averages within 0.01 V, except when 5-nitro-l.lO-phenanthroline was present as a ligand. To explain this observation, the calculated and observed potentials for ternary complexes in which at least one molecule of l.lO-phenanthroline is present are summarized in Table 3. The calculated values agree with the experimental results within 0.01 V for all complexes involving ligands which form binary complexes with reduction potentials within 0.19 V of that of the tris(l,lOphenanthroiine)iron(III, II) couple. Only when the 5-nitro. 4-hydroxy, or 4,7-dihydroxy substituted phenanthrolines are present do the observed values differ
significantly from the calculated values. These three have substituents which are either strongly electronaccepting or electron-donating, and form binary complexes with reduction potentials which differ by more than 0.20 V from 1.06 V. Thus it appears that the weighted average provides a good prediction of the reduction potential of a mixed-ligand iron-phenanthroiine complex when the ligands are similar in electronic character. Such agreement is not expected, however, when two ligands are present which have substantially different electronic properties. The relative magnitudes of the diffusion currents for the individual mixed complexes are listed in Table 2 as an indication of the equilibrium concentrations. Conversion of these values into concentrations is hindered by the inability to prepare a standard solution of a single mixed complex. Furthermore, there is no assurance that true equilibrium conditions were obtained, especially for systems in which ligand C was present. Trace amounts of oxygen leaking into the system would immediately be reduced by C,Fe(II), causing a continuous disturbance of the equilibrium. The detection of C,Fe(III) in some systems demonstrates this to be the case. Other species, particularly BC,Fe(II), B2CFe(II), and AC,Fe(II) might be oxidized if &Fe(H) is absent. Several cases in Table 2 deserve special comment. In the A,B system, some A,Fe(II) was probably present, but was not observed because of the onset of oxidation of water. In the ABC system five separate waves were observed. shown in Fig. 2. Assignments were made on the basis of E,,z-values observed in simpler systems, but it was impossible to distinguish between A,CFe(II) and AB,Fe(II), or between AC,Fe(II) and B,CFe(II). It was also assumed that negligible amounts of the binary iron(I1) derivatives would be formed in the ABC system, and the wave at 0.39 V was regarded as due entirely to ABCFe(I1). even though the potential corresponds exactly to E,,, of B,Fe(II).
Mixed-ligand Table
3. Calculated
Complex*
complexes and
1011
of iron and hydroxy-l,IO-phenanthrolines
observed reduction potentials phenanthroline complexes Binary complexest EL - 1.06 EL
(V) of ternary
Ternary
iron-
complexes$
E,,,,
E Obr
E ohs -
L,,
;;;4jWSWIA,
1.28 1.28 I .23 I.16 1.16 I .08 1.09
0.22 0.22 0.17 0.10 0.10 0.02 0.03
1.21 I.13 I.17 I.13 I 09 I .07
I .24
(~-NDz)AI (5-SO&A (5-Cl)2A (5-Cl)A2
1.09 1.16 I.13 I.10 I .07
+ 0.03 - 0.04 -0.01 0.00 +0.01 0.00
(5-CH,jA, (5-CH&A (4,7-diCHS)Az $di$,)zA
I.01 I.01 0.87 0.39 0.87
-0.05 - 0.05 -0.19 -0.67 -0.19
1.04 1.02 1.00 0.93 0.84
1.04 I .02 0.99 0.78 0.93
0.00 0.00 -0.01 -0.06 0.00
0.39 -0.10 -0.10
-0.67 - 1.16 - I.16
0.61 0.67 0.29
0.60 0.58 0.21
-0.01 -0.09 -0.08
(5-N%@
(4-OH),; (4.7-hOH)A* (4.7-diOH)2A
* The iron(III.11) complex couples are designated by the ligands only, with the following abbreviations: A (I,lO-phenanthroline), 5-NO2 (5-nitro-l,lO-phenanthroline), 5-Cl (5-chloro-l.I@phenanthroline), 5-SO3 (I,lO-phenanthroline-5sulphonic acid), 4,7-di(q%SO,) (4,7-diphenylsulphonato-l,lO-phenanthroline), 5-4 S-CHJ (5-methyl-l,l@phenanthroline), 4,7(5-phenyl-l,lO-phenanthroline), diCH, (4,7-dimethyl-l,lO-phenanthroline, 4-OH (4-hydroxy-l,lO-phenanthroline), 4,7-diOH (4,7-dihydroxy-l,lO-phenanthroline). t EL is the reduction potential of the binary complex with the substituted ligand. potential of the mixed complex predicted from the 5 EC,,, is the reduction weighted average. Eobr is the observed reduction potential of the mixed complex. Values for 4-OH and 4,7-diOH are from this work; others are taken from reference 16.
