OXIDATION-REDUCTION INDICATORS

CHAPTER 4 O X I D A T I O N - R E D U C T I O N I N D I C A T O R S

O N E OF the earliest analytical uses of 1,10-phenanthroline was in the form of its iron(II) chelate as an oxidation-reduction indicator. Introduced by Waiden, Hammett, and Chapman in 1931[ΐ·2] the use of tris(l,10-phenanthroline)iron(II) as a high potential redox indicator soon gained many advocates, proving to be highly satisfactory for a variety of titrations involving strong oxidants. The availability of this indicator was a significant factor in the development and promotion of cerate oxidimetry, previously unpopular because tedious Potentiometrie end point detection was necessary.t^J Gleu, an early proponent of the indi­ cator, suggested the trivial name ferroin as an abbreviation for, in his words, "this most important and best oxidimetric indi­ cator, . . . the Phenolphthalein of oxidimetry .. . " t ^ l The accent mark was so frequently neglected that the term ferroin is now more commonly used. The oxidation-reduction potential of ferroin can be modified by introduction of various substituents into the 1,10-phenan­ throline groups, as evidenced by the higher potential of the 5-nitro derivative introduced by Waiden, Hammett, and Edmonds.t^^ Recognizing the analytical implications of this, G. F. Smith of the University of Illinois undertook a systematic investigation of substituted ferroins. Over a period of about 15 years in conjunction with F. H. Case of Temple University, Professor Smith and his students studied the iron and copper complexes of some 150 different substituted phenanthrolines and related compounds. A series of redox indicators covering the potential range 0.87-1.33 V, as well as a number of outstand­ ing colorimetric reagents, evolved from these studies. In addition to ferroin and substituted ferroins, many other metal chelates of a similar type have been investigated as redox 102

OXIDATION-REDUCTION INDICATORS

103

indicators. These are described in the present chapter. Use of the complexes as indicators in titrations other than redox titrations is considered in the next chapter. FERROIN

Ferroin very nearly meets all the requirements of the ideal indicator. Its orange-red color is sufficiently intense so that only a single drop of 0.025 molar solution is required for 50 ml of titrate solution. On oxidation to the ferric complex, ferriin, the color changes to a very pale blue, so slight in most cases as to be colorless. The color change is sharply discernible, rapid, and can be reversed repeatedly without adverse eff'ect. A correction for the amount of oxidant consumed by ferroin is significant only in microtitrations. Ferroin is remarkably stable in most solutions. Dissociation in the presence of strong acids occurs very slowly at room temperature but rapidly at elevated temperatures. Since the Perchlorate salt is very slightly soluble, ferroin is not appli­ cable as an indicator in perchloric acid solutions. Although less stable than ferroin, the oxidized complex, ferriin, is extremely slow to dissociate in the presence of strong acid, decomposing rapidly only if heated to 50''C or above. Solutions of ferroin at pH 2-9 are stable for years on storage under ordinary conditions. Unlike most indicators, oxidation of ferroin does not change the organic part of the complex. Redox indicators of the purely organic type, e.g. diphenylamine and triphenylmethane deriva­ tives, tend to be susceptible towards irreversible oxidation, particularly in the presence of excess oxidant. The highly reversible behavior of the ferroin-ferriin redox couple is a matter of considerable theoretical interest. When ferroin is oxidized to ferriin the phenanthroline ligands are not chemically altered; instead, it is the central ferrous ion that ultimately gives up an electron. Since it is very unhkely that an oxidant can penetrate the complex to come into direct contact with the central ferrous ion, the mechanism of the electron transfer probably involves direct electron tunnelling through the aromatic ligands. The dissociation rate for ferroin and the rate of exchange of iron between radioactive ferrous ion and ferroin are both slow;^®'^^ however, the rate of electron exchange between ferroin and ferriin is extremely fast {k > 10^ mole"* sec~* at

104

A N A L Y T I C A L APPLICATIONS OF 1,10-PHENANTHROLINE

25°C).t^J Hence a dissociation step prior to electon transfer is quite improbable. Clearly, the presence of the phenanthroline Hgands is conducive to rapid electron transfer, since the rate of electron exchange between aquated iron(III) and iron(II) is considerably slower {k = 0.87 mole"' sec'^ at OX)t»J than that between ferroin and ferriin. Kinetic studies of electron transfer reactions between ferroin and the following oxidants have been carried out by the rapid-mixing and flow technique: manganese(III),l^«l

peroxydisulfate

ion,[^il

cerium(IV),ti2> i3] ^^^1-

lium(III),t^21 and cobalt(III).l^^l In general the results are consistent with the Marcus theoryí^^-^^í for outer-sphere electron-transfer reactions. The high rate constants and low activation energies are indicative of very small solvent and ligand rearrangement barriers to electron transfer—i.e. the ferroin and ferriin are so very similar that very little rearrangement is necessary and electron transfer can proceed rapidly.t^^J Substituents in the phenanthroline ligands exert pronounced influences on the rates of electron transfer reactions,^^®'^^'^^'^®*^^! primarily as a con­ sequence of their eflfects on electron densities rather than for steric reasons. Formal Potentials The use of formal potentials is a practical convention for des­ cribing the behavior of redox systems. It avoids the problem of knowing activity coefficients, hydrolysis constants, dissociation and formation constants, etc., that would be required if standard potentials were to be used. Although applicable only for a given system of given composition, formal potentials are not difficult to measure and enable reliable predictions to be made by greatly simplified calculations. For the case at hand, the redox couple consisting of ferroin and ferriin involves the following halfreaction Fe(phen)3-'3-h^ = Fe(phen)3-^2

and the ratio of the formal concentrations of the two species conforms to the potential Ε predicted by the Ν ernst equation

* [Fe(phen)3+2]

