Langmuir 1988,4,433-438
433
Infrared Spectra of Methane Adsorbed by Ion-Exchanged ZSM-5 Zeolites Tatsuya Yamazaki, Isao Watanuki, Sentaro Ozawa, and Yoshisada Ogino* Department of
Chemical Engineering, Faculty of Engineering, Tohoku University, Aramaki-Aoba, Sendai, 980 Japan
Received J u l y 17, 1987.
I n Final Form: September 30, 1987
Infrared spectra of methane adsorbed by five kinds of ion-exchanged zeolites (HZSM-5, LiZSM-5, NaZSM-5, KZSM-5, RbZSM-5, and CsZSM-5) have been measured, and the results have been compared with adsorption characteristics obtained from the adsorption isotherms measured separately. The shifts of the IR u1 peak position (Aul), the IR absorption coefficient (A), and also the isosteric heat of adsorption (q,?) have been found to be intimately related to the ionic radius of the cation exchanged. Furthermore, the electric field ( E )of the cationic site has been found to decrease according to the sequence Li+ > Na+ > K+ > Rb+ > Cs+. From these results, it has been postulated that the interaction between the cationic site and a methane molecule plays a dominarit role in the adsorption. In addition, the IR and adsorption data have revealed that there are two kinds of adsorption sites over the zeolite surface, i.e., the silicalite-like site (site 1)and the cationic site (site 2). The difference between the adsorption energy for site 1and for site 2 has been estimated to be qs2- 20.92 kJ/mol, which has been almost identical with the chemical potential difference evaluated from IR data.
Introduction
Since the publication' of the characteristic catalytic activity of the ZSM-5 zeolite for the methanol conversion, extensive studies on the catalysis over various ZSM-5-type zeolites have been carried out around the world.2 On the other hand, only a limited number of studies on the adsorptive properties of the ZSM-5 zeolite have been published."' Considering this situation, we have attempted to clarify the adsorptive characteristics of this important adsorbent in more detail. Thus, in an earlier report: we have presented information about methane adsorbed on NaZSM-5 zeolites with different Si02/A1,03 ratios. The effects of ion exchange upon the adsorptive properties of the ZSM-5 zeolite have been studied in the present work. For this purpose, infrared spectroscopy has been adopted as the methodology. According to Cohen de Lara et al,+l4 infrared spectroscopy is a powerful tool in studying the nature of the cationic adsorption site in the NaA-type zeolite, and Papp et aL4have applied this method to study the behavior of methane adsorbed by an HZSM-5 zeolite. Thus we expected that valuable information about ad(1) Argauer, R. J.; Landolt, G. R. U.S.Patent 3702886, 1972. (2) For instance: Proc. 8-th Intern. Congr. Catal.;July 24,1984, West Berlin; Veralg Chemie: Weinheim, 1984. (3) Flanigen, E. M.; Bennett, J. M.; Grow, R. W.; Cohen, J. P.;Patton, R. L.; Kirchner, R. M. Nature (London) 1978,271, 512. (4) Papp, H.; Hinsen, W.; Do, N. T.; Baerns, M. Thermochim. Acta 1984, 82, 137. (5) Caro, J.; Bulow, M.; Schirmer, W.; Krager, J.; Henke, W.; Pfeifer, H.; Zdanov, S. P. J. Chem. SOC.,Faraday Trans. 1 1985,81, 2541. (6) Wu, P.; Ma, Y. H. Proc. Sixth Intern. Zeolite Confer.;July 10-15, 1983,Reno, NE; Olson, D., Bisio, A., Eds.; Butterworths: Guilford, 1984; p 251.
(7) Richards, R. E.; Rees, L. V. C. Langmuir 1987,3, 335. (8)Yamazaki, T.; Watanuki, I.; Ozawa, S.; Ogino, Y. Nippon Kagaku Kaishi 1987, 1535. (9) Cohen de Lara, E.; Delaval, Y. J. Chem. Phys. 1974, 78, 2180. (10) Cohen de Lara, E.; Delaval, Y.; Tsakiris,J. J.Chim. Physiq. 1976, 73, 387. (11) Cohen de Lara, E.; Kahn, R. J.Physiq. (Les Ulis, Fr.) 1981,42,
1029. (12) Cohen de Lara, E.; Seloudoux, R. J.Chem. SOC.,Faraday Trans.
