3838
ROBERTW. MARSHALL AND GEORGE H. NANCOLLAS
ization of nuclei by relaxation processes in free radicals (CIDNP) as developed so far2*J1is very probably now inadequate as an explanation of all features nuclear polarization of reaction products in alkyllithiumalkyl halide and other reactions where the nmr multiplet effect occurs. The CIDNP concept is still able to explain most of the polarization phenomena in reaction products of phenyl and substituted phenyl radicals rather satisfactorily, and it seems too early to reject it completely. However, during alkyl radical reactions, another polarization mechanism mustat leastin part, if not generally, be operative which, as ~OIIOWSfrom the
consideration of lifetimes, induces nuclear polarization during molecule formation in the reaction transition state^.^^^^^
Acknowledgment. It is a pleasure to acknowledge stimulating discussions with Dr. R. W. Fessenden of this laboratory on many topics of this paper. This work was supported in part by the United States Atomic Energy Commission. (33) NOTEADDED IN PROOF.While this paper was in press similar results have been published by G. A. Russell and D. W. Lamson, J . Amer. Chem. ~ O C . ,91, 3967 (1969).
The Kinetics of Crystal Growth of Dicalcium Phosphate Dihydrate by Robert W. Marshall and George H. Nancollasl Chemistry Department, State University of New York at Buffalo,Buffalo,New York
(Received April 3, 1969)
The kinetics of crystallization of dicalcium phosphate dihydrate (DCPD) has been studied by following the changes in calcium, phosphate, and hydrogen ion concentrations when stable, supersaturated solutions are inoculated with seed crystals of the salt. After a brief initial surge, the rate of growth of the crystals follows an equation that is second order with respect to concentration, ([Caz+][HP02-J), over a wide range of calcium/phosphate ratios, suggesting a predominantly surface reaction controlled process. Experiments at 15, 25, and 37" have been made in supersaturated solutions over a range of calcium and phosphate concentrations and the en. ergy of activation corresponding to the crystal growth process is 10.5kcal mol-l. Addition of sodium pyrophosphate to the supersaturated solutions has a striking inhibitory influence upon the rate of crystal growth of DCPD. The effect is analyzed in terms of adsorption, following the Langmuir isotherm, of pyrophosphate ions at the available crystal growth sites. Dicalcium phosphate dihydrate (CaHP04. 2H20; hereafter written DCPD) has long been recognized as an important product in the application of fertilizers to soils.2 The main interest in this sparingly soluble salt, however, stems from its participation in the physiological formation of calcium phosphates. Thus DCPD has been proposed as an intermediate in the formation of bones and teeth,3 but there is still much discussion as to the exact nature of the precursor in the formation of hydroxyapatite (&,(OH) (Po&; hereafter written HAP) under physiological conditions. DCPD may also play an important part in dental caries production. Caries is believed to result from the attack upon the dental enamel, mainly HAP, of organic acids produced by the metabolic action of the bacteria living in the vicinity of the enamel ~ u r f a c e . ~These organic acids, with pK values in the range 5-7, produce acidic conditions with pH from 4 to 6.4 In aqueous solutions of calcium phosphate under such conditions, the stable solid phase consists of dicalcium phosphate, and the The Journal of Physical Chemistry
complex species present in the solution are CaH2POdf and CaHP04.6 It is clear from the above brief historical outline that there is considerable interest in the mechanism of growth of DCPD crystals from supersaturated solutions of the salt. Previous studies6,' have been concerned with spontaneous crystallization processes in which both nucleation and growth are occurring spontaneously. Such processes are very difficult to analyze kinetically, (1) T o whom all correspondence regarding this paper should be addressed. (2) (a) W. E. Brown and J. R. Lehr, Soil Sei. SOC.Amer. Proc., 23, 7 (1959); (b) J. R. Lehr and W. E. Brown, ibid.,22, 29 (1958). (3) A. E. Sobel, M. Burger, and S. Nobel, Clin. Orthop., 17, 103 (1960). (4) J. A. Gray, J . Dent. Res., 41, 633 (1962). (5) A. Chughtai, R. Marshall, and G. H. Nancollas, J . Phys. Chem., 72, 208 (1968). (6) T. Hlabse and A. G. Walton, Anal. Chim. Acta, 33, 373 (1966). (7) A. G. Walton, 13. A. Friedman, and A. Schwartz, J . Biomed. Matter Res., 1,337 (1967).