REFERENCES
CONCLUSION
The possibilities for application of the mixed complexes as redox indicators are limited by the fact that several mixed complexes will always exist in equilibrium with the desired complex. Perhaps the ABCFe(I1) system, in which at least six individual mixed complexes exist in equilibrium, with a continuous range of accessible potentials from 0.2 to 0.8 V, will lead to some novel applications. A solution of the ABCFe(I1) system could serve as a wide-range redox indicator system. in which absorbance is a measure of the reduction potential of a solution, much the same as the colour of an acid-base indicator shows the pH. Another possibility is the use of the mixed complexes as electrode mediators in coulometric titrations. Finally, it should be possible to extend the upper potential limit of this system by incorporating another ligand such as .5-nitro-IJO-phenanthroline in the mixed-ligand complex. The reduction potential of tris(4-hydroxy-l,lO-phenanthroline)iron(II) has not been previously reporied except as a passing note in a recent paper.” Its reduction potential of 0.39 V is significantly different from any previously reported for the binary iron-phenanthroline complexes, and it may find use as a redox indicator as well as a reagent in the spectrophotometric determination of traces of oxidants and reductants.
G. H. Walden
Jr., L. P. Hammett and R. P. Chapman, J. Am. Chem. Sot., 1931,53, 3908. 2. G. F. Smith and F. P. Richter, Phenanthroline and Substituted Phenanthroline fndicarors, G. F. Smith Chemical Co., Columbus, Ohio, 1944. 3. W. W. Brandt, F. P. Dwyer and E. C. Gyarfas, Chem. Reo., 1954, 54, 959. 4. A. A. Schilt, Analytical Applications of I,lO-Phenonthroline and Related Compounds. Pergamon, New York. 1969. W. W. Brandt and G. F. Smith, Anal. Chew 1949. 21. 1313. G. F. Smith and F. P. Richter, Ind. Eng. Chem.. Anal. Ed., 1944, 16. 580. M. N. Hale and M. G. Mellon, J. Am. Chem. Sot.. 1950. 72. 3217. P. George, G. I. H. Hanania and W. A. Eaten, unpublished work. in Hemes and Hemoproteins, B. Chance. R. W. Estrabrook, and T. Yonetani. eds., p. 269. Academic Press, New York, 1966. 9 D. P. Poe and H. Diehl, Talanta. 1974, 21. 1065. 10. Idem, ibid., 1976, 23, 147. Anal. I I. A. A. Schilt, G. F. Smith and A. Hiimbuch, Chem., 1956, 28, 809.
12. D. P. Poe and H. Diehl, Talanta. 1976, 23, 141. 13. W. M. Clark, J. Appl. Physics, 1938, 9, 97. 14. I. M. Kolthoff and V. A. Stenaer. Volumefric Analysis. Vol. I, 2nd Ed., pp. 105-141. interscience, New York. 1942. Iowa State University. 15. D. P. Poe. Ph.D. Dissertation, 1974. 16. P. J. Taylor and A. A. Schilt, Inorg. Chim. Acra. 1971. 5, 691.
1012
DONALD
P. POE and ALAN D. EPPEN
17. D. P. Poe and J. A. Adamczak, Talanta, 1980, 27, 1025. 18. H. R. Snyder and H. E. Freier, J. Am. Chem. SM.. 1946. 68, 1320.
