摘要:
J. CHEM. soc. DALTON TRANS. 1995 477Role of the NH,' Moiety in lron(iii)-, Aluminium(iii)- andGallium(iii)-Aminohydroxamate InteractionsEtelka Farkas,*#a Emese Kozma,a Tamas Kiss,a lmre Totha and Barbara Kurzakba Department of Inorganic and Analytical Chemistry, Lajos Kossuth University, H-40 10 Debrecen, HungaryInstitute of Chemistry, Pedagogical University, Siedlce, PolandStability constants have been determined and the bonding modes and effects caused by the side-chainNH,+ moiety in aminohydroxamic acids evaluated for complexes formed in aqueous solution in betweeniron( iii), aluminium(iii) and gallium(lll) with a-alaninehydroxamic acid (a-Alaha), P-alaninehydroxamicacid (P-Alaha), aspartic acid-P-hydroxamic acid (Asp-P-ha) and glutamic acid-y-hydroxamic acid (Glu-y-ha).The iron(iil)-, aluminium(lll)- and gallium(lIl)-acetohydroxamic acid (aha) systems were studiedas models. Co-ordination of hydroxamate oxygens occurs in the cases of aha, a- and P-Alaha, while Asp-p-ha and Glu-y-ha are co-ordinated via their hydroxamate and carboxylate oxygens. The OH- ion wasfound to be an effective ligand in these systems (especially for Gaili) causing the formation of bothbinary and ternary hydroxo complexes. The presence of NH,+ in the hydroxamic acids favours thehydrolysis to an extent which depends on the distance between the hydroxamate moiety and NH,+.These findings can be explained by the electron-withdrawing effect of N H,+ and electrostaticrepulsion between it and the co-ordinating M3+ ion.It is well known that siderophores synthesised by micro-organisms, containing primarily either hydroxamate orcatecholate groups as chelators, are able to protect the classical'hard' metal ion Fe"' against hydrolysis.' The same behaviourhas been found with interesting synthetic analogues prepared indifferent l a b o r a t ~ r i e s . ~ ~ The simple monohydroxamic acidsare also effective ligands of Fe"' and other high-oxidation-statecations. However, aminohydroxamic acids, such as hydroxamicacid derivatives of simple amino acids, exhibit significantdifferences in their complex formation as compared with themonohydroxamic acids. They are able to form chelates via theamino and hydroxamate nitrogens; this type of chelate ispreferred by many metal ions,5 except by typical hard metalions, which prefer chelation through the hydroxamate oxygens.The very few papers relating to the complexes formed with hardmetal ions reveal contradictions and unsolved problems.Inour laboratory, studies have already been performed on thecomplexes in the Fe"'- and A111'-or-Alaha,6+f Fen'-Asp-P-ha 'and Fe"'-Glu-y-ha systems.* The surprising result was obtainedthat hydrolysis of the metal ion is much more suppressed inthe systems containing Asp-P-ha and Glu-y-ha, than in the caseof a-Alaha, where numerous mixed hydroxo species are foundeven if a fairly large excess of hydroxamic acid is used.Moreover, the complexes formed with Glu-y-ha were found tobe more stable than those of Asp-P-ha. The present work wascarried out to clarify these, at first sight, unexpected results;detailed equilibrium and structural studies were performed onthe complexes of Fe"', Al"' and Ga"' formed with acetohy-droxamic acid (aha) as reference, and with a-Alaha, P-Alaha,Asp-P-ha and Glu-y-ha, using pH-metric, 'H, 27Al NMRand spectrophotometric methods. For comparison, resultspublished earlier 5-8 are included as reference