Inorganica Chimica Acta 376 (2011) 500–508

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Inorganica Chimica Acta

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Characterization of Mn(II) and Mn(III) binding capability of natural siderophoresdesferrioxamine B and desferricoprogen as well as model hydroxamic acids

Orsolya Szabó, Etelka Farkas ⇑Department of Inorganic and Analytical Chemistry, University of Debrecen, P.O. Box 21, H-4010 Debrecen, Hungary

a r t i c l e i n f o

Article history:Received 21 March 2011Received in revised form 28 June 2011Accepted 11 July 2011Available online 30 July 2011

Keywords:Hydroxamic acidsSiderophoresManganese(II)/(III) complexesSolution equilibriaRedox behaviourRelaxivity

0020-1693/$ - see front matter � 2011 Published bydoi:10.1016/j.ica.2011.07.010

⇑ Corresponding author.E-mail address: efarkas@delfin.unideb.hu (E. Farka

a b s t r a c t

Complexes formed between Mn(II) ion and acetohydroxamic acid (HAha), benzohydroxamic acid (HBha),N-methyl-acetohydroxamic acid (HMeAha), DFB model dihydroxamic acids (H2(3,4-DIHA), H2(3,3-DIHA),H2(2,5-DIHA), H2(2,5-H,H-DIHA), H2(2,4-DIHA), H2(2,3-DIHA)) and two trihydroxamate based naturalsiderophores, desferrioxamine B (H4DFB) and desferricoprogen (H3DFC) have been investigated underanaerobic condition (and some of them also under aerobic condition). The pH-potentiometric resultsshowed the formation of well-defined complexes with moderate stability. Monohydroxamic acids not,but all of the dihydroxamic acids and trihydroxamic acids were able to hinder the hydrolysis of the metalion up to pH ca. 11. Maximum three hydroxamates were found to coordinate to the Mn(II) ion, but pres-ence of water molecule in the inner-sphere was also indicated by the corresponding relaxivity valueseven in the tris-chelated complexes. Moreover, prototropic exchange processes were found to increasethe relaxation rate of the solvent water proton over the value of [Mnaqua]2+ in the protonated Mn(II)–siderophore complexes at physiological pH. The much higher stability of Mn(III)–hydroxamate(especially tris-chelated) complexes compared to the corresponding Mn(II)-containing species resultsin a significantly decreased formal potential compared to the Mn(III)aqua/Mn(II)aqua system. As a result,air oxygen becomes an oxidizing agent for these manganese(II)–hydroxamate complexes above pH 7.5.The oxidation processes, followed by UV–Vis spectrophotometry, were found to be stoichiometric onlyin the case of the tris-chelated complexes of siderophores, which predominate above pH 9. ESI-MSprovided support about the stoichiometry and cyclic-voltammetry was used to determine the stabilityconstants for the tris-chelated complexes, [Mn(HDFB)]+ and [MnDFC].

� 2011 Published by Elsevier B.V.

1. Introduction

Manganese is involved in numerous biological processes, includ-ing in oxidation of water by the manganese-containing complex,Photosystem II [1,2], in dismutation of superoxide catalysed bythe SOD enzyme [3,4]. There is a great interest in the study of man-ganese complexes which can be somehow related to the processesof manganese-containing biosystems. Since, any biological impor-tance of interaction between manganese and hydroxamate-basedcompounds was previously not expected, no any special interest to-wards such kind of complexes seemed in the past. In recent years,however, some interesting aspects came to light in this respect.Namely, (i) since significant metalloenzyme inhibitory effect ofnumerous hydroxamate-based compounds have been found [5–8]valuable results might be obtained to this question by investigationof the different factors, determining the interaction between bio-genic metals, including manganese and hydroxamic acids, (ii) givenits five unpaired d-electrons, long electronic relaxation time, fast

Elsevier B.V.

s).

water exchange and its lower toxicity compared to Gd(III), Mn(II)complexes can be potential candidates for contrast agent applica-tion in medical magnetic resonance imaging [9,10], (iii) the investi-gation on the manganese(II)/(III) and the hydroxamate-basedsiderophore, desferrioxamine B (DFB) systems, as well as the com-parison between the stability of the complexes, Mn(III)HDFB andFe(III)HDFB, by Duckworth et al., led to the conclusion that, becausethe stability constants of these complexes are similar, the Mn(III)might compete with Fe(III) during the siderophore-mediated ironuptake process [11–13]. The same conclusion was made by studyingthe complexes with two additional hydroxamate-based sidero-phores, pyoverdin and desferrioxamine E [14,15].

In spite of the above-mentioned possible importances, at pres-ent, only a few equilibrium data are known for manganese(II)/(III)–hydroxamate complexes. Because, the desferrioxamine B hasbeen used in the clinical practice for many decades [16], this sider-ophore and the cyclic analogue, plus the pyoverdin have beeninvolved into investigations, but in addition to the mentionedsiderophores, only a few additional data were published forMn(II)-complexes of simple monohydroxamic acids [17,18]. Noany date was found in the literature e.g. for the complexes with

Scheme 1.

O. Szabó, E. Farkas / Inorganica Chimica Acta 376 (2011) 500–508 501

desferricoprogen (DFC), the production of what is wide-spreadamong fungi [19–21]. That is the reason, why it was decided tostudy in our laboratory the manganese complexes formed withselected mono-, di- and trihydroxamic acids. Scheme 1 shows theformulae of the ligands. All of the systems have been studied underanaerobic conditions, but the systems containing either DFB or DFCas ligand, or one of the DFB model dihydroxamates or acetohydrox-amate were also studied on air.

