YRACTIONATIOXOF OXYGENBY XZEOTROPIC DISTILLATIOX
August#,1961
and the first equilibrium constant is The terms involving K z can be shown to be small even when the urea concentration is 2 M . At 25", 2 M urea [D]/[D,] = 1.66/1.95 an estimate of K1 is 0.096, by using an activity coefficient of 0.87 for the 2 fM ureal6 and assuming that there are two equal interaction sites for urea on each DKP molecule, so K z 2 1/4 K1. The fraction of peptide bonds associated with urea under these conditions is 0.154. The difference of the heat of solution of DKP in water and urea (2 M ) is equal to - (0.52 f 0.07) kcal./mole, and so the heat of association of urea with the cis peptide bond would be -3.4 f 0.5 kcal.1' The problem of estimating an average hydrogen bond enthalpy for this situation requires knowledge of the average number of bonds formed between urea and the cis peptide unit. A maximum number would be two, which gives a value of -1.7 f 0.2 kcal. per hydrogen bond. Schellman4has assumed that there are approximately 4/3 hydrogen bonds in urea dimers, and estimates a value of -1.5 kcal. per amide hydrogen bond. Since the heat of reaction is larger between DKP and urea and the decrease of entropy is probably not so great in (16) G. Scatchard, W. Hamer and S. Wood, J . Am. Chem. Soc., 60, 3061 (1938). (17) A less accurate value can be estimated from K I as a function of temperature giving -4.5 f 2.0 kcal.
1435
forming two hydrogen bonds as in the case of urea with itself where internal rotation is removed, one would expect a larger average number of hydrogen bonds between urea and the cis peptide bond. An estimate of the enthalpy of a peptide amide hydrogen bond formation in water would then be more in the range of - 2 than -1.5 kcal. The stronger association of urea with the cis peptide bond than with itself is reflected in an estimate of K1 as 0.096 a t 25" and 2 M urea, compared with Schellman's urea-urea association constant of about one-half this value. Even with the increased tendency of DKP to associate with urea, and likewise with itself, the fraction of DKP molecules associated in water to form dimers is still small. Assuming a self association constant of 0.1 for D W a t 25' where a saturated solution is about 0.1 M , something like 1% would be associated. The formation of a molecular complex between DKP and urea when precipitated from an 8 A4 solution, lends further evideiice to the hypothesis of significant hydrogen bonding of urea with proteins as the cause of protein denaturation. la It is interesting that, such a complex forms only at high urea concentrations which are comparable to concentrations where pronounced protein denaturation likewise occur^.^^^^^ This similarity could well be due to the similar action of urea to both cases. (18) W. H. Harrington and J. 4 . Bchellman, Compt. Rend. Z'rau. Lab. Carlsberg, Ser. Chim., 30, 21 (1956). (19) L. K. Christensen, ibid., 28, 39 (1956). (20) J. G. Foss and J. 8. Gchellman, J . P h y s . Chem., 63, 2007 (1959).
FRACTIONATION OF OXYGEN ISOTOPES BY THE DISTILLATION OF AZEOTROPIC SOLUTIONS'~2 B Y LOISNASHKAUDER, w.SPINDELAXD E. u. MONSE Department of Chemistry, Rufgers, The State University, Newark 2, New Jersey Received April 6, 1961
+
+
The possibility of using the exchange equilibrium H2Ols [hydrated] H20l8 [solvent] e H20l8 [hydrated] HzOL6 [solvent] for concentrating oxygen isotopes, by distilling azeotropic acid solutions, has been investigated. Rayleigh distillations were carried out to determine the relative fractionation factors, CY, for constant boiling HCl, HBr and HNOJ a t several temperatures and pressures as compared to a' for water. Experiments were also run in a distillation column [65 cm. long and 13 cm. i.d.1 packed with tantalum helices, at atmospheric pressure and varying flow rates. With a flow rate of 38 ml./hr.,. water showed an over-all separation [depletion] of 1.41; 6.8 m HCl, 1.83; 10.9 m HBr, 1.90; and 31 m "03, 1.28. The single stage faoctor was calculated to be 1.007 for HCl a t 110' and 1.007 for HBr a t 126" as compared to a value of 1.004 for water a t 100
.
