J . Phys. Chem. 1985,89, 1970-1973
1970
-
at elevated temperatures where the net reaction Pb2+
+ 2H20
Pb4+
+ 20H- + H,(g)
might have occurred. The electrochemical potential for the formation of Pb4+in basic solution is ca. -0.28 V, and the reaction might reasonably have had a positive potential within the zeolite a t dehydration temperatures. Because the product crystal was colorless, the disproportionation reaction 2Pb2+
-
PbO
+ Pb4+
which would have generated metal atoms which would have been expected to gather on the crystal's surface, could not have occurred. Therefore not more than nine lead ions entered per unit cell during ion exchange. The geometry of Pb2+in Pb[OP(C6H5)20]2is pseudotrigonal bipyramidal with an equatorial lone pair and axial and equatorial Pb-0 distances of 2.40, 2.44 and 2.23,2.23 A, respectively!2 The Pb2+ ion in bis(nicotinat0-N,O)lead(II) has a distorted cubic coordination geometry from two nitrogen atoms (Pb-N, 2.69 A) and six oxygen atoms (Pb-0, 2.39, 2.78, 2.90 A).43 The Pb-0 (42) Hursthouse, M. B. Mol. Srrucr. Diffr. Methods 1977, 5 , 399 (only that page). (43) Hursthouse, M. B. Mol. Srruct. Dqfr. Methods 1976, 4 , 389-392.
distances in PbTi307,where Pb2+is 7-coordinate, range from 2.36 to 3.04 In PbV206, the lead ions are 9-coordinate, with Pb-0 distances of 2.56 to 2.90 A.43 6Pb0*5B20,contains three crystallographically independent Pb2+ ions, all coordinated by three or four oxygens at distances of 2.23-2.55 A.44 In PbOHI, the Pb-OH distance is 2.60 A.45 These representative examples show that Pb2+ has a wide range of coordination geometries and bonding distances to oxygen. These distances span most of those encountered in Pb2+-exchanged zeolite A (see Table 11). Although a somewhat greater uptake of Pb2+ from aqueous nitrate than from aqueous acetate was observed in the batchmethod equilibrium experiments of Hertzenberg and Sherry, this work shows that the capacity of zeolite A for Pb" is much greater from aqueous acetate solution when only Pb2+-containingcations are present, ignoring H+, in the exchange solution (flowmethod).
Supplementary Material Available: Observed and calculated structure factors with esd's for dehydrated Pb6-A and partially dehydrated Pb9(OH)8(H20)3-A,supplementary Tables 1 and 2 (8 pages). Ordering information is available on any current masthead page. (44) Krogh-Moe, J.; Wold-Hansen,P. S.Acta Crystallogr.,Sect. B 1973, 29, 2242-2246.
(45) NSisiinen, R. Suom. Kemistil. B 1966, 39(4), 105-108.
Intercalation Kinetics Study of Alkali-Metal Ions into TIS, Uslng the Pressure-Jump Technique Minoru Sasaki, Hiroshi Negishi, Hisao Ohuchi, Masasi Inoue, and Tatsuya Yasunaga* Department of Chemistry, Faculty of Science, Hiroshima University, Hiroshima 730, Japan (Received: August 1 , 1984; In Final Form: December 1 1 , 1984)
In aqueous suspensions of alkali-metal intercalation compounds of Tis2, double relaxation were observed by using the pressure-jump technique with conductivity detection. For all intercalation compounds, the fast relaxation times decrease with particle concentration, while the slow ones are approximately constant. From the kinetic results obtained, the fast and slow relaxations are attributed to association-dissociation of the alkali-metal ions on the surface of the intercalation compounds and intercalation-deintercalation of their ions in the interlayers of the intercalation compounds, respectively. It was found that the order of association of rate constants of alkali-metal ions corresponds to that of the mean time of movement of a water molecule between their ions and that the deintercalation rate constant of a lithium ion in the interlayer is 1 order of magnitude faster than those of other alkali-metal ions.
