Jha, S. K.

Sharma, Sudhindra Swarup

Singh, Mukhtar

English

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TABLE 1 - FORMATION CONSTANTS OFLANTHANIDE-MONOTHIODIBENZOYUIETHANIDESMetal 1/" log Kl lug K. log K3 log [3.ion (A-1)La3t 0·942 4·33 3·95 3·25 11·53Ce3t 0·967 4·54 4·25 4-11 12·90Pr3t 0·987 4·75 4·39 4·33 13·47Nd3• 1·005 4·88 4·51 4·43 13-82Sm3• 1·037 5·25 4·29 4·20 13·74EuH 1·052 5·16 4·44 4·32 13-92GdH 1·066 4·99 3·90 3·67 12·56DyH HOI 5·04 HI 4·05 13-20HuH 1·118 5·11 4·25 HO 13·46ErH H35 5·29 4·29 4-15 13·73Yb3r 1·165 5·35 4·32 4·18 13-85yH 1·136 4·83 3·95 3-84 12·62on the accuracy of the earlier constants and therewill be accumulation of error in the successive steps.The practical acid dissociation constant (PKa =9'34) of HSDBM as determined by us is in excellentagreement with value (PKa = 9'38) reported byManone and Saitoll in 3:1 dioxane-water mixtureat an ionic strength of 0·1. Uhlemann et al.12 re-ported the thermodynamic acid dissociationconstant, PKo of HSDBM in 75% aqueous dioxane,and their value pKo = 11·40 is somewhat higherthan the value PKo = 11,14, reported by Chastonet alP. Fernelius and his coworkers-! have observedthat dibenzoylmethane form relatively strongcomplex with Na", so its monot hio analogue mayalso show a similar trend. Since we were primarilyinterested in the evaluation of practical stabilityconstants of lanthanicle-HSDBM complexes, noattempt was made to determine the thermodynamicacid dissociation constant.Table 1 shows the first three formation constantsof eleven lanthanide and yttrium ions. In general,for a given metal ion, log Kl > log K2 > log K3•However, the difference between any two consecutiveconstants is quite small, indicating that there is analmost equal tendency for all of the complex speciesto coexist in solution. It seems probable that whenn> 2, the concentration of the tetrakis-cornplex[LnA4J- in the equilibrium system becomes appre-ciable. In fact, both the tris- and the tetrakis-complexes of HSDBM have been isolated by USB.The formation constants of 1:1 complexes exhibitthe trend commonly observed in various complexesof Ianthanidesv-P. It is apparent from Table 1that the log Kl values monotonically increase fromLa(III) to Sm(III) with respect to the 1jr values(where r stands for ionic radius). The middle portionof the series show a decline with the Gd(IIl) valueat the minimum. After Gd(III) , there is slow in-crease in log Kl and the value of Sm(III) is almostequal to that of Er(III). It may also be noted thatlog K] of Y(III) falls in between those of Pr(Ill)to Nd(III). The change in a quoion structure anddepression of class-b character in the middle of thelanthanide series appear to govern the sequence offormation constants throughout the series-",The lower stabilities of thio-~-diketonates in com-parison with ~-diketonatel impose greater Lewisacidity or residual coordinating power in thecomplexes. This has been verified" by measuringNOTESthe formation constants of the mixed complexesof Nd(STTAh with several neutral donors whichare found to have greater values than those of theNd(TTAla adducts. Thus it is expected that thelanthanide rncnothio-Bvdiketonates might serve asbetter shift reagents than the currently popular~-jiketonates. Work with suitable reagents in thisdirection is in progress.References1. DUTT, N. K, NAG, K. & SHESHADRI, T., J. inorg. nucl.cu«; 31 (1969), 1435.2. BROWN, D., HOLAH, D. G. & RICKARD, C. E. F., J.chem. Soc., A (1970), 786.3. SIDDALL, T. H. III & STEWART, 'Y. E., J. inorg. nucl.cie«, 32 (1970), 1147.4. PINKERTON, A.A., Tnorg, nucl. chem, Lett., 10 (1974). 