7604
J. Phys. Chem. 1995,99, 7604-7612
Molecular Adsorption-Desorption Reactions of Ammonia on Alkali Halide Clusters and Nanocrystals Margie L. Homer$ Frank E. Livingston, and Robert L. Whetten*?$ Department of Chemistry and Biochemistry, University of California, Los Angeles, California 90024-1569 Received: September 30, 1994@
The initial adsorption of ammonia molecules (NH3) on larger alkali halide clusters and small nanocrystals (mainly NaF) has been investigated by flow-reactor methods as a function of cluster size and temperature. An analysis of the strong size-dependent reactivity of positively-charged NaF clusters, along with their computed structures, indicates that a particular type of defect in the nanocrystal structure facilitates adsorption. This defect is formed by removing an ion-pair from adjoining face and internal sites of a perfect crystallite, creating a basket-like opening. KF nanocrystals show very similar reactivity patterns, reflecting their corresponding structures, but LiF clusters follow a different pattern. It was established that NH3 adsorption on preformed NaF nanocrystals takes place under equilibrium conditions. The equilibrium constant for initial adsorptiondesorption increases with decreasing temperature (250-340 K) and allows one to derive heats of N H 3 adsorption near 0.2 eV for the more reactive (defective) nanocrystals. The much lower reactivity of negatively-charged clusters is ascribed to an additional kinetic-dynamic barrier to adsorption.
I. Introduction Chemical and physical processes on surfaces of ionic solids are of particular interest and have been studied e~tensively.'-~ Metal oxides are important as catalysts for both hydrogenation and amination reaction^;^,^ and the MgO surface is used as a substrate to support metal catalyst^.^ The contact charging of alkali halide surfaces is an important step in potash pro~essing.~ Surface reactions of small salt particles are important in nature as well as industry. In marine regions, salt water droplets in the ocean spray are propelled into the atmosphere, where the water evaporates, leaving behind airborne sea-salt (mainly NaCl and NaBr) particles. In the urban atmosphere these salt particles react with nitrogen oxide pollutants to release ClNO:! and BrN02. These molecules are important sources of chlorine and bromine atoms, which is turn have been implicated in catalytic ozone depletion.5 Important surface processes such as these frequently take place preferentially at defect sites, such as steps, adatoms, or vacancies. For example, the contact charging of NaCl is believed to be mediated by adsorption of H20 in surface cation vacancies: and dissociative adsorption of H20 occurs at C1vacancies on a NaCl film.6 However, an elucidation of defectcontrolled surface processes is difficult to obtain: sophisticated surface preparation and analysis techniques are needed to investigate processes at various defect sites such as adatoms' and steps.* Scanning tunneling microscopy (STM)has also been employed in identifying molecular adsorption at surface defects9 Unfortunately, many important tools of modem surface science are still better adapted to dealing with surfaces as extended, atomically perfect crystalline regions. Furthermore, electron spectroscopy methods are generally not well suited to the study of ionic or insulating surfaces.2s6,'0 STM is inapplicable to insulators, but defects in insulators have recently been imaged using atomic force microscopy (AFM)." Molecular adsorption is another approach for probing the local environment of an ionic surface.I2-l6 For example, the infrared (IR) spectrum of an adsorbed molecule can yield information about its binding site, including important defect sites.I2 By
' Present address: Jet Propulsion Laboratory, 4800 Oak Grove Drive, M.S. 302-306, Pasadena, CA 91 109. Present address: Georgia Institute of Technology, School of Physics, Atlanta, GA 30332-0430. Abstract published in Advance ACS Abstracts, May 1, 1995.
*
@
the 1970s, the difficulties in investigation of surface defects on extended crystalline surfaces, arising from the low defect density, were well e~tablished;'~,'~ the use of dispersed powders to increase the surface area often failed to ensure well-defined and characterizable surface structure. What was needed were small particles (high surface-to-volume ratio) with well-defined surfaces. Using high-resolution electron microscopy, Smart13 demonstrated that films of alkali halides deposited onto cooled (- 196 "C) surfaces are actually made up of many small (1-3 nm diameter), discrete particles, although they are also highly colored (reflecting a high density of electronic defects). Nitric oxide (NO) adsorbs preferentially at defect sites on these small particles. Infrared spectro~copy'~-'~ revealed that adsorption of NO to NaC1, NaF, or LiF films destroys the color centers; the IR spectra indicated that NO adsorbs to NaCl to form NO+, donating an electron to an electron-deficient center, whereas NO forms NO+ and NO- on the NaF and LiF films.I3 Similarly, cube-shaped crystallites of MgO (100 nm scale) were produced by oxidizing strips of Mg in air, and these cubes were used to study spectroscopic and chemical properties of various sites on the MgO surface.I6 Subsequently, Zecchina et showed that "...NaCl can be prepared in form of cubelets...", allowing them to study the "CO adsorption on the (100) facelets of the microcrystals...". The cubelets studied by Zecchina et al. were in the 1-3 nm r a r ~ g e ' ~and . ' ~ therefore contain fewer than lo3 ions. Although these experiments