Iron(III) Complexation by Hydroxyurea in Acidie Aqueous ...

CROATICA CHEMICA ACTA CCACAA 61 (1) 21-31 (1988)

CCA-1776YU ISSN 0011-1643

UDC 547.72OriginaL Scientijic Paper

Iron(III) Complexation by Hydroxyurea in Acidie AqueousPerchlorate Solution"

A. Bed1"ica, M. BiTUŠ, N. Kujundžić, and M. Pribomić

Department oj Chemistry, FacuLty oj Pharmacy and Biochemistry, University ojZagreb, Zagreb, Croatia, YugosLavia

Received Oc:tober 19, 1987

Equilibrium and kinetic studies were performed to investigatethe complexation of aqueous high spin iron(III) by hydroxyureaH2NC(O)NH(OH) in acidie solutions at 25 <c and I = 2.0 mol dm?(maintained by NaC104). Complexation has been interpreted interms of coordination of the N-O oxygen atom and the NH2nitrogen atom of the ligand to the iron(III) ion with concomitantloss of a proton yielding the complex of the molar ratio 1 :1. Theequilibrium quotient for the formation of mono(hydroxyureato)iron(III) complex is found to be Kl = 1.4. The kinetic results suggesta parallel path mechanism involving substitution on Fe(H20)63+and Fe(H20l50H2+ by the hydroxyurea, HU:

klFe (HzO)63+ + HU ;;= Fe (HzO)4 U2++ H+

u.,k'

Fe (H20)s (OH)2++ HU;;=' Fe (HzO)4 U2'k'.!

The formation of the complex occurs by the rate constantskl = 16.8 M-l S-I and k1' = 5450 M" S-I. The analogous rate eon-stants for the reverse hydrolysis reactions were obtained as k-l == 11.8 M-l S-I and k_1' = 6.3 S-I. The results are compared withkinetic data previously reported for the different mono(hydroxa-mato)iron(III) complexes.

INTRODUCTION

Hydroxamic acids, RI-C(O)N(OH)-R2' have a wide variety of applica-tion in industry, pharmacy and chemistry. They have been used as flotationreagents in extractive metallurgy, inhibitors for copper corrosion, food ad-ditives, therapeutic agents and analytical reagents. They are biologicallyactive as antibiotics, growth factors, tumor inhibitors, pigments and chelatingagents."?

The most important feature of hydroxamic acids is their ability ofiron(III) sequestration which classifies them in a group of compounds corn-monly called siderophores which are intimately associated with iron trans-port in living organisms.' The inter action between iron(III) and synthetic,

* Taken, in part, from the Master Thesis of A. B.

22 A. BEDRICA ET AL.

as well as naturally occurring, hydroxamic acids apears to be a very im-portant bioinorganic reaction currently stirring wide interest.š' "

Hydroxyurea, (H2N-CONHOH = HU), is one of hydroxamic acids whichshows antitumor activity and is introduced into cancer therapy.P It has beenshown that HU inhibits enzyme ribonucleotide reductase.t! Since some ofribonucleotide reductase contain non-he me iron(III), the complex formationreaction of HU with iron(III) may be of importance for an understanding ofantitumor activity, as well as the HU side effects.

EXPERIMENTALMaterials

Iron(III) perchlorate was prepared by dissolving freshly obtained iron(III)hydroxyde in concentrated (700/0) perchloric acid and recrystallized from diluteperchloric acid. A stock solution of iron(III) perchlorate (0.15 M in 0.1 M HCl04)was prepared and stadardized as described previously."

The hydrogen ion concentration in the stock solution was determined bypassing an aliquot through a Dowex cation exchange resin in the acid form. TheIr ion concentration was determined by titration with NaOH and corrections weremade for the iron(III) present,

Sodium perchlorate was prepared by neutralization of dry Na2CO:1 by eon-centrated HCl04, and was recrystallized from water. A stock solution of NaCl04was used to maintain constant ionie strength,

Hydroxyurea was purchased from Sigma Chem. Co. and its reagent solutionwas prepared by dissolving the solid immediately before the measurements weremade.

