2381
Solute-Solvent Interactions in Water-tertSutyl Alcohol Mixtures (13) P. Gallezot, Y. Ben Taarit, and B. Imelik, J. Phys. Chem., 77, 2556 (1973). (14) C. C. Chao and J. H. Lunsford, J. Am. Chem. Soc., 93, 71 (1971). (15) P. H. Kasai and R. J. Bishop, Jr., J. Am. Chem. SOC., 94, 5560 (1972). (16) C. Naccache, M. Che, and Y. Ben Taarit, Chem. Phys. Lett., 13, 109 (1 972). (17) M. V. Mathieu and P. Pichat, “La Catalyse au Laboratolre et dans I’lndwtrle”, Masson, Paris, 1967, p 319. (18) L. H. Little, “Infrared Spectra of Adsorbed Species”, Academic Press, London, 1966, p 84; N. G. Connelly, lnorg. Chim. Acta Rev., I, 47 (1 972). (19) Y. Ben Taarit, C. Naccache, and B. Imelik, J. Chim. Phys., 70, 728 (1 973). (20) A. Lubezky and M. Folman, Trans. Faraday Soc., 67,3110 (1971). (21) C. C. Chao and J. H. Lunsford, J. Am. Chem. SOC., 93, 6794 (1971). (22) (a) A. Zecchina, E. Garrone, C. Morterra, and S. Coluccia, J. Phys. Chem., 79, 978 (1975); (b) E. Garrone, G. Ghiotti, S.Coluccia, and A. Zecchina, ibid., 79 984 (1975). (23) C. Naccache and Y. Ben Taarit, J. Chem. SOC.,Fara&y Trans. 1, 69, 1475
(28) (29) (30) (31) (32) (33) (34)
( 1973). J. R. Norton, J. P. Collman, G. Dolcetti, and W. T. Roblnson, Inorg. Chem., 11,382 (1972). D. M. Adams, “Metal Ligand and Related Vibrations”, Edward Arnold, London, 1967. p 292. F. A. Cotton and G. Wilkinson, “Advanced Inorganic Chemistry”, Interscience, New York, N.Y., 1972, p 716. J. P. Collman, P. Farnham, and 0. Dolcetti, J. Am. Chem. SOC.,93, 1788 11971). H. iunsford, J. Phys. Chem., 72, 4163 (1968). E. Cohen de Lara and J. Vincent Geisse, J. Phys. Chem., 76, 945 (1972). P. Tarte, J. Chem. Phys., 20, 1570 (1952). M. Primet, unpublishedresults. M. Primet, J. M. Basset, E. Qrbowski, and M. V. Mathieu, J. Am. Chem. SOC., 97, 3655 (1975). H. Dunken and H. Hobert, 2.Chem., 3,398 (1963). D. W. Breck, “Zeolite Molecular Sieves”, Wiley, New York, N.Y., 1974, p 636.
2.
Solute-Solvent lnteractions in Water-terf-Butyl Alcohol Mixtures. 7. Enthalpies of Transfer for LiCl and HCI as Obtained through Dilution of an Aqueous Concentrated Electrolyte Solution in Hydroalcoholic Media Yvon Pointud, Jean-Pierre Morel, and Jean Julllard’ Laboratoire d’Etude des lnteractions Solutes-Soivants, Universite de Clermont, B.P. 45, 6317O-Aubier8, France (Received March 16, 1976) Publication costs assisted by the Universite de Clermont
Enthalpies for the process of transferring Li+, C1- and H+, C1- from water to mixtures containing up to 40% of tert- butyl alcohol are reported. They are derived from experimental heats of dilution of an aqueous concentrated solution of the electrolyte, both in pure water and in hydroalcoholic mixtures. Corresponding equations are obtained from classical thermodynamics of solution. It is shown that the enthalpies of transfer are related to experimental enthalpies, relative molal enthalpies of water, and dilution enthalpies in the two media. Data thus obtained are discussed in terms of structure of the water-tert- butyl alcohol mixtures.
