THE JOURNAL OF
PHYSICAL CHEMISTRY (Registered in
VOLUME 66
U.S. Patent Office)
(0Copyright, 1962, by the American Chemiaal Society)
NOVEMBER 15, 1962
NUMBER11
A STUDY OF THE CONDUCTANCE BEHAVIOR OF LITHIUM AND AMMONIUM IODIDES IN TL-BUTANOL'-~ BY H. V. VENKATASETTY AND GLENNH. BROWN^ Department of Chemistry, University of Cincinnati, Cincinnati, Ohio Received September 6, 1981
The conductances of lithium iodide and ammonium iodide have been measured in n-butanol a t 0, 25, and 50". The values of the dissociation constants for the ion-pairs and the limiting equivalent conductances are obtained by the method of Shedlovsky. The ion-size parameters are calculated by Bjerrum's method as well 8,s by the Denison-Ramsey method. Both of the salts are found to behave as weak electrolytes. The observed differences in the dissociation constants and the mobilities are explained in terms of the specific types of ion-ion and ion-solvent interactions.
Introduction The investigation of the properties of solutions of ionophores6 in non-aqueous systems has received corrsiderable attention in recent years with a view to understanding the nature of ion-ion and ionsolvent interactions in these systems under a wide variety of conditions. A large number of 1 : l electrolytes have been investigated in a wide variety of solvents.0 A literature survey of nonaqueous systems reveals the absence of conductance data for solutions of ionophores in n-butanol over a range of temperature. Seward' measured the conductances and viscosities of the solutions of tetra-n-butylammonium picrate in n-butanol over the entire concentration range from dilute solution to fused salt at 91'. +Butanol, being an associated solvent with low vapor pressure and moderate value of dielectric constant of 17.1 a t 25", offers an interesting solvent to study specific ion-ion and ion-solvent interactions. (1) Presented a t the 139th National Meeting of the American Chemical Society, St. Louis, Missouri, March 21-30, 1961. (2) This paper is abstracted in part from a dissertation submitted by 13. V. Venkatasetty to the Graduate School of the University of Cincbinnatiin partial fulfillment of the requirements for the Degree of the Doctor of Philosophy, 1961. (3) The conductance equipment used in this research was made available through a grant from the Research Corporation. (4) Department of Chemistry, Kent State University, Kent, Ohio. ( 5 ) R. N. Fuoss, J . Chem. Educ., 32, 527 (1955). (6) H. 8. Harned and B. B. Owen, "The Physical Chemistry of Electrolyte Solutions," 3rd Ed., Reinhold Publ. Corp., New York, N. Y.,1958, Chapter 16. (7) R. P. Sen-ard, J . A m . Chem floe., 78, 515 (1951).
Experimental Apparatus.-The conductance measurements were made using a conventional alternating current Wheatstone bridge equipped with a Wagner earthing device. The power source for the bridge waa a Leeds and Northrup oscillator with a frequency of 1000 c.p.5. The balance detection device consisted of a crystal headphone. The ratio a r m of the bridge consisted of a Kohlrausch slide wire purchased from Leeds and Northrup and in almost all the measurements great care waa taken in measuring resistances to balance the bridge at the center of the slide wire to eliminate errors due to inductance.* The conductance cell used waa designed according to the recommendations of Jones and BollingerO and had a graded seal between the Pyrex and the platinum. The electrodes were slightly platinized usin a standard rocedure.'O The cell constant was determine3 using 0.01 KC1 solution according to the procedure of Jones and Bradshaw.Io The cell constant was 0.15000 em.-' at 0.0" and the cell constants a t the other temperatures were calculatedll using the 0.0" value. The cell constant calculated for 25.0" was checked experimentally using the procedure of Ives and Sames,la with benzoic acid solutions prepared from a National Bureau of Standards sample. The cell was maintained a t 25.00 i 0.02" and 50.00 f 0.02" in an oil bath and the measurements a t 0.0" were made in an oil trough placed in a well insulated ice chest.'$ The temperature of the bath wm measured with a Beckmann thermometer which had been calibrated against a thermometer certified by the National Bureau of Standards. Materials.-A saturated solution of lithium iodide trihydrate of "National Formulary" quality, purchased from
%
(8) G. Jones and R. C. Josephs, {bid., SO, 1049 (1928). (9) G. Jones and G. M. Bollinger, ibid., 63,411 (1931). (10) G. Jones and B. C. Bradshaw, ibid., 56, 1780 (1933). (11) E. W. Washburn, ibid., 88, 2465 (1916). (12) D. J. G. Ives and K. Sames, J . Chem. Soc., 511 (1943). (13) G. Jones and M J. Prendergast, J . Am. Chem. Soc., 69, 731 (1937).
