WESLEYTV. WESDLANDT AND JOHS L. BEAR
1516
TABLE Ir AlClz A N D KCI ON COLFFTCIE~XTS ASD SOLUBILITY OF Bi
‘r11E Eb’I
( T E M P FOR CURVES I 100
I
200
1 300
C AND D )
I
I
I
I
400
500
600
700
TEMPERATURE ' C .
Fig. 2.--Differential thermal analysis curves of the aquo-
p entammine- and anionpentaniminecobalt (111) complexes.
Curve A, [Co(SH?)hH201 (NO?)?;B, [Co(NH&NO31(?jO&; C, [Co(NH&H20]I3;and D, [Co(NH3)J]Iz.
Oil further heating, the monoanionic substituted cwnplexes decomposed with thP evolution of am-
iiinnia, nitrogen and ammonium halide, along JTith a subsequent oxidation reaction, to give the CozOe oxide levels in the 190 to 690" temperature region. The curve for [Co(SH,)jHZO]C13 had a break at 350" n-hich corresponded closely to the vomposition for CoC12. As was expected,23 [Co(SH~)~HzO](K03)3 exploded a t about 280-290°, leaving a residue of Co304which was less than calt 2 1 ) AI
klila17
C o m p t rend
249, 2780 (1959)
T'ol. G5
culated, perhaps due to the explosive violeiice of the reaction. The monoanionic substituted complexes were also studied by TGX and their curves were similar to those found for aquopentaniniinecobalt complexes, after the deaquation reaction. The DTA curves for the aquopentamminecobalt complexes, in general, gave an endothermic peak (or peaks) in the 100 to 125" lemperature region which was due to the deaquation reaction. This was followed by additional endothermic peaks (except for the nitrate complex) iii the 175 to 230" temperature range, which were caused by the decomposition of the moiioaiiionic substituted complexes. Additional endothermic and exothermic peaks mere also observed which corresponded to further decomposition and oxidation reactions. Only one of the complexes, [Co(NFI,);H201Cl,. gave two eitdothermic peaks for the deaquation reaction. It is tempting to postulate a two-step reaction for the mater loss; however, there is no evidence for this from the TGA curve. From the [Co(NH3)sH20]13curve, it can be seen that the complex begaii to decompose before the deaquation reaction was completed. Fair agreement \vas observed between this work and that previously reportedg for [Co(5&)&0](X03)3. The curve shows an endothermic peak, centered at, 110", follon-ed by a large exothermic peak a t 19.5"; previously reported values were 118 and 215", respectively. The yigorous exothermic nature of the decomposition reaction can be seen from the DTA curve. Kinetic Studies of the Deaquation Reaction.The kinetics of the deaquation reaction, as shown in equation 1, were evaluated by folloming the pressure increase in the system, due to the liberation of gaseous mater, by means of an isoteniscope. It mas assumed that the number of moles of gaseous water produced upon the completion of the reaction is equal to the number of moles of reactant initially present. Therefore, the final pressure, a. represented the number of moles of complex at the beginning of the reaction, and x, the pressure after timP t . represented the number of moles of water produced up to time t . ,4 plot of log a / ( a - x) uerszis t was made for each coinplex a t three different temperatures. A straight line v a s obtained for each temperature indicating that all of the deaquation reactions followed a first-order rate law. From the first-order rate constants obtained at various tenipcratures, the activation energies m r e calculated using the Arrhenius equation. The first-order rate constants and the actii-ation energies for the complexes are given in Table I. It was not possible to study the deaquation of [Co(SH3)jHJO]Idbecause of the serondarg dc] I ~ . activation composition of [ C O ( S H ~ ) ~ I The energy obtained for [Co(SH3)SH2O)Cl7 was ill fair aqreement with the 21.47 kcal. value previously reported by Nori, et a1.' The results reported here indicated a slight anion effect in that E* inincreased in the order: chloride < bromide < nitrate. However, the effect of particle size on the rate of deaquation was not studied. Reproducible rewlts \\-ere obtained on different preparations h a r i i ~ groughly the same particle size range.
Fc'pt., l % i l
FLUOHESCESCE AX 11 PHOSPHORESCENCE OF THIFLCOROACETONE TT.4po~
1519
Heat of Deaquation by DTA.-The A H of the TABLE 1 FIRST-ORDER KATE CONSTANTS AND ACTIVATIOK ENERGIESdeaquation reaction was determined by the method of quantitative DTA. This method is FOR T E E !lQUOPENTA\IMIKECOBALT( 111) C O M P L E X E S Ttmp., C.
-log
E*,
k
Order
kcal.