provided researchers with a way to increase the number of certain types of defects in a surface, it did not provide a precise method for studying defects. An alternative approach could be provided by methods for preparing and investigating size-selected inorganic clusters that are sufficiently large to assume the crystalline structure as the bulk crystal. These could be used to examine the role of defects as well as the perfect surface of the crystal and are ideal for studying surface properties because each provides a large surface-to-volume ratio and a well-defined surface. As the size of a cluster increases, there will appear structures where all the ions can fill an ideally truncated lattice, so that there are only perfect nanocrystalline surfaces, complete with facet edges and vertices. More frequently, there will also be structures where the crystal will not have the correct number of ions to fill a small lattice, and so the surface will contain defects. Because of the periodic nature of this growth pattem, these structural characteristics repeat themselves over n in a pattem; Le., they
0022-3654/95/2099-7604$09.00/0 0 1995 American Chemical Society
al.'43'5
J. Phys. Chem., Vol. 99, No. 19, 1995 7605
NH3 on Alkali Halide Clusters and Nanocrystals
a
10-5 torr
IO-’ torr
DMCP Det&Of
I
Figure 1. Optimized structure of N ~ 3 F 6 2 + . ~ ~
provide a recurring series of perfect surfaces and surface defects which can be exploited to examine the nature of the adsorption process. Alkali halide clusters have been extensively investigated on a size-selected basis and have certain well-established structural characteristics: the domainance of rock-salt lattices and cuboidal morphology, i.e. with only (100) facets exposed, for clusters of more than about 25 ions, marking the transition from clusters to nanometer-scale crystallites (or “nanocrystals”). Figure 1 illustrates the relaxed and highly stable structure of a larger cluster (or small nanocrystal) composed of 125 ions, computed using the potential appropriate for NaF, chemical formula N%3F62+. Although experimentsI7and theoretical modeling’ are in agreement on the structure of the clusters which can form perfect cubes such as that in Figure 1, less is known from experiments about the structures of the other clusters in the series. A few experiments have suggested structures for some of the other s i ~ e s , l ~ particularly -~~ those which can be derived from perfect structures by terracing, and these are supported by calculations which predict their stability and give other structural detail^.'^.^^ If perfect cubic structures repeat, i.e. exist in distinct generations, throughout the cluster series, then structures representing different types of surface defects will also repeat. Molecular adsorption on nanocrystals, as investigated using isothermal flow-reactor methods, could be used to probe the structures of the nanocrystals and could subsequently be correlated with structure on a bulk surface. Our aims have been to establish the nanocrystal flow-reactor approach to studying surface reactivity, to explore surface phenomena on ionic crystallites,particularly by determining the role of surface defects through correlation of structure with reactivity, and to quantify this information by the extraction of precise thermodynamic quantities. In particular we have been studying the interaction of polar molecules with ionic surfaces and have examined interactions of different alkali halides with NH3 and H20.20-22 The choice of NH3 and salts is motivated also by the extent of information already available for these materials. In a recent short report, we have given a summary of our preliminary work on these reactions, as observed at ambient temperature.20*2iThe present report provides a full description of these experiments and results and adds information establishing the equilibrium control of the reactions and their temperature dependence.
11. Experimental Methods A. Apparatus. Chemical reactivity of clusters and nanocrystals is measured using a laser-ablation cluster source in conjunction with a flow-reactor.20-22 Related cluster flowreactor designs are used by several other group^.^^-^^ The experimental apparatus is depicted in Figure 2. The nanocrystalsklusters are produced in a pulsed laser-ablation source, where they are cooled and entrained in a buffer gas (He). The clusters are diluted to prevent further cluster growth or neutralization and enter the flow-reactor where they mix with a reactant gas (diluted in the same buffer gas). The entire flow
I
Flow
Pulsed General V.?lve
Rod
Vaporization Laser ca. 5 m.1 337 nm Retarding VoItilgc
Reacceleration Voltage
Pulsed Extraction Ficlcl
t
t
Liquid Nitrogen IM’fled
6“ Diffusion Pump
Pulsed Cas Valve
4” Diffusion 1’11 mp
b
k
n
Figure 2. Apparatus for flow-reactor experiments. (a) Right: source chamber. Left: mass spectrometry and detection chamber. (b) Schematic diagram (top view) of the cluster source and flow-reactor: a, Newport BVlOO pulsed molecular beam valve; b, gas reservoir for the valve; c, ceramic spacers (1.3 mm) for thermal isolation; d, stainless steel vaporization block (7.6 mm long, 1 mm diameter gas channel entrance, 2.5 mm diameter gas channel exit); e, sample target (5.03 mm diameter); f, stainless steel expansion piece (25.4 mm long, 2 mm i.d.); g, copper cooling coils; h, copper or aluminum flow-reactor (6 mm i.d.); i, stainless steel tubing (45 mm long, 1.6 mm o.d., 0.76 mm i.d.) to connect the General valve to the flow-reactor; j, General valve; k, laser vaporization channel (1 mm diameter).
