摘要:
1973 789Stoicheiometry and Kinetics of the Reaction between Thallium(ii1) andAntimony(ir1) Ions in Perchloric Acid SolutionBy Prem Dutt Sharma and Yugul K. Gupta," Department of Chemistry, University of Rajasthan, Jaipur, IndiaThe rate of the reaction between TI'I'and SblI'ions in perchloric acid solution (stoicheiometry 1 : 1 ) is measurableby conventional methods at 60 "C. After ca. 70430% reaction, a precipitate is obtained consisting mainly ofSbV ions, with small amounts of TI', and adsorbed T1'''and SbIXx ions. The rate of reaction is independent of theionic strength, and reaction appears to occur via two transition states, [TIOH,HSbO2I2+ and [TI,HSbOa13+, thefirst path accounting for ca. 85% reaction. The tentative rate law is as in (i), where K1 and K, are hydrolytic constants- d [ S b"'] = [TI'"] [ S bI'I]dt ([H+] + Kl) (- 4- kiK2)for the ions TI3+ and SbO+ respectively.The quantities k3K2 and k,'K2 were found to be 2.66 and 0.30 s-l respec-tively at 60 "C, and the overall energy and entropy of activation were 16.0 f 0.4 kcal mol-1 and -1 3.1 f 1.8cal K - l mol-l respectively. In the presence of chloride ions, the reaction is complex because of the formation ofvarious chloro-complexes with the ions TI"', Sb'I', and SbV.HALPERN has previously shown that the principle ofequivalence change has limited validity because offactors, such as the entropy of activation, pertaining toredox systems. Reductions of the ion TIIn presentsomewhat interesting features. Thus whereas non-complementary reactions with the ions Fe*,3s4 Hg2+,5,6VmJ7 etc.are much faster than complementary reactionswith U1V,*,9 H3P0,,10 and H,P03,11 the complementaryreaction of the ion As" (ref. 12) is faster than the non-complementary reaction of Fell.g No kinetic evidencefor the formation of the intermediate ions T1* and AsIVwas found in the oxidation of AsIn, although the non-complementary oxidations of the ions FeII (ref. 6)and V1* (ref. 7) suggest the formation of TP. Theoxidation6 of the ion Hg,2+ provides evidence forintermediate formation of HgO. The present study ofthe reaction between the ions T P and SbIII was carriedout in order to compare the observed rate constants withthose for the corresponding reaction with As", and todetermine whether there is any kinetic evidence for theintermediate ions TlII and SbII.The oxidation oftriphenylstibine by thallium(m) chloride has beenstudied p r e v i ~ u s l y , ~ ~ , ~ ~ but no kinetic details are available.EXPERIMENTALStock solutions of thallium(II1) perchlorate were preparedand standardised 16916 as described previously.12 A solutionof antimony(II1) perchlorate, prepared l7 by dissolvingantimony trioxide (E. Merck, Guaranteed Reagent) inhot perchloric acid (60y0), standardised l8 with bromate-ionJ. Halpern, Canad. J . Chem., 1969, 37, 148.P. A. Shaffer, J . Amer. Chem. SOL, 1933, 55, 2169.C. E. Johnson, J . Amer. Chem. SOL, 1962, 74, 969.0. L. Forcheimer and R. P. Epple, J. Amer. Chem. SOC.,A. M. Armstrong and J. Halpern, Canad.J . Chem., 1967,35,A. M. Armstrong, J. Halpern, and W. C . E. Higginson,N. A. Daugherty, J . Amer. Chem. SOL, 1966, 87, 6026. * A. C. Harkness and J. Halpern, J . Amer. Chem. Soc., 1969,J. 0. Wear, J . Chem. SOL, 1966, 6696.lo K. S. Gupta and Y. K. Gupta, J . Ckem. Soc. ( A ) , 1970, 266,l1 K. S. Gupta and Y . K. Gupta, J . Chem. SOC. ( A ) , 1971,1962, 74, 6772.1020.J . Phys. Chem., 1966, 60, 1961.81, 3626.1180.solution using Methyl Orange as indicator, and with anexcess of Cem, back titrating the latter with Fe* to