2. Experimental

2.1. Chemicals

Acetohydroxamic acid (HAha) and benzohydroxamic acid(HBha) were commercially available chemicals (Sigma andAldrich). N-Methyl-acetohydroxamic acid (HMeAha) [22], thedihydroxamic acids (H2(3,4-DIHA), H2(3,3-DIHA), H2(2,5-DIHA),

H2(2,5-H,H-DIHA), H2(2,4-DIHA), H2(2,3-DIHA)) [23,24] and des-ferricoprogen (H3DFC) [25] were prepared by reported procedures.Desferrioxamine B (H4DFB) was produced by CIBA Geigy. Thepurity of the ligands and the concentrations of the ligand stocksolutions were determined by Gran’s method [26].

The metal ion stock solution was prepared by the dissolution ofMnCl2 in tri-distilled water, which contained known amount ofHCl. The concentration of the Mn(II) stock solution was checkedby gravimetrically via precipitation of MnNH4PO4�H2O. The acidconcentration of the metal ion stock solution was determined bypH-potentiometric method.

2.2. Potentiometry

The pH-potentiometric and spectrophotometric measurementswere carried out at an ionic strength of 2.0 � 10�1 mol dm�3 (KCl)and 25.0 �C. Carbonate-free KOH of known concentration

502 O. Szabó, E. Farkas / Inorganica Chimica Acta 376 (2011) 500–508

(ca. 2 � 10�1 mol dm�3) was used as titrant. pH-potentiometrictitrations were performed throughout in the pH range 2.0–11.0(or up to appearance of precipitation). For the metal–ligand sys-tems initial volume of the samples was either 20.0 or 10.0 or8.0 ml, depending on the amount of the ligand, and the metal-to-ligand ratio ranged from 1:1 to 1:5, whereas the ligand con-centration varied in the range of 1.5 � 10�3–5.0 � 10�3 mol dm�3.The measurements were carried out under anaerobic conditionsby bubbling purified argon for ca. 15 min before starting and alsoduring the titration. A Radiometer PHM 93 pH meter equippedwith a Metrohm combined electrode (type 6.0234.100) was usedfor pH-potentiometric measurements together with a Metrohm715 Dosimat automatic burette. The electrode system was cali-brated according to Irving et al. [27] and the pH-metric readingscould be converted into hydrogen concentration. The pKw calcu-lated from strong acid–strong base titrations was 13.76 ± 0.01.During the titrations the equilibrium position was reached withinless than 2 min in all cases. The reproducibility of the titrationpoints included in the evaluation was within 0.005 pH unit. Theexperimental results were used to establish the stoichiometry ofthe formed species and to calculate the stability constants. Sincethe hydrolysis of the Mn(II) starts at pH ca. 9.5, the mono-hydro-xo species with fixed stability constant (log KMnH-1 = �10.46, at25.0 �C [28]) was always included in our equilibrium modelswhen the experimental data were fitted. Calculations were per-formed with the computer program PSEQUAD [29]. The volumes ofthe titrant were fitted and the accepted fittings were always be-low 1 � 10�2 ml.

2.3. Spectrophotometry

A Perkin Elmer Lamda 25 spectrophotometer was used to recordthe UV–Vis spectra over the range 250–850 nm. The path lengthwas 1 cm. The spectrophotometric method was utilized to studythe behaviour of the systems under aerobic condition. The Mn(III)has a characteristic band (310 nm, e = 2060 ± 20 mol�1 dm3 cm�1

[11,12]). This band was used to follow the change which mightoccur on air in the investigated systems. Any possible degradationof the Mn(III)-complexes was checked for 2 weeks.

The measurements were carried out by preparing individualsamples which contained ionic strength of 2.0 � 10�1 mol dm�3

KCl and the metal ion concentration was 1.0 � 10�3 mol dm�3

(DFB and DIHAs) or 7.0 � 10�4 mol dm�3 (DFC) or 3.0 � 10�3 moldm�3 (Aha). The metal ion to ligand ratio was 1:1 in the case ofMn(II)–DFB and Mn(II)–DFC systems, 1:1.5 at the Mn(II)-3,4-DIHA,Mn(II)-2,5-DIHA and 1:3 at Mn(II)-Aha. The samples were oxidizedby bubbling oxygen gas. The spectra were registered at pH values�7, �8, �9.5 and �10 in all cases (pH was set by KOH solution)and at stated intervals during a maximum period of 3 h ortill the oxidation process was completed according to thespectra (when the molar absorbance coefficient reachede � 2000 mol�1 dm3 cm�1).

2.4. ESI-MS measurements

The oxidized samples of Mn-3,4-DIHA, Mn-2,5-DIHA, Mn–DFBand Mn–DFC systems were analysed by ESI-MS. The ligand concen-tration was 5.0 � 10�4 mol dm�3 in the case of 2,5-DIHA, 3,4-DIHA,DFC and 1.0 � 10�3 mol dm�3 at DFB. The metal-to-ligand ratioswere 1:1.5 (DIHAs) and 1:1 (DFB and DFC). pH � 9.5 was set byKOH solution in all cases. Oxygen gas was bubbling through the tri-hydroxamic acids containing samples for the time determined tothe complete oxidation in the UV–Vis spectrophotometric measure-ments. In the case of dihydroxamic acids the samples were oxidizeduntil the maximum oxidation was indicated by the UV–Vis results.