Introduction Transition-metal dichalcogenides are interesting compounds with layered structures. The ionic and electronic properties of the compounds can be modified Bignificantly by the intercalation of various electronic donors such as Lewis bases or alkali metals which are accommodated into the interlayers.'" Lithium intercalation compounds of the transition-metal dichalcogenides have recently attracted considerable attention because of their potential use as a cathode material in b a t t i e r i e ~ . ~For , ~ cathode materials. ( 1 ) Whittingham, M. S. Prog. Solid Stare Chem. 1978, 12, 41. (2) Lerf, A.; SchBllhom, R. Inorg. Chem. 1977, 16, 2950. (3) Thompson, A. H. Physicu B+C (Amsterdum) 1980, 99B+C, 100. (4) Hibma, T. Physicu B + C (Amsterdam) 1980, 99B+C, 136. (5) Whittingham, M. S.; Jacobson, A. J. "IntercalationChemistry";Ac-
ademic Press: New York, 1982. Jacobson, A. J. Ibid. Chapter 7. Hibma, T. Ibid. Chapter 9. SchBllhorn, R. Ibid. Chapter 10. (6) Negishi, H.; Sasaki, M.; Yasunaga, T.; Inoue, M. J . Phys. Chem. 1984, 88, 1455.
0022-3654/85/2089-1970$01.50/0
high planar mobility of the donors intercalated between interlayers is k t m a r y . Kinetic investigation of the intercalation process may give important information on ionic mobility in addition to the details of the intercalation kinetics. In the previous work,6 we have successfully carried out kinetic measurements of sodium intercalation into the interlayers of Tis2 using the pressure-jump technique. Kinetic information on the intercalation-deintercalation process of Na+ in the interlayers was obtained. The purpose of the present investigation is to elucidate the mechanism of intercalation-deintercalation of alkali-metal ions (Li', K+, Rb', and Cs') in the interlayers of Tis, and to clarify the relationship between the kinetic properties of alkali-metal ions in the interlayers and their ionic properties. Experimental Section
Preparation of titanium disulfide (Tis,) crystals used and the structural data obtained by X-ray diffraction have been reported previously.6 The alkali-metal intercalation compounds of Tis2 0 1985 American Chemical Society
The Journal of Physical Chemistry, Vol. 89, No. 10, 1985 1971
Alkali-Metal Intercalation Compounds of Tis,
5 c
u)
0
c
8 9 2.5 P-
r
4
2
0
c,
'0
, 10 g dm-3
F w 1. Particle concentration dependences of the fast relaxation times
Results and Discussion Kinetic measurements were performed in the aqueous suspensions of TiS2(Li), TiS2(K), TiS2(Rb), and TiSz(Cs) using the pressure-jump technique with conductivity detection, and double relaxations were observed, as found previously for an aqueous suspension of TiS2(Na).6 The amplitude of the double relaxations increases with the particle concentration, C,, but is drastically reduced by the addition of dilute solutions of alkali-metal salts. Preliminary experiments revealed that the alkali-metal ions are released from the interlayers of TiS2(M) as a result of deintercalation. This fact suggests that the relaxations observed may be related to the intercalation-deintercalation of the alkali-metal ions in the interlayers of TiS2(M). The particle concentration, C ,dependences of the fast and slow reciprocal relaxation times, 7; P and 7s1,are shown in Figures 1 and 2, respectively. For all alkali-metal intercalates, the values of T < ~increase with C,,while those of 7s-l show slight increases and then tend to approach constant values. The magnitudes of 7r1are in the decreasing order of Li+, K+, Rb+, and Cs', respectively, which corresponds to the order of ionic radius of the alkali-metal ions. The magnitudes of T F ~however, , are approximately independent of the alkali-metal ionns except for Li+. These trends are quite different from those obtained in aqueous sus-
4
C, , l o g dm-) Figure 2. Particle concentration dependences of the slow relaxation times in the aqueous suspension of TiS2(M)at 25 OC.
in the aqueous suspensions of TiS2(M)at 25 OC.
