BiS (2,9-DIMETXI'IL-1,1
3029
@PHENANTHROLINE)CU(I) PERCHLORATE
above ,131,~. In Figure 5 it can be seen that omaxincreases with time; exfrapolation to to = 0 gives a value of @llm as 1.40 for sodium dodecyl sulfate in 0.04 M NaC1. If the area per molecule in an undisturbed film of sodium dodwyl sulfate is taken to be 40 A2 lo tRis limiting valine of @ indicates that Amin i s 28 A2. It is interesting that the A,,, value obtained in this way is nearly independent of electrolyte concentration even when tht: electrolyte considerably affects the sham of t h c aureole an?^ the value of /Arm. It is possible that when A,,, IS 28 i2 a close-packed layer of S moleeulita Pxist,3, and that aureoles with high electrolyte concentral,ions are flatter because their surface coverage befnr e compression is already higher.
So it seems that the observations of Figure 6 fit qualitatively into the same model as those of Figures 5 and 7 ; the aureoles with high electrolyte concentrations must remain flatter because they reach the limit For surface coverage A,,, at a smaller degree of cornpression @ l l m ) ~ Acknowledgments. The authors thank Dr. Karol 5. Mysels for daily guidance of their work at the R. J. eynolds Research Laboratory. The help of Dr. P. Sornasundaran and Dr. L. R. Rarriagton 1s gratefully acknowledged. (10) E. Matijevic and B. A. Pethica, Trans. Faradag f h c . , 54, 1382 (1958).
Ion Pair Formation of hyl-1,10-phenanthroline)copper(1) Perchlorate. I. cohols and Ketones suhiko Miyoshi Depwtment of Chemistry, Faculty o j Science, Hiroshima University, Hiroshima, J a p a n
(Received February 28, 1072)
Conductance of bis(2,9-dimethyl-l,lO-phenanthroline)Cu(I)perchlorate was measured in methanol, ethanol, 1-propanol, I-butanol, acetone, and methyl ethyl ketone at 2.5'. Experimental data were analyzed with the Fuoss--0nsager conductance theory (1965). The chelate salt was found to be more associated in alcohols than in ketones. This behavior was explained by the solvation of ketones to the chelate cation. Perchlorate ion does not seem to be solvated by either alcohols or ketones. The association constants obtained here were much smaller than those of Bu4NCl04 which is hydrodynamically a little smaller than the chelate salt. This fact can be interpreted in terms of the decrease In charge density on the chelate cation arising from the coordination bond between Cu(1) and the aromatic ligands.
Introduction Recently K a y and Evans, et u Z . , ' - ~ have extensively studied the conductance behavior of a number of tetraalkylarnmonium salts in aqueous and various organic solvents. They found that in water and alcohols, salts with large anions, such as perchlorate and iodide ions, are more associated than those with sinall ions. They also measured the conductance of alkali metal perchlorates in acetonitrile6 and found that salts with large oations, such as cesium and rubidium ions, are more associated than those with small ions. These observations seem to conflict with the ionic association theories.? * However, Evans and Kay4v9 discussed the dicmepancy in terms of the difference in the extent of solvation and concluded that in a solvent
with solvating power, small ions are well solvated and therefore so-called solvent-separated ion pairs are predominantly formed, but large ions are almost bare in (1) D. F. Evans, C. Zawoyski, and R. L. Kay, 1.Phys. Chem., €19, 3878 (1965). ( 2 ) R. L. Kay, C. Zawoyski, and D. F. Evans, ibid., 69, 4208 (1965). ( 3 ) R. L. Kay and D. F. Evans, ibid.,69, 4216 (1965). (4) D. F. Evans and R. L. Kay, ibid., 70, 366 (1966). (5) R. L. Kay, D. F. Evans, and G. P.Cunningham, ihid., 73, 3322 (1969). (6) R. L. Kay, B. J. Hales, and G. P. Cunningham, ibid., 71, 3925 (1967). (7) N. Bjerrum, K . Danske Videnslcab. Selskab., Skrifter, Naturoidenskab. math. Afdel., 7, No. 9 (1926). (8) R. M.Fuoss, J. A m e r . Chem. Soc., 80, 5059 (1958). (9) D. F. Evans and P. Gardam, J . Phys. Chem., 73, 158 (1969). T h e Journal ofPhysica1 Chemistry, Vol. 7 6 , N o . $1, 1972
solvents such as water and alcohols and form contact ion pairs. They also have systematically investigated the interaction OF water with the hydrocarbon portion of the tetraalkylainmoniuan~~um ions from the viewpoint of its transport proper tie^.^*'^^'^ It is the purpose of this study to investigate, by means of the coiiduatance measurement, the interaction with solventh of the chelate salt in which the cation 1s surrounded by large aromatic groups and to see whether Evans's explianaticn for the association behavior holds for thnc, systeii. Since the charge of the cation seems to be screened by a significant amount by the bulky groupF, I C is expected that there may be some interesting diffcn en-es between this complex cation and small inorganic ions frequentlv measured regarding the Interaction n i t h solvents.
