1086
J. Phys. Chem. 1984, 88, 1086-1089
relative to C F 2 0 and a further 1.6 kcal mol-' increase for CF3CF2COF relative to CF3COF. As the chain length of the perfluoroalkyl group increases correspondingly smaller increases would be expected for each additional CFz group. After accounting for the anticipated increases due to the increased inductive stabilization, the fluoride affinity of perfluoroglutaryl fluoride FCO(CF2)3COFis still roughly 6 kcal mol-' more than anticipated. This increased stabilization may be due to a number of factors: either (i) the interaction of the FCO dipole moment with the negative charge of a CFzO- group as in V, (ii) a genuine
Lewis acids with more than one acidic site. The differences in results for F and C1- have been interpreted in terms of the relative tendencies of these two species to undergo covalent or electrostatic binding. These results also have implications for the design of anioncomplexing reagents analogous to crown ethers for cations. The very strong binding energies of F and C1- to fluoroethers suggest that cyclic complexes such as VI11 may be efficient agents for ,F
p-cFz \ /
I
I
VI
V
VI11
VI I
chelation interaction in which F is bound equally strongly by two equivalent FCO groups as in VI, or (iii) an intramolecular base-induced cyclizatibn reaction forming an adduct such as VII. At present we have no basis for distinguishing which of these three situations may be corrdct; however, if chelation does occur, as in VI, the relatively small increase in stability compared to the total single-site fluoride binding energy would be a further example of weakening of a first interaction by formation of second.
Conclusion The results presented in this work demonstrate that gas-phase halide ions may undergo chelation interactions with Bronsted and
F
solubilizing ionic salts in organic media and for generation of naked cations.
Acknowledgment. Financiai support of this work by the Natural Sciences and Engineering Research Council of Canada is gratefully acknowledged. Registry No. S02F2,2699-79-8; SOF2, 7783-42-8;PF,, 7783-55-3; SO2, 7446-09-5;CF,COF, 354-34-7; CH,SiF,, 373-74-0; SiF4, 778361-1; (CZH5)3B, 97-94-9; t-C,H,OH, 75-65-0; CzFsH, 354-33-6; HC1, 7647-01-0;CF,OCFzH, 3822-68-2; (CFzH)O, 1691-17-4;COF2, 35350-4; FCO(CFZ)pCOF, 678-78-4; CF,CF&OF, 422-61-7; (CF&zCO, 684-16-2;(CFzH)zCO,360-52-1; CFZHCOCF,, 431-71-0;C1-, 1688700-6; F, 16984-48-8.
On the Existence of Na- In Llquld Ammonia Steven G. Bratsch and J. J. Lagowski* Department of Chemistry, The University of Texas, Austin, Texas 78712 (Received: June 8, 1983)
The question of the existence of alkali-metal anions in liquid ammonia is examined from a thermodynamic viewpoint using a simple semiempirical model. A fit of the Born equation for ionic solvation to the available data indicates that Na- (sodide) is the alkali-metal anion most likely to be stable under standard conditions (1 m activity) in liquid ammonia at 298 K. The lack of an observed spectral peak for sodide ion in Na/NH3 solutions is discussed in terms of activity-to-concentrationconversions and dilution effects. Possible methods of Na- detection in Na/NH3 solutions are offered.
Thermodynamic Considerations Alkali-metal anions M- are known in a variety of nonaqueous solvents such as cyclic ethers and organic amines.l*z Their existence in metal-ammonia solutions has been p o ~ t u l a t e dbut ~ * ~not conclusively p r ~ v e n . 'Thermodynamic ~~~~~ stability of alkali-metal anions in a given solvent at a specified temperature requires that the Gibbs free energy changes for the following processes both be positive: process 1: process 2:
-
M-(solv) M-(solv)
+ e-(solv) M+(solv) + 2e-(solv) M(c)
The equilibrium for process 2 lies far to the right in dilute liquid (1) Dye, J. L. J. Chem. Educ. 1977, 54, 332. (2) Dye, J. L.Angew. Chem., Int. Ed. Engl. 1979, 18, 587. (3) Golden, S.;Guttman, C.; Tuttle, T. R . , Jr. J. Am. Chem. SOC.1965, 87, 135. (4) Golden, S.;Guttman, C.; Tuttle, T. R., Jr. J . Chem. Phys. 1966, 44, 379 1. ( 5 ) Dye, J. L.; Andrews, C. W.; Mathews, S . E. J. Phys. Chem. 1975, 79, 3065. (6) Dye, J. L.; Andrews, C. W.; Ceraso, J. M. J. Phys. Chem. 1975, 79, 3076.
