1996
NOBUYUKI T'4NAK11, I. hf. KOLTHOFF AND [CONTRIBUTION FROM
THE SCHOOL OF
CHEMISTRY O F
THE
lv. STRICKS
UNIVERSITY
OF
Vol. 77
MINNESOTA]
Iron-Cysteinate Complexes BY NOBUYUKI T A N A K A I 1 I. ht. KOLTHOFF AND
w. STRICKS
RECEIVED OCTOBER 25, 1954 In solutions of a pH between 10 and 11.2, two complexes composed of iron(I1) and cysteine are found: Fe(OH)(RS)and Fe(RS)z', in which RS denotes -OOC.CH(XHz) .CH,S-. The dissociation constants, K F e ( O H j ( R 8 j - = (UF-+- X UOH- x ~ R S - ) / U F ~ O H R R - and K F e ( R R j 2 = upe++ X a 2 ~ g ' / a ~ e ( ~ ~ )were nestimated from solubility data to be equal to 1.7 x and 1.7 X 10-'2, respectively, a t 25'. In weakly acid medium (pH 5 -7) a slightly soluble compound, FeRS, is formed with a solubility product, S F ~ R (= S u p e + + X U R S - ) , of 2.6 X 10-11 a t 25". Ferric iron catalyzes the air-oxidation of cysteine to cystine. From solubility data of ferric hydroxide in ammoniacal cysteinate solutions and from spectrophotometric measurements in ammoniacal Versenate-cysteinate solutions, both extrapolated to zero time when no reduction of ferric cysteinate has occurred, it was postulated that two complexes, FeOH(RS)zZ-and Fe(RS)33-, exist a t pH 10 to 11, with the dissociation ( = ( U F ~ + + + x UOH- x U 2 ~ ~ - ) / U ~ e ~ ~ (Of~ 5~ )xr P - ) and K F e ( R S j 2 ' - ( = UFe** x a 3 R S - / a F e ( K S j ~ J - ) COnStantS KF~oH(RS)~'of 8 X at 25'.
Ferric iron when added to a n acid solution of cysCHS teine hydrochloride produces a deep indigo blue species can exist': 1 \Fe, (FeTS) in acid coo/ color which disappears in a short time and cannot CH$-FeOH, (FeOHTS-) and he restored by ~ x y g e n . ~Harris4 .~ states that the violet coloration obtained when ammonia is added medium, and 1 to a cysteine solution is due to the presence of cootraces of metal ions (generally ferric iron). He also CH2S-FeSCH2, mentions that ferrous iron yields no colored comI (Fe(TS)2') in alkaline medium. cooplex with cysteine. On the other hand, Schubert5 cooreports that ferrous iron forms two kinds of com- Considering the dipolar nature of cysteine the folpounds with cysteine, one of which is formed in lowing seven species with ferrous iron map exist acid medium (pH 4-6) and has the formula FeCHzS CHSS (OOCCHNH2CHS) .1.5&0. Upon mixing two I \ I \ CHXH3+ Fe and CH3YH2 Fe, corresponding t o FeTS moles of cysteine hydrochloride, one mole of ferrous / I / salt and six moles of potassium hydroxide, he obcoo coo served an orange color indicating the formation of a I I1 soluble ferrous cysteinate complex.6 This complex CH?S-FeOH CI&S-FeOH may have the formula Fe(SCHZNH2CHC00)2-, I I corresponding to the formula of ferrous complex CHiYH3+ and CHNH? , corresponding to FcOHTSwith thioglycolic acid formed in alkaline m e d i ~ m . ~ ,I ~ I coo In alkaline medium, the ferrous complex is eas- coo I11 IT' ily oxidized to a ferric complex by molecular oxyCHzS-Fe-SCH2 gen.3 According to Schubert5 a violet color which CHzS-Fe-SCHz I / I / develops when a trace of ferric iron is added to an CHX"3+ CHh'H?+, CHNH3+ C H S H ? and oxygen-containing solution of cysteine of pH 8 to 9 I / I / fades and disappears when the dissolved oxygen is coo- coocoo- cooused up. The cysteine is oxidized to cystine. V VI ll'hen cysteine is left in the solution, the color is CHZS-Fe-SCH? I / regenerated by shaking with air. I n order to exCHSHz CIISH?, plain the catalytic oxidation of cysteine Schubert I / postulated that the violet ferric complex which is coo- cooformed on oxidation of the ferrous complex is re1-1I duced to the ferrous complex with the formation of corresponding to Fe(TS)2=. The species I, 11, 111, the disulfide and ferrous hydroxide. IV, V, VI and VI1 will be denoted in this paper as Oxidation FeRS+, FeRS, FeOHRS*, FeOHRS-, Fe(RS)Z**, Fc(RS)? {FeOH(RS)z)2 + Fe(RS)Z RSSR 4-Fe(0H)z Fe(RS)' and Fe(RS)', respectively. For the forIn the present paper we report the results of a mation of one mole of the above species from unquantitative study of the composition and the sta- charged cysteine and ferrous ion the following bility of complexes between cysteine with ferrous number of moles of sodium hydroxide is required: and ferric iron. The formulas used refer to composi- for (I) 1 mole, for (11), (111) and (V) 2 moles, for (IV) and (VI) 3 moles and for (VII) 4 moles. Evition and do not reveal structural information. Ferrous-Cysteinate Complexes.-In the ferrous dently, the system is very complex. Experimental iron-thioglycolic acid system, the following three conditions were selected which allowed reasonable conclusions regarding the kind of complexes pres(1) On leave of absence from Tokyo University, Japan. ent in the solutions. (2) E. Baumann, 2.physiol. Ckem., 8, 279 (1883-1884). (3) (a) A. P. Mathews and S. Walkers, J . B i d . Chcm., 6 , 21 (1906); Experimental (b) I, Michaelis and E. S. G. Barron, i b i d . , 83, 191 (1928). ____f
+
(4) L. J. Harris, Biochcm. J . , 16, 739 (1922). (5) M. Schuhert. THIS J O U R N A L , S4, 4077 (1932). (6) M. Schubert, i b i d . , SS, 4563 (1933).