analytical data were as follows: found (%j: R, 4.07; C, 57.79; N, 9.66 (calcd (%): 13, 4.17; C, 58.03; N, 9.67). Reagent grade methanol mas refluxed over calcium oxide, distilled, and then refluxed over magnesium. methoxide. Finally, it was purified by fractionation and a middle fraction was collected. All procedures were done under nitroget1 gas t o a,void the oxidation of m.ethanol. Its specific conductance was 3.0 X IO--' ohm-' 6m-I. Ethanol v a s purified in the same manner as methanol. Its speciiie con,ductance was 6 X IO.-.*ohm-' em-]. Propanol and butanol were refluxed over calcium oxide and fractionatled. Their specific conductances were 6.6 X IO-* atid 3.2 X ohm-' crn-l,, respect'ively. 'Their purity was checked by measuring the densities. Observed values were 0.78665, 0.78518, 0.79954, and O.HO570 for M.eCPR, EtOH, PrOH, and BuOR, respectively, which were in good agreement with the reported vad-ues in .t,lieliterature, that is, 0.78658,25.7851.1,13 0.79PO0,1J and 0.80576,9 respectively. Ketones were refluxed with pot'assiuril permanganat,e and were distilled over activat'ed alumina. Specific conductances were 3 X 4.0+ ohm-' cm--l for acetone and 3.9 >( cm-l for methyl ethyl. ketone (abbreviated to RZEK), the densities of which were 0.78468 and 0.9955Py respectively. The corresponding values reported are 0.?843314 and 0.7845315 for the former and 0.79955l6 for the lat,ter. The value of 0.8011 report'ed by Fuoss'7 for 1'6EI.I is considerably high, which is probably due to water contamination. Water content in solvents used was examined by t'it'ration with Karl Fischer reagent and found to be less than 0.01 wt % in each case. Other physical properties of the solvents used were quoted from the literatures cited above.
Apparatins and Procedure. The Pyrex conductance cells were of erlenmeyer type with lightly platinized eleci rodes, coiitaining about 200 ml of solution, and were standardized by KC1 solutions, using the Lind, constants. Potassium chloride Zwolenik, and FUOSS was recrystallized t v e e times from conductivity water and dried at l100" The cells were cleaned with chromic acid mixturets t o elrminate traces of organic materials (solvents and complex salt) and with nitric acid. They mere washed with conductivity water fully and with steam before each run. Measurements were carried out by a contiuctometer, type MY-7, from Yanagimoto Nfg. Co , Lttl , nitl- the Wheatstone bridge (SO0 c/sec). All the resis tames were calibrated by the resistance box of Yokogawa Electric Works, Ltd., which was standardized by Shirikawa Electric Co. Cell solutions were thermostated to 25 rt 0.01" in a double water Results bath with B mercury-in-glass thermoregulator, In Equivalent conductance A and the corresponding order to hasfrn the temperature equilibrium and to concentration in equivalent per liter C are given in the avoid thc polarization effect during the measurement, microfilm editions of this journa1,l8 Th.ese data were the solution in the cell was stirred by a magnetic ana,lyzed with the Fuoss-Onsager conductance theory stirrer. All solutions were prepared gravimetrically (1965)19 in the form since the molecular weight of this salt is very high (579.222). For each run, a sample of 100-110 mg was directly aclded to tb.e solvent of known conductivity in (10) R . L. Kay and D. F. Evans, J . Phys. Chem., 70, 2325 (1966). the cell. After a constant value of resistance was (11) D. F. Evans, G. P. Cunningham, and R. L. Kay, ibid., 70, 2974 (1966). attained, 38 .id JC solution was siphoned out and equal Fuoss, J . Amer. Chem. Soc., (12) L. Lind, J. Zwolenik, and R. AM. volume of thc solient was added to the cell. This 81, 1557 (1959). procedure u'a9 continued to an appropriate