August, 1961
nECOMPOSITIOS ~ I X K L ’ I C SOF I J T H I U M 1’EHCHLOHA’l’L
TABLE I1 MOLEFRACTIONS OF CONFORMATIONAL ISOMERS, AND ENERQY DIFFERENCES BETWEEN THEY, FOR SOME CYCLOHEXANE DERIVATIVES This research Substance
trans-I,2Dichlorocyclohexane
zce
0.54
A I = Eso Ese
-
140b
Literature AE (g)
AE (bz) AE(CC1d
100
300’ - 503 4003 +1405 6504 8205 -SO3 -5003 -7 0 4 - 4004 -705 - 7005
trans-1,20.41 -330 Dibromocyclohexane -30 trans-1-Chloro- 0.49 2-bromocyclohexane 1,4Cyclo0.88 1900 hexanedione’ a z = fraction. b 811 energy values are in cal./mole.
3704
-2205
5605 13007
much the difference in AE would be expected to be about the same in all three cases. The reversal of the sign of AE(so1vent) - AE(g) in the case of carbon tetrachloride and n-heptane indicates that the above theory is not adequate. Specific solvent
1419
effects may exist which stabilize the non-polar la2a form in these solvents, contrary to the prediction of electrostatic theory, or As may not be aero. The apparent dipole moment of 1.4 D observed for 1,4-~yclohexanedionein the vapor state corresponds to about 12% of the boat form V if equilibrium between forms IV and V is assumed. The energy difference AE = E(boat) -,!?(chair) = 1900 cal./mole is computed on this latter assumption. However, these calculations have two serious defects. First, the atomic polarization, which we have neglected, may be except,ionally large for this molecule and a value P A = 10 cc./mole is not unreasonable by analogy with l14-benzoquinone.lo This would reduce the experimental moment to 1.1 D. Second, the boat forms of type VI may be present since they presumably do not differ much in energy from V. If the energy of forms V and VI is assumed equal the total contribution of these forms (neglecting atomic polarization) would be 22% and AE would become 2250 cal./mole. The value of A E does decrease on going from vapor to benzene solution as would be expected if the polar boat form is sta.bilized in solution relative to the non-polar chair form. (10) C. P. Smyth, “Dielectric Behavior and Structure,” McGrawHill Book Co., Ino., New Y Irk. N. Y., 1955.
THE DECO1CIPOSITIOK KINETICS OF LITHIUM PERCHLORATE BY MEYERM. MARKOWITZ AND DANIEL A. BORYTA Foote Mineral Co., Research and Engineering Center, Chemicals Division, P. 0. Box 615,West Chester, Pa. Received April 1, 1961
The thermal decomposition of pure lithium perchlorate and in admixtures with lithium chloride was studied over the temperature range of 392-415” by constant temperature thermogravimetry. It is shown by phase data that above 247’ any mixture of lithium perchlorate with its decomposition product, lithium chloride, always contains the perchlorate in solution. Over the temperature range covered, the kinetics follow the autocatalytic Prout-Tompkins rate law EA,^. = 52.2 f 4.1 kcal./mole) up to about 40% decomposition and then conform to first-order kinetics EA^^. = 62.0 =!= 4.1 kcal./ mole). The point of transition between the two rate laws occurs when the decomposing melt is saturated with lithium chloride. From the kinetic data, the solubility of lithium chloride in the melts was computed and a kinetically derived value for the heat of fusion of lithium chloride was obtained. The relationship of these studies to the thermal decompositions of the other alkali metal perchlorates is discussed.