(7) D. L. Leussing and I. M. Kalthoff, ;bid., 16,3904 (1963).
+
Cysteine hydrochloride ( C.P. grade by Pfanstiel Chemical Co.) was used. Stock solutions of 0.5 M ferrous chloride and of sodium hydroxide were prepared in the same way as described in previous papers.718
1997
IRON-CYSTEINATE COMPLEXES
April 5 , 1955
TABLE I PRODUCT OF ANALYSESOF MIXTURES CONTAINING FERROUS CHLORIDE, CYSTEINEAND SODIUMHYDROXIDE.SOLUBILITY FERROUS CYSTEINATE (FeRS) Volume of mixture 100 mt. if not stated otherwise. Initial concn. in mixture Fe, Cysteine, mmoles mmoles
PH
Supernatant s o h Fe, Cysteine, M X 10' M X 10'
Fe, mmoles
product Solubility of ferrous
Precipitate Cysteine, mmoles
29.4 2.11 10.04 5.11 5.83 79.3 77.2 2.22 5.02 10.00 5.84 28.0 24.9 2.18 5.00b 25.7 4.99 6.06 5 . OOb 49.0 4.67 2.94 9.99 6.45 1.6 4.78 2.20 5.00 5.00 7.40 34.6 3.00 45.3 7.53 8.02 5.03 (51.1)" 0.24 47.8 5.02 5.04 8.90 No precipitate 5.00' 10.03 9.80 27.1 50.0 2.31 5.02 5.00 10.05 (49.2)' 2.31 26.9 5.00 10.07 4.94 74.6 0.77 42.3 5.00 10.18 7.46 41.6 75.3 0.86 5.02 7.53 10.38 50.0 2.72 5.02 23.0 5.00 10.69 5.02 21.3 50.2 2.89 5.02 11.13 5. 02b (69 5)" 1.68 30.4 11.15 7.55 5.02 49.5 3.01 20.1 4.95 11.20 Figures in brac ?ts have not been ma.-: use of in numerical ci ulations. mixture 98 ml. Corrected for the formation of FeOHRS- and Fe(RS)2'.
cysteinate (FeRS)
2 . 4 X lo-" 2 . 6 X lo-" 2.4 X 3 . 1 X lo-" (4.8 X
2.17 2.28 2.25 4.60 4.84 1.57 (-0.07)"
...... (0.02)"
......
...... ......
......
( -0.10y
......
* Volume of
mixture 110 ml.
Volume of
The ferrous iron-cysteine system was studied in air-free contained a precipitate of ferrous hydroxide and media in which the amounts of iron, cysteine and sodium therefore the ferrous ion activity in these solutions hydroxide were varied. For the preparation of each mixture, the same procedure was used as in the studies of fer- could be calculated. Although five different ironrous-thioglycolate complexes.' All bottles were rotated for cysteine complexes might exist in this PH range 5 days or longer and centrifuged when a precipitate had stoichiometric considerations indicate that the speformed. Samples of the supernatant liquid were with- cies FeOHRS- and Fe(RS)t'prevail. Combinations drawn by means of a syringe for the determination of PH, of other possible complexes were considered and iron and cysteine. The determination of PH was carried out in a nitrogen sets of equations were developed. None of them atmosphere using the Beckman Model H 2 fiH meter. A satisfied the experimental results. glass electrode 1190-80 ("for general purpose") was used. The following equations can be set up. The total iron analysis of the supernatant liquid was carried out by the same procedure as described previouslyP Total [Felt = [Fef+] f [FeOH+] [FeOHRS-] f cysteine was determined polarographically by measuring [Fe(RS)z'I (1) the anodic diffusion current a t -0.48 v. vs. S.C.E. in oxygen-free 0.1 M sodium hydroxide so1ution.Q In blank ex[RSHIt = [NHa'RSHCOO-] f [NHzRSHCOO-] f periments with sodium hydroxide solutions, 0.001 M in [iYH3+RS-COO-l f [NHzRS-COO-] cysteine and 0.0005 to 0.002 M in ferrous iron, it was es[FeOHRS-] f 2[Fe(RS)z'] (2) tablished that the ferrous iron does not interfere with the polarographic determination. [OH], = [OH-] [FeOH+] f [NHzRSHCOO-] A summary of the details and the results of the experi[SHa+RS-COO-] f ~[NHzRS-COO-]f ments is given in Table I. In the lower pH region a slightly 3[FeOHRS-] f 4[Fe(RS)z'] (3) colored (brownish-yellow) solution was obtained. Upon standing a brownish precipitate appeared within 24 hours, where the quantities in brackets on the right-hand while the solution became almost colorless. From the difference in the iron and cysteine contents present initially side are the equilibrium concentrations while and from those found in the supernatant solution the ratio [Felt, [RSH]t and [OHIt are the total concentraof ferrous iron and cysteine in the precipitate is calculated tions of iron, cysteine and sodium hydroxide in the to 1: 1. This compound can therefore be presented by the supernatant liquid. [Felt and [RSHIt were deterformulas FeRS as reported by Schubert.6 A t a OH of about 8 a brownish-yellow colored solution mined analytically and [OHIt was obtained by the was obtained. On standing a greenish-white precipitate equation : separated, which appeared to be a mixture of ferrous cysteinate (FeRS) (mmole of 9 a O H added) - 2 X (mmole of ferrous hydroxide) (mole,l,) E ml. of total vol. of supernatant liq. and ferrous hydroxide. The brownish color in the equilibrium solution At pH higher than 10 [Fe++], [FeOH+] and indicates the presence of a ferrous-cysteinate complex. In the higher pH region (9-11.2) complexes were clearly [RSH*] are negligibly small. Introducing the present in the solutions. The iron concentration in the equilibrium solutions was much greater than that calculated dissociation constants Of the various charge from the solubility product of ferrous hydroxide.8 The pre- species of cysteine, KA, Kg, Kc and KD (Cf. cipitate was found to consist only of ferrous hydroxide. ref. 10) The color was white a t f i s t , then gradually turned to greenish-white. "3' 3" +
+
+
+
+
Calculation and Discussion Almost all equilibrium mixtures of pH 9 or greater (8) D. L. Leussing and I. M. Kolthoff, Tms J O U R N A L , 76, 2476
I
RSH
I
coo-
2 ASI
+H+
KA
coo -
(1963). (9) I . M. Kolthrrff and C. Barnum, ibid., 63,3061 (1940).