expands into vacuum at the end of the flow-reactor, and the resulting jet is skimmed to form a cluster beam including reaction products. The charged clusters are mass analyzed by a time-of-flight mass spectrometer, without any need for subsequent ionization. B. Laser-Ablation Source. The cluster source and flowreactor are shown in Figure 2. The ablation source has also been described e l ~ e w h e r e . ~A~ . Newport ~~ BVlOO pulsed molecular beam valve is backed with 80-100 psi of a carrier gas mixture consisting of 2-4% SF&e; the valve releases a gas pulse which flows over the alkali metal rod. The metal rod is ablated using radiation from a nitrogen laser (1 = 337 nm and -8 &/pulse) focused on the sodium rod using a 20 cm lens, timed with respect to the gas valve. The metal vapor is entrained by the carrier gas, and clusters form and equilibrate within the expansion piece. Appropriate metal rods are used
7606 J. Phys. Chem., Vol. 99, No. 19, 1995 for Na, K, and Li; an SF&Ie mixture is used as the carrier gas. Details on the production of the alkali metal rods have been published p r e v i o u ~ l y . ~ ~ To remove trace water contaminant, the carrier gas line has a water filter made from a coiled copper tube filled with molecular sieves. Before each experiment, the filter was wrapped with heating tape and then heated and evacuated with a mechanical pump overnight. For the SFdHe mixture, the filter is placed in a dewar with an ice bath; for pure He, the filter is placed in a dewar with liquid nitrogen. C. Flow-Reactor. After passing through the expansion piece, the cluster mixture further expands into the flow-reactor (see Figure 2b). This expansion terminates cluster growth. Reactant gas is synchronously pulsed into the flow-reactor through an inlet port using a General valve. This valve is mounted by angle brackets inside the source chamber and attached to the flow-reactor by a 0.15 cm diameter x 5 cm long stainless steel tube which is inserted in a hole in the reactor 0.25 cm from the beginning of the reactor. Three flow reactors of different lengths are used to vary the reaction time (1.3, 2.5, and 3.8 cm long, all with a 0.6 cm inner diameter). The total backing pressure on the General valve is 600 Torr. Most reactant gas mixtures used were from 0.2 to 4% (partial pressure) NH3 in He. Mixtures of less than 2% NH3 were diluted from a purchased 2.03% NH3Me mix, 2% nominal mix was taken straight from the purchased mixture, and anything richer than 2% was diluted from a purchased 10%NH3/He mix. All reactant gas mixtures were allowed to sit for at least 5 h after dilution to ensure uniform mixing. The reactions were normally measured at 300-335 K, except for the temperaturedependent sweeps, which spanned the 245-340 K range. A thermocouple probe is attached to the flow-reactor to monitor the temperature of the flow-reactor. The cooling coil on the expansion block is used to cool the source by flowing liquid nitrogen through it, as described for earlier experiments on cold Nan clusters, where the laser-ablation source runs continually at liquid nitrogen temperature^.^^ In order to vary the temperature, the liquid nitrogen was turned off and the flowreactor was allowed to warm up slowly. The temperature change was sufficiently slow that it was possible to record highquality 1000-shot averages ( - 2 min) with no more than a 2 K temperature increase during each accumulation. The temperature is varied either by driving a power resistor attached to the flow-reactor, which increases the temperature of the flow-reactor above ambient, or by flowing cold nitrogen through copper cooling coils thermally equilibrated with the laser-ablation source and the flow-reactor. In the latter case, the source and flow-reactor and allowed to warm gradually, and equilibrium constants are measured as the temperature changes. Experiments attempted below 243 K failed because the NH3 vapor begins to condense in the flow-reactor. The temperature scans are taken in two different ways. The first is simply to record complete mass spectra as the temperature changes. Subsequently the mass spectrum at each temperature is analyzed to yield an equilibrium constant for each cluster size. Later experiments were run by using a scanning program that had previously been used for optical spectro~copy,~’ in which only the integrated selected parent and product peaks at each temperature are saved. All of the scans were run to optimize conditions for the most reactive clusters. D. Mass Spectrometry, Cluster Detection, and Data Analysis. After the clusters pass through the flow-reactor, they expand with carrier gas into a vacuum chamber where the cluster jet is skimmed and passes into a second, differentially-pumped chamber. A voltage pulse (f2000 V) extracts the charged clusters into a field-free region of a screenless reflectron timeof-flight mass spectrometer; and the time-separated ion-packets are detected by a dual microchannel plate assembly. The wave
Homer et al. form from the microchannel plates is amplified using a 100 MHz video amplifier, digitized, averaged over 400-2000 pulse cycles, and stored. For these reactivity studies, the peaks in the mass spectrum are identified and their areas are integrated and stored for analysis including peak integration and deconvolution. In order to ensure that the reactions do not involve the breakup of the clusters, a separate summation of reactant and products is produced to compare with the abundance distribution when no reactant is added.22 Most fast-flow-reactor measurements on clusters have been carried out at ambient tem~erature.~~-~’ Generally, if the reaction is thermodynamically controlled, one can extract an equilibrium constant for each cluster, and if the reaction is kinetically controlled, then one can extract relative rate constants for the cluster reactions. A plot of the equilibrium constant at various temperatures will provide the heat of reaction, AH, and the change in entropy of the reaction, AS. Similarly, a plot of rate constants at various temperatures will provide the energy of activation for the reaction. In one example, Weiller et al. determined that the reaction of iron clusters with water, in their flow-reactor, was taking place under equilibrium conditions; they further determined that the extent of reaction, i.e. the equilibrium constant, increased with decreasing t e m ~ e r a t u r e .Although ~~ they estimated the enthalphy of adsorption of D20 on Fen clusters, a study of equilibrium constants versus temperature was not reported.24 More recently, Riley and co-workers reported further temperature dependence studies of the uptake probabilities of DZand N2 on Ni, clusters32and the uptake of DZon small Con clusters.33