All solutions were prepared using water which was double distilled fromalkaline KMn04 in an al l-glass apparatus. All other chemieals were of analyticalgrade and were used without further purification.

Methods

All experiments were performed at 25 ± 0.1 Cc in an aqueous solution of 2.00Mionic strength. The total concentrations of H+ ion in the experiments werecalculated by summation of added HCl04 and the proton released from the ironspecies present in solutions.

The spectrophotometric and kinetic measurements were performed on a UnicamSP 800 spectrophotometer, Durrum D-110 stopped-flow spectrophotometer and aDionex stopped-flow apparatus linked with a Harrick rapid-sean monochromator,all equipped with a thermostated cell compartment. A modified version of anoriginally published non-Iinear least square procedure was applied on a UNIVAC1100 computer at the University Computing center, Zagreb, for the data reductionanalysis.'!

The pseudo-first order conditions were ensured by holding one reactant inexcess over the other.

RESULTS AND DISCUSSION

By rmxing the iron(III) solution with hydroxyurea, a blue colored com-pl ex is formed which quickly decomposes. Therefore, the rapid-sc an stopped--flow technique was us ed to record the spectra of mono(hydroxyureato)-iron(fII) complex (Figure 1). Essentially the same spectra (Amax = 560 nm)were recorded both in mol ar excess of iron(III) over the ligand and viceversa. This suggests that in solution of [H+]> 0.01 in the first stage a com-plex of 1: 1 iron(III): ligand is formed. This is confirmed by the methodof continous variation applied to the iron(III)-hydroxyurea system at pH == 2.0. The plot of the absorbance vs. iron(III) fraction, X, where X = [Fe(IIl)]!

IRON(III) COMPLEXATION BY HYDROXYUREA 23

00 L---4~5~0----;50:;I;O-;::-----;:5J.=50"'----6~0=0------:6::!:5-:::-0--lWAVEL ENG TH Inm

Figure 1. Visible spectra of mono(hydroxyureato)iron(III) complex during the for-mation of the complex. All 11 spectra were taken in 0.3 s (each spectrum after 0.03 s).Conditions: [Fe(III)l,o' = 5.65 X 10-4, [HUlIO!= 7.5 X 10-3, [HCI04l = 0.4, I = 2.0 M

(HCI04!NaCI04).

010 o o

01 Q2 OJ Q4 as 06 07 OBV/\

Figure 2. Continuous variation curve at pH = 2.0, I = 2.0 M (HCI04!NaCI04),[Fe(III)],o'+ [HUl,o' = 1 X 10-2 M. The solid line represents the theoretical curve cal-culated us ing the values for Kl and El listed in the Table. t = 25°C, ). = 560 nm.

!([Fe(III)] + [HU]), showed a maximum at 0.5 indicating Fe(III): HU = 1 : 1complex stoichiometry (Figure 2). Figure 3 shows the plot of the ratio of


the total iron(III) concentration to the absorbance, vs. total iron(III) eon-centration. The data were obtained at 560 nm, where the spectrum ofthe complex shows a maximum of absorption. At this wavelength iron(III)ions do not exhibit significant absorption. The line ar relationship confirmsthat only one complex exists und er these conditions as it follows from theBenessi-Hildebrant method-" for the calculation of stability constants. Thatone H+ ion is involved in the complexation reaction is illustrated by theHill plot shown in Figure 4, where the slope is 0.98. Thus, the first step inthe reaction of iron(III) with hydroxyurea may be ascribed to the formationof mono(hydroxyureato)iron(III) complex and may be defined by equation (1)(coordinated water molecules were omitted):

K,Fe3+ + HU ~ FeU2++ H+

[FeU2+] [H+]K -

1 - [Fe3+][HU]

U- = H2N-C(=O)NHO-

(1)

(2)

Therefore, under the conditions studied the formation of other Fe(III)-hydro-xyurea complexes, such as the bis(hydroxyureato)iron(III) complex, havenot been recorded, at least, not to a measurable extent.

