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Isotopic chlorine exchange between trimethylchlorosilane and hydrogen chloride
Notes
385
precipitate of the [Pt(en)d ++ -salt formed rapidly and provided a clean separation. With an initial, careful purification of the two salts by fractional crystallization, any apparent initial or separationinduced exchange was less than 3--4 per cent. However, a slow precipitation of Pt xv and the finite solubility of [Pt(en)t][B4418 limited the concentration range which could be employed. Exchange half-times were taken from the usual plots of log (1 -- F) vs. t (where F is the fraction of exchange). In general, the plots were satisfactorily linear over a period that the exchange was followed, usually about two half-times. The experimental conditions and the results are presented in Table 1. If the rate of the exchange is given by expression (1), the rate constant, k, is related to the half-time of an exchange experiment, ~'1/~ by the expression k = In 2/([Pt(en)2 +÷] ÷ [Pt(en)~Cl2++])[Cl-]~-l/~.
(2)
The indicated values of k are included in Table 1. The results, which include a considerable range of concentration variables, confirm the form of rate expression predicted for the exchange by BASOLOet al. ~ The average value for k, 12.1 sec-x M -~, at 25 ° is 30 per cent lower than their value; however, with the statistics indicated in the table this difference may not be real. The limited number experiments at 15°C indicate values for the enthalpy and entropy of activation, AH~ = 10.1 kcal, and AS ~: = --20 e.u. Institute for Atomic Research and Department o f Chemistry Iowa State University Ames, Iowa
L. T. Cox S.B. COLUNS D . S . MARTIN,JR.
Isotopic chlorine exchange between trimethylchlorosilane and hydrogen chloride (Received 20 October 1960; in revised form 21 November 1960)
THE recently reportedwho reported a decrease in the exchange rate between CHsC1 and HCI when the hydrogens were successively replaced by more electronegative fluorine atoms. Because the exchange studies on the gas phase system showed an appreciable (but irreproducible) "zero-time" exchange, it was decided to investigate the isotopic halogen exchange in the system HCI (gas)-(CHs)sSiC1 (liquid) since the separation method in the study of the gas phase system involved a transition through the inhomogeneons state. Since the Henry's Law constant for HCI in (CH,)sSiCI has not been measured, it is not possible to evaluate the kinetic parameters in terms of concentrations in the liquid phase. In order to establish the dependence of the exchange--which proved to be appreciably rapid, even at 0°C----on the hydrogen chloride concentration, the experiments were carried out in reaction ve~eis of constant volume (39-16 crn s) and with a constant quantity of the liquid silane. Thus, assuming a linear relationship between the HCI concentration in the liquid phase and the total HC1 pressure, the order of the reaction with respect to this component could be established. el) A. F. REID and R. MILLS,3.. Chem. Soc. 703 (1960). ~s) A. F. REID and R. MILt,s, J. Chem. Soc. 708 (1960). ~8~R. H. H~aD~R,3'. Chem. Phys. 27, 653 (1957). I,~ K. CLUStUsand H. HAIMEaL,Z. Phys. Chem. 51B, 347 (1942). ¢sj j. E. B o ~ s and L. O. BaOCKWAY,3". Amer. Chem. Soc. 77, 3444 (1955).
386
Notes
A best fit plot of eight independent runs at a trimethylchlorosilane concentration of about 1.38 mmoles in the reaction vessel gives a first order (slope ,-41.9) dependence on the HCI concentration. A typical set of data are summarized in Table 1. The root-mean-square deviation of the pseudo rate constant is 4-18 per cent over a sixfold change in the HC1 concentration. Assuming a first order dependence on the silane concentration as suggested by RZID and MILLS,tg~ the pseudo second-order specific rate constant is 4.22 × 10 -2 mole -x sec-x at 0°C. Despite considerable effort in reagent purification, rigorous exclusion of moisture and light, and use of an all-glass, stopcock grease-free system, the scatter of the data was rather severe. The reason TABLE 1.--EXCHANGE DATAFOR THE SYSTEM(CHa)3SP6C1 + HC1 Temperature: 0°C Volume: 0"03916 1.
