Reaction of ironIII with hydroxyaminoacids

Reaction of ironIII with hydroxyaminoacids

REACTION OF IRON”’ WITH HYDROXYAMINOACIDS E. R. NIGHTINGALE,Jr.* and R. F. BENCK~ Esso Researchk Engineering Company, Linden, New Jersey, U.S.A. an...

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REACTION

OF IRON”’

WITH HYDROXYAMINOACIDS

E. R. NIGHTINGALE,Jr.* and R. F. BENCK~ Esso Researchk Engineering Company, Linden, New Jersey, U.S.A. and UniveMy of Nebraska, Lincoln, Nebraska, U.S.A. (&c&ed 22 JuQ 1963. Acceprcd29 septcmbct 1963) Summa~--The stoichlomctry and aqurhbrium for the reaction of ironl*I with iV,IVdthydroxycthylglycmem aqueous medium has been mvesttgated. Ironrxxis shown to form only a 1: 1 compkx with the figand; previous mports of a 2.3 complex are demonstratedto beerroneous. and probabi rcsultcd from noncqulbbrium titration measurements. Using a mod1tLd BJerrumtitration procedure, the uilibrium constant for the formation of the 1: 1 complex at 25” has?a n determined to k 3.8 x IO-#moles*/htrc~. Attempts to dH&cntiatc between altcmat!vc structures for the complex are tnconsequcntial, because the complex extsts oniy in aqueous m&turn, where the ligand hydroxyl groups exchange rapidI? with solvent water. PREVIOUSstudies ~1 our laboratories l*s have investigated the use of N,N-dihydroxyethylglycine (DHEG) as a complexing agent to control the mechanism for hydrolysis of iron”’ and permit the precipitation from aqueous solution of crystalline iron”’ oxide as /?-FeOOH. Because the role of the tetradentate DHEG @and appears to be unique in facilitating the precipitation of the crystalline oxide, the stoichiomctry and equilibrium for the reaction of DHEG with iron II* have been investigated in order to provide further info~ation concerning the reactions involved. In an aqueous soIutron of pH 2-4, iron”’ reacts with DHEG to form the compIex (I) Fc+3 + HG + 2H,O 3 Fe(OH),G + 3H+. (1)

(I) In this reaction DHEG is assumed to behave as* a monoprotic acid HG, which coordinates in four positions about the iroP ion. Formula (I) is identical with that proposed by Torcn and Kolthoff3 from polarographic studies of the reaction of iron” and iron*” with DHEG. It conflicts with the earlier work by Martell and [email protected] who assumed DHEG to react as a triprotic acid H,A, according to equation (2): Fe- + H3A + 2H,O + Fe(H,O),A + 3H+, (2) (II) * To whom mqumes should be addressed: Esso Research & Engineering Company, P.O. Box IZI, Linden, N.J., U.S.A. t Presentaddress: Nuclear Defense Laboratory,Nuclear Chemistry Division, Edgewood Arsenal,

Md , U.S.A.

241

242

E. R

NIGHTINGALE,Jr

and R. F.

BENCK

where water fills the two co-ordmatton postttons allotted to the hydroxyl tons m equation (1). Formulae (I) and (II) doffer m that Martell et al. assumed the two hydroxyl groups m the co-ordmated hgand to be more acldlc than water, whereas the conventlonal mterpretatlon assumes the co-ordmated water to be the more actdtc Because the iron”’ complex IS stable only in aqueous solutton, where the hgand hydroxyl groups should exchange raptdly with the solvent, dtfferenttatton between the two structures would be trivtal if Martell et al. had not reported the formatton, above pH 8, of a second bmuclear complex, Fe,A,-3, whtch was not observed m the other studies *13 The present work is a reinvesttgation of the stotchiometry for the reactton between iron”’ and DHEG, and the equilibrium constant for equation (1) has been determuted The prevtous evidence’ for the formation of the Fe2A3-3 complex IS demonstrated to be erroneous, and probably resulted from non-equtlibrium titration measurements EXPERIMENTAL DHEG was puritiad by recrystalhsmg twice from ethanol. Standard rron*** solutrons were prepared by dtssolving iron whe in hydrochloric acid and diluting to volume Oxrdatton of ahquots of the iron solution was effected by treatmg with hydrogen peroxrde and decomposing the excess hydrogen peroxide by boiling. The stabihty constant of the ironln-DHEG complex was determined by titration at 25” (aide [email protected]). Before titration, the acidrty of the rron*rr solutron was reduced to pH 2.0 f 0 05 by neutrahsation with gaseous ammoma to prevent the excess

actd from consuming an appreciable rtion of the added base. The DHEG complexmg agent was then added, and the solutton was dt-p” uted to volume and titrated wtth standard sodturn hydroxtde soluuon Details have been described elsewhere * NMR spectra were recorded using a Varian A-60 spectrometer. RESULTS

