CHARLES V. BANKSAND JAMES H. PATTERSON
3062 [CONTRIBUTION
NO. 128
FROM THE INSTITUTE O F
ATOMICRESEARCH AND COLLEGE ]
THE
Vol. 73
DEPARTMEXT OF CHEMISTRY,
I O W A STATE
A Polarographic Investigation of Iron-Sulfosalicylate Complex Ions BY CHARLES V. BANKSAND JAMES H. PATTERSON’ The yellow iron(III>sulfosalicylate complex ion formed in neutral and alkaline solutions was studied by the half-wave potential method. It was found t o contain three triply ionized sulfosalicylate ions for each iron(II1) ion, and have a molar hydrolysis constant of approximately 10-44 at ionic strength 1. No evidence was found for complex ion formation between iron( 11) and sulfosalicylate ions. The second molar dissociation constant, KA2, of sulfosalicylic acid was determined spectrophotometrically to be 3.23 X 10-3 a t 25’ and ionic strength 0.1.
Introduction By applying Job’s method of continuous variations to spectrophotometric data, Foley and Anderson* found that in acid solutions (to pH 2.4) iron(II1) and sulfosalicylic acid react in a 1:l ratio to form a violet complex ion. Above this pH their data seemed to indicate the formation of other complex ions having different mole ratios. Kennard and JohnsonS found evidence for three complex ions having maximum stabilities at pH values of 1.5 (violet), 5.0 (red), and 8.2 (yellow) by studying the PH dependency of the absorption spectra of solutions containing iron(II1) and sulfosalicylic acid. Babko4 found complexes of iron(111) and salicylic acid that correspond well with these pH values and colors. Although the yellow iron(II1)sulfosalicylate complex ion has been used extensively for the determination of iron, its composition and stability have not been carefully studied. Some6 have actually considered the yellow color in alkaline solution to be caused by the dissociation of the complex ion to form colloidal hydrous iron(II1) oxide.
Experimental A Sargent model XXI polarograph was used for nearly all of the polarographic measurements. The cell was similar to the H-type cell described by Lingane and Laitinen* with a saturated calomel half-cell as anode. A large capacity cell of the type described by Meites? was used for the reversibility study. The temperature of the cell was kept a t 25 f 0.2” by a constant temperature bath. The dropping mercury electrode assembly was the conventional type supplied with the polarograph. All solutions were flushed for at least 15 minutes with nitrogen t o remove dissolved oxygen before polarograms were made The spectrophotometric data were obtained with a Beckman DU spectrophotometer with a hydrogen discharge ultraviolet source One centimeter silica cells were used.
Data and Discussion In the present research the half-wave potential method8 was used for the investigation of the yellow (1) Abstracted from a dissertation submitted by James H. Patterson to the Graduate Faculty of Iowa State College in partial fulfillment of the requirements for the degree of Doctor of Philosophy. 1950 (2) R T Foley and R C .4nderson THISJOURNAL 70, 1195 (1948), 72, 5609 (1950) (3) M Kennard and C R Johnson Proc l r a i i s Teras Acad S a , 27, 45 (1944) (4) A K Babko, J Gcn Chem ( V S S R ), 16, 743 (19451, C A , 40, 7042 (1947) (5) S Lacroix and M. Labalade, Anal Chrm A d a , 4, 68 (1950) (6) J J Lingane and H A Laitinen, Ind E W EChcm , Anal E d , 11, 504 (1939). (7) L. Meites THISJot RNA1 , 71, 3269 (1919) J’ his polarographic method has been very adequately described by (a) h.z v Stackelberg and v. Freyhold. 2 Elekfiochcm , 48, 110 (1940), (b) I M Rdthaff aud f . J. Lmgane. “Polarography,” Inter-it‘ience publishers, tnc , Ne# *ark, N. Y , 1941, pp. 161-183. and ( I ) P, Soirchny and J hrchwre, Bull 806, rhtms &dHd, $89 (tDd’l),
iron(II1)-sulfosalicylate complex ion. In Table I are listed the data for the variation of the halfTABLE I VARIATIONOF HALF-WAVE POTENTIALS WITH pH“ &I*, volts Iron(II1) wave
P 13
US
S C.E. Iron(I1) wave
8.38 -0.479 -1.404 8.51 - ,488 -1.417 8.65 - .505 -1.422 8.88 - .540 -1.428 9.09 - .E1 -1.446 9.26 - .561 -1.456 9.48 - .596 -1.463 9.70 .624 -1.483 9.77 - .624 -1.490 10.05 - ,639 -1.502 10.22 - .663 -1.502 4 Solutions were 0.02 M in sulfosalicylate, 0.0001 M in iron(II1) and 0.35 M in total borate. Ionic strength adjusted to 0.6 with sodium perchlorate.
