NOTES
0.5
2225 shown in Figure IC. The esr spectra of radicals produced6 on irradiation of ONB are shown in Figure lb. The experimental results obtained indicate that (a) the intensity of the new ir bands and the melting of the solid increase with time of irradiation. The similarity of the ir spectrum with that of o-nitrosobenzanilide confirms the existence of the photochemical rearrangement according to reaction 1; (b) contrary to earlier the unphotolyzed ONB, the orange-red, and the deepyellow forms of ONB all produce after melting a similar ir spectrum, but different from the nonmelted unphotolyzed ONB crystal; (c) in the flash photolysis6 of ONB 460 nm is in heptane solutions a transient with, , ,A observed which is somewhat similar to the reflectance spectrum and is assigned to the quinoid form
I
-
X. nm
Qc=.-Q
/ \
N-OH
II
Q I 4000 3000
1 2000
'11 ' 1500
I 1200
I I 1000 900
I 1
I
800
700
CM-' Figure 1. (a) Reflectance spectra produced on illumination for 1.5 hr of o-nitrobenzylideneaniline with light of X >320 nm: top points taken immediately after exposure, and succeeding spectra taken after standing in darkness for 8, 21, 48, 64, and 140 days, respectively. (b) Esr spectra of radicals produced on photolysis of o-nitrobenaylideneaniline: I, orange-red crystalline form, 11, deep yellow crystalline form, 111, in heptane solution ( 4 x 10-8 M). (c) I r solid spectra of o-nitrobenzylideneaniline m p 69.j0,I; after 85 hr exposures to light, mp 120-140°, I1; and of o-nitrosobenzanilide, m p 171", 111.
yellow color turns to orange-red and this color persists for many hours until, after about 20 hr illumination, a sudden change takes place and the orange-red color turns to a deep yellow: deep yellow spots first appear on the surface of the orange-red powder from which the new color spreads around and covers all the material in a matter of minutes, like a phase transformation under illumination. These photo-induced changes were followed using reflectance spectroscopy, esr, and ir spectroscopy. The reflectancespectrum (using a zeiss spectrophotometer P M Q l l with reflectance attachment) of the orange-red color is shown in Figure la, using the unexposed ONB as the reflectance standard and the Kubelka-Munk equation to derive4 the spectrum, , , ,A 490 nm. The fading of the orange-red color with time is also shown (Figure la), and the kinetics was found t o follow a first-order process with k 2.5 X min-'. The infrared solid spectra of ONB, of ONR after 85 hr photolysis (deep yellow form), and of o-nitrosobenzanilide (prepared from the photolysis of ONB in benzene, extracted product has a mp 171") are
-
-
the precursor of o-nitrosobenzanilide. However, since the ir bands of the nitroso compound are observed in both the orange-red and deep-yellow forms, it would appear that the orange-red color is not associated witth the quinoid form of ONR; (d) furthermore, the similarity of the esr spectra of both the orange-red and deep-yellow forms would indicate that the radical is not connected with the orange-red color and is probably the result of a simultaneous photochemical process. (4) G . Kortum, M. Kortum-Seiler, and S. D. Bailey, J. Phys. Chem., 66, 2439 (1962). (5) We thank Dr. M. Arick for measuring the esr spectra. (6) E. Hadjoudis and E. Hayon, in preparation.
Activity Coefficientsin Mixed Solutions.
Prediction Of Harried from Ionic Entropies
by J. V. Leyendekkers Diviswn of Fisheries and Oceanography, CSIRO, (Received November 96,1969) SVdneU %@o, AU&dk
The Harned equations log rl/rlo=
-a1212
(1)
log
-a2111
(2)
Y2/Y20
=
apply to a large number of two-electrolyte systems.' Preliminary analysis of available data indicates that for a system of two electrolytes i and j with a common anion the coefficients can be estimated simply from the relationship (eq 3) Volume 74, Number 10 Maa 14, 1970
NOTES
2226 Table I : Conventional Ionic Entropies at 25” Computed Relative to So=t = 0 in the Hypothetical Standard State of 1 g-ion/kg of Waterlb Code-1
2
3
4
5
Ion
H+
Sei, cal deg-1
0
Li+ 3.4
Na+ 14.4
K+ 24.5
Cs+ 31.8
7
7
6
MgZ+
Sr2+
Cas+ -13.2
-28.2
9
8
Baz+ 3
-9.4
11
10
Ala+ -74.9
Ce3 -38.3“ +
c113.2
mol-’ a
Calculated from eq 4.