Introduction The participation of a thermally decomposing material in the phase change from the solid to t.he liquid state usually complicates the observed kinetics.l Thus, potassium perchlorate has been reported t,o decompose according to two first-order rate laws2.3; one characteristic of the solid-phase decomposition and the other of the liquid-phase. Differential thermal analysis ~ t u d i e s ~of- rubidium ~ and cesium perchlorates have demonstrated the concomitant occurrence of fusion and decomposition phenomena, thereby indicating complex kinetics for these compounds. Lithium perchlo(1) C. E. H. Baivn in “Chemistry of the Solid S t a t e , ” edited by W, E. Garner, Academic Press, New York, N. Y., 1955, pp. 254807. (2) A. E. Harvey, Jr., hl. T. Edrnison, E. D. Jones, R. -4. Seybert a n d K. A. Catto. J . Am. Chem. Soc., 76, 3270 (1954). (3) A. E. Harvey, C. J. Wassink. T. A. Rodgers a n d K. H. Stern, Ann. N . Y . dead. Sci., 79, 971 (1960). (4) hl. M. h‘larkowitz, J . P h w Chem., 62, 827 (1958). ( 6 ) M. Xf. h‘larkowitz a n d D. A. Boryta, ibid., 64, 1711 (1960). 16) M . Jf. J l a r k o a i t a , D. A. Boryta a n d R. F. Harris, ibid., 66, 2G1 (19til).
rate,4-6 on the other hand, shows a considerable temperature interval between fusion (247’) and measurable rates of reaction (392415O).’ Accordingly, it was felt that a study of the thermal breakdown of lithium perchlorate would be of interest inasmuch as the salt probably would be in the liquid phase during the entire period of decomposition. Thus, the system LiC104-LiCl was studied and the quantitative kinetic behavior of lithium perchlorate and of mixtures of lithium perchlorate with lithium chloride was investigated. Experimental Anhydrous lithium perchlorate, prepared as previously reported,* was analyzed for perchlorate content by precipitation as nitron perchlorate.8 Chloride was determined gravimetrically as silver chloride and chlorate was computed as the additional chloride produced after reduction by sulfurous acid. Analysis of product: c104-, 93.4 (calcd., (7) M. M. Markowitz a n d D. A. Boryta, Anal. Chem., Sa, 1588 (1960). (8) F. P. Treadwell a n d W. T. Hall, “Analytical Chemistry,” Vol. 11, John Wiley and Sons, Inc., New York. N. Y . , Ninth English Edition, 1942,pp. 383-385.
MEYERM. MARKOWITZ AND DANIEL A. BORYTA
1420
93.5); C1-, 0.00; (YOa-, 0.00. Reagent grade lithium chloride, dried under vacuum at 150" for 12 hours, was analyzed for chloride content. Analysis of product: C1-, 83.4 (calcd., 83.6). All materials were stored in a desiccator over phosphorus pentoxide. The System LiC!104-LiC1.-The phase relationships between lithium perchlorate and lithium chloride were determined by differential thermal analysis and by visual observations of the liquidus temperature^.^^^^^ For the former studies, sample compositions varying from pure lithium perchlorate to pure lithium chloride in 10 mole % increments were wed; for the latter studies, Sam le comLiclol positions from 100 mole % LiC10, to 90 mole were covered. Because of the thermal instability of the mixtures at elevated temperatures, these experiments were only carried out t o about 300'. The persistence of an endothermic break a t about 235' over the entire range of composition in the differential thermal analysis runs indicated the system to be of the simple eutectic type and the observations of liquidus temperatures gave the correeponding eutectic composition a t 91.0 mole % LiC10,. To extend the phase diagram to higher temperatures, the Clapeyron-Clausius equation was applied to the entire system.10 The heat of fusion of lithium perchlorate6 was taken to be 3.8 kcal./mole and that of lithium chloride,ll as 4.7 kcal./mole. The computed eutectic was found to be a t 87.2 mole % ' LiClO, and 228', in good agreement with the experimental values. A comparison of these two sets of data is presented in Table I. Clearly, a t temperatures of 247' and above, any mixture of lithium perchlorate and lithium chloride wi!1 always contain the lithium perchlorate in solution.
$
TABLE I THESYSTEMLiC104-LiC1 Mole % LiCiO,
100.0 97.5 95.0 92.5 91.25 91 .o 90.5 90.0 87.2 85.0 80.0 70.0 60.0 50.0 40.0 30.0 20.0 10.0 0.0 a By visual analysis.