I 1983, 79, 2271. (13) Cohen de Lara, E.; Kahn, R. Physiq. Lett. 1984, 45, L-255. (14) Cohen de Lara, E.; Kahn, R.; Seloudoux,R. J. Chem. Phys. 1985, 83, 2646.
0743-7463f 88f 2404-0433$01.50f 0
Table I. ZSM-5Zeolites Used in Adsorption and IR Measurements specific content of surface cationic sites, area, Si02/A1203 zeolite mo1.g-' m2.g-l mole ratio HZSM-5 1-32x 10-3 442 23.3 LiZSM-5 0.94 x 10-3" 439 23.3 1.28 X 414 23.3 NaZSM-5b KZSM-5 1.25 x 10-3 390 23.3 RbZSM-5 1.18 X 359 23.3 CsZSM-5 1.12 x 10-3 310 23.3 337 2000 silicalitec 0.02 x 10-3 "After correction for C02 contamination; even in a fresh state, this sample exhibited IR bands assigned to surface COB species (Bertsch, L.; Habgood, H. W. J. Phys. Chem. 1963, 67, 1621) occupying part of cationic sites. Thus the amount of C02 species was determined by thermogravimetry and mass spectrometry. The original zeolite was kindly presented by Mr. K. Igawa, Chemical Research Lab, Toyo Soda Manufacturing Co. Ltd. CReference8. This zeolite was kindly presented by Prof. T. Yashima, Tokyo Institute of Technology. sorptive properties of ion-exchanged ZSM-5 zeolites (HZSM-5, LiZSM-5, NaZSM-5, KZSM-5, RbZSM-5, and CsZSM-5) would also be obtained by the methodology mentioned above.
Experimental Section Materials. The main properties of the adsorbents used in the present study are shown in Table I. The source of the original zeolite (NaZSM-5)is shown in the footnote to the table. Other adsorbents were prepared by a conventional ion-exchangemethod using a predetermined alkali chloride solution as a cation source. Methane (99.95%)was commercially obtained and used as an adsorbate gas without further purification. Helium was used in determining the dead space volume of the adsorption vessel, and nitrogen was used in the measurement of the surface area 2 of the adsorbent. Both of these gases were purified by passing through a cold trap (-196 O C ) . Apparatus and Procedures. A conventional glass adsorption apparatus (a constant-pressuretype) was used in the measurement of the adsorption isotherms for methane. The adsorbent zeolite (about 2 g) was placed in the adsorption vessel and degassed at 300 "C for a 12-h period at 1.3 X Pa. The adsorption isotherms were measured under the following conditions: temperature, -80 to +30 O C ; pressure 6.7 X 102 to 6.7 X 104 Pa. The same 0 1988 American Chemical Society
Yamazaki et al.
434 Langmuir, Vol. 4, No. 2, 1988 *
f ? I+l
4
7 hv
8
9
Figure 1. Schematic drawing of the IR cell: 1, thermocouple; 2, inlet of cold nitrogen; 3, outlet of cold nitrogen; 4, microheater; 5, cover (stainless steel); 6, sample holder (copper);7 , cell body stainless steel); 8, NaCl window; 9, O-ring; 10, coolant room; 11, sample room. apparatus was used to measure the adsorption isotherm of nitrogen at -196 "C. The resulting isotherm of nitrogen was used to evaluate the specific surface area of the adsorbent; the BET theory with a finite adsorption layer (- 1.3 layers) was applied. Illustrated in Figure 1 is the cell used for the infrared (IR) spectroscopic study of methane adsorbed on the zeolite surface. An integral part of this cell is the sample holder, which is made from copper and is capable of keeping the sample at a desired temperature between -80 and +300 "C; a microheater served for heating the sample. When a low temperature of the sample was desired, a cold nitrogen stream was passed through a coolant space of the sample holder. By controlling the flow rate of the cold nitrogen, we obtained the desired low temperature. The IR spectra of the adsorbed methane were recorded on a FT-IR spectrometer (JEOL-JIR-100)with a TGS detector using a disk method. The sample disk (about 100 wm thick X 13 mm in diameter; about 15 mg) was prepared by pressing the powder of the sample zeolite at 2.2 X los Pa for 10 min. After being weighed, the sample zeolite was placed on the sample holder and degassed at 300 "C for a 12-h period at 1.3 X Pa. Two different control measurements were carried out in order to obtain reliable information about the IR absorption characteristics of the adsorbed methane. The first controlling measurement was made to correct for the IR absorption of methane in the gaseous phase. A KBr disk approximately the same size as the zeolite disk was placed on the sample holder, and IR spectra were recorded under various temperatures and pressures of methane. The spectral data thus obtained were put into the memory of the data processing computer of the spectrometerand used to correct the IR data for the sample zeolite. The second controlling measurement was made to obtain an exact temperature of the sample disk. Since the temperature measurement was made by a copper-constantan thermocouple spot-welded to the sample holder, the sample temperature was considered to be different from the thermocouple reading, in particular under the IR irradiation. Thus a separate thermocouple was attached to the sample disk, and the difference between the temperature of the sample holder and that of the sample under the IR irradiation was measured. Temperature differences thus obtained were utilized to draw a calibration curve necessary for correcting the temperature of the sample holder.