Introduction Taube and ~ o - w o r k e r s ~have .~ investigated oxygen isotope effects in aqueous solutions and have (1) This work was carried out under contract -4T(30-1)-2250 between Rutgers, The State University, and the U. S. iltornic Energy Commission. (2) Presented before the Division of Physical Chemistry a t the 138th meeting of the American Chemical Society, New Y o r k , N. Y., September, 1960. (3) J. P. Hunt and H. Taube, J . Chem. Phys., 19, 602 (1951); H. M. Feder and H. Taube, ibid., 80, 1335 (1952); A. C. Rutenberg and H. Taube, ibid., 80, 825 (1952). (4) H.Taube. J . P h y s . Chem.. I S , ,528 (1954); Ann. Rev. Nucl. Sei., 6. 277 (1956).
used such effects to study the nature of ionic hydration. I n brief they found that the bonds formed between a number of cations and the oxygen of hydrated water molecules are of sufficient strength to cause the equilibrium constant for the reaction HzO'6 (hydrated)
+ H2018(solvent) + H201a (solvent)
H2018(hydrated)
(I)
to be appreciably different from unity. The magnitude of the constant is dependent on the charge to size ratio of the cation, and as an example it may be noted that the equilibrium constant K for equation 1 is 1.040 at 25' when the hydrated
LOISNASHKAUDER, W. SPINDEL AND E. U. MONSE
1436
ion is a p r ~ t o n . Naturally ~ in aqueous ionic solutions the large excess of uncoordinated water molecules makes it impossible to realize this magnitude of isotope discrimination, but the preferential hydration by the H2018species leads to a change in the oxygen isotope fractionation for distillation of aqueous solutions above the fugacity ratio H2016/H201s.The magnitude of the change depends upon the nature of the cation, and its concentration in the solution. Taube4 indicates that salts containing cations of high charge to size ratio, ie., Al+++, Mg;++, H + increase the Separation factor above the value found in pure water, and the effect is linear with the molality of the salt over a considerable concentration range. From the magnitude of the separation factors for several ions at 25O, it appeared that the hydration effect might be favorable for concentrating oxygen isotopes, provided that the single stage factor could be multiplied by some suitable process. The distillat>ion of azeotropic solutions appears to be a reasonable way of accomplishing this, because the chemical composition of the system remains unchanged during distillation, while the H20I6species shiould preferentially concentrate in the vapor phase, to a greater extent than in the distillation of pure water. The single stage separation factor, a, is defined as the ratio (H201s/H2016) liq./(H 2 0'8/H20la) gas. It may be estimahed from the equilibrium constant K for the exchange between hydrated and solvent molecules in the liquid according to equation 1, and the fugacity ratio, a', for water, by the relationship6 (Y
= [ M ( K - 1)
+ lja'
(2)
Here, M is the mole fraction of hydrated water molecules and is obtained from the molality, m, of the cations in solution, and the hydration number, u for the cation [i.e., no. of hydrate water molecules per cation] by the relationship4 M = um/55.5. For ilf = 0, the single stage factor, a,reduces to a', the fugacity ratio in water. Using equation 2, and Taube's datal4jt is found that for 4 to 12 m acid solutions the single stage factor, CY, a t 25' increases by 0.003 to 0.009 above the value a' = 1.0083for pure water.6 Rayleigh type distillations were carried out to determine the relative magnitudes of single stage factors at temperatures other than 25O, for distilled water, and for azeotropic solutions of hydrochloric acid, hydrobromic acid and nitric acid. Studies of these systems also were carried out in a multi-stage diatillation column to measure the isotopic concentration produced by the hydration effects. Experimental Rayleigh Distillations.-A large volume (1 liter) of each solution was distilled to a small volume (5-50 ml.) a t a slow distillation rate of 1.7 ml./min. BO as to minimize bumping. The single stage separation factor, a, was calculated