0022-3654/84/2088-1086$01.50/0
ammonia solutions presumably because ammonia strongly solvates the cations M+.2 Le Chdtelier's principle predicts that the equilibrium will shift toward the left in more concentrated solutions. The standard Gibbs free energy of formation of a solvated anion M- may be derived from the following thermochemical cycle: AGfo(H+,g) step 1: 0.5H2(g) H+(g) e-
-+ - + -+ -+
step 2: M(std state)
e-
M-(g)
AGfo(M-,g)
step 3:
H+(g)
H+(solv)
AG,oI,o(H+)
step 4:
M-(g)
M-(solv)
AGsolvo(M-)
sum: M(std state)
0.5H,(g) M-(solv)
H+(solv)
AGfO(M-,solv)
AGO values for steps 1 and 2 are given in Table I. The standard Gibbs free energy of solvation of the proton, AG,,I,o(H+) (step 3), has been estimated for several solvents at 298 K; hydration (solvent HzO) and ammoniation (solvent NH3) values used in this paper are AGhrdo(H+)= -1066 kJ mol-'' and AG,,O(H+) = (7) Conway, B. E. J. Solution Chem. 1978, 7 , 721.
0 1984 American Chemical Society
The Journal of Physical Chemistry, Vol. 88, No. 6, 1984 1087
Existence of Na- in Liquid Ammonia TABLE I: Thermodynamics of Selected Gas-Phase Ions at 298 K
M
z
e F
1-
C1 Br I
H Li Na K Rb
Cs H
111I1111111+
AH;-
AH:-
(M,g)a
(Mz',g)C
79.39 121.302 111.86 106.762 217.997 159.37b 107.3Zb 89.24b 80.877b 76.065b 217.997
-254.8 -233.7 -218.9 -194.8 139.0 93.4 48.4 34.7 27.8 24.4 1536.2
S"(M,std state)e
So(Mz',g)g
AG;-~ (Mz',g)
101.342f 111.482f 76.105f 58.O7Of 65.2855 29.12 51.30 64.68 76.78 85.23 65.285f
20.87h 145.46 153.24 163.38 169.15 108.85 132.91 147.84 154.46 164.22 169.72 108.84
O.Od -261.7 -239.9 -238.7 -221.7 132.2 68.7 25.8 14.2 8.0 5.4 1517.0
O.Od
a In kJ mol-' ; reference 11, except as noted. Reference 27; reference 11 gives no value. In kJ mol-' ; calculated by the folloming: aHf"(M-,g) = aHf"(M,g) - EA (0 K) - 2.5RT; aHf"(M+,g)= aHf"(M,g) + IP(0 K ) 2.5RT. EA from ref 28; IP from ref 29. d Arbitrary convention. e In J K-' mol-' ; reference 11. Value refers to 0.5M2. In J K-' mol-' ; calcuS "In=(at wt) + R In ( W +1) + 108.745, where lated ~ Y ' ~ ~1.5R J - magnetic spin quantum number = 0 except J(e-) = 0.5. Value agrees with ref 15d. In kJ mol-'; S"(e-) accounted for in AG," = AH^' - TAS~".
+
-1 163 kJ respectively. These values lead to the following relationships a t 298 K:
AGfo(M-,aq) = AGfo(M-,g)
+ AGhydo(M-)+ 451 kJ mol-'
AGfo(M-,am) = AGfo(M-,g)
+ AGamo(M-)+ 354 kJ mol-'
(1)
(2) If the AGO value for step 4 can be estimated, AGfo(M-,solv) can be calculated.