Exp.
44 48 49 50 51 52
[HCI] (moles)
[(CHs)~SiCI] (moles)
(min)
0"697 × 10-8 0.717 0.454 2.00 1.94 2.73
1'38 × 10 -3 1'37 1.36 1"39 1.39 1"39
1.07 × 108 1.73 2.12 0.79 0.73 0.59
tl/2
R (moles min-0
R k = [HC1] (min-0
3'00 × 10 -e 4'30 × 10 -8 1'89 2-63 1"11 2"45 7"20 3"60 7"70 3"97 10.8 3.96 avg. 3-49(4-0.63) × 10-3
for this lack of replication may be that the reaction rate has a surface influenced component, as was found cs~true for the isotopic exchange reaction between CHBC1 and HC1 in which carefully outgassed (450°-500°C for 4 hr) reaction tubes resulted in very much slower exchange rates than were observed in untreated tubes. Because of this scatter it was not felt justified to pursue the present investigation in an attempt to establish more quantitative kinetic parameters. It is, however, interesting to note the qualitative implications of the observed exchange. From the results of RocHow et al., ~6~ REID and WlLKIm t~ and REID and MILLS,~} it appears clear that ionic mechanisms based on the self-dissociation of trialkylchlorosilanes are probably not important in effecting the halogen exchange between (CH8)8 SiC1 and SbC18, and it is probable that a similar situation obtains in the case of the silane as the primary solvent. A plausible exchange path thus presumably involves an activated complex in which the silicon atom exhibits a co-ordination number of five with a transition from spa to spSd bonding. Such a transition state does not, of course, uniquely characterize the nature of the adduct, which presumably is undissociated hydrogen chloride with the hydrogen atom simultaneously bonded to the incoming and leaving chlorine atom. While the methylchlorosilane liquid phase has an appreciably polar nature, it should be noted that halogen exchange between the silane and HC1 has also been observed in benzene solutions and thus the polar nature of the solvent can have, at best, only a small effect on the exchange rate. The observation, t3,s~ that no rapid isotopic halogen exchange between SIC14 and HCI or C12 is noted under similar conditions must be taken as an indication of the greater ease of formation of the quinquevalent transition state for the trimethylchlorosilane, as compared to the tetrachloride. This view is consistent with the postulate ~4~ that at room temperature the very slow exchange between SiCI~ and HC1 proceeds through an HSiC15 intermediate, but at rates very much lower than those reported for the (CH3)3SiC1 system in the present work. Experimental
Radiochlorine labelled (CH3)3SiC1 was prepared either by exchange with HseC1 or by exchange with (CH3)4Ns~C1 under anhydrous conditions as reported earlier. {9} All materials transfer was (S) K. GINGOLD,E. G. ROCHOW, D. SEYFERTH,A.C. SMITHand R. WEST,.]'. Amer. Chem. Soc. 7.-I,6306 (1952).
~7~A. F. REID and C. J. WlLKINS,d. Chem. Soc. 4029 (1955). ts~ R. H. HERBERand A. W. CORDES,d. Chem. Phys. 28, 361 (1958). to) R. H. HERaER,J. Phys. Chem. 62, 379 (1958).
Notes
387
effected by the usual high vacuum techniques in thoroughly dried glass systems. The reaction vessel, which consisted of a 10 nun Pyrex glass ampoule provided with a filling arm (subsequently sealed off) and a break-off tip, was covered with opaque paint and a layer of aluminium foil to exclude light, and pre-treated with inactive chlorceilane (subsequently pumped off) to insure anhydrous conditions. The heterogeneous exchange runs at 0°C were carried out in an ice-water bath held in a large capacity Dewar flask. Complete liquefaction of the solid (liquid N~ cooled) trimethylchlorosilanein 30 sec was noted in all runs. The reaction flask was shaken intermittentlyto insure mixing. At the end of the exchange run, the reactants were separated at --78°C and radioassayed, using the gas phase counting method described elsewhere,c1°~ Concentrations of the reactants were calculated on the basis of known gas dosage volumes, pressures and temperatures. No corrections for the solubility of HC1 in liquid (CHs)sSiCI were applied. Part of this work was carried out at the University of Illinois. The support of the U.S. Atomic Energy Commission is gratefully acknowledged.
