886
NOTES
Vol. 62
for both salts, while a maximum occurs at 49.7 wt. cis- and trans-sulfates,6*l1but in the present case it 70for the cis salt and 52.7 wt. Yofor trans salt. holds in spite of the existence of minima and maxSimilar minima and maxima in ethanol-water ima as the solvent changes. Thus, it appears that mixtures were observed by Fischero for the solubil- addition of alcohol to the solvent affects the two ity not only of sodium, potassium and barium pic- salts proportionately, and again indicates that the rates, but also for sodium dinitro-, chloronitro-, and presence of minima and maxima are concerned with p-nitro-phenolates and barium dinitrosalicylate. effects on the picrate ion and not on the cations. Fischer’s data for sodium picrate (a monohydrate), recalculated to a weight yo and molar basis, are re- T H E ALKALINE EARTH COMPLEXES OF produced in Fig. 1; the curve for potassium picT H E ADENOSINE PHOSPHATES rate, an anhydrous substance, has a very similar shape. Sodium picrate is approximately 100 BY GEROLDSCHWARZENBACH AND ARTHURE. MARTELL times, and potassium picrate 10 times, as soluble as Eidgenossische Technisch: Hochkhzsle, Zurich, Switzerland and Clark Unzversity, Worcester, Mass. the cis-cobaltamrnine picrate. We have confirmed, Received April 7, 1968 by spectrophotometric measurements, Fischer’s observation that there is a second minimum for soI n his paper on the stabilities of the complexes dium picrate between 92 and 100% ethanol. Similar of calcium and magnesium with AMP, ADP and measurements have shown that this is not the case ATP, Nanningal compares his results with the for the dinitrotetrammine cobalt picrates; for these constants obtained by Smith and Alberty2 and salts, the solubility decreases steadily beyond the with the data obtained by US.^ Our measuremaximum. ments have been carried out at 20” in a solution I n seeking an explanation of these results one of ionic strength y = 0.1 obtained by adding KCI. notes that association into ion pairs is not likely t o Smith and Alberty on the other hand had 25” and be appreciable in the low concentration range and I(. = 0.2 ebtained with a tetrapropylammonium relatively high dielectric constant range concerned salt, whereas Nanninga used mixtures of varying in the present experiments. Solvolysis t o picric acid ionic strength (NaC1 and tetraethylammonium may be shown t o be insignificant, by calculations bromide). I n order t o make a comparison of the from the known dissociation constants.’O The sul- data possible, Nanninga corrected his and our fates of the two cations in the same solvent pair’l results for Na- and K-binding, assuming that both show no unexpected behavior, and one looks for a alkali ions would, according to M e l ~ h i o r ,form ~ possible complex involving the picrate ion. Kor- complexes of exactly the same stability. In tum12has observed shifts in absorption maxima of applying these corrections, Nanninga increased picrates in the intermediate range in alcohol-water twofold the association constants obtained by us mixtures, which he attributes to changes in solva- for ATP and found that the “corrected” data are tion. Application of Walden’s rule to ethanol- much larger than those obtained by other inwater mixtures shows a maximum in Aov for many vestigators. This was explained by Nanninga by electrolytes between 14 and 22 wt. % ethan01,’~-‘~ making the statement that Martell and Schwarzbut this is attributed t o the change in the mobility enbach had implied the equality of K& and K& of the cation14; transference measurements16con- in their derivation and that their equations were firm this. There is, however, some increase in the therefore incorrect. product lo7 for picrate ion in lithium