A Bruker BIOTOF II ESI-TOF instrument equipped with a Cole Pal-mer 74900 Series pump (sample flow rate of 2.0 � 10�3 ml min�1)was used to perform ESI-MS measurements. All experiments werecarried out in the positive ion mode.

2.5. Cyclic voltammetry

Cyclic voltammetric measurements were performed on a Metr-ohm 746 VA Trace Analyser. A Pt-wire auxiliary electrode and aglassy carbon working electrode were used. The working electrodewas polished with alumina paste (particle size 5 � 10�2 lm) beforeeach measurement. The reference electrode was Ag/AgCl/3 mol dm�3 KCl (E1/2 = 209 mV versus NHE). Aqueous solution ofK3[Fe(CN)6] was used to calibrate the system (E1/2 = 0.458 V versusNHE in 5 � 10�1 mol dm�3 KCl) [30]. The scan rate was varied withinthe range of 50–100 mV s�1 during the determination of the redoxpotentials.

The metal concentration was 1.5 � 10�3 mol dm�3 and theligand concentration was 4.5 � 10�3 mol dm�3 (DFB) or 2.25 �10�3 dm�3 (DFC). The measurements were made in the presenceof KNO3 electrolyte, the concentration of what was varied in therange 5.0 � 10�3–7.5 � 10�2 mol dm�3, in the different systems.All measurements were carried out at 25 �C, under an Ar atmo-sphere and the volume of the samples were 1.5 ml in all cases.Before each measurement the oxidation of complexed Mn(II) toMn(III) was carried out by bubbling oxygen gas in all samples(details are given at the Spectrophotometric part (2.3)). Theexperimental results were used to calculate the formal redox po-tential (E00) and the stability constants for the Mn(III)-containingcomplexes. The following equation was used to calculate thestability constants:

E00ðMIIIHxL=MIIHxLÞ ¼ E0ðMIIIðH2OÞ6=ðMIIðH2OÞ6Þ þ

RTnF

lnðbII=bIIIÞ

where E00(MIIIHxL/MIIHxL) is the formal potential of the complexx = 1 in the case of DFB and x = 0 in the case of DFC E0(MIII

(H2O)6)/(MII(H2O)6) is the redox standard potential of Mn(III)(aq)/Mn(II)(aq) (1.6 V [31]) bII/bIII is the quotient of the overall stabilityconstants of the appropriate complexes.

2.6. Relaxometry

Relaxometric measurements were made on a 20 MHz MinispecBruker instrument. With this technique T1 (spin-lattice) relaxationtime was measured by inversion-recovery method. The volume ofthe samples was 4 � 10�1 ml, at an ionic strength of 2 � 10�1

mol dm�3 KCl and the temperature was 25 �C. The ligand (2,5-H,H-DIHA, 2,5-DIHA, DFB and DFC) concentration in all sampleswas 2 � 10�3 mol dm�3 and the Mn(II) concentration was variedin the range 4 � 10�4–2 � 10�3 mol dm�3. To set the pH Hepes(4-(2-hydroxyethyl)-1-piperazineethanesulfonic acid) and NEP(N-ethyl-piperazine) buffer solutions were used and the pH valueswere �7.6 and �9.3, respectively. All measurements were carriedout under inert atmosphere (Ar). Relaxivity values were calculatedby using the registered data according to this equation: R1 = 1/T1.

The relaxivity can be expressed as follows:

1=T1 ¼ RMn½MnðH2OÞ6� þ RMnL½MnL� þ RMnHL½MnHL� þ . . .þ 1=T10

where 1/T10 is the relaxation rate in the absence of Mn(II), while RMn

and RMnL ��� are the relaxivities for the existing various manganese-containing species, [Mn(H2O)6]2+ and MnL ���, respectively [32].

Individual relaxivity values of the complexes were calculatedusing the computer program PSEQUAD [29].

8

10

12

pH

2

4

6

8

10

12

-0.5 0.0 0.5 1.0 1.5

Base equivalent

pH

a

b

O. Szabó, E. Farkas / Inorganica Chimica Acta 376 (2011) 500–508 503

3. Results and discussion

3.1. H+–ligand systems

As it is shown in Scheme 1, there are mono-, di- and trihydrox-amic acids among the investigated ligands. In the case of themonohydroxamic acids the RN substituent is varied, it is eitherhydrogen or methyl or benzyl moiety. The investigated dihydroxa-mic acids are all derivatives of H4DFB. Both the H2(2,5-DIHA) andH2(2,5-H,H-DIHA) contain exactly the same linker between thetwo hydroxamic functions as the DFB does, they differ from eachother only in their RN substituents. In the other dihydroxamic acidderivatives either the chain-length is shortened (H2(2,4-DIHA),H2(2,3-DIHA)), or the position of the peptide group within thechain (H2(3,4-DIHA)) or both (H2(3,3-DIHA)) are modified.

Because all of the investigated ligands have been involved inprevious studies in our lab [23–25,33–35], their protonation con-stants have already been determined under the same conditionas used in the present work. The values obtained in the presentwork and summarized in Table 1 are all in good agreement withthe previously determined ones.