Acidic solutions containing 0.OOlM iron”’ and 0.001 1M to 0*02M DHEG were titrated with standard base. The titration curves exhtbit two end-point inflections. The first of these occurs when the ratio m, the number of moles of base added per mole of iron”’ in solution, equals three. This end-point corresponds to the titration of the three equivalents of hydrogen ion liberated accordmg to equation (1) or (2). Usmg a modified Bjerrum-type procedure, the equilibrium (1) may be calculated from the titration results, where

constant,

K, for equation

(3) If n is the number of moles of complex formed per mole of total tronnT, for n = 0.5, [FelrT] = [Fe(OH) zG] 7 and K = [H]+3/[HG]. In acid solution the concentrations may be readily calculated as [HG] =

(4)

of [OH-] and (G-1 are negligtbly small, and HG ((3 - a& - W+l}K 4[H+] + 3K, ’

where a is the number of moles of base added per mole of total DHEG, K1 is the t%sr acid dissociation constanta*4 of the protonated amino acid, H,G+, and Co is the

Reactlon of mm*l’ :wth hydroxvammoaclds

total concentration of DHEG tn durmg the tttratton, n IS given by

n=

solution.

243

Knowing [HG] as a function of the pH

C, - V-=IIW+I/~K,+ 1)

(6)

C FC

where C,, is the concentratton of total tron”’ m solutron. Plotting [H+] and [HG] as a function of n, values correspondmg to n = 0.5 permit the calculatton of K in equation (4). Using this procedure, the pH of the soluttons III the vrciruty of n 2~ 0.5 remains sufficiently acidic (pH < 3.6) to prevent Interference from the mcrprent precrpitation of Fe(OH),. TABLE I-DTTERMINATION OF rquii_teRiuhl CONSTANTFOR Fe(OH),G. (Iontc strength = 001 T = 25” C,, = 1 00 lo-‘M) IO2 K

(PHI n-0.5

105 WGl.=w

1I

3 55 360

0490 0 492

4 58 321

20 20

3 54 3 55

I 26 1 17

I.90 I 91

50 50

3 34 3 36

3 61 3 74

264 2 23

100 100

3 10 3 10

8 43 621

5.82 821 1 8)

G/C,, 11

avg. = (3 8 i pK=74

Table I indicates typical values for K calculated from titration results for varrous ratios of C,/CX,. Our value of pK = 7.4 IS in reasonable agreement with the equally precise value of 6.1 calculated from the polarographic data of Toren and Kolthoff 3 Their K,,,is related to our K by

K = Kl,,Kn2K2. m which Kw IS the autoprotolysis constant for water and K, is the second iorusatton constant for DHEG.3p4 The position of the second end-pomt mflectron in the titration curves depends upon the concentration of excess DHEG in solutlon. For Co/C,, ratios of 2, 5, 10 and 20, the end-point occurs when m equals 5, 8, 13 and 23, respectively. If three moles of base are required to titrate the acid l&rated by the formation of the non”‘DHEG complex according to equation (1) or (2), the remainder (m - 3) represents the amount of base required to precipitate Fe(OH), from the complex and to neutrahse the excess DHEG m solution. These latter reactrons may be represented by the general equation: Fe(OH),G

+ (m - 4)HG + (m - 3)OH- -

Fe(OH), + (m - 3)G- + (m - 4)H,O. (6)

As written, this reaction requires all of the ironr” in solution to be precipitated upon addition of (m - 3) moles of base. The amount of Fe(OH)s precipitated during the titrations was analysed gravimetrically by Ignition to F%03 and also by dissolving the hydrous oxide in acid and

E R

244

NIGHTINGALE, Jr and R

F BENCK

titrating with standard cenum” solutron. Because the aggregation and prectpitatlon of the colloidal lronr1r hydroxide IS very slow under these condltlons, if solutions to which m moles of base had been added were allowed to settle for l-3 days, quantltatlve precipitation of the iron was achieved m every case The ability to precipitate the Iron under such condltlons has been the basis for the procedure for the quantltatlve precipitation of crystalhne iron”’ oxide as /?-FeOOH.’ DISCUSSION