19. D. P. Poe, A. D. Eppen and S. B. Whoolery, Talanta, 1980, 27, 368. 20. R. N. Adams, Electrochemistry at So/id Electrodes. Dekker, New York, 1969.
MIXED-LIGAND COMPLEXES OF IRON AND HYDROXY-l,lO-PHENANTHROLINES DONALD P. POE and ALAN D. EPPEN Department of Chemistry, University of Minnesota, Duluth, MN 55812, U.S.A.
(Received 17 August
1979.
Accepted
7 November
1979)
Summary-The half-wave potentials of all of the possible tris-complexes formed between iron(I1) and I,lO-phenanthroline (A). 4-hydroxy-l,lO-phenanthroline (B). and 4,7-dihydroxy-l.lO-phenanthroline (C) have been measured at a rotated platinum electrode. Values on the hydrogen scale are 1.06 V for A,Fe(II), 0.78 for A,B. 0.58 for A,C, 0.60 for AB,, 0.39 for B,, 0.39 for ABC, 0.21 for AC2. 0.22 for B&Y, 0.06 for BC2. and -0.10 for C3 at pH 11.0, 25”, and ionic strength 0.2. Half-wave potentials of the parent binary complexes are constant from pH 9 to 13 and are equal to the conditional reduction potentials. The mixed complexes are stable from pH 10 to 12, and form a redox indicator system with a continuous range of accessible potentials from -0.10 to + 1.06 V vs. NHE.
The introduction of tris( 1, IO-phenanthroline)iron(II) (ferroin) as a reversible, high-potential redox indicator in 1931’ was followed by intense study of the iron l,lO-phenanthroline complexes as redox indicators. Several reviews2-4 are available which discuss their oxidation-reduction chemistry. It is notable that almost all the potentials reported fall between 0.84 V 3,4,6,7_tetramethyl- I, IO-phenanthroline for the 1.26 V for the 5-nitro-l,lOcomplex 5 and phenanthroline complex.6 Two exceptions are 0.71 V for the 3-carbethoxy-4-hydroxy-l,lO-phenanthroline compound’ and -0.1 V for the 4,7-dihydroxy-l,lOphenanthroline complex.s-10 The last named phenanthroline has several properties which are unique in the class of l,l@phenanthrolines. It forms a stable coloured complex with iron in highly alkaline solutions, and has been used as a spectrophotometric reagent for iron in such media.’ ’ In addition the iron(II1) tris-chelate is stable indefinitely in mildly alkaline solutions,‘* a property possessed by no other iron(III)-phenanthroline complex at any pH. The stability of the iron(N) derivative and the extremely low reduction potential of the iron(II&II) chelate couple have resulted in the use of the iron(I1) derivative as a reagent in the spectrophotometric determination of dissolved oxygen’ and as a redox indicator in the titration of sodium dithionite with potassium ferricyanide.” The development of additional iron-phenanthroline complexes with potentials in the intermediate range from -0.1 to +0.8 V seems highly desirable. Although a large number of redox indicators with potentials in this range are already available, such as the indophenols and indigosulphonates,13*14 these compounds are characterized by multielectron redox reactions with varying degrees of irreversibility. In addition the potentials of the latter compounds vary with pH, and in most cases the oxidized form is the highly coloured species. In contrast the ironphenanthrolines undergo rapidly reversible single-
electron reactions, the potentials are independent of pH, and the reduced forms are the highly coloured species. The low reduction potential of the tris(4,7-dihydroxy-IJO-phenanthroline)iron(III,II) couple is apparently a direct result of the powerful electron-donating effect of the hydroxy groups, which leads to stabilization of the iron(III) derivative. Thus one would look to other hydroxy-1,lSphenanthrolines in the search for low-potential indicators. To date at least 39, substituted l,lO-phenanthrolines incorporating the hydroxy group have been synthesized;” only a few of these have been studied in detail, and in general their synthesis is rather difficult. On the other hand, available ligands can be used to prepare mixed-ligand which have properties intermediate complexes between those of the respective binary complexes. This approach has been used by Taylor and Schilt,” who found that the reduction potentials of the mixedligand complexes of iron and various substituted l,lO-phenthrolines were approximately equal to the weighted averages of the reduction potentials of the corresponding binary complexes. In that study the range of potentials involved was 0X7-1.28 V. In this paper we report potentials for all the possible tris complexes formed by mixing iron(I1) with l,lO-phenanthroline, Chydroxy-l,lO-phenanthroline and 4,7-dihydroxy-l,lO-phenanthroline, covering a range of potentials from -0.10 to 1.06 V. EXPERIMENTAL Apparatus
The instrumentation used for obtaining current-potential curves at a rotated platinum electrode is described in the accompanying paper. ’ ’ Reagents
4-Hydroxy-I .IO-phenanthroline and 4.7-dihydroxy-1 ,lOphenanthroline were prepared according to Snyder and Freier’s with minor modifications.” 4-Hydroxy-l,lO-phenanthroline was recrystallized three times from water.