In. Results and Analysis A. Adsorption of Ammonia on Na,F,-I+. Reactivity Variation with n. Figure 3 compares mass spectra of Na,F,-l+ with and without a reactant gas pulse of the NH3/He mixture. The spacing between sequential cluster peaks is 42 amu, and the NH3 mass is 17 amu, so that a peak for the adsorptionreaction product, [Na,Fn-~+].NH3,appears 17 amu higher than that for its unreacted parent. An NanFn-l+ cluster will be referred to by size as an n cluster, and similarly, a [ N ~ I ~ F I ~ + ] . N H ~ product will be referred to as n.NH3. Once N H 3 is added, several striking features appear in the mass spectrum. As reported in brief earlier,20,21the clusters differ greatly in their reactivity toward ammonia. The perfect crystallites, n* = 14, 23,32, and 38, do not appear to adsorb at ambient temperature. In contrast, several clusters are considerably more reactive than others, namely those with n* - 1, or 13, 22, 31, and 37. In Figure 3, the 13 and 22 clusters are approximately two-thirds reacted away to form products with one and two NH3 molecules. Although less drastic in appearance, 31 and 37 still show a considerable decrease in parent intensity relative to neighboring sizes. Many others are reactive, but less so: 16, 19, and 34. At this (low) exposure of NH3, it is difficult to differentiate between many of the less reactive sizes simply by looking at the mass spectrum. Furthermore, smaller clusters, n < 9, appear to be almost as reactive as the n* - 1 clusters and certainly more reactive than all other large clusters. In the baseline of the unreacted mass spectrum (Figure 3, top), it is clear that the Na,F,-I+ clusters are not the only species in the beam. These extra peaks can be variously attributed to Na,F,-,+.H20, N&F,+, and Na,F,-2+, although they are in most places quite weak. It is possible to subtract their contributions to the integrated peak areas of the products, when overlapped, using a deconvolution program. Reactivity measurements for the larger alkali halide clusters are more difficult because of limited mass resolution. Also, it turns out that the larger crystals are less reactive than the small ones, so higher partial pressures of ammonia are necessary to see reactivity for large crystals. Nonetheless, it was possible
N H 3 on Alkali Halide Clusters and Nanocrystals
J. Phys. Chem., Vol. 99, No. 19, 1995 7607
a
I
A
+ -0.2-I d
!
A
2
x
2
14
-0.6
32
Reactor Length / cm
+
Figure 4. Plot of In{ [NanF,-l+]/([N~Fn-~+] [NanFn-~+]*NH3)} versus reactor length. The triangles are n = 11, the squares are n = 13, the circles are n = 16, and the diamonds are n = 22.
to extract
I
I
b
53
where t is the exposure (transit) time in the reactor. If the reaction is thermodynamically controlled, one calculates the equilibrium constant for the initial adsorption: Na,F,-
I,
53
1.4% NH3/He
Figure 3. (a) top, time-of-flight mass spectra of Na,F,-I+ taken when no reactant gas was pulsed into the 3.8 cm flow-reactor; bottom, as above, but with 1.4%NH3/He gas pulsed into the flow-reactor. (b) top, time-of-flight mass spectra of NanFn-I+ taken when no reactant gas was pulsed into the flow-reactor; bottom, as above, but with 1.4% NH3/He gas pulsed into the flow-reactor. In both a and b n denotes parent NanFn-I+and n* denotes [NanFn-1+]*NH3.Temperature = 308
K.
to determine reactivity for Na,F,-I+ to n = 50-70 (not shown). All show at least slight adsorption of N H 3 , and one of the n* 1 defect crystals, 52, is significantly more reactive than the neighboring crystals. Kinetics or Equilibrium? To extract quantitativeinformation from results such as those in Figure 3, one must determine whether the reactions as observed are governed by kinetics or if equilibrium has been established. If the reaction is kinetically controlled, one extracts relative reaction rates for the initial adsorption process: Na,F,-,+
+ NH, -NanFn-I+-NH3 kn
(1) by assuming constant [NH3]26 and expressing the measured depletion of parent as
,+iNH,
Kn
Na,F,- i+*NH3
(4)
For a kinetically controlled reaction, one may determine a rate constant by varying the concentration. However, the rate constant and the equilibrium constant have the same dependence on reactant concentration. To determine whether the reaction is in the kinetic regime, one must vary the reaction time. In the case of the fast-flow-reactor, the time of reaction depends on the velocity, vflow,of the cluster beam and the length, 1, of the reactor, which must, therefore, be varied. Rearranging eq 2
If we plot the left hand side of eq 6, versus reactor length, any variation proportional to the reactor length indicates that the reaction is taking place on the time scale of the reactor length change and implies that the reaction is kinetic. If the reaction quickly achieves equilibrium,then no time increase will change the extent of reaction. Plots obtained for several cluster sizes (11, 13, 16, and 22) show that the extent of reaction is nearly independent of reactor length (1.3-3.8 cm) and is, therefore, not kinetically controlled (Figure 4). This can also be seen simply by comparing two mass spectra taken with different reactor lengths but all other conditions held constant. For example, Weiller et al. had compared their mass spectra at two different lengths to determine whether the reactions of Fen D20 were at equilibrium. They found identical mass spectra, implying equilibrium.24 On the basis of these results, we choose to plot all reative reactivity data as ratios (eq 5 ) , rather than as fractional extents of reaction (eqs 2 and 3). Because the partial pressure of NH3 in the reactor is not known, the absolute equilibrium constants cannot be obtained directly from eq 5, although they can be obtained on a relative basis for all cluster sizes using the 'relative
+
7608 J. Phys. Chem., Vol. 99, No. 19,1995
Homer et al. 14
22 I
0
6
1
10
ki
20
30
40
14
Cluster Size (n)
I
I
23
I
I
15
25
35
45
Figure 6. Top: Time-of-flight mass spectra of KnFn-,+ taken when no reactant gas was pulsed into the flow-reactor. Bottom: As above, with 10% N H a e gas pulsed into the flow-reactor. T = 308 K.