Calculation of Gibbs Free Energies in Liquid Ammonia at 298 K The following modification of the Born equation for ionic solvation9 (see also ref 10) has been found useful:
+
AGsolvo(MZ+)= klz2/(r k 2 ) (3) where z is the ionic charge, r is the crystal ionic radius of MZ+, and k l and k2 are empirically determined by curve fitting. For the four common halides, the known AGf0(M-,aq)" and AGfo(M-,g) (Table I) may be used to back-calculate AGhyd0(M-) via eq 1. If Pauling ionic radiii2 are used, the best fit to eq 3 is AGhydo(M-)= -56.1/(r
- 0.017) f 1.3 kJ mol-'
(4)
where AGhydo(M-)is at 298 K and r is expressed in nanometers (Figure 1). Matsuda and NatoyaI3 have shown that AGsolvoof cations are roughly proportional between solvents, calculating that the solvation of cations in liquid ammonia is about 1.1 times as strong as in water at 298 K. It seems logical to assume a proportionality for anions as well, although the constant of proportionality need not be the same as that for cations. In terms of eq 3, this assumption requires a constant value for k2, which is found from the presumably accurate water expression (eq 4) to be -0.017 nm when Pauling anionic radii are used. The datai4 (adjusted to ref 11) for F, C1-, Br-, and I- in liquid ammonia at 298 K give AGamo(M-) = -50.3/(r
- 0.017)
f 8.7 kJ mol-'
(5)
(8) Izmailov, N. A. Dokl. Akad. Nauk. S S S R 1963, 149, 288. (9) Born, M. Z . Phys. 1920, 1 , 45. (10) Latimer, W. M.; Pitzer, K. S.; Slansky, C. M. J. Chem. Phys. 1939, 7., 108. ~(1 1) CODATA Recommended Key Values for Thermodynamics, J. Chem. Thermodyn. 1978, 10, 903. (12) Pauling, L. 'The Nature of the Chemical Bond", 3rd ed.; Cornell University Press: Ithaca, NY, 1960; p 514. (13) Matsuda, A.; Natoya, R. J . Res. Inst. Catal., Hokkaido Univ. 1981, -29. ~1 5,1
(14) Latimer, W. M.; Jolly, W. L. J . Am. Chem. SOC.1953, 75, 4147.
0
-500 0
I.0
2.0
3.0
,
U.0
I
S.0
I
6.0
7.0
8.0
Figure 1. Correlation of AGmlVo(M-)with rM-1 AGfo(M-,g) from Table I; AGfo(M-,aq) (M = F, C1, Br, I) from ref 11; AGo(e-,aq) from ref 16a; AGhydo(M-)from eq 1; AGfo(M-,am) (M = F, CI,Br, I, e) from Table 11; AGamo(M-) from eq 2; rM- (nm) from Table 11. Solid circles are experimental; open circles are calculated.
This is illustrated in Figure 1. Taken together, eq 4 and 5 indicate that the solvation of uninegative anions in liquid ammonia is about 0.9 times as strong as in water at 298 K. The proportionality constants for cations (1.1) and anions (0.9) are consistent with the increased Lewis basicity and decreased Lewis acidity of ammonia relative to water. The values of kl in eq 3 are predicted by the Born equation for ionic solvation9 to be -68.6 and -65.4 for water and liquid ammonia, respectively. However, Born's model does not involve readily available crystal ionic radii but ill-defined "effective" radii of ions in the solvent. In addition, the dielectric constant of the solvent is assumed to be uniform, even in the vicinity of the ions.15a In view of the problems associated with the Born equation, the level of agreement between the predicted values of k, and the empirical values given in eq 4 and 5 is about as expected. The estimated AGfo(e-,aq) = 267 kJ mol-' may be combined with eq 1 and 4 to give AGhydo(e-)= -184 kJ mol-' and re-= 0.322 nm. In a like manner, the