NOTES
1544
data in Table I, are shown in Table TI, along with corresponding data obtained previously for the reaction in several other solvents. TABLE I1 KINETICDATAFOR THE DECARBOXYLATION OF OXASILIC ACIDIN THE MOLTEN STATEAND IN SEVERAL POLAR LIQVIDS Solvent
Bis-(2-cliloroethyl) ether n-Butyl ether n-Amyl ether @-Chlorophenetoles Phenetole8 Anisolea Dibenzyl ether N,N-Dimethylaniline3 Melt n-Hexyl ether o-Et hylaniline 1 o-Toluidine’ Aniline‘
AH* kial.) Illole
21.4 25.1 28.3 31.3 32.6 35.6 36.8 37.6 40.1 40.1 45.5 47.8 49.8
AS*,
e.u./ mole
-22.4 -13.9 6.6
-
+ 0.7
3. 4.0
$11.1 f14.2 +15 3 f21.4
+22.1 +34.3 $39 9 f46.3
AF*IL.Do, kcal./ mole
30.9 31.0 31.1 31.0 30.9
31.0 30.8 31.1 31.0 30.8 31.0 30.9 30.2
Discussion of Results The parameters of the Eyring equation for the decarboxylation of oxanilic acid in the molten state as well as in twelve organic solvents are listed in Table 11. The solvents are arranged in the order of increasing enthalpy of activation, which is likewise that of increasing entropy of activation. An enthalpy-entropy plot of the data given in Table I1 yields a straight line, indicating that the reaction takes place in each case by the same mechanism.’ This mechanism is very likely similar to that of the malonic acid decomposition, an activated complex being formed uia the electrophilic polarized carbonyl carbon atom of the acid and a n unshared pair of electrons on a nucleophilic atom of s01vent.~ If this is the case, it would be expected that the A H * of the reaction would decrease as the nucleophilicity of the solvent increases.8 It has been noted previously3 that the A H * of the oxanilic acid reaction in aromatic amines is higher than would be expected on the basis of the proposed mechanism (see Table 11). Very possibly, in the reaction in ethers, unionized oxanilic acid takes part in the rate-determining step, whereas, in more strongly basic solvents, ionization of the oxanilic acid takes place and the anion is the entity involved in the reaction. The value of AS* is a measure of the relative complexity of the activated c o m p l e ~ . ~I n general, an increase in size or steric hindrance of either solute or solvent will bring about a decrease in AS*, and vice versa. It will be seen from the data in Table I1 that the reaction in the smallest ether is accompanied by the lowest value of AH*, and also the largest negative value of AS*. As the size of the solvent molecule increases it appears that the size of the complex decreases. Two factors may contribute to increase the size of the activated complex-polymerization of the solvent and polymerization of the solute. Monobasic acids exist (7) J . E Lcfflrr, J. 0 ~Chem., . 80, 1202 (1965). (8) K. I. I aidler “Chemical Icinetios,” McGraw-Nil1 Book Co., Ino., New York, N. Y.,1950, p. 138. (9) S. Glasstone, J. K. Laidler, and H. Eyring, “The Theory of Rate Prooemes.” MpGraw-Hill Book Go., Ino., New York, N. Y., 1941, p. 22.