Results and Discussion Adsorption Isotherm, Monolayer Adsorption Capacity, and Heat of Adsorption. Adsorption isotherms obtained in the present study are shown in Figure 2. The adsorption data, in particular those obtained under a high-pressure region, were found to obey a Langmuir-type adsorption equation and enabled us to evaluate an approximate maximum monolayer adsorption capacity V,,, of the adsorbent used. Furthermore, application of the Clausius-Clapeyron equation to the adsorption isotherm enabled us to evaluate the heat of adsorption qst.
The approximate maximum monolayer adsorption capacity V , and the specific surface area Z of the adsorbent zeolite are plotted in Figure 3 as a function of the ionic radiusI5 of the cation exchanged. A parallelism between V , and 2 is evident in the figure. Both V , and Z remained almost constant for H+-, Li+-, and Na+-exchanged ZSM-5 zeolites while they tended to reduce by K+-exchanging and then sharply decreased to much smaller values for Rb+- and Cs+-exchanged zeolites. The experimental results mentioned above strongly suggest that the incorporation of a cation with a small ionic radius such as H+, Li+, or Na+ has brought about little change in the pore structure of the original zeolite. With respect to this, it is interesting to discuss the V , values shown in Table I. The V , values for HZSM-5, LiZSM-5, and NaZSM-5 approximately correspond to 2.3 methane molecules/cationic site; this is equivalent to -18 methane molecules/unit structure of the zeolite. If we adopt the channel model proposed by Richards et al.,' the monolayer adsorption capacity V,,, mentioned above corresponds to -69% occupation of the channel available for the adsorption. This means that V , does not represent the total adsorption capacity (V,) of the adsorbent. It must be pointed out that V , is an approximate monolayer adsorption capacity obtained by applying a Langmuir-type equation to the high-pressure region of the adsorption isotherm, and hence we have to distinguish V , from V,. If we assume that the Vt value for the Richards model is the maximum adsorption with 1.4-1.5 layers, the abovementioned difference between V, and Vt can be reconciled. The 1.4-1.5-layer adsorption also reconciles with the 1.3layer adsorption model adopted in the surface area evaluation. On the other hand, we have to consider that an incorporation of a large cation has resulted in a change in the pore structure. A local plugging and/or local reduction of pore size appears to be only slight for KZSM-5 while that for RbZSM-5 and CsZSM-5 is significant. Although the exact position and structure of the cationic site over the surface of the ZSM-5 pore wall are still unknown, it can be pointed out that the void which is surrounded by five (or six) oxide ions constituting the pore wall cannot accept the whole body of one Rb+ or Cs+ ion. Thus, for instance, Cs+ ions incorporated occupy about 7% of the pore space. This probably results in a significant reduction in the monolayer adsorption capacity because a steric hindrance against methane adsorption is introduced by the Cs+ ions occupying the pore surface; up to a 40% reduction of the adsorption capacity would be brought about depending on the adsorption model. Plotted in Figure 4 are the heats of adsorption qat for methane adsorption on different zeolites. The differences of the qat values among different adsorbents at higher coverages are small and uncertain. However, it appears worth noting that the heats of adsorption at zero coverage, qa:, seem to decrease in the sequence Li+ > Na+ > K+ > Rb+ > Cs+. Although there remains uncertainty due to the experimental error involved in qs:, this sequence appears to suggest that one of the primary factors determining the adsorbate-adsorbent interaction is the polarizing power of the cation. It is well-known16that the polarizing power of a cation is inversely proportional to the ionic radius. As we will show in a later section, this view is supported by the IR spectroscopic data. (15) Pauling, L. The Nature of The Chemical Bond, 3rd ed.; Cornel1 University: Ithaca, 1960; p 518. (16) Cotton, F. A.; Wilkinson, G. Advanced Inorganic Chemistry, 2nd ed.; Interscience: New York, 1966; p 53.