55
Cluster Size (n) Figure 5. Relative reactivity plotted against cluster size for NanFn-I+ -I- NH3 [NanFn-~+]*NH3, 1.4% N H a e mixture. T = 308 K. (a) n = 5-39. (b) n = 20-54.
-
13
reactivity' ratios
R, = [Na,F,_,f~NH3]/[Na,F,,-,+]
(7)
which are R, = K,,[NH3]. The relative reactivity of the adsorption of ammonia on NaF,-I+ is shown in Figure 5 , which quantifies what was observed from the mass spectrum. B. Adsorption of Ammonia on &Fn-l+. The adsorption of NH3 on &F,,-I+ is found to be generally less reactive than the sodium fluoride clsuters under similar conditions. The KF clusters had to be exposed to a larger partial pressure of ammonia (>4% NH3/He reactant mixture) before addition of one molecule was seen at ambient temperature. Features of the relative reactivity pattern versus cluster size for this reaction are otherwise very similar to that of sodium fluoride. Mass spectra showing the KF clusters under reactive and nonreactive conditions are shown in Figure 6. The top half is a mass spectrum of the K,F,-I+ clusters without reactant gas. There are slight ledges in the abundance after 14,23,32,and 38. These are the familiar magic numbers indicating that potassium fluoride, like other alkalihalides, forms cuboidal rocksalt crystallites. The weak peaks interspersed in the main series arise from several different series, namely KnFn-2+,&F,,+, and (&F,,-l+)*H20, and are not difficult to subtract. The bottom half of Figure 6 is a mass spectrum of &F,,-I+ and reaction products formed when an N H a e mixture is pulsed into the flow-reactor. The mass spacing for a KF unit is 58 and 60 amu (K has isotopes 39 and 41 amu), and the peaks for these clusters are broadened (relative to those of the NaF clusters) by the K isotopes. Where the peaks overlap, they are deconvoluted and the weighted mass distribution of the integrated peak areas from the reaction mass spectrum shows that the initial concentration of parent peak is adequately represented by the sum of the parent and product peak areas. As with NaF, clusters 14 and 23 are the least reactive, and the most reactive clusters are 13 and 22. Although the product peaks are difficult to discern at much higher mass, there is definitely more of a decrease in abundance for two others from
5
1
I
I
5
I
I
I I
9
I
I
l
I I I I
: I
I
,
I
I
I
,
I
I
32 (3
13 17 21 25 29 33
37
Cluster Size (n) Figure 7. Plot of the relative equilibrium constants versus cluster size [KnFn-l+]*NH3,n = 4-38. for the reaction KnFn-1'- NH3
+
-.
the same series, namely 31 and 37,reflecting greater reactivity in relation to neighboring sizes. At lower mass, the clusters n -e 14 are considerably less reactive than are the same-numbered NaF clusters. Assuming that the KF reactions also occur under equilibrium conditions, we have calculated the R,, for several sets of data. Figure 7 shows an example of two sets of data for the relative equilibrium constants, on a logarithmic scale, versus cluster size for the adsorption of NH3 on KF clusters and small nanocrystals. C. Adsorption of NH3 on Li,Fn-l+. The adsorption reactions of ammonia on Li,F,-I+ were measured and also found to be less reactive than those on the corresponding NaF clusters. They had to be exposed to a larger partial pressure of ammonia (10%N H a e reactant mixture) before significant initial adsorption was seen at ambient temperature. Some of the gross features of the relative reactivity pattern versus cluster size for this reaction are similar to those of NaF and KF clusters, but the reaction, overall, shows very different trends. Mass spectra showing the lithium fluoride clusters under reactive and nonreactive conditions are shown in Figure 8. The top half is a mass spectrum of the Li,F,,-I+ clusters when no
NH3 on Alkali Halide Clusters and Nanocrystals
I
J. Phys. Chem., Vol. 99, No. 19, 1995 7609
4
23 I
1
c
K 0.1
Cluster Size (n) Figure 8. Top half Time-of-flight mass spectra of LinFn-I+ taken when no reactant gas was pulsed into the flow-reactor. Bottom half: As above, but with 10%NH3/He gas pulsed into the flow-reactor. T = 308 K.