data.ExperimentalThe compounds Asp-P-ha, Glu-y-ha and aha were obtained7 For example, a-Alaha = a-alaninehydroxamic acid (2-amino-N-h ydrox ypropanamide).from Sigma, while a-Alaha and P-Alaha were prepared via themethyl ester of the respective amino acid.' The purities of thehydroxamic acids and the exact concentrations of the stocksolutions were determined by the Gran method. loSolutions of Al"' and Fe"' were prepared by dissolving anappropriate amount of the metal chloride in doubly distilledwater or in hydrochloric acid of known concentration.Theirexact concentrations were determined gravimetrically viaprecipitation of the quinolin-8-olate or oxide. Gallium(Ir1)solution was prepared by dissolution of the pure (99.99%) metalin a known amount of HCl. The exact concentration of themetal was established by complexometric titration withethylenediamine-N,N,N',N'-tetraacetate (edta) and the amountof acid was determined by the method of Harris and Martell."pH-Metric titrations were performed at five or six differentmetal ion: hydroxamic acid ratios in the range 1 : 1-1 : 10. Theconcentrations of the hydroxamic acids were (1-3) x lop3 moldmP3.All pH-metric measurements were carried out at25.0 f 0.1 OC, at an ionic strength of 0.2 mol dm-3 KCl, withKOH solutions of known concentration (ca. 0.2 mol dm-3). Theionisation constant of water under these conditions is lo-' 3-756mo12 dmp6. The measurements were made with a RadiometerPHM 84 instrument with a GK 2421 C combined electrode, anda Radiometer PHM 64 instrument with a GK 2401 B electrode,throughout the range pH 2.0-1 1 .O, using 10.00 cm3 samples.The electrode system was calibrated by the method of Irving elal. l 2 so that the pH-metric readings could be converted intohydrogen-ion concentrations. Concentration stability constants(Ppqr = [MpA,Hr]/[M]p[A]q[H]r) were calculated from pH-metric data with the aid of the PSEQUAD computerprogram,13 using literature data on the metal hydroxocomplexes. ' 4-1 'Visible absorption spectra were recorded on a Beckman ActaMIV double-beam recording spectrophotometer over the rangeca.300-800 nm. The concentration of Fe"' was 5 x lo4 or1 x mol dm-3, and the metal : hydroxamic acid (aha or p-Alaha) ratio was 1 : 20 or 1 : 10.The 'H NMR measurements on the A1"'-aha system weremade on Bruker AM400 and WP 200SY FT instruments. Thealuminium(1n) concentration was 2 x mol dm-3, and th478 J. CHEM. SOC. DALTON TRANS. 1995Table 1 Overall stability constants (log p) for the hydroxamic acids and hydroxo metal ions at 298 K and Z = 0.2 rnol dm-3 (KCl)Species aha" a-Alaha P-Alahab Asp-P-ha' Glu-y-ha"HA 9.26 9.16 9.42 9.42 9.50- 16.50 17.91 17.71 18.05- - - 19.83 20.26H2AH A[MH-J2 + - 3.21CMH-21+ - 6.73CM,H-,I~ +- 4.09CM3H-41 +- 7.58CMH-31[MI 3% 21+-CMH4I ---" Ref.8. Ref 19.' Ref. 7. Ref. 15. Ref. 16. Ref. 17.-5.52 -2.46- - 5.92-7.70- 13.57- - 10.63-23.46 - 16.87- 109.1 -Al"': aha ratio was 1 : 3 or 1 : 5. The 27Al NMR spectra wererecorded on a Bruker AC200 spectrometer equipped with aselective broad-band probe and operating at 52.148 MHz. Thechemical shift reference was 0.1 mol dmP3 aqueous AlCl,. Theconcentration of the Al"' was 0.04 mol dm-, and the Al"' : aharatio was 1 : 3. All NMR measurements were made in D,O. ThepD values were calculated via expression ( l).I8Results and DiscussionDissociation constants for all of the hydroxamic acids involvedin the present work had already been determined in ourNot only the macroconstants, but also thedissociation microconstants, are available for a- and P-Alaha.The macro dissociation ccnstants are listed in Table 1, whichalso contains stability constants for metal ion-hydroxo com-plexe~.'