J. Phys. Chem. 1980, 84, 3230-3232
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(7) P. A. Barrett, C. E. Dent, and R. P. Linstead, J. Chem. Soc., 1719 (1936). (8) H. A. Szymanski, "Interpreted Infrared Spectra",Vol. 2,PlenumPress, New York, 1964. (9) W. L. Bragg, "The Crystalline State", Vol. 1, 1933,p 189;C. C. Murdock, Phys. Rev., 31, 304 (1928);B. E. Warren, 2.Krisfallogr.,
99, 488 (1938). (10) Unpublished results. (11) T. Kobayashi, T. Ashida, N. Uyeda, E. Suito, and M. Kakudo, Boll. Chem. SOC. Jpn. 44, 2095 (1971). (12) Y. Saito, T. Kobayashi, N. Uyeda, and E. Suito, Bull. Insf. Chem. Res. Kyoto Unlv., 49, 256 (1971).
Chemisorption of Methanol on Magnesium Oxide. Observations by Carbon-I 3 NMR Ian D. Gay Department of Chemistry, Simon Fraser University, Burnaby, British Columbia V5A lS8,Canada (Received: December 27, 1979; In Final Form: June 18, 1980)
The adsorption of methanol on magnesium oxide is studied by high-resolution solid-state 13CNMR techniques. It is found that adsorption to coverages of less than about half a monolayer produces a rigidly bound methoxide-like species. Higher coverages produce in addition an isotropically rotating methanol species.
Carbon-13 NMR has become a nearly routine method for studying weakly adsorbed species on solid s~rfaces.l-~ In order to obtain useful results by conventional techniques, one must use adsorbed molecules that undergo isotropic rotation on a timescale of less than lo4 s, in order that the C-H dipolar interactions be averaged to small values. Such techniques are therefore not applicable to strongly chemisorbed species. Recent advances in highresolution solid-state NMR techniques4 have shown that this dipolar coupling may be suppressed by high-power irradiation of the protons, and high-resolution 13Cspectra obtained. The decoupling may readily be combined with cross polarizationH to produce a valuable increase in signal to noise. Surprisingly, these newer techniques have been little used in surface studies. There exist to date only two studies7i8in which results have been achieved at natural abundance, an equal number7v9with enriched compounds, and a few studies1&12of tightly bound aprotic molecules. One reason for this is that 13C shielding anisotropies can be rather large. This aggravates the signal-to-noise problem, and precludes surface studies on complicated molecules, where the necessity for polycrystalline samples would lead to intractably overlapping powder patterns. An obvious solution to this problem is the use of selectively enriched compounds. The utility of this approach has recently been demonstrated by SefcikSg In the present work we present the results of studies of methanol adsorbed on MgO, at natural isotopic abundance. The magnesium oxide was prepared from reagent grade magnesium hydroxide (Matheson Coleman and Bell). The hydroxide was decomposed under vacuum at 300 "C in a thin bed. This gave a product having a surface area of 250 m2/g. The initial oxide was then loaded into 12-mm NMR tubes and slowly heated to 500 "C in vacuo. The oxide samples had areas of 200-210 m2/g after this treatment. Measured amounts of methanol vapor were then allowed to adsorb from the gas phase, with the oxide at room temperature. The samples were then sealed off and allowed to stand for several days at room temperature before NMR measurements. 0022-3654/80/2084-3230$0 1.OO/O
TABLE I molecule 6 la 8 1IU HOCH, 73 10 KOCH, 78 18 Mg(OCH3 I* 74 10 Ca(OCH3l2 73 20 In parts per million downfield of liquid tetramethylsilane.
