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Vol. 67
TABLE I PHOTOCONDUCTIVITY AND SEMICONDUCTIVITY OF PURINES AND PYRIRZIDINES Photocurrent (Ai) X 1011
Compound
Dark current X 1011,amp.
air
vac.
air
vac.
air
vac.
Adenine Uracil Thymine Cytosine Guanine Hypoxanthine
0.26 .50 .55 .50 .22 .25
2.88 5.75 4.45 4.25 2.74 5.75
1.24 4.10 3.20 2.80 1.09 3.50
0.71 .10 .70 .70 .60 2.24
0 0 0 0 0 0
0.56 .55 .55 .60 .47 .55
0.90 1.38 1.09 1.40 1.10 1.10
290 ma
, 2.6
I
WAVELENGTH,
mp
Fig. 1.-Comparison of absorption spectrum of 0.02 M solution of adenine in water (solid line) and photoconductive action spectrum of adenine powder in air (dashed line). Ideally these measurements should be made on single crystals but we have not been able to grow suitable crystals of these quite insoluble compounds. Therefore we used thin films of finely powdered samples. The bases were chromatographically purified preparations from the Pabst Laboratories; repeated recrystallizations did not alter their spectral characteristics. The electrodes were quartz plates (Corning Glass Works) rendered conductive by an adherent stannic oxide layer. Blank measurements on sucrose showed essentially zero conductance and photoconductance, indicating that the preparation of the films did not abrade conductive material from the plates. The plates showed a transmission of only 15% a t 260 mp rising sharply t o IO070 from 320 mp. I n drawing the action curve, the photocurrent was corrected for the transmission characteristics of the quartz and the spectral intensity of the light. The electrical measurements were made by the standard d.c. technique6 for such “sandwich cells” with the use of a vibrating-reed electronieter. Monochromatic light was isolated from a hydrogen or mercury arc lamp by means of a Bausch and Lomb 250 nmi. grating monochromator. Intensity of light was measured with an IP28 phototube calibrated against an oxalic acid dosimeter. Absolute intensity data have little significance since much of the incident light is lost by scattering from the powdered sample, and from 270 to 400 mp no light passes completely through the conductivity cell. The sample thickness was set by Teflon spacers a t 0.0086 cm. The incident light intensity at, 290 mp was 120 uw./cm.2; a t 360 mw, 265 wRr./cm.a. Measurements of dark conduction were made in the temperature range 30 to 70”. The dark conduction was the same in air and in vacuo of torr, but the photocurrent was markedly dependent on the ambient atmosphere. The electric field across the electrodes was 2500 volts/cm. Ohm’s law was followed over the range 600 to 9000 volts/cm.
Results and Comments The experimental results are summarized in Table I. The reproducibility of current measurements on duplicate samples was about 10%. The photoconduc(6) B. Rosenberg, J . Chem. Phye., 81, 238 (1969).
E, e.v.
360 ma
tive action spectrum followed the absorption spectrum of a concentrated water solution of the base. An example is shown in Fig. 1. This correlation is evidence that the measured photoconductivity is a molecular property of the crystals and is not simply a surface effect in the powdered samples. On the other hand, the importance of surface effects is emphasized by the markedly higher photocurrents in the samples exposed to air. In air, appreciable photocurrent was detected in the range 320 to 380 mb, but this was practically absent in vacuo. I n air, the photocurrent was higher when the sample was illuminated through the positive electrode than when it was illuminated through the negative electrode. In vacuo, the photocurrent was independent of the direction of illumination. The last two columns in Table I give the activation energy E for the semiconduction, computed from c = UO exp(-E/kT). The activation energies were markedly lower in air, although the dark currents at 30” were about the same. Since the band gap for the onset of the photoconduction process is about 3.5 e.v., it is likely that it is not related to the semiconductivity observed from 30 to 70’. The higher photocurrent in air when the sample is illuminated through the positive electrode is in accord with a mechanism in which the predominant current carriers are positive holes. It is possible that adsorbed oxygen molecules at the crystallite surfaces are then acting as trapping centers for electrons. However, the question of the detailed mechanism of conduction in these materials is likely to be solved only by ail exhaustive study of monocrystalline specimens of ultrahigh purity, in which surface effects and trapping by impurity centers can be more readily isolated. The present results are a t least consistent with exciton mechanisms in irradiated nucleic acids and suggest that adsorption of electron acceptors may influence the electronic properties. Acknowledgment.-This work is part of a program supported by the Office of Naval Research. TYe are indebted to Henry Mahler for helpful discussions. ~
A LBNGMUIR MEASUREMENT OF T H E SUBLIMATION PRESSURE OF MANGAKESE (11) FLUORIDE BY REKATO G. BAUTlSTA .4ND JOHN L. %TARGRAVE* Department of C h e m i s t r y , U n i n e r s f t y of Wzsconszn, lMadzson, W i s c o n s i n Received Fehruarg 16. 1966
The sublimation pressure of a single crystal of MnF2 has been measured by the Langmuir free-evaporation technique. Until this work was carried out, no other
*
Rice Univereity, Houst.on, Texas.