OXIDATION-REDUCTION

INDICATORS

105

The quantity £ ° ' , which is the formal potential, corresponds to the experimental value of Ε for a given system containing equal formal concentrations of ferroin and ferriin. The magnitude of E°' varies with the composition of the solution and depends on the relative influences of solution composition on activity coefficients and stabilities of the complexes. Potentiometrie determination of formal potentials is a straightforward, uncompli­ cated procedure. In essence it consists of measuring the potential of a platinum electrode versus a suitable reference electrode (usually a saturated calomel electrode) immersed in a solution containing equal molar amounts of ferroin and ferriin, prepared from standard solutions or by titrating a known amount of ferroin with a standard solution of a strong oxidant. The ñrst measurement of the ferroin potential was made by Waiden, Hammett, and Chapman who reported a value of 1.14 V.1^1 They titrated a mixture of ferroin and ferrous sulfate potentiometrically with eerie sulfate, in a medium of 1 Μ sulfuric acid, using a quinhydrone half-cell in 1 Μ sulfuric acid as the reference electrode. The mid-point potential of the ferroin was 0.38 V higher than that of the ferrous-ferric couple and 0.29 V lower than that of the cerous-ceric couple. Assuming 0.75 and 1.44 V to be the standard potentials of the iron and cerium couples, respectively, they obtained an average value of 1.14 V for the standard potential of ferroin. As pointed out later by Hume and Kolthoff,^^^^ this result is incorrect because the standard potential of 0.75 V for iron is not applicable for 1 Μ sulfuric acid solutions. Moreover, the experimental data of Waiden, Hammett, and Chapman for solutions containing excess eerie sulfate were found to be in error by nearly lOOmV.^^^l When the known potential of 0.696 V for the quinhydrone electrode in 1 Μ sulfuric acid is used to calculate the ferroin potential directly from the mid-point potential observed by Waiden and coworkers, a value of 1.06 V is obtained for the formal potential. Hume and Kolthoff repeated the titrations of Waiden et al, using both the quin­ hydrone and the mercury-mercurous sulfate electrode in 1 Μ sulfuric acid as reference electrodes. They found a value of 1.06 V for the formal potential of ferroin in 1 Μ sulfuric acid. From visual titrations performed in 1 Μ hydrochloric acid media they con­ cluded that the formal potential in hydrochloric acid is essentially 1.06Valso.t2oi

106

ANALYTICAL APPLICATIONS O F 1,10-PHENANTHROLINE

Waiden, Hammen, and Edmonds found that the formal poten­ tial of ferroin decreases with increasing concentration of sulfuric acid.l^^i Taking advantage of this, they demonstrated that ferroin could serve as a practical indicator for the titrimetric determina­ tion of vanadium(V) with standard ferrous sulfate, provided that the titration is carried out in 5 Μ sulfuric acid. At lower concen­ trations of acid, end points occur prior to the equivalence point. Their data for the relative molar potentials, referred to the quinhydrone electrode at the same acidity, are given in Table 8. Included in the table are estimates of the formal potentials, referred to the standard hydrogen electrode and based upon a value of 0.696 V for the quinhydrone electrode in 1 Μ sulfuric acid (neglecting the effect of higher acid concentrations on its potential). The importance of considering the effects of acid concentration on formal potentials and indicator end points is clearly evident. TABLE 8. SOME FORMAL REDOX POTENTIALS Molar cone. ofH2S04

E°' vs. Quinhydrone electrode, V V(V)/V(IV)

1 3 5

0.33 0.39 0.44

Ferriin/Ferroin Fe(III)/Fe(II) 0.36 0.30 0.23

-0.01 -0.03 -0.05

E°' vs. Hydrogen electrode, V Ferriin/Ferroin 1.06 -1.00 -0.93

Successful application of ferroin as a redox indicator requires a knowledge of its formal potential as a function of acid concentra­ tion. This information is provided by the experimental data compiled in Table 9. A further practical consideration in the selection of any redox indicator is its transition potential—It, the potential at which the color change is first detectable. This characteristic is closely related to the formal potential, but its estimate varies somewhat from one individual to another depend­ ing on the observer's ability to perceive colors. In general, since the orange-red color of ferroin is so much more intense than the pale blue color of ferriin, approximately nine-tenths of the ferroin must be converted to ferriin before most observers can detect the color change. Accordingly, the transition potential is approxi-

107

OXIDATION-REDUCTION INDICATORS

mately 0.05 V greater than the formal potential, as predicted by the following: (^)transition

^ £^' +0.059 log ^^^""'"^ [ferroin] +0.059 log

= Ε°' + 0.05

Ideally the transition potential of the indicator should coincide with the equivalence point potential in the titration, but either or both can be influenced by solution composition. The equivalence point potential is determined by the formal potentials of the titrate and titrant species, which in turn are dependent on the nature and concentration of the solution. Thus in order to apply ferroin successfully as an indicator, it is necessary to consider the formal potentials of the titrate and titrant species as well as those of ferroin. By proper adjustment of solution conditions, within limits dictated by the variations in formal potentials, the transi­ tion potential of ferroin and the equivalence point potential of the titration can be brought into coincidence in certain cases. TABLE 9 . FORMAL REDOX POTENTIALS FOR FERROIN IN SULFURIC ACID SOLUTIONS H,SO, molarity

(volts)

Refs.

H2SO4 molarity

(volts)

Refs.

0

1.141

22

1.0

1.06

26

0

1.14

23

2.0

1.028

24

0

1.120

24

2.0

1.03

25

0.005

1.112

24

2.5

1.015

24 24

0.05

1.102

24

3.0

0.9%

0.05

1.10

25

3.0

1.00

25

0.25

1.086

24

3.5

0.977

24 25

1.0

1.06

20

4.0

0.96

1.0

1.057

24

6.0

0.89

25

1.0

1.06

25

8.0

0.76

25

Practical Titrations If ferroin is to be employed as a redox indicator, the titration must involve a strong oxidant, because the ferroin transition potential is quite high. Cerium(IV) salts are the most commonly used oxidants in conjunction with ferroin. Other oxidants include

108

ANALYTICAL APPLICATIONS O F

1,10-PHENANTHROLINE

permanganate, dichromate, and vanadate. All four oxidants are most effective in strong acid solutions, where their oxidation strengths are enhanced and their tendencies to undergo hydroly­ sis or other undesirable reactions are discouraged. Their formal potentials are dependent on both the nature and concentration of acid used. Data of this type are given in Tables 10, 11, and 12 TABLE 10. FORMAL POTENTIALS OF CERIUM(IV)CERIUM(III) (Reference: G. F . Smith and C . A . Getzi^^i)

Acid concentration, Ν

1 2 4 6 8

V o l t s v s . Normal Hydrogen electrode HCIO4

HNO3

H2SO4

HCl

1.70 1.71 1.75 1.82 1.87

1.61 1.62 1.61

1.44 1.44 1.43

1.28

— 1.56

—

1.42

— — — —

TABLE 11. FORMAL POTENTIALS OF CHROMIUM(VI)CHROMIUM(III) AND OF IRON(III)-IRON(II) IN H2SO4 AND IN H C l (Reference: G. F . Smith and F . P . Richterl^eJ) Volts v s . Normal H y d r o g e n electrode Acid concentration Μ

Cr(VI)/Cr(III) H2SO4

1 2 3 4 6 8

— 1.11

—

1.15 1.30 1.35

Fe(III)/Fe(II)

HCl

H2SO4

HCl

1.09 1.11 1.19 1.15

0.68 0.68 — 0.68 0.68 0.68

0.69 0.68 0.67 0.66

— —

— —

for cerium(IV), dichromate, and vanadate, respectively. Similar data are not available for permanganate. The latter behaves irreversibly and can be reduced to a variety of products, depend­ ing upon the pH or the presence of complexing agents. In acid solution, the standard potential for the reduction of permanganate to manganous ion is 1.51 V (calculated from other half-reactions that are reversible). Approximate formal potentials, observable