Nuclear Science Center Rutgers, The State University New Brunswick, New Jersey
Departmentof Biochemistry
ROLF~ H. H~RBER
SHIH-CH~NCHANG
University of Pittsburgh Pittsburgh, Pennsylvania (le) R. H. HERBER,Rev. Sci. Instrum. 28, 1049 (1957).
Reaction of osmium tetroxide with iodide ions (Received 26 July 1960; in revised form 14 October 1960) Tim reduction of osmium tetroxide by iodide ions in the presence of hydrochloric acid yields a variety of products. Earliercl,j~ it had been ~eported that four equivalents of iodine are produced, but the other reaction products were not identified; although one expects to obtain Os(IV) e.g. as OsClsI-, the formation of improbable derivatives such as HsOs!, had also been suggested. ~8~ After the removal of the iodine from the solution a dark green-blue solution containing the Os(IV) compounds remains. This investigationset out to identify the nature of these reduction products. The results are of interest because unusual mixed crystals with potassium hexachloro-osmate of a new chloro-iodo osmium IV complex are formed. The reaction between osmium tetroxide and potassium iodide was carried out in hydrochloric acid solution to minimize hydrolysis and the iodine liberated was extracted with chloroform (see flow sheet). The yield of four iodine atoms per osmium tetroxide molecule wasconfirmed.~l,~) Extraction of the aqueous solution with ether removed a deep green ether-soluble compound but all attempts to isolate this as a solid material were unsuccessful. Decomposition occurred on evaporation under partial vacuum, even in the cold. Also, the green colour was destroyed when the ether solution was treated with drying agents; this ether solution contained osmium and iodine in the ratio 1:5, but no potassium. It was concluded that the green solute was probably H[OsIs.H20] or perhaps Ht[OsIsOH] but owing to its instability it was not further investigated. By concentrating the aqueous solution until crystallization just began, two different crystalline substances were isolated, one soluble--subsequentlyshown to be potassium hexa-iodo-osmate (TV)--and the other insohible in acetone. Small amounts of potassium chloride and iodide were removed by fractionation and the two substances then separated by treatment with acetone. The acetone insoluble portion, analysed as KjOsCls.TsIe.isand at first a mixed crystal of KmOsCle and KsOsIs was suspected. X-Ray powder photographs showed that the compound was cubic as for (x~ D. J. RIAirrscmKov,Y. Appl. Chem. Russ. 17, 326 (1944). (s) F. KaAUHand D. WXLKEN,Z. Anorg. Chem. 137, 352 (1924). (s~ E. P. ALV,~EZ, C. R. Acad. Sci.,Paris 140, 12.54, 1905; Chem. News 91, 173 (1905). 13
385
precipitate of the [Pt(en)d ++ -salt formed rapidly and provided a clean separation. With an initial, careful purification of the two salts by fractional crystallization, any apparent initial or separationinduced exchange was less than 3--4 per cent. However, a slow precipitation of Pt xv and the finite solubility of [Pt(en)t][B4418 limited the concentration range which could be employed. Exchange half-times were taken from the usual plots of log (1 -- F) vs. t (where F is the fraction of exchange). In general, the plots were satisfactorily linear over a period that the exchange was followed, usually about two half-times. The experimental conditions and the results are presented in Table 1. If the rate of the exchange is given by expression (1), the rate constant, k, is related to the half-time of an exchange experiment, ~'1/~ by the expression k = In 2/([Pt(en)2 +÷] ÷ [Pt(en)~Cl2++])[Cl-]~-l/~.