NOTES
July, 1963 experimental data on the sublimation rate of MnFz had been reported in the literature. Brewer, Somayajulu, and Brackettl have estimated the heat of sublimation to be equal to 75 kcal./mole. Experimental Apparatus and Techniques The experiment was crzrried out with a microbalance built inside a vacuum system. The apparatus has been previously described in detail.2 It consists basically of a resistance furnace, a beam balance, and an electrical control circuit for the balance to measure the change of voltage with the change in weight. The MnFl single crystal was suspended from a holder made of tungsten wire inside a glass envelope contained in the furnace. The temperature was measured by means of a calibrated chromelalumel thermocouple placed inside the glass envelope. Several readings were taken of the rate of weight loss of the sample a t a given temperature and the time-weighted average of all of the weight loss readings was calculated and reported as one pressure point. These d.ata were used to derive a log P us. l / T equation with a Fortran program on a Control Data-1604 digital computer.
Discussion of Results MnFz(g) was considered the only important vapor species. The free energy functions for both solid and gaseous MnFp, were .taken from the data compiled by Brewer, Somayajulu, and Brackett.l The experimental results are summarized in Table I. A third law calculation from these data gives AH0298 = 76.4 =t 1 kcal./mole for the heat of sublimation while t'he second law least-square plot of log P vs. 1/T gives AHo = 73.5 :j= 0.5 kcal./mole at Tavg = 939.0'K. When corrected to 298'K. using the heat content values compiled by Mah3 and the molecular constants estimated by Brewer, Somayajulu, and Brackett,' the 1.0 kcal./ second law approach gives AHo298=; 76.0 mole. TABLE I VAPORPRESSURE DATAON MnFz T , OK.
P
924.1 939.0 952.3 963.8 969.9 982.9 964.3 951.7 887.1
2.49 X 4.73 x 8.54 x 1.35 x 1.69 X 2.78 X 1.40 X 8.34 X '4.94 x
AH%,,, koal./mole
10+ 10-9 10-9 10-8
lo-* lo-*
IO-$ 10-10
76.46 76.46 76.38 76.39 76.43 76.44 76.36 76.38 76.34
-.
76.40 & 0 . 5 kcal./mole
Using the above sublimation and other available thermochemical data,4 one calculates the atomization energy of MnF2(g)to be 217 kcal./mole, which is in good agreement with the 220 kcal./mole estimate of Brewer, et a2.l This indicates 108.5 kcal./mole for the average bond energy in MnlF2 while Gayd.on5 has estimated DO(NnF) as 81 kcal./mole from molecular spectra via a linear Birge-Spooner extrapolation. (1) L. BreNer, G. Somayajulu, and E. Braokett, University of California Radiation Laboratory Report KO.9840, September, 1961; C h e m . Rez., 63, 111 (1963). (2) L. H. Dreger a.nd J. L . Margrave, J . P h y s . C h e m . , 6 4 , 1323 (1960); R. C. Paule, Ph.D. Thesis, University of Wisconsin, 1962. (3) A. D. Mah, U. S. BuTeau of Mines, Report of Investigation 5600 (1960). (4) D. R. Stull and G. C. Sinke. "Thermodynamic Properties of the Elements," Advan. in Chem. Series, ACS, 1956; F. D. Rossini, D. D. Wagman. W. H. Evans, 5.Levine, and I. Jaffe, National Bureau of Standards Circular 500 (1952); JANAF Interim Thermochemical Tables, edited by D. R. Stull, The Dow Chemical Company, Midland, Michigan, 1960. ( 5 ) A. G. Gaydon, "Dissociation Energies," Chapman and Hall, Lt,d., lY53.