Table I1 : Experimental Values of Harned Coefficients a t 25” I, m
2
a
I,
System 2
1
ai2
0 .0O5lc 0.031 0.0575 0.099 0.058 0 . 0653l’-0. 06175m 0.063 0.0908 0.03168 - 0.03768 0.067 -0.032 0.1115b -0.0165b 0 .O23O6O - 0. 00846” 0 . 0446d - 0.008’d -0.0135e1‘ - 0 , 0056g,h 0.01 1 1 6 ’
HC1-LiCI HCI-NaCl HC1-KCl HC1-CsCl HC1-SrC12 HC1-BaC12 HCl-AlCL HC1-CeCl3 LiC1-NaC1 LiCI-KC1 LiC1-CsCl NaC1-KC1 NaCl-CsC1 NaC1-MgC12 NaC1-CaC12 NaC1-BaC12
System 1 2
m
a21
2.0
0.1 0.2 0.5 0.7
6
a 21
Dl12
KCl-CsC1 KC1-CaClp KC1-BaC12 NaC1-CaClz NaC1-CaC12 NaC1-CaC12 NaC1-KC1 NaC1-CsC1 NaC1-MgC12 NaC1-CaC12 NaC1-BaC12 NaC1-LiC1 NaC1-KCl NaC1-CsC1 NaCl-MgC12 NaC1-CaCh
- 0 .00g5b
0.0135b
- 0.016‘j - 0.0025k 0 .06Q6‘ 0.036 0.011 0. 02450 0.0515d -0.0195es‘ -0.0169 0 .00825i - 0 . 0335* 0.02550 0 . 04376d
- 0.0096‘ 0 .00 15C,h
0 . 04258 - 0. 00335d
Estimated from graph, ref 5j. This system does not obey Harned’s rule.
aij
= a
+ bS”ij
(3)
where a and b are parameters characteristic of a given electrolyte a t a given ionic strength, I . Soijis a function of the conventional ionic entropies, which have been computed from experimental data for a large number of ions.Ib These entropies can be estimated from the empirical relationshiplc,z
So = 3/2R In W
Soil= zlSo,+ z,so2 + z3S03
(5)
the suffixes 1, 2, and 3 refer to the ions i+, j + , j-, respectively, and
+ 37 - 2701z//r2
(4) where R is the gas constant, W the atomic weight, 1x1 the absolute value of the valency, and r the effective radius of the ion in solution (taken as 1.0 more than the crystal radius for anions and 2.0 A more for cations). While Soij is an empirically derived function, its form was suggested by the theory of Friedman3 and general thermodynamic considerations, the original idea being initiated by a paper of Wood, et al., on enthalpies of m i ~ i n g . The ~ simplest possible function was adopted, the weighted sum of the conventional entropies of the three ionic species in the mixture. The weights were chosen with the restriction that all possible pairs of ions contribute for a charge-asymmetric system whereas for a charge-symmetric system only cation-cation pair terms are nonzero. The The Journal of Physical Chemistry
ionic charges were considered the most appropriate quantities to use for the weights, and a suitable entropy function was found to be
= (3x1
2 1
- z3)(22 - 21)/4(Zg - Z 1) 2,
2 3
= 3(2i -
= 1/x22
x3)(&
- Zi)/Zi2(Zz - 23)
x being the charge on the ion, e.g., for the system NaC1-CaCl,, z1 = +1, x2 = + 2 and 23 = -1, 21 =
-l/z,
ZZ
=
l/4,
and Z3
=
2 , so that
(1) (a) R. A. Robinson and R. H. Stokes, “Electrolyte Solutions,” 2nd ed, Butterworth, London, 1959; (h) G. N. Lewis and M. Randall, revised by K. S. Pitser and L. Brewer, “Thermodynamics,” 2nd ed, McGraw, New York, N. Y., 1961, p 400; (c) See ref lb, p 523; (d) H. S.Harned and B. B. Owen, “The Physical Chernistxy of Electrolytic Solutions,” 3rd ed, Reinhold, New York, N. Y., 1963. (2) R. E. Powell and W. 1LI. Latimer, J . Chem. Phys., 19, 1139 (1951). (3) H. L. Friedman, ”Ionic Solution Theory,” Interscience, New York, N. Y., 1962, eq 18.36. (4) R. H. Wood, J. D. Patton, and Mi. Ghamkhar, J . Phys. Chem., 73, 346 (1969).
