XOTES
July, 1961 only about one-third the affinity shown by serum a l b ~ r n i n . ~Likewise it shifts the pKa of a covalently-linked acid-base substituent in t'he same direction as is observed with corresponding protein ~onjugat'es.~ Furthermore, with the PVP conjugate, as with those of proteins, the addition of urea to the a,queous solution decreases markedly the shift in acidity c ~ n s t ' a n t ~ . ~ Urea is a well-known protein denaturant and produces mar'ked changes in macromolecular configurat'ion of proteins. The intrinsic viscosity of ovalbumin (molecular weight 44,000) for example, changes from 0.043 (g./100 c,c.)-I for t'he native protein6 to 0.21 for t'he protein in 7.5 M urea.' It sl3emed of interest, t'herefore, to examine t'he effect of urea on the viscosity and henre configuration, of polyvinylpyrrolidone. Experimental Viscosities were measured in a standard Ostwald viscometer whose water outflow time was 110 see. Before each measurement, the viscometer was cleaned with warm chromic acid solution, rinsed throughly with distilled water and dried with a stream of dry air. It was mounted carefully with the same chmp in the same position in a water-bath a t 25' and its vertical alignment was checked with a plumb line. Outffow times for a given solution checked within 10.03 sec. or betti?r. Stock solutions were prepared containing 4y0 polyvinylpyrrolidone in water and in 8 M urea, respectively. These were diluted with corresponding solvent to prepare more dilute so1ut:ions of the polymer. All solutions were filtered before use. Five-ml. samples were added carefully a t the bottom of the bulb of the viscometer to prevent the formation of bubbles. Densities of the solutions were measured with a Westphal balance. Measured viscosities, Q, were converted to reduced viscosities n r e d , according to the equation (S/QO) - 1 Sred =
Discussion It is of interest to note first that the intrinsic viscosity l : ~ ] , where =
4 ;
3 0.10 i
2 3 4 (PVP). Fig. 1.-Reduced viscosity of polyvinylpyrrolidone as a function of concentration in g. per 100 ml.: 0, in water; 0 , in 8 .If urea in water. 1
for a decrease in the degree of hydration of the macromolecule. In regard to comparisons of the behavior of proteins with that of polyvinylpyrrolidone, these viscosity experiments lead to two interesting conclusions. First it is evident that interactions such as binding of ions or masking of the reactions of conjugated groups can occur with a macromolecule having a relatively random configuration such as is found for PVP. Secondly urea may perturb these interactions of PT'P even though it does not produce a major change in macromolecular configuration. These observations are directly relevant to the molecular interpretation of corresponding interactions of protein macromolecules.6
C
where qo is the viscosity of the solvent free of polymer and c is the concentration of polymer in g. per 100 ml. of solution. Reduced viscosities for polyvinylpyrrolidone in water and in urea are plotted in Fig. 1.
h1
1275
lim
c+o
(Qrtd)
for polyvinylpyrrolidone in water is 0.225 (g,/ 100 ml.)-I Jvhich is of the same order of magnitude as that of denatured ovalbumin.' The average molecular weiglit of the sample of polyvinylpyrrolidone used is 40,000,* compared to 44,000 for ovalbumin. Hence even in water, polyvinylpyrrolidone is in a relatively loose configuration. It is not surprising therefore that urea has very little effect on t'he intrinsic viscosity of polyvinylpyrrolidone, the value dropping slightly, to 0.215. Such a drop is small and probably within experimental error, but it is of interest t'o note that the change wi.th urea is in the direction to be expected M. Klotz and J. Ayerc, unpublished experiments. ( 5 ) I. ,If. IClotz and Y. H. Stryker. J . A m . Chem. Soc., 82, 5169 (1960). (6) A. Polson, Kolloid Z., 88, 5 1 (1939). (7) H. K. Frensdorff. &I. T. Watson and W. Kauemann, J. An. Chem. Soc., 76, 5167 (1953). (8) "PVP, Polyvinylpyrrolidone," Bulletin P-100, General Aniline (4) I.
and Film Corp., New York, N. Y.,1951.