Vol. 66
largely as dimers. Ethers associate to a certain extent through dipole-dipole interaction. The extent of association would tend to decrease as the number of carbon atoms increases. These two sets of conditions would help to account satisfactorily for the data involving the ethers shown in Table II. I n the case of the amines we find a large A H * accompanied by a large positive value of AS* (see Table 11). The aromatic amines have fewer carbon atoms than the ethers and show little association. Furthermore, if the anion of oxanilic acid is the entity reacting it would not dimerize. These two factors would contribute to decrease the size of the activated complex, giving rise to a large positive value of AS*. The slope of the enthalpy-entropy plot of the data in Table I1 turns out to be 423. This corresponds to the isokinetic temperature, that is, the temperature a t which the rate of reaction is the same in all solvents.’ This is also the temperature at which A F * = AHo*, where AHo* is the enthalpy change corresponding to zero change of entropy. The experimental enthalpy change corresponding to AS* = 0 is 31.0 kcal./mole. The isokinetic temperature, 423’K., is 150°, which is very nearly the melting point of oxanilic acid. It will be seen for the decarboxylation in Table I1 that the AF1500* of molten oxanilic acid, as well as for the reaction in all the solvents listed (aniline excepted), turns out to be 31.0 kcal./mole within plus or minus one or two tenths, The discrepancy in the case of aniline is probably due to some error in the data. Acknowledgment.-The support of this research by the National Science Foundation, Washington, D. C., is gratefully acknowledged.
THE EFFECT OF ELECTROLYTES QK THE SOLUTION CHROMOTROPISM OF BIS-(meso$2,3-DIL4MIXOBUTAT\’E)-NICKEL (11) IONS BY D. L. LEUSSING, J. HARRIS,^
AND
P. WOOD
National Bureau of Standards, Washington, D. C. Received January 86, 1.961
The bis-C-methylethylenediaminecomplexes of nickel(I1) have been found to exhibit a yellow diamagnetic form in solution in addition to the normal blue paramagnetic form.2 Increasing substitution favors diamagnetism so that the C,C’-TetraMeen3 complex is present completely in the diamagnetic form but with mego- and racemic-bn both forms are present in appreciable amounts at room tempbrature. Ahmed and Wilkins4 report that the addition of sodium perchlorate to nitrate solutions of ;“rTi(r~c-bn)~++ causes the proportion of the yellow form to increase. Similar behavior has been reported by Jorgensens with nickel(I1) and trien (1) NBS Summer Student, 1961. (2) F. Basolo, Y. T. Chen, and R. K. Murmann, J . Am. Chem. Soc., 76, 956 (1964). (3) en = ethylenediamine, bn = m-2,3-diaminobutane, pn = propylenediamine, C,C’-TetraMeen = C,C’-tetramethylethylenediamine, trien = triethylenetetramine. (4) A. K. Shamsuddin Ahmed and R. G. Wilkins, J . Chem. Sor., 2901 (1960). (5) C. Klixbull Jorgensen, Acta Chem. Scand., 11, 399 (1957).
August, 1962
NOTES
and by Sone and Kato6for hot solutions of Nien2++ and N p n e+ +. I n order to learn more about this interesting property me have examined the spectra and determined the formation constants of the nickel(I1) complexes with meso-bn a t 25O in various salt solutions. Experimental meso-2,3-Diaminobutane dihydrochloride was prepared Pccording to the method of Dickey, Fickett, and Lucas.' Analysis by potentiometric titration with standard silver nitrate solution indicated a purity of 99.2%. The formation const,ants were calculated from the_r e s u h of the usual pH-titration experiments which yielded 12 values and the corresponding free ligand concentrations at 25.0'. The primary medium in all experiments was a solution 0.080 &f in bn.2HC1 and 0.013 iM in NiClp, which also contained added e1ectrolyt)e. A standardized 1.OO M solution of sodium hydroxide was the tit-rent. I n those experiments where the concentration of added electrolyte was 1.00 M or greater the titrant also was made up to contain the same concentration of electrolyte as the test solution. Equilibrium was attained within the time of mixing. With the perchlorate solutions, a saturated sodium chloride salt bridge was used to connect the solution being titrated with a.n external saturated calomel electrode. In the perchlorate medium it was found impossible to obtain useful results regarding the stability of the tris complex,,since the perchlorat,e of this ion is only slightly soluble. Spectrophotometric measurements were made, using a Gary Model 14 spectrophotometer with a cell compartment thermostatecl a t 25.0'. In the solutions where the added electrol te: way 0.10 M sodium nitrate the spectra of Nibn++, Sibnz+q and Nibna+-"were obtained on solutions in which ?a, the average number of bn molecules bound per nickel(I1) ion, was 0.35, 2.04, and 3.05. For the first case a solution containing 0.10 M potaseium nitrate and a concentration of nickel chloride equal. to that of the uncomplexed nickel(I1) in the sample solution was used as the reference solution to subtract out the contribution of the free nickel ions. Similar measurements wcre made using en to obtain the spectra of the complexes with this latter ligand. The extinction coefficients calculated for the maxima corresponding to those of the paramagnetic species are given in Ta,ble I.