Calcd.
Liquidus temp., OC. Expt1.n
Eutectic temp., "C., exptl. b
247.0 247 242.8 245 239.9 239 236.2 238 234.4 237 234.0 236 (eutectic) 233.4 239 232.5 252 228.0 (eutectic) . . 245.5 .. 280.1 .. 338.5 .. 387.6 .. 431.6 .. 472.1 .. .. 510.1 546.3 .. 581.3 .. 615.0 .. observation. b By differential
..
.. .. *.
.. ..
.. 240
.. ..
242 235 232 233 236 232 232 230 *.
thermal
The Products of Decomposition of Lithium Perchlorate.Earlier studies had shown that the over-all decomposition of lithium perchlorate could be re resented by the equation*JJ2-14: LiClO,( -c LiCl 282. During the present study, chemical analyses were performed on the decomposi-
+
(9) M. M. Markowitz and R. F. Harris, J . Phys. Chem., 6% 1519 (1959). (lo) W. C. McCrone, Jr., "Fusion Methods in Chemical Microscopy," Interscience Publishers, Ino., New York. N. Y., 1957, pp. 156-157. (11) National Bureau of Standards Report No. 6297,U. S. Department of Comniercc, Washington, D. C.,Jan. 1, 1959, pp. 31, 72. (12) G . G. Marvin and L. B. Woolaver, Ind. Eng. Chem., Anal. Ed., 17,474 (1945). (13) T. W. Richards and H. H. Willard, J . Am. Chem. Soc., 32, 4 (1910). (14) T. W. llicharda and M. W. Cox, ibid., 36,819 (1914).
T-OL
65
tion residues from two-gram samples of lithium perchlorate which had been maintained at 413 f 3" for varying periods of time. I n this way it was hoped to determine the existence of appreciable amounts of reaction intermediates and side reactions. Small quantities of a lower oxyhalide (assumed to be chlorate) were present in the partially decomposed materials. The alkalinity of the residues, determined as % Li20 by titration with standard acid, indicates some slight loss of chlorine or chlorine oxides during decomposition. Chloride was determined volumetrically with silver nitrate by Mohr's method and chlorate was fou,nd as additional chloride after reduction of the sample with sulfurous acid. Lithium perchlorate was computed by difference from 100.00% of the sum of the three separate analyses. The data are summarized in Table 11.
TABLE I1 COMPOSITION (IN WT. yo) OF LITHIUM PERCHLORATE SAMPLES HEATED AT 413 3" %
Time, nun.
LiCl01 (by diff.)
%
LiClOs
%
Liz0
% LiCl
98.28 0.60 0.00 1.12 93.45 1.34 .02 5.19 91.10 1.35 .03 7.52 79.85 1.72 .04 18.39 78.74 1.80 .04 19.42 69.36 2.03 .05 28.56 58.70 1.28 .08 39.94 45.99 0.53 .10 53.38 37.40 0.85 .12 61.63 5.25 0.00 .14 94.61 % ' LiClO, from weight loss data 11. 100 % LiC104from chemical analyses
50 80 90 105 110 115 130 140 160 260
[
%
Dev.a
$0.3 $1.0 $1.0 $0.8 $0.9 $0.8 -1.5 -4.0 -4.7
...