2

4

6

-1.0 0.0 1.0 2.0 3.0

Base equivalent

2

4

6

8

10

12

-1.0 0.0 1.0 2.0 3.0 4.0 5.0

pH

c

3.2. Mn(II) (M)–Ligand (L) systems under anaerobic condition

Representative pH-potentiometric titration curves obtained forthe ligands and for Mn(II)–ligand samples with a selected mono-hydroxamate (Bha), dihydroxamate (3,4-DIHA) and siderophore(DFB) are presented in Fig. 1a–c, respectively.

As it is clear from the curves, there is no measurable interactionbetween these ligands and the Mn(II) ion below pH 5–6. It is alsodemonstrated that at high pH the hydrolysis of the metal ion can-not be hindered by a monohydroxamate even at 1:5 metal to li-gand ratio (Fig. 1a). However, if a minimum of 50% excess of adihydroxamate compared to the metal ion (Fig. 1b) or minimumequimolar amount of a trihydroxamate (Fig. 1c) is present, thehydrolysis is completely hindered up to ca. pH 11. The equilibriummodels and the stability constants of the formed complexes werecalculated by fitting of the titration curves with the aid of PSEQUAD

computer program. The results are summarized in Table 2.In agreement with the conclusion drown form the titration

curves (Fig. 1a), Table 2 shows the formation of mixed hydroxo

Table 1Overall (log b) and stepwise (log K) protonation constants of the investigated ligands(I = 0.2 mol dm�3 KCl; t = 25 �C).

Ligand log b1 (log K1) log b2

(log K2)log b3

(log K3)log b4

(log K4)

Aha 9.25(1) (9.25) – – –MeAha 8.63(1) (8.63) – – –Bha 8.69(2) (8.69) – – –2,3-DIHA 9.25(1) (9.25) 17.65(1)

(8.40)– –

2,4-DIHA 9.26(1) (9.26) 17.70(1)(8.44)

– –

2,5-DIHA 9.23(1) (9.23) 17.69(1)(8.46)

– –

2,5-H,H-DIHA

9.70(1) (9.70) 18.49(1)(8.79)

– –

3,3-DIHA 9.30(1) (9.30) 17.67(1)(8.37)

– –

3,4-DIHA 9.31(1) (9.31) 17.70(1)(8.39)

– –

DFB 10.85(1)(10.85)

20.37(1)(9.52)

29.33(1)(8.96)

37.66(1)(8.33)

DFC 9.82(3) (9.82) 18.66(3)(8.84)

26.59(4)(7.93)

–

Standard deviations are given in parentheses.

Base equivalent

Fig. 1. (a) pH-metric titration curves for the Bha (�) and for Mn(II)-Bha samples at1:5 (h), 1:3 (N), 1:2 (s) and 1:1 (�) ratios at clig = 5 � 10�3 mol dm�3. (b) pH-metrictitration curves for the (3,4-DIHA) (�) and for Mn(II)-3,4-DIHA samples at 1:3 (h),1:2 (N), 1:1.5 (�) and 1:1 (s) ratios at clig = 2.5 � 10�3 mol dm�3. (c) pH-metrictitration curves for the DFB (�) and for Mn(II)–DFB samples at 1:1.37 (N) and 1:1(s) ratios at clig = 2 � 10�3 mol dm�3 (negative base equivalent values mean acidexcess).

complex, [MH�1L], at high pH with monohydroxamates, followingthe formation of mono- and bis-chelated complexes. On the otherhand, there is no measurable existence of mixed hydroxo speciesneither with dihydroxamates, nor with the two natural sidero-phores, DFB and DFC. Most probably the hydrolysis is hindered bythe formation of tris-chelated complexes at high pH in thesesystems.

If the corresponding numerical values of the stability constantsare compared and evaluated the following conclusions can bemade: (i) With monohydroxamates the trend of the log stabilityconstants (Mn-Aha > Mn-Me-Aha �Mn-Bha) follows the trend ofthe ligand basicities. At the same time, the moderate stability con-stants obtained for the complexes [ML]+ and [ML2], are as expectedon the basis of the Irving–Williams trend (valid for the high-spinoctahedral complexes of the 3d M(II) transition metal ions) [36].

MO

O

O

HO

MO

O

OO

M

O

O

O

O

O

OM

O

O

O

O

O

O

M

O

O

O

O

O

ONH3

+

M

OO

O

OH3N+

OHO

MO

OH3N+

OHO

O OH

(I) [MHL]+ (IV) [MH3L]2+

(II) [ML]

(III) [M2L3]2-

(V) [MH2L]+

(VI) [MHL]

Scheme 2.

Table 2Overall stability constants (log b) of the complexes formed in Mn(II)-monohydroxamate, -dihydroxamate and -trihydroxamate systems (I = 0.2 mol dm�3 KCl; t = 25 �C).