01

Fe(IIIJ.tat

A

Q2

[Fe(J]J)7iot

Figure 3. The ratio of total iron(III) concentration to absorbance vs. total iron(III)concentration at 560 nm. Conditions: [HU] = 1 X io=, [HCl04]= 1 X 10-2, I = 2.0 M

(HCl04/NaCl04),25 -c.

0025 005


r-::;-, 5"{

I "-'>.; "{ .4"{E~ ..s 3

2

o

-3 -1In[Hj

-1 L-~ ~ ~ ~ ~

-2

Figure 4. Determination of the number of hydrogen ions involved in the equilibrium.Conditions: [Fe(III)]tot= 2 X 10-2, [HU].o.= 1 X 10-3, 1= 2.0 M (NaCl04/HCl04), 25 -c.Molar absorptivity used in calculation of Am" is obtained from the plot depicted in

Figure 3.

The absorbance data from Figures 1-4 were treated all together usinga non-line ar ·least square procedure to fit the function

(3)

where A is the absorbance at )"= 560 nm, and c, and Ei are concentrationsand molar absorptivities at 560 nm of each species present in the solutions.Concentrations c, are dependent on the experimental total concentrations ofiron(III), hydroxyurea, proton, and refined value of K1. Since the molarabsorptivities of non-chelated iron(III) species, HU and proton at 560 nmare negligible, only the molar absorption coefficient of the mono(hydroxy-ureato)iron(III) complex and Kl had to be refined during the calculations,The calculated values of Kl and El are given in the Table.

The influence of the electron donor-acceptor ability of the --C and -Nsubstituent of the hydroxamate functionality on the stability of mono(hydro-xamato)iron(III) complexes has been throughly discussed by Crumbliss eta1.5,9,11 They found that increasing inductive electron donor strength of R2,

for example, when R2 = CH3, enhanced the relative contribution of resonanceform II by delocalization of the N atom lone pair of electrons into the carbonylfunctionality and thereby increased the negative charge density on Oh whichwould be expected to enhance the iron(III)-carbonyl oxygen bond strength


Fe Fe Fe

/~ /~ /~o, °2 O' °2 Q' °2II I [' I. I' IC--N C=N C--N

/ '" / -, ,/ '"R R R R H2N' H

I Il III

(Fe-c-Oj). A buildup of negative charge density on 01 is expected also throughresonance form III for the hydroxyureatoiroruffl) complex. This means thatthe stability constant of the hydroxyureatoiron(III) complex should be higherthan that of for example, acethydroxamatoiron(III) complex, with R2 = H inboth complexes.

Since the obtained values are opposite to those expected, that is KJ == 1.42 for the hydroxyureato- and KJ = 80 (ref. 18) and KJ = 109 (ref. 5)for the acethydroxamatoiron(III) complexes, a different mode of coordinationshould be proposed, which will be discussed later.

TABLE

Equilibrium, Kinetic and Spectrai Data for the Mono(hydroxyureato)iron(III)Complex at 25 DC, 1=2.0 mol dm-3 (H/NaCL04)"

pK." k' 1 E(Amaxl

M-! cm'!

9.0· 1.42 16.78 5450 11.82" 6.33d 400 (560)

8.520 1.3'

• Estimated values of standard deviations are not shown throughout this Table sincethey do not exceed 100/0 of the reported parameters.

b For the following reaction: HU~ U-+ H'c Determined potentiometrically under the same experimental conditions.đ Hydrolysis rate constants were calculated by expressions k-l = kl/K), k'-l =k'l Kh/Kl'

o At 0.1 Mionic strength and 20 ce (lit. 17).r Kinetically determined under the same experimental conditions.