207 5
2076
H. V. TENKATASETTY AND GLENNH. BROWN
the Mallinckrodt Chemical Works, was prepared in conductivity water and freed from anionic impurities by passing it through the anion-exchange resin Amberlite IRA-400. After evaporating the solvent from this solution under vacuum, and cooling, lithium iodide crystallized. These crystals were dried in a vacuum oven at about 120’ and stored over PtOs in a vacuum desiccator for several weeks. The solid lithium iodide was placed in a specially designed steel bomb and heated under vacuum in a furnace a t about 650” for about 2 hr. and then the tem erature was raised to about 850” for 1.5 hr. to completely &hydrate the salt. The salt showed no alkalinity and analysis of the sample by Fajans’ method showed it to be better than 99.96% pure. The ammonium iodide was Baker analyzed “Reagent Grade” and a saturated solution was made in de-ionized water freed from dissolved oxygen by boiling. This eolution was passed through a column containing the cationexchange resin Dower 50-X, equilibrated with 2 M ammonium nitrate solution, and then washed free of nitrate; the cation-exchanger was followed by an anion exchange resin IRA-400 which had been equilibrated with . 2 potassium iodide solution and then washed free of iodide ion. The efiuent was collected in a colored bottle to protect the ammonium iodide from photochemical action. The solution of ammonium iodide was evaporated under reduced pressure, the salt recrystallized from an ethanolether mixture, and dried to constant weight in vacuo a t 6065”. Analysis by Fajans’ method showed a purity of 99.98 %. Reagent grade n-butanol was refluxed for several hours with lime and distilled three times. In the final distillation the fraction boiling between 117.3 and 117.5’ under atmospheric pressure was collected. The refractive index of the purified solvent was 1.3992 compared to the literature value of 1.3991.14 The specific conductance of the solvent reported in the literature’s is 9.12 X 10-9 mho at 25”. The dielectric constant of n-butanol calculated according to Circular No. 514 of National Bureakt of Standards‘Bis 20.44 at 0”, 17.1 a t 25’, and 14.10 a t 50 , The viscosities of nbutanol a t different temperatures reported in the literature are 0.05186 poise a t O 0 , I 7 0.0246 poise a t 25’,’8 and 0.01411 poise a t 50”17; these values check well with the extrapolated values of viscosity temperature functions of liquids found in the literature.19 The densities of n-butanol determined in thw work are 0.8246 g./cm.S at O’, 0.8057 g./cm.a at 25”, and 0.7875 g./cm.8 a t 50’; these values are in agreement with the values reported in the literature.20 The purified solvent was stored in sealed bottles and any transfer of the solvent was made in a drybox under a positive pressure of dry nitrogen. Preparation of Solutions.-All solutions were prepared in a drybox b transferring to a container a known weight of the salt anTdissolving i t in an exact volume of the solvent. Dilute solutions were prepared by further dilutions, using calibrated pipets and burets. Concentrations were established by analysis. These solutions were preserved in bottles which had been coated on the outside with black paint and were sealed with serum caps. n-Butanol is found to have no effect on these serum caps. Transfers of solutions always were made using the proper sized hypodermic syringe. All weighings were corrected for the buoyancy effect of air. Procedure for Making Measurements.-The cell was first cleaned with potassium dichromate-sulfuric acid cleaning solution, washed several times with distilled water, and findlv with de-ionized water. It then was rinsed well with “chekically pure” acetone, dried, and stored in a drybox. From this oint on, all handling of solutions was carried out in a dryiox. Before each filling of the cell, it was rinsed well a number of times with the solution under study. The (14) “International Critical Tables,” Vol. 7, McGraw-Hill Book Co., Ina., New York, N. Y.,1930, p. 36. (15) L. Scheflan and M. B. Jacobs, “The Handbook of Solvents,’’ D.Van Nostrand Co., Inc., New York, N. Y., 1953, p. 159. (16) National Bureau of Standards Circular No. 514, 1951. (17) “Handbook of Chemistry and Physics,” 40th edition, Chemical Rubber Publishing Co., Cleveland, Ohio, 1958, p. 2158. (18) Reference 14, p. 41. (19) A. N. Nissan. Phil. Mag., 33, 441 (1941) (20) J. Timmermann, “Physico-Chemical Constants of Pure Organic Compounds,” Elsevier Publishing Co., New York, N. Y., 1950, p. 819.
