NOTES
Dec., 1963
plex ions are compared with that of potassium thiocyanate; l mM of these complex ions gave an effect on the drop-time comparable to that given by 1 M potassium thiocyanate (Fig. 1). Coordination of NCS- ion to Cr(II1) seems to increase the affinity of the sulfur atom in NCS- ion to mercury. The electrocapillary curves of [Cr(nTCS)(NHJb]2+, C ~ ~ - [ C ~ ( N C S ) ~ ( N Hand + ] + ,trans-[Cr(NCS)z(en)z]+ ions clearly indicate that these cations are strongly adsorbed in the positive branch of the electrocapillary curve. In this respect, they exhibit the character of capillary-active anions. The adsorbability of these complex ions seems to depend on the local charge distribution rather than on the total charge of the complex ion. Figure 1 also shows that trans- [Cr(NCS)*(en)?]+ ion has much higher adsorbability than trans- [Cr(NCS)z(pn)z]f ion. This difference may be due to a kind of steric effect. The current-potential curves of these complex ions also were obtained, which indicated clearly that anomalies on the electrocapillary curves are closely related to the reduction or oxidation of these complexes. The current-potential curve of trr~ns-[Cr(NCS)~(en)~]+ ion showed a sudden increase in reduction current a t -0.92 volt vs. s.c.e., where the electrocapillary curve shows an anomaly as seen in Fig. 1. On the current-potential curve of [Cr(NCS)6l3- ion a small anodic wave appeared at -0.15 volt us. s.c.e.,6where an anomalous change is observed also on the electrocapillary curve in Fig. 1. These relations suggest, in turn, that the adsorption of reacting species plays an important role in the electrode reactions. The present results provide not only interesting examples of electrocapillary phenomena of inorganic ions, but also some important experimental data pertinent, to the discussion of the structure of complex ions in solution and their electrode reactions. Acknowledgment.-The authors thank the Japan Society for the Promotion of Science for the financial support granted for this research.
2707
However, the reaction under study may be controlled by vapor phase diffusion. Naphthalene and picric acid were purified by sublimation in vacuo and fractional crystallization with ethanol, respectively. The melting points were 80.3 and 121.4O, respectively. The reaction in the solid state was studied in the following way. Experimental Clean thick-glass capillaries ( 5 in. long) with uniform diameters (internal 3 mm. and external 9 mm.) were taken. Finely ground naphthalene was introduced in the capillary with a clean metal rod from one side while picric acid waa introduced from the other side in a similar manner and the position of the interface wa.a noted. The capillaries were sealed from either side with a paste which hardened after some time. These capillaries were kept in air-thermostats maintained at suitable temperatures but below the eutectic temperature. The temperature fluctuations were of the order of &lo. A scale was attached to these capillaries. The start of the reaction was indicated by a change in color a t the naphthalene-picric acid boundary. The distance through which the boundary moved was noted a t different time intervals. The induction period at room temperature was found to be only 7 to 10 min. It also was observed that picric acid did not diffuse through product layer whereas naphthalene did. This is schematically shown below.
1
I 1
Naphthalene AB Picric acid - A B
1
t
In this case it appears that the phase-boundary processes are so rapid that equilibrium is established at the boundaries during the entire course of reaction. The diffusion in the product layer is alone rate determining as happens in the case of tarnishing reactions so that if 5 is the thickness of the diffusion layer2 kt (1) where k is a certain constant and t is the time. Modification has been made in the above relationship to account for the heating at the interface due to poor thermal conductivities of the solids. For such a case it has been shown that (2
=
t2 = 2k,t exp(-Pi)
(2)
where
KINETICS OF i m i c r m ~ BETWEEK NAP€ITHSI,Eh’E AND PICRIC ACID IN THE SOLID STATE BYR. P. RASTOGI,* I’ARMJIT S.BASSI,*A N D S.L. CHADHA Chemzstry Department. Panjab lln?zcrs?t$i,Chandauarh. I n d i a
Racezved lllareh ID, 1962
Reactions in the solid state are a class by thcmselves. Because of difficulties in analysis of the composition of the solid phase, studies have been confined mainly to those cases where the course of reaction could be followed by X-ray crystallographic methods or by measuring the amount of gas evolved in suitable reactions.‘ The present note describes a new technique for studying the kinetics of reaction between naphthalene and picric acid in the solid state by following the movement of the colored interface which apparently gives worthwhile results. * Chemistry Department, Gorakhpui University, Gorakhpur, India. ( 1 ) S. W. Benson, “The Foundations of Chemical Kinetics,” XIcGraw-Hill Gook Co., Inc , New York, N. Y..1960,p. 616.
l;, C
= = =
E P
T,,,
= =
C exp(-E/RT,,,) certain constant energy of activation proportionality constant maximum temp. attained instantaneously in the mixture
In this derivation it is assumed that Ti - T = k’t where k’ is another constant so that P = k’E/ RTiT; Ti is the initial temperature and T is any temperature intermediate between T , and Tma,. Equation 1 did not satisfy the data. For testing eq. 2 log [/t was plotted against $. Straight lines were obtained a t all temperatures of observations, justifying the validity of eq. 2. This is shown in Fig. 1. It is interesting to note that all the curves have approximately the same slope, indicating thereby that P has the same value, which should be the case. Further, with increase in the value of T i , k i also (2) G. Cohn, Chem Reu , 42, 527 (1948).
2708
SOTICS
Irol. 66
I
ing X=0l6 absorptions. The ultraviolet absorption intensity of a number of 018-labeled X=O compounds also appears to be measurably different from the parallel normal compounds.ld92 Since the absorption intensity of molecules, both in the infrared and in the ultraviolet region of the spectrum, is a function of their respective transition dipole moment, this (unexpected)lb change in the absorption intensity of X=O1s compounds made it interesting to measure their permanent dipole moment in comparison with that of the corresponding normal substances. Theoretically, however, 3.0 dipole moment of two isotopic modifications \ a the of the same compound should differ only very 2.0 slightly, mainly as the result of the change in the , ', , zero-point energy of the molecules. The calculated -0.2 -0.1 0 0.1 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 dipole moment of HD is thus only3 5.67 X loF4D. log ( C Z / t ) , mm.2/hr. (as compared with zero for Hz). Fig. 1.-Estimation of ki and p . No dipole moment measurements of Ols-compounds seem to have been reported yet, although would increase. This is confirmed by the fact that some comparative values for various pairs of the intercepts of the straight lines increase with normal and deuterated moledes were reported4 rise in temperature. The values of k i and P were and the effect of Cl13 and S34on the moment was determined by the method of least squares. These estimated5in one case. It thus was observed that are given in Table I. even for ammonia-& the increase in the electric moment is only small4 (0.03 D.), although the relaTABLE I tive difference in all three isotopic masses is PARAMETERS OF EQUATION 2 here very high and in spite of the fact that X-H Temp., O C . ki (mm.'/hr.) P (Dun.-*) bonds show the largest anharmonicities'j (which 26 f 1 1.08 0.206 are the main cause for the observed changes in the 31 i 1 1.47 .205 dipole moment of labeled compound^^^^^). 37 f 1 4.225 ,218 The electric moment of the following pairs of 45 f 1 4.68 ,223 X=O compounds therefore has been measured 54 f 1 5.77 ,234 accurately: (a) normal and 90 atom yo OL8-benzo59 & 1 10.8 .225 phenone, (b) normal and 54 atom % OWriphenylIf T i / T m a ,
