J. Phys. Chem. 1982, 86, 1035-1037
1035
Ionic Equilibria in Pyridine-Iodine Solutions S. Aronson,' P. Epsteln, D. B. Aronson, and 0. Wleder Department of Chemistry, City University of New York, Brmkwn College, Brwklyn, New York 11210 (Received: August 23, 1981; In Final Form: October 23, 198 1)
Electrochemical cells of the type, Ptl12,KI in PyllI, in PylPt, in conjunction with electrical conductivity measurements have been used to study ionic equilibria in 0.16 to 1.6 M solutions of iodine in pyridine. The experimental data are discussed in relationship to a mechanism for the ionization of a pyridineiodine molecular complex. The results indicate that 8-1870 ionization of the iodine occurs.
Introduction Iodine reacts with pyridine in organic solvents to form a stable molecular complex.'-3 The chemical reaction for the complexation can be written Py I, = PyI, (1) The equilibrium constant at room temperature is in the range of 2 X lo2L/moP4 and the heat of complexation is about -8 kcal/m01.~3~ An interesting and unusual aspect of the interaction of iodine with pyridine is the observation first made by Audrieth and Birr' that solutions of iodine in pyridine are good electrical conductors with conductivities of the same order of magnitude as aqueous salt solutions. Several investigators have tried to elucidate the nature of these ionic solution^.'*^^^ Mulliken'O suggested a reaction sequence for the ionization of PyI, as follows:
+
PyI, = PyI+I(2) PyI+I- = PyI+ + I(3) Equations 2 and 3 show first the formation of an ion pair and then the dissociation into solvent-separated ions. Reid and Mdiken3were unable to elucidate the nature of the ionic complex through a spectrophotometric study of dilute to lo* M. solutions, with I, in the range of The behavior of halogens and halogen complexes in pyridine has recently been reviewed by Nigretto and Jozefowicz? These same investigators performed a voltammetric analysis of pyridine-iodine solutions in a series of pH buffers8 with I, concentrations in the range of M. They deduced apparent disproportionation constants, K , for the following equilibria: I, = I+ IK = 1.8 X M (4) 21, = I+ I,K = 1.4 X M (5) I,-
+ + = I, + I-
K = 1.25 X M (6) Their disproportionation constant for eq 4 is in good agreement with a value calculated from voltammetric analysis of iodine solutions in acetonitrile in the presence of pyridine7but is in poor agreement with a value obtained (1) L. F. Audrieth and E. J. Birr, J. Am. Chem. SOC.,55,668(1933). (2)K.Hartley and H. A. Skinner, Trans.Faraday SOC.,26,621(1950). (3)C. Reid and R. S. Mulliken, J. Am. Chem. SOC.,76, 3869 (1954). (4)W. J. McKinney and A. I. Popov, J. Am. Chem. SOC.,91,5215 (1969). (5)G.Kortum and H. Wilski, 2.Phys. Chem., 202, 35 (1953). (6)R.Zinaaro, C. A. Vander Werf, and J. Kleinberg, J. Am. Chem. SOC.,73,88 0951). (7)G.Pezzatini and R. Guidelli, Electrochim. Acta, 16,1415 (1971). (8)J. M. Nimetto and M. Jozefowicz. Electrochim. Acta.. 19, . 809 (1974). (9)J. M. Nigretto and M. Jozefowicz, in 'The Chemistry of Nonaqueous Solvents", J. J. Lagowaki, Ed., Academic Press, New York, 1.78, Chapter 5. (10)R.S. Mulliken, J. Am. Chem. SOC.,74,811 (1952). 0022-365418212066-1035$01.25/0
from conductivity meas~rements.~ With the exception of the above-mentioned studies, little quantitative information is available on ionic equilibria in pyridine-iodine solutions. The reliability of data obtained in these solutions has been questioned because slow, spontaneous changes in solution properties have been observed especially in very dilute solution^.^^^^^ In the present investigation, we have employed electrochemical concentration cells to obtain information on ionic equilibria in pyridine-iodine solutions with high iodine concentrations. A mechanism for the ionization is suggested and the experimental results are discussed. Experimental Section Chemicals. Pyridine (Fisher ACS grade) was dried and distilled over CaH,. A middle fraction was retained. Iodine (Fisher ACS grade) was resublimed twice. Potassium iodide (Fisher ACS grade) was dried and stored in a desiccator. Picric acid (Fisher ACS grade) and tetraethylammonium bromide (Fisher Reagent grade) were used as received. Tetraethylammonium picrate, (C2H5),NPi,was formed by the metathesis of equimolar quantities of picric acid, HPi, and tetraethylammonium bromide, (C2H5)4NBr.A solution of (C2H5)4NBrin water was added to