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Kinetics of solvolysis of thetrans-dichlorotetrapyridinecobalt(III) ion in water and in water + methanol |
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Dalton Transactions,
Volume 1,
Issue 12,
1980,
Page 2405-2409
Christine N. Elgy,
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摘要:
1980 2405Kinetics of Solvolysis of the trans-Dichlorotetrapyridinecobalt( 111) Ion inWater and in Water + MethanolBy Christine N. Elgy and Cecil F. Wells," Department of Chemistry University of Birmingham Edgbaston,P.O. Box 363 Birmingham B15 2TTRates for the first-order solvolysis of the trans-dichlorotetrapyridinecobalt(lll) ion have been measured in waterand in mixtures of water and methanol for a range of temperatures. A non-linear dependence of log (rate constant)with reciprocal of dielectric constant suggests that the effects of solvent structure are important and this is con-firmed by the comparison of the variation with solvent composition of the transition-state parameters with thevariation with solvent composition of a range of physical properties of the solvent mixture.The application of afree-energy cycle relating the process initial state+transition state in water to that in water + methanol shows thatchanges in solvation of the transition state have a dominant effect on the rate.IN a survey of the kinetics of SN1-type solvolysis re-actions of octahedral complexes in mixtures of water +co-solvent the variation of rate with respect to di-electric properties and solvent structure was investi-gated.l By applying a free-energy cycle to the longstretching of the metal-ligand bond it is possible todetermine whether changes in solvation with solventcomposition in the initial or in the transition state havea dominant effect on the rate of reaction. A highdegree of unanimity was found for a large number ofcomplexes for a range of co-solvents changes in solv-ation in the transition state have a dominant effect onthe rate.I t is therefore of interest to explore the vari-ation of rate with solvent composition for other com-plexes. Although the trans-dichlorotetrapyridine-cobalt(II1) ion was one of the complexes used to showthat basic hydrolysis requires a mechanism different fromhydrolysis under acidic or neutral conditions,2 littlefurther kinetic work has been done with it. It has beenshown that at high concentrations of OH- and for theaddition of other bases like azide or acetate ions theloss of pyridine from the complex OCCU~S,~ contrary tothe results in acidic solutions and in a borate bufferof pH 9.18 where only a chloride ion is lost.The effectof changes in dielectric constant on the rate of solvolysishas also been investigated by adding sucrose to thereaction mix t ~ i r e . ~EXPERIMENTALMaterials .-General purpose reagent CoCl,%H,O andlaboratory grade pyridine nitric acid and hydrochloricacid were iised in the preparation of trans-dichlorotetra-pyridinecobalt( 111) nitrate. The method used was based onthat proposed by Werner and Feenstra involving oxidationby chlorine gas. However our observations during thepreparation differed froni those of Werner and Feenstraa t one point. A considerable number of experiments wasperformed in order to achieve the optimum conditions forthis preparation. When the chlorine was passed throughtlie mixture used previously (20 g of CoC1 in 30 g of waterand 30 g of pyridine) a pink solid was formed and not a bluesolution as previously reported with a low yield of productdue to poor contact between tlie solid and the chlorine gas.Addition of a large excess of pyridine (py) to dissolve thispink solid caused the subsequent absorption of greaterthan molar proportions of chlorine and no [Co(py),Cl,]CIwas precipitated.The addition of an excess of water todissolve the pink precipitate resulted in only a small yieldof the product. ,4 combination of a slightly higher temper-ature with small volumes of water was found to give optimumyields of the product. The following final conditions wereemployed pyridine (30 g) was added to a solution con-taining cobalt(I1) chloride (20 g) dissolved in water (40 g),and the precipitate formed redissolved at 35 "C.Afterpassing chlorine gas through the solution i t was cooled andallowed to stand for a few hours. The precipitate was thenfiltered off and washed successively with small volumes ofice-w ater ethanol and dic t h yl ether. The tva ns-dic hloro-tetrapyridinecobalt(r~r) chloride was purified by re-pre-cipitation