~-'~ The tabulated data reveal a strong tendency tohydrolysis in the case of these metal ions, and especially Ga"'.For this reason different hydroxo species were included in ourspeciation models.Initially they were included with fixedstability constants, but if their concentrations were found to besignificant in a system these values were refined too {as was thecase with the gallium(rI1)-hydroxo complexes and [Al(OH),] - } .Metal Ion-aha Systems.-The model aha (CH,CONHOH)contains only one hydroxamate moiety for co-ordination.Threesuch ligands, each forming a five-membered chelate, saturatean octahedral co-ordination sphere in stepwise processes. Inaddition to pH-metry, various spectroscopic measurements(UV/VIS for Fe"'-aha, and 'H and 27Al NMR for Al"'-aha)were made to establish the best equilibrium model.The absorbance spectra were registered as a function of pH atFe"' : aha ratios of 1 : 10 and 1 : 20. The spectra exhibit maximumabsorbance in the ranges pH 5-7.5 and 5-8, respectively, withE, = ca. 2800 dm3 mol-' cn-' at h,,, = ca. 420 nm. This E,,,value suggests the co-ordination of three hydroxamates periron(II1) in this pH region.Above this pH range adecrease in the absorbance is observed, indicating the formationof complexes containing less than three hydroxamates (mixedhydroxo complexes). Decreasing the ligand excess favourshydrolysis. If the hydroxamic acid : metal ion ratio is 1 : 5 or less,precipitation occurs at below or ca. pH 7.In the calculations on FeI'I-aha the experimental pH-metricpoints were used taken from pH ranges where'measurablehydrolysis did not occur. The models yielding the best fitting,together with the related constants, are given in Table 2.Evaluation of the experimental pH-metric data on Al"'-ahademonstrated that two different speciation models yieldedacceptable fittings. One involved only the species [MA12 +,Table 2 Stability constants (log p) for the complexes of Fe"', Ga"' andAl"' formed with aha at 298 K and Z = 0.2 mol dm-3 (KCl)MSpecies Fe"l * Gall' Al"'[MA21 +CMA2H-,I - - 1 0.40( 3)CMA2H-21- - - 1.04( 5 )[MA12+ 1 1.42 9.51(5) 8.15(2)21.10 18.42(4) 15.77(2)[MA31 28.30 26.21(5) 21.5(2)* Ref.8.[MA2]+ and [MA,], in addition to the binary hydroxocomplexes (see Table l), while the formation of mixed hydroxospecies ([MA2H-J and [MA,H-,]-) too was postulated in thesecond model. To differentiate between the two models, 'H and27Al NMR measurements were carried out. The 'H NMRmeasurements were performed in the same concentration rangeas used in pH-metry. For comparison, the pH dependence of thespectrum of aha (containing only one singlet resonance line)was also measured.When the isotope effect is taken intoaccount,21 the pK, calculated from the pD dependence of thechemical shifts for aha correlates well with the value determinedby pH-metry. The A1"':aha ratios were 1 : 3 and 1 :5. Somerepresentative spectra recorded at different pH values at a ratioof 1 : 3 are depicted in Fig. 1.As Fig. 1 shows, resonance lines with different chemical shiftsare to be found for the aha species throughout the wholemeasured pH range. The separated signals of free aha (6 1.94and 1.80 for HA and A-, respectively) clearly