INDIAN J. CHEM., VOL. 17A, MAY 1979

6·8

6·4

8·9

~ 11K

Fig. I-Plots of Log Kl against Kp for the mental com--plexe3 of substitued _methoxybenzoic acids with dioxoura-:nium(VI) (0), Be(H) (e) and Cu(H) (.1.). [(1), 2-methoxy-benzoic acid; (2), 4-methoxybenzoic acid; (3), 2,3-dimeth?xy-benzoic acid; (4), 2,4-dimethoxybenzoic acid: (5), 2,6-dime--tboxybenzolc acid; (6), 2,4,5-trimethoxybenzoic acid; (7),

2,4,6-trimethoxybenzoic acid]

9·2

group has very little contribution towards chela-'tion, when singly substituted in position 2. .

For all oxygen donors the Basolo-Pearson stabi-. 1ity order is expected to run parallel to the respec-tive z2Jr values, i.e. dioxouranium(VI) (20'00»Be(II)(11'43) >Cu(II) (5'56)10. The higher stability ofthe Be (II) chelates compared to that of Cu(II)chelates may be due to the smaller radius of theformer and its greater affinity for the oxygen do~~t-.ing ligands. As expected the values of stability.constants reported in the Table 1 show that the-order of stabilities is maintained. Similar observa-·tions have been reported by Mahadevan et alP.

The authors thank Prof. D. D. Khanolkar and·Prof. K. A. Thakar for their interest and encourage-ment. One of the authors (B.D.M.) is thankfulto the authorities of Marathwada University for'the award of a junior research fellowship.

.References1. YASUDA,M., YAMASAKI,K. & OHATAKI,H., Bull. Chem.

Soc. Japan, 33 (1960), 1067.2. MAY, W. R. & JONES, M. M., ]. inorg. nucl, Chem., 24

(1962), 511.3. BERNSTROM,B. H., Acta Chem. scand., 10 (1956), 174.4. VOGEL, A. 1., A text book of practical organic chemistry

(Longmans, Green, London), 1956: 177. .'5. VOGEL,A. 1., A text book of quanhtatwe inorganic analysts

(Longmans, Green, London), 1973.fl. VAN UITERT, L. G. & HAAS, C. G.,]. Am. chem, Soc., 75

(1953), 451.7. IRVING,H. M. & ROSSOTTI,H. S., J. chem. Soc., (1954),

2904.8. PETHE, L. D. & MALI,B. D., Indian]. Chem., 16A (1978),

364.9. JONES,J. G., POOLE,J. B., TOMKINSON,J. C. & WILLIAMS,

R. J. P., ]. chem, Soc., (1958), 2001.

.532

,.

(

10. SIENKO,M. J. & PLANE, R. A., Physical inorganic chemis-try (W. A. Benjamin Inc., New York), 1963, 124.

11. MAHADEVAN,N., SHETTY,S. Y. & SATHE,R. lVI:., Indian]. Chem., 9 (1971), 258.

Stability Constants & Compositions ofComplexes of .1-Anthraquinonyloxamic Acid &

- Its Nitro Derivative with Ce(III), Th( IV) &U(VI)

K. A. IDRISS*, M. M. SELEIM & A. Z. ABU ZUBRIDepartment of Chemistry, Faculty of Science

, Assiut University, Assiut, Egypt

Received 6 May 1978; revised 12 December 1978;accepted 5 ] anuary 1979

l-Anthraquinonyloxamic acid and its 4-nitro deri-vative form 1 : 1and 1 :2 complexes with Ce(III), Th(IV)and U(VI) as revealed by spectrophotometric methodsand conductance measurements. The values of theapparent stability constants of the different complexeshave been determined by continuous variation moleratio, straight line and limiting logarithmic methods.Solid chelates have been isolated and their structuresestablished by PMR and IR spectral data.

IN continuation of earlier workers=" from theirlaboratories on the coordination compounds of

amino-substituted anthraquinones, we report inthis note the results of spectrophotometric andconductometric studies on the reaction of Ce(III),Th(IV) and U(VI) with l-anthraquinonyloxamicacid (AOX) and its 4-nitro derivative (NAOX). .