concentra(13) D. F. Evans and P. Gardam, J . Phys. Chem., 72, 3281 (1968). tion. A dryhox urth dry nitrogen gas was used for the (14) D. F. Evans, J. Thomas, J. A. Nadas, and S. M. A. Matesich, ibid., 75, 1714 (1971). preparation aR solutions because the solvents used were (15) G. Pistoia and G. Pecci, ibid., 74, 1450 (1970). rclatively hygroscopic. The concentration range was (16) 8 . R. C. Hughes and D. H. Price, J . Chem. SOC.A , 1093 (1967). mithin the limit of the applicability of the extended (17) F. M .Sacks and R . M. Fuoss, J. Amer. Chem. Soc., 75, 5172 conductance theory ( m < 0.2). The solubility of the (1953). sal1 in butanol was 90 small that a sample of 50 mg was (18) These data will appear following these pages in the microfilm edition of this volume of the journal. Single copies may be obtained used in this case. from the Business Operations Office, Boolra and Journals Division, iMnterzabs. BSis{2,9-dimethyl-l,10-phenaJnthroline)-American Chemical Society, 1155 Sixteenth St., N.W., Washington, D. C. 20036, by referring t o code number JPC-72-3029. Remit CuiI) perchlorate was prepared by Dozin Chemicals, check or money order for $3.00 for photocopy or $2.00 for microLtd. Its p m t y was checked by analysis. The fiche. ~
The Journal g f P h y s i c u 1 Chemistry, Vol. 7 6 , No. % I , 1971
AQ
Acetone MEK MeOEI EtOH PrOH BUOW PhNQ2
1 7 6 . 7 4 r t 0.02 137.68st0.04 107.95 rt 0.07 50.86A0.04 26.90A0.04 18.77 3 ~ 0 . 0 4 81.865 If 0.004
R -- BO -. S(C‘)”/”
a,
A
4.87 st 0.01 4.97 rt 0.04 3.27 rt 0.25 3.88i0.18 3,923 rt 0.35
3.1 st 0.8 4 . 3 3 st 0.04
.+ E’C In (6E1’C) + LC
KA
L
0 12.8 i 1 . 6 3 . 5 rt 1 . 9 30.3 rt 3 . 0 82 f 10 119 Jr 34 0
2929 2344 626.6 482.0 247 8 -138.2 264.0
@A
~
0.04
0.025 0.024 0.007
0.008 0.018 I)I008
(1)
for unassociated electrolytes and in the form A = A,,
-k E‘G? In (GEi’Gr)
-+ LCr - K ~ C y f , t ~ h(2)
for associated electrolytes. The symbols used here have their usual meaning. Although this large electrolyte is expected to affect the solution viscosity to a , viscosity correction was made, considerable ex ; e ~ t , no since it is well known1 that such a correction has no effect on AO cr K a and only results in small changes in All the calculations were carried out by an electronic computer, TOSBAC 3400, using the leastsquares method proposed by Kay.21 I n acetone a standard deviation when treated with eq 1 was smaller than that wilh sy 2 and a negative K Avalue of smaller magnitude t hail the standard deviation was derived. Therefore it should be concluded that this salt is completely dissooiated ~n acetone. On the other hand, in methanol it was reasonable to treat the data with eq 2 since a positive K A value of larger magnitude than the standard de7,iarion was derived. These facts are the criteria, generaliy accepted as a good evidence that an electrolyte is asFociated in any solvents.z2 All the parameters obtained are summarized in Table X together with those in nitrobenzenez3for comparison. Thr precision of those in butanol is somewhat poor because of the low d u b i l i t y of the chelate salt in this solvent.
iscussion Figure 1 shows the plot of the log K A vs. 1/D where the arrow correspondE to acetone in which this chelate salt is complesely dissociated. Although log K A is expected to be linear in the reciprocal dielectric constant according to the ionic association theoryS log R
A
=
log Kao 4- e2/akDlcT
(3)
a curvature is observed in this case. This behavior, also found in dioxane-water has been explained generally by some kinds of solute-solvent interactions. But it is not clear what kind of the interaction works 111 this case because of the lack of the data in other and mired solvents. Such an investigation is under ~5 ay in our laboratory.