495.5. NAG, K. & CHAUDHURY, M_, J. inorg. nucl. Chem., 38(1976). 309.6. NAG, K. & CHAUDHURY, M., J. inorg. nucl. Chem., 39(1977), 1213.7. NAG, K. & CHAUDHURY, M., Inorg. nucl, chem. Lett.,12 (1976), 307.8. NAG, K. & CHAUDHURY, M., Inorg. Chem., 15 (1976), 2291.9. CHASTON, S. H. R., LIVINGSTONE, S. E., LOCKYER,T. N., PICKLES, V. A. & SHANON, J. S., Aust. J.Chem., 18 (1965), 673.10. VAN UITERT, L. G. & HASS, C. G., J. Am. chem. Soc.,75 (1953), 451.11. MANONE, A. & SAITO, K., Bull. chem, Soc. Japan, 41(1968), 2419.12. UHLEMANN, E. & SUCHAN, 'Y. W., Z. anorg, allg. cu,«,342 (1966), 41.13. CHASTON, S. R. H. & LIVINGSTONE, S. E., Aust. J.Chem., 19 (1966), 2035.14. VAN UITERT, L. G., FERNELIUS, W. C. & DOUGLAS,B. E., J. Am. chem, Soc., 75 (1966), 2739; FERNELIUS.,Y- C. & VAN UITERT, L. G., Acta chem. scand., 8(1954), 1726.15. MOELLER. T., MARTIN, D. F., THOMPSON, L. C., FERRUS,R., FEISTEL, G. R. & RANDALL, W. J., cu»: Rev.,65 (1965). 1.16. SINHA, S. P., Europium (Springer-Verlag, Berlin), 1967.17. NIEBOER, E., Structure &- bonding (Berlin). 22 (1975), 1.Influence of Concentration & Nature ofSupporting Electrolytes on the Kinetics ofReduction of Zn(II) at Dropping MercuryElectrodeSUDHINDRA SWARUP SHAR~rA, S. K. JHA & MUKHTAR SINGHChemistry Department, Agra College, AgraReceived 29 December 1976; accepted 25 March 1977The polarographic behaviour of Zn(I1) has beenstudied in a variety of supporting electrolytes of varyingconcentrations, viz, NaCI, NaCIO•• NaN03, KCI, KCIO.,KN03, KN02, KBr, KI, K2SO., KH2PO., KCNS, pot. ace-tate, pot. oxalate, pot. tartrate, pot. citrate, pot. benzo-ate, NH.CI, NlJ.N03• Mg(N03)., Ca(N03)2' Sr(N03). andBa(N03)2' It has been found that the reduction ofZn(lI) at d.m.e, in the presence of pot. oxalate, pot.citrate and O·5M pot. tartrate is totally irreversiblewhile in other supporting electrolytes it is quasi-reversible. The kinetic parameters, viz. the transfercoefficient (a) and standard rate constant (hl,h) ofthe quasi-reversible electrode process have been cal-culated by Gelltngs' method. The influence of con-centration and nature of supporting electrolytes on thekinetics of quasi-reversible electrode process of Zn(II)has been discussed in terms of the values of kl,h and 'a'.923INDIAN J. CHEM .• VOL. 15A. OCTOBER 1977THE nature of the reduction of Zn(II) .at d.m.e.is still regarded a vexed problem Inasmuchas that many polarographic studies, under somewhatsimilar conditions, do not unambiguosly illustrateif the reduction is reversible, quasi-reversible orirreversible in the same supporting electrolyte",This behaviour has also been observed in fluoridesupporting electrolytes's". According to Bond- theconflicting results obtained by various workers onthe polarographic behaviour of Zn(II) could becaused by differences in supporting electrolyte ordrop time, or the presence or absence of a particularsurface-active maximum suppressor.Mukherjee and Chakravarti", and Nath andBhattacharya" have studied the influence of thenature and concentration of some supporting elec-trolytes on El/2 and id of Zn(II). While Mukherjeeand Chakravarti held that viscosity and complexformation are the two important parameters whichaffect El/2 and id, Nath. and Bhattacharya emp~a-sized the role of adsorption at the mercury-solutioninterface in affecting EI/2 and id·From a survey of the literature1-4,7-9 it is evidentthat much remains to be investigated with regardto the nature of the electrode reaction of Zn(II)at d.m.e. in a variety of supporting electrolytesparticularly in complexing supporting electrolytes.With this aim in view. the present study dealingwith the polarographic behaviour