3839
CRYSTAL GROWTH OF DICALCIUM PHOSPHATE DIHYDRATE however, since the implicit assumption that only homogeneous nucleation takes place is of doubtful validity. Nucleation probably occurs heterogeneously on impurity particles which offer available sites for crystal growth. Added complicatioiis arise in deciding whether nucleation and growth occur simultaneously or consecutively. In the physiological environment, homogeneous nucleation is unlikely to occur. The process of calcification is probably initiated a t suitable sites on the bone or tooth mineral or on the organic matrix. In the latter case, it is likely that the organic matrix offers a suitable conformation of binding ligand atoms to the calcium ions in the initial tissue calcification step.' It has frequently been stated that normal blood serum is supersaturated with respect to the calcium hydrogen phosphate activity product { Ca2+){HP042-j, and it is clear that very special epitaxial conditions have to be met for calcification to take place. A study of the growth of added seed crystals in supersaturated solutions of calcium phosphate therefore appears to be particularly pertinent for a better understanding of the mechanism of calcification. The present study is concerned with the kinetics of growth of DCPD crystals from solutions supersaturated with respect to this electrolyte. The existence of a welldefined metastable limit for spontaneous crystallization has made possible the preparation of supersaturated solutions stable for periods of days. The rate of growth of added seed crystals could then be measured in a highly reproducible manner in contrast to the results of experiments involving spontaneous precipitation. A temperature study has made possible the determination of the activation energy for the growth process.
Experimental Section Materials. Reagent grade chemicals and grade A glassware were used throughout. DCPD crystals were prepared by the slow addition a t 30" of 250 ml of 0.02 M disodium hydrogen phosphate to a mixture of 125 ml of 0.40 M calcium chloride and 250 ml of 0.18 M potassium dihydrogen phosphate. Nucleation commenced at pH 5.5 which was maintained approximately constant during precipitation by controlling the addition of the disodium hydrogen phosphate solution. The seed crystals were washed by decantation ten times with conductance water and stored as a slurry a t a DH of 5.5. The regular, monoclinic crvstals had an average size of 30 p X 12 p . Analysis of a portion of the solid, dried at 40", gave calcium, 23.78%, phosphate, 54.46% (theoretical for CaHP04.2HzO: Ca, 23.29%; P, 55.77%). The X-ray diffraction pattern and infrared 'Pectrum were to those previously reported for the Seed crysta1s were aged for at least 6 weeks before use. The solid phase analysis was checked periodically for more than a year, and no be detected* 'Odium change in pyrophosphate solutions were prepared from the Baker
Analyzed reagent (Lot No. 33159) without further purification. Growth Experiments. The crystallization cell consisted of a 500-ml round-bottom flask with five standardtapered necks. The glass stirrer, passing through a gland in the central neck, consisted of two glass swingout links; the speed could be adjusted within the range 50 to 700 rpm. A nitrogen gas lead-in tube was sealed to the bottom of the flask, and additional turbulence could be effected by the bubbling of nitrogen through the cell to remove carbon dioxide. The cell was supported in a water thermostat, and growth experiments were made a t 15 f 0.02, 25 f 0.02, and 37 The pH was measured by using the elecf 0.03'. trode system contained in the cell glass electrodelsolution under studylsatd KCI; Hg2C12, Hg together with a Corning Model 12 pH meter and Beckman Type 41263 glass electrode. The meter output was displayed on a Sargent Model SR recorder making possible continuous monitoring of pH during the crystal growth. The electrode systems were standardized before and after each experiment with NBS standard buffer solutions prepared according to Bates;l0 0.05 m potassium hydrogen phthalate, pH 4.008 at 25", 3.999 a t 15", and 4.028 at 37" ; 0.025 m potassium dihydrogen phosphate 0.025 m disodium hydrogen phosphate, pH 6.865 at 25", 6.900 at 15", and 6.841 at 37". Solutions, supersaturated with respect to DCPD, were prepared by the slow addition of disodium hydrogen phosphate solution to a mixed solution of potassium dihydrogen phosphate and calcium chloride contained in the stirred cell. Each solution was allowed to equilibrate, with nitrogen bubbling, for at least 30 min before the addition of seed crystals in the form of a slurry containing from 18 to 180 mg of DCPD. Crystal growth commenced immediately in all cases and was followed by monitoring the pH and also by the removal of aliquots of supernatant solution for chemical analysis. Calcium analyses were made using a Perkin-Elmer Model 303 atomic absorption spectrophotometer, and phosphate was determined spectrophotometrically as the phosphomolybdate. " I n addition, periodic microscopic examinations of the growing crystals were made using a Leitz Wetzlar Dialux-Pol polarizing microscope.