OXIDATION-REDUCTION

INDICATORS

109

in Potentiometrie titrations with permanganate in strong acid solution, tend to lie within the range 1.2-1.5 V. TABLE 12. FORMAL POTENTIALS OF VANADIUM(V)-VANADIUM(IV) (Reference: G . F. Smith and W. M. Banick, Jr.t^O H2SO4, F E^' (volts) H2SO4, F 0.10 0.30 0.40 0.50 0.75 1.0 Í.3 1.5 1.8 2.0 2.3 2.5 2.8

0.910 0.934 0.940 0.975 0.993 1.008 1.018 1.Ö30 1.047 1.056 1.Q72 1.089 1.095

3.0 3.3 3.5 3.8 4.0 4.3 4.5 4.8 5.0 5.3 5.5 6.0

(volts) 1.103 1.110 1.120 1.132 1.143 1.160 1.182 1.189 1.193 1.206 1.211 1.226

Since they have high formal potentials, both cerium(IV) and permanganate are satisfactory oxidants for use in conjunction with ferroin, Dichromate and vanadate are less satisfactory since they are weaker oxidants. It should be emphasized, however, that the formal potentials of all four oxidants increase with increasing acid concentration, while the opposite trend occurs in the potentials of ferroin. Thus there is a range of acid concentrations for each oxidant where its formal potential is appreciably greater than that of ferroin. Depending on the formal potentials of the reductant, titrations are therefore possible. For practical titrations involving dichromate or vanadate, the acid concentration should be approximately 5 Μ or greater if ferroin is to be suitable as the indicator. Various reducing substances can be determined by direct titratOn with cerium(IV) or other strong oxidant using ferroin

as the indicator. Most of these are listed in Table 13, with litera­ ture references and a summary of appropriate solution conditions. A number of reducing agents react too slowly or are too unstable for direct titration. These can be determined indirectly by adding a measured amount of cerium(IV) and back-titrating the excess

Pu(III) Sn(II) T1(I) Ce(IV) U(IV) V(IV) Μηθ4Hydroquinone

Crfir Fe(II) V(V) Fe(II) Fe(CN)6-^ H2O2 Ce(IV) ICe(IV) Mo(V) C204=

Fe(II)

Sb(IIl) As(III)

Determination of

Conditions

Refs.

Ce(IV) HCl, ICl catalyst, 50°C MnOr H2SO4, OSO4 catalyst Ce(I V) H2SO4, OSO4 catalyst Ce(IV) HCl (3-4 Μ), ICI catalyst, 50 C Μηθ4H2SO4 Ce(I V) H2S04(5 M, so that V(I V) does not interfere) HCl 33 H2S04(5M) 21 Ce(IV) HCl HCl 30 HCl, ICI catalyst 35 Ce(IV) HClandH3P04 Ce(IV) HCl, ICI catalyst, 50°C Ce(IV) HCl, Mn(II) catalyst, 20-25°C Ce(I V) HCl (1 Μ), BaCl2 as scavenger for S04=, 20-25X Ce(IV) H2SO4 Ce(IV) HCl HCl, ICI catalyst, 50°C 30 Ce(IV) H2SO„50X Acetate buffer, 50°C 34,43 Ce(IV) H2SO4

Titrant

TABLE 13. SOME DIRECT REDOX TITRATIONS USING FERROIN INDICATOR

44,45

30

36 30 37 38 39 40,41,42

30

30 4 4 30,31,32 2,34 2,5

110 ANALYTICAL APPLICATIONS O F 1,10-PHENANTHROLINE

Summary of procedure

Refs.

Al(III), Mg(II) Pptn. as oxinate; xs. Ce(IV); Fe(II) back-titm. As(0) Ce(IV) in H2SO4; As(III) back-titm. 48 Ca(II) Pptn. as oxalate; Ce(IV) titm. 49 Cr(III) Ce(IV) in H2SO4, hot; oxalate or nitrite back-titm. 50,51 Cr(VI) Meas. xs. Fe(II) in 5 Μ H2SO4; ΚΜηθ4 back-titm. 34 52 Cu(0) Meas. xs. Ce(IV) in H2SO4; Fe(II) back-titm. Cu(I) Excess FeOII) in H2SO4; titm. of Fe(II) with Ce(IV) 53 CIO3Meas. xs. As(III); OSO4 catalyst, Ce(IV) back-titm. 54 Pptn. as oxinate; xs. Ce(I V), Mn(II) and Ag(I) catalyst; Fe(ll) back-titm. Ga(III), In(III) 55 Hydrazoic acid Ce(I V) in HNO3; Fe(II) back-titm. Hydroxylamine Ce(IV) in H2SO4; hot; OSO4 catalyst, As(III) back-titm. 8-Hydroxyquinoline Ce(IV), hot, Mn(II) and Ag(I) catalyst; Fe(II) back-titm. Hypophosphite Ce(IV) in H2SO4; Fe(II) back-titm. Phosphite Ce(IV) in H2SO4, hot; Fe(II) back-titm. 57,58 Hg(I) Ce(IV) in H2SO4; Fe(II) back-titm. 59 Te(IV) Ce(IV) in HCIO4, Mn(II) and Ag(I) catalyst, hot; add H2SO4, Fe(II) back-titm. V(IV) Excess ferricyanide in NaOH (0.5-1 M); H2SO4, titm. of ferrocyanide with Ce(IV) Glycerol or other polyhydric alcohols Ce(IV) in H2SO4, Mn(II) and Ag(I) catalyst, 95X; Fe(II) back-titrn.

Determination of

TABLE 14. SOME INDIRECT DETERMINATIONS USING FERROIN INDICATOR

46

60

57

56 46 57,58

46

46,47

OXIDATION-REDUCTION INDICATORS

111

112

ANALYTICAL APPLICATIONS OF

1,10-PHENANTHROLINE

with Standard ferrous sulfate. Another indirect procedure consists of adding an excess of iron(III) to the reductant, followed by titration with cerium(IV) of the iron(II) formed. Examples are given in Table 14. The use of ferroin as a redox indicator is very common in cerium(IV) oxidimetry. For details and procedures in the use of cerium(IV) oxidants, a review article by Youngt^^i and the book by Smith entitled Cerate Oxidimetry^^^ may be consulted. SUBSTITUTED