from an aqueous solution by adding HCI thisprocess was repeated. To prepare the nitrate the chloridewas dissolved in water a t 45 "C and nitric acid added. Thenitrate was purified by re-precipitation with added HNO,,followed by repeated washings as described above for thechloride. The nitrate was then dried in air producing veryfinc green crystals.Large green needles were obtainedafter recrystallization from water containing a little aceticacid. The dry crystalline nitrate showed no sign ofdecomposition or loss of pyridine when stored for severalmonths. For the kinetic work the very fine crystals wereused to obtain rapid dissolution.For the kinetic investigation A.K. nitric acid was used.Laboratory grade NaCl and K[OH] were used. A buffersolution of pH 8 was prepared by mixing 25 cm3 of 0.2 moldm- Na[H,E'O,] (A.K.) with 46.8 cm3 of 0.2 mol dm-,K[OH] and the resulting solution diluted to 100 cm3 withdistilled water. A.R. methanol was used and water wasdistilled once in an all-glass still.Procedure.-The solvolysis kinetics were followed in thetherinostattcd cell compartments of various spectrophoto-meters a Unicaiii SP800 with scale expansion unit and amillivolt recorder ; a Unicain S1'500 series 2 syectrophoto-meter equipped with SP505 programmed controller toallow tlie recording of optical densities at pre-set timeintervals on a flat bed recorder; and a manually operatedUnicam SP500 spectrophotometer. 5 mm and 10 mmstoppered cells were used anti the results obtained fromthese several spectrophotometers agreed well amongst them-selves.RESULTS AND DISCUSSIONBoth the chloride and nitrate showed a peak absorptionat 295 nm which disappeared completely when the sol-volysis was finished.Addition of methanol to aqueouJ.C.S. Daltonsolutions produced no spectral shifts in this peak. Allthe kinetic results were obtained using the nitrate a t295 nm with an initial concentration ca.8 x moldm-3. The decay of optical density was measured inthe following conditions a t the same tempcrature:5 x mol dm-3 NaCl; 5 x lW3 and 5 xmol dm-3 HNO,; 1 x niol dm-3 K[OH]; pH 8phosphate buffer and pH 8 borate buffer. The plots oflog (optical density) against time for pure water and theNaCl solution gave straight lines with the same slope.For the solutions containing HNO the plot of log(optical density) against time for 5 x mol dm-3 wascurved whereas that for 5 x mol dm-3 was linearwith a slope equal to that for pure water and the NaClsolution. The plots of log (optical density) againsttime for the KCOH] solution and the buffer solutionswere curved.The detailed solvolysis kinetics weretherefore followed in pure water and in the mixtureswater + methanol without any adjustment of the pHor the ionic strength this absence of an effect of acidityconfirms the earlier observation.2Variation of Rate with Temnjxmztttre in Pzbre Water.-The rates were followed for very long periods a t 30,35 40 42.5 45 and 49 "C. Good linear plots for log(optical density) against time were always obtained.First-order rate constants k derived from the slopes ofTARLL 1First-ordcr rate constants for tlic solvolysisof tvarts-[Co(py),Cl,]+ in pure waterO,/"C30. ou30.0030.1535.0035.0040.004 0.0039.9542.5042.5042.4545.05105k 1s-10.921.181 .A72.032.074.004.483.926.75.75.89.50J"C45.0545.0045.0045.0048.6548.8548.6548.7048.6548.6548.701Osk/s^J0.39.58.09.011.212.311.515.713.812.211.8these plots are given in Table 1.A plot uf log Fz againstreciprocal of absolute temperature was linear ; valuesfor the enthalpy AH and entropy AS of activationobtained by the application of the least-squares pro-cedure are given in Table 3.Variation of Rate with Temperatwe in Water +hIethaizoZ.-Rates were measured in 5 10 20 30 40,50 66.7 80 and 93.774 v/v methanol for temperatures35 40 45 and 50 "C. Values for the first-order rateconstants obtained from the slopes of the plots of log(optical dcmsity) against time are collected in Table 2.In addition values of k obtained in pure methanol areTABLE 2First-order rate constants ( lo5 1zls-l) for the solvolysis of tmns-[Co(py),Cl,] ' in water + mctliaiiolMole fraction of methanole,/oc34.9535.0035.0035.0035.0035.0035.0535.0535.1035.1535.5539.8040.0040.0040.0040.0540.0540.1044.8044.8544.9044.9545.0046.0045.0545.0545.1048.5048.6048.6048.6548.7048.7548.8548.8548.9049.049.10.023 0.0473.303.303.223.573.693.773.123.153.387.37.47.77.87.ti12.018.5 lti.816.522.7 25.322.727.70.1003.773.723.728.I8.38.58.78.519.719.233.033.531.50.1604.304.274.184.459.59.521.651!1.231.821.235.335.00.2295.04. ti04.8310.311.011.0024.225.326.240.036.841.50.308 