Complex Monohydroxamate Dihydroxamate Trihydroxamate

Aha MeAha Bha 2,3-DIHA 2,4-DIHA 2,5-DIHA 2,5-H,H-DIHA

3,3-DIHA 3,4-DIHA DFB DFC

log b

[MH3L] – – – – – – – – – 32.6(1) –[MH2L] – – – – – – – – – 25.51(2) 22.56(8)[MHL] – – – 12.66(7) 12.92(6) 12.8(1) 13.43(7) 12.70(5) 12.72(1) 17.39(3) 16.09(2)[ML] 3.83(2) 3.36(1) 3.52(1) 5.57(2) 5.75(2) 6.25(1) 6.77(1) 4.80(2) 4.93(1) 6.81(3) 8.28(3)[ML2] 7.05(2) 5.89(3) 5.90(5) – – – – – – – –[M2L3] – – – 17.54(7) 17.87(8) 18.0(2) 20.43(8) 15.6(1) 14.4(2) – –[MH-1L] �6.38(3) �6.89(3) �7.7(7) – – – – – – – –pM 5.00 5.00 5.00 5.00 5.00 5.01 5.01 5.00 5.00 5.00 5.08Number of fitted

data146 167 136 80 74 74 152 79 77 127 109

Fitting parameter(Dml)

8.4 � 10�3 1.0 � 10�2 1.0 � 10�2 2.1 � 10�3 2.3 � 10�3 4.7 � 10�3 7.5 � 10�3 1.8 � 10�3 9.9 � 10�4 1.1 � 10�2 2.8�10�3

Standard deviations are given in parentheses. Charges of the complexes are omitted. pM values calculated for various Mn(II)-complexes at pH 7.4 andcMn(II) = clig = 10�5 mol dm�3.

504 O. Szabó, E. Farkas / Inorganica Chimica Acta 376 (2011) 500–508

(ii) With dihydroxamic acids, as it shown in Table 2, the interactionstarts with the formation of [MHL]+, and this species most probablyhas the bonding mode depicted as (I) in Scheme 2.

Since only one of the two hydroxamates of each ligand is coor-dinated to the manganese(II) ion in this species, any measurableeffect of the structure and length of the connecting chain on thecomplex stability cannot be expected. The obtained results agreewell with this expectation, namely the stability constants of themono-chelated [MHL]+ complexes are almost identical (exceptthe value for the complex with 2,5-H,H-DIHA). The higher stabilityof [MHL]+ with 2,5-H,H-DIHA, corresponds to the higher basicity ofthis ligand. However, significant differences can be realized if thecorresponding stability constants of the [ML] and [M2L3]2� formedbetween the manganese(II) ion and different DIHA-s (see Table 2)are compared with each other. The fact that the length and struc-ture of the connecting chain have measurable effect on the stabil-ity of the [ML] type complexes supports the coordination mode

(II), in which the two hydroxamate functions of a DIHA are coordi-nated to the same metal ion resulting in the formation of a mono-nuclear bis-chelated complex. The trend in the stability constantsof these complexes (log K[ML]) is the same as those foundpreviously for the corresponding complexes with numerous othermetal ions [23,24] and it is the following: log K[Mn(2,5-H,H-DIHA)] >log K[Mn(2,5-DIHA)] > log K[Mn(2,4-DIHA)] > log K[Mn(2,3-DIHA)]� log K[Mn(3,4-DIHA)]

� log K[Mn(3,3–DIHA)]. Again, the difference between thelog K[Mn(2,5-H,H-DIHA)] and log K[Mn(2,5-DIHA)] corresponds clearly tothe basicity difference between the two ligands, but additionalfactor, most probably sterical one, results in the unfavouredformation of [M2L3]2� (bonding mode III) with the 2,5-DIHA(RN = –CH3) compared to the corresponding H,H-derivative (thiscan be clearly seen in Fig. 2, if the concentration distributioncurves relating to these two systems are compared).

Probably, the poorly adequate length of the spacer link to thesize of the manganese(II) is responsible for the lower stability of

Mn(II)

[MHL]+

[ML]

[M2L3]2-

0.0

0.2

0.4

0.6

0.8

1.0

5.0 6.0 7.0 8.0 9.0 10.0

pH

Fra

ctio

n of

Mn(

II)

Fig. 2. Concentration distribution curves for the complexes formed in Mn(II)-2,5-DIHA (solid lines) and in Mn(II)-2,5-H,H-DIHA (dotted lines) systems at 1:1.5 ratio(clig = 2 � 10�3 mol dm�3).

O. Szabó, E. Farkas / Inorganica Chimica Acta 376 (2011) 500–508 505

the complexes formed with 2,4-DIHA and 2,3-DIHA compared to2,5-DIHA. On the other hand, although, the length of the connect-ing chain in 3,4-DIHA and 2,5-DIHA is completely the same, muchless stable complexes are formed with 3,4-DIHA (the complex for-mation is shifted to the direction of higher pH) than with 2,5-DIHA(see Table 2). Significant effect of the position of the peptide groupin the linker on the orientation of the two chelating functions isknown from our previous Molecular Mechanics (MM) calculations.Those results related to the bis-chelated mononuclear complexeswith Fe(III) ion and supported the energetically favoured coordina-tion of the 2,5-DIHA over the 3,4-derivative [23]. The resultsobtained in the present work support the same situation withMn(II). (iii) Numerous protonated complexes are formed with thetwo trihydroxamate-based siderophores, DFB and DFC. The possi-ble bonding modes in the different complexes of DFB are depictedin Scheme 2 (structures IV–VI). The coordination mode in [MHL]and [ML]� is the same, only the non-coordinating amino-end is stillprotonated in the former species. Because, the completely proton-ated forms of the two investigated siderophores contain differentnumber of dissociable protons, only the stability constants of theircompletely deprotonated metal complex, [ML]�, can be compareddirectly. If we do this, significantly higher stability of [M(DFC)]�

compared to [M(DFB)]� is found. This difference in the stabilityresults also in the significant difference in the concentration distri-bution curves (Fig. 3) obtained with the two ligands. For example,one can see in Fig. 3 that at physiological pH (pH 7.4) the fraction ofthe free Mn(II) is only 0.1 if the ligand is DFC, while it is still ca. 0.5with DFB. Most probably, the more effective manganese-binding