Kinetics

The three first order observable processes in the reaction of hydroxy-urea with iron(III) ion are illustrated in Figure 5. After the complex isformed it decomposes through the stages which involve redox processes andwhich are rather complicated. This paper deals with the first reaction.Figure 6 shows that a dependence of kobs on the total iron(III) concentrationat constant H+ ion concentration is line ar. The calculated value of K, fromthis kinetic data agrees well with that obtained from equilibrium measure-ments (see Table). The acid dependence shown in Figure 7 is similar to thatobserved for the complexation of a series of mono(hydroxyamato)iron(III)


complexes, suggesting the paralleI path mechanism which is well known andtypical of many ligation reactions of ferric ions. The reaction path whichwould involve interaction of iron(III) ion with hydroxyureato anion, U-, maybe ruled out by the same arguments as presented before." Thus, the reactionby which the iron(III) and hydroxyurea form the mono(hydroxyureato)iron-(III) complex may be described by the Scheme:

kHU + Fe3+...--!.- FeU2+ + H+11.,

KhJ ~H+ + HU + FeOH2+

k' 1K'---

1 - k'-I

(5)

When iron(III) is present in a molar excess over hydroxyurea and when[H+] »Kh, Kh = 1.0 X 10-3 M at 2.0 Mionic strength, 25°C,19 the observedrate constant is defined by eq. (6).

Khkobs = (kl + k'i [H+]) [Fe]tot + k_1 [H+] + k'_1 (6)

Equation (6) requires a linear dependence of kobs on the total iron(III) eon-centration at constant proton concentration, as shown in Figure 6. In addition,eq. (6) requires a three parameter function when the system goes in an

0.2

01 0.2 10 20. 30. 1.0.ttne]»

50. 10.0. 20.0.

Figure 5. Illustration of the three first order observable processes in the reaction ofhydroxyurea with iron(III) ion. Conditions: [Fe(III)]tot= 2 X 10-2, [HU]tot= 1 X 10-3,

I = 2.0 M (HCI04!NaCI04),25°C, .l = 560 nm, [HCI04]= 0.05.


007 002 003 001.[Fe(1I1jj

tot

Figure 6. Observed first order rate constants for the formation of mono(hydroxy-ureato)iron(III) complex plotted as a function of the total iron(III) concentration.Conditions: [HU] = 5 X 10-', A= 560 nm, 250C, I = 2.0 M (NaCI04/HCI04), [HCI04] =

= 0.05.

kObJŠ'

78o

71.

727086

4

2

02 06 08al.

•

Figure 7. Plot of kob, VS. [H+] for the mono(hydroxyureato)iron(III) complex for-mation reaction at 25°C. The solid line represents a least square fit of the data to

eq. (7). Conditions: [Fe(III)]wt = 2 X 10-2, [HU] = 1 X 10-3, (e), and 5 X 10-4 (O).

IRON(III) COMPLEXATION BY HYDROXYUREA 29equilibrium (Figure 7). The non-linear least square procedure was appliedto refine the parameters of rearranged eq. (6) into eq. (7) using the formationkinetic data points and spectrophotometrically determined Kl value.

Kh kl [H+] + k'l . KhkOb' = (kl + k'l [H+] ) [Fe]tot + x, (7)

Similarity of the kinetic expressions of the interaction of ferric ion withhydroxyurea and the other monohydroxamic acids allowed us to test thereaction mechanism by the already us ed kinetic relationship.š In Figure 8kinetic data for hydroxyurea are plotted together with the data for differentmonohydroxamic acids taken from references 5. and 9. The linear relation-ship between In k'_l and In k_l is usually interpreted to me an that the acid--independent (k'_I) and acid-dependent path (k_l) of the hydrolysis exhibit asimilar mechanism which is, in addition, common to all the complexes studiedin a particular series. Therefore, on the basis of this plot it is reasonable tosuppose a common reaction mechanism for all monohydroxamic acids includinghydroxyurea. Obviously, this is in disagreement with the conclusion presentedabove, based on the equilibrium data, where a different mode of coordinationhas been proposed.