Introduction Enthalpies of transfer of electrolytes from one solvent to another are generally obtained by comparing enthalpies of solution in the two solvents. Such a method can only be employed in the case of dry solid electrolytes sufficiently soluble in the two solvents to obtain noticeable thermic effects. Recently, the reverse method in which enthalpies of transfer are derived from enthalpies of precipitation of insoluble compounds has been proposed and developed by Bright and Jez0rek.l As far as gaseous solutes are concerned direct calorimetric measurements of enthalpies of solution although possible2 are rather difficult and transfer enthalpies are generally obtained from Gibbs free energy changes with temperature. In the course of a continuing study3 on thermodynamic quantities related to the transfer process of electrolytes from water to water-tert- butyl alcohol (TBA) mixtures, it appeared useful to record enthalpies of transfer of both a gaseous electrolyte (HCl) and an hygroscopic solid electrolyte (LiCl). Therefore the dilution of a concentrated aqueous electrolyte solution into both water and water-alcoholic solvents has been considered. Using a formal thermodynamic treatment it can
be shown that enthalpies of transfer can be derived from the experimental enthalpies thus obtained, changes in partial molal enthalpies of water, and dilution enthalpies in the two media. Such a treatment is here presented in the first part of this paper. This method is thus applied to the determination of the transfer of HCl and LiCl from water to water-TBA media (with a TBA content varying from 10 to 40 wt %). Data concerning the other alkali metal have been reported earlier.4 All these results concerning thermodynamic functions of transfer (C, H,S ) can then be considered.
Experimental Section Heat measurements were performed using an LKB solution calorimeter (Type 8 700). The general experimental procedure has been described elsewhere.5 Ampoules containing weighed amounts (500 mg) of a stock aqueous solution of HC1 or LiCl were crushed into 100 mi of either the mixed solvent or pure water. Final concentrations thus obtained are of the order of 2X mol kg-l. KC1 and LiCl concentrations in stock solutions are determined through conductometric titration of a chloride ion by silver nitrate. HCl stock solutions have been prepared from The Journal of Physical Chemlstry, Vol. 80, No. 21, 1976
Y.
a Merck "titrisol" and their concentrations checked through precise indicator titration. Purification of both solvents (water and TBA) has been done as previously r e p ~ r t e d Solvents .~ properties used in the calculations have been tabulated in a former paper.4
Theoretical Small amounts of a concentrated aqueous solution of the electrolyte were poured into both water (process A) and the hydroorganic mixtures (process B). Process A is classical and well defined; it is a dilution. Process B is, on the contrary, more complicated since a change in the composition of the solvent media also occurs. In order to establish rigorously the relationship between the experimental heats and the standard enthalpy of transfer, we use the mathematical formalism relative to the functions of three variables, since a ternary mixture is here involved. At constant temperature and pressure the state of the system is defined by the number of moles of the three components: water, cosolvent, and electrolytes, n,, n,, ne, respectively. The intensive nature of the partial molar quantities, which are homogeneous functions of the zero degrees, allows the use of the reduced variables
n,--1, -n,-- r ,