[MnH2DFC]+

[MnHDFC]

[MnDFC]-Mn(II)

[MnH3DFB]2+

[MnH2DFB]+

[MnHDFB]

[MnDFB]-

0.0

0.2

0.4

0.6

0.8

1.0

5.0 6.0 7.0 8.0 9.0 10.0

pH

Fra

ctio

n of

Mn(

II)

Fig. 3. Concentration distribution curves for the complexes formed in Mn(II)–DFB(solid lines) and in Mn(II)–DFC (dotted lines) systems at 1:1 ratio(clig = 2 � 10�3 mol dm�3).

ability of DFC, compared to DFB, can be attributed to the longerconnecting chains and favoured electronic delocalization betweenthe hydroxamates and double bonds in b-positions in the formerligand. Similar difference in Pb(II)-binding ability of the two sid-erophores was observed previously [37].

3.3. Relaxivity of Mn(II)–DFB and Mn(II)–DFC complexes

As it is well known, vast majorities of the MRI (magnetic reso-nance imaging) investigations use contrast agents, first of all Gd-complexes [38]. One of the key requirements for a MRI contrastagent is the marked ability to enhance the relaxation rate of solventwater. ‘‘Given its five unpaired d-electrons, long electronic relaxationtime, and fast water exchange, Mn(II) is a potential candidate forcontrast agent application in medical magnetic resonance imaging.Nevertheless, the design of chelators that ensure stable Mn(II) complex-ation and optimal relaxation properties remains a coordination chem-istry challenge’’.[9] Despite this fact and also the much less toxicityof Mn(II) compared to Gd(III), still less attention is given to theMn(II) complexes from this respect. According to our best knowl-edge, no any date has been published in the literature relating therelaxivity of Mn(II)–hydroxamate complexes. That is the mainreason that it was decided to investigate the relaxometric propertiesof manganese complexes of DFC and DFB. Although, the relaxivityvalue for [Mn(H2O)6]2+ has already been published in the literaturepreviously (7.4 mM�1 s�1 at 24 MHz) [39], but as a first step, it wasalso redetermined under our condition in this work. The value ob-tained in the present work (8.1 mM�1 s�1 at 20 MHz) is in accept-able agreement with the previously determined one [39].

Because the stability of the Mn(II)–hydroxamate complexes ismoderate (the calculated pM values are ca. 5, see in Table 2), highpercentage of the metal ion can be in non-complexed (‘‘free metalion’’) form in these systems at physiological pH. Taking this intoaccount, relaxometric studies were performed only on Mn(II)–siderophore (DFB, DFC) complexes with the highest stability. Themeasurements were made at two different pH values (7.6 and9.3) as given in Section 2. As it is shown in Fig. 3, a mixture ofthe various (bis- and tris-chelated) complexes exists in these twosystems at ca. physiological pH, while the tris-chelated complexalmost predominates at pH 9.3. The individual relaxivity valueswere calculated for the complexes formed at the investigatedpH-values, by using the measured relaxivities, the equilibriummodel and stability constants. The results obtained are shown inTable 3.

As it can be seen in Table 3, the individual relaxivities for the tris-chelated complexes, [Mn(DFC)]� and [Mn(HDFB)], are much lowerthan that found for the [Mn(H2O)6]2+, but still high enough to indi-cate involvement water molecule in the inner-sphere [9]. On theother hand, the values for the bis-chelated complexes [Mn(H2DFB)]+

(12.7 mM�1 s�1) and [Mn(HDFC)] (11.8 mM�1 s�1) are surprisinglyhigh (much higher than that for the [Mnaqua]2+ species). Because, inthese latter two complexes only two of the three chelating functionsare coordinated and the third is still protonated. This significantincrease in the relaxivity most probably belongs to the effect ofwater mediated intra- or/and intermolecular prototropic exchangeon the coordinated hydroxamate and protonated, non-coordinatedhydroxamic functions. Similar behaviour was found when relaxo-metric properties of some gadolinium–hydroxypyridinonate com-plexes were investigated [40]. To support this assumption, becausethere is no possibility for the above mentioned prototropic exchangeprocess in their bis-hydroxamate type complexes, two modeldihydroxamates (2,5-DIHA and 2,5-H,H-DIHA) have also been in-volved into the study. The obtained values are much lower thanthose for the bis-chelated complexes of the siderophores (Table 3).This supports that the very high values for the [Mn(H2DFB)]+ and[Mn(HDFC)] cannot be assigned simply to the relaxation effect of

Table 3Measured and calculated individual relaxivity for some dominant complexes formed in Mn(II)-2,5-DIHA, -2,5-H,H-DIHA, -DFB and -DFC systems.