IR Spectra of Methane Adsorbed by ZSM-5
Langmuir, Vol. 4, No. 2, 1988 435
I
I
f
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0
1
2
3
4
5
6
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1
2
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5
6
.c,
c
5
60
60
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40
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a
'
0
1
2
3
4
5
6
'
Pressure I
0
1
2
3
4
5
6
l o 4 Pa
Figure 2. Isotherms for methane adsorption onto various ion-exchanged ZSM-5 zeolites.
2 d 0
ul
I
0
0.5 1.0 Ionic radius (r,)
1.5 I
A
Figure 3. Surface area and monolayer adsorption capacity as a function of the ionic radius of cation exchanged.
It must be pointed out that the qs: value for the HZSM-5 seemingly contradicts the view mentioned above. The heat of adsorption for this zeolite is expected to be
10 20 Amount of adsorption I
30 ml.g.1
Figure 4. Isosteric heats of adsorption as a function of the amount of methane adsorbed 0 , LiZSM-5; A, NaZSM-5; A, KZSM-5; O , RbZSM-5; H, CsZSM-5; 0,HZSM-5.
the largest among those obtained in the present study, if we take the formal ionic radius of H+." Thus the small q,: value for the HZSM-5 appears anomalous. This dis(17)Shannon, R.D.;Prewitt, c. T.Acta Crystallogr., Sect. B: Struct. Crystallogr. Cryst. Chem. 1,969, B25, 925.
Yamazaki et al.
436 Langmuir, Vol. 4, No. 2, 1988
-a--~-=Y--A-A0 0 -
D-D-0-D-
2
0
10
20
Amount of adsorption ? . ;*a.
3200
,
3100 3000 2900 Wavenumber
/
2800 2700 cm-'
Figure 5. Examples of IR spectra of adsorbed methane at -47 OC:
-, CH4-Li,NaZSM-5;- - -, CH4-silicalite.
crepancy, however, disappears if we take the report of Jacobs et al.18 into account. According to this report, the cationic site of the HZSM-5 zeolite has the character of a hydroxyl group. This view undoubtedly explains a large apparent effective ionic radius of the adsorption site and rationalizes the small q a t value. It must be pointed out that factors other than the electric field or the polarizing power of the cation incorporated also affect the magnitude of the heat of adsorption. Namely) contributions from the dispersion force and the average electric fieldlg exerted by Si4+,A13+,and 02-constituting the zeolite structure are included in the qat obtained. It is reasonable to consider that the magnitude of contributions from these factors would be approximated by the heat of adsorption of methane onto silicalite (20.92 kJ/mol).8 Since this value must be common for every zeolite used, the difference q a t - 20.92 (in kJ/mol) would probably be due to the interaction between methane and the cationic site. It appears worth noting that the above-mentioned q s t value for the methane-silicalite system can be reconciled with the data reported by Richards et al.' Namely, they reported that the heat of adsorption of ethane onto silicalite is 30 kJ/mol and q s t increases approximately by 10 kJ/mol/additional CH2 group. According to this view, it is apparent that the q s t for methane is 20 kJ/mol, in good agreement with our value (20.92 kJ/mol). The above discussion implies that there are at least two kinds of adsorption sites over the zeolite surface, Le., cationic sites and silicalite-like sites. This deduction is important, and the two-site model is available for a qualitative explanation of the decrease in qst with the surface coverage (Figure 4). IR Spectra of Adsorbed Methane. Exemplified in Figure 5 are the IR spectra of adsorbed methane. The absorption band a t 2885 cm-l is assigned to the v1 vibration of the methane molecule. It must be pointed out that this vibration is IR inactive under ordinary conditions. Therefore, gaseous methane does not show this absorption band. It is reported by Cohen de Lara et aL9Jothat the v1 absorption develops when adsorbed methane molecules are distorted by a strong electric field of cations in the zeolite pore. Indeed, as we can see in Figure 5 (dotted line), methane adsorbed by silicalite, which has little cationic site, exhibits negligibly small IR absorption for the vl vi(18)Jacobs, P.A.;Ballmoos, R. J. Phys. Chem. 1982, 86, 3050. (19)Barrachin, B.; Cohen de Lara, E. J.Chem. SOC.,Faraday Trans. 2 1986,82, 1953.
