adding to this 0.27 g of 2,6-pyridinedicarboxaldehydedissolved in 25 ml of water. The mixture was well shaken, kept for 30 min in a water bath (about SOo). and then stored overnight in the refrigerator. The yellow precipitate that formed was filtered, washed with water, and recrystallized from methanol. The yield was 78%. The Schiff base was stable for a t least one year in a dark brown bottle. A 0.4% solution of the reagent in chloroform was used for the tests. Saturated aqueous thiourea and 20% aqueous tartrate solutions were used for the U(V1)test. The buffer solution used for the Sb(II1) and Bi(II1) tests was prepared by mixing three parts of 9% aqueous ammonium acetate solution and one part of 6h‘ HCl (pH 0.235). Procedures. To Test for U(V1).One drop of thiourea solution and one drop of glacial acetic acid were added to one drop of the test solution. The mixture was shaken, and three drops of the reagent were added, followed by 1 ml of water. The red color produced was extracted into the chloroform layer by vigorous shaking. Limit of identification: 0.05 pg U(VI), limit of dilution: 1: 1,000,000. To Test for Sb(II1). Four drops of the buffer solution and four drops of the reagent were added to one drop of the test solution. The red color produced was extracted in the chloroform layer. When Bi(II1) was present, a drop of saturated thiourea was used to block it. Limit of identification: 0.05 pg Sb(II1); limit of dilution: 1: 1,000,000. To Test for Si(111).Four drops of the buffer solution and four drops of the reagent were added t o one drop of the test solution. The red color produced was extracted in chloroform. When Sb(II1) was present, its interference was avoided by first treating one drop of the test solution with two drops of 15% hydrogen peroxide. The mixture was kept in a water bath for one minute. The buffer and reagent were then added as above to produce a red color which was extracted in the chloroform layer. Limit of identification: 0.5 pg Bi(II1); limit of dilution: 1:100,000.
RESULTS AND DISCUSSION U(V1). The addition of a solution of (I) in chloroform to an aqueous solution of U(V1) salt results in the formation of an intense red precipitate. In the presence of glacial acetic acid, this precipitate could be extracted into chloroform. Aqueous solutions of the following ions were tested in the presence of glacial acetic acid: Li, Na, K, Rb, Cs, Be, Mg, Ca, Sr, Zr(IV), Cr(lI1). M O W ) , W(VI), Mn(II), Fe(III), Co(II), Ni(II), Pd(IV), Pt(IV), Cu(II), Ag(I), Au(III), Zn, Rh(III), Cd(II), Hg(1, 11), Al, Tl(I), Sn(II,IV), Pb(II), As(III,V), Sb(III,V), Bi(III), La(III), Ce(III,IV), Th(IV), and V(I1). Of all the above-mentioned ions only U(V1) and Bi(II1) gave a red color, while with the others only the yel-
low color of the reagent was observed in the chloroform layer. To remove the interference of Bi(II1) with U(VI), saturated thiourea solution was used, which also improved the limits of the identification of U(V1) in the presence of other cations. Sb(II1) also blocked the color formation; its interference was removed completely by the addition of tartrate. The test for U(V1) was satisfactory in the presence of the following anions: F, C1, Br, I, Nos, SO4, S, s&3,CN, COa, Po4, citrate, tartrate, malonate EDTA, CNS, Cr04, Mn04, IO3, C103, Br03, and C104; but it failed in the presence of oxalate. This test for U(V1) made it possible to identify 0.1 g U(V1) in 500 g Be, Zr(IV), Mn(II), Ag(I), Cd, Hg(1, 11), Sn(I1, IV), and Bi(II1). Moreover, 0.5 g of U(V1) could be detected in 500 g of V(V) and Ce(1V). When Sb(II1) was present, two drops of 20% sodium, potassium tartrate was substituted for the thiourea of solution. I t was thus possible to detect 0.5 g U(V1) in 500 g Sb(II1). Sb(II1). The addition of a solution of Schiff base(1) in chloroform to an aqueous solution of Sb(II1) buffered by ammonium acetate produced an intense red color at pH 0.235. Aqueous solutions of the cations mentioned in the test for U(V1) were tested with the reagent in buffer; none produced a red color in the chloroform layer besides Sb(II1) and Bi(II1). However, the color due to Bi(II1) disappeared when a drop of saturated aqueous thiourea solution was added. The test was satisfactory in the presence of the following anions: F, Cl, Br, I, Nos, SO4, s&3,CN, C03, Po4, citrate, tartrate, malonate EDTA, CNS, oxalate, IO3, BrOa, C103, C104, Mn04, S 2 0 3 , and CrzO:. Sulfide reduced the limit of the sensitivity of the test to 50 g Sb(II1) in 500 g sulfide. It has been possible by this test to identify 0.2 g Sb(II1) in 500 g U(V1) and 0.5 g Sb(II1) in 500 g Ce(IV), V(V), Au(III), Hg(I), Pb(II), and Bi(II1). The remaining cations mentioned under U(V1) did not affect the limit. Bi(II1). The test for Bi(II1) works the same as for Sb(II1) except for the fact that thiourea blocks the Bi(II1) reaction. Thiourea produces a yellow color with Bi(III), which is the same color as that of the reagent itself. RECEIVEDfor review January 16, 1975. Accepted April 10, 1975.