Dihydroxamate Trihydroxamate

2,5-DIHA 2,5-H,H-DIHA DFB DFC

Relaxivity (mM�1 s�1) Relaxivity (mM�1 s�1) Suggested bonding mode Relaxivity (mM�1 s�1) Suggested bonding mode

Measured pH �7.6 7.6 ± 0.3 7.9 ± 0.3 9.1 ± 0.3 � 9.6 ± 0.5 �pH �9.3 � � 5.7 ± 0.3 � 5.7 ± 0.6 �

Individual [MH2L] � � 12.7(1) 2-Hydroxamate-(O,O) � �[MHL] � � 4.8(1) 3-Hydroxamate-(O,O) 11.8(3) 2-Hydroxamate-(O,O)[ML] 7.4(2) 7.9(1) � � 5.4(2) 3-Hydroxamate-(O,O)

The relaxivity value for [Mn(H2O)6]2+: 8.1(1) mM�1 s�1.Standard deviations are given in parentheses. Charges of the complexes are omitted.

506 O. Szabó, E. Farkas / Inorganica Chimica Acta 376 (2011) 500–508

inner-sphere water molecules, but also to the effect of the prototro-pic exchange processes.

3.4. Mn(II)–ligand systems under aerobic condition

As it is known from previous works of Faulkner and co-workers[15] and also of Duckworth and co-workers [11], the Mn(II)–DFBsystem becomes sensitive for air-oxidation above pH 7. It wasfound in these studies that stable Mn(III)-containing complex canbe formed in this system within the pH-range 7–11.3. To obtainresults under our condition, we have also studied the Mn(II)–DFBsystem. Moreover, because practically no result for other Mn(II)–hydroxamate systems was found in the literature, it was decidedto characterize the redox behaviour of various complexes formedwith some additional hydroxamates. In these investigations, onemonohydroxamate (Aha), two dihydroxamates (2,5-DIHA and3,4-DIHA), as well as, in addition to DFB, DFC have also been in-volved. The Mn(II)–hydroxamate systems were exposed to air atvarious pH-s and, first of all, UV–Vis spectrophotometry was usedto follow the changes. To obtain information about the possibleformation of Mn(III)-complexes, the change of a characteristic UVband (kmax is at ca. 310 nm, e is ca. 2060 mol�1 dm3 cm�1)[11,12] was followed within a 3 h period.

According to the experimental findings obtained, indication wasnot found for the measurable oxidation of Mn(II) in its hydroxa-mate complexes below pH ca. 7.5. At and above this pH, no matterwhat the ligand was, some change in the UV-spectrum around300 nm was observed, but, the rate and extent of changes werefound to depend both on the denticity of the ligand and on thepH in a great extent. Only low extent of the oxidation of theAha-containing complexes below the pH of the hydrolysis was

Fig. 4. Absorbance spectra recorded as a function of time for Mn–DFB system at pH8 (cDFB = 1.1 � 10�3 mol dm�3, metal-to-ligand ratio 1:1). Inset: The molar absorp-tion coefficient values at kmax = 310 nm registered as a function of time for Mn–DFBsystem at pH �9.5 (cDFB = 1 � 10�3 mol dm�3, metal-to-ligand ratio 1:1).

found, and this was the reason, why a more detailed investigationwas not made on this system. Also with the DIHA containing sys-tems (most probably because the tris-chelated complex is not pre-dominant even at pH ca. 10), the oxidation was not stoichiometricunder the investigated conditions. The appearance of the charac-teristic UV-band supported the formation of Mn(III) containingcomplex(es), but the measured absorptions did not reach the max-imum values expected on the basis of the molar absorption coeffi-cient and concentrations. What is more, at pH P 9.5, after ca. halfan hour, the absorbance of the characteristic peak (indicating thebeginning of some additional new processes) started to decreaseif the ligand was 3,4-DIHA. Direct evidence for the not completeoxidation was obtained by the ESI-MS investigations, where amixture of various Mn(II) and Mn(III) complexes (including aMn(III) containing mixed hydroxo species, [Mn(2,5-DIHA)(OH)]K+

at m/z = 398.06 or the Mn(II) containing [Mn(2,5-DIHA)]K+ atm/z = 381.06) was always found.

Out of the investigated systems, only the air oxidation of theMn(II)-complexes with the two siderophores was found to resultunambiguously in the stoichiometric formation of the correspond-ing Mn(III) complexes.

As representatives the spectra registered for the Mn–DFBsamples at pH 8 are shown in Fig. 4.

Fig. 4 shows a quite slow oxidation of the DFB-chelated Mn(II) toMn(III) at pH 8 (it almost completes in ca. 3 h), but, as it can be seenin the Inset in Fig. 4, the process becomes really fast at pH 9.5 (com-pletes in ca. 10 min). In this latter case, after 10 min, according tothe ESI-MS result, the Mn(III)-containing complex, [MnHDFB]+ (m/z = 613.28) unambiguously predominates. This complex is reallystable, any degradation (any change) of it was not found within2 weeks (followed by UV spectrophotometry). The behaviour ofthe DFC-containing system is completely the same. Also, existenceof a very stable Mn(III)–DFC ([MnDFC]K+, m/z = 859.27) complexwas supported by ESI-MS following the oxidation (as it is shownby the representative ESI-MS spectrum in Fig. 5).

Because the stability of the siderophore containing complexeswith Mn(III) can be close to those with Fe(III), manganese mayinfluence the siderophore-mediated uptake of iron(III) by microbes

a b 859.265 859.268

860.267 860.271

861.266 861.271 862.268 862.273

Fig. 5. Measured (a) and calculated (b) ESI-MS spectra assigned to [Mn(III)–DFC]K+

at pH 9.5 at 1: 1 metal-to-ligand ratio, cDFC = 5 � 10�4 mol dm�3.