0.4715.25.15.96.45.85.66.8l't.313.313.02ti.227.728.846.5ti.66.915.215.529.228.041.044.20.ti41 0.870 1 .0007.6 8.8 10.57.7 8.88.38.17.08.00.38.36.8 20.721.518.832.740.337.233.055531u.26471980 2407TABLE 3Values of tlie enthalpy and entropy of activation a t 25 "Cin water + for the solvolysis of Irans-[Co(py),Cl,]methanolMole fractionof methanol AHt/k J mo1-I A.S/ J I<-* mol0 101 f 11 - 3 % 210.047 124 f 4 74 f 120.I60 165 & 2 81 f 40.308 126 f 8 83 f 250.870 115f 7 54 f 190.02 3 121 7 64 f 240.100 128 f 5 88f 150.229 128 & G 88 f 190.47 1 116 rJr 8 51 & 2G0.641 115IJt 7 61 f 29also included in Table 2.Linear plots of log 12 against1/T K were obtained for all solvent compositions andvalues for AH; and AS obtained by tlw application oftlie least-squares procedure to these plots arc given inTable 3.The Efcct of Solvent Combositiota on the Solvolysk.-Tables Z and 2 show that the rate constant tends toincrease smoothly with increasing methanol content ofthe mixture. Extrapolation of the data in Table 3 to25 "C for pure water produces k = 5.8 x lW5 s-l whichcompares reasonably well with the value reported forthis temperature k = 8.2 x s-l. Figure 1 showsthat a plot of log k against the Grunwald-Winstein Yvalues is linear which shows that the reaction is of anSN1 type with considerable extension of tlie Co-C1 bondin the transition state Y-values were interpolated fromthe collection of Wells.' However m = -0.13 obtainedfrom the slope of Figure 1 is considerably different fromm = 0 .1 8 4 . 3 5 obtained8 from similar complexes withammonia or 1,2-diaminoethane as complexing ligandsinstead of pyridine in water + methanol.The relationship between log k and dielcctric constantD for the solvolysis of a complex C". involving the exten-sion of a MXM * Xzhf- is expressed 1 by equation (I),2.303RTlog (k;) - - Y(&- - A)AGP(M)ri + AGF(X)n - AGP(C)ri (1)where subscripts w and s indicate pure water and mixedsolvent respectively 2 represents charge Y = radius,e = electronic charge G is related to dipole moment,N = Avogadro's number and AGt*(i)a includes allchanges in free energy resulting from structural changesin the solvents when species i is transferred from waterinto the mixture.A linear plot for log k against D,-Ican only be obtained if equation (2) holds. However,AGbe(C)n - AGt,*(lc.l)n + AGt,"(X)u ( 2 )Figure 2 shows that a plot of log k against D,-l iscurved with a positive slope values for D wereinterpolated from the data of Akerlof and of Martin aridBrown.1° The curvature is accentuated wlieii pointsat mole fractions of mctlianol ',0.3 1 are iltclu(lec1.0 71 I I I I-1 0 1 2 3 4YFIGURE 1 I'lot o f log k at 25 "C in water + methanolagainst Y valuesFigurc 2 also shows a. plot o f log k and D,? for thesolvolysis in 1120 -1- sucrose values for D were inter-polated from tlie data of Akcrlof s and Malmberg andMaryott.ll Tliis is also a curve contrary to the figureshown in the original ref~rence,~ where rate data forpure water arc not given; but it has a negative slope andthat in water + inethanol has a positive slope.For asolvolysis like tliis where 2~ = +1 Z ~ I = 4-2 ZX =-1 and Y~~ - YC plots of log kS against D;l shouldhave a riegativc slope if the term in G in equation (1)can bc ignored and if equation (2) holds. The non-linearity of the plots in Figure 2 combined with thedivcrsity of the slope suggests that equation (2) doesnot hold since it is not expected that tlie term in G willvary greatly between mixtures with different co-solvents.I t therefore seeins likely that structural changes inducedin the mixture by the addition of the co-solvent areimp or t an t .The effect of solvent structure on the rate is shownmore clearly in Figure 3 where plots of AH$ and AS:show extrema a t a mole fraction of methanol ca.0.2 -0.3 the errors which are not abnormal are given inTable 3 and Figure 3 and show that the extrema arereal. Properties of water + methanol mixtures suchas the excess thermodynamic functions of mixing,12J3decrease in the partial molar volume of methanol,l4*15increase in the structural contribution to tlie change inthe temperature of maximum density,16 and viscosity l7show that the addition of methanol to water firs2408 J.C.S. Daltonincreases the structure in solution. The comparisonlof these properties with those of other water + alcoholmixtures shows that compared with other alcohols,methanol has only a small structural effect.Thus thedeviation of the