THE DISSOCIATION PRESSURE OF GALLIUM ARSENIDE BY V. J. LYONS ASD V. J. SILVESTRI I B M Research Laboratory, Poughkeepsze .Veu York Receaved January 7, 1961
Gallium arsenide thermally dissociates to the constituent elements and two previous investigations of the reaction equilibria have been reported. The solid-liquid-vapor equilibria were studied by van den Boomgaard and Schol' over the range 781 to 1237'. Goldfinger and Drowart2 have reported a mass spectrometric study of the vapor species resulting from thermal dissociation of the compound in the range 758 to 863'. The purpose of this work was to re-examine the reaction in a temperature range below the compound melting point because of the inconsistency in the reported data. Since the vapor pressure of arsenic is several orders of magnitude greater than that of gallium the dissociation pressure of GaAs is essentially equal to the arsenic pressure in equilibrium with the compound. The method used to measure the equilibrium arsenic pressure was the visual observation of the arsenic dew-point in a sealed tube containing solid Gails. The applicability of the method to decomposing solids has been demonstrated, and the more general details of the experimental procedures have been described (1) J. van den Boomgaard and K. Schol, Phillips Res. R e p . , 12, 127 (1957). (2) P. Goldfinger and J. von Drowart, J . G'hem. Phps., 68, 721 (1968).
1276 103
couple well; thus the film did not interfere with the dewpoint observations. After equilibration, the radiation shield was removed and the temperature of the cooler furnace was lowered. Dew-points then were observed in the manner described in earlier papers. The dissociation pressure of the coompound thus was measured over the range 1051 t o 1196 The data, which are presented in Table I and plotted in Fig. 1, can be fitted to the equation
0 0 0
P
102
I 0
I
.
00
I
10'
Vol. 65
NOTES
I
log P,,
I
I I
10-1
GaAs T,'C.
"A A A
10-4
A A
10-5
0.65 0.70 0.75 0.80 0.85 0.90 0.95 1.0 Fig. 1.--Comparison of three determinations of the dissociation pressure of GaAs: (1) circles are data of Boomgaard and Schol; (2) triangles are data of Goldfinger and Drowart; (3) vertical lines represent this work. elsewhere.3~4J For this experiment, ten grams of high purity monocrystalline GaAs was broken into small pieces and sealed in an evacuated quartz tube 34 cm. long and 20 mm. in diameter. The end of the tube containing the compound was inserted into a 6 cm. well in a cylindrical stainless steel block 8.5 cm. long. A thermocouple well drilled through the opposite end of the block permitted a thermocouple to be positioned adliacent to the reaction tube. This assembly was placed in the center of a 32 cm. long Hoskins furnace. The opposite end of the tube was enclosed by a separately controlled furnace fabricated by winding nichrome ribbon on a quartz tube. This furnace extended into the Hoskins furnace to ensure against cold spots occurring a t the junction of the two heating elements. The entire construction provided (1) isothermal regions over the GaAs sample and the thermocouple well a t the opposite end of the tube, (2) independent temperature control of the two regions, (3) a monotonically increasing temperature gradient in the region between the thermocouple well (dew-point temperature) and the GaAs sample, and (4) visibility of the thermocouple well. The temperature of the Hoskins furnace was maintained to within zt3' by a proportioning Wheelco controller while a Variac operating from a constant voltage transformer controlled the dew point furnace. Temperatures were measured with Pb-Pt,lOyo Rh calibrated thermocouples and the potentials were read on a Leeds & Northrup Potentiometer, Model 8662. Each experimental point was measured in the following manner. The decomposition furnace was heated to a temperature in the range 1051 to 1196', while the dew-point furnace was hea1,ed to approximately 650'. Equilibration times of 15-71 hours were allowed. In ractice, the area around the therinocouple well was heatex to temperatures slightly higher than 650', through the use of a radiation shield, because of the gradual condensation of a GaAs film a t the color end of the tube. Identification of this film was made by X-ray diffraction powder analysis. By using the shield, the GaAri film was caused to condense on the tube wall between the decomposition furnace and the thermo(3) R. Weiser, J Phy.3. Chem., 61, 513 (1957).
v
(1)
TABLE I
10-2
3. Lyons, J . Phy8. Chsn., 6 3 , 1142 (1959). (5) V. -1. Lyons and V. J. Silvestri, ibid., 64, 266 (1980). (4)
- 347500 ___ + 25.8 T'K.