1545
light beam passes were mwked BO that the light beam did not strike the meniscus of the Bolution. The solution then was titrated with base dispensed with z1 microburet. The spectrum in the region of 442 mp WBB scanned for each increment of base. The extinction coefficients at 442 mfi are given in Table 11. Because finite increments of base were added, the possibility exists that measurements were made either slightly bpfore or slightly after the point corresponding exactly to the formation of n'ibnz++. For this reason the extinction coefficients for these experiments tend to be low but because the titration points were sufficiently close the maximum relative error in the extinction coefficientsis 10 to 15%. This error is acrep table for the present comparison purposes and is indicated in Table I1 for the experiments so affected.
Results and Discussion I n Table I the low extinction coefficients of Xibnz++ a t its two long wave length maxima are apparent. The low results are due to the presence of an appreciable proportion of the diamagnetic isomer which does not absorb in this region. The extinction coefficients of only the spin-free isomer, Esf, were estimated by assuming that esfg35 is equal to the average of eg65 and €915 for the mono and tris bn complexes and, similarly, €670 for the bis is equal to the average of the extinction coefficients for the second peaks of the lower and higher complexes. Applying this procedure to the en system gives a calculated value ol 6.4 for Cg27 of &'ien2++ compared to 6.5 observed and a calculated value of 5.5 for E568 compared to 5.4 observed, so this method appears to give reasonably good values. The ratio €sf9650be/~sf966calo for Nibnz++ is 0.73 and the ratio Esf570oba/€sf670calc is 0.75. Thus, the results a t both wave lengths are in good agreement and indicate about 74% of the his complex to be spin-free. Using this fraction the extinction coefficient of the spin-paired form of Sibnz++ a t 442 mp, Esp442, was calculated from the result given in Table II for 0. LO M NaN03. In this calculation the valine of E442sf was assumed to be 1.4, which is TABLE I the value for Nicnz++ a t this wave length. This THE POSSTIONS AND EXTINCTIONCOEFFICIENTS CORREI- choice is justified because of the close similarity of the spectra of the en and bn nickel(I1) complexes SPONDIKG TO THE ABSORPTION MAXIMAOF THE PARAMAG(with the exception of the 442 nip peak) and, owing NETIC NKCKEL(II) COMPLEXES 0.10 M KNOa, 25' to its low value, the results are not sensitive to X,mp e A, m@ s X,mp Nien + + 970 5.6 616 4.3 370 7 . 6 relatively large variations in the assumed value of Nibn ++ 965 5.7 623 4.2 372 7 . 3 Esf442. The calculated value of Esp442 is 88. This Sien2++ 927 6.5 568 5.4 354 11.4 value is somewhat kigher than the value 66 for Kibn2++ 935 5.1 570 4.5 360 12.4 Ni(C,C'-TetraMeen)2++ a t its 434 mp maximum. Niena++ 884 7.1 543 345 12.0 Both of these values lie well within the range re6.8 Xibn8++ 915 557 8.3 7.9 353 14.0 ported by Holm9 for various diamagnetic nickel(11) salicylaldimine complexes. For the other salt solutions only the absorbance of the bis Finally, the relative proportions of the spinspecies a t 442 mp was obtained. This wave length corresponds t o the ligand-fieldmaximum of the diamagnet,icspecies. paired and spin-free forms of Nibnz++ in the various In those experiments where potentiometric data were ob- media were calculated using the observed extinctained (see Table 111)the absorbance was determined when E tion coefficients of 442 mp, and the values assigned was within 1-2% of 2.00.8 For the remaining solutions a spectrophot'ometric titrat,ion was conducted. An aliquot of above to the isomers.10 These results also are given a solution cont,aining known amounts of nickel chloride and in Table 11. bn.2HC1 was placed in a cylindrical optical cell, which in These results show that not only do perchlorate turn was placed in the sample compartment of the spectro- ions affect the equilibrium between the two forms photometer. The aliqu.ot was of such a volume that the cell was not completely filled, leaving room for additional liquid. but other salts do as well. Even chloride ions in The openings in the sample cell holder through which the high concentration bring about an appreciable increase in the fraction of the diamagnetic species. 6
(6) K. Sone an'd M. Kato, 2. anorg. albem. Chem., 801, 277 (1959). (7) F. H. Dickey, W.Fickett, and H. J. Lucas, J . Am. Chem. SOC., 74, 944 (1952). (8) Bn exception was macle with perchlorate because of preoipitation in this region of G. Spoctra were obtained on solutions in which E was 1.53 and the absorbance of Xien++wai taken into account in the calculations.