30
40
/
ml.g'
Figure 6. Shifts of IR v1 peak position as a function of the amount of methane adsorbed at -47 "C: 0 , LiZSM-5; A, NaZSM-5; A,
KZSM-5; 0,RbZSM-5; W, CSZSM-5;0,HZSM-5.
bration? Neither impurities nor a solid-like adsorbed state of methane can explain the v1 absorption mentioned above. If the band at 2885 cm-l is ascribed to contaminations due to vacuum pump oil or other impurities, the same absorption should be observed for the methane-silicalite system. Since solid-state methane does not exhibit the IR v1 absorption,20the absorption peak observed cannot be ascribed to a solid-like adsorbed state. Thus we can expect that the behavior of the v1 band reflects the property of methane adsorbed and hence the nature of the cationic adsorption site. On the other hand, the large absorption band at 3000 cm-l is assigned to the v3 vibration of methane molecule.13J4 Since this vibration is IR active, the v3 band is observable both for gaseous methane and for adsorbed methane. Thus it is not easy to derive information about the adsorbed methane from the behavior of the v3 vibration. It is worth noting, however, that an IR study a t much lower temperatures and low coverages enables us to derive information about the adsorbed state of methane. A splitting of the degenerated v3 absorption band is reported to be available for this purpose.13J4 Since the experimental temperature adopted in the present study was not sufficiently low to elucidate this point, the following discussion is limited to the behavior of the vl band. Peak Shift. Peak shifts of the vl bands for methane adsorbed on different adsorbents are shown in Figure 6; the original position is 2914 cm-l according to the report on the Raman spectroscopy of gaseous methane.21 If we take the preceding discussion into account, the peak shift of the v l band ( A v , ) would be related to the interaction energy between methane molecules and cationic adsorption sites. This view is supported by the fact that Avl decreases according to the sequence Li+ > Na+ 9 K+ > Rb+ Cs+. Since the accuracy of the peak position is better than 2 cm-', this sequence is significant and deserves special attention. It must be pointed out here that the sequence mentioned above is almost completely in parallel with the sequence observed for the heat of adsorption q a t . Thus we may expect that the sequence of q,: is not meaningless and probably reflects the sequence of the adsorptive force of the cationic site. Intensity of v1 Absorption. On the basis of the above discussion, it has been assumed that the total amount of adsorption ( N , molecule/g) is divided into two parts, Le., the amount of molecules entrapped in the zeolite pore (Nl, molecule/g) and that residing on the cationic sites (N2, (20) Ewing, G. E. J. Chem. Phys. 1964, 40, 179. (21) Herzberg, G.Infrared and Raman Spectra of Polyatomic Molecules; Van Nostrand: New York, 1945;p 306.
Langmuir, Vol. 4, No. 2, 1988 437
IR Spectra of Methane Adsorbed by ZSM-5
Table 11. Adsorptive Properties of Ion-Exchanged ZSM-5 Zeolites
v,,
zeolite HZSM-5 LiZSM-5 NaZSM-5 KZSM-5 RbZSM-5 CsZSM-5 silicaliteC a
mL.g-l 64 68 66 63 54 44 57
1 esu = 3 X lo4 V.m-'.
Vm,
Vm,
molecules/ cation 2.2 2.3 2.3 2.2 2.1 1.8
molecules/ unit 16.5 18.0 17.6 16.9 15.5 13.4 14.7
4.2, io3 J-mol-' 23 32 30 27 25 24
A, m.mo1-l 1.4 x 103 5.4 x 103 5.2 x 103 4.7 x 103 4.0 x 103 3.2 x 103
E,O 106esu 0.8 1.6 1.6 1.5 1.4 1.2
103 J-mol-' 5.1 9.9 7.2 5.7 5.1 4.9
PA
Kz 15 192 45 21 15 13
21
K1 = 4.1 X 10" Pa-' for every adsorbent. Reference 8.