of Zn(II) incomplexing and non-complexing supporting electro-lytes of varying concentrations has been under-taken.ZnCl2 and the supporting electrolytes were eitherof AR (BDH) or GR (E. Merck) grade. Theirstock solutions were prepared in conductivity water.The resultant concentration of Zn(II) was main-tained at 1·0 X 10-3M in each case.A manual polarograph (Toshniwal, CL02) inconjunction with a polyflex galvanometer (Toshni-wal, PL50) was used. The d.m.e., used in conjunc-tion with SCE, had the following characteristics:hcorr=61·32 em, m=2·075 mg/sec, t=3·45 see (in0,1]1.1 KCl, open circuit); and m2/3t1/6 =2·005 mg2/3sec-1/2•All the measurements were made at 25°±0·1°.Purified hydrogen was used for removing dissolvedoxygen. Since no maxima appeared in the caseof Zn(H) in the present study, addition of a sup-pressor was avoided.Zn(H) gives a single well-defined wave ineach of the supporting electrolytes of varying con-centrations. The limiting current is diffusion-con-trolled in each case as evidenced by the linearityof id vs V hplots. The plots of Ed.e. VS log i/(id-i)are linear in all the cases (representative Fig. 1).Applying different criteria10,l1 of reversibilityand irreversibility of the electrode process, it hasbeen found that the reduction of Zn(II) in potas-sium oxalate, potassium citrate and in 0'5M potas-sium tartrate is totally irreversible while in othersupporting electrolytes used in the present study,reduction is neither totally irreversibile nor does itfulfil the requirements of reversibility. The slopevalues of log plots in these cases lie in the range0,040-0,050 V (Table 1). These slope values suggest!924A234,r234to:I'0 2w3I4I234ecDIo.02volt-0.8 -0.4 0f--- LOG'0.4 0.8~/id-iJ--)Fig. 1 - Log plots of Zn(II) in (1) 1M. (2) M/2. (3) MilO and(4) MI20 concentration of supporting electrolytes (A)Ba(N03Jt. (B) Sr(N03)•• (C) Ca(l'\03l. and (D) Mg(N03).the quasi-reversible nature of the electrode processof Zn(II} at d.m.e.Gellings' methodt! as also used by Gaur andcoworkers" has been used to evaluate E~/2 (thereversible half-wave potential) and the kineticparameters ('Qt' and ks,h) of the quasi-reversibleelectrode process.The calculated values of k. hare listed in Table 1.These values are of the sa~e order of magnitudeas suggested by Delahayt" for quasi-reversibleelectrode processes. Thus, except for potassiumoxalate, potassium citrate and 0'5M potassiumtartrate, the quasi-reversible nature of the polaro-graphic reduction of Zn(II) in all the other support-ing electrolytes used in this study is ascertained.It is also evident that the value of the standardrate constant (ks.,,) for the electrode reaction ofZn(II) at d.m.e. decreases with an increase in [sup-porting electrolyte] in all the cases (Table 1). Thissuggests that with increase in [supporting electrolyte]the electrode reaction is gradually pushed towardsirreversibility. This may be due to the increasedadsorption of the ions of the supporting electrolyteson the electrode surface at higher concentrations'<.A decrease in the value of 'Qt' with an increase in[supporting electrolyte J lends further support tothis conclusion-",Taking different sets of supporting electrolytessuch that when cation was varied anion remainedidentical and vice versa, the effect of anions andcations at a certain concentration (O·1.M) on thevalue of kS•h has been attempted.NOTESTABLE 1 - VALUES OF POLAROGRAPHIC AND KINETIC PARAMETERS FOR THE REDUCTION OF Zn(II) IN DIFFERENTSUPPORTING ELECTROLYTES[SUp- Slope -Ei/2 IX k.,h X 10' [Sup- Slope -E~/2 IX ks,h X 10'porting V V em/see porting V V em/seeelec- (SCE) elec- (SCE)trolyte] trolyte]NaCIKH2PO.MI20 