Calculations based on the experimental conditions showed that 0.05 t o 1.0% of the sample had dissociated during the measurements. Evidence of a liquid phase on parts of the sample was observed after removing the GaAs from the reaction tube. These data show that the pressures were measured along the three-phase line and therefore describe, in part, the boundary of compound stability.
10"
10-3
=
Total pressure,a mm.
0.5-0.6 1051 2.0-2.6 1079 3.5-4.2 1100 10-13 1124 22-28 1142 43-47 1158 88-92 1176 185- 195 1196 a From dew-points. Calculated pressure.
Pas4 b
Phstb
0.53 0.03 2.06 0.24 0.44 3.41 2.36 9.1 17.5 7.6 29.9 15.1 53.3 36.4 97.6 92.4 from average total
The measured arsenic pressures were considered equivalent to the total arsenic pressure over solid G a h for the following reason. The reaction tube mas heated by a two-zone furnace so that approximately one-third of the tube was heated to the decomposition temperature and one-third was heated to the dew-point temperature. Hence, a significant length of the tube (approximately 10 cm.) was maintained a t the dew-point temperature. The tabulated data of Stull and Sinke6 show that the species As4 would predominate in the dew-point range of temperatures while two species, .4s4 and AS^, would be measurably present in the decomposition temperature range. From considerations of the temperature distribution in the reaction tube and the mean free path of the arsenic molecules it was concluded that, a t equilibrium, the Asz molecules would completely recombine to give only As4 molecules a t the dew-point temperatures. Thus, the measured pressures and the values of the equilibrium constant for the reaction As4 $ 2Asz were used to calculate the pressure of each species in equilibrium with solid GaAs. The over-all reaction equilibrium a t each experimental point may be represented by the chemical equation
where a is the As4 fraction of the vapor. The pressures of As4 and As2 calculated from each experimental point are given in Table I. A comparison of this work and that of van den Boomgaard and Schol shows good agreement a t the higher temperatures. It is thought that the (6)
D. R. Stull and G. C. Sinke, "Thermodynamic Properties of the
Elements." Advsnces in Chemistry Series. No. 8, 1956.
NOTES
July, 1961
1277
experimental difficulties encountered in attempting to extend the triple-point method to lower pressures and more dilute solutions ( i e . , the slow establishment of the equilibrium) led to their observation of higher pressure values a t the lower temperatures. Extrapolations of the Goldfinger-Drowart data and t,his work show reasonably good agreement when considering the range of pressures measured. The authors express their thanks to Dr. K. Weiser and Dr. G . A. Silvey for valuable discussions. 4690
THE SORET EFFECT AS A SOURCE OF ERROR IN CONDUCTANCE MEASUREMENTS BYR. H. STOKES Department of Physical and Inorganic Chemiatry. UniuerailU of Xew England, Armidale, N.S. W., Australia Receiiied Januarzi 20, 1061
Agar,' iii a paper on the rate of attainment of Soret equilibriunn, recently predict'ed that "there will be small transient changes of composition in a conductivity cell when it' is transferred from the ambient temperature into a thermostat." The reality and. considerable importance of this effect are demonstrated by the findings now reported. Two conductance-cells were used. Cell A consisted of two electrode-bulbs ahoiit 2.5 cm. in diameter, joined by 6 cm. of tubing of 6 mm. inkrnal diameter; its cell constant was 20.71 cm.-l. Cell 13, of constant 4.595 cm.+ was of similar constriictioii rxccpt that the central section was only about 1 cm. long. Filling tubes were attached to both electrode-bulbis, and the leads to the electrodes were carried through glass side-arms well separated from the filling-tubes and each other. After the detection of the effect reported below, the filling-tubes were fitted with mixing-bulbs to facilitate the mixing of the cell contents without removal of the cl.4 from the thermostat. The electrodes were lightly coated with platinum-black; frequency-dependerice in the range 1-4 kc./sec. was less than 0.002% with cell A and 0.008% with cell B. This small residual dependence was extrapolated out by an R us. f plot. The lead resistances mere measured by filling the cells with rnercury, and calibrations were made with the Jones and Rradshnw 0.1 dema.1 potassium chloride standards a t 25'. The cells were filled with 0.001 M hydrochloric acid and heated in a thermostat to 50 ; a t this temperature their contents were thoroughly mixed. They were then transferred quickly to a (hermostat held a t 25 & 0.001'. Here the resistances were measured over a period of several hours by a calibrated Jones bridge. I t was established by a prev; ions study that the contents of the cell must bp within 0.001 of the thermostat temperature after 20 minutes; hence any changes occurring after this time cannot be ascribed to tempuature changes. Results such as those shown in Table I n-erc consistently observed.