+
Results and ~ ~ S c u s s i o n At the pH of the crystallization experiments both [Hap041 and [Pod3--]concentrations are negligible, (8) p. W. Amold, Trans, Faraday
46,1061 (1950).
(9) E. E. Berry and C. D. Baddiel, Spectrochim. Acta, 23A, 2089 (1967). (10) R. G.Bates, "Determination of pH," John Wiley & Sons, Inc., New York, N. Y., 1963. (11) D. N. Fogg and N. T. Wilkinson, Analyst (London), 83, 406 (1958).
Volume 78, Number 11 November 1060
ROBERTW. MARSHALL AND GEORGE H. NANCOLLAS
3840 and the following equilibria in the solution phase may be formulated HzP04H+ HP04'ha
+
+ HzP04- CaH2P04+ K + Ca2+ + HP04'- E CaHP04 K"
Ca2+
IC, represents the thermodynamic dissociation constant (values at 15, 25, and 37", respectively, 5.888 X 1 0 4 , and K is the 6.339 X lo+, and 6.562 X thermodynamic association constant for the formation of the calcium complexes? The concentrations of ionic species in the solutions were calculated as described previ~usly;'~f2, the activity coefficient of the divalent ions was obtained from the extended form of the Debye-Huckel equation proposed by Davies.I4 The thermodynamic solubility product, K,, = [Ca2+][HP042-]j~2, was obtained by allowing crystal growth experiments to proceed to equilibrium; at 25" the mean value of 21 determinations was 2.10 f 0.09 X mol2 This may be compared with values in the literature quoted by Bjerrum16 (KSp= 2.7 X 10-7 a t 37"), Fa+ (2.18 X mol2 1.-2 at 18" and 2.4 X mol2 law2, obtained after review of the literature), and Strates, Neuman, and Levinskas" (2.68 X 10-7 mol2 lo-') who did not include corrections for the concentrations of calcium phosphate ion pairs. More recently, Moreno and Brown and their coworkers1~~'9 mol2 a t 25" have obtained a value of 2.77 X by an extrapolation method and 2.19 X lo--' mol2 1.-2 at 37" after allowing for the presence of CaHzPOa+ and CaHP04. The results of some crystallisation experiments are summarized in Table I and typical curves of concentration against time are shown in Figure 1. Reproducibility is illustrated by the agreement between the results for experiments 10 and 11 and between 26 and 27. Throughout all the crystallization experiments, the decrease in total calcium concentration in the supersaturated solutions was within 3% of the corresponding total phosphate concentration change. The resulting indication that DCPD was indeed the growing phase was confirmed by petrographic examination of the crystals. In addition, the increase in crystal size, measured microscopically, was close to that calculated from the changes in calcium and phosphate concentrations on the assumption of equal growth rates in all three dimensions. The crystal growth curves in Figure 1 are characterized by a brief initial fast period followed by a smooth decrease in concentration with time. During the latter period, which constituted the major portion of the total crystal growth, the process follows the equation (l), where the rate is expressed as a change
either in total calcium or total phosphate concentration The Journal of Physicul Chemistry
10
20
30
40 50 Time, min.