The present studies have confirmed the stolchlometry of reaction (I) [or (2)] in which iron”’ forms only a 1: I complex with DHEG In acid medium They contrast sharply with those of Martell et al. (lot. cir.), who reported an unspecified partial precipitation of iron which led them to suggest a second binuclear complex Fe,Asw3*. We conclude that the latter complex is not formed, for it IS not thermodynamlcally stable, and iron”’ IS quantitatively precipitated under the experimental conditions The propriety of formula (I) or (II) in solution cannot be determined. Intultlvely, structure (I) IS preferred, because the hgand hydroxyl groups should have approxlmately the same acldlty as in ethanol, and should be considerably less acidic than water (K3 and K, for DHEG are too small to be measured in aqueous medium) However,the relative acldltles of HOCH,CH,and H,O when co-ordmated by iron”’ have not been determined On the basis of polarisabihty and hgand-field strength, it is predicted that H,O will dlssoclate to a considerably greater degree than will CH,CH,OH when co-ordmated by iron”‘; and hydrolysed species such as Fe(OH)-2 and Fe(OH),+ are well known NMR spectra of solutions of DHEG in 99.5% D,O show, m addition to the HOD singlet, two triplets and one singlet, each of relative area 2, which are characterlstlc of the three methylene groups In the hgand molecule The absence of a more detailed spectrum, mcludmg a trlplet of relative area 1 at the proper chemical shift for the ethanolic -OH group, indicates that the ligand hydroxyl groups do exchange rapidly with the solvent water. Because the exchange IS rapld, differentiation between structures (I) and (II) IS mconsequentlal. Zusammenf~ung-Stochlometrle uncj

Glelch$wlcht

der Reaktlon

van Etsen(Iil)

mlt N,N-dthydroxyathylplycm m wassrtgem Medium wurden untersucht Es wlrd gezelgt dass Elsen(lll) nur emen 1.1Komplex blldet; fruhere Angaben uber emen 2 3-Komplex srnd falsch und baslerten wahrschemhch auf Messungen ausserhalb des Gleichgewichts Mlt einer modlfizlerten Tltratlonsmethode nach BJeMm wurde eme Glelchgewlchtskonstante von 3,8 10-8 mol*/l* fiir dte Bddung des 1 1-Komplexes be1 25’C erhalten. Strukturangaben fur den Komplex smd nocht smnvoll. da er nur m wassrlgem Medium exlstlert, wo die als Llganden auftretenden Hydroxylgruppen rasch mit dem Ldsungsmlttel ausgetauscht werden RCum&--La stoechlom&rle et I’6qulhbre de la reactlon du fer(II1) avec I’aclde N,N’ dlhydroxyCthyl ammo-adtlque en mlheu aqueux ont et6 &dies II est ttabh que le fer(II1) forrne seulement un complexe 1 1 avec le complexant; II est dCmontr6 que le complexe 2 3 d&rrt pr&%demment n’exlste pas et rCsultalt probablement de mesures faltes avant que I%qullrbresoit r&ah& En utrbsant une merhode de BJetTum modlfi& la constante d’tqurhbre de formation du compfexe 1 1 d&ermm& A 25°C est 3,80. 10-O moles*/l*. Les structures du l Martell er ~1.4 do not report If approxrmately one-tlurd of the total Iron was prccrpitated as required by the stoichlometry of the reaction

Reactron of Iron(l~I)

with hydroxyamIn~cld~

245

complexe sont sans mtC*t pu~squ’~lexiste seulement en mlhtu aqueux o& les hatsons entre fes groupes hydroxyles s’tchangent rapldemenr avcc I’eau du solvant REFERENCES 1 E, * R. * P. 4 S. * R.

R Nightmgale, Jr and R F Benck, Analyr Chem , 1960,32,566 F. Benck, M. S. thesis, Untvtrs~ty of Nebraska, 1958 E. Toren and 1. N. Kolthoff, J Amer Chem. Sot , 1955.27,2061. Chabwck, Jr., R. C Courtney and A. E. Marteli, &id, 1953,7S, 2185 C. Courtney, R I.. Gustafson, S. Chaberek, Jr and A E. Martell, r&d,

1958,80,2121.