demonstrate a'slow exchange regime' between the free and co-ordinatedhydroxamic acid. At lower pH values [Fig. l(a)-(c)] more thanone signal of the complexes can be seen.It can be explained by'slow exchange' on the 'H NMR time-scale not only betweenthe free hydroxamic acid and the complexes but also for theligand exchange between the different complexes. At higher pH[Fig. l(c)-(e)] some broadening of the signal of the complex orcomplexes can be seen. This cannot be attributed to a mutualexchange with the free aha, because it does not broaden. Thebroadening observed might be related to some chemicalexchange between different complexes, but the appearance oftwo signals with almost identical chemical shifts cannot beruled out. A detailed study of the ligand-exchange reactions isunder progress.The above results and fittings give greater support to thesecond speciation model. Representative concentration distri-bution curves are depicted in Fig.2. From a comparison of theresults in Figs. 1 and 2, the species [MA]", [MA,]' and[MA,H-,] - can easily be identified, but in the pH range wherethe parallel existence of [MA,] and [MA,H-,] can be assumeJ. CHEM. soc. DALTON TRANS. 1995 4791 I I I I I2.2 2.0 1.8 2.2 2.0 1.86Fig. 1 The 400 MHz 'H NMR spectra recorded for Al"'-aha at 1 : 3ratio at different pD values: 3.46 (a), 4.43 (b), 6.37 (c), 7.66 (d), 8.47 (e),and 1 0 . 2 2 ( f )" 2.0 4: 0 60 8.0 10.0 12.0PHFig. 2 Concentration distribution curves for complexes formed in tneAl'"-aha system at 1 : 3 ratio at cAl = 2 x mol dm-3from pH-metry only one NMR peak is observed at 298 K.However, at 273 K (where chemical exchange is even slower)it is split into two lines (a representative example is given inFig.3), indicating the existence of at least two differentaha-containing species.Aluminium-27 NMR spectroscopy requires much higheranalytical concentrations than those sufficient for 'H NMR,and this causes significant changes in concentration distribution(the hydrolytic reactions are appreciably hindered in the lattercase), as is seen on comparison of the distribution curves inFigs. 2 and 4. The NMR spectra show that Al"' can be detectedin two extremely different chemical environments in the rangepH 5-9.5. As an illustration, a "A1 NMR spectrum is shown inFig. 5. The relatively sharp band at 6 36.7 indicates Al"' in afairly symmetrical complex; this can probably be ascribed to theoctahedral tris complex [AlA,], while the very broad band (ca.3000 Hz) with maximum at around 6 56 is characteristic ofmixed hydroxo complexes having much lower symmetry.I 1 I I2.25 2.00 1.75 1.506Fig.3 Temperature dependence of 200 MHz 'H NMR spectrarecorded for A1"'-aha at 1 : 3 ratio, at pD 6.4: (a) 298, (b) 273 KI I , I 12.0 4.0 6.0 8.0 10.0 12.0Fig. 4 Concentration distribution curves of complexes formed in theAl"'-aha system at 1 : 3 ratio at cA, = 4 xPHmol dm-3Formation of the tetrahedral, very symmetric species[Al(OH),] - can also be detected by the observation of a sharpsignal (ca. 5 Hz) at 6 80.8 at pH > 10. These conclusionsare in good agreement with the speciation model listed inTable 2.All the above findings show that aha is an effective chelator ofthese metal ions and suppresses their hydrolysis to an extentdetermined by the metal : aha ratio, the analytical concentrationand the pH.The highest stability constants were foundfor iron(1rr) complexes and the lowest those for A]"'. Thesame stability sequence was observed earlier for otherhydroxamates.'Metal Zon-ac-Alaha [CH,CH(NH, +)CONHOH] and -p-Alaha (NH, +CH,CH2CONHOH) Systems.