60
70
Figure 1. Plots of total calcium (a) and total phosphate (b) concentrations as 8 function of time. One ordinate scale division corresponds to a concentration change of 2 X 10-4 M
with time, dT/dt. Since there is an appreciable change in crystallite size during the experiments (approximately a factor of 2 in the surface area), the correction factor (Wi/'W)'/* is introduced to correct the rate of growth a t time t to that corresponding to the initial surface area of the seed crystals. wi is the weight of seeds crystals added initially, and w is the weight present a t time t. The correction is made on the assumption of a regular three-dimensional crystal growth, and this is supported by the good agreement between the extent of (12) R.G.Bates and 5. F. Acree, J . Res. Nut. Bur. Stand., 30, 129 (1943). (13) G. H. Nancollas, "Interactions in Electrolyte Solutions," Elsevier Publishing Co., Amsterdam, 1966. (14) C. W. Davies, "Ion Association," Butterworth & Co., Ltd., London, 1962. (15) N. Bjerrum, "Selected Papers," Munksgaard, Copenhagen, 1949,p 245. (16) T. D. Farr, "Phosphorus," Chemical Engineering Report No. 8, T.V.A., Wilson Dam, Ala., 1950,p 52. (17) B. S. Strates, W. F. Neuman, and G. J. Levinskas, J. Phys. Chem., 61,379 (1957). (18) E.C.Moreno, W. E. Brown, and G. Osborn, Soil Sci. SOC. Amer. Proc., 24, 94 (1960). (19) E. C. Moreno, T. M. Gregory, and W. E. Brown, J . Res. Nat. Bur. Stand., A , 70, 645 (1966).
3841
CRYSTAL GROWTH OF DICALCIUM PHOSPHATE DIHYDRATE Table I : Crystallization of DCPD from Supersaturated Solution" Total concentrations Calcium Phosphate Expt no.
x
x
102
102
[Cas+I
x
PH
10'
[HPO@-1
x
104
crystals, mg
IC, 1. mol-' min-1 (mg of seed) - 1
Seed
Temp 15'
47 48
1.210 1.199
1.184 1.173
5.662 5.655
34 35 26 27 28 29 30 31 9 10 11 19 49 50 7 8
0.248 0.247 0.481 0.495 0.720 0.741 1.236 1.188 1.204 1.196 1.205 1.229 1.205 1.212 1.214 1.213
1.258 1.280 1.185 1.212 1.151 1.187 1.164 1.124 1.114 1.104 1.128 1.137 1.170 1.189 1.121 1.119
6.120 6.188 5.939 5.910 5 * 744 5.744 5.543 5.573 5.477 5.537 5.519 5.528 5 I557 5.563 5.513 5.507
44 45
1.209 1.204
1.159 1,155
5.417 5.419
1.059 1.050
4.674 4.560
27.5 27.5
0.16 0.17
0.178 0.173 0.377 0.377 0.596 0.611 1.066 1.026 1.048 1.038 1.045 1.065 1.037 1.041 1.054 1.054
13.46 15.74 8.680 8 * 377 5.614 5.814 3.756 3.855 3.109 3.502 3.447 3.543 3.891 4.010 3.378 3.329
23.0 23.0 23.0 23.0 23.0 23.0 23.0 23.0 18.0 18.0 18.0 18.0 27.5 27.5 90.0 90.0
0.28 0.26 0.30 0.30 0.28 0.26 0.30 0.27 0.29 0.25 0.24 0.25 0.31 0.32 0.30 0.29
1.024 1.020
2.895 2.894
27.5 27.5
0.60 0.60
Temp 25"
Temp 37'
a
Stirring rate 120 rpm; concentrations in moles per liter.