(2)
The indicated values of k are included in Table 1. The results, which include a considerable range of concentration variables, confirm the form of rate expression predicted for the exchange by BASOLOet al. ~ The average value for k, 12.1 sec-x M -~, at 25 ° is 30 per cent lower than their value; however, with the statistics indicated in the table this difference may not be real. The limited number experiments at 15°C indicate values for the enthalpy and entropy of activation, AH~ = 10.1 kcal, and AS ~: = --20 e.u. Institute for Atomic Research and Department o f Chemistry Iowa State University Ames, Iowa
L. T. Cox S.B. COLUNS D . S . MARTIN,JR.
Isotopic chlorine exchange between trimethylchlorosilane and hydrogen chloride (Received 20 October 1960; in revised form 21 November 1960)
THE recently reported
386
Notes
A best fit plot of eight independent runs at a trimethylchlorosilane concentration of about 1.38 mmoles in the reaction vessel gives a first order (slope ,-41.9) dependence on the HCI concentration. A typical set of data are summarized in Table 1. The root-mean-square deviation of the pseudo rate constant is 4-18 per cent over a sixfold change in the HC1 concentration. Assuming a first order dependence on the silane concentration as suggested by RZID and MILLS,tg~ the pseudo second-order specific rate constant is 4.22 × 10 -2 mole -x sec-x at 0°C. Despite considerable effort in reagent purification, rigorous exclusion of moisture and light, and use of an all-glass, stopcock grease-free system, the scatter of the data was rather severe. The reason TABLE 1.--EXCHANGE DATAFOR THE SYSTEM(CHa)3SP6C1 + HC1 Temperature: 0°C Volume: 0"03916 1.
Exp.
44 48 49 50 51 52
[HCI] (moles)
[(CHs)~SiCI] (moles)
(min)
0"697 × 10-8 0.717 0.454 2.00 1.94 2.73
1'38 × 10 -3 1'37 1.36 1"39 1.39 1"39
1.07 × 108 1.73 2.12 0.79 0.73 0.59
tl/2
R (moles min-0
R k = [HC1] (min-0
3'00 × 10 -e 4'30 × 10 -8 1'89 2-63 1"11 2"45 7"20 3"60 7"70 3"97 10.8 3.96 avg. 3-49(4-0.63) × 10-3
for this lack of replication may be that the reaction rate has a surface influenced component, as was found cs~true for the isotopic exchange reaction between CHBC1 and HC1 in which carefully outgassed (450°-500°C for 4 hr) reaction tubes resulted in very much slower exchange rates than were observed in untreated tubes. Because of this scatter it was not felt justified to pursue the present investigation in an attempt to establish more quantitative kinetic parameters. It is, however, interesting to note the qualitative implications of the observed exchange. From the results of RocHow et al., ~6~ REID and WlLKIm t~ and REID and MILLS,~} it appears clear that ionic mechanisms based on the self-dissociation of trialkylchlorosilanes are probably not important in effecting the halogen exchange between (CH8)8 SiC1 and SbC18, and it is probable that a similar situation obtains in the case of the silane as the primary solvent. A plausible exchange path thus presumably involves an activated complex in which the silicon atom exhibits a co-ordination number of five with a transition from spa to spSd bonding. Such a transition state does not, of course, uniquely characterize the nature of the adduct, which presumably is undissociated hydrogen chloride with the hydrogen atom simultaneously bonded to the incoming and leaving chlorine atom. While the methylchlorosilane liquid phase has an appreciably polar nature, it should be noted that halogen exchange between the silane and HC1 has also been observed in benzene solutions and thus the polar nature of the solvent can have, at best, only a small effect on the exchange rate. The observation, t3,s~ that no rapid isotopic halogen exchange between SIC14 and HCI or C12 is noted under similar conditions must be taken as an indication of the greater ease of formation of the quinquevalent transition state for the trimethylchlorosilane, as compared to the tetrachloride. This view is consistent with the postulate ~4~ that at room temperature the very slow exchange between SiCI~ and HC1 proceeds through an HSiC15 intermediate, but at rates very much lower than those reported for the (CH3)3SiC1 system in the present work. Experimental
Radiochlorine labelled (CH3)3SiC1 was prepared either by exchange with HseC1 or by exchange with (CH3)4Ns~C1 under anhydrous conditions as reported earlier. {9} All materials transfer was (S) K. GINGOLD,E. G. ROCHOW, D. SEYFERTH,A.C. SMITHand R. WEST,.]'. Amer. Chem. Soc. 7.-I,6306 (1952).