-
The lithium perchlorate contents of the samples of Table I1 were compared with the compositions derived from t,he weight losses of the samples. Close agreement, between the two sets of values wae found. This is indicated in the last column of Table I1 by the per cent. deviation of the lithium perchlorate content, determined by weight loss measurement from that determined by chemical analysis of the decomposition residues. Accordingly, i t was concluded that a convenient means of following the progress of reaction would be by constant temperature thermogravimetry . Constant Temperature Thermogravimetric Analysis of Lithium Perchlorate and Lithium Perchlorate-Lithium Chloride Mixtures.-Differentia1 thermal analysis studies of the alkali metal perchlorates show that a t the temperatures of rapid decomposition, the reactions are exothermic .6*s Accordingly, it was necessary to determine the maximum temperature at which lithium perchlorate could be decomposed without self-heating and consequent departure from isothermal conditions. Constant temperature differential thermal analyses were run on two-gram samples of lithium perchlorate a t temperatures below those of rapid decomposition in a dynamic differential thermal analysis experiment (472'). The maximum temperature at which the decomposing masa remained at the furnace temperature through to complete decomposition of the sample was about 420'. It was necessary t o carry out these experiments to complete decomposition because the heat release from the sample increases as solid lithium chloride forms and the heat of decomposition is augmented by the heat of cryst,allization of the lithium chloride. Thus, the thermal decomposition of a molten alkali metal perchlorate must usually be represented by two general equations,6Jb mz., (a) x1cIo4 (solution) + XW1 (solution) 202 (gas), AH., and (b) x x c l o 4 (solution) -,XIC1 (solid) -I- 202 (gas), AHb, and xiC1. In the present studies, m b - AH, = - AHfUeion isothermal experiments were carried out to a maximum temperature of about 415" using sample weights containing 0.5-0.8 g. of lithium perchlorate. The instrumentation attendant to and the arrangement of the recording balance and the furnace have already been described.' A new furnace was constructed for the investigations reported here. Thie 20" vertical furnace consisted of a Is/," Vycor tube uniformly wound with 'asbestos-
+
(15) M. M. Markowits, J. Phys. Chem., 61, 505 (1957).
DECOMPOSITION KINETICS OF LITHIUM PERCHLORATE
August, 1961
covered nichrome ribbon. The exterior of the furnace was heavily insulated with asbestos tape and pipe lagging. A two-inch zone of constant temperature waa found in the furnace interior by movement of a thermocouple through the capped bottom of the furnace. During a run the sample waa contained in a small test-tube suspended by a platinum wire from the balance above the furnace. The wire entered the furnace through a ceramic cap with a radial slit along the length of the cap. The sample was consistently placed about from the measuring thermocouple in the isothermal region of the furnace. Close temperature regulation ( =k0.2') was achieved with an electronic controller by using the seriea-connected output from three thermocouples placed about the circumference of the Vycor tube and sandwiched between the tube and the heating elements. The cold junctions of the thermocouples were kept in a constant temperature bath. In each experiment the sample was preheated for ten minutes a t 275" and then suspeuded quickly from the balance into the furnace below. By using a stepping-switch, furnace temperature and weight changes were alternately recorded on a potentiometric strip-chart recorder of known speed. Weight loss-time measurements were made on three tvpes of samples: pure LiC104, LiCIOp (95 mole %)LiCl ($mole %), and LiClO, (50 mole %)-LiCI (50 mole %), and over the temperature range from about 392 to 415'.
Results Pure Lithium Perchlorate.-A
typical decornposition-time curve for pure lithium perchlorate is shown in Fig. 1. A long induction period, folTime, minutes, refers to pure LiClO,. -if3 336 396 456 516 576 636 696 756 816
' ) 31
r
loo
t
-
r
1
LiCLO4 (100 MSLE %); 395.3"C. o LiCLOi ( 9 5 ) ,LICL (5);395.99;. e -
LiGL04(5@) , LICL (50);395.3"C:
1
0 60 120 180 240 300 360 420 480 540 Time, min., refers to LiClOd-LiCl mixtures. Fig. 1 .-Typical decomposition-time curves for lithium
perchlorate and for lithium perchlorate-lithium mixtures.
chloride
lowed by rapid acceleration and slow decay is evident. The presence of the induction period gives rise to the sigmoid curve usually characteristic of the autocatalytic decompositions of many solids.I6 Because the decomposition of the lithium perchlorate proceeds from the liquid phase over the entire extent of decomposition, it was thought that a single rate law might be applicable. The data for a representative set of results as obtained (16) P. TI'. M. Jambs and F. C. Tompkins in "Chemistry of the Solid State." edited by W. E. Garner, Academic Press, New York. N. Y.,1955, pp. 184-211.