1007
1008
DONALD
P. POE and
complexes
and
electrolysis
EPPEN
RESULTS
Electrode ptetreatment At the beginning of each day, the rotated platinum electrode was soaked in sulphuric acid-dichromate solution for 5 min and rinsed thoroughly with distilled water. Before each run, the potential was kept at the foot of the oxygen evolution wave for 2 min, then at the foot of the hydrogen wave for 2 min, then at +0.3 V relative to the hydrogen wave for 2 min. according to a procedure suggested by Adams.” mixed
D.
general procedure are described in the accompanying paper. ’ 7 Cyclic voltammograms were obtained at pH 11.O and 2.5”, in ammonia-ammonium chloride buffer (ionic strength = 0.02) and dilute sodium hydroxide (ionic strength = 0.1) media, in presence and absence of I.0 x IO-‘M tris(4-hydroxy-I,lO-phenanthroline)iron(II).
yielding the anhydrous reagent (99.9% pure by potentiometric titration). 4,7-Dihydroxy-l,IO-phenanthroline was recrystallized twice from 6M hydrochloric acid, yielding a crystalline compound corresponding to the formula ClzHsN,02.0.31 HCI. Solutions of these two phenanthrolines were prepared by adding IOM sodium hydroxide dropwise to an aqueous suspension of the crystals until dissolution was complete. Solutions of I.lO-phenanthroline were prepared by adding the minimum required amount of concentrated sulphuric or hydrochloric acid dropwise to an aqueous suspension of the solid. Solutions of iron(H) were prepared by dissolving FeSO,. 7H,O (99.9% pure by titration with standard dichromate) in deaerated I M sulphuric acid. Ammonia buffers were prepared by adding SM ammonium chloride to 5M ammonia to give pH IO.
Preparation of the
ALAN
pro-
cedure
Solutions of the binary complexes were prepared, containing 5.0 x 10m5mole of iron and 5.0 x 10m4 mole of ligand in a total volume of 50 ml. Solutions of the mixed complexes contained 5.0 x 10e5 mole of iron(I1) and 2.5 x 10m4 mole of total ligand in 50 ml. The typical procedure is illustrated here for bis(l,lO-phenanthrolinet(4hydroxy-l,IO-phenanthroline)iron(II). IJO-Phenanthroline monohydrate (0.0335 g, 1.69 x 10e4 mole) was dissolved in a few ml of dilute hydrochloric acid and transferred to the electrolysis cell. Then 0.0162 g (0.83 x IO-” mole) of 4-hydroxy-l.lO-phenanthroline was added directly to the electrolysis cell, followed by IO ml of 5M ammonia buffer and enough water to give a total volume of 50 ml. The pH was adjusted to 11.0 by dropwise addition of IOM sodium hydroxide, the solution purged with nitrogen for 20 min. a final pH-adjustment made, and a current-potential curve obtained. Next 0.50 ml of O.lOOM FeSO, was added, the solution was stirred and left for 40 min to come to equilibrium, a final pH-adjustment was made, and a currentpotential curve for the mixed complexes was obtained. Potentials were scanned from 0.6 or 0.7 V us. SCE to that for the hydrogen-evolution wave and back again. Scans initiated at more positive potentials showed almost complete suppression of anodic currents. The ionic strength of the solutions was approximately 0.2. Reversibility studies The reversibility of the oxidation of tris(Chydroxy-l,lOphenanthroline)iron(II) was studied bv cyclic voltammetry at a stationary pla&um electrode. -Thk equipment anh
AND
In this discussion the following symbols will be used: A = l,lO-phenanthroline, B = 4-hydroxy-l,lOphenanthroline, C = 4,7-dihydroxy-l,lO-phenanthroline, a, b, c = stoichiometric coefficients for A, B, C respectively. In the formation of a stable mixed-ligand complex A,B&,Fe(II), the sum a + b + c is equal to 3, where a, b, and c can take integral values from 0 to 3. The formation of A,Fe(II) is normally considered complete over the pH-range 2-9. In the presence of sodium dithionite, or in the absence of oxygen, formation of this complex as determined from its absorbance at 510 nm is quantitative up to pH 1I, with only a slight decrease in absorbance occurring at pH 12 and 13. Previous investigations