Study of Liquid-Powder Interfaces by Means of Solvent Strength Parameter Measurements Claude H. Eon Laboratoire de Chimie-analytique-physique, Ecole Polytechnique, 17, rue Descartes Paris
Since the Forties, the need for optimizing adsorption chromatography has given rise to many attempts to classify solvents (and adsorbents) leading to scales of “polarity” often referred to as eluotropic series. Most of them, arbitrarily derived by simply ranking parameters like dielectric constants, dipole moments, heat of wetting, solubility in water . . . ( I ) can be considered as only very rough guidelines for they have very little theoretical foundations. A far more realistic approach, described by Snyder (2), consists in measuring the “strength” of the solvent from
50,France
chromatographic behavior; a proper correction for the solute effect leads to a solvent strength parameter t o which exhibits the remarkable property of being only a function of the pair solvent/adsorbent. Despite the fact that these series have quickly become popular, the rationale of this classification still escapes many analysts for t o is often considered to be an empirical parameter. In fact, t o is a rather fundamental parameter which can be related to the adsorption energy of the solvent ( 2 ) . It will be shown here that this view is consistent with the thermodynamics of the in-
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terfacial monolayer and that co is an unambiguous measure of the interfacial tension solvent-adsorbent relative to a reference system. Hence, we strongly suggest the use of chromatography as a way of studying interfaces between liquids and powdered solids.
RELATIONSHIP BETWEEN THE SOLVENT STRENGTH AND THE SURFACE TENSION OF THE SOLID/LIQUID INTERFACE As we are here mainly concerned with an estimation of surface energies, it is first necessary to reflect upon the theoretical foundations of the solvent strength parameter. As in the original theory ( 2 ) ,we consider the action of a change of the solvent nature upon the retention volume of a solute. Our starting point will be the classical relationship among the retention volume (VR), the dead volume of the column (Vo),and the capacity factor (k’)( 2 ) : where the subscript 1 stands for the solvent and 2 stands for the solute. The capacity factor, ratio of the number of moles of solutes in the adsorbed and mobile phase, can be written in its classical form:
where wBb is the standard chemical potential, in such case defined for one mole of pure solute in the liquid state and 78 is the bulk activity coefficient. By analogy with the bulk relationship, the surface chemical potential can be written ( 4 ) :
where the surface standard chemical potential is related to the bulk standard chemical potential through ( 3 , 4 ) :
In this Equation, the subscript 3 stands for the adsorbent; thus U2/3 stands for the solvent/adsorbent interfacial tension and UM/3 for the solution/adsorbent one. Combining Equations 6 , 7 , and 8 and taking into account that the solute is infinitely diluted, it is clear that:
Compared to the classical expression log K O( 2 ) : log K O= log V,
+ a (So - Azo + Aeas to)
(10)
Equations 5 and 9 suggest the following equalities: where n2” is the “adsorbed solute concentration” expressed in moles per gram of adsorbent, Cz is the solute molar concentration in the mobile phase, and ma is the mass of adsorbent. Here it is, however, more appropriate to express the concentrations in mole fractions; a rather simple derivation ( 3 ) leads to: (3)
where as is the specific area, VI” and Ai0 are, respectively, the solvent molar volume and molar area while X z S and Xzb are the solute mole fractions, respectively, in the surface and in the bulk regions. Thus, the product as V1O/A1’ is the volume of a monolayer film of solvent adsorbed by 1 gram of adsorbent; this term is referred to as Va in Snyder’s terminology. Notice that the factor Xzb in Equation 3 is almost always small compared to XzS so that this correction [for thermodynamic consistency with the Gibbs theory ( 4 ) ] ,can be omitted in common practice. Thus combining Equations 2 and 3 leads to:
kz’ = V,
@(I)
ma VO,
(4)
where @ ( I ) , the partition ratio, is in solvent 1. At this stage it must be pointed out that a change in solvent nature only induces a change in the product Va@(l).For further reference, note that this product is equivalent to the factor K Odefined in ( 2 ) .Thus: As V, is roughly constant for any given system, the problem has become finding the variation of CP induced by a change in the nature of the carrier; e.g., for a shift from a reference solvent 1’ to a solvent 1. The way to sort things out is to set the chemical potentials of the solute in the bulk and at the interface equal. The bulk chemical potential of the solute in the solvent 1’ is given by the well known equation:
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aso= log (x&)/Xpl,)) ato =
(“1’/3
(11)
- ‘J1/3)
2.3 RT
Aeas = log Y h Y%’) Y%)
(13)
YZb(1,)
The elegance of Snyder’s concept is immediately recognized, for each of the three preceding parameters are related to completely dissociated factors making up the final retention of the solute 1. First, the selectivity parameter So times the adsorbent activity a. is only a function of the solute elution with the reference solvent. Second, the solvent strength parameter to times the adsorbent activity is related to the change of interfacial tension solvent-adsorbent from the reference solvent to the other solvent; it is usually the main effect governing the retention volume change. Third, the so-called secondary effect factor Aeas is related to the activity coefficient ratio; it has been found to be usually small compared to the main effect. As far as this article is concerned, Equation 12 is of first importance for it confirms that the solvent strength is a measure of the relative interfacial tension. Because the unit of surface area used in Equation 2 is 8.5 A2, it turns out that:
-
ai~/3 a1/3 = &r
= 111.4 to
(14)
with the u’s expressed in dynes per centimeter and the temperature of 3OOC.
