NOTES
580
17 (*5-15) cal/deg mole, obtained by Hepler and coworkers,2o from their thermal data and Kolthoff and Tomsicek's standard oxidation potential. Since the values of -y* used in this study are probably uncertain to about *0.005, a recalculation of the entropy of ferricyanide ion is not warranted at this time. Hepler and co-workers used T * = 0.122 for a saturated solution of KsFe(CN)a,*whereas the value obtained by the approximate methods discussed in this investigation is T * = 0.118. This difference leads to a difference in AGO of solution for potassium ferricyanide of 0.06 kcal/mole and a difference of 0.2 in the entropy of ferricyanide ion. Accordingly, we take 3" = 22.8 f 0.3 cal/deg mole as the entropy of ferrocyanide ion. The entropy of ferrocyanide ion can be combined with the heat of solution20 of K4Fe(CN)6*3HzO(s) and the free energy of solution of this salt (calculated using m (satd) = 0.8559,21T*(satd) = 0.0478 (based on the solute standard state established in this work), U ~ , O (satd soln) = 0.9623 (calculated from 9 = 0.498 given in ref 8) to yield AGoBoln = -RT In (4m)%. -yf6ua203 = 6.243 kcal) to give ASoaoln = 23.16 cal/ deg mole, from which we computelg So = 147.8 cal/deg mole for K4Fe(CN)6.3H20(s)at 25". of K4Fe(CN)6.3HzO(s) From the heats of and K4Fe(CN)6(s) and AH" at 25' for the process1g HzO(l) = H,O(g), we compute for the reaction: K4Fe(CN)6.3H20(s) = K4Fe(CN)&) 4- 3HzO(g), M" = 35.10 kcal. Schottkyz2 reports 7.35 mm at 20" for the dissociation pressure of the trihydrate. From the above AH" for the reaction we calculate 10.30 mm for the dissociation pressure at 25". From this dissociation pressure we compute A G O 2 9 8 = 7.65
kcal for the above reaction, and hence ASo = 92.08 cal/deg. This yields So = 105 cal/deg mole for the entropy of K4Fe(CN)6(s) at 25". Although this is not an unreasonable value for the entropy of this salt considering that Stephenson and Morrow report2a So = 100.4 cal/deg mole for K8Fe(CN)6(s)at 25", it may be too low a value since, if one considers that the average entropy per mole of hydrated water is 9.4 cal/deg mole,24the entropy of K4Fe(CN)&) may be estimated from the entropy of K4Fe(CN)6-3H20(s) as 119.6 cal/deg mole. The most likely source of error in the calculation of the entropy of KaFe(CN)e(s) is the dissociation pressure of K4Fe(CN)6*3Hz0, since the authorz2emphasized the difficulty in obtaining reliable dissociation pressures in his measurements. From the available thermal data1*and the entropies of the ion^,'^^^^ together with the entropy of ferrocyanide ion determined in this investigation, we compute the standard free energies of formation from the elements of ferrocyanide and ferricyanide ions at 25" as 167.1 and 175.6 kcal/mole, re~pectively.~~ (21) R. H. Vallance, J . Chem. Soc., 1328 (1927). (22) H. Schottky, 2. Physik. Chem., 64, 415 (1909). (23) C. C. Stephenson and J. C. Morrow, J. A m . Chem. SOC.,78, 275 (1956). (24) See ref 19, p 364. (25) NOTE ADDEDIN PROOF. R. H. Busey, J . Phys. Chem., 69, 3179 (1965), has recalculated the entropy of &Fe(CN)&) and obtained S o = 101.8 cal/deg mole at 25'. This entropy is 1.4 units larger than the previously reported value and leads t o an increase of 1.4 cal/deg mole in the entropies of Fe(CN)c'-(aq), K'Fe(CN)v 3HzO(s), and K'Fe(CN)+j(s) reported in this paper. The revised values for the entropies are 24.2, 149.2, and 106 cal/deg mole, respectively. AGt" and AH!' values reported here are unaffected by
this change.