selected AGfO(e-,am)= 186 kJ mol-' 17a may be combined with eq 2 and 5 to give AGamo(e-)= -168 kJ mol-' and re- = 0.316 nm. Thus, we estimate the crystal ionic radius for e- to be about 0.32 nm. This value may be compared with Peer's spectroscopically derived estimate of 0.345 nm17band also with estimates of the hydrated ionic radius, 0.25-0.30 nm.16b Plambeck18has proposed an alternative modification of the Born equation wherein AGmIvois calculated as a power series in r". The approach used in the present paper has the advantage of requiring only one parameter change (k, in eq 3) when a different solvent is considered. Table I1 lists calculated and experimental Gibbs free energies for selected ions in NH3 at 298 K. The sources of the AGO data and ionic radii are also indicated in Table 11. The thermodynamic calculations for the alkali-metal anions M- are directed toward the evaluation of the standard-state properties of these ions, where (by definition) the only interactions of the solute ions are those with the solvent.'9a Thus, the predictions do not take into account the possible existence of the ion triplets e-M+e-. The latter species cannot be treated by the present method because both r and AGfo(e-M+e-,g) are unknown. It is possible that the ion triplets e-M+e- are thermodynamically more stable than the genuine alkali-metal anions M- in liquid a m m ~ n i a . ' * ~ > ~ AGO values (kJ mol-') for processes 1 and 2 identified in the preceding section may be derived from the calculated AGfo(M-,am) and the experimental AGfo(M+,am) and (15) (a) Johnson, D. A. "Some Thermodynamic Aspects of Inorganic Chemistry", 2nd ed.; Cambridge University Press: New York, 1982; pp 122-124. (b) Ibid., p 251. (c) Ibid., p 8. (d) Ibid., p 25. (16) Hart, E. J.; Anbar, M. "The Hydrated Electron"; Wiley-Interscience: New York, 1970) (a) p. 63, (b) p. 62. (17) (a) Peer, W. J. Ph.D. Dissertation, The University of Texas, Austin, TX, 1979; p 260. (b) Ibid., p 315. (c) Ibid., p 313. (d) Ibid., pp 301-315. (18) Plambeck, J. A. Can. J . Chem. 1969, 47, 1401. (19) (a) Day, M. C., Jr.; Selbin, J. "Theoretical Inorganic Chemistry", 2nd ed.; Van Nostrand-Reinhold: New York, 1969; p 530. (b) Ibid., p 528.
1088 The Journal of Physical Chemistry, Vol. 88, No. 6, 1984
Bratsch and Lagowski
TABLE 11: Gibbs Free Energies for Selected Ions in Liquid Ammonia at 298 K
M
rM-a
e F
0.32 0.136 0.181 0.1 95 0.216 0.208 0.23 0.263 0.308 0.316 0.352
c1 Br I H Li Na
K Rb cs
AG fO(M-,am),,lcdc
AG.,O(M-)~ -166 -422.7 -306.7 -282.6 -252.8 -263.4 -236 - 204.5 -172.9 -168.2 -150.1
188 -330.4 -192.6 -167.3 -120.5 222.8 187 175.3 195.3 193.8 209.3
AG
fO-
(M',am)expt
ACf"(M+,am)expt
1
2
Oe -225e -183=jf -198e9f -198e*f -205e,f
-37h -1 11 -9 -8 - 23
149 -40 14 -21 -20 -42
AGde/
186" -343e - 185e -168e -121e
a In nm, re- back-calculated from thermodynamic data in H,O and NH,; halide and hydride radii from ref 12; r ~ i estimated; other radii from ref 17c; values claimed to be coiisistent with Pauling radii.I2 In kJ mol-'; equation 5. In kJ mol-' ; equation 2. Reference 17a. e Reference 14; adjusted to ref 11. f From AC,,~,"(MCl) given in ref 30 and/or in ref 31, using aGf"(MC1,c) from ref 27 adjusted to ref 11. g AGO (kJ mol-') for anion decomposition processes 1 and 2. For the process H- + 0.5H2 + e'.