Vol. 66
transfer of the solution from the bottle to the cell was conveniently accomplished by the use of a syringe fitted with a suitable hypodermic needle that waa inserted through the stopcock into the filling tube of the cell. The conductance cell was placed in the appropriate bath and allowed to come to temperature. Readings were taken every 20 min. until the resistance was constant. Since the specific conductance of n-butanol is 9.12 X 10-9 mho at 25’, solvent correction is not necessary even for the most dilute solutions.
Results and Discussion The data for the equivalent conductance of lithium iodide and ammonium iodide corresponding to different concentrations a t 0, 25, and 50’ are recorded in Table I. The maximum experimental error inherent in these data is 0.1%. TABLE I EQUIVALENT CONDUCTANCE O F LiI Temp.
0.0” c
x
104
Temp. 25 i 0.02’ c x 104
N
Aobsd
N
nobed
8.21 16.42 20.50 24.62 41.00 61.50 82.10 102.60 123.10 143.70 164.20
6.37 5.74 5.44 5.27 4.66 4.32 3.87 3.66 3.59 3.28 3.05
8.02 16.04 20.05 24.06 40.10 60.10 80.20 100.20 120.30 140.40 160.40
12.13 10.69 10.17 9.75 8.69 7.79 7.02 6.73 6.30 5.98 5.91
EQUIVALENT CONDUCTAWE O F NHJ c
x
n-BUTANOL
IN
Temp.
50 f 0.02O
c
x
104
N
7.84 15.68 19.60 23.52 39.20 58.70 78.40 97.90 117.60 137.20 156.80 IN
hobad
19.39 16.57 35.64 15.02 13.02 11.64 10.71 10.06 9.53 9.15 8.81
n-BUTANOL
Temp.
Temp.
Temp.
0.00
25 & 0.02’ 104
50 i 0.02°
104
0
N
Aobsd
6.79 13.57 16.90 23.74 33.90 50.90 67.80 84.80 101.80 118.70 13F.70
6.86 6.07 5.84 5.43 4.97 4.42 4.07 3.94 3.78 3.57 3.49
x
N
6.60 13.26 16.57 23.20 33.15 49.70 66.30 82.90 99.40 116.00 132.60
c Aobsd
12.23 10.99 10.59 9.84 8.98 7.97 7.28 7.10 6.77 6.33 6.32
x
104
N
6.50 12.96 16.20 22.68 32.40 48.60 64.89 81.00 97.20 113.40 129.60
hobid
19.25 16.53 15.84 14.43 13.04 11.67 10.77 10.08 9.51 9.07 8.55
The plots of equivalent conductance us. the square root of concentration for both LiI and NHJ a t the three temperatures studied show marked deviation from linearity. This indicates that these salts behave as weak electrolytes in n-butanol with definite ion-pair formation. The limiting equivalent conductance (Ao) and the ion-pair dissociation constant ( K ) for the different solutions were evaluated by the Shedlovsky method.21*22Figure 1 shows the Shedlovsky plot for ammonium iodide. Comparable curves are obtained for lithium iodide solutions. The values of Roand K are summarized in Table 11. The conductance of ions to a large measure depends on the size of the ions and also on the vi-cosity of the medium through which they move. (21) T. Shedlovsky, J . Franklin Inst.. 336, 739 (1938). (22) R. M.Fuosa and T. Shedlovsky, J . Am. Chen. Soc., T i , 1496
(1949).
KOV.,
CONDUCTANCES OF L1THIUX.I
1962
ANC
AMMONIUMIODIDES I N %BUTANOL r---
TABLE TI SUbIAfARY OF
I)ATAFOR
2077
SOLUTIONS O F
LiI
AND
"41
IN
n-BUTAA O L Electrolyta
Temp.,
LII LiI LiI NHJ NHJ NHJ
0.0 25.0 50 0 0.0
"C.
