3695
NOTES in most nitro-substituted aromatics to be due to an n - t T * transition, originating on the nitro group, which determines the lowest singlet-singlet absorption and provides a pathway for internal quenching. However, 5-nitro-8-methoxyquninoline in 1 M perchloric acid, which may be considered a model of the protonated compound in which the phenolic hydrogen cannot dissociate, is weakly fluorescent. Its emission maximum is a t 450 mp (Figure 4). This suggests that
Calorimetric Determination of the Distribution Coefficient and Thermodynamic Properties of Bromine in Water and Carbon Tetrachloride
by John 0. Hill, Ian G. Worsley, and Loren G. Hepler Department of Chemistry, University of Louisville, Louisville, Kentucky .40$08 (Received April 3, 1968)
Because of discrepancies in thermodynamic data for bromine in water and carbon tetrachloride, as discussed later in this note, and a desire to develop a method for calorimetric investigation of distribution of a solute between immiscible liquids, we have determined heats of solution of liquid bromine in carbon tetrachloride and two-phase mixtures of water-carbon tetrachloride. The results show that the distribution coefficient can be obtained from results of calorimetric measurements.
Experimental Section 200
300
400
500
600
WAVELENGTH Mp
Figure 4. Emission spectrum of 5-nitro-8-methoxyquinolinium M ) in perchloric acid (Ho= -6.7). ion
protonated 5-nitro-8-quinolinol would be fluorescent if it formed. The lowest energy absorption maximum of 5-nitro-8-quinolinol is red shifted upon increasing the solvent polarity. It occurs at 27,950 cm-l in 75% ethanol, 28,400 cm-l in chloroform, and 28,760 cm-I in carbon tetrachloride. Furthermore, the molar absorptivity of the lowest singlet-singlet transition of the protonated form is 7550. These data are inconsistent with the typical solvent polarity dependence and intensity of an n 4 T * transition. It was considered possible that a low-energy n T * transition is buried under the envelope of the “lowest” singlet-singlet absorption band, but the fluorescence of the protonated 5-nitro-8methoxyquinoline seems to contradict this. The present evidence leads to the conclusion that the failure to fluoresce of 5-nitro-8-quinolinol in acid solution is due to the failure of the excited protonated specie to form because of its extremely high acidity. The lowest singlet-singlet transition in the absorption spectrum of 5-nitro-8-quinolinol is T + T * and can be assigned to the lL, class.”
The calorimetric apparatus used for these measurements was patterned after one previously described. Temperatures in the calorimeter were followed with two 100-ohm thermistors wired in parallel and connected to a Mueller G-2 bridge with a Liston-Becker Model 14 breaker amplifier. The thermistors and a manganin heater (300 ohms) were contained in a glass spiral filled with mineral oil. All of the calorimetric data, reported here in terms of the defined calorie, were determined with 950 ml of solvent in the dewar a t 25.0 zt 0.2’. The carbon tetrachloride used was Spectrograde and bromine was 99.5% minimum assay, both from Baker and Adamson.
Results and Calculations Results of our determinations of the heat of solution of Br2(l) in CCld(1) are given in Table I. Since the
-f
Acknowledgment. The authors are grateful to Dr. Quintus Fernando for allowing fluorescence spectra to be taken on his Aminco-Bowman spectrophotofluorimeter. (11) L. Morpurgo and R. J. P. Williams, J . Chem. Soc., A , 73 (1966).
Table I : Heats of Solution of Brt in CCl4 Mol of
Br2/950 ml of
cc14
AH, koal/mol of Brx
0.01537
0.707
0.01681
0.713 0.698
0.01797
Av 0.705 & 0.007
earlier ca,lorimetric investigations of Blair and Yost2 led to AH = 0.712 * 0.010 kcal/mol for this heat of solution, we confidently adopt the following. (1) W. F. O’Hara, C. H. Wu, and L. G. Hepler, J . Chem. Educ., 38, 612 (1961). (2) C. M. Blair, Jr., and D. M. Yost, J . Amer. Chem. Soc., 55, 4489 (1933).