molecule/g). It must be pointed out that the v1 absorption intensity (S) relates only to N2 and not directly to the total amount of adsorption N, S is an integral IR absorption intensity defined by S = [ l / (W/a)]Jlog(Io/l) dv where W, a, I,, I, and v denote the weight of the sample disk, the sample diameter, the intensity of incident light, the intensity of transmitted light, and the wavenumber, respectively. Thus, we have to clarify the relation between N a n d N,, before discussing the IR absorption intensity. If we assume that both N l and N 2 are expressed by Langmuir-type adsorption isotherms, the following relations should hold: p1 = pl0 + RT In [ e , / ( l - e,)] (1) p2
KiKz,b 10-'Pa-' 0.62 7.9 1.9 0.9 0.6 0.5
= p Z 0 + RT In [ e 2 / ( i- e,)]
(2)
where p, denotes the chemical potential of methane adsorbed over the j site 0' = 1 , 2 ) , the superscript o indicates the standard state, and 0 . denotes the methane coverage for the j site. At equilibrium, we obtain the following relation by equating eq 1 with eq 2: a p 2=~RT In [e,(i - e l ) / e l ( l - e,)] (3)
Pressure
/
103pa
Figure 7. Linear p/S-p relationships proving the validity of theoretical treatment: 0 , LiZSM-5; A, NaZSM-5; A, KZSM-5; O , RbZSM-5; H,CSZSM-5; 0,HZSM-5. ,35
I
where Ap20 = p Z 0 - pIo. Putting el = N l / b l ,8, = N2/b2, and K 2 = exp[-Ap,'/RT] into eq 3 gives (4) N2 = b&&i/[bi + (k2 - 1)NiI where bl and b2 denote the adsorption capacities of site 1 and site 2, respectively. On the other hand, the following equation also holds a t equilibrium: N1 b i K $ / ( l + KG) (5) Combining eq 5 with eq 4 yields
Nz = b z K i K G / ( l + K ~ K G )
(6)
The total methane pressure p involved in eq 5 and 6 is related to the total amount of adsorption N because the adsorption isotherm has already been determined (Figure 2 ) . Thus eq 6 can be regarded as the relation between N and N2. The integral IR absorption intensity S must be related to the amount of adsorption N z by S = AN,, where A is the intrinsic molecular extinction coefficient of the IR absorption. Combining this relation with eq 6 gives S = Ab&'K$/(l+ K,K$) (7) Rearrangement of eq 7 yields p/S = l/AKlK2b2 + p/Ab2
(8)
which shows that plS should be linearly related to the equilibrium methane pressure p . Presented in Figure 7 are the relations between p/S and p for every adsorption system studied in the present work. Evidently the linear relation required by eq 8 is satisfied, indicating the validity of the theoretical consideration mentioned above. It is important to note that the linear
-
-
>
- 20
1
OO
,
,
05
1.0
Ionic radius ( r o 1 I
a
,.Ij I5
A
Figure 8. Parallelisms among heats of adsorption qs:, IR peak 0 , 4,: - 20.92; shifts Avl, and chemical potential change 0,-Akza.
relation enables us to evaluate b2A and K1K2. The slope of the straight line shown in Figure 7 gives the value of b2A. On the other hand, the number of cationic adsorption sites b2 is known (Table I). Thus we can readily obtain the value of A for each adsorption system. Values of A thus obtained are summarized in Table I1 (in mmol-'). I t must be pointed out that the A values decreases according to the sequence Li+ > Na+ > K+ > Rb+ > Cs+ > (H+). This indicates that the intensity of v1 absorption is determined by the electric field exerted by the cation site. The A value obtained above enables us to evaluate the strength of the electric field ( E ) of the cation site, because A is related to E by the relation', A = ( A/ 3c)E2(aa /aQ) (9)
438
Langmuir 1988, 4, 438-445
where c is the velocity of light, and &/dQ is the average of the first derivative of the polarizability a of methane with respect to the normal coordinate Q of the vibration. If we adopt the literature valuez2for (da/dQ, Le., 2.27 X 10-ls cm2.g-1/2.p/2 (Ris Avogadro’s number), values of E summarized in Table I1 are obtained. It must be pointed out that these E values are of a reasonable order of magnitude compared with the values (1.6-2.6 X lo5 esu) reported for Na+ in the NaA zeolite.12 It is also important to recognize that the field strength E decreases according to the sequence Li+ > Na+ > K+ > Rb+ > Cs+. This proves the consistency of the present discussion about q,:, vl, and A. Finally, we discuss the chemical potential difference Ap20. As already mentioned, the slope and the intercept a t the ordinate of the straight line shown in Figure 7 give the value of KlK2. It must be pointed out that K1(4.105 X Pa-’ at 226 K) obtained from the adsorption data for silic&te* can approximately be regarded to be common
for every adsorption systems studied. Thus we can readily evaluate K2 and the corresponding Ap20 value for each adsorption system. Table I1 also summarizes the values of K 2 and ApzO. The chemical potential difference Ap20 is expected to have an intimate relation to the heat of adsorption qs:. Since the enthalpy of adsorption of methane onto the cationic sites from the silicalite-like site can be approximated by AHo = -(q,O - 20.92 kJ/mol, we expected that information about the corresponding entropy of adsorption ASo would be obtainable by comparing Ap20 with q,:. Unfortunately, however, as we can see in Figure 8, the difference between Aw20 and qs: - 20.92 is too small to be discussed; i.e., a theoretical entropy analysis appears meaningless. Thus we can obtain little information about the difference between the motional state of methane over the silicalite-like site and that over the cationic site. It is only suggested that the motional-state difference mentioned above would not be large.