1421
10 8 6
LiGLO4; 391.8'C.
4l
i:::[ /i7.-....'..2
h
0.2
0.03 396
Fig. 2.-Integrated
516
636 756 876 Time, min. rate plots for pure lithium perchloratc at 391.8'.
a t 391.8' are given in Fig. 2 in terms of a firstorder plot (In (Wt/Wo) versus time, rate law: d(Wt/Wo)/dt = -kz(Wt/Wo), where Wt = weight of LiCIOcpresent at time t and Wo= original sample weight of LiC104) and a P r o u t - T ~ m p k i n s ~ ~ - ~ ~ plot (ln(Wo - TVt)/Wt versus time, rate law: dWt/Fo)/dt = -ki(wt/Wn)((wo - wt)/Wo)>. Neither rate law satisfies the experimental data over the complete course of reaction. Nevertheless, it appears that each law is valid over a limited region of decomposition and that the transition from the autocatalytic expression to the first-order equation occurs a t about 40% decomposition. The functions (Wt/Wo)and (Wo - Wt)/W+.are, of course, equal at Wt/Wo = 0.618, but this coincidence is irrelevant to the observed departures of the two rate laws from linearity in this region of decomposition. Equivalent results were obtained a t thc other temperatures studied. Table 111 contains a summary of the rate constant computed for the acceleratory period ( I C ~ from Prout-Tompkins plot) and for the decay period (kzfrom the firstorder plot). Lithium Perchlorate (95 mole %)-Lithium Chloride (5 mole %) Mixtures.-Ik characteristic decomposition-time curve for this type of sample is seen in Fig. 1 for a run a t 395.9'. The induction period is eliminated leaving only the acceleratory and decay sections. Rate law plots of the data for an experiment a t 402.5" (Fig. 3) show the adherence t o an autocatalytic law up to about 400/, decomposition and then a transition to first-order kinetics. A tabulation of the corresponding values of kl and Lz is given in Table IV. Lithium Perchlorate (50 mole %)-Lithium Chlo(17) E. G. Prout and F. C. Tompkins, Trans. Faraday Soc., 40, 488 (1944); 44, 468 (1946). (18) P. J. Herley and E. G . Prout, . I Inorg. . and Nuclear Chsm., 16, 16 (1960). (19) P. J. Herley and E. G. Prout, J . Phys. Chem., 64, 675 (1960).
MEYERM. MARKOWITZ AND DAXIEL A. Boltma
1422
(kz)
PERIODS OF
0.9
TABLE I11 ACCELERATORY (kl) AND DECAY D E C O M P O S I T I O N O F P U R E LITHIUMP E R -
RATECONSTANTS FOR
Vol. 65
0.8
THE
0.7
CHLOR4TE
Temp.,
kl,
OC.
391.8 395.3 399.8 404.7 406.8
a
0.00378
0.0106 ,0137 ,0178 ,0222 ,0261 ,0312 ,0402
410.5 414.2
Icdkl,
k2. min.-'
Inin.-'
.0053G ,00675 ,00948 ,0112 ,0137 ,0187
XLiCi.
%
0.6
mici
calc.
0.357 ,391 ,380 ,426 ,443 ,438 .4i2
0.409
-12.i
,417 ,427 ,438
-6.2 -11.0
0.5
-2.7 +3.2 -2.7 f2.9
0.4
,429 ,451 ,459
der.n
1
i""'-.1-
100
[ZLiCI,
calo.
,--.
G 0.3 b
v Lc
TIBLE
1i-
R A T E C O N B T A S T S F O R THE -4CCELERATORY
(ki) AND DECAY
( k 2 ) PERIODS OF DIXOMPOSITION OF LITHIUM PERCHLORATE I N A LITHIEMPERCHLORATE (95 MOLE y o ) - L CHLO~ ~ ~ ~ RICE
Tynp.,
kl. Inin.
C.
-i
( 5 MOLE
kr/ki. zi,ici
calc.