1.o
2.o
(CCa2t1CHP042-l- K S p / f z 2x)1 0 6 I
5.5
I
6.0 6.5 7.0 Log ( L C O ~ + I [ H P ~ ~-~KsP/~z') -]
Plots of log (wi/w)'//" (-dT/dt) against log ( [Ca2+] [HPOl*-] - K,,/fZ2): expt 31, 0 ;expt 34, A; slopes of the lines are as indicated.
Figure 2.
Plots of -dT/dt against ([Caa+l [ H P O I ~ - ]KsP/fz2):expt 49, 0; expt 28,Q; expt 31, 0 ; expt 19, A. Figure 3.
crystallization calculated from solution concentration changes and from microscopic measurement of crystal size during the experiments. Volume 78, Number 11 November 1060
3842
ROBERT W. MARSHALL AND GEORGE H. NANCOLLAS
Typical plots of log [(Wi/W)*'a( -dT/dt) ] against log ([Caz+J[HP04z-] - Ksp/fzz)are shown in Figure 2. For the major portion of the crystal growth, the slope 3.0 of the resulting line is unity, indicating a second-order dependence given by eq 1. During the initial surge, which was always less than 15% of the total reaction, 0 a slope greater than unity may be interpreted in terms x of the surface nucleation process used to explain a x similar phenomenon observed in the growth of barium 2.0 -0 sulfatez0 and strontium sulfatez1 crystals from their I supersaturated solutions. s In the region following the initial surge, the linearity 5i of the plots of (Wi/W)e'a(dT/dt) against ( [Ca2+]. L E [HP02-] - K s p / f z 2is ) seen in Figure 3, and the ex1.0 cellent reproducibility of the experiments is exemplified by the results of experiments 28 and 31. Equation 1 is closely followed over the wide range of [Caz++l/ [HP02-] indicated in Table I. Several experiments were made with different stirring rates; changes from the normal 120 (Table I) t o 600 rpm resulted in about 1.o 2.0 3.0 20% change in the value of the rate constant k in eq 1. ([CO~+][HPO,~-] - K S p / f 2 ' )X lo6 If diffusion were rate determining, the rate would be Figure 4. Plots of - dT/dt against ( [Ca2+I [HP02-] expected to be appreciably dependent on the rate of K.,/f22) at different temperatures: 0, expt 44 at 37'; El, expt stirring, and it is clear that within the stirring condi49 at 25"; A, expt 48 at 15". tions used in the present work, mass transport of ions to the crystal surfaces plays only a very small part in the over-all growth process. A preliminary analysis of marked inhibiting influence on the precipitation of some of the results had indicated a diff usion-controlled pressparingly soluble salts from s o l ~ t i o n . Their ~~~~ ~ crystal However, no correction was made ence, at low concentrations, in both plasmaz5and urinez6 for the changes in crystallite size during the experihas led a number of investigators to propose that they ments. The conclusions here reported are based on a are also effective in preventing the precipitation of more extensive body of data covering a wider range of calcium phosphate from metastable solutions in vivo.z7 concentration and stirring conditions. By assuming These observations suggest that pyrophosphate may be a surface-controlled process, the rate of crystallization the effective regulator of calcification processes in the of DCPD may be expressed in terms of the bulk conbody, and it is therefore of interest to study its effect centrations of the lattice ions upon the rate of crystallization of DCPD. The results of crystallization experiments made in rate of crystallization of DCPD = Ics[Ca2+][HP02-] the presence of small concentrations of sodium pyroThe opposing rate of solution is proportional to the phosphate are summarized in Table 11. Some data available surface area are