The solutions (10-3M) of the ligands and of per-chlorates of Ce(III) and U(VI) (10-2M) and thoriumnitrate (10-2M) were prepared in absolute ethanol.The metal ion concentrations were determined byconventional methods". The absorption spectra ofsolutions Were recorded on a UNICAM S.P. 8000spectrophotometer using 1 ern matched quartzcells. Conductometric measurements Were carriedout using a Pye conductance bridge and a dipping-type conductivity cell.

The solid complexes Were prepared by refluxing amixture containing the metal ion solution (1 mol)and the organic ligand (2 mol) in ethanol for ,.....,15min. The solid complexes which separated oncooling are characterized by a brownish yellow,pale and dark green crystals in Th(IV), U(VI) andCe(III) complexes respectively.

The visible spectrum of AOX in ethanol exhibitsmaximum absorption at 395 nm. The band under-goes an apparent blue shift on adding the metalion solution. This is presumably due to the block-ing of the resonance between the donor (-NH)group and the anthraquinone nucleus by the metalion.

AOX-Ce(III) complex exhibits maximum absor-ption at 365 nm and slight blue shift (,.....,10nm) isobserved on raising the concentration of the organicligand in solution. The spectra of Th(IV) andU(VI)-AOX complexes show absorption bands at360 and 375 nm respectively.

The NAOX itself, (in ethanol), shows an absorp-+ion band at 495 nm, the position of which suffersan apparent shift to shorter wavelength on theaddition of the metal ions. The spectrum of'Ce(1II)-NAOX exhibits two maxima at 360 and,470 nm. At higher concentrations of ,the compo-nents .an apparent blue shift of the t,,:o bands is-observed, indicating the probable formation of morethan one type of complexes. Th(1V)-NAOX showsabsorption bands at 350 and 450 nrn whereas theU(V1)-NAOX shows two maxima at 355 and 475nm.

The stoichiometry of the complexed ligands in'solutions was determined by Job's continuousvariation method? and also by applying molar ratio",straight line", slope rat io-? and .lin:iting logarithmic:1'methods. The continuous variation and mole ratiomethods indicate the formation of two types of,complexes foi the metal ions investigated withstoichiometric ratios 1: 1 and 1: 2 (M: L).

The apparent stability constant ~n is determinedirom the results of the molar ratios, straight Iine",,continuous variation 7, and limiting logarithmic'!methods using the relation:

A/Am -~n = (1-AJAm)n2CifL

where A = absorbance, Am = absorbance at satura-tion, CHLconcentration of the ligand and n = num-ber of ligand groups in the complex.. .

For the straight line method, the relation appliedfor the determination of ~n is as follows:

~ = (log C,!+.CHL)/(log A-n log V)where V is the volume of the mixed solution and-Cx+ is the concentration of the metal ion.

mThe mean values of log ~ as well as the valu~s of~G+ are given in Table 1. The results obtainedindicate that the complexes of AOX and NAOXexhibit moderate stability, their stability orderbeing Ce(1II)> Th(1V)> U02(II). .

The compositions of the complexes forme~. 111

'solution were further confirmed by conductivitymeasurements, which show a progressive Increasein specific conductance <l:ssociated wit~ a decre~sein pR during the ritratron o~ metal Ion solution(10-4M) with that of the ligand (10-3M). The

'TABLE 1 - ApPARENT STABILITY CONSTANTSOF AOX ANDNAOX COMPLEXES

Complex log -~Gt~ !fcal mol-1

Ce(III)-AOX 4·9 6·7Ce(III)-2 AOX 9·6 13·1Th(IV)-AOX 4·5 6·1Th(IV)-2 AOX 8·8 11·9U(VI)-AOX 4·1 5·6U(VI)-2 AOX 8·4 • 11·4

Ce(III)-NAOX 4·4 6·0Ce(III)-2 NAOX 8·9 12·1Th(IV)-NAOX 4·07 5·5Th(IV)-2 NAOX 8·3 11·3U(VI)-NAOX 3·9 5·3U(VI)-2 NAOX 7·8 10·6