"
TABLE 111: Predicted Compositions of Na/NH, Solutions at 298 K activity,b m
concn, m tot [Nal, m 0.0010 0.0030 0.010 0.030 0.10 0.30 1.0
1.5 x 10-7 2.9 X loT6 5.6 x 10-5 6.2 x 10-4 0.0054 0.028 0.15
--
Na'
e-
yia
Na-
Na+
e-
[Na-] ,c %
0.0010 0.0030 0.0099 0.029 0.095 0.27 0.85
0.0010 0.0030 0.0099 0.029 0.089 0.24 0.70
0.77 0.65 0.48 0.32 0.17 0.085 0.038
1.2 x 10-7 1.9 X 2.7 X 2.0 x 10-4 9.2 x 10-4 0.0024 0.0057
7.7 x 10-4 0.0020 0.0048 0.0093 0.016 0.023 0.032
7.7 x 10-4 0.0020 0.0048 0.0093 0.015 0.020 0.027
0.015 0.097 0.56 2.1 5.4 9.3 15
Na-
These values are highly speculative: There is n o thermodynamic method of assigning activities to single ions, and the taba Equation 6. ulated values are calculated by assuming yion = yt for a 1: 1 electrolyte. Percent of total sodium present as Na-.
AGt(e-,am); these are given in Table 11. Of all the alkali-metal anions, only Na- (sodide) is calculated to have a positive AGO for both processes 1 and 2 so that, a t standard conditions (Le. 1 m activity), Na- might be expected to exist in significant quantities in liquid ammonia.
Sodium in ammonia may be viewed as a solution of a 1 : l electrolyte containing Na+ as the cation and varying proportions of e- and Na- as the anion, with the ionic strength p equal to mNa+. For an unsaturated Na/NH, solution the appropriate equilibrium is
Correction for Activity Effects: Predicted Compositions of Sodium/Ammonia Solutions at 298 K Mean activity coefficients y+ in liquid ammonia are known to be extremely low, even for "strong" 1:l electrolytes. The experimental values of yi for NHpCl, NaCl, and KC1203' are related to the ionic strength p by
NaNa+ 2e(7) for which AGO is calculated to be +14 kJ mol-' and K = 0.004. Using this equilibrium constant and the concentration/activity conversion (eq 6), one may calculate the compositions of various Na/NH, solutions at 298 K (Table 111). The percent sodium present as sodide ion is predicted to become significant only above about 0.01 m total N a concentration.
log y* = -3.678p1/'/(1
+ 1.51p'/2)
(6)
with a mean absolute deviation of 2.1%. Equation 6 provides a uniform means of estimation of the mean activity coefficient for a solution of a 1:l electrolyte in liquid ammonia up to at least 1.O m (eq 6 is actually useful in NH4Cl solutions up to about 5 m ) . It is possible that ion pairs and higher aggregates are important in concentrated solutions of electrolytes in liquid ammonia. However, from a thermodynamic point of view, it is not necessary to relate the activity coefficient to any particular mechanistic behavior. 19b*22 (20) Ritchey, H. W.; Hunt, H. J. Phys. Chem. 1939, 43, 407. (21) Sedlet, J.; De Vries, T. J. Am. Chem. SOC.1951, 73, 5808. (22) Jolly, W. L. J. Chem. Educ. 1956, 33, 512. (23) Farrow, M. M.; Burnham, R. K.; Eyring, E. M. Appl. Phys. Lert. 1978, 33, 735. (24) Lloyd, L. B.; Riseman, S.M.; Burnham, R. K.; Eyring, E. M.; Farrow,M. M. Rev. Sci. Instrum. 1980, 51, 1488. (25) Teng, Y.C.; Royce, B. S. H.Appl. Opr. 1982, 21, 77. (26) Ceraso, J. M.; Dye, J. L. J . Chem. Phys. 1974, 61, 1585. (27) NBS Technical Note 270-8; National Bureau of Standards: Washington, DC, 1981. (28) Hotop, H.; Lineberger, W. C. J. Phys. Chem. Re$ Data 1975,4, 539. (29) Moore, C. E. NSRDS-NBS 34; National Bureau of Standards: Washington, DC, 1970. (30) Jolly, W. L. Chem. Reu. 1952, 50, 351.