reactant gas is pulsed into the flow-reactor. The cluster distribution is smooth, and while some earlier experiments on excess electron LiF clusters suggested that the clusters form cubic structure^,^^ other experiments demonstrated that LiF clusters do not always conform to this.34 The bottom half is a mass spectrum of Li,F,-I+ and its reaction products when a 10% NH3/He mixture has been pulsed into the flow-reactor. The mass spacing for a LiF unit is 26 amu, and the peaks for these clusters, like the KF clusters, are broadened by the metal atom isotope distribution. The main products, (LinF,-1+)*NH3, are 17 amu heavier than the parent peaks. When the peaks begin to overlap in the mass spectrum, the nth product peaks is overlapping with the (n+l)th parent peak. These peaks are deconvoluted, and the weighted mass distribution of the integrated peak areas from the reaction mass spectrum shows that the initial concentration of parent peak is adequately represented by the sum of the parent and product peak areas. Again, assuming that the LiF reactions occur under equilibrium conditions, the relative equilibrium constants are calculated from eq 7 for the adsorption of N H 3 on positively-charged LiF clusters and are potted on a logarithmic scale against cluster size for several different partial pressures of ammonia in Figure 9. The numer density of the ammonia can be increased somewhat by increasing the intensity of the gas pulse; this puts more reactant gas in the flow-reactor, but since the intensity of the cluster gas pulse is so much larger, it does not appreciably change the total pressure in the flow-reactor. The most striking feature of the reaction mass spectrum for Li,F,-I+ is the nearly monotonic decrease in extent of reaction for n -= 38. This is in marked contrast to the highly selective reactivity of both the NaF and KF clusters. Overall, the mass spectrum suggests that the smallest clusters are most reactive, and the reactivity tends to decrease with size. From observing a sharp decrease in parent crystal, it appears that Li22F21+ is more reactive than its neighboring sized clusters. Compared to the potassium and sodium fluoride nanocrystals, there are decreases in the reactivity of the Li,F,-,+ clusters after 13,22, 32, and 37; but these decreases are not the most significant ones. They may correspond to ideal structures which are capable of adsorption, but are less reactive than the others. There are two very distinct reactivity minima at n = 28 and 34, whereas for KF these are local maxima and, for sodium fluoride, n = 34 is a local maximum.
-
Figure 9. Plot of the relative equilibrium constants versus cluster size for the reaction Li,F,-I+ -I- NH3 [Li,F,-l+]*NH3, n = 9-39.
D. Temperature Dependence of Adsorption Reactions. By varying the temperature of the reactor and source walls with otherwise unchanged conditions, we first determined that the extent of the NaF cluster reaction with NH3 increases strongly with decreasing temperature. This is also consistent with equilibrium being established. We have further attempted a careful study of the temperature dependence of the relative equilibrium constants for this reaction in order to obtain estimates for the AH of reaction. If the heat of reaction is independent of temperature, then it can be determined by measuring the equilibrium constants of the reaction at different temperatures and relating them by the van’t Hoff equation:35
AH
ln(K,) = -RT
+-AS R
(8)
If one plots ln(Keq)versus 1/T for a reaction, then the slope of the line is -AH/R and the y-intercept is ASIR. Because the equilibrium constants calculated for the nanocrystals in these experiments are relative equilibriumconstants, R,, as the partial pressure of NH3 has been left out, it is necessary to hold all the experimental conditions constant while varying temperature, so that the ammonia pressure can be factored out. The partial pressure of the ammonia will change only the y-intercept of the van’t Hoff plot if the natural log of the relative equilibrium constant is plotted versus the inverse of the temperature. The change of entropy of the reaction cannot be extracted from this data. The flow-reactor part of the experiment is carried out as described earlier. Figure 10 gives two eamples of temperature scans, the first for a highly nonadsorbing cluster, 14, and the second for a strongly adsorbing one, 22. Although there is considerable scatter in the data, it is clear that the 14 cluster shows only a small increase in adsorption with decreasing temperature over the plotted range. In fact, substantial adsorption occurs only below 273 K, and the lowest temperature in this plot is 285 K. The 22 cluster shows nearly a factor of 3 change in equilibrium constant over the same range. Several runs were attempted, and Table 1 gives results (negative slopes from the van’t Hoff plots) from the two runs with the least scatter in the points. E. Reactivity on Negatively-Charged Clusters. Several attempts to measure the reactions of ammonia with negativelycharged alkali fluoride clusters have been carried out. Only at the lowest temperatures, approaching the condensation temperature of NH3 vapor, was any substantial adsorption observed.
7610 J. Phys. Chem., Vol. 99, No. 19, 1995
Homer et al.
13 -1.5
=
-e u
-3 305
315
325
335
E
0 -
22,21
31
32
37
38 - -
.=
=
-0.6 -0.9
23
345
Q)
-5
22
-
-1.24 305
'
'
315
'
'
325
'
'
335
'
'
345
1
10% (1/K) Figure 10. van't Hoff .plots for clusters 14 and 22.
TABLE 1: Heats of Reaction for the Adsorption of NHJ to NanFn-1' -AHIeV
cluster
n,n-1 11,lO 12,ll 13,12 14,13 22,21 23,22
run 4
run 5
0.20 0.15 0.24 0.07 0.16
0.27 0.21 0.26
0
0
0.04
0.18
At ambient temperatures, an upper bound on the adsorption reactivity of Na,-IF,- clusters is estimated to be at least 2 orders of magnitude lower than that of the corresponding positivelycharged clusters. This is in spite of the evidence that there structures are very similar.34 F. Other Reactions. The adsorption reactions ,with ammonia for five different alkali halide clusters have been measured: NaF, LiF, KF, NaCl, and CsI (NaC1 and CsI not shown in this paper). The sodium fluoride clusters are most reactive. Lithium fluoride, potassium fluoride, and sodium chloride show roughly the same extent of reaction under similar conditions but are less reactive than sodium fluoride. Even under extreme conditions, it was nearly impossible to see addition of ammonia to cesium iodide clusters. Further reactions on NaF were done with water, acetic acid, and diethylamine. Water is much more reactive than ammonia, although many of the reactivity patterns are similar. The same is true for diethylamine. No reaction products were detected with acetic acid.