(22) Kahn,R.; Cohen de Lara, E.; MBller, D. J.Chem. Phys. 1986,83, 2653.
Registry No. CHI, 74-82-8;Li, 7439-93-2;Na, 7440-23-5;K, 7440-09-7; Rb, 7440-17-7.
Oxygen Quenching Studies of Nonaqueous Dispersions of Poly(viny1 acetate) Labeled with Phenanthrene Groups’ Luke S. Egan and Mitchell A. Winnik* Department of Chemistry and Erindale College, University of Toronto, Toronto, Ontario, Canada M5S 1 A l
Melvin D. Croucher Xerox Research Center of Canada, 2660 Speakman Dr., Mississauga, Ontario, Canada L5K 2Ll Received January 20, 1987. I n Final Form: August 7, 1987 This paper describes experiments on nonaqueous dispersions of poly(viny1acetate) (PVAc) particles sterically stabilized with poly(Zethylhexy1methacrylate) (PEHMA). Particles were prepared with trace amounts of phenanthrene (Phe) groups covalently bound to either the PVAc core phase or the PEHMA stabilizer. Phe fluorescence was very similar in degasaed samples of the two materials, but strong differences were noted in the presence of oxygen. Detailed oxygen fluorescencequenching studies provide information about particle morphology, about swelling of the stabilizer phase by the dispersion medium, particularly for PEHMA trapped inside the particle, and, most importantly, about the PEHMA-PVAc interface. In the presence of the dispersion medium (here cyclohexane), the interface is very diffuse: the presence of 7 w t % of PEHMA in the particle transforms nearly half the PVAc component into a phase swollen with cyclohexane and extremely permeable to oxygen.
Introduction In a previous paper2 we described the synthesis and characterization of nonaqueous dispersions (NADs) of poly(viny1 acetate) (PVAc) particles sterically stabilized with poly(2-ethylhexyl methacrylate) (PEHMA) in which fluorescent g r o u p were covalently attached to either the PVAc or PEHMA molecules. In these materials, PVAc is the major component, representing 90+ % by weight of the ca. 300-nm-diameter particles. I t is insoluble in the hydrocarbon media in which the particles are colloidally dispersed and is referred to as the “core polymer”. The stabilizer, PEHMA, comprises the rest of the material. It (1) Luminescence Studies of Polymer Colloids 17. For the previous paper in this series, see: Winnik, M. A.; Egan, L. S.; Tencer, M.; Croucher, M. D. Polymer 1987,28, 1553-1560. (2) Egan, L. S.; Winnik, M. A.; Croucher, M. D. J . Polym. Sci., Polym. Chem. Ed. 1986,24, 1895-1913.
is very soluble in hydrocarbons. A surface shell of PEHMA, grafted to the PVAc, confers colloidal stability to the particles. As our studies indicate, the remainder of the PEHMA is buried in the particle interior. Our objective is to use fluorescence quenching experiments to elucidate the morphology of the particles, which we take to be a prototype typical of other NAD systems of industrial importance. LPhe
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0743-746318812404-0438$01.50/0 0 1988 American Chemical Society
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