+
Uncertainties Although the above calculations appear promising, Na- has not yet been identified in liquid a m m ~ n i a . ' The - ~ ~discrepancy ~~~ may be traceable to one or more of the following: 1. The electrostatic model for Gibbs free energies of solvation (eq 3) may be too simple for general quantitative predictions. Neglect of nonelectrostatic contributions (e.g. polarization effects) may lead to serious errors. 2. The ionic radii used may be unreliable. Peer's alkali-metal anionic radii"" have been used because they are believed to be most consistent with Pauling radii. However, Peer has derived re- in the same way (from spectroscopic correlations) as for the alkali-metal anionic but his result (0.345 nm) is not compatible with AGmk0(e-). DyeZhas derived alkali-metal anionic radii by assuming that the bond lengths in the hypothetical ionic solids M+M- are the same as in the metals M. Dye's radii are about 0.01 nm larger than Peer's values. For example, Dye estimates rNa-= 0.272 nm, which leads to AGfo(Na-,am) = 182.5 kJ mol-', about 7 kJ mol-' more positive (less stable) than the value in Table 11. 3. Because the alkali-metal anions are probably larger than the halide ions, the predictions for standard Gibbs free energies (31) Strong, J.; Tuttle, T. R., Jr. J . Phys. Chem. 1973, 77, 533.
J. Phys. Chem. 1984,88, 1089-1094
1089
Peer’s method can be used to predict E,,,(Na-,am) = 1.85 f 0.08 eV or A,, = 670 f 30 nm. Sodium/ammonia solutions are intensely absorbing at total N a concentration greater than 0.1 m, and Fourier transform photoacoustic s p e c t r o ~ c o p y may ~~-~~ be the only viable technique for Na- detection in these systems. 3. 23Na N M R in liquid ammonia’.2s5,6,26 may prove to be another useful detection technique. The method has been used by Dye in various solvents to differentiate between Na- and eNa+e-. Experiments are being constructed in our laboratory to investigate this possibility.
of solvation of metal anions are all extrapolations. 4. The concentration/activity conversion (eq 6) may become increasingly inappropriate as concentrated M/NH3 solutions begin to experience deviations from “strong” electrolyte behavior. Perhaps the only statement one may make with confidence is that Na- is probably the least unstable alkali-metal anion in liquid ammonia at 298 K; that is, given the proper experimental conditions, Na- stands the greatest chance of being detected.
Possible Detection of Na- in Liquid Ammonia 1. A detailed study of the physical and chemical properties of Na/NH, solutions may reveal anomalies with respect to other alkali-metal/”, solutions, which may be best explained in terms of Na-. 2. Peer’7d has noted that charge transfer to solvent (CTTS) absorption band maxima are linearly related between solvents.
Acknowledgment. We acknowledge the generous financial support of the Robert A. Welch Foundation (Grant No. F081). The reviewers’ comments on the original manuscript have also been helpful. Registry No. Na-, 19181-13-6; NH3, 7664-41-7; Na, 7440-23-5.
Formation of Vertically Oriented Aromatic Molecules Chemisorbed on Platinum Electrodes: The Effect of Surface Pretreatment with Flat Oriented Intermediates Manuel P. Soriaga* and Arthur T. Hubbard* Department of Chemistry, University of California, Santa Barbara, California 93106 (Received: June 10, 1983; In Final Form: August 8, 1983)
The adsorption of aromatic compounds on Pt electrodes, pretreated with a layer of flat oriented intermediates at fractional or full coverages, has been studied as a function of concentration. Measurements of packing densities were based on thin-layer electrochemical methods. Four aromatic compounds, previously shown to adsorb on clean electrodes in flat and edgewise (vertical) orientations depending upon the adsorbate concentration, were studied: hydroquinone (1) 1,4-naphthohydroquinone (2), 2,3-dimethylhydroquinone (3), and 2,5-dimethylhydroquinone(4). The same edge orientations formed on the clean electrode were also obtained on the pretreated surface, but the adsorption profiles (total packing density vs. concentration) of 1-3 showed that complete formation of the edge orientation was retarded (severely for 1 and 2 slightly for 3) when the electrode was pretreated with a full monolayer of flat oriented species; the same retardation was observed even when the electrode was pretreated with less than half a monolayer. When the electrode was precoated with a layer of aromatic molecules in transition between flat and vertical orientation, the adsorption profile was remarkably similar to those for clean electrodes. The adsorption profile of 4 (which displays hindered reorientation due to the blocking effect of the methyl groups) was unaffected by surface pretreatment. These results suggest that (i) adsorption of 1-3 at concentrations above 1 mM on clean Pt leads primarily to direct attachment of edge-oriented species; (ii) for 1-3, complete reorientation from flat to vertical structures is less facile than direct adsorption in the vertical orientation; (iii) adsorption of 4 in the edge orientation involves formation of flat oriented species as an intermediate step; (iv) adsorption of aromatic molecules at high concentrations on a sparsely pretreated surface involves completion of the flat oriented layer prior to formation of edge-bonded species; and (v) reorientation of nearly isolated flat adsorbed intermediates is at least as difficult as that of closepacked molecules. Similarities and differences between results from solution and vacuum studies are discussed.