0·0400 1·024 0·35 6·42M/20 0·0420 0·987 0·37 6·40 MilO 0·0440 1·037 0·34 5·04M/lO 0·0500 0·988 0·36 5·56 M/5 0·0454 1·051 0·33 4·96M/2 0·0526 0·986 0·35 5·32 M/2 0·0500 1·068 0·33 4·56M 0·0542 0·995 0·34 4·45POT. ACETATENaCIO, M/20 0·0500 1-011 9·44 5·38stn» 0·0525 1·030 0·43 5·15M/20 0·0373 0·974- 0·44 6·42 MI2 0·0512 1·052 0·41 4·86M/I0 0·0384 0·973 0·43 6·08 M 0·0470 1·066 0·40 4·25M/2 0·0408 0·973 0-42 5·82M 0·0555 0·974 0·41 5·24 POT. TARTRATENaNOaMI20 0·0512 1·052 0-47 4·68MilO 0·0526 1·063 0·45 3·76M/20 0·0510 0·958 0·47 6·40 M/5 0·0555MilO 0·0501 0·970 0·46 6·16 M/2 0·0714M/2 0·0520 1·001 0·44 4·56 POT. BENZOATEM 0·0490 1·002 0·44 4·20MI40 0·0440 1·026 0·42 5·76KCI MI20 0·0454 1·034 0·41 5·42MilO 0·0500 1·032 0·41 4'95MI20 0·0500 0·987 0·44 7·94 M/5 0·0512 1·030 0·40 4·32MilO 0·0444 0·991 0·41 6·60MI2 0·0416 0·997 040 6·08 KCNSM 0·0454 1·013 0·37 4·36 M!20 0·0370 0·997 0·32 6·20KCIO.MilO 0·0408 1·012 0·31 6·04MI2 0·0440 1·035 0·30 5·08MI40 8-12M 0·0465 1·033 0·30 4·530·0400 0·975 0·38MI30 0·0434 0·978 0·37 8·07 NH1CIMI20 0·0454 0·976 0·36 7·68MflO 0·0465 0·977 0·34 7·32 M!20 0·0434 0·979 0·37 6·84MilO 0·0454 0·984 0·34 6·64KNOa MI2 0·0487 1·022 0·32 5·84M 0·0500 1·021 0·31 5·00MI20 0·0435 0·962 0·41 7·28MIlO Q'0444- 0·987 0·40 6·60 NH.NOaM/2 0·0512 0·996 0·39 6·40 M!20 0·0374 0·988 0·33 5·48M 0·0521 1·000 0·38 5·92 MilO 0·0425 0·992 0·32 4·98MI2 0·0444 0·997 0·31 4·78KNO, M 0·0476 0·996 0·30 4·59MI20 0·0424 1·030 0·36 5·68 Mg(NOaJ.MilO 0·0488 1·028 0·35 HIM/20 0·0425 1·015 0·35 6·71M/2 0·0500 1·039 0·34 4·36M 0·0555 1·056 0·35 4·36 MilO 0·0434 1·014 0·34 6·19M/2 0·0454 1·011 0·33 5·67KBr M 0·0500 0·998 0·32 4·48MI20 0'0370 1·012 0·34 7·28 Ca(N03)2MflO 0·0380 1·018 0·33 6·32 M/20 0·0454 1·014 0·43 8'88M/2 0·0400 1·028 0·32 5·00 MilO 0·0465 1·016 0·41 7·75M 0·0440 1·025 0·30 3-60 MI2 0·0476 1·021 0·40 7·35M 0·0512 1·010 0·38 6·04KISr(NOa)2MI20 0·0384 1'018 0·32 5·80 MI20 0·0372 1·018 0·33 6·51MilO 0·0400 1·014 0·32 5·48 MilO 0·0408 1·099 0·32 6·25MIS 0·0416 1·025 0·31 H6 MI2 0·0425 1·020 0·32 5·72MI2 0·0434 1·023 0·30 4·84 M 0·0440 1·020 0·31 5-44K2S04 Ba(NOa).MI20 0·0434 1·024 0·35 6·52 MI40 0·0416 1·005 0·34 6·14MilO 0·0440 1·028 0·34 5·38 MI20 0'0440 1·007 0·32 5·65MIS 0·0454 1·029 0·33 5·06 MilO 0·0454 1·016 0·31 5·24M/2 0·0459 1·035 0·32 4'85 MIS 0·0514 1·020 0·30 4·80925INDIAN J. CHEM., VOL. 15A, OCTOBER 1977The order of k. h value is as follows:(i) For univalent' cations (nitrates): K+>Na+>NH~(ii) For bivalent cations (nitrates): Ca2+>Sr2+>Mg2~>Ba2+(iii) For anions (as potassium salts): CI04> N0:i>Cl-> Be> CNS- > 1- > SO~-> CH3COO- > N02 >H2P0"4> C6HsCOO-> tart>Thanks are due to Dr S. N. Dubey, Principal,Agra College, Agra, for providing necessary facilitiesand also to CGC for the award of a contingencygrant to one of them (M.S.).References1. BOND, A. M., J. electroanal. cse«; 20 (1969), 109.2. MESARIC, S. S. & HUME, D. ~., Inorg. Chem., 2(1963), 1063.3. SANCHO,]., AREVALO,A. & GUZMAN,R., Anales. Real.Soc. Espan. Fis. Quim (Madrid), 50 (1954), 237.4. HEADRINDGE, ]. B., HAMZA,H. G., HUBBARD,D. P. &TEYLOR, :\1. S., Proceedings of the third internationalconference, Vol. I, edited by J. Hills (Macmillan,London), 1966, 625.5. MUKHERJEE, S. N. & CHAKRAVARTI,A., J. Indian chem,Soc., 38 (1961), 995.6. NATH, A. & BHATTACHARYA,A. K., J. Indian chem. Soc.,42 (1965), 586.7. GAUR, J. ~. & JAIN, D. S., Rev. Polarogr., 14 (1967), 206.8. GAUR, J. N., JAIN, D. S. & KUMAR,A., Indian J. Chem.,13 (1975), 165.9. RANDLES, J. E. & SOMERTON,K. W., Trans. FaradaySoc., 48 (1952), 95l.10. MEITES, L., Polarographic techniques (Interscience, NewYork), 