(IO) W.Stricks and I. hl. Kolthnff, ibid., 73, 4560 (1951).
1998
NOBUYUKI TANAKA, I. M. KOLTHOFF AND NH3
lTH2
I
RSH
J_
+ H'
ASH
coo-
coo -
?Ha
THq
i-
I
J_
RS-
I
coo -
RSI
TABLE I1 CALCULATED CONCEXTRATIONS OF SPECIES AND DISSOCIATIOS CONSTANTS OF THE FERROUS-CYSTEINE COMPLEXES PRESENT I N ALKALISE SOLUTIONS(PH 10.69 TO 11.20)
Kc
coo -
NH,
[RS-I,
.\I x 103
YH,
+ H-
RSH
R5-
coo -
coo -
I
KD
+ RS=,
and using the simplified symbols RSH', RSH- and RS= for the various charge species of cysteine, rearrangement of equations 1, 2 and 3 yields [RS-] = [OH]$ - 2[FeIt - [RSHIt [FeOHRS-]
=
2[Fe], - [RSHIt
-
[OH-]
(4)
+
+
The dissociation constants of the two ferrous cysteinate complexes are I'eOHRS-
FeTr
+ OH-
-I-RS=; KF~OHRS= X aRS(7) aFeOHRSK F ~ ( R S= )~-
f l ~ ~ +X+ OH-
I'e(RS)?=
If F e + -
+ 2RS',
flFe*
x
VOl. 77
equations cannot be applied when this term yields a negative value or a value close to zero. The results of the calculations are given in Table 11.
KB
+H+
Wr.STRICKS
U'RS-
UF~RS)~-
(8)
[FeOHRS-I, .If x 103
8.9 17.3 19.4 18.4 av. values of
[Fe(RS)z-I,
M
xi03
KF~(RS)Z-
KF~OHRS-
9.5 13.5 8.4 12.9 15.2 15.2 8.0 13.1 the constants are:
2.8 X 1.4 X 1.3 X 1.4 X
2.3 X 1.9 X 1.1 X 1.6 X
lO-'3 lO-'3
lo-'' lo-'' lo-'* lo-''
K F ~ O H R=S 1.7 - X 10-13 (4.2 X K F ~ ( R S=) ~ 1.7 - X lo-'' (1.1 X lo-")
The values between parentheses are the corresponding dissociation constants of ferrous-thioglycolate complexes which have been recalculated from data given by Leussing and Kolthoff.' The stabilities of ferrous-cysteinate and of ferrous-thioglycolate complexes are of the same order of magnitude, indicating that the uncharged amino group has little effect on the stability of the iron cysteine complex. In calculation of dissociation constants the results of experiments with pH 8.9 to 10.38 were not used. However, if the complexes FeOHRS- and Fe(RS)2' are the predominant species in these systems, the total iron concentration calculated from the above dissociation constants (obtained from experiments with pH 10.69 to 11.2) must agree with the total analytical iron concentration. In a test of this relation equation 2 was used. Considering equations 7' and s',equation 2 can be rearranged as
Since the solutions are in equilibrium with solid ferrous hydroxide equations 7 and 8 can be written and
where S F ~ ( is ~ the H )solubility ~ product of ferrous hydroxide. In calculating the constants, values of activity coefficients for mono- and divalent ions a t an ionic strength less than 0.1 were estimated from Kielland's tables." At ionic strengths greater than 0.1 mean activity coefficients were used from other data.l2>l3 Concentrations of various cysteinate anions were calculated from the constantslo a t 25", using values of ~ K A~, K BpKc , and ~ K ofD8.66, 8.60,10.45 and 10.51, respectively. A value of 3.5 X was taken for the solubility product of "green" ferrous hydroxide.7*a The expressions for [REF], [FeOHRS-] and [Fe(RS)2'] in eq. 4, 5 and 6 involve the term, ([OHIt - 2[FeIt - [RSHIt - [OH-]). These (11) J. Kielland, THISJOURNAL, SS, 1675 (1937). (13) H . S . Harned a n d B . B. Owen, "The Physical Chemistry of Electrolytic Solutions," 2nd ed., Reinhold Publ. Corp., N e w York, N. Y . , 1950. (13) As for t h e procedure of evaluation: cf. H. Borsook, E. L. Ellis n n d €I. M. Huffman, J . R i d . Chem., 117, 281 (1937).
where y1 and y z represent the activity coefficients of mono- and di-valent ions, respectively, and YRS-
X a ~X + SF~(OH)~ Kw X KF~OHRS-
YF~OHRS-X
c y =
8 =
Y2RS-
x
Y F ~ ( R S - ) ~X
a'H+
x
SFe(OH)z
K2, X
KFe(Rsjr-
Solving the equations, the total iron concentration [Felt can be obtained as [Felt = [FeOHRS-]
+ [Fe(RS)?=]= a[RS=] + P[RS']' (10)
Total iron concentrations calculated from four experiments are given in Table 111, and compared with the experimental values. TABLE 111 OBSERVEDASD CALCULATED VALUESOF TOTAL IRONCONCESTRATION IN SUPERNATANT LIQUID
PH
Observed[RSHIt xU 103, .
10 05 10 07 10 18 1 0 38
50 19 74 75
r-
--
[Felt 103. If
x
5 0
27 1 26 9 42 3
.'3
31 6
0
[FeOHRS-I x 103.
M
14 13 18 17
0
7 1 8
Calculated---[FeIRS1-21 x 103.
