207
NOTES
~onfiguration.~-QThe purpose of this note is to reinterpret the results of ref 3 (where it was assumed that each AOH molecule was bent) in terms of a linear model and t o compare values of AHOo obtained with those given previously.* Table I shows measured4-9 and estimated2vibration frequencies and molecular moments of inertia for AOH (linear). Corresponding quantities for AOH (bent) are tabulated in ref 3. Application of elementary statistical thermodynamics to the equilibrium of reaction (1) at 2475OK (the temperature of the thirdlaw calculations of ref 3) provides 1
2
3
4
5
'1C
Figure 2. Calibration curve for the correction to Stokes law radius of some ions in methanol.
the other solvents it is on the straight-line plot with the other anions. The radii of the monovalent cations in solution are decreasing from Lif to Cs+ in almost all the solvents; the behavior of the Li+ ion in acetonitrile4 and in sulfolane5 is exceptional in the sense that it has a value smaller than that of the Na+ ion. Acknowledgement. We are grateful to Dr. Joseph Steigman, Polytechnic Institute of Brooklyn, for a helpful comment. (4) S. Minc and L. Werblan, Electrochim. Acta, 7, 257 (1962). ( 5 ) R. Fernandez-Prini and J. E. Prue, Trans. Faraday SOC.,61, 1257 (1966) : M. Della Monica, U. Lamanna, and L. Jannelli, Gam. Chim. Ital., 97, 367 (1967).
K1 =
7.8
x
10-"(mAOH/mA)
$/a
x
(fvibfrot)AOH exP(- AHoo/RT) (2) I n eq 2, m represents a molecular mass andfvibfrot the product of vibrational and rotational partition functions. K1 is in units of ml molecule-1. The constant in (2) is independent of the configuration of AOH and comes directly from JANAF data.l Table I1 shows calculated zero-point standard enthalpy changes for reaction 1 for both linear (data of Table I) and bent (data of ref 3) models. The equilibrium constants used stem from the flame experiments on all five metal^.^ For all metals except lithium, AHOO (linear) and AHoo (bent) are closely similar: increases in fvib and decreases in trot which take place as the configuration is changed from bent to linear (cf. values in Table I and in ref 3) approximately compensate for one another. Recommended values of AHoo for these metals are: AHOO,N& = -325 f 16 kJ mol-l; AHOO,K= -339 9 kJ mol-I; AHoO,Rb = -347 f 9 kJ rnol-l, and AH0o,cs = -380 12 kJ mol-'. The dispersions in these values are due largely to the experimental uncertainties in values of K1.3 For lithium, AHoo (bent) exceeds AH0o (linear) by 14 kJ mol-I. However, the estimated vibration frequencies for LiOH (linear) may well be too highnote that measured frequencies9 are considerably lower than estimated frequencies2 for NaOH (linear)-and it would seem premature t o recommend decreasing the value of AHOOfor lithium to this extent at present. A flame value of AHoo for lithium of -430 i 12 kJ
*
*
Molecular Structures and Enthalpies of Formation of Gaseous Alkali Metal Hydroxides
by D. E. Jensen AeroChcm Research Laboratories, Inc., Princeton, N e w Jersey (Received J u n e 19, 1969)
09640
Reliable values for the zero-point standard enthalpy changes AHooof the reactions A@
+ OH(d zAOHk)
(1) (where A is an alkali metal atom) are frequently needed in thermochemical calculations on high-temperature systems. Values of AHOO available in the literature (e.g., ref 1-3) stem largely from third-law calculations based on experimentally determined equilibrium constants and consequently require knowledge of the rotational and vibrational contributions t o the total partition functions of AOH. Both linear and bent configurations for AOH have been proposed; the most recent work, however, indicates a linear or nearly linear
(1) "JANAF Thermochemical Tables," Dow Chemical Company, Midland, Mich., June 1968. (2) H. Smith and T. M . Sugden, Proc. Roll. SOC.,A219, 204 (1953). (3) D. E. Jensen and P. 6. Padley, Trans. Faraday Xoc., 62, 2132 (1966). (4) R. L. Kuczkowski, D. R. Lide, and L. C. Krisher, J . Chem. P h y s . , 44,3131 (1966). (5) D. R. Lide and R. L. Kucnkowski, ibid., 46, 4768 (1967). (6) N. Acquista, S. Ahramovitn, and D. R. Lide, ibid., 49, 780 (1968). (7) C. Matsumura and D. R. Lide, ibid., 50,71 (1969). (8) D. R. Lide and C. Matsumura, ibid., 50, 3080 (1969). (9) N. Acquista and 8. Abramovitz, to be published.