have shown that the optimum conditions for formation of B,Fe(II) and &Fe(II) are from pH 11 to 2M sodium hydroxide” and from pH 9 to concentrated sodium hydroxide solution,’ ’ respectively. Spectrophotometric measurements of solutions of the mixed-ligand complexes indicated that maximum stability of the complexes is obtained from pH 10 to 12. Therefore pH 11 was selected as the optimum pH for study of the mixed complexes. This pH represents the upper limit of stability of tris( 1,lO-phenanthroline)iron(II) in the presence of excess of phenanthroline. In order to ensure complete formation of ‘the mixed complexes in solutions containing l,lO-phenanthroline, a 5: 1 molar ratio of ligands to iron was used. Half-wave potentials of the binary complexes
The half-wave potential of the A,Fe(III,II) couple is 1.06 V (all potentials will be quoted us. NHE) from pH 0 to 11 at ionic strength = 1, and is equal to the conditional reduction potential, as demonstrated in the accompanying paper.” The reduction potential of the B,Fe(III,II) couple was recently reported as 0.39 V, with no supporting data.” The E ,,,-values obtained in this work at pH 9, 11, 13 are all 0.39 V. A potentiometric measurement of the reduction potential was not attempted because
Table 1. Cyclic voltammetry of tris(4-hydroxy-l,lO-phenanthroline)iron(II) Electrolyte NH,, NH&I NH,, NH&I NaOH NaOH
DISCUSSION
Ionic strength
Scan-rate, mVfsec
AE,, V*
E ,,2, V vs. SCEt
0.02 0.02 0.1 0.1
5 100 5 50
0.060 0.080 0.067 0.080
0.135 0.130 0.151 0.140
* AE, = difference between anodic and cathodic peak potentials. was calculated as the mid-point between the peak potentials. anodic peak current, i, = cathodic peak current.
at pH 11
i&3 0.83 0.74 0.92 1.19
1009
Mixed-ligand complexes of iron and hydroxy-l,IO-phenanthrolines
1
a
0.40
,
I
I
0.00
-0.40
I
-0.80
E va SCE, volts Fig. 1. Linear-scan vohammetry of BC,Fe(II) system in pH 11 ammonia buffer at a rotated wire electrode: 1.00 x IO-‘M total iron added, scan-rate 4 mV/sec. a stable iron(W) tris-chelate is not formed under these
conditions. In order to relate El,,-values to the reduction potential, cyclic voltammetric studies at a stationary platinum electrode were performed. A summary of the results appears in Table 1. Background currents due to surface reactions of the platinum electrode were approximately equal to the faradaic currents for oxidation and reduction of the complex, making measurements of peak potentials and currents subject to considerable uncertainty. In the ammonia buffer system the peak separation indicates that the electrode reaction is reversible at a scan rate of 5 mV/sec. The calculated E,,, is therefore equal to the conditional reduction potential, 0.38 V, in good agreement with E,,r = 0.39 V determined by linear scan voltammetry at the rotated platinum electrode. The reduction potential of the C,Fe(III,II) couple has been reported three times previously as -0.13,
platinum
- 0.11, and - 0.06 V.s-’ ’ The ionic strength was not given in any of these reports. The E,,,-values found at ionic strength = 0.16 in this work were -0.10 V from pH 9 to 13. Potentiometric titration of a solution of sodium dithionite and C,Fe(II) at pH 10.6, 25” and ionic strength = 1.0 gave E” = -0.11 V. Because of the good agreement with potentiometric and cyclic voltammetry results we conclude that the E,,,-values found are equal to the conditional reduction potentials of the binary complexes, and that the reduction potentials do not vary over the pH range 9-13. Presumably the E,,, -values reported below for the mixed complexes are also equal to the conditional reduction potentials. Half-wave potentials of the mixed complexes
Current-potential curves for solutions of BC,Fe(II) and for ABCFe(I1) are shown in Figs. 1 and 2 as
I
E vs
SCE. volts
Fig. 2. Linear-scan voitammetry of ABCFe(II) system in pH I1 ammonia buffer at a rotated platinum wire electrode: 1.00 x IO-‘&f total iron added, scan-rate 4 mV/sec.