Two possible modes of coordination of HU with iron(III) ion are depictedby formulas IV and V, in which cases a stable five-membered ring may be

4 .--------------------------------,Ink.,

2 •

o

-2.o

o-4

oo

-6 o

-6 -4 -2 o 2

Figure 8. Plot of In k'-l VS. In k_1. The data marked by (O) were collected from theCrumbliss et al. papers, (.) values for the mono(hydroxyureato)iron(III) complex.


'" /-Fe-/ '"O O

H" II IN-C--N-H

H/

'" /-Fe-

H" / -,N O

H/r IC---N-HI;

O

IV V VI

formed. If structure V iz effective, one might expect that the hydroxyureadata point (e) should lie outside the line in Figure 8. However, several fac-tors may be invoked to explain the position of the obtained point. This pointlies on the line, but outside the region of other data points of most hydroxa-matoiron(III) complexes for which structure IV has been proposed. Theexplanation may be that structure V enhances the rate in both acid catalizedand spontaneous hydrolysis by approximately the same factor, thus ensuringthe same slope for the hydroxyureatoiron(III) complex as for other complexes.

Furthermore, the position of the HU data point on the line requiresthe same intercept as other hydroxamatoiron(III) complexes. According tothe Asher and Deutch explanation'", the dominant factor for the intercept isthe net charge of the complex which is the same for structures IV and V.Asher and Deutsch demonstrated that positively charged ligands lie belowthe line, which we observed for the positively charged betaine hydroxamicacid."!

The molecular structure of hydroxyurea obtained by neutron diffractionanalysis shows that the O atom bonded to N is oriented as depicted in for-mula VI.22 It has been found that the C=O bond is an about 80010 doublebond in character and both the N-C and C-N bonds are about 10010 doublebonds in character. In addition, IR-spectral data showed that when hydro-xyurea is dissolved in water, the CO bond is substantially asingle bond incharacter and the remaining electron density is distributed across N-C-Nbonds producing a barrier to rotation.P For all these reasons, structure Vas a mode of coordination in mono(hydroxyureato)iron(III) complex is pre-ferred.

It is also supported by different Amax and Emay; Amax - 500 nm, Emax '""'

- 1000 M-i cm-i, for most of monohydroxamatoiron(III); Amax = 560 nm,ES60 = 400 M-i cm-i for mono(hydroxyureato)iron(III).

These findings provide a good basis for studying the subsequent decom-position reactions which are the subject of our current interest.

Acknowledgement. - This work received Iinancial support from the CroatianCouncil for Research and the U.S. - Yugoslav Joint Board on Scientific and Techno-logical Cooperation which is gratefully acknowledged, We thank Professor N a n c yR o w a n Go r d o n from The American University, Washington, D. C., USA, forher interest in this work and for sending us some relevant data.

The authors should like to thank to Professor R a 1p h G. W i 1k i n s (NewMexico State University, Las Cruces, NM, USA) for the use of laboratory facilities.

IRON(III) COMPLEXATION BY HYDROXYUREA 31REFERENCES

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4. K. N. Ray m o n d and T. P. Tu fan o in The BioLogicaL Chemistry of Iron,H. B. D u n for d et aL Eds., D. Reidel Publishing.

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SAŽETAK

Kompleksacija željeza(III) sa hiđroksiureom u vođenoj otopini perklorne kiseline

A. Bedrica, M. Biruš, N. Kujundžić i M. Pribanić

Zeljezo(III) i hidroksiurea stvaraju mono(hidroksiureato)željezo(III) kompleks.Na 250C i u kiseloj otopini ionske jakosti I = 2,0 (HCI04/NaCI04) ravnotežni kvoc-jent stvaranja kompleksa iznosi Kl = 1,4, a konstante brzine stvaranja jesu: kl == 16,8 M-l S-l iz Fe3+ i k'l = 5450 M-l S-l iz FeOH2+. Analogna konstante brzina reak-cija za povratnu reakciju hidrolize su k-l = 11,8 M-l S-l i k'-l = 6,3 S-l.

Iron(III) Complexation by Hydroxyurea in Acidie Aqueous ...

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