plotted according to eq 1 in Figure 5 , and it is seen that even in the presence of pyrophosphate ion, rate of solution = kls this equation accurately represents the experimental I n a saturated solution these rates are equal, and so results. If diffusion played an important part in the kl = kK,,/fzz. In a supersaturated solution, if diffugrowth process, the addition of an additive, known to sion plays an insignificant part in the growth process, be heavily adsorbed on the surface of the crystals, the net rate of crystal growth will then be given by eq might be expected t o have a marked effect upon the 1. Crystal growth experiments (Table I) were made also at 15 and 37") and the resulting kinetic plots are (20) G . E. Nancollas and N. Purdie, Trans. Faraday SOC.,59, 735 shown in Figure 4. It is seen that eq 1 is closely fol(1963). lowed at all temperatures, and the energy of activation, (21) J. R. Campbell and G. H. Nancollas, J. Phys. Chem., 73, 1735 (1969). obtained from the excellent linear plot of In k against (22) G . H. Nancollas, J. Crgstal Growth, 34, 335 (1968). T-I, is 10.5 =t 0.8 kea1 mol-'. This is appreciably (23) J. R. Howard, G. H. Nanoollas, and N. Purdie, Trans. Faraday larger than the value,z34.5 kcal mol-l, to be expected Soc., 56, 278 (1960). on the basis of pure mass transport control again point(24) M. Muira and H. Naono, Bull. Chem. SOC.Jap., 38, 492 (1965). ing to the relative unimportance of the latter mechanism (25) H. Fleisch and 5. Bisaz, Nature, 195, 911 (1962). in the growth of DCPD crystals. (26) H. Fleisch, Amer. Physiol., 203, 671 (1962). Pyrophosphate ions have long been known to have a (27) H. Fleisoh and W. F. Neuman, ibid., 200, 1296 (1961). YI
v
Y
The Journal of Physical Chemistry
3843
CRYSTAL GROWTH OF DICALCIUM PHOSPHATE DIHYDRATE Table I1 : Crystal Growth in the Presence of Sodium Pyrophosphatea
50 64 63 62 61 65 60 a
x
x
loa
PH
X 107
5.563 5.533 5.551 5.577 5.532 5.546 5.509
3.36 6.60 9.70 33.0 67.5 98.0
10'
4.010 3.721 3 * 795 3.946 3.641 3,846 3.405
1.041 1.070 1.052 1.027 1.051 1.051 1.044
k', 1. mol-1 min-1
[PY1
[HPOnZ-]
[Cas +]
Expt no.
(mg of
0
k:
k'
k
seed) -1
...
0.32 0.27 0.23 0.17 0.08 0.06 0.03
6.4 3.6 2.2 1.4 1.2
1.1
Stirring rate 120 rpm; 27.5 mg of seed crystals.
0.3
d .
0.2
.L
0.1
1
I
I
2.0
4.0
6.0 8.0 [Py] x lo6
10.0
Figure 6. Plots of k', in the presence of pyrophosphate, against [pyrophosphate].
6.0 1.o ![Ca2+][HP0,'-]
2 .o
3.O
- KSp/f2*)x IO6
Figure 5 . Kinetic plots in the presence of sodium pyrophosphate: expt 64, 0 ; expt 63, 0 ; expt 62, 0 ; expt 65, &
. -s .$
I 4.0
2.0
kinetics of the reaction. That such is not the case provides additional evidence for a surface-controlled process represented by eq 1. The striking effect of the additive in reducing the rate of crystallization is illustrated in Figure 6 in which the rate constant in the presence of pyrophosphate, k', is plotted as a function of adsorbate concentration. If we assume that the retarding action of the pyrophosphate ions is a result of their absorption a t growth sites on the crystal surfaces, the Langmuir adsorption treatment should be applicable. Suppose that a fraction a of the surface is covered by adsorbed foreign ions of molar concentration [Py]. The rate of adsorption may be written kl[Py]. (1 - a), and the rate of desorption is k2a, where kl and k2 are the corresponding rate constants. At equilibrium these rates are equal, and a = kl[Py]/
1.o
2 .o
3.O
[Pyl-' x
Figure 7. Langmuir isotherm plot of k / ( k [pyrophosphate] -l.