-The system480 J. CHEM. SOC. DALTON TRANS. 1995I I I I120 80 40 06Fig. 5 A 27Al NMR spectrum recorded for A1"'-aha at 1 : 3 ratio at ca. pD 7.07 6350 450 550 6503JnmFig. 6 Visible spectra registered in the Fell*-P-Alaha system at 1 : 10ratio, at different pH values: 1.97 (l), 2.94 (2), 4.03 (3), 5.03 (4), 6.01 (5),6.94 (6), 8.00 (7) and 10.0 (8). cFe = 1 x low3 mol dm-'Fell'-a-Alaha and Al"'-a-Alaha were studied in our laboratoryearlier.6 It was found that both Fe"' and Al"' prefer chelation atthe hydroxamate binding site and, due to the high hydrolytictendencies of these metal ions, mixed hydroxo complexes toowere formed in the pH range where the unco-ordinated NH,+group of the ligand is still protonated.This leads to theformation of protonated mixed hydroxo complexes.6 Withregard to the above, only the Gar"-containing system wasstudied in the present work with a-Alaha. The results showedthat the formation of gallium(II1) hydroxo complexes is favouredeven in the acidic pH range. Besides the binary hydroxo species,a stability constant could be calculated only for [GaAH13 +.On the basis of this result Ga"' was not included in our studieson P-Alaha.To choose the appropriate equilibrium models for the studiedsystems, we again used spectrophotometry for the Fe"'-P-Alahaand 'A1 NMR spectroscopy for the Al"'-P-Alaha system.Additionally, our previous work on Al"'-cr-Alaha wassupplemented with NMR measurements.The pH dependence of the visible spectrum of the iron(n1)system at a metal: hydroxamic acid ratio of 1 : 10 can be seen inFig. 6 .The value of cmax calculated from the spectral data in therange pH 5-6 at h,,, ca. 420 nm is ca. 2600 dm3 mol-' cm-',suggesting the formation of complexes containing threehydroxamate chelates in quite high percentage. 2o The decreasein absorbance above this pH might be due to the lowering of theaverage number of hydroxamate groups co-ordinated by anOH- displacement reaction.The equilibrium model based on the pH-metric resultsyielded a good fitting and correlated acceptably with thespectrophotometric results.These stability data are listed inTable 3. The model relating to Al"'-P-Alaha (see Table 3) waschosen with the aid of 27Al NMR measurements. A comparativestudy of the A1"'-a-Alaha and -P-Alaha systems revealed littleTable 3 Stability constants (log p) for the Fell1-, Ga"'-" and Al"'-a-Alaha and Fe"'- and Al"'-P-Alaha complexes at 298 K and I = 0.2 moldm-3 (KCI)MFe"'17.1513.92-28.3621.9914.5433.90-Al"'14.3522.2117.5912.635.85- 2.4416.7-__9.62-0.1619.95(6) 16.72(2)38.30(4) 32.07(2)- 27.04(2)- 19.87(2)22.1 l(7) 10.74(2)12.12(8) 0.04( 32)16.99(5) -55.47(4) -50.24(5) -42.75(7) --a Only one complex, [GaAH13+, was formed in measurableconcentration with a-Alaha.The log P value is 16.4(4). Ref. 6.differences in NMR behaviour between the two hydroxamicacids. In both systems only very broad, highly overlapping bandswere observed in the range pH 3-9, with a dominatingmaximum at around 6 60-65, indicating octahedral Al"' in lowsymmetrical environments. In contrast with the Al"'-ahasystem, no sharp signal characteristic of Al"' in a highlysymmetric co-ordination sphere was observed. Thus speciationmodels which included tris complexes involving hydroxamate-type co-ordination (0-, =O) were