RESULTS AND DISCUSSION The structural identity of Equations 9 and 10 and, consequently, the merit of Equation 14 is backed by some experimental evidences found from our earlier results ( 5 ) . Figure 1 represents the logarithm of the retention volumes resulting from the adsorption of some solutes at the organic solvent/water interface as a function of the interfacial tension of the two phases. Here heptane is taken as a reference solvent. Usually, however, chemists are more concerned with solvent/porous adsorbents systems (for which
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t
nixEo
(+/ncm”) e.
L
/
(en-, cm-’1
0
c
0
25
53
75
1CO
Figure 2. Variation of the solvent strength parameter on silica with the adhesion tension. Numbers as in Table I
I
od. Their results are given in terms of an adhesion tension 7.4 which is defined as:
\
Figure 1. Variation of the logarithm of the retention volume as
a
function of the interfacial tension between the two liquid phases
+
System A = heptane-water, B = 1.1 heptane toluene-water, C = cytolueneclooctane-water, D = toluene-water, E = 1.1 nitrobenzene water, F = nitrobenzene-water, G = olive oil-water, H = amylacetatewater. Solutes: (1) = amyl acetate. (2) = butyl acetate
+
Table I. D a t a for Silica e o silica SO.
Solute
(2 )
TA(dynrs crn-l) (6)
Carbon tetrachloride 0.11 39.5 0.20 43.2 Carbon disulfide 0.25 51.2 Benzene Toluene 0.23 53.4 Chloroform 0.26 58.7 Aniline 0.48” 73.8 Ethyl acetate 0.45a 76.1 Amyl alcohol 0.47“ 77.5 Estimated from e o Alumina with the help of Equations 6-8 in
1 2 3 4 5 6 7 8
reference 2.
the interfacial tension cannot conveniently be measured); in such cases, we can expect Equation 14 to lead to meaningful estimation of relative interfacial tensions as demonstrated by the analysis of Bartell and Scheffer data (6). These authors have studied the free energy of wetting of powders by means of the “pressure of displacement” meth-
7A
=
g3
- g1/3
(15)
where u3 is the surface tension of the adsorbent ( 7 ) . In the absence of water a = 1 (2); thus, Equation 14 pre. dicts that there must exist a linear relationship of unity slope between A 0 and T ~ .Figure 2, established for silica from the data in Table I, substantiates the exactness of Equation 14. Notice that the slope 1.07 is in fair agreement with the theoretical one. This introduces the possibility of chromatography as a way of studying liquid solid interfaces free energy. I t is justified because it is impossible to make any direct experimental determination of the interfacial tension of such systems because of the lack of mobility of the interface. ACKNOWLEDGMENT I am indebted to L. Snyder for his comments and suggestions. LITERATURE C I T E D (1) W. Trappe, Biochem. b,305, 160 (1940). (2) L. Snyder, “Principles of Adsorption Chromatography”, MarceCDekker. New York, N.Y., 1968. Chapters 6 and 8. (3) C. Eon and G. Guiochon, J. Coiioidlnterface Sci., 45, 521 (1973). (4) J. Eriksson, Ark. Kemi., 26, 46 (1966). (5)C. Eon, B. Novosel, and G. Guiochon, J. Chromatogr., 83, 77 (1973). (6) F. E. Bartell and H. J. Osterhof, J. Phys. Chem.. 37, 543 (1933). (7) R. Johnson and R. H. Dettre, “Surface and Colloid Science”, Volume 11, E. Matijevic, Ed., Wiley-lnterscience, New York, N.Y., 1969, p 85.
RECEIVEDfor review February 27, 1975. Accepted June 3, 1975.
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