NOTES
932
ilt low pressures of t-butyl nitrite and nitric oxide, no exchange of nit,rite was observed in blank experinients carried out in a darkened room. The nitric oxide did show dark exchange, yet the N16 did not appear in the nitrite. Evidently the nitrite contains an impurit,y (0.5 t,o 274) which undergoes fast dmk exchange with nit,ric oxide but not with t-butyl nitrite. The impurity could not be removed by dist2illation and was not, cJr1,erl~edhy gas chromatography. In the experiments :~1, 2,537 A . , a small correction pas necessary for exchange due to tho small amount, of 4047 A. light transmitted by the filter.
Results A fsw experiments were carried out at’ 25’ and 2537 A, using pressures of nitrite and nitric oxide of 14.5 mm. The quantum yield of combination of t-butoxy radicals and nitric oxide should be 0.16 under these conditions if the proposed mechanism6 is correct. The observed quant’umyields of exchange vere 0.3, 0.7, and 0.7. The excess may be due t’odark exchange beOween t-butyl nit,rite and some reaction product, nit’rosomethane, for example, derived from methyl radicals formed by decomposition of excit’edt-butozy. The quant’um yield of acetone at 3660 A. (temperature, 25’; nitric oxide, 14.5 mm.; nitrite, 14.5 mm.) wa,s found to be 0.03, 0.04,0and 0.04. This may be compared with 0.84 at, 2537 A. under these conditions.6 Dark exchange with roeaction products may thus be a small effect a t 3660 A. The observed photocohemical exchange under different, conditions at 3660 A. is included in Table I. TABLE I PHOTOLYSIS OF (CH&COXO IN THE PRESESCE OF W50 A, 3660 T , 25’; incident light intensity, -5 X
w.;
quanta/sec. Psiitrite,
Time,
% Exch.
mm.
mm.
8ec.
found
% Exch. theoret./q
L*
14.5 14.5 14.5 14.5 14.5 40.5 44.0
14.5 15.5 17.5 32.5 33.5 14.5 16.5
540 600 900 1200 1200 540 900
1.68 1.65 2.44 2.07 2.14 1.65 2.58
1.72 1.85 2.44 2.15 2.09 1.66 2.61
0.98 .89 1.00 0.96 1.02 0.99 .99
Pxo,
Discussion Ignoring for the present the small amount of aceto2e formed, the main reactions in the system a t 3660 A. are
+ h~ +(CH3)3CO + 3 0 (CH3)sCO + NO + (CH,)&OXO
(CH,),COKO
(1)
(2)
Re-formation of nitrite accounts for the observation that prolonged illumination of t-butyl nitrite with 3660 8.radiation led to no net decomposition.1° Decomposition of primary and secondary nitrites did occur,1o presumably because disproportionation is possible between primary or secondary alkoxy radicals and nitric oxide. Based on reactions 1 aiid 2, the quotient % exchanged (theoretical)/p may be calculated aiid divided into the observed yo exchange to give the primary quantum yield. The mean value of 9 is 0.98 with an uncertainty not easily estimable. Possibly the acetone quantum yield, 0.04, should be added to the photo-exchange quantum yield. But if the view is correct that rapid (IO) H. I\-.T h o m p s o n and F.
(1937).
S Dainton, T r a n s . Faraday
Soc., 88, 1646
T‘ol. 67
dark exchange takes place between t-butyl nitrite and nitrosomethane (or derivatives thereof). the small portion of the primary process ultimately yielding acetone may already be counted in the photo-exchange. KO firm conclusions can be drawn about the source of the small amount of acetone, In consistency with the low wave length results,Gthe acetone may be ascribed to decomposition of excited t-butoxy radicals. This does not require that excited radicals be formed by absorption of light in the bands; the few excited radicals may be formed by absorption in the weak continuum underlying the bands. The results on t-butyl nitrite niight be summarized by noting that the primary process following light absorption in the bands or in the “continuum” is dissociation to a t-butoxy radical and nitric oxide, with a probable primary quantum yield of unity. The net photochemistry is different in the two regions because of the increased importaiice of decomposition of excited t-butoxy radicals at lower wave lengths.