(9) R. H. Holm,
J. Am. Cham. SOC.,83, 4683 (1961). (IO) The intrinsic values of raf and ea,, are most likely insensitive to the medium changes judging from the behavior of the en and TetraMeen complexes. In 3.0 M NaC1, e m 4 and a 4 s for Niena++were found to be 7.2 and 6.7, respectively. and in 1.9 M NaNOa, e m for Ni(TetraMeenla++ was found to be 70.
1546
Vol. 66
NOTES
TABLE I1 1%TIYATION 0%RELATIVE PROPORTIONS O F THE SPIS-FREE
AND
SP~S-PAIRED FORUS o s Nibna+-k T = 25"
TABLE I11 THE F O R V A T ICOWSTAWS O~ OF
THE
Ni(II)-m~ao-2,3-
I)IAMINOBUTANE COMPLEXLS
T = 2 5 O , initially: 0.08 171 bn.2HC1, 0.13 LIT X C L
Obsd.
Added electrolyte
pK1,
pKza
Log Qi
Log Qz
Log Qa
Added electrolrto
extinction anefficient, of Nibns -1 4at, 442 ma
Hpinfree
Spinpaired
6 92 7 03 6 92
9 99 9 96 9 97
0.10 M NaCl 1 . 0 f%IKaC1 3 . 0 M NaCl Searly satd. KaCl 0.10 M NaK03 1.25 M Purar\T03 0.10 M NaC104 1 . 0 M NaC104 1.0 M KI 1 . 0 *ti' Na2804
24 21-24 33 43-49 24 30 22-26 30 35-40 22-25
0.74 .i7-0.74 ,64 .52- .45 (.74) .67 ,76- .73 .67 .61- .55 .76- .73
0.26 ,23426 .36 .48-.55 (.26) .33 .24- .27 .33 .39- .45 .24- .27
0 10MNaX0, 1 25 lk! NaxO! 0 10 JZ SaCl 3 0MNaCl 1 0114 NaC104
7 09
10 13 9 87
6.97 7 23 6 99 7 67 6 97
12 72 13 20 12 83 14 08 13 66
15 I6 15 17
EIrR ct,ion
However, the fact that an increase in the chloride in concentration from 0.3 to 3.0 M produces a relative increase of only 50% in the fraction of the spin-paired form shows that the interaction is weak. The large monovalent anions have a greater effect while the divalent sulfate ions have a negligible effect, at least up to 1.0 Jd concentrations. We attribute this behavior either to actual or to incipient ion-pair formation. In the classical cases of ion-pair formation between oppositely charged multivalent ions the electrostatic forces are strong enough to overcome the individual ionic hydration energies. On the other hand, the solvation energy is low along the axis perpendicular to the coordination plane of the diamagnetic nickel(I1) complexll and it is reasonable to expect that in this case weak association also can occur with oppositely charged ions provided their solvation energies are small enough so that the sum of the ionic hydration energies (which promotes dissociation) is less than the electrostatic attractive force. Thus, weakly solvated cations would interact more stroiigly with the large anions than with the small anions which have higher hydration energies. This argument is similar to that discussed regarding the solubilities of salts12 and, indeed, those salts having the greatest effect. Clod- and I-, also have the greatest tendency to form precipitates with the nickel(I1)-bn complexes. I n connection with the trien system, Jorgensen5 also has invoked ion-pair formation. The formation constants obtained for various media are given in Table 111. It is seen that all the &-values increase appreciably as the salt concentration increases. The changes can be accounted for to a large extent by considering changes only in the activity of the free ligand. The ratio of the (*onstantsfor the formation of a giren complex, *\[A,++, in two different media is given by
(11) D. Dyrssen and R f . Hennicks, Acta Chem. Scand., 16, 47 (1961).