aFerroin end point.1° Fresh solutions of antimony(xx1)perchlorate were prepared daily, since such solutionsdeteriorate over a period of time.A solution of cerium(xv) perchlorate was prepared bydissolving cerium(1v) ammonium nitrate (B.D.H., AnalaR)in perchloric acid (60%) and standardised l9 against iron@)ammonium sulphate to the Ferroin end point.Lithiumperchlorate solution was prepared by neutralising per-chloric acid with lithium carbonate to pH 6.5-7. Allsolutions were prepared in doubly distilled water, thesecond distillation being from potassium permanganate.Determination of the Stoicheiowetry .-Reactants of suitableconcentrations were mixed in a measuring flask (250 ml)and allowed to react on a water-bath a t 60 "C for ca. 2 h. Aprecipitate was obtained, the colour of which varied fromgreen (excess of Tlnl) to yellowish-white (excess of SbIII).The remaining liquid in the flask was made up (to 250 ml)with distilled water and the precipitate allowed to settlecompletely.The precipitate was then filtered off andwashed three or four times with water. The filtrate andthe washings were collected in a measuring flask and madeup to 250 ml with water [solution (A)]. The precipitatewas dissolved in ~ M - H C ~ on heating and the resultingsolution made up to 250 ml with water so as to yield asolution 2.5111 in HC1 [(B)]. The concentrations of thevarious species in (A) and (B) were determined separatelyas follows.(a) The ion SblI1 was determined using cerium(1v)ions as described under Kinetic Procedure. (b) The ionsSblI1 and Tll together were determined bromometricallyin 2.!5-3-5~-HCl using Methyl Orange as indicator.(c) The ions Tlnl and SbV were reduced by sulphite ions 2ol a P.D. Sharma and Y . K. Gupta, J.C.S. Dalton, 1972, 62.l3 A. E. Goddard and D. Goddard, J . Chem. SOL, 1922, 121,l4 A. E. Goddard, J . Chem. SOL, 1923, 125, 1161.l6 I. M. Kolthoff, R. A. Belcher, U. A. Stenger, and G. Mat-suyama, ' Volumetric Analysis,' Interscience, New York, 1967,vol. 3, p. 370.H. G. S. Sengar and Y. K. Gupta, J . Indian Chem. SOC.,1966, 43, 223.l7 Fr. Fichttr and E. Jenny, Helv. Chim. Ada, 1923, 6, 226;J . Chem. SOL, 1923, 124, 246.1* Ref. 16, p. 44.l9 H. Rathsburg, Ber., 1928, 61, 1663.2o W. M. Macnabb and E. C. Wagner, Ind. and Eng. Chem.(Analyt. Edn.), 1930, 2, 261; J. Knop, Analyt. Chem., 1923,63, 8.(A).483790 J.C.S. Daltonand the excess of the latter then removed as sulphur di-oxide by boiling the solution.The total amount of SblI1(that present initially plus that obtained by reduction ofSbv ions) was determined using cerium(1v) ions as in (a).( d ) The total amount of T1I and SblI1 in the reducedsolution (c) was determined bromometrically.18(B). (e) The total amount of the ions Sbnl and T1I wasdetermined bromometrically l8 without interference fromthe ions TIU1 and Sbv. (f) The ion SbV was reduced bysulphite ions and the excess of the latter removed byboiling the solution. The total amount of the ion SblI1thus obtained was determined iodimetrically 21 in a hydro-gencarbonate-ion medium. (g) The ion Tlm was alsoreduced in (f) and hence the total amount of the ions SbIIIand TlI, which is equivalent to the initial total amount ofthe ions Sbl*I, Sbv, TIU1, and Tll, was determinedbromometrically.18 (h) The ion TllI1 was precipitated asthallium(r1I) oxide by sodium hydroxide, centrifuged,washed, and determined iodometrically.16916 It was notpossible to determine the ion SblI1 in solution (B) usingcerium(1v) ions, because the reaction between CeIV and T1Iions is catalysed z2 by chloride ions.Kinetic Procedure .