(

NOTES

( X-2)+

R R

o o

curves obtained are characterized by breaks at themolar ratios 1: 1 and 1: 2 (metal/ligand). Theinitial increase in conductance is probably due torelease of hydrogen during the formation of thecomplex. Further increase in conductance afterthe break may be due to the partial acid dissocia-tion of the free reagent molecules, the latter isassumed to be enhanced by the masking of thedissociated ligand in the complex. The formationof 1: 1 and 1: 2 complexes is in accordance withthe elemental analyses data of the solid complexes.

The R NMR spectrum of the ligands show a D20exchangeable signal in the range 9-12 ppm, ascrib-able to the proton of the carboxyl group. Thissignal is not present in the spectra of the complexesindicating that the proton of the COOR group isreleased during the formation of the complex.

The 1R spectra of the ligands (in KBr) exhibitbands in the vicinity of 3450 crrr ', due to thehydrogen bonded \lNR. These bands are onlyslightly altered When the ligand is bonded withthe metal ions. The bands at 3150, 1740 and 1630cm-1 in the spectra of AOX or NAOX have beenassigned to \lOR, \lC=O of the -COOR group and\lC=O of the quinonoid structure respectively.The spectrum of NAOX possesses bands in thevicinity of 1560 and 1370 cm! which can be assignedto the stretching vibrations of the C-N02 group.

It is worthy to mention that the 1R spectra ofthe metal chelates of both the ligands show bands.-,,1630 which can be assigned to \lC=O of theanthraquinone nucleus. The broad bands .-,,3150due to \lOR of the ligands are no longer presentin the spectra of the complexes.

From the foregoing results we can propose themode of coordination of the metal ions with theanthraquinonyloxamic acids in the 1: 2 type ofcomplexes as represented in structure (I).

Reference's

1. IDRISS, K. A., Issx , 1. M. & SELEIM, M. M., J. appl. Chem,Biotechnol., 27 (1977), 549.

2. ISSA, 1. M., IDRISS, K. A. & SELEIM, M. M., Mh. Chem.,108 (1977), 1461.

3. IDRISS, K. A., Issx , 1. M. & SELEIM, M. M., Indian J.Chem., 15 (A) (1977), 918.

4. IDRISS, K. A., IssA, I. M., SELEIM, M. M. & KH.UIL,M. M., J. apPl. cu-«, (accepted).

5. IDRISS, K. A., IssA, 1. M. & SELEIM, M. M., u»: Chem.,(accepted).

6. SCOTT, W. & FURMAN: H., Standard methods of chemicalanalysis (Van Nostrand, New York), 1962.

7. JOB, P., Ann. Chim., 10 (1928), 113; SHERlEF, F. G. &AWAD, A. M., J. inorg. nucl, Chem., 24 (1962), 179.

\\

INDIAN J. CHEM., VOL. 17A, MAY 1979

8. YOE, G. H. & JONES, A. L., Ind. Engng Chem, Analyt,Edn, 16 (1944), 111.

9. ASMUS, E. Z., Analyt. Chem., 178 (1960), 104.10. HARVEY, A. E. & MANNING, D. L., J. Am. chem, Soc.,

72 (1950), 4488.11. BENT, A. E. & FRENCH, C. L., J. Am. chem, Soc., 63

(1941), 5681.

Mechanism of Oxidation of Acetoin &Benzoin by Hexacyanoferrate(III) in Alkaline

Medium

R. SINGH & E. B. SINGH*

Department of Chemistry, Magadh University. Bodh Gaya

Received 10 July 1978; accepted 18 November 1978

Kinetics of oxidation of acetoin and benzoin byhexacyanoferrate(III) in the presence of aqueousammonia show zero-order dependence on hexacyano-ferrate(III) and first-order dependence on substrateand hydroxyl ion in both the cases. The thermo-dynamic parameters have been calculated. A suitablemechanism best fitting the kinetic data is proposed.