(12) D. L. Leussing, "Treatise on Analytical Chemistry." I. hl. Kolthoff and P. J. Elving, Ed., Vol. 1, Pact. 1, Interscience, Publishere, Ins, NQW Uosk, N. Y,, 1B6Ob
6 91
52 36 66 90
where the y represent the activity coefficients of the various species. If to a first approximation the quotient ( y ~ 4 a , , z ( y ~ I ) l / ( y l l I a n ) l ( y is ~ ~equal ) 2 to unity, then R, is equal to (y"A)l/(y"L4)2. From this it results that R1 = RZ1/2= R3'/a, Comparing 0.10 and 1.25 ,7/1 nitrate media the observed values are 0.56, 0.33, and 0.14 for R1, R2, and RS,respectively. The square and cube roots of the last two quantities, 0.57 and 0.52, are in satisfactory agreement with the observed value of R1. For 0.10 and 3.0 X chloride, the observed ratios, R1, R2, and R3, are 0.26, 0.057, and 0.006, while R2'/2 and R3'/a are 0.24 and 0.18. I n these figures no exceptional increase in the value of Q2 with increasing salt concentration is manifested as would be the casc if strong ion-pair formation occurred. Thus, as was concluded from the spectral results, only weak ion-pairs are formed. Conductirity measurements on Nienz++ and Xi (C,C '-TetraMeen) ++ solutions also indicate that association of anions with the latter diamagnetic species possibly occurs. I n 0.10 M solutions at 22' the molar conductivities of the chlorides were found to be 223 for Nieiiz++ and 139.5 for Ki(C,C'-TetraMeen)z++. The former value is typical of those for di-univalent salts while the latter is closer to those for uiii-univalent salts.
THE HEAT CAPACITY OF THE SILVER CHA4LCOGEYIDES, Ag1.&, Agl.&e, AND Agl.,,Te FROM 16 TO 300°K.1 BY PATRICK PI'. KALSH, EDWARD W. ART,AKD DAVID WHITE Cryogenic Laboratory Deparfment o f Chemzstry, The Ohzo State Unzverszty, Columbus, Ohzo Recezied J a n u a r y 80, 1961
The heat capacities of the sulfide, selenide, and telluride of silver have been measured from the liquid hydrogen region to room temperature to determiile whether any macroscopic reordering processes which would affect the electrical properties of these semiconductors occur in this temperature range. The magnitudes of the heat capacities of these compounds at low temperatures are also of considerable interest because of their possible significance in thermoelectric cooling. Prior to this work, the only investigation of the heat capacities of the d y e r chalcogenides a t low temperatures was that of Gu1'tyae.i- and Petrov.* They measured the heat capacities of AgsS and (1) This xork vias supported by the Viright Ail Development D ~ v i sion, Air Research and Development Command, United States Air Force. (2) P. V. Gul'tyaev and A. V. Petrov Sovzet Phva, Bblid .%#le (riz. Tverd Tela), 1, 830 (1060),