DIEFUSIOS OF TRITIATED WATER ( ~ 3 I N AGAR GEL AKD WATER BY F. S. XAKAYAXA.4XD R.D. JACKSO&
~ 1 0 1 9
U . 5’. W a t e r Conservation Laboratory, T e m p e , A r i z o n a Received October 86,1968
The “self-diffusion” coefficients of liquid water are of interest in that they may be used to ascertain the structural properties of liquid .ivater.l These coefficients are by necessity determined by isotopic tracer techniques. Wang, et a1.,2 by use of a diffusion capillary technique, determined the diffusion coefficient of tritiated water (H3H1016)in ordinary water at 25’. set.-' is Their value of 2.44 f 0.057 X lov5 widely accepted, although it is the only reported value known to the authors. We have determined the coefficient using a different technique and obtained a see.-’, which agrees value of 2.41 =I= 0,055 X with Wang’s value within the error of the experiment. Experimental The method consisted of determining the diffusion coefficient of tritiated water in low concentrations of gel and getting the coefficient in liquid water by extrapolating to infinite dilution of gel. Working u-ith gel material minimized error caused by non-diffusional movement of the liquid resulting- from mechaniral shock m d vibrations. Four low concentrations 10.3. 0.5. 0.75. and 1.0%- bv ” weight) of steam-sterilized agar-agar solutions in duplicate were allowed t o set in cylinders 1.9 cm. i.d. hy 1 2 cm. long. The cylinders were constructed so tshat.incremental 1-em. sections could be separated easily a t the end of a diffusion run. A 1.9-cni. filterpamperdid< was placed in contact with one end of the sample and treated with 0.01 ml. of a 100 pc./ml. t,ritiated water solutfion. The sample was settled and stored a t 25 f 1’ for 4s hr., and then sectioned. Water was ext,racted from the gel by vacuum dehydration, and the activity of the HZH1016 in the exbracted m-ater was determined by the liquid scintillatlon t e ~ h n i q u e . ~ The diffusion coefficient was calculated by obtaining the ratio of the activity (itx) in the column from z = 0 t o x = x to the tot,al activity (AT) in the column. Assuming a semi-infinite column, this ratio is equal to the integral of the instantaneous plane 8ource solution of the general diffusion equation. That) is
A,:AT
=
erf z(UIt)-’/*
J. H. Wang, J . A m . Chern. Soc., 73, 510 (1951). (2) J. H. W a n g , C. V. Robinson, and I. A. Edelman, ibid.. 75, 466 (19.53). (3) F. E. Kinard, ‘‘A Liquid Scintillator for t h e Analysis of Tritium in W a t e r , ” Atomic Energy Comm. Rept. DP-190. 1956. (1)
April, 1963
DIFF USION C OE FFIC IE N T, CM?/SEC.
to5
AGAR CONCENTRATION % ,
0
93s
XOTEH
0.5
I .o
Fig. I.---Diffusion coefficient of tritiated water in agar gel a t different concentrations. The assumption of a semi-infinite column was validated in that tracer artivity was not present in thr untreated end within the 3-8-hr. exposure prriod. An average diiiusiun Coefficient vias taken from the calruiation for the first five 1-cm. sections of rach column.