- k ' ) against
+
(kz kl[PyJ). Now k' = k (1 - a),and substituting for a we have
k2 - -k -1+k - k' kl [PYI In Figure 7, the left-hand side of eq 2 is plotted against [Pyl-', and it is seen that the Langmuir isotherm satisfactorily describes the marked inhibiting effect of
F. MEYER
3844 the pyrophosphate ions in terms of a monomolecular “blocking” layer of foreign ions at the crystal surface. Acknowledgments. This work is supported by Contract N00014-66-C0227 (NR 105-419), between the
Office of Naval Research, Department of the Navy, and the State University of New York at Buffalo. We also thank Union Carbide Corporation for the award of a research fellowship to R. W. M.
Plane Specificity in the Reaction of Methanol and Ethanol Vapor with Clean Germanium by F. Meyer Philips Research Laboratories, N.V. Philips’ Gloeilampenfabrieken, Eindhoven, Netherlands
(Receined April 4, 2060)
The adsorption of methanol and ethanol vapor on clean germanium powder and the subsequent desorption of the decomposition products as a function of temperature have been studied. The results are interpreted in terms of a dissociative adsorption mechanism and covalent adsorption complexes are proposed. The adsorption and desorption are found to be plane specific. Twice as much is adsorbed on the (100) plane as on the (111)plane, which is in accordance with the number of dangling bonds per germanium surface atom on these planes. The desorption products from these two types of surface are different.
Introduction The clean surfaces of germanium and silicon are highly reactive to many gases, which is due to the uncompensated (dangling) bonds of the surface atoms. This was demonstrated‘ for a number of hydrides of general formula H,A with 2 = 1, 2, 3, i e . , HC1, H2S, “I, etc. These gases showed a fast adsorption up to a coverage of one molecule per 22 surface atoms, and a dissociative adsorption mechanism, compensating all dangling bonds a t the surface, was suggested. However, hydrides with x = 4 such as CHI and SiH4 did not adsorb chemically, and neither did organic compounds with unsaturated bonds such as ethene or benzene. Therefore, the nature of the A atom seems to be of prime importance for the adsorption. It was the purpose of this work to study in detail the interaction between a clean germanium surface and two oxygen-containing organic compounds, namely methanol and ethanol. The results from earlier work2 with ellipsometry on silicon single-crystal surfaces showed very clearly that adsorption can be dependent on the crystallographic orientation of the surface plane. I n the present study gas volumetric measurements also indicate that the adsorption of alcohols on clean germanium powders is plane specific.
Experimental Section The adsorption and desorption measurements on powders were performed in an all-glass apparatus. A The Journal of Physical Chemistry
vacuum of lo-’ Torr was obtained with a mercury diffusion pump provided with a liquid nitrogen cold trap. Pressure readings of methanol and ethanol were taken with a McLeod manometer. The germanium powder, obtained by crushing a high ohmic single crystal in air, was cleaned by heating to 650” a t lo-’ Torr for 15 hr.l After cooling to room temperature, the methanol or ethanol vapor was admitted and the pressure decrease recorded. Because the alcohol adsorbs reversibly on the glass walls of the apparatus, blank experiments were performed a t different alcohol pressures to correct for the amount removed in this way. Some typical adsorption curves for methanol are shown in Figure 1. The excess alcohol was pumped off a t room temperature until no further methanol or ethanol could be detected by the mass spectrometer connected to the adsorption apparatus. Then the temperature was raised by using an electric furnace around the reaction tube and the temperature was measured with a thermocouple imbedded in the powder. No alcohol was desorbed as such; only decomposition products were detected. The desorption products were analyzed with an Atlas M86 mass spectrometer. Calibration curves in the appropriate pressure range for the different gases (1) A. €1.Boonstra and J. van Ruler, Surface &i., 4 , 141 (1966); A. H. Boonstra, Philips Research Reports, Suppl. No. 3, 1968. (2) G. A. Bootsma and F. Meyer, Surface Sci., 14,52 (1969).