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wave potentials with pH at constant sulfosalicylate and iron(II1) concentrations and constant ionic strength. Table I1 shows the variation of the TABLE I1 VARIATIONOF HALF-WAVE POTENTIAL WITH SULFOSALICYLATE
CONCENTRATION^
ex mole/l.
log Cx
El/,, volts US. S C E. Iron(II1) wave Iron(I1) wave
-0.520 - 1 392 0 00486 -2.313 .568 -1.308 01043 -1.982 - .586 -1,416 -1.689 .0204 - .647 -1.441 -1.302 .0499 - ,764 -1,502 -0.999 .lo02 - .830 -1.536 .2067 -0 685 Solutions were 0.5 M in total borate and 0.001 M in iron(II1). Ionic strength was brought to 1.0 with sodium perchlorate, and PH was 9.0.
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half-wave potentials with sulfosalicylate concentration a t constant pH, ionic strength and iron(111) concentration. All of the solutions were buffered with boric acid-sodium borate buffer, and the ionic strength was kept constant by adjusting the concentration of the sodium perchlorate supporting electrolyte. The reversibility of the iron(II1) and iron(I1) waves was tested by the criterion of T~rneS,~i.e., El/, - E8/, should be 0.056/ n volt a t 25’. For all the iron(II1) waves this value was between 60 and 87 millivolts. Only the waves for the highest PH had values over 80 millivolts, indicating a decrease in the reversibility with increasing pH. Waves for the reduction of iron(1I) to inetallic iraii had values of 34 to 105 f ‘t‘emei, roll.rfron C w r h o d w , Chgm, Cmtmnnr, 9, 81 (lDn??,
POLAROGRAPHY OF IRON~ULFOSALXCYLATE COMPLEX IONS
July, 1951
Potentiol us Soturoted Colomel
Elrctrodc
3063
, volts.
Fig. 1.-Proof of reversibility of iron( 111)-iron(11) couple in alkaline solution: curve 1, polarogram of solution containing iron(II1) ion, but negligible iron(I1) ion; curve 2, polarogram of solution from curve 1 with iron(I1) ion added.
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millivolts for E l / , Ea/, as compared to 28 milliThe partial derivatives of the half-wave potential volts expected for a two electron reduction, thus, of the iron(II1) complex ion wave with respect to this reduction must be irreversible. The reversi- pH and log CX are, respectively, -0.0985 and bility of the iron(II1) wave was also tested by ob- , -0.1660. Although these are somewhat less taining a polarogram of a solution containing both than the values of -0.1182 and -0.1773 theoretiiron(I1) and iron(II1) in the presence of sulfo- cally expected, they indicate that three sulfosalicylsalicylate at pH 9. Air oxidation of the iron(I1) ate ions and two hydroxyl ions are required for the was prevented by the addition of iron(I1) perchlo- formation of the iron(III)-complex ion from FeOH+ rate solution by microburet into the large polaro- with which it is in equilibrium. From KA, = graphic cell containing a solution of iron(II1) 3.23 X lo-* for sulfosalicylic acid and the fact and sulfosalicylate that had been flushed for 30 that Meek12 has shown that the third hydrogen of minutes. Figure 1 shows the polarograms ob- this acid is not appreciably ionized below PH 11, tained before and after addition of iron(I1) solu- it is apparent that -03SCsHs(OH)COO- is the only tion. The smooth curve exhibiting only one point important sulfosalicylate species present in solution of in0ection for the combined cathodic and anodic in the PH range used in this study. Therefore, it waves of the iron(II1)-iron(I1) oxidation-reduction seems probable that the process of forming the system indicates that the reduction of iron(II1) iron(II1)sulfosalicylate complex ion involves the to iron(I1) is essentially reversible.8a removal of the third hydrogen from the participatThe irreversibility of the wave for the reduction ing sulfosalicylate ions thus giving rise to a chelate of the iron(I1) ion to metallic iron makes the data structure like that ascribed by Babko4 to the for this wave useless for study by the half-wave closely analogous iron(II1)-salicylate complex ion. potential method. It was assumed that the sulfoThe reaction for the dissociation of the iron(II1)salicylate had negligible complexing action on iron- sulfosalicylate complex ion may be written (11) ions,l0 and that the principal iron(I1) species Fe [03SGHa(O)COO]3-6 + 3H10 + in these solutions is FeOHf. According to LindFef3 4-3[03SCeHs(OH)C00]-2 + 30H- (1) strand" the constant for the formation of this ion That three equivalents of hydroxyl ion are reby the hydrolysis of ferrous ion is 1.2 X a t 20'. quired for the formation of iron(III)-sulfosalicylate (IO) A. Thiel and 0. Peter, Z. anal. Chcm., 103, 161 (1935). (111 P, liadatroacl, Surfirk Kern. T i d . , 60, 181 (1044).