-Most kinetic experiments werecarried out at 60 "C in stoppered flasks of Corning glass.The reaction was followed by determining the Sbm ion intrations of the reactants that no precipitate appeared, orappeared only when more than 70% reaction had occurred.The order of mixing of the reactants had no effect on therate of the reaction.Rate constants for identical runswere reproducible within f1-4yo depending on the errorin the analyses of the aliquot portions.RESULTSStoicheiometry.-Although stoicheiometric experimentswere complicated by formation of the precipitate andadsorption of TIIU and T1I ions, all the four species TlI,TlIUa Sbnl, and Sbv could be determined and the results(with a maximum deviation of &loyo) show that 1 mole ofSblI1 reacts with 1 mole of TllI1 to form 1 mole of Sbv and1 mole of TlI.The concentration of the different species inthe filtrate and precipitate were calculated as followsFiltrate (A) Precipitate (B)[SbIII] = a [Tl'II] = h[TlI] = b - a [TlI] = g - f - h[Sbv) = c - a [SbIn] = e - g + f + h[TlIII] = d - c - b + a [Sbv] = g - e - h(results are shown in Table 1).In view of the complicatedway in which the determinations have been made, thestoicheiometric results are more than satisfactory. It isTABLE 1Concentrations * of the species TlI, TlIII, SblI1, and Sbv obtained in the filtrate (A) and precipitate (B) on reaction ofthe ions Tlnl and Sbnl2.003-004.005.001.002.003.004-002.003.004.005-001.002.003.004.002.003.004.005.002.003.004-005.002.602.666-256.0037-526.621.216.047.561.772.575.080.08'7-691.791-5090.088.393.897.010.010.08.338-7512.515.012.511.211.610.010.07.5042.627.517.512.05.04.23.123.17.508.333.754.006-268-338.5012.507.506-004.5010.085.087.590.092.636.050.061.265.090.087.091.289.0* Expressed as a percentage of the concentrations of the reactants TlIII and SbnI.the reaction mixture a t various times in the following way,which is based on the reaction 22 between SblI1 and CeIV ionsin perchloric acid.Aliquot portions (5 or 10 ml) of thereaction mixture were added to a known excess of cerium(1v)perchlorate solution, diluted to yield [H+] = 0.15-0.2~,and set aside for 8-10 min. The excess of Cem was backtitrated against a standard solution of iron(I1) ammoniumsulphate to the Ferroin end point; H2S04 was added(0.5-1~) to improve the end point. The ions FeII andTllI1 do not interfere with this method.23 The method isnot suitable in the presence of chloride-ion concentrations> 0 .1 ~ for the very slow reaction between Cem and T1I ionsis catalysed 24 by chloride ions, and a slow reaction 26between Cem and chloride ions occurs yielding a complexwhich decomposes to give chlorine.All kinetic runs were carried out under such concen-a1 A. Blanchetiere, Bull. SOC. chim. France, 1920, 27, 477.22 S. K. hlishra and Y. K. Gupta, J . Chem. SOC. ( A ) , 1970,23 K. G. Ashurst and W. C. E. Higginson, J . Chenz. SOC.,260.1953, 3044.obvious that the precipitate consists mainly of the ion Sbvwith a small quantity of TlI. It also has adsorbed TllI1 orSblI1 ions, whichever was in excess in the original mixture.It may be of interest to report that simple mixing of theions SbIII and Sbv in perchloric acid solutions results in awhite turbidity as is also found 26 in the exchange reaction.Kinetics.