formulated as M,(H20),(OH) T i l . o ~ ~ - ~ (M: y O ~Li',y K+, Rb', and Cs+), hereafter called Ti&(M), were prepared by the same procedures as that reported previously for TiS2(Na).6 The values of x and y in these compounds were determined to be 0.22 and 0.06 for TiSz(Li), 0.23 and 0.05 for TiS2(K), 0.25 and 0.04 for TiS,(Rb), and 0.20 and 0.07 for TiS2(Cs), respectively, from base titration and the saturated amounts of alkali-metal ions released. The sizes of the particles of TiS2(M) were less than 1 pm. In an aqueous suspension of the TiS2(M), small amounts of alkali-metal ions were released from the TiS2(M) crystals to the bulk solution. The bulk concentrations of the alkali-metal ions in the supernatant solutions of the TiS2(M) suspensions were measured by an isotachophoresis apparatus (Shimazu Ind. Corp., IP-2A Type),' where two kinds of reference solutions are used; the leading and terminal solutions contain the cations having high and low mobilities compared with the sample cation, respectively. The leading solutions were 2 X loV3N HCl for K', Rb', and C d and lo-, mol dm-3 CH3COOK (pH 4.8) for Li+, where the pH was adjusted with CH3COOH. The terminal solutions were 2 X mol dm-' U 0 2 ( C H 3 C 0 0 ) 2containing 2 X N HC1 for the former case and caminocaproic acid (pH 4.8) for the latter case. The pressurejump and ultramicroelectrophoresis apparatus used are the same as those reported previously$ and the experiments were carried out anaerobically.
2
0
3 K
2
0
4
Cp , 10 g dm-3 Figure 3. Dependences of the amounts of alkali-metal ions released on the particle concentrations at 25 OC.
pensions of alkali-metal intercalates of a-zirconium phosphate.s As described above, the alkali-metal ions are released from TiS2(M), and thus the amounts of the alkali-metal ions released, [M+],, are measured, and the results are shown in Figure 3. [M+], increases nonlinearly with C, for each ion, and there is no trend among these dependences. In the present constant pH condition (pH ~ 9 . 9 the , values of [M+], are much larger than that of the concentration of OH-, and thus cointercalation of M+ The values of the electrostatic poand OH- can be negle~ted.~ tential \k@defined by Davis et ale9which are calculated from (-potential measured are also independent of C,,; $8 = -30.9, -34.5, -31.8, and -30.0 mV for TiS,(Li), TiS,(K), TiS,(Rb), and TiS2(Cs), respectively. This result indicates that the adsorption and desorption processes between the surface of TiS2(M) and bulk phase are governed by the constant electrostatic potential on the surface of TiS2(M). According to the mechanism of intercalation-deintercalation of Na+ in the interlayers of TiS2(Na) proposed in the previous work$ the mechanism for TiS2(M) is described similarly as TiS2(M)i
k1
k-I
TiSz(M)s
k2 -q-$ TiS2-
(I)
M' KI
KZ
where the subscripts i and s denote the interlayer and the surface, (8) Mikami, N.; Sasaki, M.; Yasunaga, T.; Hayes, K.F.J . Phys. Chem. 1984,88, 3229.
(9) Davis, J. A.; James, R. 0.; Leckie, J. 0.J . Colloid Interface Scf.1979, (7) Haruki, T.; Akiyama, J. Anal. Lett. 1973, 6,985.
63, 480.
1972 The Journal of Physical Chemistry, Vol. 89, No. 10, 1985
Sasaki et al.
TABLE I: Kinetic Parameters of Association-Dissociation of the Alkali-Metal Ions on the Surface and Intercalation-Deintercalation of Their Ions in the Interlayers of TiS2(M) at 25 O C
10-4kp,
d, A 8.60 8.66 8.83 8.92 9.14
compd
Tis2(Li) TiS2(Na)" TiS,(K) TiS,(Rb) TiSz(Cs)
k,, s-l 11 f 4 2.8 f 0.6 3.2 f 0.5 4.8 f 0.8 5.5 f 0.5
lO-Ik-,, s-I 1.3 f 0.4 1.3 f 0.4 1.0 f 0.4 0.9 f 0.4 1.9 f 0.4
mol-' dm3s-l
10-1k7 in1. s-l 12f2 3.2 f 0.9 1.7 h 0.3 2.4 f 0.3 0.8 f 0.2
3.1 f 0.4 1.3 f 0.1 0.82 f 0.04 0.59 f 0.03 0.67 f 0.07
Reference 6.
. . 2
in
c1 -
i."