112) H. V, Meek, Bewtaral Dissertation, I e n o Otots College, 1860.
3064
CHARLES V. BANKSAND
complex ion from iron(II1) ion was confirmed by titrating a mixture of bis-(2,4-pentanediono)-iron(111) and sulfosalicylic acid (100% excess) with sodium hydroxide solution. The vertical line in Fig. 2 is the calculated equivalence point for the reaction
+
+
Fe(CsH,O& GHO~SC~HS(OH)COOH 12NaOH + NasFe [OaSCeH~(O)C00]3 3C6Hk02 3Naz[OaSCeH3(OH)C00] (2)
+
+
JAMES
VOl. 73
H. PATTERSON
Lindstrand“ Kl? is estimated to be about In these K’ a t 2 5 O , so that K : is then values the activities of the hydroxyl ions and the molar concentrations of the other species are used. The second molar dissociation constant, KL2,of sulfosalicylic acid was determined by a spectrophotometric method13 based on the fact that the molar absorbancy indexes, U M ~and UM,, for -03SCsHa(0H)COOH and -O~SCBH~(OH)COOare appreciably different a t 317 mpL.12Figure 3 shows the results of this investigation. Under the experimental conditions, the values for U M , and uhrz were found to be, respectively, 1568 and 708. ____I__-
ILi,--
I
- -
--
.
- + \L
\
i 81
Fig. 2.-Titration of 0.316 millimole of bis-(2,4-pentanediono)-iron( 111) plus 1.898 millimoles of sulfosalicylic acid.
The molar equilibrium constant, K:, for reaction (1) may be obtained from the approximate equation ( E I / ~-) ~ (El/*). = (0.0591) log K:/K[ (0.1773) log Cx - 0.1182(fiH
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14) (3)
where (Bl/Je - (El/& is the difference of the half-wave potentials of the iron(II1)~ulfosalicylate complex ion, Fe [OSSC~HS(O)COO]~~, and the simple iron(II1) ion, and KI is the instability constant of FeOHf. The half-wave potential for the reduction of the simple iron(II1) ion is a t f i s t approximation equal to the standard reduction potential (0.51 volt us. S.C.E.) of this ion. From this reduction potential and the polarographic data in Tables I and I1 an average value of -33 is obtained for log K:/Ki. From the data of
Fig. 3.-Absorbancy of sulfosalicylic acid with varying PH: sulfosalicylic acid, 6 X lod4 M , p = 0.10, 0.1 AI HClOd adjusted t o desired $H with NaOH, 25’. Absorbancies measured on Beckman DU spectrophotometer, A = 317.0 mp, band width 0.97 mp: 0, observed points; calculated points for K ’ A = ~ 3.23 X lO-a(fiK’~,= 2.49); X, points calculated for K’A, = 3.23 X f, points calculated for R’A, = 3.23 X U H + [-03SCeHa(OH)C00-] K’A2 [-OaSCeHs( OH)COOH]
a,
Acknowledgments.-The authors wish to thank Mr. Richard E. Ewing and Miss Doris V. Stage for their translation of the very extensive article by Souchay and Faucherre from the French language.& Acknowledgment is also made to the Ames Laboratory of the Atomic Energy Commission for underwriting the expenses for this investigation, and providing the necessary facilities for the conduct of the research. AMES,IOWA RECEIVED OCTOBER 9, 1950 (13) C.E. Crouthamel, H. V. Meek, D. S. M i r t h and C. V. Banks, THISJOURNAL. 71, 3031 (1949).