Volume 72, Number 10 October 1968
3696
NOTES
Br2(l) = Br2(in CCL)
(AH1" = 0.71 kcal/mol)
We define the distribution coefficient as K = m / X , in which m and X represent the molality of Br2 in the aqueous phase and the mole fraction of Brz in CC14. The numbers of moles of Brz in the aqueous phase (naq)and in the organic phase (no)are given by n,, = mW and no = Xnccl,/(l - X ) E Xnccl,, in which W is the mass of water (kg) and n c c ~ ,is the number of X nccl, is moles of CC14. The approximation no justified for dilute solutions for which X is small. The measured molar heats of solution (AH,) are related to the molar heats of solution in water (AH,,) and in CCL (ANI) by
(1)
Thus we also have AHf" = 0.71 kcal/mol for Brz (in CCL), which is the same as the value cited in the National Bureau of Standards technical note.3 We combine the AHf" = 0.71 kcal/mol for Br2 in CC14 with AH" = 7.387 kcal/mol for vaporization of Brz(l)S to obtain the heat of vaporization of Brz from solution in CC1, as Brz(in CC14) = Brz(g) (AH2" = 6.68 kcal/mol)
(2)
This value is more reliable than the AH2" = 6.92 kcal/ mol that is derived from the equilibrium measurements at several temperatures made by Kelley and Tartar.4 It might be noted that combination of this AH2' = 6.92 kcal/mol with the AH" of vaporization of Brz(l)a leads to AH1" = 0.47 kcal/mol as compared with the more reliable AHlO = 0.71 kcal/mol that has been found in two independent calorimetric investigations. We combine our AHf' of Brz(in CCl,) with AHf' = -0.20 kcal/mol for Brz(aq) based on the heat of solution of Brz(l) in water5 to obtain the heat of transfer of Brz from CC14 to H2O solution as
(naq4- no)AH, = n,,AH,,
(5)
(3)
The distribution coefficients determined by Kelley and Tartar4 at several temperatures lead to AH," = - 1.57 kcal/mol for this transfer. Because of the well known uncertainties associated with calculating AH" values from equilibrium data at different temperatures, we regard the AH3" = -0.91 kcal/mol based on calorimetric data as more reliable than the value from Kelley and Tartar.4 It might be noted that the AHf" for Br2(aq) that is cited in the National Bureau of Standards technical note3 was adopted6 before the AHfO we have used was publi~hed.~ I n order to develop and test a calorimetric method for the determination of the distribution coefficient for a solute between two immiscible liquids, we have determined heats of solution of Br2 in the two-phase HzOCcl, system. The aqueous layer was 0.01 M HClOd to repress hydrolysis of Brz. Experimental results are reported in Table 11.
Table I1 : Heats of Solution of Brz in H20-CClaTwo-Phase Systems AHIlh
Ha0
Mol of CClr
koal/mol of Bra
0.7104 0.7578 0.8525
2.445 1.956 0.9782
0.624 0.604 0.499
Mol
Kg of
of Br2
0.01487 0.01353 0.01333
The Journal of Physical Chemistry
K (eq 5 )
0.38 0.35 0.34 Av 0 . 3 6 2 ~ 0.02
(4)
Elimination of n,, and no from eq 4 in terms of K , W , and nccl, gives
Br2(in CC1,) = Brz(aq) (AH3' = -0.91 kcal/mol)
-t- noAH1
t
We substitute our calorimetric data from Table I1 into eq 5 to obtain an average K = 0.36, with an average deviation of 0.02. This value is in satisfactory agreement with K = 0.37 reported by Kelley and Tartar4 and Lewis and Storch' on the basis of analyses of equilibrium mixtures at 25". The activity data4,' justify our taking all activity coefficients equal to unity in deriving eq 5 for dilute solutions. Equilibrium data already cited permit calculation of the activity of Brz(aq) in saturated solution and thence the standard free energy of solution of bromine as follows. The data of Lewis and Storch' lead to the Henry's law constant k , = 0.539 atm for Br2 in CC14. Combination of this value with the distribution coefficient K gives the Henry's law constant k,, = 1.45 for Brz in H2O. Taking the usual solute standard state based on the hypothetical 1 m solution, we have P o = 1.45 atm and combine with the vapor pressure of Brz(l) to obtain a = P/P" = 0.28/1.45 = 0.193 for the activity of Br2 in saturated aqueous solution. Similar calculation with the data of Kelley and Tartar* leads to ko = 0.519, IC,, = 1.40, and the desired a = 0.200 in saturated solution. These activities lead to AGO = 0.97 and 0.95 kcal/mol for the standard free energy of solution of Brz(l) in HzO. For subsequent calculations, we use AGf" = 0.96 kcal/mol for Br2(aq),which agrees well with AGf" = 0.94 kcal/mol cited in the National Bureau of Standards technical note.3 (3) D. D.Wagman, W. H. Evans, V. B. Parker, I. Halow, 8. M. Bailey, and R. H. Schumm, National Bureau of Standards Technical Note 270-3,U. S. Government Printing Office, Washington, D. C. (1968). (4) C. M. Kelley and H. V. Tartar, J . Amer. Chem. Soc., 78, 5752 (1956). (5) C.H. Wu, M.M. Birky, and L. G. Hepler, J . Phys. Chem., 67, 1202 (1963). (6) W.H. Evans, 1968, personal communication. (7) G. N. Lewis and H. Storch, J . Amer. Chem. Soc., 39, 2544 (1917).