IV. Discussion A. Sodium Fluoride. The most striking features in the relative reactivities plots (Figure 4) are the drops in reactivity going from the n* - 1 to the n* clusters. These features can be related to the nanocrystallite structures. It is well established that the n* clusters are near ideal cuboids with (100) facet^.^^-^* These nonreactive crystallites failed to adsorb ammonia because Na+ ions at edges and vertices are not favorable adsorption sites. Here we assume that the ammonia will adsorb to the alkali fluorides with the nitrogen attracted to the metal ion, as in ammonia adsorption on the MgO surface.2 The most reactive clusters, 13,22,31, and 37, are all one NaF unit away (n* - 1) from being an ideal crystallite with six (100) faces. Without knowing the structure of the n* - 1 clusters (these structures being less well established), it is
Figure 11. Recurring structures for some perfect and defect nanocrystals, Na,F,-I+, for n = 13,14,22,23,31,32,37,and 38. Calculated potential energies with respect to separated ions (in eV): U(13) = -110.6, U(14) = -120.6, U(22) = -194.2, U(23) = -204.2, U(31) = -277.7, U(32) = -287.6, U(37) = -333.8, U(38) = -344.1.
reasonable to assume that they have related structural features on the basis of removing one NaF unit from a perfect crystal, and it is that defect which facilitates the adsorption reaction. Figure 11 shows the lowest energy structures found for the perfect crystals and the defect crystal^.'^,^^ The thermal stabilities of the structural isomers for Nal3Fl2+ have been compared using Monte-Carlo simulations. The simulations show that the basket isomer is the lowest energy isomer and that if the other isomers are heated, their structures collapse into the basket s t r u ~ t u r e . ' ~(An * ~ ~extensive discussion of isomers can be found in refs 19 and 22.) The heating of Na22F21+ isomers has also been investigated and shows similar results: nonbasket isomers decay to the basket isomer upon heating.22 We propose that it is this lcwest energy basket feature which promotes ammonia adsorption on the nanocrystal. The basket structure, with its cation-anion vacancy, has several features which make it a conceivable structure for facile adsorption of an ammonia molecule. First, the negative end of the ammonia molecular dipole would be attracted to the positively charged crystallite. Second, the ammonia molecule appears to fit easily in the basket vacancy with the lone pair electrons in the anion vacancy and the nitrogen (with its partial positive charge) near the cation vacancy, with the hydrogen atoms near the top of the basket and pointing out. Figure 12 illustrates how an ammonia molecule might adsorb to Nal3F12+ where the ammonia molecule (bond lengths to scale) is placed rigidly within the basket to indicate the possible adsorption arrangement. The ion coordinates are taken from the calculations in ref 22, but the sizes of the ions are not to scale. Previous experiments on alkali halide clusters have established that the integrity of the lattice is easily maintained while integrating various species into the lattice: an excess e l e ~ t r o n ? ~a .carbon ~ ~ or sulfur atom?9 an F2-,37 or two paired electrons.38 For Nal3F12+ the average Na-F bond length is 2.1-2.2 which is very close to the computed Na-N bond length for the linear, gas-phase complex
NH3 on Alkali Halide Clusters and Nanocrystals
Figure 12. Possible structure for [Nal3F12+]*NH3.The dimensions for Nal3F12+ are taken from ref 22, and the NH3 (to scale) is placed within the crystal to illustrate the possible adsorption site.
28
4
Figure 13. Lowest energy structures calculated for Na,F,-I+, for n = 17,20,26,28,43. Energy (in eV) with respect to separated ions: U(17) = -148.0, U(20)= -175.4,U(26) = -231.5, U(28)= -250.1, U(43) = -390.0.
Na+*NH3.39The measured heat of adsorption for the Na+*NH3 complex is 1.26 eV?O which is small compared to the 3.3 eV gained by adding a NaF to Nal3F12+. This value includes the rearrangement energy and is calculated using the binding energies from ref 22: E(14,13) - E(13,12) - E(NaF). It is also possible to correlate the nanocrystal structures with the relative reactivities of other clusters. There are a series of peaks which represent local maxima in the relative reactivity pattern: n = 16,19,25,27,34,42,and 46. These clusters can be viewed as the n** - 1 series related to a slightly less-thanperfect crystallite structure pattern: n** = 17, 20, 26, 28, 35, and 43. These less-than-perfect crystallites are terraced, relaxed ring structures.22 The lowest-energy structures calculated for n** = 17,20,26,28, and 43 (see Figure 13) are structures that have added an even number of complete rows of ions to a perfect cube; for example, 17 is a 3 x 3 x 3 cube with a 2 x 3 terrace added on. The terrace and part of the perfect cube relax to a
J. Phys. Chem., Vol. 99, No. 19, 1995 7611 ring.22 Although none of these crystallites form a completed lattice, they each have seven completed (100) faces and a monatomic step; the defect is not on a molecular scale. In some of these crystals the terrace relaxes away from the face to form a ring. The reactivity of the n** - 1 nanocrystals must be related to the removal of a NaF unit from these terraced crystals. What kind of an absorption site this would create is unclear because there is no clear repetition of a particular type of lowenergy structure among these crystals. These incomplete terraces form a step which could act as an adsorption site; the n** crystals do not really have a step site between the terraces, since they relax into rings. The relative reactivities may also be related to the binding energies of the parent clusters and the stability of those clusters against loss of an NaF unit. The calculated energies to lose a NaF unit for an Na,F,-I+ cluster show that certain clusters are less likely to fragment: 14, 17, 20, 23, 26, 28, and 32.22*29*36 Comparison of these numbers to the drop in the relative reactivities (see Figure 5) after clusters 13, 16, 19, 22, 25, 27, and 31 suggests that the clusters which have the most to gain from adding an NaF