FERROIN

DERIVATIVES

The formal potential of ferroin can be appreciably altered by group substitutions in the 1,10-phenanthroline ligands. Nucleophilic substituents tend to increase and electrophilic substituents tend to decrease the formal potential. A direct relationship exists between formal potentials and ligand p/C^ values of substituted 1,10-phenanthroline iron(II) complexes.t^^.62] ¡j^ case of derivatives with substituents in the 5-position, the relationship is a linear one.i^*^ Also it is observed that the effects of methyl substituents on the pKa values of phenanthrolines are additive.t^^l Their effects on formal potentials of the iron(II) trischelates are also additive.í^^í These relationships are reasonable, considering that the ligand pA^« is a measure of the electron density about the nitrogen atoms. An electron-withdrawing substituent lowers the stability of the complex in both oxidation states; however, the stability of the iron(III) form predictably is decreased more than that of the iron(II), thus the formal potential is lowered also. The converse is true for electron-releasing substituents. Formal potentials are compiled in Table 15 for a number of substituted ferroin derivatives, as a function of acid concentration. Certain ferroin derivatives are not sufficiently soluble to permit direct measurement of their formal potentials by ordinary methods. Smith and Banicki^«] devised a colorimetric procedure whereby these could be characterized. Equal amounts of the ferrous complex to be tested are added to equal volumes of a graded series of potentiopoised solutions of known redox potentials. Upon visual examination of the resulting solutions, to find in which solution the color transition is completed, an estimate of the transition potential of the complex can be made. After confirming the validity of their procedure with soluble complexes of known formal potentials. Smith and Banick deter-

OXIDATION-REDUCTION

113

INDICATORS

T A B L E 15. F O R M A L R E D O X P O T E N T I A L S O F S O M E SUBSTITUTED FERROIN DERIVATIVES

1,10-Phenanthroline derivative 3-Methyl 5-Bromo 5-Chloro 4-Methyl 5-Nitro 5-Nitro-6-methyl 3,4-Dimethyl 3,8-Dimethyl 4,5-Dimethyl 4,6-Dimethyl 4,7-Dimethyl 5,6-Dimethyl 3,4,6-Trimethyl 3,4,7-Trimethyl 3,5,7-Trimethyl 3,5,8-Trimethyl 3,4,6,7-Tetramethyl 3,4,6,8-Tetramethyl 3,4,7,8-Tetramethyl -3-Sulfonic acid -5-Sulfonic acid Bathophenanthrolinedisulfonate

Formal Potentials in H2SO4 Refs. 0.1 Μ 0.5 M 1.07

1.0 M

2.0 M

4.0 M

1.03 1.13

1.26 0.97 1.03 0.95 0.95 0.88 1.00 0.92 0.88 0.93 0.99 0.84 0.89 0.85

1.11 1.02 1.25 1.23 0.93

1.10 l.OO 1.22

1.04 0.93 1.12

0.87 0.97

0.89

1.23(1.21inMHC104) 1.20(1.16inMHC104) 1.09(1.01inMHClO4)

25 26 26 26 26 26 25 25 25 25 25 25 25 25 25 25 25 25 25 64 64 65

mined the potentials of a number of slightly soluble ferroin deri­ vatives. For potentiopoised solutions, they employed sulfuric acid solutions containing equal molar amounts of vanadium(V) and vanadium(IV). The redox potentials of the graded series of potentiopoised solutions ranged from 0.91-1.23 V, as governed by the sulfuric acid concentration (see Table 12). Transition potentials of the complexes were converted to formal potentials using the relationship E"' = ^transition — 0.06. The results are compiled in Table 16. Although numerous substituted ferroin derivatives have been examined as possible redox indicators, only a few have proven to be sufficiently different from ferroin to be of special or unique value. The most important of these is probably the 5-nitro deri­ vative (commonly referred to as nitroferroin), which because of its high formal potential is particularly well suited for titrations using cerium(IV) in perchloric or nitric acid solutions. The

1.0

1,10-Phenanthroline derivative

3,7-Dimethyl 1.4 3,8-Dibromo 5.4 4,7-Dimethoxy 0.1 4,7-Diethyl 0.9 4,7-Diphenoxy 2.9 4,7-Diphenyl 4.6 4,6-Diphenyl 4.5 5,6-Cyclohexeno 1.4 5,6-Diethyl 1.6 5,6-Dimethoxy 2.4 5,6-Dichloro 4.4 3,5,6-Trimethyl

H2SO4 E°' formality (volts)

5-Fluoro 3.2 1.05 3-Chloro 4.4 1.11 3-Phenyl 2.4 1.02 3-Ethyl 2.1 1.00 4-Bromo 2.9 1.04 4-Phenyl 1.9 0.99 4-Ethyl 1.4 0.96 4-n-Propyl 1.4 0.96 3,4-Cyclopenteno 1.4 0.96 3,4-Cyclohexeno 0.9 0.94 3,4-Cycloocteno 1.0 0.95 (3,4)(7,8)-Dicyclohexeno 0.3 0.87

1,10-Phenanthroline derivative

0.95

H2SO4 E^' formality (volts) 0.96 1.21 0.85 0.94 1.04 1.13 1.12 0.96 0.98 1.02 1.11

(Reference G. F. Smith and W. M. Banick, Jr.l^si)

TABLE 16. COLORIMETRIC FORMAL REDOX POTENTIALS OF SOME SUBSTITUTED FERROIN DERIVATIVES AS ESTIMATED FROM TRANSITION POTENTIALS

114 ANALYTICAL APPLICATIONS OF 1,10-PHENANTHROLINE

OXIDATION-REDUCTION INDICATORS

115

sulfonated derivatives are outstanding because of their solubiHty in systems containing Perchlorate ion, where other ferroin deriva­ tives generally are quantitatively precipitated. Certain of the methyl derivatives are especially useful in titrations involving dichromate because of their lower formal potentials. The more important derivatives and their application are described below. Nitroferroin Use of the iron(II) complex of nitrophenanthroline as a redox indicator was first proposed by Hammett, Waiden, and Ed­ monds, The earliest practical application was made by Smith and Getz,t®®l who demonstrated that the high transition potential of nitroferroin is ideally suited for titrations involving cerium(I V) Perchlorate or cerium(IV) nitrate in perchloric or nitric acid solutions. Ferrous sulfate, arsenic(II), and sodium oxalate were sucessfully titrated. The determination of arsenic(lll) in perchlo­ ric acid with osmic acid as catalyst, cerium(IV) Perchlorate as oxidant, and nitroferroin as indicator proved more satisfactory than the corresponding titration in sulfuric acid solution using cerium(I V) sulfate and ferroin. Nitroferroin is most commonly used as an indicator in con­ junction with cerium(IV) Perchlorate as an oxidant. By taking advantage of the higher oxidation potential of cerium(lV) in perchloric acid. Smith and Duke found conditions whereby a variety of organic substances could be oxidized stoichiometrically.t^^l Results obtainable using cerium(IV) in sulfuric acid solutions for the same p u φ o s e are generally quite empirical. Smith and Duke demonstrated that a high degree of accuracy can be obtained in the determination of glycerol, glucose, sucrose, cellulose, biacetyl, acetylacetone, tartaric acid, malonic acid, citric acid, and malic acid by the following general procedure: known amount of the compound is added to an excess of perchlorato-cerate in 4 Μ perchloric acid at the given reaction temperature. After sufficient time, the excess perchlorato-cerate in 4 Μ perchloric acid is diluted to 2 Μ by the addition of an equal volume of water and is back titrated using standard oxalate solution with nitroferroin as indicator, the reaction being carried out at ordinary temperatures.''^^^1 The oxidation products, stoichiometry, and appropriate reaction temperature and time for each compound are given in Table 17.