minimum in the excess enthalpy ofmixing from mole fraction = 0.5 is small for methanolcompared with the other alcohols particularly thosewith a branched carbon chain; l 2 3 l 3 9 l 8 the decrease in thepartial molar volume of methanol P2 - V2* is onlysmall compared with the decrease found for otheralcohols; l49I5 methanol produces no real elevation ofthe absorption of ultrasonic waves unlike the highabsorption found with branched chain alcohols.19 How-ever the maximum of the effect on P2 - V2* 14,15 andviscosity l7 in water + methanol occurs in tlie sameregion of composition as that found for the extrema inAH3 and AS3 in Figure 3.Application of the free-energy cycle,l referred toearlier concerning the process initial state+transitionstate in water and in water + alcohol to this presentsystem produces the scheme shown where subscript wAGWt[Co(py),C1,]+ w- [Co(py),C112+w -t Cl-,ACtG(CI-) I 1 AGte{[Co(py),CilP+1 I - AGt*{[CdPY),-Cl*l+1[Co(py),C1,1 +S % [C0(py),C1l2+s + c1 -sand s refer to water and water + methanol respectively,and AGt*(i) is the free energy of transfer of species i from140pj 2 O 1c' 3 E' 3XS aI I I I I0.2 0.4 06 0.8 1.0-10'Mole fraction of methanolFIGURE 3 Plots of AH$ (kJ mol-l) (0) antl A S (J Ti-' niol-')(0) for 25 "C against solvent composition for water +methanol5 -12-141L a 1 I0.1 0.2 0.3 0.4 0.5Mole fraction of methanolFIGURE 4 Plot of the left-hand side of equation (3) for 25 "Cagainst solvent composition in water + inethanol the value ofAGte (C1-) for mole fraction = 0.471 is extrapolated from theexpcriniental data in ref.20water into water + methanol. Equation (3) can bededuced from this. Values for AGLe(C1-) are avail-k Wk S2.303RTlog- - AGto(C1-) =G*( [Co(py),C112+l - AGt"(iCo(py),C1,1+> (3)able 20,21 for mole fractions of methanol up to ca. 0.35at 25 "C and values for k and ks at 25 "C can be cal-culated from the transition-state parameters in Table 3.In general for water + co-solvent mixtures AG,*(i) fori = a cation are negative,20-22 except for some uni-positive alkali metals with small positive values inwater + methanol,20 and AGte(i) for i = bipositivecation M2+ are always negative with AGt*(M2+) morenegative than AGt*(M+). Therefore if the left-handside of equation (3) is negative changes in solvation onthe bipositive transition state have a dominant effecton the rate; alternatively if the left-hand side ofequation (3) is positive changes of solvation on theunipositive initial state have a dominant effect.Figure4 shows that a plot of the left-hand side of equation (3)for trans-[Co(py),Cl,]+ at 25 "C in water + methanolis negative; and therefore changes in solvation on thetransition state have a dominant effect on the rate.This is comparable with the conclusions reached fromthe application of a cycle of the above type to the solv-ation of a range of transition-metal complexes in water +co-solvent with a variety of co-solvents.[0/294 Received 20th February 19801REFERENCESC .F. Wells J.C.S. Faraday I 1977 73 1851.K. G. Pearson 13. E. Meaker and F. Basolo J . Inorg.R. G. Pearson P. M. Henry and F. Basolo J . Amev. Chem.V. D. Panaqyuk antl A. V. Arkharov Russ. J . Chem. 1965,A . Wcrncr ant1 R. I;ecnstra Rer. 1906 39 1538.E. Grunwalcl ant1 S. Winstcin J . Aimr. Chem. SOC. 1948 70,1'. l<. Wclls Chem. Hen. 1903 63 171.C . H. Langford Im)i,g. Chem. 1964 3 228; J . Burgess J .Chem. Soc. ( A ) 1970 2703; J . Burgess and M. G Price ibzd.,1971 3108.Nuclear Chem. 1955 1 341.Soc. 1957 79 5382.10 852.8461980G. Akerlof J . Amer. Chem. SOC. 1932 54 4125.lo A. R. Martin and A. C. Brown Trans. Faraday SOC. 1938,34,l1 C. G. Malnibergantl -4. A. Maryott J. Res. Nut. Bur. Stand.,l2 A. G. Mitchell and W. F. K. Wynne-Jones Discuss. Faradayl3 R. F. Lama and B. C.-Y. Lu J . Chem. and Eng. Data 1965,lo K. Nakanishi Bull. Chem. SOC. Japan 1960 33 793.l5 V. S . Griffiths J. Chem. SOC. 1954 860.742.1950 45 299.Soc. 1953 15 161.10 216.Is G. Wada and S. Umeda Bull. Chem. SOC. Japan. 1962 35,646.l7 F. Winkler and H.-H. Emons 2. CJaem. 1963 9 73.1* J . Kanttamaa E. Tommila and 3%. Martti Ann. Acad.M. J . Blandainer ' Introduction to Chemical liltrasonics 'Scient. Fennicae 1959 no. 93.Academic Press London 1973 ch. 11.2o C. F. Wells. J.C.S. Faraday I 1973 984.21 C. F. Wells Adv. Chem. Ser. 1979 177 54.22 C. F. Wells J . C . S . Favaday I 1974 694; 1!)75 1868; 1976,601 ; 1978. 636 1569; and unpublished work
ISSN:1477-9226
DOI:10.1039/DT9800002405
出版商:RSC
年代:1980
数据来源: RSC
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