-
Vol. 65
from the measured separation factors in these distillations by the well known Rayleigh formulaT ( h T / N o ) a / a - l = Wo/W ( 3) where N and N O(both assumed to be < < 1 ) refer to the mole fraction of the rate isotope in the final and initial samples, respectively, and W and WOare the final and initial volumes. Column Experiments.-A vacuum jacketed column 65 cm. long and 13 mm. i.d., packed with tantalum "Helipalc" spirals (0.035" X 0.070" X 0070", Podbielniak, Inc., Chicago, Ill.) was used. The system was operated with a large reservoir of boiling liquid a t the lower end and a minimum holdup a t the upper end. Samples of condensate were analyzed for oxygen-18 depletion as a fiinction of time, and flow rate. The column was preflooded before operation by pouring boiling liquid down the column while heating the reservoir a t a much higher rate than required to produce the desired throughput. As the voltage of the heater was reduced gradually, a column of liquid slowly descended through the exchange column while refluxing continued. Contamination of the vapor a t the top by atmospheric water vapor was prevented by closing the outlet at the top with a balloon or polyethylene bottle. The constancy of flow rate was checked by measuring the drop rate every hour. Isotope Analysis.-Samples of the original and final solutions (0.6-1.0 ml.) were equilibrated overnight with measured amounts of COn and analyzed for 01*content.* The acid solutions were first neutralized with anhydrous ammonia since highly acid solutions do not reach equilibrium even after 20 hours of contact. Another method of isotope analysis was used in some cases. This involved a rapid equilibration with 0 2 in an electrical discharge? About 0.05 cc. of H2O or acid was added to a 100-cc. discharge tube which had been filled with dry N2or air. After degassing the liquid, about 6 cc. atm. of O2was metered in. A 7500 volt transformer was connected, and the system was arced for 20 minutes in an oven at 70-80". Before admitting the gases to the mass spectrometer, the liquid was frozen with a Dry Ice-trichloroethylene mixture. No oxygen was added for analysis of nitric acid, and O2and the latter since it decomposes partially into can be analyzed without isotopic dilution.
Nz
Results Rayleigh Distillations.-A summary of the results is shown in Table I. The single stage separation factors shown are the ratios (0l6/O1*) solution/ (016/018) gas phase. The absolute values of a obtained by Rayleigh distillation of water at the three temperatures, 42, 63 and 100" are somewhat higher than the values given by Dostrovsky and Raviv? This probably is due to some refluxing in the flask. TABLE I SINGLESTAGEFACTORS WOR OXYGENFRACTIONATION AS DETERMINED BY RAYLEIGH DISTILLATION Solution
Molality
Temp. ("'2.)
1120
42 63 100 125 45 110 65 126 121
RCl
8.2 6.8 HBr 13.4 10.9 "01 31 a Extrapolated value.
Pressure (om.)
Separation fnrtor, a
(rtO.001)
-7
1.009
-9
1.007 1.005 (1.004)"
76 -7 76 -9 76 76
1.014 1.009
1.016 1.010 1.006
P
-
a'
0.005 .OO4 ,009
.006 .002
(5) E. U. Monse, W. Spindel, L. N. Kauder and T. I. Taylor, J . Chem. Phys., 83, 1557 (1960).
Rayleigh distillations often yield erroneous single-stage separation factors if refluxing occurs
(6) I. Dostrovsky and A. Raviv, "Proc. Int. Symp. on Iaotope Separation," Interscience Publishera, Inc.. N e w York. N. Y.,1958, p. 336.
Aten, Phys. Rev., 40. 1 (1932). (7) H. C. Urey and A. (8) M.jCohn and H. C. Urey, J . Am. Chem. Soc., 60, 679 (1938). (9) T. I. Taylor, personal communication.
.
FRACTIONATION OF OXYGENBY AZEOTROPIC DISTILLATION
August, 1961
OXYGEN ISOlOPE SEPARATION I N
TABLE I1 DISTILLATION COLUMN AT ATMOSPHERIC PRESSURE 5
1
Soh.
&O
F . 8 M HCl
2
ml./hr.