41
14 14 24 23
8 2 2 6
Fel;
x tf10, 28 27 42 41
8 9 3
4
IRON-CYSTEINATE COMPLEXES
April 5, 1955
1999
Measurements were made in ammoniacal buffers of various p H and in the presence of a large excess of cysteine. Suitable amounts of ammonium hydroxide, ammonium nitrate and cysteine hydrochloride solutions were placed in the absorption cell and oxygen was removed by bubbling for 40 to 60 minutes with nitrogen which previously had been washed with a buffer of the same concentration of ammonia and ammonium nitrate as used in the experiment. The cell was immersed in a thermostat of 25" and an air-free ferric ammonium sulfate solution was injected with the aid of a syringe. The absorption of the solution was measured as a function of time. A cell filled with distilled water was used as reference. The pH of the solutions was S F ~ R B (a~~ )e + +X REcalculated from the activities of ammonia and amThe last column of Table I gives the values of the monium ion and determined experimentally with a solubility product calculated from experimental Beckman Model H 2 pH meter after the spectrodata, assuming the activity coefficient of RSH* to photometric measurement. be unity. A set of involved equations was derived Reciprocals of optical density plotted against which allowed by a series of approximations the time yielded straight lines as shown in Fig. 1, which calculation of concentrations of FeOHRS - and indicates that the disappearance of ferric cysteinFe(RS)2-. These concentrations increase with in- ate is a second-order reaction. Under identical concreasing pH but are negligibly small in the calcula- ditions solutions of different concentrations in ferric tion of the solubility product up to a p H of 7.4. ion gave approximately the same rate of disappearEven a t this p H the value of the constant corrected ance. The rate at a given concentration of cysteine for these concentrations hardly differs from the became smaller with increasing pH and a t constant uncorrected value (4.8 X lo-" as compared to 5.4 p H with increasing concentration of cysteine. The x lo-"). From the data of experiments with p H experimental data and the results of calculations 5.83 to 6.45 an average value of S F e R S ( s ) of 2.6 X are summarized in Table IV. It is interesting to is calculated. From the data of Leussing and note that the product of the rate constant and the Kolthoff a solubility product of ferrous thioglyco- concentration of uncombined NHzRS-COO- (last late of 1.6 X is calculated. column of Table IV) is practically constant. This At the intermediate PH of 8.02 (see Table I) the indicates that the rate determining step in the solution is in equilibrium with the two solids, fer- oxidation of cysteine by ferric iron is not the direct rous hydroxide and ferrous cysteinate. From the 9.0 I following equilibria I
The calculated values of total iron concentrations are in good agreement with the observed ones, indicating that in the PH region from 10.0 to 10.4 two species, FeOHRS- and Fe(RS)Z', are predominant. However, when the same procedure was applied to expts. with p H 8.02 and 8.90, the calculated values of [Felt and [OHlt were greatly different from the experimental ones, indicating that under the particular experimental conditions another complex must be present in addition to FeOHRSand Fe(RS)a. I n the lower pH range (5.8 to 7.4) a precipitate is present composed of FeRS with a solubility product
I
I
I
I
4
6
8
10
I
+ 20H+ RS'
1
Fe++ FeRS I_ F e + +
Fe(OH)*
it is seen that a o A - / a R s - must be constant. From ( ~ ) "green" the values of S F e R S ( s ) and of S F ~ ( O H ) ~for ferrous hydroxide,a aoA - / a R s - is calculated to be 1.35 X 10-5.14 Ferric-Cysteinate Complexes. Rate of Reduction of the Ferric-Cysteinate Complex.-Ferriccysteinate is formed upon addition of ferric iron to either acid or alkaline cysteinate solutions, the complex being reduced to the ferrous complex with formation of cystine. The rate of reduction of the ferric-cysteinate complex was investigated spectrophotometrically by following the rate of disappearance of the violet color of the complex. A Beckman Spectrophotometer Model DU was used with absorption cells designed to keep the solution free from oxygen during the m e a s ~ r e m e n t .Length ~ of the light path was calibrated for each cell. The absorption curve cannot be measured accurately because the violet color of the ferric-cysteinate complex disappears rapidly in air-free solutions. In air-saturated solutions the absorption maximum was found a t approximately 580 mp. All measurements were performed a t that wave length. (14) When ferrous hydroxide alone is in equilibrium with the solution a&-/oRsmust be greater t h a n 1.35 X 10-6, while, when ferrous cvsteinate (FeRS) alone is in eauilibriurn with t h e solution. the ratio milst be smaller than I.%; X 10-5.
~
z
TIME,
12
14
MINUTES.
Fig. 1.-Disappearance of ferric-cysteinate complex in air-free ammoniacal buffers in the presence of a large excess of cysteine. (See Table I V for experimental conditions; the numbers in the figure correspond to the sequence of the experiments listed in Table TV.)
2000
NOBUYLJKI TANAKA, I. AM. KOLTHOPF AND
w.STRICKS
vO1. 77
TABLE IV SPECTROPHOTOMETRIC MEASUREMENT OF RATEOF DISAPPEARANCE OF FERRIC-CYSTEINATE IN AIR-FREE AMMONIA BUFFERS (["a] = 1.01 to 0.91 M,["I+] = 0.22 to 0.12 M in final solution). Initial concn. Fe(NH4)
RSH-HCI, M
(so,)~
M
x
10'
Ji
PH
Optical density att = 0
Molar ext. coeff. a t t = 0, mole-1 cm.-l
R a t e of decrease in concn. of ferriccysteinate complex k, [RS-Iuncomb. k X mole-1 sec.-l M X 102 [ R S ' l u u c ~ m b
0.0100 1 .oo 0.13 10.33 0.305 3.05 x 103 28.0 4.26 0.12 .0100 2.00 .13 10.35 .600 3.00 x 103 27.5 4.18 .12 ,0200 1.00 ,156 10.21 ,305 3 . 0 5 x 103 18.4 7.96 .15 .0200 2.00 .156 10.22 ,606 3.03 x 103 16.8 7.S8 .13 .(I500 2.00 .233 10.03 .617 3 . 0 8 x 103 10.9 16.3 .18 .0500 2.00 .143 10.31 ,617 .18 7.94 23.3 3.08 x 103 .0500 4.00 ,143 10.30 1.220 .14 6.46 22.1 3 . 0 5 x 103 Concentration of Fe(NH4)(S04)zadded corresponds to the concentration of ferric-cysteinate a t t = 0 ([Fe(111) cornplex]~) .