2 I
I
I
8
I
Figure 1. Plot of log K A us. 1/D.
It seems worthwhile to compare the results in alcohols with those in ketones since Figure 1 shows that this salt is more dissociated in ketones than in alcohols. The same pattern has been observed for tetraand alkylarnmonium salts by Evans, et u1.,5,9,13,14329 Pistoia.16 Evans, et al., discussed the association behavior of these salts in several alcohols and proposed &hetwo-step aasociation mechanism involving solventseparated and contact ion pairs
(4) where K1 refers to the solvent-separated ion pair forma(19) R. M . Fuoss, I,. Onsager, and J. F.Skinner, J. Phys. Chcm., 69, 2581 (1965). (20) Though in principle, the 1965 equation demands that the parameter a be derived from L and K A terms, this makes the estimation of the standard deviation quite difficult. Therefore a was derived from only the L term since the magnitude of the standard deviation is essential to the discussion described below. I n practrce this treatment is equivalent t o that used in the 1957 equation except that the activity coefficient is expressed by the Debye-IfUckel limiting form. (21) R. L. Kay, J. Amer. Chem. Sac., 82, 2099 (1960). (22) J. L. Hawes and R. L. Kay, J. Phys. Chem., 69, 2420 (1965). (23) K. Miyoshi and T. Tominaga, unpublished data. (24) 9. M . A. Matesich, J. A . Nadas, and D. F. Evans, J . Phys. Chcm., 74, 4568 (1970). The Journal of Physical Chemistry, Val. 78, K O .21, 1971
3032
MATSUHIKO l/IrYosra
tion and K S to the contact ion pair formalion. They found that, while K 1 is nearly independent of a k over some finite region of ion sizes, K , becomes larger as the anion size is increaschd and therefore the overall association constant is greater for large anions than for small ones which. prefer the solvent-separated ion pairs. Thur, bis(2,g.-djmetliyl-l,10-phenanthroline) Cu(1) perchlorate mag be assumed to form a contact ion pair in aIcoliols. Table I1 sho~vsthe ionic Walden product of the chelate cation An inspection of Table I1 clearly indicates thac the values obtained in alcohols are almost constant with the change in solvent and are nearly equal t o the value in nitrobenzene which has no acid-base properties and so (cannot solvate electrolytes to a significant extent. This fact is used by Kay, et uL,l for the eqilanation 3f the abnormally large association constants of horganic salts in this solvent.
Table ET: Ion Mobility, Walden Product, and Ionic Walden Product AQ +q b.0 -1.
doll
MeOH EtOH
37.08 0.588 19.3(18.8)b 0.551
PrQH Acetone
10.5. 7.51P 58. 5 e
R4EK PhNOz
lP,OJ
BuQW
0,525 0.486 0.536 0.520 0,586
A0
+q
0.201 0.209(0.204) 0.205 0.195 0.177
(BUN ")
0,212g 0.213h 0,209h 0 . 203i 0.201i
0.202
A"-' = 70.94, ref 5 . b Xo- = 31.6 or 32.1: E. D. Copley, et al., J . Chem. Soc., 2492 (1930); A. R. Gordon, et al., J. Amer. Chem. SOC.,79, 2352 (1957). Xo- = 16.42, ref 13. Xo- = 11.22, ref 9 . Xo- = 118.35, ref 14. Xo- = 20.9; C. A. Kraus, et al., *J. Amer. Chem. Soc., 69, 2472 (1947). g Reference 2. * Referensx 13. Reference 9. i Reference 14.