IN recent years potassium hexacyanoferratef l Hjv+!has been widely used as an oxidant in kinetic

studies. Marshall and Watersl2.13 have studied thekinetics of n-glucose, acetoin and benzoin oxidationby Cu(II) and have confirmed the results obtainedby Singh et at.1" in case of n-glucose, In thepresent study. the re.sults obtained in the oxidat~onof acetoin and benzoin by hexacyanoferrate(III) IOnin mild alkali (aqueous ammonia) are reportedand a suitable mechanism proposed.

Acetoin solution was prepared in water byweighing. while benzoin (E. Merck) solution wasprepared in 80% ethyl alcohol. All other reagentswere prepared as usual-'. The progress of. thereactions was observed by the method descnbedearlier+'. Two moles of hexacyanoferrate(III) were·required for one mole of ac~toin or benzoin .to giveacetil or benzil as the reaction product, which wasalso observed by Marshall and W.atersI2.~3:.

Ten-fold (in the case of aceto~) van<l:tI~n and. five-fold (in the case of benzom) variation of

hexacyanoferrate(III) ion concentration at fixedconcentration of other reactants give constant(-dc/dt) values (calculated from the slopes of line;;.obtained on plotting unreacted ferricyanide againsttime) and K, (standard velocity constant) valuesindicating, clearly zero-order dependence of oxidationof both acetoin and benzoin on hexacyanoferrate(III)ion (Table 1). In the case of benzoin. ko valuesdecrease at the end which might be due to transform-ation of oxidation product benzil to benzilic acid.

Ten-fold variation of acetoin concentration gives.nearly constant values of ks/[Acetoin] (Table 2)indicating first-order dependence on acetoin. Inthe case of benzoin. isolation method was usedand kl values (Table 2) obtained are practicallyconstant showing thereby that the order of thereaction is unity with respect to benzoin.

Variation of ammonium hydroxide concentrationshows that ks/[NH40H]t values (Table 3) are fairlyconstant indicating thereby that rate is directlyproportional to square root of [NH"OH] in theoxidation of both acetoin and benzoin. The dis-sociation constant (Kb) of NH"OH is related to[OH-] in the following manner at equilibrium when.[NHt] = [OH-] ; [OH-] = (Kb [N,H"OH])t.

This along with the results presented in Table 3~thus indicate that rate of oxidation of both acetoinand benzoin is directly proportional to [OH-]. The'values of catalytic constants (kOH- and kNHOH)'obtained from the plots between ks/[substrate] ~nd[NH40H] in different ratios are approximatelysame at the same temperature indicating generalbase catalysis in the oxidation of both acetoin andbenzoin. .

The kinetic data described above lead us tosuggest the following mechanism (Eqs. 1-3).

oII

R.CCH(OH)R+OH(S)

~R-C = ~(OH)R+H20

(E) ... (1)

k, slow-'"..,--k,

where R = CH3 or C6HS' S = acetoin or benzoin,E = enediol anion

6k, I

E+Fe(CN)~- --7 [Fe(CN)~-]+R-C = C(OH)Rfast (R') ... (2,

1·00 1·50

TABLE 1_ EFFECT OF VARYING INITIAL [FERRICYANIDE] ON THE REACTION RATE AT 30°

2·00 3·00 4·00 10·005·00 7·00

[ACETOIN] = 5·00 X 10-2M, [NH.OH] = 5·00 X lC-1M; (L = 0'10M

5·86 5·89 HI 5·55 5·31Ks X 106 mol litre-1 min-1 (calc.)

( _~) X 106 mol litre-1 mirr+

(graphical)

5·60 5·60

5·10

5·30

4·07

5·3()5·50 5·50 5'50

[BENZOIN] = 5·00 X 10-3M; [NH.OH] = 5·00 X lo-2M; (L = 0'10M

1.20 1·26 1·24 1·20 1·23 1·14K. X 106 mol litre-1 mirr+ (ca.lc.)

( _ ~) X 106 mol litre-1 mirr+

(graphical)

1·25 1·28 1·25 1-19 1-13. 1·25

534

(

"