ne and-=p n, n, n, that is to say, the molar ratio relative to water. The enthalpy of the system is then H(n,,n,,n,) = nwRw(l,r,p)+ nsRs(l,r,P)
The difference between these two enthalpies can be identified as the integral heat of dilution for the solute in the mixture of composition (l,rf,O):
= -fldli[(l,rf,Pf)
(1,rf,0)] (4)
Process A The aqueous concentrated solution of composition (l,O,p,) which contains nw*moles of water and ne* moles of solute is here poured into N,* moles of water. The experimental molar enthalpy is thus G
A
=
- &dil[(l,O,Pf*)
&A'
-
(1,0,0)]
(5)
and the molar enthalpy of the process at infinite dilution: G A B=
(~/PI)[Rw(~,O,O) - Rw(l,O,Pi)]
+ [Re'(l,O,O) - Re(l,O,~i)] (6) Enthalpy of Transfer. The combination of the previous equations yields G
B
-
+ Hdil[(l,rf,Pf)
(1,rfyO)l - M
A
- fldll[(l,O,Pf*) (1,0,0)] Re'(l,rf,O) - Res(l,O,O) + (l/pl)[Rw(l,r~,o) - R,(l,o,o)l + Nw -[H,(l,rf,O) - Rw(l,rl,o)l
(1)
The infinitely diluted state of the solute is noted as by the superscript 8. So, the expected standard enthalpy of transfer is = Re'(I,r,O)
-
+ [Re(l,rf,pf) - R e ' ( ~ , r f , ~ ) ]
+
+ nZe(lrr,P)
&ts
Pointud, J.-P. Morel, and J. Juillard
ne
, +N - [H,(l,rf,O)- Rs(l,rl,o)l
- R,'(I,o,o)
Process B The aqueous concentrated solution of composition (l,O,pi) which contains n, moles of water and ne moles of solute is poured into an initial mixture of composition ( l,rl,O)which contains N, moles of water and N , moles of cosolvent. A ternary mixture (l,rf,pf)which contains (N, n,) moles of water, N, moles of cosolvent, and ne moles of solute is then obtained. Under these conditions, the experimental enthalpy per mole of solute is
+
ne
ne
The first two terms of the right-hand side correspond to the standard enthalpy of transfer for the electrolyte from water to the final mixture. This enthalpy of transfer is therefore
-
atB = A H B - A H A - A T c + md,l(mixture)
- Aiiddwater)
m,,
(8)
where the heat of dilution of n, moles of water in ( N , N,) moles of the mixture divided by the number of moles of solute ne, represents the three last terms in eq 7. can be obtained breaking ampoules of water in the mixture in such a way that the number of moles n,, N,, and N , be identical (or as near as possible) with the number of moles used in process B. If the electrolytes final concentrations are low enough, the molar enthalpies of dilution can both be calculated using the Debye-Huckel relationship. The corresponding complete equation has recently been established within the scope of the Bjerrum theory for 1-1 electrolytes fully dissociated.*p6 Equation 8 is, as far as the physical meaning of its various terms is concerned, identical with an equation previously derived from the MacMillan-Meyer formalism during the study of the experimental process of mixing dilute solutions of electrolytes and nonelectrolytes (eq 11 in the paper by Desnoyers et aL7). It is worth noting that if n, is negligible as compared to N,, the last two terms in eq 7 are then equivalent to
+
mc
+
(l/ne)[N, dRw(l,rl,O) N, d R s ~ l , r l , O ~ ~ The Journal of Physical Chemistry, Vol. 80, No. 21, 1976
(7)
2383
Solute-Solvent Interactions in Water-tertSutyl Alcohol Mixtures
TABLE IV: Enthalpies of Transfer for H+, C1- following Various Authors AHte(kJ mol-')
TABLE I: Relative Molal Enthalpy of Water [Bw(mixture)- Rw(water)]in J mol-' X" 5 10
15
20
-twb
Xa
-t,c
11
8
67 242 502
96 243 506
-t,b
25 30 35 40
-t,c
647 711 750 777
636 703 736 761
Weight percent of TBA. Our data. Interpolated from ref 8 by H. Gillet, L. Avedikian, and J. P. Morel, Can.J . Chem., 53, 455 (1975). a
X"
Das
RovC
Our data
10
6.96 7.32 4.18
3.18 5.49 0.99
2.6 3.4 -2.0
20 40
a Weight percent of TBA. Calculated from 4G't and 4St given in ref 12. c From ref 11.