BUAN ion for comparison. Despite the fact that hydrodynamically this complex cation is only slightly larger than Bu4Ni, the association constants of the complex salt are much smaller than those of Bu4XC104, which are reported to be 769 f 6,1J 2200 h 20,9 and 90 f 214 in PrOH, BuOH, and acetone, respectively. This may be explained by the asiumptiorr that thc charge on the chelate cation is well distribclted over the ligands through the coordination bend bet n-een Cu(I) and the aromatic ligands. Therefore its charge density is considerably reduced, whereas m Bu4X'-its charge is localized chiefly on the nitrogen atom. These features seem to cause the difference between the complex salt and Bu4NC104in the M A values. On the other hand, ketones are good hydrogen bond acceptors but have no donor hydrogens, and they are probably not as associated as alcohols. The direction of the dipole suggests that ketones have an affinity for cations but not for anions. The iomc Walden product for the cation obtained in acetone is found to be much smaller than those in alcohols and nitrobenzene from the Table 11. Though the limiting mobility of Clodin MEK is not available in the literature, the Welden product is smaller than that in acetone. This means that the ionic Walden product in bK&X may be smaller than those in alcohols. Of course the relative size of the cation to solvent molecules is large enough t o satisfy the hydrodynamic conditions for the Stokes law to hold. As Fu0ssZ5and ZwunzigZ6pointed out, there may be a contribution from the dipole relaxation of the solvent molecules t o the mobility iormulateed as
5
1.
From the above discussions, it may be reasonable to assume contact ion pairs in this case. Stabilization of ions by alcohols does not seem to work here due to the large size of the component ions and low number of active monomeric alcohols available for solvation. Kay, et aZ.,2 pointed. out that it is possible for alcohols to stabilize the ion pairs by hydrogen bonding for bulky eled.rolytes. 'Their assumption holds also for this chelate salt in alcohols. However, it may be possible that since thr: dielectric constants of alcohols are larger than expected from their dipole moments owing to the structure arising from hydrogen bonding, local breaking of the bonding by the dissolution of electrolytes decreases the effective dielectric constant of alcohols, so that the ion pair formation becomes more favorable. The fact that this salt is more dissociated in dioxanewater mixturesz3than in pure alcohols suggests the possibility of the decrease in effective dielectric constants. 113 Tab1 2 li 1 arc included the ionic Walden product of The Journal of Ph,g/sicul Chemistry, Vol. 7 6 , N o . 21, 1972
A07 =
F 2 / N ( 6 n r 4- B / r 3 )
(5)
where B is a function of the dielectric constants and viscosity, and so a decrease in the ionic Walden product is expected with decreasing dielectric constant. The fact that these values are almost constant in alcohols while the dielectric constant varies from 32.62 to 17.45 leads us to assume that the contribution is decreased considerably owing t o the large size of the cation. Then the solvation of ketones to the chelate cation appears to retard the ionic migration and promote the dissociation of the chelate salt, On the other hand, the Walden product of Bu4N+ in acetone is nearly equal to those in alcohols. From the discussions described above, this chelate salt prefers a solvent-separated ion pair with the solvent held to the cation in ketones. This is a reasonable explanation for the association behaviors shown in Figure 1. However association constants of lithium and sodium perchlorates in acetone were reported t o be 5300 and 4300, r e s p e c t i ~ e l y . ~Such ~ abnormally large (25) (26) (27) 1257
R. M. Fuoss, Proc. Nut. Acud. Sei. 1J. S., 45, 807 (1959). R. Zwanzig, J . Chem. Phys., 38, 1603 (1963). R. Fernandez-Prini and J. E. Prue, Tians. Faraduy Soc., 62, (1966).
KIS(~,~-DXMETBYL-~, WPHENANTHROLINE) Cv(1) PERCHLORATE values make u:, assume that Li+ and Na+ as well as Clod- are not ($0 strongly solvated and therefore form contact ion pairs. BLSO the dielectric saturation may f7 'The complex cation has specific -11 is surrounded by large aromatic groups and iitF: c h q y may be screened considerably. Therefore the intwaa:r,ion between the complex cation and ketones muy br: different from that between alltali metall ions and ket>ones. Nay, et al.," reported that for h g e ions, the degree of solvation is determined predominantly by this dipole moment of the solvent molecules, whereas for small ions, it is determined predominant iy by the acid-base properties of the solvent molecules. This may be one of the answers to the problem oi the difference in K A values between this compXex and alkali metal perchlorates. The same cnucluenon niay be derived from the consideration o f tFle values of the distance of the closest approadeh parametw a. From Table I1 it is seen tlial ( I value^ nblairtrd in ketones are about 5 8, which are greater by 1 than those in alcohols and nitrobenzene. 'The incream in u values in ketones may be attributed t o solvation, However, it is well known that in order t o obtain definite a values from the curvature o; the phoreogrzlrn, n very high accuracy in data is reuired and that a values show a considerable depenence upon