NOTES
3697
Although the principles of an earlier calculation6 leading to AG" of solution of Brz(l) in HzO are correct and the arithmetic of that calculation has been verified, careful consideration of the uncertainties associated with the solubility data and the error from neglect of solubility of HzOin Brp indicates that the earlier AGO = 1.38 kcal/mol is far less reliable than the values cited above. Combination of our AGO = 0.96 kcal/mol with AH" = -0.20 kcal/mol6 for solution of Brz(l) in HzO gives A#' = -3.89 gibbs/mol. Further combination of this AX" with the entropy of Brz(l)8gives so = 32.5 gibbs/mol for the standard partial molal entropy of Brz(aq). The entropy cited earlier6 is in error because of use of an incorrect AG" of solution, while the entropy cited in ref 3 is in error because of an incorrect AH".
Acknowledgments. We are grateful to the National Science Foundation for support of this research, which was suggested by Pitzer and Brewer.* We also thank W. H. Evans for his helpful comments. (8) Pitzer and Brewer, "Thermodynamics," revised by K. S.Pitzer and L. Brewer, McGrsw-Hill Book Co., Inc., New York, N. Y., 1961, problem 20-4.
where a(iu) and ~ ( i u )are, respectively, the polarizability of the adsorbed molecule and the dielectric constant of the continuum solid at an imaginary frequency iu. Equations 3 and 4 have been numerically integrated by Johnson and Klein4 to obtain estimates of 11 and 1 3 . These authors used an approximation for a(iu) which reproduces atom-atom dispersion energies, together with experimental data for the dielectric constant of graphite. The theoretical interaction energy els(z) = -11/z3 gives, for particular atom-surface potential models, values of so, the atom-surface collision parameter,6 and since experimental values exist for the product A , Y ~it, ~ is possible, if the surface area, A , is known, to compare 11 with experiment. Comparison of I3 with experiment is possible via the estimation of an experimental value for 133. In this note these procedures are applied to the inert gas-graphite (P33) system.
Second-Order Interactions If so is to be determined from experimental values of Aso, then A must be evaluated by a method independent of so. Such a method6 gives A = 10.7 m2 g-' (correction given in ref 3) for a monolayer adsorption model or A = 8.32 m2 g-l,' when an allowance for three dimensionality is madea8r9 Experimental values of so determined from Asovalues6 and theoretical values of so are given in Table I. The latter are calculated by equating eq 1 to the long-range limit of the atom-surface 3-p potential
Applications of McLachlan's Theory to Physical Adsorption
by J. D. Johnson' School of Chemical Sciences, University of East Anglia, Norwich, England
giving
Accepted and Transmitted by The Faradau Society (September 3, 1967)
According to McLachlan,2the van der Waals energy between an atom and a continuum surface (distance x from the surface) is given by
while the perturbation to E(r), the bulk-gas interaction energy, can be written for a monolayer as3
where p = r / u , Pm = Z m / g , ~ ( r=) 0, and Em is the distance of the monolayer from the surface. Here I1
)
= I"Jma(iu)( e(iu) - 1 du 2* 0 €(iU) 1
I3 =
+
'> +
I"Jm az(iu)(e(iu) 2n
0
€(iU)
1
du
(3)
(4)
where els* is the atom-surface interaction energy.6 The theoretical values are much smaller (>20% smaller) than experimental values suggest. If the atomsurface potential models considered' are realistic, then the theoretical I1 estimates are too small. (1) This work was supported by the Science Research Council. (2) A. D. McLachlan, Mol. Phys., 7, 381 (1964). (3) J. D. Johnson and M. L. Klein, Trans. Faraday Soc., 63, 1269 (1967). (4) J. D. Johnson and M. L. Klein, unpublished calculations. (5) J. R. Sam,, G. Constabaris, and G. D. Halsey, J. Phys. Chem., 64, 1689 (1960). (6) J. D. Johnson and M.L. Klein, Trans. Faraday SOC.,60, 1964 (1964). (7) D. H. Everett, Discussions Faraday SOC.,40, 177 (1965). (8) A monolayer model gives exactly A / B * = - 2kl21[ k2 - klB/RT]. A three-dimensional model gives A / ( B * - a) = -2k12/[kz 3klB/ R T ] (see ref 9 for details). (9) J. A. Barker and D. H. Everett, Trans. Faraday Soc., 58, 1608 (1962).
+
Volume 72,Number 10 October 1968