1010
DONALD P. POE and ALAN D. EPPEN Table 2. Half-wave
. Molar A
ratio L:Fe B C
E ,,*, Vvs. NHE calculated observed
3 2
0 I
0 0
2
1 0 0
0.67 0.6 1
0
0 2 3
I
I
I
0.45
I 0 0 0
0 2 1 0
2 I 2 3
0.29 0.23 0.06
I
potentials
0.84
I .06 0.78
of the mixed-ligand
AE,,?,
V*‘
complexes
at pH
I1
Species observedtg
- 0.06
1.0 A, 1.0 AIB: 0.5 AB,: 0.0 AJ
0.58 0.60 0.39 0.39
- 0.09 -0.01
1.0 A&: 0.2 AC,: 0.2 C,(III) 1.0 AB2: 0.7 A,B: 0.1 B,
0.21 0.22 0.06 -0.10
- 0.08 -0.01 0.00
I.0 8, -0.06
1.0 ABC: 0.7 (A,C + AB,): 0.6 (AC2 + B,C): 0.2 A2B: 0.1 C,(W) 1.0 AC,: 2.6 A$: 0.1 AC2(III) 1.0 B,C: 0.5 B,: 0.3 BC, 1.0 BCZ: 0.8 B,C: 0.2 C,: 0.3 CJII) I .o c,: 0.2 C,(U)
*AE I:Z = Eobrcrved - Eca,cu,n,cd. t Relative magnitude of the diffusion current for each observed species is given, normalized so that the value for the desired species equals I .O. $ All species indicated are the iron(I1) derivatives. except for those followed by a romal numeral III, used to indicate the iron(II1) derivatives. $ It is likely that some A3Fe(Il) was present, but could not be observed (see text).