picrate solution I n Table I, the data of the three groups of inon addition of the first 12 wt. yo ethanol, followed vestigators are collected without making any corby a lesser increase and a very slight maximum at rections. The comparison shows very good agreeabout 52 wt. %.I6 Comparison with the solubility ment of the groups Smith-Alberty and Martelldata shows that the first rise corresponds to a de- Schwarzenbach. It seems that the potassium crease in solubility, yet the maximum appears near salt used by us has little influence on the apparent the maximum solubility. stabilities, a fact which is in harmony It has been noted above that for the cis- and complex the experience obtained with other complextrans-dinitrotetrammine cobalt picrates the trans with ing agents. Thus it was found6 that potassium salt is more soluble than the cis. It is of interest only very weakly bound by the anion of uramildithat this holds throughout the solvent range and is is acetic acid, which forms a sodium complex with a such that a plot of the logarithm of the solubility of stability constant of 108. Van Wazer and Campathe cis salt against the logarithm of the solubility of nellas also found a considerable difference between the trans salt is a straight line. Relationships of this Kf and Naf when complexed with the polyphossort have been noted before16 and hold also for the phates. The somewhat smaller stability constants (9) W. M.Fischer, 2. physak. Chem., 92, 581 (1918). obtained by Smith and Alberty readily are ex(10) P. S. Danner, J . A m . Chem. Sac., 44,2832 (1922); I. M. Koltplained by the higher ionic strength used by these hoff, J. J. Lingane and W. D. Larson, &id., 60,2512 (1938). authors. Binding of Kf would on the other hand 62,358(1958). (11) H.L. Clever and E”. H. Verhoek, T H I s JOURNAL, (12) G. Kortum, 2. physik. Chem., B88, 1 (1937). (13) L. C. Connell, R. T. Hamilton and J. A. V. Butler, Proc. Roy. Soc. (London), 6147, 418 (1934). (14) A. F. H. Ward, J . Chem. Soc. 522( 1939). (15) G. Kortum and A. Weller, Z. Naturforsch., 6 6 , 451, 590 (1950); G. Kortum and H. Wilski, 2. physik. Chem. (FrankfuTt), 2, 256 (1954). (16) J. N. BrZnsted, “Chemistry at the Centenary Meeting of the British Association for the Advancement of Science,” Cambridge, W. Heffer and Sons, Ltd., 1932, p. 39.
(1) L.B. Nanninga, THISJOURNAL, 61, 1144 (1957). (2) R. M. Smith and R. A. Alberty, J . A m . Chem. SOC.,‘78, 2376
(1956). (3) A. E.Martell and G . Schwarzenbaoh, Helv. Chim. Acta, 89,653 (1956). (4) N. Melchior, J . B i d . Chsm., 208,615 11954). (5) G. Schwarzenbach, E. Kampitsch and R . Steiner, Relv. Chim. Acta, 29,364 (1946). (6) I. R. Van Water and D. A. Campanella, J . Am. Chem. SOC.,72, 655 (1950).
88i
XOTES
July, 1958
TABLE I
LOQ OF EQUILIBRIUM CONSTANTS OF ADENOSINE MONO-,DI- AND TRIPHOSPHATE (Represented by AMP2-, ADPa- and ATPd-, respectively) Reaction
Martell-Schwaraenbacha p = 0, 1 (KCI); 20''
+ AMP2- Ca-AMP + ADPs- & Ca-ADPICa2+ + ATP4- i=? Ca-ATP2(Cas+ + HATPa- e Ca-HATPH++ Ca-ATP2- e Ca-HATPiH+ + ATP4- e HATP3Mg2+ + AMP2- e Mg-AMP Mg2+ + ADPs- e Mg-ADPMg2+ + HADP2- & Mg-HADP {Hf + ADPa- F? HADP2Mg2+ + ATP'- F? Mg-ATP2Mgz+ + HATP3Mg-HATP'H+ + Mg-ATPZ- e Mg-HATPH + + ATP4- F? HATPa-
1.41 f 0.03 2.78 f 0.03 3 . 6 0 & 0.03 1.8 f O . l 4.7 i o . l 6.50 f 0.01 1.69 f 0.02 3.11 f 0.05 4.7 f0.2 6.35 i 0.03 4.00 f 0.04 2.0 * O . l 4.5 f0.1 6.5 1 0 . 0 1
(Caz+
Ca2+