Results The diffusion coefficients of H3H1OI6measured a t four agar gel concentrations are presented in Fig. 1. A linear regression equation was calculated for this relation from which a value D = 2.41 f 0.055 X cm.? sec.-’ was determined for the diffusion coeficieiit of tritiated water in ordinary water. This value compares favorably with the results of Wang, et aL2 THERMODYX AMIC PROPERTIES O F T H E ATiMOSPHERIC GASES I X AQUEOUS SOLUTIONS‘ BY CORNELIUS E. NLOTSAND BRUCEB. BEXSON
qualitative solubility data.4 h precision of 0.1% in the solubility measurements permittrd the graphical evaluation of these data with an estimated error of 1% for the enthalpies and entropies and 10% for the heat capacities Relative values for different solutes are, however, considcrahly niow accwnl c. ‘I‘hiis, for r\nniplr, the identical d u e s of the partial molal heat capaci t ics reflect a relationship revealed in this work and exhibited in Fig. 1 and 2 . Plots of ln[K(S2)/K(Ar,02)1 = In [ai(Ar,Oz)/a(Nz)] us. l/T, where the K’s are the Henry’s law constants and the ai’s are the Bunsen solubility coefficients, give excellent straight lines. The standard deviations are 0.16% in the case of oxygen and 0.10% for argon. The nitrogen solubility coefficients employed for this purpose mere the result of thirty-two absolute measurements throughout the full temperature range. Discussion The observed entropies of solution and the enormous partial molal heat capacities give substance to the iceberg picture of aqueous solutions3 which envisages this heat capacity as arising largely from the melting of an ice-like structure surrounding each solute molecule. This 17iew has proved a useful one5 and has been extended to larger molecules of biological importance.6 Recent 3.m.r. investigations reveal, however, that the situation is not so simple.89 The present results therefore are of interest as they suggest an approach to this intriguing and important problem of the structural modifications in aqueous solution. The slopes of Fig. 1 and 2 are related to A(AH) = AH(Sz),,lu - AH(Oe,Ar),,ln and imply that these quantities are constant throughout the temperature range studied despite the fact that the individual enthalpies of solution display a pronounced temperature dependence. This suggests that the “icebergs”
’
TAnLE
Department of Phystcs, Amherst College, Amherst, Massachusetts Kecezved A p r d 18, 1961
Recent precision determinations of the solubility coefficients of nitrogen, oxygen, and argon in distilled water2 permit straightforward evaluations of certain thermodynamic properties of these solutions. Their magnitudes are of interest as an indication of the “iceberg” structures surrounding the solute molecules in water.3 A previously unreported relationship among these properties has now been obtained which supports previous ideas and offers a means of testing detailed models of these structures. Experimental The solubility measurements covered the temperature range 2-27”. Both manometric determinations of the absolute solubilities and mass spectrometric measurements of solubility rattos for a pair of gases were made, the two techniques giving identical results. These methods and the solubility coefficients are reported elsewhere.2
Results The derived partial molal enthalpies and entropies of solution and the partial molal heat capacities in solution are given in Table I. They are in good agreement with previously reported values based upon more (1) This work was supported by t h e National Science roundatlon under Grant NSF-G9437. (2) C. E. Klots and B. B. Benson, J. Marine Res., 21, 48 (1983). (3) (a) H. 6. Frank and M. W. Evans, J . Chem. Phys., 13, 507 (1945); (b) W. F. Clausseri and M. F. Polglase, J . Am. Cfiem. Soc., 74, 4817 (1952).
I
HEATSAKD ENTROPIES O B SOLUTION AS FUNCTIOXS OF TEVPERATURE (BASEDOK A HYPOTI-ZETICAL STAND.4RD STATE I N SOLUTION OF UNITMOLEFRACTION) AC, = Cp(M)so1n. - C,(M)gas T,‘(2. Nz
-AH (cal./mole)
0 2
Ar Kz
-AS
(cal./mole-deg.) 02 4r
ACp(oal./mole-der:)
2 5 10 20 25 15 4190 3650 3260 2950 2698 2520 4560 4020 3630 3320 3065 2890 4510 3970 3580 3018 2840 3270 36 9 34 9 31 6 31.0 33 5 32 5 36 8 34 8 33 4 32 4 31 5 30 9 36 4 34 4 33 0 32 0 31 1 30 5
:: 0 2
161
120
76
59
46
33
for these solutes are virtually identical at any given temperature. Small differences in the enthalpies (and entropies) of solution are then due solely to the process of introducing the solute molecule into its ice-like cage. Thus one may write: AH(M)soln. = AH(iceberg) AH(R4),,,ity. Here the first term on the right represents the heat of formation of the iceberg and is strongly temperature dependent; the second term arises from the introduction of the solute molecule (AI)
+
(4) D. M. Himmelblau, J. Phys. Chem , 6S, 1803 (1959). (5) (a) H. 8. Frank and W. Y . Wen, Dzseussions Faraday Soc., 24, 133 (1957); (b) H S Frank, Proc. R o y . SOC.(London), 8247, 481 (1958). (6) I. &I Klots, Sczence. 128, 875 (1955). (7) I. AI Klots and S.W. Luborsky, J Am. Chem. Soc., 81, 5119 (1959). (S) E. A. Balazs, A. A. Bothner-By, and J. Gergely, J , M o J , BzoJ,, 1, 147 (1959). (9) F. A. Bovey, Nature, 192, 324 (1961).