0.399 ,329 ,354 ,387
0.409
0.0114 ,0149 ,0205
0.00392
402.5
,0233
,00904
409.5
,0332 ,0435
,0126
413.6
00492 ,00724
,0175
,381 ,403
XLGI,
,418 ,425 ,443 ,449
,458
Q
dev.a
-
2.4 -21.3 -16.7 -10.6 -15.1
-12.0
ride (50 mole %) Mixtures.-These
mixtures exhibited decomlposition-time curves which were completely deceleratory (Fig. 1). From the phase data, this composition is saturated with lithium chloride at the decomposition temperatures studied here. Interpretation of the weight loss results in terms of a first-order law appears unequivocal as judged from inspection of Fig. 4 for a run a t 406.6'. 10
~ ~
5%) MIXTURE
k?, inin. -:
391.5 395.9 398.8
0.2
r-
I
LiCLO4 195) , LiGL ( 5 ) ; 402.5"G.
LiGL04(50), Li CL (50); 406.6"C 0.1
The Prout-Tompkins rate law gives an integrated function which is considerably bowed when compared to the clearly defined straight line of the first-order plot. The latter rate constants (k2) are listed in Table V. TABLE V RATEC O S S T A N T S ( k z ) FOR T H E DECOMPOSITIO\ O F L1THIL.M PERCHLORATE IN A LITHIUMPERCHLORATE (50 MOLE %)LITHIUMCHLORIDE (50 MOLE %) Mixture Temp., OC
4
2
c
30
BO 90 120 144 Time, min. Fig. 4.-Integrated rate law plots for a lithium perchlorateo(50 mole yo)-lithium chloride (50 mole %) mixture a t 406.6 .
6
Mixture k2,
mln.
Temp
k2, inin - 1
O C
-1
391 2 395 3
0.00381 .00549
406 6 408 9
0 0112
399 8 402 5
.00724 ,00946
413 0
0181
0131
Discussion '1'he autocatalytic and first-order rate laws applied to the decomposition of lithium perchlorate appear to coincide a t about 40% decomposition. Reference to the phase data shows that over the iiarrow temperature range covered in these studies, this extent of decomposition corresponds to saturation of the melt with the reaction product, lithium chloride. At saturation then the two rate laws should be equal, and hence Transposing terms, it is found t,hat
84 114 144 171- 204 Time, min. Fig. 3.-Integrated rate law plots for a lithium perchlorate (95 mole :%)-lithium chloride ( 5 mole %) mixture
24
at N 2 . 3 O .
54
9 ki
=
(3g5) eat,,
Thus, the ratio (kn/kI) should be equal to the niolr fraction solubility of lithium chloride in the saturated melt, Z L ~ C ~ . Tables I11 and I T - contain com-
DECOMPOSITIOR' KINETICS OF LITHIUMPERCHLORATE
August, 1961
100 90 80 70 60 50
parisons of this ratio with the mole fraction solubility of lithium chloride as computed from the Clapeyron-Clausius equation. The disparity between these two sets of results corresponds to an average deviation of 10.8%. Carrying equation 2 one step further
(EZJ?!)
1423
0
L ~ C L O(9.51, ~ L;CL ( 5 )
e
LiCLQ (SO), LiCL (50)
40
30
= aatn
xL1cl =
constant exp( -AHruslon,L,cI/RT) ( 3 )
Thus, a plot of the logarithm of xL,cl versus 1/T should yield a straight line, the slope of which is the heat of fusion of lithium chloride divided by -R. This, of course, assumes ideal solution behavior of the melts. The kinetic data result in a value of 6.6 f 0.5 kcal./mole for the heat of fusion of lithium chloride, compared with 4.7 kcal./mole for the calorimetrically derived value. l1 ,Isa consequence the present study is of interest in that it allows for the partial determination of a phase diagram and for solubility values from kinetic data. Knowledge of the saturation solubility of lithium chloride in-the decomposing melt at any given temperature makes possible the interrelation of the two rate laws over the entire range of decomposition. Thus
s 0'
20
x
h
,s *
1;
K
v
%
S 7 6 5 4
3 I
1.5
1
I
1-15
146
I
I
147
148
I
149
I
150
I
151
152
1jT x 105. Fig. 5.--Arrhenius plots for aceeleratory ( k ,) and decelerstory ( k 2 )rate constants.