NOTES Temperature Dependence of Absorption of Liquid Water in the Far-Ultraviolet Region
conflicting results have been Only a few data are available in these reports on the absorption a t other temperatures. For photochemical studies in dilute aqueous solutions in the far-ultraviolet region,'
by M. Hahnann and I. Platzner Isotope Department, The W e i m a n n Institute of Science, Rehouoth, Ierael (Received September 16, 1966)
For the absorption spectrum of liquid water in the 200- to 180-mp region at room temperature, several The Journal of Physical Chemistry
(1) J. Barrett and J. H. Baxendale, Trans. Faraday Soc., 56, 37 (1960). (2) J. Barrett and A. L. Mansell, Nature, 187, 138 (1960). (3) J. L. Weeks, G. M. A. C. Meaburn, and 8. Gordon, Radiation Res., 19, 559 (1963). (4) M. Halmann and I. Platsner, J . Chem. SOC.,1440 (1965).
NOTES
58 1
Table I Temp., OC. 24.0
21.5
k k k k
(crn.-l), (cm.-l), (cm.-I), (cm.-I),
ref. 1 ref. 2 ref. 3 this work
1
I
1.67
25.0
30.7
30.5
1.2 1.46 1.80
2.09
2.50
1.47
l
1
I
I
l
80.00
45.5
62.5
80.0
2.63
4.55
7.05
the emission a t this wave length from low-pressure mercury lamps, the absorption data listed in Table I were obtained. Our results agree with that a t 25.0" by Barrett and Mansel12 and are considerably lower than those reported by Weeks, et aL3
Phosphorus-31 Chemical Shifts of Quaternary Phosphonium Salts
by Samuel 0. Grim, William McFarlane, Edward F. Davidoff, and Tobin J. Marks' Department of Chemistry, University of Maryland, College Park, Maryland (Received July 12, 1966)
182
186
184
188
mP
Figure 1. Absorption of liquid water (in k, crn.-I) a t several temperatures: 0, ref. 1; A, ref. 2; 0, ref. 3; 0, present work.
it was necessary to obtain reliable values for the absorption over a wider range of temperature. Experimental Section Absorption spectra were measured with a Zeiss PMQII far-ultraviolet spectrophotometer, in which the whole optical path was flushed with nitrogen (99.9% pure; 2 1. min.-' flow rate). Absorption spectra of triply distilled water (redistilled from alkaline permanganate and then from phosphoric acid in an allglass still) in Suprasil quartz cells (1.00-mm. path length) were measured against empty cells in a temperature-controlled cell holder ( zkO.5"). Results The absorption spectrum of water a t different temperatures is presented in Figure 1, which includes available literature data. For 184.9 mp, which is particularly important in photochemistry because of
It has been shown that the phosphorus chemical shifts of tertiary and secondary phosphines can be predicted very accurately from empirically determined additive group contributions ,)'a( which have been assigned for phenyl and various alkyl groups2 This type of behavior had been suggested previously for trivalent phosphorus compounds but the predictions were only moderately successful, and similar attempts to predict the phosphorus chemical shift in "quadruply connected" phosphorus compounds, such as phosphine oxides, were unsuccessful. Groenweghe, Maier, and Moedritzer4 observed a regular variation in the P3' chemical shift of a tertiary, secondary, or halophosphine, or a "quadruply connected" organophosphorus compound upon stepwise substitution of the organic groups by other organic groups. In addition, two very recent theoretical treatises5v6 (1) National Science Foundation Undergraduate Research Participant, 1963-1965. (2) S. 0.Grim and W. McFarlane, Nature, in press. (3) J. R. Van Wazer, C. F. Callis, J. N. Shoolery, and R. C. Jones, J . A m . Chem. SOC.,78, 5715 (1956). (4) L. C. D. Groenweghe, L. Maier, and K. Moedritzer, J. Phys. Chem., 66, 901 (1962).
Volume 70, Number 2 February 1966