Samples of potassium, magnesium, and calcium methoxides were used for comparison. These were prepared by reaction of the appropriate metal with dried methanol, followed by heating in vacuo at a sufficient temperature to decompose s0lvates.~~9~~ These samples were kept sealed under vacuum. Spectra were measured at 15.08 MHz on a modified TT-14 spectrometer. Single-contact Hartmann-Hahn cross polarization4 was used, with rf field strengths of 70 kHz (yH1/27r) at both frequencies, and a contact time of 2 ms. The proton TIvalues of the adsorbed species lay in the range of 50-200 ms, permitting a relatively high data rate. Sample temperature during the measurements was 35 f 5 "C. The spectra were not changed on repetition after some months storage of the samples at room temperature. The solid methoxides were all found to give an axially symmetric shielding pattern, which could be well fitted by the theoretical powder pattern convolved with a Lorentzian line shape. The results of these computer fits are given in Table I, together with our measurement of solid methanol. The uncertainties in these data are estimated as f 2 ppm, based on the results of least-squares fitting and on reproducibility of different measurements on the same sample. In the case of magnesium methoxide, reproducibility was also checked on samples from different preparations, and agreement within the above limits obtained. Our result for solid methanol agrees with the values of 74 and 11 f 6 ppm found by Pines et al.16 Figure 1 shows the spectra obtained for the various adsorbed samples. The vertical lines in this figure represent the shielding components of magnesium methoxide, as given in Table I. At low coverage, it can be seen that 0 1980 American Chemical Society
The Journal of Physical Chemistty, Vol. 84, No. 24, 1980 3231
Chemisorption of MeOH on MgO
I
I
IO0
0 PPm
Flgure 1. Spectra of methanol adsorbed on magnesium oxide. Coverages are (,a) 3.0, (b) 4.9, (c) 6.0, {d) 7.4, and (e) 10.0 pmoi/m2. Each spectrum is average of 2 X 10 to 4 X I O 5 single contacts. Contact time 2.10 ms. Spectra have been normallzed for differing numbers of scaris before plotting, but other factors affecting intensity have not been closely controlled. Thus absolute intensities are only approximately CIDmparabb.
the adsorbed methanol produces a pattern which is approximately axially symmetric, with shielding components in agreement with the values of either methanol or magnesium methoxide, these two compounds not being distinguishable to the precision of the present measurements. As can be seen, however, when the methanol coverage exceeds 6 pmlol/m2 a new peak appears, and grows with increasing coverage. It is clear at the higher coverages that this is an isotropic peak superimposed upon the underlying powder pattern. The chemical shift of this peak is 49 i 1 ppm, which coincides with that of liquid methanol, within experiimental error. Our picture of methanol adsorption on MgO is then one of adsorption producing immobile species up to coverages of 5-6 pmol/m2, followed by production of an effectively isotropically rotating species at higher coverage. Because of the coincidence of the chemical shifts of methanol and magnesium miethoxide, it is difficult to be sure of the chemical nature of these species. The view which makes best chemical sense m d which is in general agreement with previous infrared i n v e s t i g a t i ~ n swould ~ ~ J ~ be that the immobile specieri is a surface methoxide, produced by dissociative chemisorption, and the mobile species a hydrogen-bonded methanol molecule. This molecule need not rotate isotr0pic:ally in the sense that all rotational diffusion coefficients are equal, but only in the sense that all directions in space are equally probable, and achievable in times short compared with the reciprocal of the chemical shift anisotropy, i.e., shorter than 1ms. The present experiment permits some conclusions about the motions of the nondissociated methanol. Since our apparatus implements alternating spin temperature reversal,16the observed signals arise only via cross polarization from protons. On the other hand, the observation of an isotropic spectrum means that the molecule is rotating. This combination of circumstances implies that either (a) the methanol molecules rotate at just the right rate (a few
kilohertz) to average the chemical shift anisotropy while not averaging the stronger C-H dipolar coupling to zero, (b) the molecules rotate but do not translate significantly, so that intermolecular cross polarization (as in solid adamantane) can take place, or (c) that cross polarization takes place via the C-H scalar coupling. The last is improbable, since it would require an implausibly precise setting of the Hartmann-Hahn matching condition, in view of the magnitude of the rf fields used. A few experiments with varying contact times showed that the cross-polarization time constants are about 0.5 and 0.8 ms, respectively, for the rigid and isotropic species. This indicates that scalar coupling is not important in the cross polarization. We observe no changes in the spectra up to temperatures of 70 "C, which tends to support mechanism b over a. Careful comparison of the low-coverage spectra with that of magnesium methoxide suggests that the surface species do not have a spectrum which corresponds accurately to a rigid axially symmetric powder pattern. The upper and