rejected.The formation of[Al(OH)J- was detected in both systems at pH > 10.0.Comparison of the results calculated for the metal(m)-a-Alaha and -P-Alaha systems both with each other and withthose obtained for metal(rr1 )-aha reference systems led to thefollowing findings. The stability constants for the reaction M +HA MAH, characteristic of the interaction of themetal(m) ion with the hydroxamate binding site, were derivedfrom the overall stability data (log KMAH = log PMAH - logKHA). Since HA contains the amino group in protonated formand the two deprotonation processes of the free hydroxamicacids overlap, the corresponding microconstants had to be useJ. CHEM. soc. DALTON TRANS. 1995 48 1Table 4 Stability constants (log P) for the complexes of Fe", Ga"' and Al"' formed with Asp-P-ha and Glu-y-ha at 298 K and Z = 0.2 mol dm-3 (KC1)Fe"' Ga"' Al"'Species Asp-P-ha Glu-y-ha Asp-P-ha Glu-y-ha Asp-P-ha Glu-y-ha[ MAH] ' + 18.82 18.92 18.33(4) I8.5( 1) 16.27(5) 16.65(5)[MA2H2] + 36.35 36.65 3 5.53(8) 35.0(2) 3 1.76(7) 32.79(7)CMA2HI 3 1.63 31.70 31.61(1) 30.8( 1) 26.80(4) 27.62(5)24.25 24.10 24.23(6) 24.2(3) 1 9.26( 3) 19.89( 5)[MA,H..,]Z- 15.32 14.70 15.6(3) - 10.36(4) 9.9(5)[MA21 -[MA2H-J 3- - - - - - 1.10( 15)a Ref.7. Ref. 8.in these calculation^.^^ The log KMAH values (7.99, 10.53, 5.19and 7.30 for Fe"'-a- Alaha, Fe"'-P-Alaha, Al"'-a-Alaha andAl"'-P-Alaha, respectively) show that the complexes formedwith P-Alaha are more stable than those formed with a-Alaha,and both are less stable than those formed with aha (see con-stants for [MAI2 + in Table 2).This means that the presence ofthe NH, + group in the side-chain decreases the stability of thehydroxamate chelate, and the extent of the decrease depends onthe distance between the NH3+ group and the hydroxamatemoiety. As to the origin of this effect, the following conclusioncan be drawn: at least two effects result from the presence of theNH, + group. First, an electron-withdrawing effect, whichmakes the hydroxamic acid group more acidic, as is known fromthe acid-base properties of these ligands. l9 Secondly, there is anelectrostatic repulsion between the NH3+ group and the co-ordinating M3 + ion. These two superimposed effects result inthe significant difference in the equilibrium models discussedabove for the a-Alaha-and P-Alaha-containing systems.Hydrolysis of the metal ions is more favoured in the presence ofthe former hydroxamic acid.Metal Ion-Asp-f3-ha [HO,CCH(NH, +)CH2CONHOH] and-Glu-y-ha [HO,CCH(NH , + )CH,CH,CONHOH] System.-Besides the chelate-forming hydroxamate moiety, thesehydroxamic acids contain an extra carboxylate group which is anew potential site for co-ordination. Our earlier results on Fell1-Asp-f3-ha and -Glu-y-ha showed that the hydroxamate oxygensare the main donor atoms, but that co-ordination of thecarboxylate group can also be assumed.697 It was also foundthat the non-co-ordinating amino group containing adissociable proton in the protonated complexes deprotonates ina process overlapping with ionisation of the co-ordinated watermolecules. The complexes of Ga"' and Al"' with Asp-P-ha andGlu-y-ha were studied in the present work.The speciationmodels yielding the best fitting of the pH-metric experimentaldata were the same as for the Fell'-containing system^.^.