examples. The curves for each mixed-Egand system showed two or more separate waves. The assignments in Table 2 were made by matching observed halfwave potentials with weighted averages calculated from the E,,,-values of the binary complexes by means of the equation : E i.ca1c =
a(1.06) + b(0.39)
+ c( -0.10)
3
In making the assignments the magnitudes of the diffusion currents and the probable equilibrium concentrations of individual species were also considered. In all except the AC2 system the largest diffusion current corresponded to the complex in which the ligand-toligand ratio matched the ratio in which the ligands were added to the system. Appearance of the waves at the predicted potentials in several different systems lends validity to the assignments. Some significant differences between the observed and calculated half-wave potentials, listed as AE,,, in Table 2, are apparent. In every case the observed value is equal to or less than the calculated value, in contrast to the findings of Taylor and Schilt.16 They determined the reduction potentials of I2 mixedligand iron-phenanthroline complexes, incorporating the methyl, chloro. phenyl. sulpho and unsubstituted derivatives of 1.10-phenanthroline. and found that the reduction potentials agreed with the weighted averages within 0.01 V, except when 5-nitro-l.lO-phenanthroline was present as a ligand. To explain this observation, the calculated and observed potentials for ternary complexes in which at least one molecule of l.lO-phenanthroline is present are summarized in Table 3. The calculated values agree with the experimental results within 0.01 V for all complexes involving ligands which form binary complexes with reduction potentials within 0.19 V of that of the tris(l,lOphenanthroiine)iron(III, II) couple. Only when the 5-nitro. 4-hydroxy, or 4,7-dihydroxy substituted phenanthrolines are present do the observed values differ
significantly from the calculated values. These three have substituents which are either strongly electronaccepting or electron-donating, and form binary complexes with reduction potentials which differ by more than 0.20 V from 1.06 V. Thus it appears that the weighted average provides a good prediction of the reduction potential of a mixed-ligand iron-phenanthroiine complex when the ligands are similar in electronic character. Such agreement is not expected, however, when two ligands are present which have substantially different electronic properties. The relative magnitudes of the diffusion currents for the individual mixed complexes are listed in Table 2 as an indication of the equilibrium concentrations. Conversion of these values into concentrations is hindered by the inability to prepare a standard solution of a single mixed complex. Furthermore, there is no assurance that true equilibrium conditions were obtained, especially for systems in which ligand C was present. Trace amounts of oxygen leaking into the system would immediately be reduced by C,Fe(II), causing a continuous disturbance of the equilibrium. The detection of C,Fe(III) in some systems demonstrates this to be the case. Other species, particularly BC,Fe(II), B2CFe(II), and AC,Fe(II) might be oxidized if &Fe(H) is absent. Several cases in Table 2 deserve special comment. In the A,B system, some A,Fe(II) was probably present, but was not observed because of the onset of oxidation of water. In the ABC system five separate waves were observed. shown in Fig. 2. Assignments were made on the basis of E,,z-values observed in simpler systems, but it was impossible to distinguish between A,CFe(II) and AB,Fe(II), or between AC,Fe(II) and B,CFe(II). It was also assumed that negligible amounts of the binary iron(I1) derivatives would be formed in the ABC system, and the wave at 0.39 V was regarded as due entirely to ABCFe(I1). even though the potential corresponds exactly to E,,, of B,Fe(II).
Mixed-ligand Table
3. Calculated
Complex*
complexes and
1011
of iron and hydroxy-l,IO-phenanthrolines
observed reduction potentials phenanthroline complexes Binary complexest EL - 1.06 EL
(V) of ternary
Ternary
iron-
complexes$
E,,,,
E Obr
E ohs -
L,,
;;;4jWSWIA,
1.28 1.28 I .23 I.16 1.16 I .08 1.09
0.22 0.22 0.17 0.10 0.10 0.02 0.03
1.21 I.13 I.17 I.13 I 09 I .07
I .24
(~-NDz)AI (5-SO&A (5-Cl)2A (5-Cl)A2
1.09 1.16 I.13 I.10 I .07
+ 0.03 - 0.04 -0.01 0.00 +0.01 0.00
(5-CH,jA, (5-CH&A (4,7-diCHS)Az $di$,)zA
I.01 I.01 0.87 0.39 0.87
-0.05 - 0.05 -0.19 -0.67 -0.19
1.04 1.02 1.00 0.93 0.84
1.04 I .02 0.99 0.78 0.93
0.00 0.00 -0.01 -0.06 0.00
0.39 -0.10 -0.10
-0.67 - 1.16 - I.16
0.61 0.67 0.29
0.60 0.58 0.21
-0.01 -0.09 -0.08
(5-N%@
(4-OH),; (4.7-hOH)A* (4.7-diOH)2A
* The iron(III.11) complex couples are designated by the ligands only, with the following abbreviations: A (I,lO-phenanthroline), 5-NO2 (5-nitro-l,lO-phenanthroline), 5-Cl (5-chloro-l.I@phenanthroline), 5-SO3 (I,lO-phenanthroline-5sulphonic acid), 4,7-di(q%SO,) (4,7-diphenylsulphonato-l,lO-phenanthroline), 5-4 S-CHJ (5-methyl-l,l@phenanthroline), 4,7(5-phenyl-l,lO-phenanthroline), diCH, (4,7-dimethyl-l,lO-phenanthroline, 4-OH (4-hydroxy-l,lO-phenanthroline), 4,7-diOH (4,7-dihydroxy-l,lO-phenanthroline). t EL is the reduction potential of the binary complex with the substituted ligand. potential of the mixed complex predicted from the 5 EC,,, is the reduction weighted average. Eobr is the observed reduction potential of the mixed complex. Values for 4-OH and 4,7-diOH are from this work; others are taken from reference 16.