-The reaction was found to be first order in[SbIII] and [Tlnl] from the log-log plots of rate againstconcentration.Second-order rate constants (Table 2) werecalculated from conventional log[Tlm]/[Sbm] against timeplots (Figure 1). In a few cases, second-order rate con-stants were calculated from pseudo-first-order rate con-stants obtained from linear plots of log [SbIn] against timewith an excess of T1" ions.The hydrogen-ion concentration was varied with per-chloric acid at a constant ionic strength of 3~ (LiClO,).24 F. H. Duke and C. E. Borchers, J . Amer. Chem. SOC., 1953,75, 5186.25 K. K. Sengupta, P. K. Kaul, and S. P. Moulik, J . IndianChem. SOC., 1963, 40, 429.26 C.H. Brubaker and J. A. Sincius, J . Phys. Chem., 1961,65, 8671973 791The rate of reaction decreased with increasing hydrogen-ionconcentration. Results a t 50, 55, and 60 "C and I = 3 . 0 ~TABLE 2Second-order rate constants for the reaction between TllI1and SbIII ions a t 60 O C , [H+] = 1 . 5 ~ ~ and I = 1 - 5 ~1 03[T1111] /M 103[SW11]/~ k / l mol-1 s-11.00 1.00 0-771-50 1.50 0.802.00 1.35 0.762.00 1.50 0.772-00 2.50 0.772.00 3.00 0.772.00 4-50 0.8 12-00 5-00 0-8 12.50 1.35 0-763-00 1.35 0.763.00 2.00 0.773.00 2.50 0.773.50 2.60 0.833.50 3.50 0.804.00 1.35 0-764-00 4.00 0.804.50 2.50 0.7 75-00 0.50 0-77 *5.00 2.00 0.805.00 2.50 0-786.50 0.50 0.76 *6.00 0-50 0.77 *6.50 0.50 0.81 *7.00 2.00 0.77Average 0.78 f 0.02* Obtained from pseudo-first-order rate constants in runsin which TPII was in excess.o*r P /I 1 4 8 12Time/ minO00FIGURE 1 Second-order plots at 60 "C for the T1I1I-SbIt1 reac-tion: (A), [TllI1] = O-OO~M, [SblI1] = 0.002M; (v), [TlIII] =0-0045~, [SbI'I] = 0.002M; (e), [Tl? = 0*0045~, [SbIII]= 0.0026~; (a), [TII'I] = O.O036M, [Sbur] = 0.0026~; (O),[TlIIl] = 0.002M, [SbI*] = 0.0025~are shown in Table 3.A log-log plot of rate againsthydrogen-ion concentration yielded a negative order of2' D. H. Irwin, J . Chem. SOL, 1957, 184.2E G. Jander and H. J. Hartmann, 2. anorg. Chem., 1966, 339,29 S. K. Mishra and Y . K. Gupta, Indian J. Chem., 1968, 8,30 G. Jander and H. J. Hartmann, 2. anorg. Chem., 1965, 339,239.757.256.ca.1.3 with respect to [H+]. Although the ionic strengthwas changed from 0.8 to 3 . 0 ~ (LiClO,), no change in therate of reaction was observed. The products of reaction,TlI ( < 0 . 0 0 5 ~ ) and SbV (<0.002~), were also found to haveno effect on the rate of reaction.The energy and entropy of activation of reaction werecalculated to be 16.0 f 0.4 kcal mol-l and - 13.1 f 1.8cal I<-1 mol-1 respectively.TABLE 3Second-order rate constants for the reaction between TlIIIand SbIII ions a t different hydrogen-ion concentrationsand temperatures. I = 3 . 0 ~ ~ [TlIII] = [SblI1] = 2 x1 0 - 3 ~Temp./"C [HC104]/~60 1.251.501.752.002-252.502.75R/1 mot* s-10.840.740.500.420.3550.3030,26055 1.00 0-671.25 0.501.50 0.401.75 0.3162.00 0.2602.25 0.21050 1.00 0.5001.25 0.3661.50 0.2902-00 0.1912-50 0.1332.75 0.117DISCUSSIONThe hydrolysed species T10H2+ has been consideredto be reactive in most kinetic studies 9 9 2 7 involving theTlIII ion.Daugherty 7 assumed that the second hydro-lysed species Tl(OH),+ was the reactive one in theoxidation of VrI1 ion but there is no evidence for thisspecies in the present investigation. The ion SbIII pre-dominantly exists as SbO+ in perchloric acid solutions asindicated 28 by high-voltage ionophoresis, ion-exchangestudies, and diffusion measurements. The oxide Sb203or its hydrated form HSbO, may also exist 29930 in suchsolutions, but not to the extent of more than 5% inEquilibria (1) and (2) generally determine thehydrogen-ion dependence of reactions involving the ionsT P and SbIII.Although various values of K , haveK1K ,been reported,3l~~~ the most accepted value based one.m.f. measurements 33 is 0.073 mol 1-1 at 25 "C and I =3 . 