116

ANALYTICAL APPLICATIONS OF

1,10-PHENANTHROLINE

T A B L E 1 7 . C E R I U M ( I V ) P E R C H L O R A T E O X I D A T I O N S IN 4 Μ P E R C H L O R I C A C I D

(Reference: G. F. Smith and F. R. Dukei«^!) Compound Glycerol Glucose Sucrose Cellulose Biacetyl Acetylacetone Tartaric acid Malonic acid Citric acid Malic acid

Equivalents

Conditions

T e m p , °C Time, min C e ( I V ) p e r mole

Products, moles/mole H C O O H CH3COOH C 0 2

45

15

8

3

0

0

26

45

12

6

0

0

24

45

26

6

0

1

27

120

12

6

0

0

24

5

2

0

2

0

25

10

6

1

2

0

26

10

6

2

0

2

26

10

6

1

0

2

10

30

14

2

0

4

25

15

8

2

0

2

Procedures for the micro determination of oxalic acid, iron, and arsenic have been reported by Smith and Fritz,í^^í and of calcium in blood by Salomon, Gabrio, and SmithJ^^l In each case, titration with cerium(IV) Perchlorate in 2 M perchloric acid using nitroferroin as indicator gave precise, stoichiometric results. Mercury(I) can be determined oxidimetrically using cerium(I V) Perchlorate and nitroferroin.^^^ The titration of mercury(I) is initially carried out at the boiling temperature in a solution of perchloric acid containing manganese(II) and silver(I) as catalysts. Prior to the equivalence point, the solution is cooled to 50-60°C, nitroferroin is added, and the titration is continued to the end point. According to Rao, nitroferroin is a suitable indicator for the titration of oxalic acid with cerium(IV) sulfate in 0.5 Μ hydro­ chloric, perchloric, or nitric acid at room temperature.t^^J Sulfuric acid solutions are unsatisfactory for the titration because nitro­ ferroin is oxidized faster than the oxalic acid, which in turn is very slow to be oxidized by the oxidized nitroferroin. The reduc­ tion of oxidized nitroferroin is much more rapid in hydrochloric, perchloric, or nitric acid solutions. Oxidized nitroferroin is more rapidly reduced than oxidized ferroin by oxalic acid in any acid solution. Methyl Derivatives Dimethylferroin, the ferrous complex of 5,6-dimethyl-l,10phenanthroline, is reported to be a suitable indicator for the

OXIDATION-REDUCTION INDICATORS

117

titration of ferrous iron in 1 - 2 Μ hydrochloric or sulfuric acid with potassium dichromate.t^^i In hydrochloric acid the end point corresponds to a change in color from orange to green, and in sulfuric acid from red to yellow-green. Although both the formal potential and the transition potential of the indicator are some­ what greater than the apparent equivalence point potential, satisfactory results were nevertheless obtained. Interestingly, the reverse titration of dichromate with ferrous can be performed accurately in 1-2 Μ sulfuric acid using ferroin as the indicator, whereas the direct titration of ferrous iron with dichromate fails with ferroin.^^^^ The irreversible behavior of the dichromatechromic ion system is clearly responsible. For the same reason, it is possible that the actual equivalence point potential in the dichromate titration of ferrous iron may be greater than that observed, so that the transition potential of dimethylferroin may actually be quite close to that of the equivalence point. Titration of ferrous iron with dichromate in solutions of low acidity requires an indicator with a formal potential that is lower than that of either ferroin or the 5,6-dimethyl derivative. Smith recommended the use of tris(4,7-dimethyl-1,10-phenanthroline) iron(II) as the indicator (£°' = 0.88 V) for the titration in 0.5 Μ hydrochloric or sulfuric acid, and for 0.1 Μ acid solutions the use of tris(3,4,7,8-tetramethyl-l,10-phenanthroline)iron(II) as the indicator = 0.85 V).í^^> In either case the transition potentials are somewhat higher than the Potentiometrie inflection point potentials, but not to an extent that an appreciable differ­ ence exists between visual and Potentiometrie end points. Sulfonated Derivatives The trischelated ferrous complexes of l,10-phenanthroline-5sulfonic acid and of l,10-phenanthroline-3-sulfonic acid are somewhat more sensitive to acid than the parent ferroin, but the 5-sulfonic derivative is sufficiently stable for use as a redox indicator in strong acid solutions.t^^l Both complexes give vivid sharp color changes on oxidation and have the advantage over ferroin that they can be used in perchloric acid solutions without risk of precipitation. Accoring to Blair and Diehl they are ideal for titrations with cerium(IV) in either sulfuric or perchloric acid; however, no practical applications were reported.^^^ Formal potentials are given in Table 15.

118

ANALYTICAL APPLICATIONS O F

1,10-PHENANTHROLINE

Application of ferrous bathophenanthrolinedisulfonic acid as an indicator in the cerium(I V) titrimetric determination of iron in ores was made by Blair and DiehlJ^^i The titrations, performed in mixed perchloric, phosphoric and sulfuric acid solutions, gave vivid end points, sharply discernible by a color change from red to green. Their analytical results for three different standard iron ores were very precise but significantly lower than previously determined values. No explanation for the discrepancies could befound.t«^i

BIPYRIDINE AND TERPYRIDINE OF FERROIN

ANALOGS

The tris(2,2'-bipyridine)iron(II) cation (bipyridine ferroin) undergoes a color change from intense red to faint blue on oxida­ tion. Its properties as a redox indicator closely parallel those of ferroin, except that it is much more prone to dissociate in strong acid solutions. As a substitute for ferroin the only advantage it affords is its lower cost. The formal redox potential of bipyridine ferroin, like that of ferroin, decreases with increasing acid concentration. Values are given in Table 18. At one time the generally accepted value for the formal potential in 1 Μ sulfuric acid was 0.97 V. This posed an apparent anomaly: bipyridine ferroin would be expected to function as well as 5,6-dimethyl ferroin in titrations with dichro­ mate; however, it was found to behave exactly like ferroin. ^^^i anomaly was explained by Belcher and co-workers. On re­ determining the formal potential, they found it to be 1.023 V. TABLE 18. FORMAL REDOX POTENTIALS FOR TRIS(2,2'-BIPYRIDINE)IRON(II) SULFATE IN SULFURIC ACID SOLUTIONS H,SO, molarity

(volts)

0 0.01 0.01 0.11 1.0 1.0

1.0% 1.084 1.069 1.062 0.97 1.023

Refs.