REFERENCES
CONCLUSION
The possibilities for application of the mixed complexes as redox indicators are limited by the fact that several mixed complexes will always exist in equilibrium with the desired complex. Perhaps the ABCFe(I1) system, in which at least six individual mixed complexes exist in equilibrium, with a continuous range of accessible potentials from 0.2 to 0.8 V, will lead to some novel applications. A solution of the ABCFe(I1) system could serve as a wide-range redox indicator system. in which absorbance is a measure of the reduction potential of a solution, much the same as the colour of an acid-base indicator shows the pH. Another possibility is the use of the mixed complexes as electrode mediators in coulometric titrations. Finally, it should be possible to extend the upper potential limit of this system by incorporating another ligand such as .5-nitro-IJO-phenanthroline in the mixed-ligand complex. The reduction potential of tris(4-hydroxy-l,lO-phenanthroline)iron(II) has not been previously reporied except as a passing note in a recent paper.” Its reduction potential of 0.39 V is significantly different from any previously reported for the binary iron-phenanthroline complexes, and it may find use as a redox indicator as well as a reagent in the spectrophotometric determination of traces of oxidants and reductants.
G. H. Walden
Jr., L. P. Hammett and R. P. Chapman, J. Am. Chem. Sot., 1931,53, 3908. 2. G. F. Smith and F. P. Richter, Phenanthroline and Substituted Phenanthroline fndicarors, G. F. Smith Chemical Co., Columbus, Ohio, 1944. 3. W. W. Brandt, F. P. Dwyer and E. C. Gyarfas, Chem. Reo., 1954, 54, 959. 4. A. A. Schilt, Analytical Applications of I,lO-Phenonthroline and Related Compounds. Pergamon, New York. 1969. W. W. Brandt and G. F. Smith, Anal. Chew 1949. 21. 1313. G. F. Smith and F. P. Richter, Ind. Eng. Chem.. Anal. Ed., 1944, 16. 580. M. N. Hale and M. G. Mellon, J. Am. Chem. Sot.. 1950. 72. 3217. P. George, G. I. H. Hanania and W. A. Eaten, unpublished work. in Hemes and Hemoproteins, B. Chance. R. W. Estrabrook, and T. Yonetani. eds., p. 269. Academic Press, New York, 1966. 9 D. P. Poe and H. Diehl, Talanta. 1974, 21. 1065. 10. Idem, ibid., 1976, 23, 147. Anal. I I. A. A. Schilt, G. F. Smith and A. Hiimbuch, Chem., 1956, 28, 809.
12. D. P. Poe and H. Diehl, Talanta. 1976, 23, 141. 13. W. M. Clark, J. Appl. Physics, 1938, 9, 97. 14. I. M. Kolthoff and V. A. Stenaer. Volumefric Analysis. Vol. I, 2nd Ed., pp. 105-141. interscience, New York. 1942. Iowa State University. 15. D. P. Poe. Ph.D. Dissertation, 1974. 16. P. J. Taylor and A. A. Schilt, Inorg. Chim. Acra. 1971. 5, 691.
1012
DONALD
P. POE and ALAN D. EPPEN
17. D. P. Poe and J. A. Adamczak, Talanta, 1980, 27, 1025. 18. H. R. Snyder and H. E. Freier, J. Am. Chem. SM.. 1946. 68, 1320.
19. D. P. Poe, A. D. Eppen and S. B. Whoolery, Talanta, 1980, 27, 368. 20. R. N. Adams, Electrochemistry at So/id Electrodes. Dekker, New York, 1969.