interaction between ferric ion and the amino acid. The molar extinction coefficient of the ferric-cysteinate complex was found by extrapolation to zero time and estimated to be 3.05 X lo3 mole-' ern.-' a t 580 mp (see Table IV). Solubility of Ferric Iron in Ammoniacal Cysteinate Solutions.-When ferric iron is added in excess to an ammoniacal cysteinate solution a brown precipitate is obtained. Although unlikely, the formation of an insoluble ferric cysteinate complex cannot be excluded.'j Under our experimental conditions the precipitate appeared to be composed of hydrous ferric oxide which was free of cysteine. The anodic diffusion current of air-free A4 cysteine solutions was measured in ammonia buffers before and after addition of enough air-free ferric ammonium sulfate solutions to give a precipitate. Extrapolated to zero time the diffusion current was not affected by the addition of iron. The anodic current decreased on standing while a cathodic current of cystine appeared, the decrease of the cysteine current being almost equal to the diffusion current of the cystine formed. The solubility of ferric hydroxide was determined in air-saturated and air-free buffers containing a relatively large concentration of cysteine. I n air-saturated solutions the iron remained in the form of ferric because of the rapid reoxidation of the ferrous complexes by oxygen. In air-free solutions, the ferrous complexes formed were not reoxidized. The rate of reduction of the ferric complex is so great that solubility can be found only by extrapolation of the ferric concentration to zero time in both series of experiments. The procedure was to add a slight excess of a ferric ammonium sulfate solution (so that the solubility of Fe(OH)3 would be exceeded) to a solution of a known concentration of cysteine in a known volume of an ammonia-ammonium nitrate buffer. The solution was then analyzed for iron as a function of time by suitable means. Air-saturated Solutions.-As the cysteine is oxdized, the solubility of iron decreases and more ferric hydroxide precipitates. Cysteine hydrochloride was weighed into dry bottles of 250-ml. or 500-ml. capacity. Known volumes of water and of standard solutions of arnnionia and ammonium nitrate were then placed into the bottles. Air, previously washed with a buffer of the same concentration of ammonia and ammonium ( l i ) 1 V i h l i s k y .rnd
V Siiidelor, Mikrochim A c t a , 3, 228 (1038)
nitrate as used in the experiment, was bubbled vigorously through the solution and finally the desired volume of standardized 0.1 Mferric ammonium sulfate in 0.05 1cI sulfuric acid was added rapidly. This is taken as zero time. Samples of the solutions were removed a t various periods of time and filtered through glass wool contained in a syringe of 20- or 50-ml. ~ a p a c i t y . ~Known volumes of filtrate were analyzed volumetrically for iron.' At pH less than 10.6 the iron concentration versus time plot gave a straight line yielding the iron concentration a t zero time. Experimental data and values of the calculated constants are listed in Table V. From Fig. 2 (airsaturated solutions) it is seen that the iron concentration decreased linearly with time in solutions of pH 10 to 10.5. This is a t variance with results obtained in the thioglycolic acid-ferric iron system7 which gave straight line plots of the reciprocal of the iron concentration versus time in the p H range 9-10. The lines (Fig. 2) presenting experiments 1 to 3 (at various initial cysteine concentrations) are parallel, thus indicating that the rate of disappearance of iron in air-saturated solutions is independent of the cysteine concentration. Air-free Solutions.-Difficulties were encountered in analyzing the solutions. The spectrophotometric method cannot be used in the presence of a precipitate of ferric hydroxide. The polarographic method, applied in an analogous study of thi~glycolate,~ gave results which were not easily interpreted when dealing with the cysteine system. The reason is that the ferrous iron formed by reduction of the ferric complex catalyzes the polarographic reduction of cystine at a potentia! where the ferric complex yields a diffusion current while cystine alone (absence of Fe(I1)) is not reduced. This pecuIiar catalytic polarographic effect will be discussed in a subsequent paper. It was decided to extrapolate the total diffusion current measured a t -1.35 v. (at lower PH) or -1.40 v. (at higher pH) to a reaction time of zero. The measured diffusion current is the sum of that of the ferric-cysteinate complex and that of the cystine formed on autodecomposition of the coinplex. The value extrapolated to zero time should be that of the ferric complex. In separate experiments with a large excess of cysteine and in the absence of a precipitate of ferric hydroxide it was found by extrapolation to zero time that the tliffusion curreii t of ferric cysteinate is 5(,6 sinallcr
IRON-CYSTEINATE COMPLEXES
April 5 , 1955
2001
TABLE V DISSOCIATION CONSTANTS OF FERRIC-CYSTEINATE COMPLEXES IN AIR-SATURATED AND AIR-FREESOLUTIONS
Exp. no.
RSH.HCIa added, M x 103
1 2 3 4 5 6
50.06 40.25 30.15 40.08 40.06 40.13
-
Corn osition of mixture at t 0 [NHaY, ["I+], fiH
[Fe '1, M X 10' + +
M
0.93 0.94 0.94 1.44 1.90 1.96
16.52 12.22 8.50 12.10 11.60 10.77
M
calcd.