^ Theyinvolve neither polynuclear species nor any complex with ametal : hydroxamic acid ratio of 1 : 3, as can be seen in Table 4.The results obtained for GaI'I-Asp-P-ha, GaI'I-Glu-y-ha, All1'-Asp-P-ha and Al"'-Glu-y-ha are in complete agreement withthose found for FeI'I-Asp-P-ha and Fe"'-Gl~-y-ha.~ Thetridentate co-ordination via the hydroxamate function and thecarboxylate oxygens is the preferred binding mode in both the1 : 1 and 1 : 2 complexes.Since the appropriate dissociation microconstants are notavailable, the equilibrium constants for the reaction M +HA MAH, characteristic of the interaction between themetal(rrr) ion and the tridentate hydroxamic acids, cannot bederived and hence their comparison with those for a- and p-Alaha is not possible.A simple comparison of the data in Table4, however, clearly indicates that the complexes of Glu-y-ha aresomewhat more stable than those of Asp-P-ha. These differencescan again be explained if the effects of the side-chain NH3+group (electrostatic repulsion and electron-withdrawing effects)are taken into account. These effects result in the formation ofmore stable complexes with Glu-y-ha, in which the NH3+group is situated farther away from the hydroxamate moiety,than those with Asp-P-ha in which the two positively chargedcentres are nearer each other.AcknowledgementsThis work was supported by the Hungarian Academy ofSciences (Projects OTKA 04 09604, OTKA 1647 and US-Hungarian Joint Fund No.182/92b).References1 A. Chimiak, R. C. Hilder, A. Liu, J. B. Neilands, K. Nomoto andY. Sugiura, Struct. Bonding (Berlin), 1984,58.2 C. Yuen Ng, S. J. Rodgers and K. N. Raymond, Znorg. Chem., 1989,28,2062.3 I. Dayan, J. Libman, Y. Agi and A. Shanzer, Znorg. Chem., 1993,32,1467; P. Yakirevitch, N. Rochel, A.-M. Albrecht-Gary, J. Libmanand A. Shanzer, Znorg. Chem., 1983,32, 1779.4 M. A. Santos, M. A. Esteves, M. Candida, T. Vaz and M. L. S.Goncalves, J. Chem. SOC., Dalton Trans., 1993,927.5 B. Kurzak, H. Kozlowski and E. Farkas, Coord. Chem. Rev., 1992,114, 169.6 E. Farkas, J. Szoke, T. Kiss, H. Kozlowski and W. Bal, J. Chem. Soc.,Dalton Trans., 1989,2247.7 E. Farkas and P. Buglyo, J. Chem. Soc., Dalton Trans., 1990, 1549.8 E. Farkas, D. A. Brown, R. Cittaro and W. K. Glass, J. Chem. SOC.,9 C. R. Hauser and W. B. Renfrow, jun., Org. Synth., 1943,2,67.Dalton Trans., 1993,2803.10 G. Gran, Acta Chem. Scand., 1950,4,559.1 1 W. R. Harris and A. E. Martell, Znorg. Chem., 1976, 15,713.12 H. M. Irving, M. G. Miles and L. D. Pettit, Anal. Chim. Acta, 1967,38,475.13 L. ZkkPny and 1. Nagypal, in Computational Methods for theDetermination of Stability Constants, ed. D. Leggett, Plenum, NewYork, 1985.14 G. H. Khoe, P. L. Brown, R. N. Sylva and R. G. Robins, J. Chem.SOC., Dalton Trans., 1986, 1901.15 C. F. Baes and R. E. Mesmer, The Hydrolysis of Cations, Wiley,New Ymk, 1976.16 L. 0. Ohman and W. Forschlinn, Acta Chem. Scand.. Ser. A. 1981.171819202122-35, 795.I. Toth, E. Brucher and L. ZCkany, Magy. Kim. Foly., 1984,90,149.R. G. Bates, Determination of pH, Wiley, New York, 1964.E. Farkas, T. Kiss and B. Kurzak, J. Chem. SOC., Perkin Trans. 2,1990, 1255.S. J. Barlcay, P. E. Riley and K. N. Raymond, Znorg. Chem. , 1984,23,2005; S. J. Barclay, B. H. Huynh and K. N. Raymond, Znorg. Chem.,1984,23,2011.R. F. Jameson, G. Hunter and T. Kiss, J. Chem. Soc., Perkin Trans. 2,1980, 1105.A. Evers, R. D. Hancock, A. E. Martell and R. J. Motekaitis, Znorg.Chem., 1989,244,2189.Received 10th August 1994; Paper 4104923
ISSN:1477-9226
DOI:10.1039/DT9950000477
出版商:RSC
年代:1995
数据来源: RSC