TABLE I .4PI'AREST
25"
SPECIFIC C O N D U C T A N C E S I N
AT YARIOUS
Time, hr.
rrIP,fssFOLLOWISG CELLSTO 50' Cell A
OHM-' cM.-'
PREVIOUS
% low
AT
HEATING OF
Cell B
pZ low
0.58 0.0043969 0.304 0.0044067 0.077 1,58 ,0043974 .295 ,0044075 .059 5.42 .0013091 ,254 ,0044082 .043 Aftcrmixing:tt 25" ,0044103 ,000 .0044101 ,000
The results from the two cells a t all times differed far more than could be ascribed t'o measuring or cdibration errors. But when the cell contents (1) J. N. Agar, Trans. Faradau Sac., 66, 776 (1060)
I
I
,
I
10 20 30 Time, hours. Fig. 1.-Change of resistance with time in cell A containing approximately 0.01 N hydrochloric acid. Cell thermostated a t 50', then transferred to 25' thermostat a t zero time. Temperature equilibrium is known to be attained before the time of the first point plotted. 0
were thoroughly mixed at 25', agreement was immediately established, and no further drift in readzngs then occurred f o r several days. Furthermore, the results after the mixing were in agreement with Shedlovsky's values2 within O.Olyo. Exactly the opposite effect was observed 011 moving the cells from a 25 to 50" thermostat: the results again differed considerably, but this time both cells showed higher conductances than the filial true values attained after mixing of the cell contents. I n aiiother experiment, cell A was heated to 50" in a water-bath, and its contents were mixed a t this temperature. The bath and cell were then allowed to cool slowly to 2 5 O , oyer a period of about 1.5 hours. The cell was then transferred to the 25' oil-thermostat for measurement. On reaching temperature equilibrium, the resistance was 4696.4 ohm, and on mixing the cell (iontents this value changed only to 4696.1 ohm, i.e., 0.0077,. The effect, then, evidently arises froin szrdden changes in cell temperature, and is therefore not to be ascribed t o differences in adsorption of hydrochloric acid at the electrodes a t different temperatures. An explanation in terms of adsorption also seems to be ruled out hy the fact that the effect is of much the same relatire magnitude for a wide range of hydrochloric acid concentrations. It seems clear that the effect ariws from thermal diffusion of hydrochloric acid into or out of the central portion of the cell during the period of rapid temperature change-w., the first few minutes at the new temperature, while there are large temperature gradients between adjacent parts of the cell. The persistent nature of the resulting concentration-disturbance is at fkst sight surprising, though once again it is in accordance with Agar's predictions.' In Fig. 1 is plotted another series of measurements on cell A, which were followed this time for 24 hours before the cell contents were mixed; after this time, isothermal diffusion had only reduced the original disturbance to half its initial value, and it seems likely that at least a week would be needed to get a result correct within O . O l ~ o ,unless provision is made for thorough mixing of the cell contents as soon as temperatureequilibrium is reached, The initial disturbance is