~7~A. F. REID and C. J. WlLKINS,d. Chem. Soc. 4029 (1955). ts~ R. H. HERBERand A. W. CORDES,d. Chem. Phys. 28, 361 (1958). to) R. H. HERaER,J. Phys. Chem. 62, 379 (1958).
Notes
387
effected by the usual high vacuum techniques in thoroughly dried glass systems. The reaction vessel, which consisted of a 10 nun Pyrex glass ampoule provided with a filling arm (subsequently sealed off) and a break-off tip, was covered with opaque paint and a layer of aluminium foil to exclude light, and pre-treated with inactive chlorceilane (subsequently pumped off) to insure anhydrous conditions. The heterogeneous exchange runs at 0°C were carried out in an ice-water bath held in a large capacity Dewar flask. Complete liquefaction of the solid (liquid N~ cooled) trimethylchlorosilanein 30 sec was noted in all runs. The reaction flask was shaken intermittentlyto insure mixing. At the end of the exchange run, the reactants were separated at --78°C and radioassayed, using the gas phase counting method described elsewhere,c1°~ Concentrations of the reactants were calculated on the basis of known gas dosage volumes, pressures and temperatures. No corrections for the solubility of HC1 in liquid (CHs)sSiCI were applied. Part of this work was carried out at the University of Illinois. The support of the U.S. Atomic Energy Commission is gratefully acknowledged.
Nuclear Science Center Rutgers, The State University New Brunswick, New Jersey
Departmentof Biochemistry
ROLF~ H. H~RBER
SHIH-CH~NCHANG
University of Pittsburgh Pittsburgh, Pennsylvania (le) R. H. HERBER,Rev. Sci. Instrum. 28, 1049 (1957).
Reaction of osmium tetroxide with iodide ions (Received 26 July 1960; in revised form 14 October 1960) Tim reduction of osmium tetroxide by iodide ions in the presence of hydrochloric acid yields a variety of products. Earliercl,j~ it had been ~eported that four equivalents of iodine are produced, but the other reaction products were not identified; although one expects to obtain Os(IV) e.g. as OsClsI-, the formation of improbable derivatives such as HsOs!, had also been suggested. ~8~ After the removal of the iodine from the solution a dark green-blue solution containing the Os(IV) compounds remains. This investigationset out to identify the nature of these reduction products. The results are of interest because unusual mixed crystals with potassium hexachloro-osmate of a new chloro-iodo osmium IV complex are formed. The reaction between osmium tetroxide and potassium iodide was carried out in hydrochloric acid solution to minimize hydrolysis and the iodine liberated was extracted with chloroform (see flow sheet). The yield of four iodine atoms per osmium tetroxide molecule wasconfirmed.~l,~) Extraction of the aqueous solution with ether removed a deep green ether-soluble compound but all attempts to isolate this as a solid material were unsuccessful. Decomposition occurred on evaporation under partial vacuum, even in the cold. Also, the green colour was destroyed when the ether solution was treated with drying agents; this ether solution contained osmium and iodine in the ratio 1:5, but no potassium. It was concluded that the green solute was probably H[OsIs.H20] or perhaps Ht[OsIsOH] but owing to its instability it was not further investigated. By concentrating the aqueous solution until crystallization just began, two different crystalline substances were isolated, one soluble--subsequentlyshown to be potassium hexa-iodo-osmate (TV)--and the other insohible in acetone. Small amounts of potassium chloride and iodide were removed by fractionation and the two substances then separated by treatment with acetone. The acetone insoluble portion, analysed as KjOsCls.TsIe.isand at first a mixed crystal of KmOsCle and KsOsIs was suspected. X-Ray powder photographs showed that the compound was cubic as for (x~ D. J. RIAirrscmKov,Y. Appl. Chem. Russ. 17, 326 (1944). (s) F. KaAUHand D. WXLKEN,Z. Anorg. Chem. 137, 352 (1924). (s~ E. P. ALV,~EZ, C. R. Acad. Sci.,Paris 140, 12.54, 1905; Chem. News 91, 173 (1905). 13