unit also have the most to gain (energetically) from adsorbing an NH3 molecule. The extension of these recurring patterns to larger clusters and bulk surfaces is discussed in ref 22. B. Potassium Fluoride. Like the sodium fluoride clusters, the largest difference in reactivities is between the n* = 14, 23,32, and 38 clusters and the n* - 1 clusters. This confirms that like the case of sodium fluoride, the (100) faces (as well as the edges and corner) of the perfect potassium fluoride crystals do not adsorb ammonia. The large reactivity for the defect crystal is also similar to what is seen for sodium fluoride, although the 31 and the 37 crystals do not appear to be significantly more reactive than some of the other sizes. It is likely that these defects are also basket-like. There is another series of reactive potassium fluoride crystals which can also be related to the results for sodium fluoride. For sodium fluoride, the defect crystals from the relaxed-terrace crystals were relatively reactive: n = 16, 19, 25, 27, 34, 42, and 46. The local maxima in relative equilibrium constants here are nearly the same series: n = 16, 19, 25, 28, and 34. The only difference in the series is that K28F27+ is a local maxima rather than Na2+26+. The similaritiesin reactivity are probably due to similarities in the crystal structures. The largest difference in reactivity toward ammonia between the sodium and potassium fluoride nanocrystals is in the clusters. The small Na,F,-I+ clusters are more reactive than their nanocrystalline counterparts whereas the small L F n - 1+ clusters are less reactive. The radius ratio varies quite a bit from LiF to NaF to KF, and this difference will manifest itself primarily in the structures of the small clusters. When the number of ions is too few for the cluster to form completed lattices, the structures vary quite a bit.22 Calculations for CsnIn-l+clusters have found different low-energy isomers than for the Na,F,-I+
C. Lithium Fluoride. As seen in Figure 8, the overall trend in reactivity for the Li,F,-I+ clusters and nanocrystals is the decrease in reactivity with increasing cluster size. It is unlikely, then, that the reactivity of the lithium fluoride crystals corresponds to the same structures seen for sodium and potassium fluoride. From the comparison of the alkali-halide classical potentials and the dependence of bond lengths on these potentials, it was found that shorter bond lengths tend to give ringlike structuresF2 The bond lengths for LiF nanocrystals will be much shorter than for NaF, and it may be that the structures for LiF are based more on ringlike or cage structures than cubiclattice structures.
7612 J. Phys. Chem., Vol. 99, No. 19, 1995
The relative reactivity of the lithium fluoride crystals does not change dramatically simply from one size to the next. The abundance patterns of excess-electron lithium fluoride clusters, (Li,F,-I+)*e-, show that the stability against loss of a LiF unit varies less from one size to the next than for (Na,F,-I+)*e- and ( N a , C l , - ~ + ) . e - . ~ Fragmentation ~~~~ patterns from the collision of Li,F,-l+ clusters with a surface show that there is relatively little difference in binding energy from one size to the next.33 If the stability of Na,F,-i+ clusters against the loss of NaF can be correlated with the reactivity (Le. the preference of the cluster to add a NaF unit), then perhaps that can explain the overall reactivity trend for Li,F,-I+ clusters. There is little size differentiation in adding or losing a LiF unit, and therefore there is little size differentiation in adsorbing an NH3 molecule. D. Negatively-Charged Clusters. As stated here and in previous reports,20x21 the negatively charged clusters and nanocrystals are much less reactive than positively-charged ones. If we assume that the adsorption is between the nitrogen on the ammonia and the Na+ ion and we view the ammonia approaching a cluster as a dipole approaching a charged cluster, one can imagine that as ammonia approaches the positively-charged cluster, the dipole will orient with the N atom pointed toward the cluster, facilitating the adsorption; whereas with the negative clusters, the ammonia dipole will approach the cluster with the hydrogen atoms pointing toward the cluster and will never absorb. This nonreactivity of the negatively-charged clusters and crystals might also suggest that the electronic character of the sites may be important. Another factor is the ability of the NH3 molecule to act as a Lewis base. This basicity was determined to be important for the adsorption of ammonia on MgO surfaces.2 This basicity could also explain experiments on the relative reactivity of Si,* toward a m m ~ n i a . For ~ ~these ,~~ experiments, the structural analysis of some of the reactivity patterns suggests that the alkali fluoride crystal’s ability to adsorb ammonia is related to the crystal’s ability to accommodate the lone pair electrons in some kind of anion vacancy.
V. Conclusion We have shown that one can use recurring patterns in small nanocrystals, both patterns of defects and perfect crystallites, to provide another approach for examining surface adsorptiondesorption processes. In particular, for the reaction of Na,F,-I+ clusters with ammonia, we have determined that a basket-like defect facilitates the adsorption process and that perfect crystallites do not adsorb, except under relatively high partial pressures of ammonia. Clusters, II < 13, are more reactive than nanocrystals. The NaF ammonia system shows adsorptiondesorption equilibria at near-ambient temperatures, and the extent of reaction decreases with increasing temperature. Negatively-charged clusters are much less reactive. The KF nanocrystals show very similar reactivity patterns, suggesting that the structures for NaF crystallites are valid for the KF crystals. The LiF clusters have a very distinct reactivity pattern compared to those of the sodium and potassium fluoride crystallites, suggesting that LiF structures may be based on ringlike structures rather than cubic lattices.
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