H2SO4 molarity

(volts)

24 24 73 73 26 72

1.0 2.0 4.0 4.0 6.0 10

1.026 1.00 0.95 0.92 0.88 0.8

Refs. 73 73 73 26 73 73

OXIDATION-REDUCTION

INDICATORS

119

Apparently the error in the earlier value was due to a failure to make allowance for the appreciable rate of dissociation of the complex in sulfuric acid solutions.t^^'^^J There are very few references in the literature regarding the use of bipyridine ferroin. Waiden, Hammett, and Chapman reported it to be less stable than ferroin.i^l Cagle and Smith found it to perform satisfactorily in the titration of iron(II) with cerium(IV), provided that it did not remain in contact too long with excess oxidant.t^^i Nieuwenburg and Blumendal determined small amounts of iron by cerate oxidimetry employing 2,2'-bipyridine to form the indicator in situP^^ Sakai, using the iron(II) bipy­ ridine complex as indicator for the cerimetric determination of hydroquinone, reported that an indicator blank correction was necessary and that the final concentration of sulfuric acid should be 1 M.t^ei The bis chelate of iron(II) with 2,2',2"-terpyridine undergoes a vivid color change from violet to green on oxidation. Dwyer and Gyarfast^^i found formal redox potential values of 1.076, 1.060, 1.054, and 0.927 V in 0.1, 0.2, 0.5, and 1.0 M sulfuric acid, res­ pectively. They surmised that the instability of the oxidized form would be a serious deterrent to the application of the complex as an indicator. The terpyridine ferrous chelate closely resembles that of bipyridine. Neither one has much to recommend it over the use of ferroin, which covers the same potential range with much less instability towards strong acids.

OTHER

METAL CHELATES OF THE

FERROIN

TYPE

Ruthenium and osmium, as would be expected from their close relationship to iron, form complexes with 1,10-phenanthroline or 2,2'-bipyridine that possess many of the same desirable indicator characteristics as ferroin. None of the other metal ions, however, are particularly outstanding in this regard. Their complexes either lack sufficient stability in solutions appropriate for redox titrations, or the color intensities of their various oxidation states are too similar or too weak. Some of the more likely complexes that have been or that might be considered as possible indicators are those of chromium, vanadium, cobalt, and copper. These, as well as the complexes of ruthenium and osmium, are described below.

120

ANALYTICAL APPLICATIONS OF 1,10-PHENANTHROLINE

Ruthenium The ruthenium(II) complexes of 1,10-phenanthroline, 2,2'bipyridine, and 2,2',2''-terpyridine are highly colored (orange-red to orange-yellow) and very resistant towards decomposition by strong acids. They can be reversibly oxidized to the respective ruthenium(III) complexes by treatment with very strong oxidants in acid solution, giving light-green colored solutions. Their formal potentials are listed in Table 19. Increasing acid concentrations cause the formal potentials to decrease, similar to the trends observed for ferroin and its analogs. Dwyer determined the formal redox potentials of some ruthenium(II) complexes of some substituted derivatives of 2,2'T A B L E 19. F O R M A L R E D O X P O T E N T I A L S O F R U T H E N I U M C O M P L E X E S O F T H E FERROIN T Y P E

Cone, moles/ liter

r°c

HNO3

0.002 0.1 0.3 0.5 1.0 2.0 3.0 5.0 1.0 2.0 5.0

H2SO4

H2SO4

Acid

E ° O f R u ( I I I ) / R u ( I I ) , volts Trisphen

Refs.

Trisbipy

Refs.

0 0 0 0 0 0 0 0 15 15 15

1.31 1.30 1.29 1.28 1.26 1.24 1.22 1.19 1.29 1.26 1.22

77 77 77 77 77 77 77 77 78 78 78

1.303 1.288 1.279 1.270 1.257 1.240 1.222

77 77 77 77 77 77 77

2.5 3.55 5.0 6.0 6.75

15 15 15 15 15

1.22 1.205 1.16 1.15 1.105

78 78 78 78 78

0.01 0.05 0.11 1.0 2.0 4.0 6.0 10 12

25 25 25 25 25 25 25 25 25

1.274 1.273 1.262 1.236 1.20 1.15 1.09 0.88 0.76

73 73 73 73 73 73 73 73 73

Bisteφy

1.281 1.202 1.175

Refs.

79 79 79

OXIDATION-REDUCTION INDICATORS

121

bipyridine and 1,10-phenanthroline in nitric acid solutions.i^^J Methyl groups lower the potentials, analogous to the trend for the ferroin derivatives. The complex with 5-bromo-1,10-phenan­ throline exhibited the highest formal potentials: 1.41 and 1.36 V in 0.002 and 0.1 molar nitric acid, respectively. Relatively little use has been made of the various ruthenium complexes as redox indicators, perhaps because nitroferroin, which serves approxi­ mately the same potential range, gives a much more vivid color change. Steigman, Birnbaum, and Edmonds^^^i found that tris(2,2'bipyridine)ruthenium(II) nitrate is satisfactory as an indicator for titrating sodium oxalate in 2 Μ perchloric acid with cerium(IV) nitrate. Carried out at ordinary temperatures, the titrations are rapid, precise, and accurate. The formal potential of 1.33 V that they reported for 1 Μ sulfuric acid is probably in error (refer to Table 19). Kratochvil and Zatko^^^i investigated the ruthenium complexes of 2,2'-bipyridine, 1,10-phenanthroline, and several methyl derivatives as reversible fluorometric indicators for redox titrations. Of the complexes studied, tris(2,2'-bipyridine)ruthenium(I I) chloride proved the best for titrations with cerium(I V) sulfate or permanganate. The end point behavior of the indicators, marked by the disappearance of the orange-red fluorescence of the ruthenium(II) species, is rapid, sensitive, and reversible. A significant advantage over ferroin and nitroferroin is that the fluorescent indicators can be used in highly colored solutions, provided that absorption is not so excessive in the regions 450-465 m ^ (activation maxima) and 575-590 m ^ (fluorescent maxima) that fluorescence is obscured. Substances that form precipitates in the course of the titration interfere with the detection of fluorescence, since any turbidity causes light scattering and dimunition in visual perception of the fluorescence. Satisfactory results were obtained in titrating sodium arsenite, sodium oxalate, Oesper's salt (ferrous ethylenediammonium sulfate tetrahydrate), vanadyl sulfate, and hydrogen peroxide with cerium(IV) sulfate using tris(4,4'-dimethyl-2,2'-bipyridine) ruthenium(II) as fluorescent indicator. The titration of sodium oxalate or sodium arsenite with cerium(IV) Perchlorate was satisfactory using tris(2,2'-bipyridine)ruthenium(II) fluorescent indicator.