The solubility data of Tables I, 111 and IV indicate the occurrence of a small positive deviation ____ d(wt/'T'-o) = -kl(zLiC1)( wt/Wo);melt satd. (5) from ideality in the system LiC104-LiC1. Though dt the numerical values of the saturation solubilities or most generally for the latter rate law are not greatly affected by this non-ideality, the temperature coefficient of solubility is strongly in'(WtlWo) = -kl antilog fluenced. Thus, the high values for the heat of dt fusion of lithium chloride derived from the kinetic data (6.6 and 9.8 kcal./mole us. 4.7 kcal./mole for where Tois the melting point of lithium chloride in the known value) are consistent, with this type of departure from ideal solution behavior. "K. The Prout-Tompkins rate expression as applied Arrhenius plots of In IC versus 1/T yield satist o the decomposition of solids was originally defactory straight lines for kl as obtained from pure lithium perchlorate and from the LiC104 (95)- rived on the basis of branching reaction nuclei, LiCl ( 5 ) mixtures and for kz as obtained from usually structural defects on the surface of the these materials and from the LiCIOl (50)-LiCl solid possessing a low energy of activation for de(.io) mixtures (see Fig. 5 ) . The corresponding composition.16317Acceleration of reaction is energies of activation and frequency factors are pictured to occur as these branching chain nuclei given in Table VI and were obtained from least form reactive planes in the material giving rise i o mechanical stresses and physical breakdown squares treatment of the data. of the solid crystals. Eventually these planes 'FABLE 1'1 interfere with one another and the liiiwtics assunic first-order hehavior. The liquid-phase decompo4I < n . ~ ~ C'ONSIANTS rc FOR THE LIQUID-PHASE THER~~IAI, tion of lithium perchlorate bears r? DECOMPOSITION OF LITHIUM PERCHLORATE scmblance to this pictiire. Clcarly. the function ki k2 (acceleratory) (deceleratory) of lithium chloride must iiivolre sotile traiisfcr f i ' : ~,~kcal./mole t 5 2 2 f 4 1 62 0 f 4.1 inechanisni facilitating rupture of thc C1-0 bond, A , set.-' 2 7s x 10'3 1 55 x 10'6 viz., LiC104(solution) LiCl (solutiotk) + LiC10, (solution) l, ?OL (ga.) LiCl (solutioii). The The ratio iX;2/kl) in terms of the exponential functions should also yield the heat of fusion of miall concentration of lithium chlorate> in the rcacation residues indicates that the dwomposition lithium chloride, uix. of the chlorate (LiClO? isolutionj -+ LiCl (Folutiou L A = A eup( - E J R T ) / A ' exp( --E,/R2') or solid) '2 0 3 (gas,) must be qmtr rapid coin= constant exp( -(E2 - E l ) / R T ) pared to the ratc of perchlorate drcomposition. = constant exp ( - AHrusion, L,cI/RT) ( 7 ) Thus, up to about 40% decompositioii the deconiUsing the data of Table VI, a value of 9.8 f 8.2 posing perchlorate generates more nriclei t o ackcal./mole was obtained for the heat of fusion celerate its o~vnbreakdown. Hovever this acof lithium chloride. celerating autocatalytic effect can oiily proceed followed by
+
+
+
+
1424
MEYERM. MARKOWITZ AND DANIEL A. BORYTA
to the point of saturation of the melt with lithium chloride; this represents the point of maximum rate achievable in these studies and thus, in a sense, a chainbreaking step. At saturation, the ratio of lithium perchlorate to lithium chloride is perforce a constant and thus formal first-order kinetics are observed. I n this regard, it is of interest to note that the efficacy of lithium chloride as a catalyst only occurs in homogeneous solution with lithium perchlorate. The decomposition kinetics found for lithium perchlorate in the present study are in contrast to those currently accepted for pot’assium perc h l 0 r a t e . 