86 90 95
3 37 3 24 3.10
1
19f2
[ C O ( S H : ) ~ H ~ O ] B ~ ~ 85 90
1
25&2
93.7
3 32 3 12 2 94
[ C O ( ; I ; H ~ ) ~ H ~ O ] ( X O82 ~)~ 85 88
3.77 3 60 3 45
1
31k3
Compound
[Co(NHa)sHzO]CIA
Dissociation Pressure Measurements.--illthough Lamb aiid nlarden5 concluded that the transformation illustration in equation 1 was noiireversible in the solid state, hIori, et aZ.,'assumed that they obtained equilibrium dissociation pressures for the chloride complex in the 21 to 48" temperature range. They reported a dissociation pressure of 2.15 mm. a t 25" and 6.93 mm. a t 48". Previous rough measurements5 yielded values of approximately 5 , 4 and 4 mrn. for the chloride, bromide and nitrate compkxes. respectively, at 25". I n this investigation, ~ i attempt i TVRS made to confirm the results of Nori, et al.,s and establish the reversibilitg of the reaction. For the [Co(SH3)5&O]Cl3 complex, 2.fter 100 hr. a t 28.0", a dissociation pressure of 38.0 mm. was obtained. The temperature then was lom-ered to 25.0" and after 500 hr. the preseure did not change from its initial Yalue. Similar results were obtained for the bromide and nitrate complexes. ilIoriZ4stated that the system reached an equilibrium dissociation pressure after about 12 hr. a t any specific temperature. Howes~er, it is apparent from the above results that the system is not reversible. (24) RI. 3Iori private communication.
based upon the premise that under certain experimental conditions, the heat of reaction can be evaluated by integration of the differential curve peak of peaks. From the various theories on quantitative DTX, that of Speil, et a1.,25 states that the peak area is
where tl and t:! are the time limits of the peak, 0 is the differential temperature, AH is the heat of reaction involved in the chemical change, M is the mass of reactive sample present, X is the thermal conductAvity of the sample, and g is a constant dependent, on the furnace and sample holder geometry. To det'ermine the above constant's, the apparat,us was calibrated by studying reactions of known hhermal effects. The heat's of deaquation, a t 108", obt'ained by use of the above method were 6.1 and 7.8 kcal. mole-' for [Co(XH3)5H*O]C13 and [CojKH3)5HzO]Br3, respectively. Because of the many variable factors involved, no great accuracy is claimed for the above results. However, the values obtained are probably indicative of the true heats of deaquatioii. Acknowledgments.-The assistance of T. D. George, W. Robinson and P. Ruhii is gratefully acknowledged. This work was supported iii part by t'he U. S. Air Force, Office of Research aiid Development, t'hrough Coiit'ract X o . AF-49(638)787. (2,5) S. Speil, L. H. Berkelhamer, J. A. Pask and B. Davis. Bureau of Mines, Tech. Paper 664 (1945). (26) 31. J. Vold. A n d . Chem.. 21, 683 (1949). (27) S. L. Roersma, J . Am. Ceyam. Soc., 38, 281 (1955).
I?, S.
THE FLUORESCEKCE AKD PHOSPHORESCESCE OF TRIFLUOROACETOSE VAPOR1 BY P. A c s ~ o o sAND E. MVRAU* iYatzonal Bureau o j Standards, Washington, D.C. Recezued Febi u a r y 28, I 9 6 1
The fluorrscence and phosphorescence of trifluoroacetone has been investigatd at 2652, 2804, 3025, 3130 and 3341 A. The effect of concentration and temperature on the yields of triplet and single-state emissions are comparable to those obP P ~ V Rfor ~ acetone The emissions from 2-butanone and 2-pentanone have been investigated briefly Both compounds phosphoresce v e x TI eakly and their fluorescence yields are nearly identical with those observed for acetone and trifluoroacetone.
Introduction Recent studies3$ on t,he fluorescence of hexafluoroac:et,onc: have c l ~ n r l yyhown that the emis11) 'I'liis I C Q C ~ ~ C I I Has supliortril iii ljai't tJy a g i a n t from tile U. R. l'Lil,lic llriiltli , S ~ r v i v v , I ~ v ~ ~ a , t , i t< ~> f~ Irl ri ut l t l i , I G I i i c ~ t i o n ,an(l \VPIfaie.
( 2 ) Natiorial Academy of Sciences-National Research Council Postdoctoral Research Associate 19.59-1960. (3) H. Okabe and E . W.R. Steacie, Can. J . Chem., 36, 137 (1958). (4) G . Oiacometti, II. Okabe and E . W.R . Steacie, Proc. R o y . SOC. (London), A260, 287 (1959).
sioii from this compouiid differs considerably froin that reported for a c e t , ~ n e . ~For # ~ inst'ance, iii the case of hexafluoroacetone, emission was observed from t h e upper singlet-stat,e, arid t,he yield pas: st,ror~glydrpendrnt, on tempr,rat,ure and pressure. 111 coiit i~ast, \vit,li t,liis I x h v i o r , t,he p}iohphorrsc.c.l,c.~~ of acetone niay be quite strong, depending 011 (5) For a review s e e : IT. A . N o y e s , Jr., G. B. Porter and J. E. Jolle>-, Chem. Reus., 6 6 , 49 (1956). (6) J. ISeicklen, J . Am. Chem. Soc., 81, 3863 (1959).