-
nr
Dissociationb constant of Fe(OH)m(RS).(m +*n - 3 ) 0.
m-I,
n=2
n = 2 K1.n
KO,:
Ir
In air-saturated solutions 0.30 10.00 0.40 .28 10.03 .37 .26 10.05 .35 .28 10.21 .37 .28 10.33 .37 10.53 .27 .18
x
x
1030
8.4 9.0 8.4 4.5 3.2 1.7
load
2.0 2.3 2.3 1.8 1.7 1.4
In air-free solutions 0.97 2.75 0.21 10.10 0.24 12.50 6.6 2.0 .06 10.63 .08 1.2 .97 2.30 1.4 12.50 .06 1.8 .97 2.92 10.63 .08 1.5 12.50 .04 .98 10.75 .07 1.5 1.5 0.56 6.25 .04 10.75 .07 1.6 1.6 .98 0.55 6.25 10.18 .18 .16 5.5 .96 2.70 1.9 12.50 .16 .96 1.87 10.19 .18 1.7 4.9 9.38 10.19 .18 .16 1.7 4.9 .96 1.43 7.81 .97 1.02 10.21 .18 .15 1.7 4.6 6.25 .14 .97 0.29 10.25 .17 2.1 5.2 3.13 concentrations of Fe(II1) were 20.06 X M in exp. 1-3, 20.02 X M in exp. 4-5, b K,.o = aFe* X POHX a%exp. 6, 6.25 X M in exp. 7-15 and 3.13 X M in exp. 16. 7 8 9 10 11 12 13 14 15 16 Initial
m=2, n = l
x
m=O,
n - 3
K1.i
x
108'
Ko,: 1038
1.4 1.6 1.7 1.5 1.3 1.1
0.0058 1.1 3.4 3.0 0.84 3.0
2.4 1.8 2.1 3.7 3.8 2.3 2.5 2.8 3.2 5.9 19.95
7.3 4.4 7.5 7.0 7.3 7.2 6.4 6.8 6.7 7.6
x
10-8 A f in
~ F ~ ( O H ) ~ R B ~ ( ' " + ~ - * ') -
I
I
time are given. It is gratifying that the values of the constants calculated in Table VI11 on the basis of these results are in close agreement with those obtained in air-saturated solutions.
I--=
10
= = = = = = TIME ,MINUTES.
Fig. 3.-Current-time curves measured a t (7), - 1.35 volt us. S.C.E.; (8) and (lo), -1.40 volt, in contact with ferric hydroxide. (See Table V for experimental details.)
Dissociation Constants of Complexes.-The dissociation constant of the ferric-cysteinate complex T I ME
, MINUTES.
Fig. 2.-Decrease of iron( 111)-concentration in airsaturated ammoniacal cysteinate solutions. (See Table V for experiniental conditions.)
than that of an equivalent solution of cystine. In Fig. 3 some examples of the extrapolation to zero
corresponding to the Fe(OH),(RS),(m+zn-3)-
Fc+++
+ mOH- + nRS'
was calculated, using different values of m and n (see Table V). Introducing the solubility product
NOBUYUKI TANAKA, I. M. KOLTHOFF AND W. STRICKS
2002
VOl. 77
TABLE VI SPECTROPHOTOMETRIC MEASUREMENTS OF FERRIC-CYSTEINATE IN AIR-FREE A 4 SOLUTIOXS ~ \VHICH \YERE ~ INITIALLY 0.97 MIS "4OH, 0 10 hf I N h"4XO3, 2 x 10- .If I N VERSESATE ( Sa!H2Y), .If I T Fe(III) A V D O F vAR\'IUG CYSTEINE CONCENTRATIONS CALCULATION OF THE DISSOCIATIOX COKSTAST K F L ( R ~ ) ~ ~ RSH.HCI added, M x 103
iJH
w
11.92 9.93 7.94 6.68 3.34
10.21 10.24 10.26 10.24 10.29
0.15 .14 .14
.I4 .13
Optical density at 580 mw at t 0 (cor.)
[Fey(OH)r*-l; M x 10
[Fe(RS)aa-], .M x 1 0 6
0.204 ,128 ,072 ,0516 ,0160
0.93 .96 .98 .98 .99
6.69 4.20 2.36 1.69 0.525
-
[Yr-], x 103
M
0.x5 .507 ,510 ,497 517
[Fe++-],
.M x
1029
5.8 5.2 4.8 5.4 4.2
[RS-1, M x 103
4.30 3.78 3.12 2.56 1.36
KF~(ltS13~-
x
1033
2.2 2.2 2.0 (1.8) (0.7)
of ferric hydroxide,I6 S F ~ ( O H ) ~ of ( ~ )6 X when the solution is in equilibrium with solid ferric hydroxide] gives
Since the apparent dissociation constants of ferric Versenate complexes are known, l 7 it is possible to calculate the dissociation constant of the ferric-cysteinate complex if the concentration of the latter can be determined in the mixture of both complexes. The concentration of the ferric-cysteinate comIntroducing the concentration of total uncom- plex was determined by measurement of the extincbined cysteine ([RSH],,co,~.), the ionization con- tion of the violet color a t 580 m p . The values were stants of cysteinate ions ( K c and KD'O)and the ac- extrapolated to zero time. The experimental details tivity coefficients yields and results are given in Table VI. The optical density data a t zero time in Table VI are corrected for the very small absorption of the ferric Versenate complex. The molar extinction coefficient of the ferric cysteinate complex which forms in ammoniacal buffer of p H 10.3 in the presence of a large excess of cys[RSH],,,,,~, = [RSI-IItotal- n[Fe(OH),(RS),(m+2"-3)-] teine had previously been determined to be 3.05 X (14) l o 3 mole-' cm.-l a t 580 mp. Using this value the where [RSH]total is the total initial concentration concentration of the ferric-cysteinate complex a t 271 - a - ] is zero time was calculated, Since all experiments of cysteine and [Fe(OH),(RS),(" iequal to the total iron concentration in the solution were carried out with varying concentration of a t zero time. cysteine keeping other conditions practically conThe inconsistent K values (see Table V) obtained stant, the ratio of [Fe-complex]/ [NH2RS-for m = 0 and n = 2 from experiments a t different C O O - I n u n c o m b , must be independent of the concenfiH indicate that the formation of the complex tration of uncombined cysteine if only one species is FefRS)2- can be excluded. From the series of ex- present in the solutions. It was found, however, periments 12 to 16, in which the concentration of that using either n = 1, or 2 or 3, the