1
0 2
0
21M'I vs. the concentration terms in eq 3. I
Figure 4. Plots of
f["S
0.4
0.6
08
4 mol dm"
Figure 5. Plots of
7;I's
vs. the concentration terms in eq 4.
The plots of 7c1and 7g-l vs. the concentration terms in eq 3 and 4 are shown in Figures 4 and 5, respectively. As can be seen from these figures, the experimental plots fell on the straight lines for all systems. These linearitis lead to the conclusion that the double relaxations observed can be attributed to mechanism I. The four
kinds of rate constants were evaluated from the slopes and the intercepts of the straight lines and are listed in Table I, where the intrinsic values were calculated by using the values of qo. In order to obtain the relationship between the layered structure of TiS2(M) and the kinetic parameters obtained, the values of interlayer distance, d, of TiS,(M) crystals were measured by X-ray diffraction, and the results are listed in Table I, showing the monolayer hydrates of the alkali-metal intercalation compounds. The results for TiS2(Li) and TiS2(Na) are different from those reported by Lerf and Schollhorn.2 Their hydrated intercalation compounds were prepared by alkali-metal treatment and then by hydration, while the present intercalation compounds were prepared by base treatment. It is known that the formation of monolayer or bilayer hydration is dependent on the interaction energy between the charged alkali-metal ion and polar water molecule. In the present alkali-metal intercalation compounds, OH- also cointercalates into the interlayers together with M+. By N M R study, it has been reported that the cointercalated M+ and OH- interact with each other. The effecive charge densities of the alkali-metal ions existing in the interlayers of TiS2(M), therefore, are reduced by the neutralization of M+ with OH-. Consequently, the formation of monolayer hydrates for TiS2(Li) and TiS2(Na) in the present systems may result from the reduction of the charge density of the intercalated alkali-metal ions. Except for TiS2(Li), the values of kl increase with d, while the value of kl for TiS,(Li) is 1 order of magnitude larger than those for other layered intercalation compounds. The former linear dependence is mainly due to two-dimensional steric hindrance. However, the latter difference could not explain such an effect. The above result indicates that the mobility of intercalated Li+ is exceptionally large compared with that of the other alkali-metal In the present stage, ions, as suggested by Berthier et al." however, the reasonable explanation on the abnormal dynamic property of Li+ is not available because of a lack of further detailed information. On the other hand, the values of k-, are independent of d . If the deintercalation process is governed mainly by the oxidation of the intercalated alkali-metal ions, the deintercalation rate constant is independent of d because the oxidation potentials
(10) Sasaki, M.; Negishi, H.; Inoue, M.; Yasunaga, T. J . Phys. Chem. 1984, 88. 3082.
1980, 99B+C, 107.
respectively, Tis2- denotes the vacant site, and the other symbols stand for their customary meanings. The equilibrium constants, K 1 and K2, are represented by6
= Kzintexp(
$)
where the superscript int denotes intrinsic and the other symbols have their customary meanings. When the first step is faster than the second step, only a single relaxation for the second step should be observed.'O If the first step in mechanism I is slower than the second step, the fast and slow relaxation times are derived as 7;'
= k2
+ k-2([TiS2-] + [M'])
= k2
+ 2k-,[M+]
(3) (4)
with
$:;(-
k2 = k p exp
k-2 = k-Pt exp(
-2kBT 5)
(1 1 ) Berthier, C.; Chabre, Y.;Segransan, P. Physica B+C (Amsrerdam)
J. Phys. Chem. 1985,89, 1973-1976
Y
0
1
2
1
3
tilt
Figure 6. Relationship between log
and f i / r .