TABLE 11: Comparison of Transfer Enthalpies for K+, C1- from Water to TBA-Water Mixtures as Obtained through Different Methods ( AHte,kJ mol-')
X"
b
c
X"
b
C
10
2.82 7.72
2.77 7.60
30 40
8.72 7.64
8.31 7.27
20
X is the weight percent of TBA. Previous data from heats solution of the solid electrolyte. These measurements. TABLE 111: Enthalpies of Transfer for H+,C1- and Li+, C1- from Water to Water-TBA Mixtures Xa 5 10
15 20 a
H + , C P b L i + , C k b Xa 1.31 2.58 3.62 3.38
0.91 1.81 3.36 4.79
25 30 35 40
,
H + , C F b Li+,C1-b 1.99 0.72 -0.72 -1.99
5.96 5.87 5.53 5.01
Weight percent of TBA. 4Htein kJ mol-l.
Figure 2. Entropies of transfer in the molal scale (same notations as
in Figure 1).
Then, if the relative molal enthalpy t, is known (generally using the same approach) can be calculated using this approximation, which is justified for the experimental process involved here.
mc
\.
\
%w
\
\
\ '\
Flgure 1. Enthalpy of transferfrom water to water-TEA mixtures: NaCI, 1; KCI, 2; RbCI, 3; CsCI, 4, from ref 4; LiCl (0),and HCI ( 0 )this
work.
which is null in accordance to the Gibbs-Duhem relationship. It yields
iw, 1 [H,(water) - R,(mixture)] N
Pi
-
= -E, Pi
Results a n d Discussion In Table I are summarized data concerning changes of partial molal enthalpies of water when going from water to water-TBA. These results have been obtained from dilution of water in the mixtures. They agree fairly well with previous values calculated by Kentamaas from heat of mixing of water and TBA. Since, the method used here is expected to be more precise and has the advantagge to be coherent with the following experiments, these new data will be prefered in the forthcoming calculations. Enthalpies of transfer of KCl from water to some TBA mixtures have been previously derived from enthalpies of solution in these media.4 As a test of the method proposed here, new values were recorded. As can be seen in Table 11, they compare favorably to the previous ones, discrepancies being a t worst 0.4 kJ. Experimental heat measurements for dilution of LiCl and HCl concentrated aqueous solutions in water and water-TBA mixtures are available as supplementary material to this paper. (See paragraph at end of text regarding supplementary material.) Only standard enthalpies of transfer estimated as described in the theoretical part are given in Table 111. Their variations vs. the TBA content are shown in Figure 1 in addition to former results concerning Na, K, Rb, and CsCl in the same media. The Journal of Physical Chemistry, Voi. 80, No. 2 1, 1976
L. J. Gerenser, M. G. Mason, and P. J. Trotter
Using Gibbs free energies of transfer of LiCP and HCllO in the same media, obtained from galvanic cell measurements, entropies of transfer have been calculated and are shown in Figure 2. Such entropies have been estimated for HCl, by two groups of authors,11J2from the temperature coefficient of the standard potential of the Agl AgCl electrode in water-TBA mixtures. The reported AH are compared to our values in Table IV. Great discrepancies are observed. As a matter of fact, A S values obtained from AG at various temperatures are very sensitive to small variations of AG. In support of this opinion it must be noted that both Royll and Bose12 using the same standard potential data recorded by Roy obtain quite different values for the enthalpies of transfer. Therefore, it is our feeling that both of these results are in error. In this perspective, a systematic comparison of data obtained by our method for HC1 to previous AH derived from galvanic cell in various hydroorganic mixtures is underway. From Figures 1 and 2 it appears that the behavior of LiCl in water-TBA of both enthalpies and entropies of transfer are situated at 0.08 mole fraction of TBA (27 wt %). With the exception of the media richer in water ( x < 0.04), the importance of the effect of transferring an alkali metal ion from water to water-TBA increases following the sequence Li+ < Cs+ < K+ < Na+. This order, observed here for enthalpies and entropies, has already been noted for Gibbs free e n e r g i e ~molar ,~ volumes, and heat capacities.13We are working on qualitative and quantitative models able to explain all the thermodynamic property changes of these solutes in these media. More extensive discussion is thus delayed and will appear in a future paper. As far as the transfer of HC1 is concerned, it is worthwhile to note that the nature of the effects is quite different from those observed for alkali metal chlorides. Transfer entropies
as well as enthalpies are a t maximum for a mole fraction of t-BuOH of 0.05 which is supposed to correspond to a maximum structuration of the TBA mixtures, a model which is supported by some thermodynamic and spectroscopic arguments.14-16 Supplementary Material Available: Additional experimental heat data for dilution of LiCl and HC1 concentrated aqueous solutions in water and water-TBA mixtures (2 pages). Ordering information is available on any current masthead page.