ratio does not cysteine was varied keeping other conditions prac- remain constant over the entire range of cysteine tically constant, the formation of Fe(OH)p(RS)- concentrations. A plot of log [Fe-complex] against ( m = 2, n = 1) is considered unlikely, because of log [NH*RS-COO-] uncomb. indicates that Fe(RS)33the systematic change in the calculated constant is predominant a t higher concentrations of cysteine with change in concentration of cysteine. (see Fig. 4), while the number of cysteine molecules The data of Table V combined with results from coordinated to iron decreases with decreasing conexperiments with Versene, described in the follow- centration of uncombined cysteine. ing section, indicate that under our experimental I n a solution of FH 10.3, FeY(OH)z3-is the most conditions the two species Fe(OH)(RS)z2- (m = 1, predominant species of ferric Versenate. l7 From the n = 2) and Fe(RS)s3- (m = 0, n = 3) coexist in apparent formation constants reported by Schwarsolution according to the equilibrium zenbach and HellerI7for the equilibria Fe(OH)(RS)r*-
+ RS- J_ F e ( R S h -
f OH-
F e + + + f Y 4 - zF e y - ,
I-FeY(0H)' Evaluation of Dissociation Constants of FerricCysteinate Complexes in Ammoniacal Versenate FeY(OH)z3 FeY(0H)- + O H Solutions.-It was observed that the violet color of the ferric complex can be developed in alkaline Ver- the apparent dissociation constant senate solutions. Upon addition of air-free cysteine solution to an air-free ferric Versenate solution of pH 10, a violet color develops which fades on stand- for the equilibrium ing. Under our experimental conditions the soluFeY(OH)2a-J 'F e + + + f 2 0 H - + Y l tion of ferric Versenate was pale yellow, its mixture a t 20" and 0.1 with cysteine becoming colorless after the reduction was calculated to be S.25 X of ferric-cysteinate complex was completed to form ionic strength. ferrous-cysteinate and ferrous-Venenate Complexes. This constant has been used although our esFeY- f OH-
i I(i1 \ I'XI I.atimer, "Oxidation Potrnfials."2nd ed , Prentice-Hall. I n r , b r a . Vork, N V , 1BR2, p 224.
(17) (; fl961).
Sd~warzenbarhand J
Heller, Hui;' rhi8,i .411n, 3 4 , 57G
~
~
April 5, 1955
2003
IRON-CYSTEINATE COMPLEXES
periments have been carried out at 25' and at ionic strength of 0.14 to 0.16. The concentration of ferric Versenate, FeY(OH)23-, was obtained as [FeY(OH)P-] = [Felt
-
[Fe(III)-cysteinate complex]
The concentration of uncombined Versenate is [Ver~enate],,~~b.= [Versenate]totsl - [FeY(OH)z3-]
Introducing the ionization constant of Versenate ion (pK4 = 10.26),18the concentration of Y4- was obtained from the concentration of uncombined Versenate, neglecting the concentrations of HzY-, H3Y- and H4Y. From [FeY(OH)z3-], [Y4-], [OH-] and KFeY(0H)ns-j the concentration of ferric ion, [Fe+++],was calculated (Table VI). The concentration of RS= was calculated from the ionization constants and activity coefficients of the various species of the cysteinate ions and the concentration of uncombined cysteine assuming the formation of Fe(RS)33-. Introducing the activity coefficients for Fe+++, Fe(RS)s3- and RS-, the dissociation constant of Fe(RS)z3-
for the equilibrium Fe(RS)2-
Feff+
Kw ---
a = 0 (15)
y3
Koomp.
aH+
where x denotes the concentration of FeOH(RS)z2and k, the term
Applying eq. 15 to the experiments in Table V, the concentrations of FeOH(RS)z2- and Fe(RS)33were calculated as given in Table VII.
-4.5
-
5
2 w d
+ 3RS-
5
was calculated (Table VI). Considering that Fe(RS)33- is the predominant species only a t higher cysteine concentrations (see Fig. 4) , the average value of the dissociation constant was found to be from experiments with initial equal to 2.1 x RSH- concentrations of 11.92 to 7.94 X M (see Table VI). The predominance of a complex with two molecules of cysteine may be considered in solutions of lower concentration of cysteine (see Fig. 4). Referring to the solubility studies in ammoniacal buffers (see Table V), this complex probably is FeOH(RS)z2--,with a dissociation constant
0 0
-W I
-
-5.0
lL Y
(3
0 J
-5.5
-2 5
-3.0
Fromtheresults of an experimentwith an initial RSH concn. of 6.68 X M (Table VI) a value of is calculated for the constant of this 1.4 X complex. The ratio of K F e O H ( R S ) 1 z - / K F e ( R S ) a a - ( = Kcomp.) corresponding to the equilibrium FeOH(RS)22-
+ RS'
Fe(RS)33-
+ OH-
yields a value of 6.5 X We have concluded that under the experimental conditions listed in Table V both FeOH(RS)Z2- and Fe(RS)s3- are present in ammoniacal cysteinate solutions. Introducing Kc,,,., the ratio of FeOH(RS)2-/Fe(RS) 3 3 - is given by
Since the total concentration of ferric complexes (a), that of cysteine (b) and the hydrogen ion concentration are known, the following equation was derived (18) C, Schmarzenhach a n d H Ackermann, Helo Chim A l l n , SO, 1798 (1947)
LOG CNHzRS- C O O - 3
( MI.