Introduction The orientation or mode of attachment of aromatic compounds adsorbed on platinum surfaces has been the subject of numerous investigations.’-28 The first systematic studies on the chemi(1) Soriaga, M. P.; Hubbard, A. T. J . Am. Chem. SOC.1982, 104, 2735. (2) Soriaga, M. P.; Hubbard, A. T. J. Am. Chem. SOC.1982, 104, 2742. (3) Soriaga, M. P.; Hubbard, A. T. J . Am. Chem. SOC.1982,104, 3937. (4) Soriaga, M. P.; Wilson, P. H.; Hubbard, A. T.; Benton, C. S. J. Electroanal. Chem. 1982,142, 317. ( 5 ) Chia, V. K. F.;Soriaga, M. P.; Hubbard, A. T.; Anderson, S. E. J . Phys. Chem. 1983,87, 232. (6) Soriaga, M. P.; White, J. H.; Hubbard, A. T. J . Phys. Chem. 1983, 87. 3048. (7) Soriaga, M. P.; Stickney, J. L.; Hubbard, A. T. J . Electroanal. Chem. 1983,144, 207. ( 8 ) Soriaga, M. P.; Stickney, J. L.; Hubbard, A. T. J. Mol. Catal. 1983, 21, 211. (9) Soriaga, M. P.; Hubbard, A. T. J . Phys. Chem., in press. (10) Soriaga,M. P.; Hubbard, A. T. J. Electroanal. Chem. 1983,159, 101.
0022-3654/84/2088-1089$01.50/0
sorption of b e n ~ e n e , ’ ~based - ’ ~ on a wide variety of experimental techniques such as radiotracer methods and deuterium exchange (11) Stickney, J. L.; Soriaga, M. P.; Hubbard, A. T.; Anderson, S. E. J . Electroanal. Chem. 1981, 125, 73. (12) Moyes, R. B.; Wells, P. B. Adu. Catal. 1975, 23, 121. (13) Bond, G . C. “Catalysis by Metals”; Academic Press: New York, 1962. (14) Garnett, J. L.; Sollich-Baumgartner, W. A. Adu. Cutul. 1966,16, 95. (15) Palazov, A. J . Catal. 1973, 30, 13. (16) Primet, M.; Basset, J. M.; Mathieu, M. V.; Prettre, M. J . Catal. 1973, 29, 213. (17) Haaland, D. M. Surf. Sci. 1981, 102, 405. (18) Haaland, D. M. Surf. Sei. 1981, 111, 555. (19) Fischer, T. E.; Kelemen, S. R.; Bonzel, H. P. Surf. Sei. 1977.64, 157. (20) Richardson, N. V.; Palmer, N. R. Surf. Sei. 1982, 114, L1. (21) Netzer, F. P.; Matthew, J. A. D. Solid State Commun. 1979,29, 209. (22) Lehwald, S.; Ibach, H.; Demuth, J. E. Surf. Sci. 1978, 78, 577. (23) Gland, J. L.; Somorjai, G. A. Adu. Colloid Interface Sei. 1976,5, 205. (24) Stair, P. C.; Somorjai, G. A. J . Chem. Phys. 1977, 67, 4361. (25) Gavezzotti, A,; Simonetta, M. Surf. Sei. 1982, 116, L207.
0 1984 American Chemical Society