1965, 217.11. GOTO, R. & TACHI, I., Proc. intern. polarograph Congr.,Vol. I, 1951, 69.12. GELLINGS, P. ]., Z. Elektrochem. n». Bunsenges. Phys.Chem., 66 (1962),477, 481, 799; 67 (1963), 167.13. DELAHAY, P., J. Am. chem. Soc., 75 (1953), 1430·14. FRUMKIN,A., Doklady Akad. Nauk. 555R., 109 (1956), 585.15. JHA, S. K. & SRIVASTAVA,S. N., Z. Naturf., 30b (1975),859.Kinetics of the Reaction between FerricChloride & ThioureaS. N. KHODASKAR& Y. B. VIBHUTEDepartment of Chemistry, Yeshwant MahavidyalayaNandedandM. H. JAGDALE*Department of Chemistry, Shivaji University, KolhapurReceived 23 November 1976; accepted 7 April 1977Kinetics of the reaction between ferric chloride andthiourea have been studied at pH 1·5. The total orderof reaction is found to be 1·5, one with respect to thio-urea and 0·5 with respect to ferric chloride. Theenergy of activation, frequency factor and entropy ofactivation have been determined. Urea is found asone of the products of the reaction. A probablereaction mechanism has been proposed.REDOX reactions of ferric chloride in acetone,with some compounds of thiourea type havebeen reported previously by Kovacovu". Howeverthe reaction has not been studied in aqueousmedium. The present paper reports the results·For correspondence.926of a systematic kinetic study of the reaction betweenferric chloride and thiourea in water.All the chemicals used were of CP grade and thesolutions were prepared in doub ly dist illed water.Ferric chloride solution was standardized by pot as-sium dichromate method. The pH of solutionswere adjusted to 1·50.The course of the reaction was followed atdifferent time intervals by measuring the absorptionat 420 nm employing an ERMA 600 photoelectriccolorimeter. This wavelength gins the maximumabsorption for ferric chloride.It is observed that at constant [thiourea J the rateincreases with increasing [FeC13J (Table 1). Theorder of reaction with respect to ferric chloride wasdetermined using Vant Hoff differential equationtaking concentration in terms of initial opticaldensity. It is found that the order of reaction withrespect to ferric chloride is 0·50. At constant [FeCI3],the rate increases with increasing [thiourea] (Table2). The plot of log x against log dxldt is linear witha slope of unity indicating that the rate followsfirst order kinetics with respect to thiourea. Hencethe total order of reaction with respect to ferricchloride and thiourea is 1·5.It is observed that the rate constant increasesas the ionic strength increases (Table 3), indicatingTABLE 1 - EFFECT OF VARYING [Fef'L] ON THEREACTION RATE{[CS(NH2)2J = O'l1IJ; temp. = 30°; pH = I-S}[FeC13] -dx/dt x 10' [FeCI3] =- dx idt x 103M rnin' ' M min ?0·05 10·0 0·0125 6·40·0375 8·0 0·005 4·350·025 7·5TABLE 2 - EFFECT OF VARYING [CS(NH')2] ON THEREACTIONRATE[FeC13 = 0·0375M; temp.= 30°; pH = 1·5J[CS(NH.hJ -dx/dtxl02 Kr x10s Kn xl0·M mirr ! mirr? mint !0·05 1·3 1·363 1·690·1 3·0 3·60 4·300·15 6·9 5·71 9·600'20 10-0 10·5 18·00·25 12-4 14·0 28·0Kr is the first order rate constant and Kn is 1·5 orderrate constant.TABLE 3 - EFFECT OF ADDED SALT ON REACTION RATE[FeCls = 0'0375M; CS(NH2)! = 0·05M; temp. = 30°;pH=1'5][NaCl]M0·000·05Krr xl0·min'?Kn X101rnin"![NaCl]M0·150'222·024'01·696·6Kn is 1·5order rate constant.

Influence of Concentration & Nature of Supporting Electrolytes on the Kinetics of Reduction of Zn(II) at Dropping Mercury Electrode

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Influence of Concentration & Nature of Supporting Electrolytes on the Kinetics of Reduction of Zn(II) at Dropping Mercury Electrode

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