22 21 17 16 22 36 21 42 21 33 32 22 42
36 48 50 56 59 63 65 72 74 76 102 -1 40
27 27 27
18 38 42
26 23
43
31
771,
HSOa
77
n for
4 S
a = 1.004
1.41 1.40 1.38 1.36 1.34
1.38 1.36 1.30 1.34 1 33 1.35 1.29 1.32
60 G3
1.89 1.83 I . 74 1.i9 1.79 1 .on 1.71 1.63
:1 s 38 37 38 37
3s 48 60 67 86
1 .!I0 1.91 1 .ti9 1.75 1.62
43 43 41 29
43
1 .as 1.22 1.21
26 26 27
10.9 m HBr
3
Run
51 56
1437
71 74 -92
in the pot, or if liquid droplets are carried over in the vapor. By distilling all liquids under the same conditions, such sources of error should cancel out to yield good relative magnitudes of a for pure water and the acids. The agreement between the increase of the single stage factor for 6.8 m HC1 a t 45', 8.2 m HC1 a t l l O o , 10.9 m HBr a t 12G0, and 12.4 m HBr a t 6.5' (as compared to pure water a t these temperatures), with the values expected for those concentrations a t 25°,4 is gratifying. The difference (a - a') increases with the concentrations of the acids, and is relatively insensitive to an increase in temperature. In the case of nitric acid, which has the highest molality, an a of 1.006 a t 121' was found. The exchange of oxygen between nitrate ion and water, and the large fraction of oxygen present as HNO, in the gas phase decrease the expected single stage factor appreciably. Column Experiments.-Table I1 summarizes the results of the column experiments. Column 4 of Table I1 is the final steady state separation obtained a t the various flow rates. The separation values shown were all obtained from samples taken from the column after operation for 20 hours, since jt was found that steady state was reached within this time. It is obvious that appreciably higher
1.19
6 n from eq. 4
7 a,calcd. froin columns 4 and G
86 84 81
77 74 81 77 66 74 72 76 64
70
.. .. ,. .. .. ,.
.. .. .
I
..
..
..
!U
I . 0070
85 84
1.OO71 1 ,0066
84 81 7!) 77 i6
s5 82
,.
l i
I*
iD
..
(is
.. ..
st-
..
..
73 72 66
1.0070 1.0072 1 ,0067 1.0070 1.0067 Av. 1.0069 f 0.0002 1 ,0075 1.0079 1.006S 1 ,0075 1.0071 Av. 1.0073 -C 0.000:1
1,0030 1 ,0027 1.0027 1.0026 Av. 1.0028 f 0.0002
separations were found under comparable conditions of flow rate and pressure for HC1 and HBr as compared with HzO and "08. In addition to yielding an over-all comparison of the various solutions for concentrating oxygen isotopes, the distillation column studies also provided an additional method for determining relative single-stage isotope separation factors, CY. The over-all separation, S , in a distillation column operating a t total reflux, is equal to an, where n is the number of separating stages. Determination of either a or n from a measurement of S is only possible if the other variable is known. The number of stages in a distilling column (or the height equivalent to a theoretical plate, HETP) is primarily influenced by the column geometry, the type of packing used and the interstage gas and liquid flow rates. A single packed column should be relatively insensitive to the liquids distilled, provided that the wetting characteristics are similar (ie., surface tension, viscosity, density). Thus, in the present column, since a is known for waters (1.004 a t loo'), n, the number of stages, can be evaluated as a function of flow rate from the steady-state separation obtained by wat,er dist,illation, as shown in column 5, Table 11. From these values the following empirical rc-
MICHAEL O'KEEFFEAND WALTERJ. MOORE
1438
lationship between the number of stages, n, and the flow rate, L , in ml./hr. can be obtained n =A
- BL
(4)
Here A and B are constants, A = 99; B = 0.36 ? 0.01 ml.-1 hr. Equation 4 is valid for flow rates between approximately 20 and 100 ml./hr. The value of n calculated according to equation 4 (col. 6) at the flow rate. L, (col. 3) and the overall separation, S, (col. 4) measured for the acid solutions, has beeii used to calculate the single stage factor, CY for these systems. The factors are listed in col. 7 . The results for the average values of LY obtained in this way are: HC1 at l l O o , a = 1.0069 I 0.000'7; HBr at 126') a = 1.0074 $. 0.0007; and HXO? at l % O , = 1.0028 2 0.0005. The error is due>to the scattering of the measured values of n and the over-all separation. Discussion.-Single stage factors a may be compared with factors a' for water distillation at corresponding temperatures6 Thus the increase for the HC1 solution is a - a' = 0.0034 ==I 0.0007 at 110' and for the HBr solution, ac - a' = 0.0049 f 0.000'7 a t 12,j'. These values are in agreement with those obtained by Rayleigh distillation (Table I). They indicate that (a -1) for HBr solution is three times as large as the corresponding value
Vol. 65
for H20 (0.0025) at 126', and (a -1) for HC1 is twice the value for HzO (0.0035) a t 110'. The value of a for "3,does not differ from that for HzO, as expected from the Rayleigh distillation. It may be of interest to calculate the equilibrium coiistant K for the exchange of oxygen isotopes between hydrated and solvent molecules from the measured data. Using equation 2 , the measured values for a' and the data for HC1 ( g = 1. m = 6.8, and a = l.O07), the equilibrium constant at 110' becomes K = 1.024 f 0.008. From the data for HBr (u = 1, m = 10.9, a = 1.00'7) the equilibrium constant at, 126' becomes K = 1.025 f 0.008. This compares to the value of 1.040 calculated by Taube2 a t 25'. Since the separation factors observed for distilling azeotropic acid solutions are significantly larger than those found for water distillation a t a comparable pressure, distillation of such solutions may be useful for coilcentrating oxygen isotopes. Experiments are in progress on oxygen exchange in a packed column, between concentrated aqueous lithium chloride aiid water vapor. Recent resulJs by K. E. Holmberg (Acta. Chem., Scand. 14, 1660 (1960)), just called t o our attention, on physical properties of isotopic forms of certain inorganic acid-water azeotropes, confirm our findings within experimental error.