lower limits of shielding seem to agree with the methoxide values, but there seems a deficit of intensity at both the low-field peak arid at the high-field shoulder. A possibl+, explanation might be rotation about the Mg-0 bond of a nonlinear Mg-O-CH3 species. A few calculations were carried out by the method of S p i e d 7for a model in which the M g - 0 4 angle is tetrahedral, and 90' hops occur about the Mg-0 axis, with the unique axis of chemical shielding assumed to be along the 0-C bond. With a hopping frequency of 200-400 Hz, this model produces effects of the right qualitative nature, but does not accurately reproduce the experimental spectra. Possibly a more elaborate motitonal model would do so, but this does not seem warranted with the present data. An alternative explanation might be the presence of more than one type of surface methoxide, leading to overlapping powder patterns. It is unlikely that the effect is caused by angular dependence of the cross-polarization time. If rapid rotation of the CH3 group is assumed, cross polarization should be least effective when the 0-C bond makes the magic angle with the magnetic field. This would lead to a lack of intensity at 6 = 1/3611+ 2/a6L, which is not observed. Our conclusions regarding the nature of the adsorbed methanol are qualitatively in agreement with infrared results of previous workers.18J9 There is, however, a quantitative disagreement. Tench et al.18 propose that the first adsorbed layer contains about 15.6 pmol/m2 of methoxide together with 3.1 pmol/m2 of hydrogen-bonded methanol. Our results show that the mobile species appears after a coverage of only 5-6 pmol/m2 of chemisorbed species. Our samples might contain as much as 3 pmol/m2 of surface hydroxyls.20 If this is assumed to block sites for methanol chemisorption, a total chemisorbed coverage of 8-9 pmol/m2 might be expected on a clean surface. This would correspond to about 50% coverage of surface Mg2+ ions by methoxide. (The concentration of Mg2+on MgO(100) is 18.6 pmol/m2.) This seems reasonable, as the Mg-Mg distance on this plane is 3.0 A, and no plausible assumption about the size of methoxide would permit adsorptionon adjacent Mg2+ions. A 4 2 X 2) arrangment, involving next-nearest Mg2+ ions seems geometrically possible, and would produce 50% chemisorbed coverage. Conclusions of this type are more reliable with NMR than with infrared, first because the samples are of a form more amenable to accurate adsorption measurements, and secondly because infrared workers must often make the dubious assumption of invariance of extinction coefficient upon adsorption. Indeed, a recent paper12 exhibits the presence of a large adsorbed population which was com-
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J. Phys. Chem. 1980, 84, 3232-3238
pletely undetected by infrared. It is of interest to consider the data acquisition time necessary to achieve a reasonable signal-to-noiseratio. The present spectra required 8-16 h of spectrometer time each. This seems forbidding, but, in fact, the experiment was by no means optimized for sensitivity. The biggest improvement would be achieved by increasing sample density. The present samples were studied as loose powders having a bulk density of 0.6 g/cm3. Since the crystal density of MgO is 3.6 g/cm3, pressed pellets of at least twice the density could be used. In addition, it would be certainly feasible to use a spectrometer of twice the field, to use twice the sample volume, and to fit quadrature detection or single-sideband filtering to the spectrometer. Each of these expedients would produce an increase of d2 in sensitivity. Thus a sensitivity improvement of 4 4 2 could be achieved, leading to acquisition times of less than 0.5-h per spectrum. In addition, some further advantages might be gained by operating at lower temperatures, if relaxation phenomena permitted. References and Notes (1) I. D. Gay and S. H. Liang, J . Catal., 44, 306 (1976). (2) J. F. Kriz and I. D. Gay, J. Phys. Chem., 80, 2951 (1976).
(3) D. Michei, W. Meiler, H. Pfelffer, and H. J. Rauscher, J. Mol. Catal., 5, 263 (1979). (4) A. Pines, M. 0. Gibby, and J. S. Waugh, J. Chem. Phys., 59, 569 (1973). (5) S. R. Hartmann and E. L. Hahn, Phys. Rev., 128, 2042 (1962). (6) F. M. Lurle and C. P. Slichter, Phys. Rev. A , 133, 1108 (1964). (7) S. Kapian. H. A. Resing, and J. S. Waugh, J. Chem. Phys., 59 5681 (1973). (8) J. J. Chang, A. Pines, J. J. Fripiat, and H. A. Resing, Surf. Sci., 47,
661 (1975). (9) M. D. Sefclk, J. Am. Chem. Soc., 101, 2164 (1979). (10) E. 0.Stebkal, J. Schaefer, J. M. S. Henis, and M. K. Tripodi, J. Chem. Phys., 61, 2351 (1974). (11) M. D. Sefcik, J. Schaefer, and E. 0. Stejskal, ACS Symp. Ser., 34, 109 (1976). (12) T. M.'Dun&n, J. T. Yates, Jr., and R. W. Vaughan, J. Chem. Phys., 71, 3129 (1979). (13) N. J. Turowa, 6.A. Popowski, and A. W. Nowosoiowa, Z . Anorg. Allg. Chem., 385, 100 (1969). (14) E. Weiss, Helv. Chim. Acta. 46, 2051 (1903). (15) A. Pines, M. G. Gibby, and J. S. Waugh, Chem.'Phys. Leff ., 15, 373 (1972). (16) E. 0.Stejskal and J. Schaefer, J. Magn. Reson., 18, 560 (1975). (17) H. W. Spiess, NMR: Basic Prin. Prog., 15, 55 (1978). (18) A. J. Tench, D. Giles, and J. F. J. Kibblewhite, Trans. Faraday Soc., 67, 854 (1971). (19) R. 0.Kaaei and R. 0. Greenler. J. Chem. Phvs.. 49. 1638 (19681. (20) P. J. Anderson, R. F. Horiock, and J. F. Oliver, -Tram. Faraday Soc:,
61,2754 (1965).