122

ANALYTICAL APPLICATIONS O F

1,10-PHENANTHROLINE

Osmium The tris chelates of osmium(II) with 1,10-phenanthroline and 2,2'-bipyridine possess suitable properties for use as indicators. Their redox behavior is reversible, and the color change on oxidation is sharply discernible from green to pale red. The bis chelate of osmium(II) and 2,2',2"-terpyridine, according to Dwyer and Gyarfas,^^^! is unsatisfactory as an indicator because of the instability of its oxidized form. No such limitation exists for the phenanthroline or bipyridine derivatives. The formal redox potentials, listed in Table 20, suggest that the complexes should be especially suitable as indicators for redox titrations involving dichromate or vanadate as the oxidant. Surprisingly little is to be found in the chemical literature on the use of the complexes as indicators. TABLE 20. FORMAL REDOX POTENTIALS OF OSMIUM COMPLEXES OF THE FERROIN TYPE H2SO4 Moles/liter 0 0.01 0.1 0.11 0.2 0.5 1.0 1.0 1.5 2.0 2.5 4.0 6.0 10 12

£ « O f O s ( I I I ) / O s ( I I ) , volt Trisphen

Refs.

Trisbipy

Refs.

0.877

80

0.859

80

0.877 0.858 0.855 0.841

80 73 80 73

0.822

80

0.819 0.802 0.803 0.775 0.777 0.727 0.723 0.63 0.44 0.37

80 80 73 80 73 80 73 73 73 73

Bisteφy

Refs.

0.951

79

0.941 0.925 0.907

79 79 79

0.884

79

Dwyer and Gibsoni^^l found that tris( 1,10-phenanthroline) osmium(II) Perchlorate gave distinct end points in titrating solutions of iron(II) in 0.5 Μ sulfuric acid containing some phosphoric acid with potassium dichromate as the oxidant. They reported that the bipyridine analog was unsatisfactory for the titrations but gave no reasons.

OXIDATION-REDUCTION INDICATORS

123

The indicator response of tris(2,2'-bipyridine)osmium(II) sulfate in 2 M sulfuric acid, evaluated by means of simulated titrations, was found to be rapid and reversible on alternate addition of cerium(I V) sulfate and ferrous sulfate, hydroquinone, or potassium ferrocyanideJ^^l Further study of the use of osmium complexes as indicators should prove rewarding, particularly since their formal potentials encompass an intermediate range not covered by other common indicators. Vanadium Shaeffert^^l explored the possibility of using the intensely blue-violet vanadium(II) complex of 1,10-phenanthroline as a redox indicator. In dilute acid solution the color is bleached by oxidizing agents at about 0.0 V vs. the saturated calomel electrode. Although the vanadium(II) complex has not been identified as yet, its maximum molar absorptivity is approxi­ mately 8000 at 645 mμ. As an indicator in the titration of ferric chloride with chromous chloride, it gave precise and accurate results. The color change from green (due to the chromic ions) to blue is rapid in titrations carried out at room temperature and coincides closely with the Potentiometrie end point. Chromium, Cobalt, and Copper The red complex produced by adding 1,10-phenanthroline to a solution of chromous acetate undergoes a sharp and rever­ sible color change to red-violet on treatment with oxidants in acid solution.t^l Hammett et al. estimated that the potential of the system is not appreciably diflferent from that of the ferricferrous p o t e n t i a l . I t is not considered to be of value as an indicator, since the color intensity is not high and the color change on oxidation is not pronounced. The cobalt(II) and cobalt(III) complexes of 2,2'-bipyridine and 1,10-phenanthroline lack sufficient color to serve as effective redox indicators. Otherwise, their redox behavior is satisfactory. A value of 0.37 V for the redox potential of the phenanthroline species has been reported.i®^^ The cobalt(II) complex of 2,2',2"terpyridine, which is more strongly colored than the others, might be of some value as an indicator, although it may lack adequate solution stability.

124

ANALYTICAL APPLICATIONS O F 1,10-PHENANTHROLINE

Formal redox potentials have been measured for a number of copper complexes of 1,10-phenanthrohne, 2,2'-bipyridine, and derivatives as a function of pH and ligand concentration in aqueous and in 50% dioxane-water mixturesJ^^'^^i The values range from 0.09-0.6 V in aqueous solutions. Although reason­ ably colored, the copper(I) complexes lack promise as indicators since they are not appreciably soluble in aqueous solutions and they decompose rapidly in acid or strong base solutions. MIXED L I G A N D COMPLEXES OF IRON, R U T H E N I U M , AND

OSMIUM

Another way of modifying the properties of ferroin and its analogs, besides introducing substituents into the aromatic groups, is to replace one or two of the bidentate ligands with other coordinating molecules or ions. This approach has been particularly successful using cyanide ions to replace 1,10phenanthroline or 2,2'-bipyridine.^^^^ All three ligands are suffic­ iently alike in their coordinating abilities, crystal field strengths, and tendencies toward π-bonding that it is relatively simple to prepare their mixed ligand complexes of iron. Both ferroin and ferrocyanide are diamagnetic, non-labile, extremely stable, and highly resistant to acid attack at ordinary temperatures. The former has a high redox potential, the latter a low potential (Ε°' ^ 0.4 V). Thus it is not surprising that the intermediate complexes dicyano-bis(l,10-phenanthroHne)iron(II) ([Feipheug (CN)2])

and

tetracyano-mono(l ,10-phenanthroline)ferrate(II)

([Fe(phen)(CN)4]"2) have redox as well as other properties that are intermediate to those for ferroin and ferrocyanide. The trivial name ferrocyphen was given to dicyano-bis (1,10-phenanthroline)iron(II) when it was found to have definite analytical value as an indicator, not only for redox but surprisingly for non-aqueous acid-base titrations as well.i^^i Similarly, the bipy­ ridine analog is referred to in abbreviated form as ferrocypyr, and the corresponding iron(III) species as ferricyphen and ferricypyr. The formal redox potentials of ferrocyphen and ferrocypyr in sulfuric acid solutions of various concentrations are given in Table 21, together with values for the ruthenium and osmium analogs dicyano-bis(2,2'-bipyridine) ruthenium(II) and dicyano-

0.01 0.05 0.11 0.50 1.0 2.0 4.0 6.0 10 12

moles/liter

H2SO4

0.806 0.818 0.852 0.925 0.990

0.786 0.787 0.820 0.901 0.979

^ Fe(phen)2(CN),]

1.13 1.16 1.22 1.26 1.30

0.781 0.777 0.776 0.776

1.12 1.11 1.10 1.12

[Fe(bipy )2(CN

0.836 0.89 0.95 1.01 1.02

0.778 0.783 0.791 0.804

[ Ru(bipy)2(CN),]

Formal potential of complex, volts

(References: 73 and 89)

0.810

[Osibipy^CN),]