0 ~ . Since the enthalpyN935 of hydrolysis is 11.0 kcal31 G. Mattock, Acta Chem. Scand., 1954, 8, 777.S2 T. E. Rogers and G. M. Waind, Trans. Faraday SOC., 1961,33 G. Biedermann, Arkiv Kemi, 1964, 6(5), 627.34 K. A. Kraus and F. Nelson, J . Amer. Chem. SOC., 1956, 77,35 R. H. Betts, Canad. J . Chem., 1955, 33, 1775.1 *5M-HCIO,.TP+ + H20 T10H2+ + H+ (1)SbO+ + H,O + HSbO, + H+ (2)57, 1360.3721792 J.C.S. DaltonFurthermore if the reaction path TlOH2+ + SbO+ isassumed instead of that of TP+ + HSbO, for the minorcontribution to the rate law, k,'K, will be replaced byK,"K, in (3), where k," is the rate constant for thecorresponding path.However, values of K3/' calculatedfrom the intercept in this way were found to varyirregularly at the three temperatures and hence thernol-l, Kl has the values 0.516,0.39, and 0.295 mol 1-1 at60, 55, and 50 "C respectively. The hydrolysis con-stant s for the ion SbO+ is ca. 0.02 mol 1-1 at 25 "C,and hence Sbm largely exists in this form in the acidsolutions of the present investigation.With T13+ or T10H2+ and SbO+ or HSbO, as thereactive species, the redox reaction could take place byone or more of four paths. The hydrogen-ion depend-ence, which could throw light on this aspect, is somewhatcomplicated and inconclusive. However it does ruleout the possibility of a reaction between the ions T13+and SbO+.The overall entropy change for the reactionis about the same as the normal value of -12 cal K-lmol-l for a bimolecular process, and hence highlycharged transition states formed from the ions T13+ andSbO+ and T10H2+ and SbO+ are less probable. Tworeaction paths, TlOH2+ + HSbO, and TP+ + HSbO,,remain with rate constants k3 and k,' respectively,giving the rate law (3). The observed rate constant k isthus given by (4).(4)Plots of k(K, + [H+]) against 1/[H+] gave straightlines with intercepts (Figure 2) in accordance withequation (4). From the gradients, values of k3K2 werefound to be 2.66, 1.70, and 1.33 s-l at 60, 55, and 50 "Crespectively.Similarly from the intercepts, corre-sponding values of k3'KZ were found to be 0.35, 0.30,and 0-25 s-1 respectively. The heat of hydrolysis forreaction (2) is not known and hence k3 and k3' cannot becalculated, but the enthalpy appears to have a negativevalue because the difference in k3K2 or k,'K, values atdifferent temperatures is small. If free energies offormation36 for the species SbO+, HSbO,, and H,O areany guide, this enthalpy change would be small andnegative. The quantity k3/k3' indicates that the con-tribution from the first path (T10H2+ + HSbO,) isca. 85%. If, therefore, the overall energy effects haveany significance, they refer to this major step.0 0.2 0.4 0.6 0.8 1.01 H+l-'/ M-'Plots of K(K, + [H+]) against l/[H+] for the TlIILSbIII reaction: [ T P ] = [SbIII] = 0 . 0 0 2 ~ ; I = 3M; (0)'FIGURE 260; (0), 5 5 ; (A), 60 "Creaction path between the ions T10H2+ and SbO+ isagain excluded.A comparison of the rate constant (1.48 s-l at 35 "C)for the corresponding oxidation of As* by Tlm (ref. 12)shows that the rate of oxidation of Sb* ion is slower inspite of its lower energy of activation (E, for Asm =21.5 kcal mol-l). The much larger entropy change(3-10-0 cal K-l mol-l) for the former reaction more thancompensates the high energy of activation.[2/580 Received, 13th March, 1972186 W. M. Latimer, ' The Oxidation States of the Elements andtheir Potentials in Aqueous Solutions,' Prentice Hall, New York,1938, p. 109
ISSN:1477-9226
DOI:10.1039/DT9730000789
出版商:RSC
年代:1973
数据来源: RSC