References and Notes L. L. Bright and J. R . Jezorek, J. Phys. Chem., 79,800 (1975). C. E. Vanderzee and J. D. Nutter, J. Phys. Chem., 67, 2521 (1963). Part VI, N. Dollet and J. Juillard, J. Solution Chem. , 5, 77 (1976). Y. Pointud, J. Juillard, L. Avedlklan, J. P. Morel, and M. Ducros, Thermhim. Acta, 8, 423 (1974). (5) L. Avedikian, J. Juillard, J. P. Morel, and M. Ducros, Thermochim. Acta, 6, 283 (1973). (6) J. Juillard, J. P. Morel, C.R. Acad. Sci., Ser. C, 277, 825 (1973). (7) J. E. Desnoyers, G. Perron, L. Avediklan, and J. P. Morel, J. Solution Chem., in press. (8) J. Kentamaa, E. Tommila, and M. Marti, Ann. Acad. Sci. Fenn. All, 93, 3 (1959). (9) Y. Pointud, J. Juillard, J. P. Morel, and L. Avedikian, Electrochim.Acta, 19, 229 (1974). (10) J. P. Morel and J. Morin, J. Chim. Phys., 67, 2018 (1970). (11) R. N. Roy, W. Vernon, and A. L. M. Bothwell, J. Chem. SOC.A, 1242 (1971). (12) K. Bose. A. K. Das, and K. K. Kundu, J. Chem. SOC.,Faraday Trans. 1, 1838 (1975). (13) L. Avedikian, G. Perron, and J. E. Desnoyers, J. Solution Chem., 4, 331 (1975). (14) M. C. R . Symons and M. J. Blandamer, "Hydrogen-Bonded Solvent Systems", A. K. Covington and P. Jones, Ed., Taylor and Francis, London, 1968, p 211. (15) E. K. Baumgartner and G.Atkinson, J. Phys. Chem., 75, 2236 (1971). (16) F. Franks and D. S. Reid, Water: Compr. Treatise, 2, 356 (197% and following. (1) (2) (3) (4)
Bonding in Silver Thionamides Studied by Infrared, Laser-Raman, and X-Ray Photoelectron Spectroscopy L. J. Gerenser, M. G. Mason,' and P. J. Trotter Research Laboratories, Eastman Kodak Company, Rochester, New York 14650 (Received January 19, 1976: Revised Manuscript Received Ju/y 8, 1976) Publication costs assisted by the Eastman Kodak Company
Silver thionamide compounds have been structurally characterized. Laser-Raman, infrared, and x-ray photoelectron (ESCA) spectroscopies were employed in a complementary manner for analysis of tautomeric forms and bonding characteristics in these compounds. ESCA binding energies have been interpreted with the aid of CNDO and extended-Huckel molecular orbital calculations. Silver thionamides were shown to form mercaptide structures of the type -N=C-S- -Ag. Thionamide silver nitrate complexes exhibited thione-silver ion coordination, >C=S- - -Ag+.
Introduction Several spectroscopic studies of silver complexes utilizing the infrared method have been rep0rted.l-8 Thionamide silver complexes were found to be particularly difficult to characterize by means of ir alone.3~6This difficulty is due primarily The Journal ot Physical Chemistry, Voi. 80, No. 21, 1976
to large uncertainties in assigning thione C=S vibrationsgin H-N-C=S containing ligands.10 In silver complexes of thionamides, the question arises whether an Ag- -N-C=S complex is formed leaving the thione group i n t a ~ t , or ~,~ whether silver may form a mercaptide complex of the structure -N=C-S- -Ag. Since the precise structure and position