DIFFUSIOIK OF OXYGEN IN SINGLE CRYSTALS OF NICKEL OXIDE BY
MICHAEL O'KEEFFEAND
n'.4LTER
J. M O O R E '
Chemical Laboratory, Indiana University, Bloomington, Indiana Received Aprzl 19, 1961
The diffusion of 180has been measured in monocrystalline S i 0 by following the exchange of gaseous oxygen enriched exp( -54 in l80with the NiO crystals. At an oxygen pressure Poz = 50 mm., from 1100 to 1500°, D'N,o = 1.0 ,X kcal./RT) cm.2 sec The D'N,o increases with PO?. The most reasonable mechanism for the oxygen diffusion is believed to be by way of interstitial oxygen atoms.
Several aut,hors have suggested that diffusion of oxygen may be a contributing mechanism in the growth of oxide films on iron2 and nickel.3 No experiment'al da,ta have been available on the diffusion of oxygen in cryslals of FeO, Xi0 or COO. The present work provides some information on the '*O-KiO system and allows us to conclude that Dgi0, the diffusion coefficient of oxygen in NiO, is at' least several orders of magnitude less than D$io, t,he diffiision casefficieiit of nickel in nickel oxide, at all temperatu:res. Experimental Methods
which was connected via a Kovar-glass seal t o a conventional vacuum syatem. After installation of the sample and pumping out, enriched oxygen was admitted to a known pressure, and the tube closed off by means of a vacuum stopcock. The volume thus sealed off was approximately 50 cc. After heating for a known time a t a constant temperature between 1100 and 1500" a sample of oxygen was drawn off and analyzed mass spectrometrically . If A Band A t are the abundances of oxygen-I8 in the gas phase initially and a t time t , respectively, the ratio of the uptake by the solid a t t to the uptake a t infinite time is given by
The diffusion coefficients were measured by following t h r oxchange of oxygen in the sample with a surrounding atmosphere of oxygen enriched in 1 8 0 . The enriched oxygen was obtained by electrolysis of water obtained from the Weissmann Institute and containing 0.437% 1 7 0 and 6.6270 180. The nickel cixide sample was in the form of a cube lvith (100) faces and 3 mm. cube edge cleaved from a single crystal boule grown by the Verneuil method a t the Tochigi Chemical Industry Company. To eliminate the possibility of exchange of oxygen with the container, the NiO sample was supported on platinum foil in a 90% Pt-lO% Rh tube
where 0 = X g / N s , the ratio of numbers of atoms of oxygen in gas and in crystal, and A is the natural abundance of oxygen-18 in the nickel oxide. The diffusion coefficient, D, is then given by
(1) Work done under Contract AT-(ll-1)-250, U. S. Atomic Energy Comniission. (2) H. Pfeiffer and B. Ilsohner, 2. Elektrochem., 60, 424 (1956). 13) R. Lindner and A. Akerstrom, Disc. Faraday Soc., 88, 133
(1957).
(2) where 01 = @/3and e erfc Z = eZ2(1 - erf a). The righthand side of equation 2 is the solution for a slab of thickneKq the cube of this expression is the corresponding solutio11 for a cube .s (4) J. Crank, "The Mathematics of Diffusion." Clarendon Press. Oxford, 1856, p. 54. (6) S. C . Jrtin, PTOC.R o g . SOC.(London), A243, 359 (1957).