2227
NOTES 0.12-
0.10
-
-
0.08
0,Ob-
/ /
0.04-
/ /
0,021 0.00
- 0.02
1
I
-0.04
Figure 1 . Harned coefficients (Table 11)of five electrolytes us. the ionic entropies to give the H lines characteristic of the electrolyte at Z = 2, and a t 25'. The number pairs ij refer to the electrolytes i and j as per the cation coding (Table I ) ; e.g., 39 represents the system NaC1-BaC12.
S'NaCl(CaC1z)
=
-l/ZS"Na+
+
1/4SoCa2+
+
2S°Cl-
and for HC1-AlCls S'HCl(A1Cla)
since S O H +
=
=
3 S " ~ l -
+
1/9soAls+
0
Values of Soijhave been computed for five electrolytes (the chlorides of H, Li, Na, K, and Cs) to cover a wide range of chloride systems. The entropies are listed in Table I. For the system HC1-CeCl8, SoCelt was computed from eq 4. sij"values were then plotted against the experimental values of the Harned coefficients (Table 115) for the respective electrolyte. Figure 1 shows the results for I = 2. The effect of any added chloride on a particular electrolyte is neatly summarized by the characteristic line (the H line) of the electrolyte a t the given ionic strength. The deviations from the line indicate a deficiency in the simple entropy functions, although eq 1 and 2 might not apply and experimental errors could also contribute. The relationships should, however, prove useful since only the ionic entropies, which are usually readily available or can be estimated from eq 4, and d(1og yij) for a few "interfering" cations are required to construct the H line. The Harned coefficients for the electrolyte with any other chloride can then be predicted with reasonable accuracy from this line. At least, this is what the pre-
liminary results suggest and extension to other more complicated systems seems feasible.* An important advantage is that information on the single electrolyte is not required for the prediction of the Harned coefficients; frequently it is only the change of the activity coefficient with solution composition that is required rather than its absolute value. The point on the H line representing the limiting case of the single electrolyte solution might be expected to have the coordinates (S'i t,O). However, this is not the case and it is necessary to apply a correction to the standard-state entropy. This shift in the coordinate is interpreted as arising from solvent modification by i, the (5) (a) R. A. Robinson and C. K. Lim, Trans. Faraday SOC.,49, 1144 (1953); (b) R. M.Rush and R. A. Robinson, J . Tenn. Acad. Sei., 43, 22 (1968); (0) R. A. Robinson, J . Phys. Chem., 65, 662 (1961); (d) R. A. Robinson, J . Amer. Chem. SOC.,74, 6035 (1952); (e) Y. C. Wu, R. M. Rush, and G. Scatchard, J . Phys. Chem., 72, 4048 (1968) ; (f) R. F. Platford, ibid., 4053 (1968) ; (gj R. A. Robinson and V. E. Bower, J . Res. Nat. Bur. Stand., Sect A , 70, 313 (1966) ; (h) R . D. Lanier, J . Phyla. Chem., 69, 3992 (1965); (i) R. A. Robinson and V. E. Bower, J . Rea. Nat. Bur. Stand., Sect A , 69, 19 (1965): 0) R. A. Robinson and A. K. Covington, ibid., 72, 239 (1968); (kj R. A. Robinson and V. E. Bower, ibid., 69, 439 (1965); (1) E. W. Moore and J. W. Ross, J . Appl. Physiol., 20, 1332 (1965); (m) H. S. Harned and R. A. Robinson, "Multicomponent Electrolyte Solutions," Pergamon, London, 1968, p 66. (6) (a) C. M. Criss, R. P. Held, and E. Luksha, J . Phy3. Chem., 72, 2970 (1968); (b) F. Franks and D. S. Reid, ibid., 73, 3152 (1969). Extension to mixed solvents seems feasible too in view of the work of these authors. Volume 74, Number 10 May 1d3 1070