Luminescent Photoelectrochemical Cells. 4. Electroluminescent Properties of Undoped and Tellurium-Doped Cadmium Sulfide Electrodes Hoiger H. Streckert, Bradley R. Karas, David J. Morano, and Arthur B. Ellis' Department of Chemlstty, Unlverslty of Wisconsin-Madison, Madison, Wisconsin 53706 (Received: March 19, 1980; In Final Form: Ju/y 16, 1980)
Single-crystal,n-type, undoped CdS and Te-doped CdS (100 and 1000 ppm CdS:Te) electrodes exhibit electroluminescence (EL) in aqueous, alkaline peroxydisulfate electrolyte. Addition of Te to CdS introducesintraband gap states which drastically alter the EL spectral distribution. The EL spectrum of undoped CdS consists of a sharp band near the band gap energy (-2.4 eV) with A, -510 nm and a weaker, broader band at lower energy; the EL spectra of 100 and 1000 ppm CdS:Te also exhibit a sharp band at -510 nm but are dominated -600 and -620 nm, respectively. Photoluminescence (PL) spectra, by broad bands with uncorrected A,, obtained with 457.9-nm excitation, aid in the assignments of the emissive transitions; PL spectra are generally similar to their EL counterparts. The potential dependence of the EL spectra was examined between --1.2 V (onset) and -2.0 V vs. SCE. Changes in potential can affect both the relative and absolute intensities of the bands present in the EL spectra. Lower-limit,order-of-magnitudeestimates of instantaneous EL efficiency, dEL,have been made, Under steady-stateconditions 4m is ?lo4 for undoped CdS and for CdS:Te. Results are discussed in terms of interfacial charge-transferprocesses and the excited-state manifolds of CdS and CdSTe.
Introduction An understanding of electrochemistry at semiconductor electrodes requires knowledge of the role that intraband gap states play in mediating interfacial charge transfer. Reference to surface states is especially prevalent in the photoelectrochemicalcell (PEC) 1iterature.l The existence of intraband gap states has been supported in part by electroluminescence (EL) studies. Electrolyte species capable of hole injection into the valence bands of n-type, semiconducting Ti02,2SrTi03,3CdS,4 GaP,4pSZn0,4and GaAs4i6yielded luminescence at subband gap energies. Since the valence band edges of these materials are generally at very positive potentials, valence band hole injection requires a potent oxidizing species. Aqueous alkaline peroxydisulfate (OH-/S20,"-) electrolyte has been the medium of choice for observing the effect; the sequence of half-reactions occurring at the n-type semiconductor 0022-3654/80/2084-3232$0 1 .OO/O
-
electrode is reported to be the f~llowing:~
+ eCB- sot- + sod-. SO4-. S042-+ hVB+
S2Oa2-
-
(1) (2)
The symbol ecB-represents an electron in or near (surface state or trap, e.g.) the conduction band; similarly, hvB+ denotes a hole in or near the valence band. Once h+ has been injected by the strongly oxidizing sulfate radical, radiative recombination with electrons can occur. Estimates for the redox potentials of eq 1and 2 are 1+0.6 and L+3.4 V vs. NHE, respe~tively.~~ Establishing the presence of intraband gap states by EL-induced, subband gap emission demonstrates the utility of the technique. Extending EL studies beyond this stage would be greatly facilitated by knowledge of the energetic location and properties of the intraband gap 0 1980 American Chemical Society