Ao Exp.
8.58 15.42 30.30
K X 10'
anj," om. X 108
aDRpb
om. X lo6
A043
20 9,
3.54 4.81 0.444 13.i4 3.71 4.94 .429 7.41 3.83 5.05 ,427 9.09 17.39 3.34 4.67 .471 ,394 25.0 16.00 19.53 4.15 5.21 50.0 29.85 6.78 3.76 4.98 .421 4 ~ B is J the ion-aize parameter from Bjerrum's theory. * a D R is the ion-size parameter from the Denison-Ramsey equation.
The interaction between two oppositely charged ions depends on the charge, size, structure, and '0 2 4 6 8 IO 12 14 16 polarizability of the ions as well as the interaction of the ions with the solvent molecules. n-Butanol CAS(r)f?x IO3. being an electron donor solvent with the dipole for ammonium iodide at 0, 25, moment of 137 Debye units is expected to solvate Fig. 1,Shedlovsky plotand 50'. the cations much more strongly than the anions. The values of A. and K for LiI and NHJ a t 0, 25, and 50' are comparable to each other. Since the ammonium ion, with a tetrahedral Therefore, the cations Li+ and NH4+ are either configuration, is large relative to the Li+ ion and of comparable size or there must exist specific since it has a lower charge density than the Li+ interactions between these cations and the polar ion, it cannot exert as strong an attractive force solvent molecules which affect the motion of these on the solvent dipoles to solvate itself as is the ome ions through the solvent medium. with the Li+ ion. However, the experimentally The model proposed to explain the small values observed low values of A, (Table 11) may be exof A. for LiI solutions assumes that the Li+ ions, plained, in part, as due to the electrostatic attracbecause of their small size and high charge density, tion between the charge on the NH4+ ion and the can exert a strong attractive force on the polar solvent dipoles. This interaction can produce solvent molecules and thus create a tight solvent a distortion of the quasi-crystalline structure atmosphere in the first solvent shell around the of the solvent and thus interfere with the motion cation. This model is supported by experimental of the ions. This coulombic ion-dipole interaction evidence from other measurementszs of A. and may be reinforced by hydrogen bonding between from X-ray data on the structure of s o l ~ t i o n s . ~the ~ ~nitrogen ~~ of the ammonium ion and the oxygen The ion-size parameter calculated using Bjerrum's of n-butanol, possibly a t all the corners of the theoryz6 of ion-pair formation and the Denison- tetrahedron. Even though these interactiqns do Ramsey method (Table 11) is larger than the not give a tightly held solvent atmosphere around crystallographic radius of 2.76 A. (Pauling) for the ammonium ion, they can account for the low LiI. These results, together with the nearly con- values of A, and the large values of K. This stant Walden product (Table 11)at all the tempera- model of the ammonium ion in n-butanol is suptures studied, indicatc that the Li+ ions are sur- ported by the reasonable agreement between the rounded by a single layer of solvent molecules calculated Bjerrum ion-size pararnpter (Table 11) and that there is little change in the size of the and the crystallographic interionic distance of solvodynamic unit over the entire temperature 3.63 A. (Goldschmidt). Burgess and Kr@us2* range. The large values of the dissociation con- evaluated A. and K for NHJ in pyridine at 25'. stant ( K ) can be explained by realizing that strongly Their values of A,, = 95.2 and K = 2.4 X coordinated solvent molecules around the Li+ compared to the present values of A. = 16.00 and ions hinder the close approach of iodide ions to K = 19.53 X lo-' suggest that there are specific form stable ion-pairs. The values of A. and K ion-solvent interactions in n-butanol and a very obtained for LiI in this study compare favorably small amount of such interactions, if any a t all, with the values obtained by Ogstonz7for LiI 3H20 in pyridine. in ethanol. (23) Reference 6,pp. 098-704. (24) G. W. Brady, J . Chem. Phys., 27, 304 (1957). (2.5) M. Strauss, Dortoral Dissertation, University of Cincinnati, lOGO
(26) N. Bierrum, KO!.Danske Videmkab. 7, No. 9 (1928). (27) A. 0.Ogston, T r a m . Faraday Soe., 8% 1679 (1938). (28) D. 9. Burgem and C. A. Kraus, J . Am. Chem. SOC.,TO, 706 ( 1 949).