Volume 74, Number 1 January 8 , 19YO
NOTES
208 Table I : Molecular Parameters of the Alkali Metal Hydroxides0 Vibration frequenoiea
P
Molecule
Y1
Y I
LiOH NaOH
(800) 43 1 (400) 354 336
(410) 337 (340) 309 310
KOH RbOH CsOH
va
-
(3800) (3650) (3610) (3610) (3610)
Partition functions at 2476'K 10%.4-
0
(1.52) (1.94) (2.27) 2.31 2.40
10881
fv ib
10 -afrot
0.22 0.66 1.09 1.34 1.55
68 161 171 226 238
1.36 4.1 6.7 8.2 9.5
a Vibration frequencies are in om-', internuclear distances rA-o in ern and molecular moments of inertia I in g em2. Vibration frequencies and moments of inertia in parentheses are estimates;a those not in parentheses stem from recent experiments. 4 - g
Table I1 : Calculated Zero-Point Standard Enthalpy Changes" K1,247p5
Metal
Li Na K Rb cs a
ml molecule-1
7.7 x 8.3 x 2.2 x 4.0 X 2.2 x
10-18 10-18 10-17 10-1' 10-16
AXoa (bent),
AH"a (linear),
kJ mol-]
kJ mol-1
-423 - 322 - 338
-437 - 325
-348 - 382
- 339 - 347 -380
K1 is calculated from equilibrium constants for the reaction
+ H20 Ft AOH + H given in ref 3, the equilibrium constant for the reaction HzO Ft OH + H being taken to be 3.0 X lo1' A
molecule ml-l at 2475°K (JANAF' data).
mol-' (dispersion due approximately equally to uncertainties in K1 and in fvibfrot) appears to be the best compromise until reliable experimental values for the molecular parameters of LiOH become available.
Acknowledgment. The support of the Office of Naval Research under Contract Nonr-3809 (00) is gratefully acknowledged.
n-T
* Transition in the Dimethylthioacetamide-
Iodine and Thioacetamide-Iodine Complexes by Arthur F. Grand and Milton Tamres Department of Chemistry, University of Michigan, Ann Arbor, Michigan 48104 (Received September 3, 1969)
Donor-acceptor interactions can be investigated by measuring either a property of the complex,or a perturbed property of one of the components. In the specific case of the iodine interaction with thioacetamide, CHG(S)NHZ, thermodynamic data for the complex in dichloromethane' have been obtained from a study of the charge-transfer (CT) band at 286 mp. In addition, on the low wavelength side of the CT band, the n + n* band for thioacetamide was found shifted from 269 to 248 mp. It was located in the spectrum of the complex because of its high intensity which was comparable to that of the unperturbed transition. The iodine complex of dimethylthioacetamide, CH&T h e Journal of Physical Chemistry
(S)N(CH&, in carbon tetrachloride has been investigated utilizing the blue-shifted iodine band in the visible region.2 The same complex has also been characterized by studying the charge-transfer band in the ultraviolet region. Unlike other thione-iodine complexes, however, a pronounced shoulder was observed on the high wavelength side of the CT band. This shoulder is attributed to a blue-shifted n -+ n* transition in the donor, with an intensity greatly enhanced over that in the free donor. A similar band is also found to exist in the thioacetamide-iodine system that was not recognized previously.
Experimental Section The purification of carbon tetrachloride, iodine, and dimethylthioacetamide have been described previously. Thioacetamide was obtained from Eastman Organic Chemicals and was recrystallized from anhydrous ethyl ether. Its ultraviolet spectrum agreed with that in the literature.1 Reagent grade dichloromethane was dried and distilled prior to use. The procedure to obtain spectrophotometric data and the mathematical treatment to determine spectral and thermodynamic properties also have been described. Results and Discussion The composite band of the dimethylthioacetamideiodine complex in carbon tetrachloride is shown in Figure 1 (curve A). It was obtained by subtracting the absorbances of the equilibrium concentrations of free donor and of free acceptor from the total absorbance of the mixture. The band has a maximum at -310 mp and a distinct shoulder at -340 mp. In addition, the curvature at the low wavelength indicates a contribution from the tail of another band, possibly the perturbed R 3 T * band of the donor. One way to test for the presence of more than one species in a complex band is to use the Liptay m e t h ~ d . ~ When this was done for dimethylthioacetamide-iodine (1) R. P. Lang, J . Amer. Chem. SOC.,84, 1185 (1962). (2) R. J. Niedzielski, R. S. Drago, and R. L. Middaugh, ibid., 8 6 , 1694 (1964). (3) A. F. Grand and M.Tamres, Inorg. Chem., 8,2495 (1969). (4) W. Liptay, 2.EZeMrochem., 65, 375 (1961).