TABLE 21. FORMAL REDOX POTENTIALS OF MIXED LIGAND COMPLEXES IN SULFURIC ACID SOLUTIONS AT 25°C

OXIDATION-REDUCTION INDICATORS

125

126

ANALYTICAL APPLICATIONS OF 1,10-PHENANTHROLINE

bis(2,2'-bipyridine)osmium(II). Two trends are clearly evident: (1) in dilute acid solutions all of the dicyano derivatives have lower formal potentials than the corresponding tris chelates, and (2) the formal potentials of the dicyano complexes increase with increasing acid concentration, just the opposite of the trend shown by the tris chelates. Both of these observations are reasonable. The first was predictable a priori from the known potentials of ferroin and ferrocyanide, also from the rationale that the removal of electrons (oxidation) should be more difficult from divalent cations (as in the case of the tris chelate) than from uncharged species (as in the case of the neutral dicyano complex). An explanation for the influence of acid concentration on the formal potentials is found in the observation that all of the metal(II) dicyano complexes have appreciable dibasic character, whereas the corresponding metal(III) complexes are much weaker bases.t^J With increasing concentration of strong acid the extent of protonation of the metal(ll) complexes increases, and the species take on greater positive charge. Since removal of electrons from positively charged species is more difficult than from uncharged species, the oxidation process requires more positive potentials. Also the process necessitates loss of protons to the medium because the metal(III) complexes are considerably weaker bases than the metal(II) complexes. Thus, increased acidity, in discouraging loss of protons as well as electrons, gives rise to more positive formal potentials. With the exception of the ruthenium complex, which does not behave reversibly on oxidation and reduction, all of the mixed ligand complexes listed in Table 21 have desirable redox in­ dicator properties. They are remarkably stable in acid solution and undergo distinct and reversible color changes on oxidation. In dilute acid solutions the colors of [Fe(phen)2(CN)2], [Fe(bipy)2 ( C N ) 2 ] , [Ru(bipy)2(CN)2], and [Os(bipy)2(CN)2] are orange, red, yellow-orange, and red-brown, respectively. In concentrated acid solutions the respective colors are yellow, yellow, pale yellow, and orange. The colors of the metal(Il) complexes are influenced by acid concentration because of their weak dibasic c h a r a c t e r . T h e metal(III) complexes are much weaker bases; and their colors are much less dependent on acid concentration. All three are pale violet in most solutions, except in very con­ centrated sulfuric acid where they are green. Thus, even though

OXIDATION-REDUCTION INDICATORS

127

the color prior to the end point is dependent on acid concentra­ tion, the end point color on oxidation of the various complexes is usually pale violet, and in any event the color transition is readily discernible. As redox indicators, the mixed ligand complexes of iron(II) and osmium(II) are rather unique in that their formal potentials increase with increasing acid concentration. Thus they are ideally suited for use in conjunction with certain oxidants-e.g. vanadate or dichromate-that also exhibit increasing formal potentials with increasing acid concentration. The control of acid concentration to favor coincidence of the indicator transition potential with that of the equivalence point is much less critical in such cases. It is also significant that the mixed ligand complexes have sufficiently low potentials to be practical as indicators for use with vanadium(V), chromium(VI), and other oxidants that have formal potentials of 0.9 V or higher. Ferrocyphen and ferrocypyr exhibit very similar indicator characteristics, and presumably they can be used interchangeably. Ferrocyphen has been used more extensively, since its colors are somewhat more intense. The osmium(II) analog of ferrocypyr has not been employed as an indicator in practical titrations, as yet,.but certainly merits consideration. Ferrocyphen has been used successfully as an indicator in a variety of titrations. A summary of these is given in Table 22. One of its most remarkable attributes is its ability to serve as a reversible internal indicator for diazotization titrations of aro­ matic amines with sodium nitrite. Prior to the introduction of ferrocyphen, such titrations were tedious and time consuming, requiring use of either an external indicator or Potentiometrie end point detection. The basis for the reversible behavior of ferrocyphen in titrations involving nitrous acid (from sodium nitrite) has been elucidated. i^^J Another attribute of ferrocyphen is its ability to function satisfactorily over a broad range of acid concentration, a consequence of the fact that its formal potentials as well as those of the titration species tend to be influenced similarly by acid concentration. The addition of phosphoric acid in titrations of iron(II) is necessary in order to lower the formal potential of the iron(III)-iron(II) couple, so that the equivalence point and indicator transition potentials will closely approximate one another.

128

ANALYTICAL APPLICATIONS O F

1,10-PHENANTHROLINE

T A B L E 22. T I T R A T I O N S U S I N G F E R R O C Y P H E N AS I N D I C A T O R (References: 89 and 91) Titrant K2Cr20; HVO3 Ce(S04)2 HV03 Ce(S04)c NaN02 NaN02 NaN02 NaNOz NaN02 NaN02 NaN02 NaN02 NaN02 NaN02

Titrate FeS04 FeS04 FeS04 Hydroquinone Hydroquinone Aniline p-Bromoaniline o-Chloroaniline 2,4-Dichloroaniline Sulfanilamide Sulfamic acid Sulfanilic acid Hydroxylamine Sodium azide Sodium azide

Solution and conditions

Rel. std. dev.,%

Η,Ρθ4; l-6MH2S04or2-4MHCl H 3 P O 4 ; l - 6 M H 2 S 0 4 o r 1-2 Μ H C l H 3 P O 4 ; 1 - 6 Μ H 2 S O 4 or 1 - 2 Μ HCl 0.5-4 Μ H , S 0 4 o r 1-2 Μ H C l l-6MH2Sb4or 1 Μ HCl 4 - 9 Μ HCl 6 Μ HCl 6 Μ HCl 6 Μ HCl 6 Μ HCl 6 Μ HCl 6 Μ HCl 6 Μ HCl 6 Μ HCl ( a t O X ) 6 Μ HCl (at 25°C)

0.2 0.1 0.2 0.09 0.6 0.12 0.2 0.1 0.1 0.30 0.33 0.45 1.4 0.15 0.81

The possibiHty of using the tetracyano derivatives [Fe(phen)and [Fe(bipy)(CN)4]"2 as redox indicators merits con­ sideration. Their formal redox p o t e n t i a l s t 2 2 - ^ i are approximately 0.2 V lower than those of the corresponding dicyano complexes. One discouraging feature, however, is that the color transitions on oxidation of the tetracyano complexes in acid solution are not particularly distinct. Neverthelesss, their ability to respond reversibly to lower potentials could be used to advantage by employing them as redox indicators in photometric titrations. The redox potentials of a variety of mixed ligand osmium complexes have been measured by Buckingham, Dwyer, and Sargeson.t*^^ The osmium(II) and osmium(III) complexes that were studied contained combinations of the following ligands: pyridine, 2,2'-bipyridine, 2,2',2"-terpyridine, chloride, bromide, iodide, and 2,4-pentanedione. The effects of charge, conjugation, and substitution in the ligands on the potentials were investigated. Standard redox potentials ranged in value from 0.1530 V, for [Os(bipy)2(2,4-pentanedione)](C104), to 0.8847 V, for [Os(bipy)3]( 0 1 0 4 ) 2 . Further study of some of the complexes for possible use as redox indicators appears to be warranted. (CN)4]"2

OXIDATION-REDUCTION

INDICATORS

129

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