a solution of HPi in hot methanol. The orange precipitate, (C2H5)4NPi,was recrystallized twice from 95% ethanol. Electrochemical and Conductivity Measurements. The electrochemical measurements were made either in pyrex H-tubes in which the two compartments were separated by a fine-frit pyrex disk or in three-compartment cells with separation effected by two fine-frit disks. In the latter type of cell, a pyridine solution of (C2H5),NPiwas placed in the middle compartment. All compartments in both types of cell were filled to the same level, the top of the fine-frit disk, to eliminate intermixing due to hydrostatic pressure. Pieces of platinum foil attached to platinum wire leads were inserted into the end compartments and served as electrodes. Voltage measurements were made with a Keithley Model 160B digital multimeter. Conductivity measurements were made with a standard conductivity cell with platinum electrodes and a YSI Model 31 conductivity bridge. The cell constant was determined to be 0.305 cm-' by using a 0.1OOO M solution of NaCl at 25.0 "C. All electrochemical and conductivity measurements were made in a constant temperature water bath regulated to 25.0 f 0.1 "C. Results and Discussion In our treatment of data from electrochemical cells of the type PtlIz(cl), KI in Pyll12(cl)in PylPt (7)
where c1 represents the nominal concentration" of ele0 1982 American Chemical Society
1036
The Journal of Physical Chemistry, Vol. 86, No. 6, 1982
Aronson et al.
mental iodine in pyridine, we will use the notation of Nigretto and Jo~efowicz.~*~ We will obtain information on the equilibrium reaction in pyridine 21, = I+
+ 1,-
(5)
We will assume that the following reaction goes to completion: KI
+ 1 2 = K+ + 1,-
TABLE I: Emf Data on the Cell PtiI, (0.315 M), KI ( C ) I 11, (0.315 M ) i R U e, M 0.00573 0.0109 0.0157 0.0201 0.0241
(8)
This assumption will later be shown to be reasonable. The chemical symbols in reactions 5 and 8 do not necessarily represent actual molecular species. In particular, I2on the left in reactions 5 and 8 is most likely entirely in the form of the molecular complex Py12 as discussed above. The positive ion, I+, may be in the form Py21+or PyI+ with the former more likely in pure pyridine.' An equilibrium expression for reaction 5 is K = a1+a1,-/(a1J2
a
E,mV 6.11 11.12 15.10 18.12 20.36
y,M 0.009 38 0.007 89 0.00698 0.00646 0.006 18
K',MZ 0.000 142 0.000 148 0.000158 0.000 172 0.000 187
x, M
0.0119 0.0122 0.0126 0.0131 0.0137 0.0127 t 0.0006 Oxidation occurs on the side of the cell containing KI.
w
, , , /
20.
(9)
If the concentration of I2 is kept approximately constant, then K' = K ( u ~=) aI+a13~
(10)
4l i I
Virtual electrochemical half-cell reactions for cells of the type described by line-diagram 7 can be written as I2 (1) - 2e = 21+ (1)
(11)
21+ (2) + 2e = Iz (2)
(12)
where half-cell (1) has KI and half-cell (2) contains no KI. The concentration of I2 is assumed to be the same in both half-cells. The emf in volts of such a cell at room temperature can be written as
We designate the equilibrium activity of I+ by the symbol x when no KI is present and by the symbol y when KI is present. We designate as c the nominal concentration of added KI and assume that concentration can be substituted for activity in the equilibrium expression. Assuming the dissociation in reaction 5 to be small, we can write K ' = yCy
+ c) = x 2
(14)
We can solve for x , y, and K'using eq 13 and 14. When KI is present in the half-cell on the right in nominal concentration c' as well as in the half-cell on the left, the equation K ' = yCy
+ c) = y'Cy'+
c?
(15)
can be combined with eq 13 to determine y, y', K', and x . Measurements were made at 25 "C on electrochemical cells of the type described by line-diagram 7 with the same nominal concentration of I2 in both half-cells and with KI present in either one half-cell or in both half-cells in different concentrations. Measurements made by using the two-compartment H-cells or the three-compartment cells with 0.3 M (C2H5),NPi in the middle compartment gave essentially the same results. This indicates that the magnitude of the liquid junction potential is not a serious (11) The term "nominal Concentration"used in connection with either I2 or KI is the number of moles of I, or KI added to a liter of pyridine solution. The dissociation and interactions of I, and KI in solution are not considered.