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Acknowledgmen.ts,-The authors are pleased to acknowledge the financial support of the National Science Foundation and of the United States Atomic Energy Commission.
ISOTOPIC FRACTIONATION I N T H E OH--HzO EXCHANGE REACTION BY MICHA.EL GREENAKD HESRYTAUBE' George Hevbert Jones Laboratory, T h e University o/ Chicago, Chicago, I l l i n o i s
Received X a r c h 16, 1863
It is import'ant to know with some certainty the equilibrium constant in the system HzO"iiq
+ 018H-aq2
HzO'*iiq
+ O1'HFaq
in order to interpret the fractionation of oxygen isotopes during the hydrolysis of certain complex ions. Values of this constant, K , a t 25' have been calculated. Hunt and Taube%have obtained a result of 1.035, from vibrational frequencies of H20 vapor and of OHin 10 M solution. More recently Thornton3has refin'ed the value to 1.0385 by using the parameters of liquid H20 and inc1udin.g librations. I n both calculations approximate meth.ods were employed in deriving the frequencies for HzO1*. We have therefore performed a direct experiment to measure this quantity, which we find to be 1.045 A: 0.003 a t 15'. When an allowance is made for temperature, the agreement with calculated values is quite goodl. The actual numerical value of K is an essenti.al feature in arguments a b w t the mechanisms of hydrolysis which we have developed el~ewhere.~ Experimental The isotopic composition of a stock quantity of redistilled water was determined by equilibration with carbon dioxide.& Freshly cut sodium was held under this water to make a solution approximately 3 M in sodium hydroxide. This solution was held a t approximately 15" A measured volume of liquid was drawn off a t a rate of 2-4 ml./hr. under reduced pressure as vapor which was condensed and equilibrated with carbon dioxide. The volume and molarity of the remaining solution were measured. The vapor was drawn off slowly and was therefore assumed to be in isot,opic equilibrium with the solution. A value of 1.009 was taken from the graph of Dostrovsky and Ravivo for 01, the distillation separation factor of H2016 relative to H2O1*. The relative fugacities of HzO'B and HzO18are unaffected by Na+ ions.? Values of 1.044 i 0.004 and 1.046 & 0.004 were obtained for K in two experiments where the respective mean concentrations of sodium hydroxide were 4.8 and 4.3 M . I
The method of Hunt and Taube3 gives a value ,of 1.044 for K a t Oo, which leads to 1.039 a t 15' by interpolation. However their procedure really refers to a m equilibrium between aqueous OH- and HzO vapor. When an allowance is made for this by multiplying by a, K becomes 1.048 at 15'. Inclusion of the libration of OH- reduces this quantity to 1.046, which agrees well with the values observed. If it is assumed that Thornton's valueS has the same temperature depend(1) Department of Chemistry, Stanford University, Stanford, Californ:a. (2) H. R. H u n t and H. Taube, J . Phys. Chem., 63, 124 (1959). (3) E. R. Thornton, J . A m . C h e m . Sac., 8 4 , 2474 (1962). (4) &I. Green and H. Taube, submitted for publication in I n o r g . C h e m . ( 5 ) C . A. Mills and H. C. Urey, J . Am. C h e m . Soc., 62, 1019 (1940). ( 6 ) "Proc. of International Symposium on Isotope Separation," NorthHolland Publ. Co., Amsterdam, 1958, p. 337, ed. by J. Kistemaker, J. Bigeleiaen, m d A . 0. C. Nier. (7) H. M. Feder and H. Taube, J. C h e n . Phys., 2 0 , 1335 (1952).