2 ~ ~For ~ ~ ~the latter compound first-order kinetics prevail over the entire decomposition region, indicating no autocatalytic effect by the potassium chloride reaction product. Earlier work on potassium perchlorate by Glasner and Weidenfeld21was interpreted by them in ternis of a Prout,Tompkins mechanism despite the obvious phase change of the salt from solid to liquid during decomposition.22 The mechanism stipulat’ed exchange of oxygen betmeen chloride and perchlorate as a result of the postulated rquilibrium: KC1 4(0) KC104. Bosch and A t ~ n ?studied ~ the distribution of radioactive chlorine between NaC1* and .KaCIOa at elevated temperatures. They concluded that, the low activity of the labeled chlorine in the perchlorate formed invalidated the earlier interpretations of Glasner and Weidenfeld. However, it seems possible, on the basis of .the studies reported here, that the thermal deconipositioiis of the alkali met’al helates may involve differences in the type of mechanism and not merely differences with t,he same fundamental kinetics. If this be so, t’hen the results of exchange ;studies based on the sodium salts cannot be applied reliably to the mechanisms of the corresponding potassium and lithium systems. It might be expected that. the differciice in the
+
(20) K. 13. Stern and 11. Bufalini, J . Phys. Chem., 64, 1781 (1960J. (21) A. Glasner a n d L. Weidenfeld, J . Am. Chena. Sac., 74, 2464, 2467 (1952). (22) L. :L. Bircumsliaw a n d T. R. I’liillips, J . Chem. Sac., 703 (1953). (23) A. ‘17. Bosch a n d A. IT. W. Aten, J . Am. Chem. SOC.,76, 3835
11953).
T’oi. 65
activation energies between the liquid- and solidphase decompositions of potassium perchlorate would be equal to the heat of fusion of the compound. This latter value has been computed to be small (1.7-2.6 kcal./mole)ls because of the appreciable heat of crystallographic transition absorbed a t 300” (3.29 k ~ a l . / m o l e ) . ~ Harvey ~~?~ and c o - ~ v o r k e r sreport ~ ~ ~ identical values, within the limits of experimental error, for the heats of activation in the two decomposition regions of potassium perchlorate. This would indicate the occurrence of essentially the same reaction within the two phases and hence little influrnce of the state of aggregation of the perchlorate salt on the strength of the C1-0 bond. Such a circunistance appears reasonable because of the crystalline transition, the effect of which might be a significant contribution to the collapse of the crystalline lattice and an enhancement of the mobility of the lattice units. The kinetic law prevailing during the induction periods in the decomposition of lithium perchlorate could not, be resolved with any reasonable degree of certainty. However, an energy of activation for this process was detrrmnied by a plot of the logarithm of the time to 0.570 decomposition against 1 / T ; a value of 58.3 f 4.3 kcal./mole way obtained. It seems plausible that this energy of activation might apply to the unimolecular decomposition of lithium perchlorate and is thus a measure of the strength of the C1-0 bond in the coml;ound n-hile in the molten state. Once sufficient lithium chloride is formed ill thc nlelt, thc autocatalytic route to decomposition EA^^. = 52.2 kcal.,/mole) is obviously mow favored on energetic grounds. Acknowledgment.-Thanks are extcnded to I’rofessors H. A. Taylor of New York University and to M. Barth of LaSalle College for e number of helpful discussions during the perforiiiniice of this work. (24) F. 11. Rossini D. n.Wagman, JV. II. 121ans. S T,eiinp xnrl J. Jaffe, “Selected Values of Chemical ThermodT iiatn,c I’ropertiPs,” National Bureau of Standards Circular 500, U. S. CorernInent Printing Office, Washington, D. C., 1952. (25) D. Vorlaender a n d E. Kaasclit, Ber., 66, 1157 (19’23)