NOTES
2228
L-
872
Figure 2. The H lines of NaCl for different ionic strengths. The lines have been spaced out for clarity as indicated by the respective zero ordinate. The temperature is 25". The numbers represent the added cation chloride as coded in Table I.
water structure is altered and hence there is a change in the entropy of solution of all the j electrolytes. Below a certain ionic strength (say 0.5 &I) this effect would probably be small. A reasonably accurate prediction of Soliwould be useful when only one or two experimental points are available; hence I have chosen what seems an appropriate variable , the "structure-breaking ~,~ entropy," A p t , defined by Frank and E v a n ~ ,(Table 111), and a simple linear form to derive the equation Table I11 : Values of the "Structure-Breaking Entropy"'" ---c1-
7
asat
T,i--
H+
Na+
K-
cs
-1.1
0
$4
$12.0
flL7
So.. - 3". - ,-(Apt. 11
1-
1-
+
+ Aptl+)
$10.2
(6) where c, for the data used here, has a value of 0.2 and is negative for structure makers (Li+, H+, ?Sa+) and positive for structure breakers (K+, Cs+). Figure 2 illustrates the effect of changing ionic strength on the slope of the H lines of sodium chloride. The 0.7 ;1.I line should be of particular interest in the oceanographic field. The lines at I < 0.7 .!If have been derived from only one experimental system since, as far The Journal of Physical Chemistry
as I know, there are no other suitable data available. I n addition, the assumption that Soll= S"i + was made. If these lines represent the true situation then it appears that at very lox- ionic strengths the Harned coefficients become either very small ( e . g . , for the chlorides of Jig, Li, and H) or very large (e.g., for the chlorides of Ca, Ba, I 0.35: they gave Vm values of 6.7 (Ar) and 8.0 (CO) relative to closely 7.8 om3 (STP)/g for nitrogen, assuming the three adsorbates have the same molecular density in the multilayer region. However, argon may not be able to pack so closely as do the other two adsorbates, since the specific surface areas obtained were (mz/gl : 33.8 (Nd, 25.6 (Ar), 36.1 (CO). These were calculated
Figure 1. Adsorption isotherms for nitrogen, argon, and carbon monoxide on ice a t 77°K. Nitrogen: 0 , adsorption; A, desorption; M, adsorption following 90.1 OK nitrogen isotherm; 8, adsorption following 90.1 OK argon isotherm; 8, adsorption following 90.1 "K carbon monoxide isotherm. Argon: a, adsorption; C), desorption. Carbon monoxide: 0, adsorption; A, desorption.
using the B E T u, values and equivalent spherical cross sections (based on the adsorbate liquid densities) of 16.2 (N2),13.8 (Ar), and 16.0 (CO) i2. Of more direct interest to the purpose of the investigation is the relative behavior in the submonolayer region. Figure 2 shows the variations with surface coverage of the isosteric heats and partial molar entropies of adsorption, as determined from pairs of isotherms at 77.3 and 90.1"1