-4.0

-2.0

0.0

2.0

4.0

6.0

8.0-0.4-0.20.00.20.40.6

E [V]

I [

mic

ro A

]

Fig. 6. Cyclic voltammogram registered for Mn(III)–DFC at pH 9.45 (GC workingelectrode and Ag/AgCl/3 mol dm�3 KCl reference electrodes; scan rate 100 mV s�1,cKNO3 = 7.5 � 10�2 mol dm�3, cDFC = 2.25 � 10�3 mol dm�3).

Table 4Formal potentials and overall stability constants (log b) of the complexes formed inMn(III)–DFB and -DFC systems at pH �9.5.

L e0 (V) log b

[MnHL]+ [MnL]

DFB 0.47 ± 0.01 36.5 ± 0.2 �DFC 0.43 ± 0.01 � 28.1 ± 0.3

O. Szabó, E. Farkas / Inorganica Chimica Acta 376 (2011) 500–508 507

[11–13]. This previous assumption initiated our cyclic voltammet-ric (CV) measurements to determine the stability constants for thetris-chelated Mn(III)–DFB and -DFC complexes.

The measurements and the calculations were performed asdetailed in the Experimental section and one representative vol-tammogram relating to the Mn(III)–DFC system is shown in Fig. 6.

Fig. 6 demonstrates the presence of both the anodic and catodicpeaks in the voltammogram of [Mn(DFC)]. This result and also theratio of the peak currents (ia/ic = 0.96) presumably corresponds toreversible reduction of the complex. The situation is different withDFB, where a catodic peak (Epc) at ca. 230 mV and a large anodicpeak (Epa) (dependent upon the concentration of DFB) at ca.400 mV versus Ag+/AgCl appears. This finding is in complete agree-ment with the previous one in Ref. [15], where irreversible electro-chemical behaviour of the Mn(III)–DFB from these results wasassumed. Taking this into account, in this case the Epc was used tocalculate the approximate formal potential and the stabilityconstant. The calculated values are shown in Table 4.

The stability constants in Table 4 (especially the value of the[Mn(DFC)]) are, in fact, comparable with those of the correspond-ing Fe(III)-complexes [12,25]. The numerical value for the overalllogarithmic constant of [Mn(HDFB)]+ also involves the protonationconstant of the terminal amino moiety, so if we take this intoaccount, the constant for the [Mn(DFC)] is higher with ca. 3 logunits compared to that for the Mn(III) + HDFB equilibrium. Theonly constant (log K), what can be found for the Mn(III) + HDFBequilibrium in the literature [11], is also higher (28.6 at 0.1 M NaCland 298 K) than our value in Table 4, but this can be explained ifthe different conditions and also the above-mentioned electro-chemical behaviour of the Mn(III)–DFB is taken into account.

4. Conclusion

Although all of the investigated hydroxamates show well mea-surable Mn(II) bonding ability, with monohydroxamates, followingthe formation of mono- and bis-chelated complexes, the hydrolysisstarts at high pH. On the other hand, formation of tris-chelated

species hinders the hydrolysis if the ligand is either a dihydroxa-mate or a trihydroxamate based siderophore. The stability of thecomplexes corresponds in all cases to the Irving–Williams trendand, at the same time, follows the basicity of the coordinating do-nors. In the case of dihydroxamates, the length and structure of theconnecting chains are also determinant to modify the stability ofthe complexes, however, the most interesting difference can beseen between the stability of the tris-chelated complexes of DFCand DFB (log bMnL values are 8.28 and 6.81, respectively, see Table2). The more effective Mn(II) binding ability of DFC compared toDFB might be attributed both to the bit longer connecting chainsand also the possibility for some conjugation between thehydroxamates and the double bonds in the former ligand.

The relaxivity values for the tris-chelated complexes formedwith siderophores are significantly lower than that for the[Mn(H2O)6]2+ species, but still high enough to assume presenceof water molecule in the inner-sphere. Surprisingly, the corre-sponding values for the bis-chelated complexes (these speciesdominate at physiological pH) are very high, much higher thanthe value for the [Mn(H2O)6]2+. Most probably, water mediated in-tra- or/and intermolecular prototropic exchange processes areresponsible for the large increase of these relaxivities.

The oxidation of Mn(II) to Mn(III) is facilitated in hydroxamatecomplexes under basic condition (pH > 7.5), but becomes stoichi-ometric only in the tris-chelated species formed with siderophores.These Mn(III)-complexes of high stability formed this way do notshow any degradation for days (it was monitored by spectropho-tometry for 2 weeks). However, again, there is a strange differencebetween both the stability and redox behaviour of the Mn(III)–DFCand Mn(III)–DFB complexes. The former one is more stable (closeto the stability of the Fe(III)–DFC complex), and its reversiblereduction to the corresponding Mn(II) species is unambiguouslyindicated by the cyclic voltammetric results. However, the processseems irreversible with DFB and the reason of this difference isgoing to investigate in our laboratory.

Acknowledgements

The work was supported by the Hungarian Scientific FundOTKA-NKTH CK77586 and TÁMOP 4.2.1./B-09/1/KONV-2010-0007. The authors say thank Prof. István Pócsi for providing desfer-ricoprogen, and Dr. Gyula Tircsó for valuable discussions concern-ing the relaxometric studies.

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