of the alkali-metal ions in bulk solution are in the narrow range of -3.0 to -2.7. The experimental fact is consistent qualitatively with this prediction. Both values of k? and k-? do not depend linearly on the diffusion coefficients of the alkali-metal ions in bulk solution, and are relatively small compared with those governed by the diffusion
1973
process. These facts suggest that the adsorption and desorption processes of the alkali-metal ions between the surface of TiS,(M) and bulk phase are governed not by the diffusion but by an activation energy. As can be seen from Figure 6, the values of k-? obtained are proportional to the mean time, ti, during the movement of a water molecule from the vicinity of an ion to the immediate neighborhood of its nearest ion, where ti is normalized with the mean time, t, of the self-diffusion of a water molecule in water as tilt.', This fact suggests that the adsorption process of the alkali-metal ions on the surface of TiS2(M) is mainly governed by the release rate constants of a water molecule from these ions. The present study reveals that Li+ has an extremely high mobility in the interlayers of TiS2(M) and the lithium intercalation compound of Tis2 will be expected to be used as a solid electrolyte or cathode materials. Registry No. Tis2, 12039-13-3; Li, 7439-93-2; K, 7440-09-7; Rb, 7440-17-7; CS,7440-46-2. (12) Horne, R. A. "Water and Aqueous Solutions: Structure, Thermodynamics, and Transport Processes'; Wiley-Interscience: New York, 1972; Chapter 14, p 598.
Reinvestigation of Triplet-Sensitized Cis-Trans Photoisomerization of Cyclooctene. Alkene-Concentration and Sensitizer-€, Dependence of Photostationary Trans/Cis Ratio' Yoshihisa Inoue,* Tomokazu Kobata, and Tadao Hakushi Department of Applied Chemistry, Himeji Institute of Technology, 2167 Shosha, Himeji, Hyogo 671-22, Japan (Received: August 13, 1984)
Reinvestigation of the photochemical cis-trans isomerization of cyclooctene effected by aromatic sensitizers revealed that the photostationary translcis ratio, (t/c),, is evidently dependent both upon the concentration of cyclooctene and, contrary to the previous work, upon the triplet energy of the sensitizer employed; the (tlc), ratio increases with decreasing alkene concentration in the high concentration range and with increasing sensitizer E T in the range 80-84 kcallmol. Interestingly, the (tlc), ratios obtained in the liquid phase are consistent with those in the vapor phase. The alkene-concentration dependence is attributed to the involvement of the excited singlet state of aromatics affording cross-adducts particularly at the higher alkene concentrations where the short-lived aromatic singlet can be quenched predominantly by the strained trans isomer. Stern-Volmer studies showed that the excitation ratio, i.e., the relative rate of quenching by the cis and trans isomer, increases with increasing sensitizer ETand is responsible for the ETdependent (t/c), ratio, while the decay ratio, from the perpendicular triplet cycloocteneto the cis and trans isomer, is invariant over the E T range. This is due to the endothermicity of the triplet-energy transfer process and the considerable difference in ET of cis- and trans-cyclooctene.
Introduction Previously' we have reported that the direct and triplet-sensitized cis-trans photoisomerizations of cyclooctene give distinctly different photostationary trans/& ratios, (tlc),, which have been explained in terms of the spin multiplicity of the excited state involved and the deformed potential curves for the excited and ground electronic states due to the considerable strain in the trans isomer (Scheme I). Also reported was the fact that the solution-phase triplet photosensitizations, using sensitizers whose triplet energies, ET, are greater than 80 kcallmol, give low, but invariant, (tlc), ratios around 0.05 over the ET range 80-84 kcal/mol, whereas the vapor-phase photosensitizations with the same sensitizers result in &dependent (tlc), ratios ranging from 0.09 to 0 . 2 . The low (tlc), values in the solution-phase photosensitization have been
ascribed not to the preferential excitation of the trans but to the predominant decay from the perpendicular triplet of cyclooctene to the cis, since at that time the triplet energies of simple alkenes were believed to be well below 8 2 kcal/mol,2 and therefore the energy transfer to each isomer was considered to be diffusion controlled. In this context it is not so curious to have been concluded that the ET-dependent (tlc), ratios in the vapor phase originate from the vibrational activation of the triplet state of cyclooctene.
(1) For the previous paper on this subject, see: Inoue, Y.; Takamuku, S.; Sakurai, H. J . Phys. Chem. 1977, 81, 7.
(2) Reid, C. J . Chem. Phys. 1950,18, 1299. Evans, D. F. J . Chem. SOC. 1960, 1735. McDiarmid, R. J . Chem. Phys. 1971, 55, 4669.
0022-3654/85/2089-1973$01.50/0
SCHEME I
0 1985 American Chemical Society