Fig. 4.-Plot of log [Fe(III)-complex] against log indicates theoretical lines for [NHzRS-COO-]; Fe(RS)n(2n-3)-. -.-*-e-
At a given $H the ratio of FeOH(RS)z2- to Fe(RS)33- increases with decreasing concentration of cysteine while a t a given concentration of cysteine i t increases with increasing pH. The dissociation constants calculated for FeOH(RS)z2- and Fe(RS)33- from these concentrations of FeOH(RS)Z2- and Fe(RS)33- and the corresponding concentrations of uncombined cysteine gave consistent values for both species (Table VII). This substantiates our previous conclusion that two complexes FeOH(RS)Z2- and Fe(RS)33- coexist in ammoniacal cysteinate solutions under our experimental conditions. The average values for K F ~ o H ( R S and )~~K F ~ ( R Swere ) ~ B -calculated to be 1.9 X and 2.9 X (Table VII), respectively, compared with 1.4 X and 2.1 X obtained spectro-
2004
NOBUYWITANAKA. 1. M. KOLTHOFF AND
x
Concn. of Fe(RS)II -
x
10'M
7.18 4.96 3.59 5.17 4.94 4.78 1.46 1.42 1.22 0.40 0.40 1.43 1.09 0.87 0.65 0.215
lo', M
9.28 7.26 4.91 6.93 6.66 5.99 1.29 0.88 0.80 0.16 0.15 1.27 0.78 0.56 0.37 0.071
KFeOH(RB)s2-
x
10"
1.0 1.7 2.2 1.3 1.4 1.5 2.5 1.5 2.1 2.1 2.1 2.4 2.2 2.2 2.2 2.6 1.9
Vol. 77
donlg yields 3 X 4 f ~the r value of K F e O H ( R S ) t * and 5 X for that of K p e ( ~ ~ ) , * - . These are in better agreement with the constants calculated from the experiments with Versene than the constants calculated with a solubility product of 6 X As probable values we propose
TABLE VI1 CALCULATION OF THE CONCENTRATIONS OF FeOH(RS)zzAND Fe(RS)r8- AND OF K F ~ o H ( R SAND ) , ~ - KFs(RB)*t- I N AMMONIACAL CYSTEINATE SOLUTXONS OF TABLE V All figures are listed in the sequence corresponding to the exp. No. in Table V. Concn. of FeOH(RS)zs-
iv. S T R I C K 5
KFe(RB)aPx 1062
KF,OH(RS),P= 5 X K F ~ ( R B )=~ 8' - x
1.6 2.5 3.3 2.0 2.1 2.3 3.7 2.3
Schubert5 postulated a dimer formula { FeOH(RS)zjZ4- for the ferric-cysteinate complex. The diffusion coefficient of the ferric-cysteinate complex in solutions of ionic strength 0.12-0.31 containing no gelatin was extrapolated to be 5.7 X cm.* set.-', which is of the same order as that of cystine,20and also of that of iron(II1)-Versenate 3.2 complex, FeY(0Hj-n (5.4 X sec.-1).21 3.1 This strongly suggests that the complex is a mono3.4 mer. 3.6 In studies of the ferric-thioglycolate complex 3.3 Leussing and Kolthoff7concluded that the predomi3.4 nant species was FeOH(TS)Z similar to one of the 3.3 ferric-cysteinate complexes postulated in the pres3.9 ent study. Recalculating the dissociation conMean 2.9 stant of the thioglycolate complex7 yields a value This value was obtained in ammophotometrically in ammoniacal Versenate solutions. of 9.4 X In the calculation of the latter values the disso- niacal buffers and is based on a solubility product of Comparison with ciation constants of ferric Versenate complexes ferric hydroxide of 6 X of the corresponding ferric-cysteinate measured a t 20" and 0.1 ionic strength were used 1.9 X instead of those a t 25" and ionic strength 0.14-0.16 complex, FeOH(RS)Z2-, indicates that the stability in the present experiments. This may be partly re- of both complexes is of the same order of magnitude. Acknowledgment.-This investigation was supsponsible for the fact that one set of values is about ten times greater than the other set. The main ported by a research grant (c-721,C~)from the source of the difference in the two sets of values National Cancer Institute, U. S . Public Health probably is the uncertainty in the solubility product Service. of ferric hydroxide which is involved in the calcu(19) K . Jellinek and H. Gordon, 2 . p h y s i k . C h c m . , 112, 236 lation in Table VII. The value of the solubility (1924). (20) I. M. Kolthoff and C Barnum, THISJ O U X N A L , 63, 520 (1941). product depends upon the method of preparation (21) I. M. Kolthoff a n d C. Auerbach, ibid., 74, 1452 (1952). and the age of the precipitate-using a value of 1 X for S F ~ ( O H ) , ( ~reported ) by Jellinek and Gor- MINNEAPOLIS, MINNESOTA [CONTRIBUTION FROM
THE SCHOOL OF
CHEMISTRY OF
THE
UNIVERSITY
Oxidation of Ferrous-Cysteinate Complex by Cystine. Cystine-Cysteine System
OF
MINNESOTA]
Oxidation Potential of the
BY NOBUYUKI T A N A KI. AM. , ~ KOLTHOFF AND W. STRICKS RECEIVED NOVEMBER 4, 1954 Upon addition of ferrous iron to an air-free ammoniacal solution containing cystine and cysteine some ferrous iron is oxidized by cystine to a violet ferric cysteinate complex. From spectrophotometric determinations of the ferric complex -RSSR- 3 2Fe(III)(RS);- and 2Fe++ *RSSR* the equilibrium constants for the reactions 2Fe(II)(RS)z2H+ Fi 2Fe+++ 2RSH* were calculated to be equal to 2 . 5 X 10-8 and 5.3 X 10-24, respectively, a t 25". From the latter value the oxidation potential of the cystine-cysteine system was calculated to be +0.08 volt us. N.H.E. at 25".
+
In a study of ferrous- and ferric-cysteine complexesZ we observed the development of a violet color upon addition of an air-free ferrous iron solution to an air-free ammoniacal mixture of cystine having a large concentration in cysteine. On standing the color intensity increased to a constant value. Other conditions being the same the color became ( I ) On leave of absence from Tokyo University, Japan ( 2 ) N. T a n a k a , I. M. Kolthoff and W Stricks, THISJ O U R N A L , 7 7 , 1996 (1955).
+
+
+
more intense with increasing cystine concentration. Spectrophotometrically the color was found to be similar to that of ferric cysteinate which is formed on the addition of ferric iron to a cysteine solution.2 These observations indicate that under proper experimental conditions cystine can oxidize ferrous cysteinate to the ferric complex. Our observations are a t variance with those of Michaelis and Barrens who found no indication for the oxidation (3) L Michaelis and E S G Barron. J BIOI Chcm., 83, 191 (1929)