make the constants measured in KCl smaller than those measured in tetrapropylammonium chloride. If it is assumed that the differences of the numbers found by Martell-Schwarzenbach and SmithAlberty are due mainly t o differences in ionic strength, it also becomes clear that the differences increase from AMP t o ADP and ATP because of the increase of charge of the anions. The large difference found for the stability of Mg-ATP seems t o be the only unexplained result. In our paper,3 the individual constants K& and KgH2(reactions H f Z F? HZ and H + MZ F? MHZ, Z denoting the anion of ADP and ATP) were given to show that these constants are different and that we have by no means implied in our derivation that they must be equal. The criticism of Nanninga is due t o a misunderstanding because he did not take into consideration the fact that the metal ion concentration which we employed was in considerable excess over that of the organic complexing agent. The experimental method which we used was described in detail as early as 1950,' and has been used with success also in other laboratories.8~9 The data for log K:z and log KgEzare the pKvalues for the proton donators HZ and MHZ. The large difference in comparison t o the data given by Smith-Alberty are due here to differlences in the definitions of these pK's. Our data are "apparent" dissociation constants, [H+J being introduced as a concentration whereas SmithAlberty used "mixed" constants in which [H+] is introduced as an activity. The following equations, in which C and a denote concentrations and activities respectively, make this difference clear
+
pk"
+
CB*W
= -log---
CAoid
x
ap
( 7 ) G . Schwarzenbach, Helu. Chim. Acta, 33, 947 (1950). (8) H. B. Jonassen and L. Westerman, J. A m . Cham. Soc., 19, 4275 41957). (9) 8, M, h m b e r t and J. I. Wettew, $aid,,79,5606 (1057).
Smith-Albertyz
p
= 0, 2 (n-[Ca%r]4NC, Cl-); 25
1.43 f 0.03 2.81 i 0.05 3.30 f 0.06 1.61 i 0.1 (5.3) (6.95) 1.69 f 0.02 3.00 iz 0.04 (5.1) (6.68) 3.47 f 0.03 1.49fO 1 (5.0) (6.95)
Nanninga' u = 0, 1; 23'
3.14
2.0 3.03
3 61
MICROWAVE ABSORPTION OF MIXTURES OF DIPOLAR LIQUIDS I N SOLUTION I N NON-POLAR SOLVENTS' BY P. K. KADABA Department of EEectrdcal Engineering, Newark College of Eneineerine, Newark, New Jersey Received April Ii, 1968
The investigations of Schallamach2 on the dielectric relaxation of binary mixtures of nonassociated polar liquids a t a frequency of 9 Mc./ see. in the temperature range -150 to -50" indicated a single absorption peak with the possibility of a single relaxation mechanism. By way of explanation, it was suggested that dielectric relaxation involved the disturbance of relatively large regions in the liquid. The use of a comparatively low frequency of 9 Mc./sec. required the use of liquids of a somewhat complex nature. It was the purpose of this investigation to extend the measurements to much simpler liquids by the use of a suitable frequency in the microwave region. With a view t o estimate the size of the regions involved in the relaxation process ternary mixtures containing two non-associated polar liquids in a suitable non-polar solvents were investigated. Based on Schallamach's observation it was reasonable t o expect that at low concentrations two peaks in the loss factor (e") us. temperature curves should be observed as the separation of the dipolar molecules of the two types is large. As the concentration is progressively increased the peaks might move closer and a t sufficiently high concentrations when the molecules can interfere with each other one peak only should be found. The critical concentration corresponding to the single peak would then give an idea of the size of the region involved in the dielectric relaxation. Mixtures of nitriles, aliphatic nitro compounds and nitrobenzene were chosen because of their fairly good solubility. Also, their large dipole (1) The principal work oonneated with the investigation was done at the Department of Eleatrical Engineering, University of Kentucky. Lexington, Ky. (2) A. Schallamach, Trans. Faraday Soc., 4aA, 180 (1946).