,
I
I
21 5
1
I
30
60
90
I20
I50
180
TIME (min )
Flgure 1. The upper curve gives the emf data on the cell RII, (0.315 M), K I (c)ll12(0.315 M)(R. The dashed lines refer to additions of K I (seeTable I). The lower curve is the electrical resistance of a 0.946 M solution of iodine in pyridine.
problem.12 The data presented in this paper were all obtained with three-compartment cells. Data obtained on a typical cell are presented in Table I and Figure 1. The nominal iodine concentration in both half-cells was 0.315 M. One half-cell contained a nominal 0.00573 M KI concentration. The emf was measured intermittently over a period of 1 h and was found to have a steady value of 6.11 f 0.05 mV. The nominal KI concentration was then increased three additional times and emf measurements were made. The calculated values of K' and x , although showing an upward trend, can be considered constant within experimental error. Data obtained on a cell with the same nominal 1, concentration containing KI in both half-cells is presented in Table 11. The values of K'and x , although showing a downward trend, can again be considered to be constant within experimental error. Use of a computer program to improve the internal agreement of the data by accounting for changes in I2 concentration due to dissociation, i.e., permitting I2to vary in eq 10, did not improve the internal consistency. The average value of x in Table I, 0.0127 M, is in good agreement with the average value in Table 11, 0.0130 M. Data were obtained on cells with nominal I2 concentrations ranging from 0.158 to 1.576 M. The nominal KI concentrations used in an experiment were between 2 and (12) D. A. MacInnes, "The Principles of Electrochemistry",Dover, New York, 1961, Chapter XIII.
The Journal of Physical Chemistry, Vol. 86, No. 6, 1982
Ionic Equilibria in Pyridine-Iodine Solutions
1037
TABLE 11: Emf Data on the Cell PtlI, (0.315M),KI (c)I 11, (0.315 M),KI (c’)IPta ~~
c, M
0.0157 0.0201 0.0241 0.0278 0.0312
c’,
M
0.0105 0.0100 0.00963 0.00926 0.00892
E , mV 4.21 8.62 12.03 15.48 18.88
Y ,M
K’, M Z 0.000204 0.000 178 0.000172 0.000158 0.000139
0.008 44 0.006 65 0.00575 0.004 85 0.00395
a Oxidation occurs on the side of the cell containing a higher
0.158 0.158 0.315 0.315 0.315O 0.630 0.630 0.946 0.946 0.946a 1.260 1.260 1.576 1.576
1,
I+ activity, M
ionization, %
0.0066 * 0.0002 0.0079 * 0.0002 0.0133 c 0.0004 0.0127 * 0.0006 0.0130 ir 0.0006 0.0315 * 0.001 0.0286 * 0.002 0.0648 * 0.002 0.0592 * 0.003 0.0579 k 0.002 0.0948 r 0.004 0.1007 i 0.003 0.1452 * 0.005 0.1386 * 0.011
8.35 9.87 8.44 8.06 8.25 10.0 9.07 13.7 12.5 12.2 15.0 16.0 18.4 17.6
M
KI concentration.
TABLE 111: Summary of Data from Electrochemical Cells
I, nominal concn, M
X,
0.0143 0.0133 0.0131 0.0126 0.0118 0.0130 k 0.0006
4.4
t
KI present in both half-cells.
8% of the nominal I, concentration. The calculated activities of I+ and the percent ionization of I, on the basis of eq 5 are presented in Table III. The ionization increases from 9% at 0.158 M to 18% at 1.576 M. Electrical conductivity measurements were made on iodine-pyridine solutions with the nominal I, concentrations listed in Table 111. Because degradation of 1,pyridine solutions has been reported,lv5p9in some experiments the conductivity was followed over a period of time. The resistance of a 0.946 M solution was followed for 3 h. The data are presented in Figure 1. An initial drop in resistivity and then a long period in which the rate of change of resistivity is very small was observed. The conductivity data in Figure 2 are from measurements taken 1 h after the solution was placed in the conductivity cell. If the percent ionization of I, into I+ and I, in eq 5 were constant and independent of 1, concentration, one might expect a linear relationship between conductivity and concentration. The fact that in Figure 2 the conductivity is observed to increase more rapidly than the concentration corroborates the conclusion drawn from the data in Table I11 that the percent ionization increases with increasing Iz concentration. We made the assumption above that reaction 8 goes essentially to completion. The equilibrium constant for reaction 8 can be estimated from the equilibrium constant for the ionization of KI in pyridine: 2.1 X and the M. equilibrium constant for reaction 6 above? 1.25 X The calculated value, 17, is sufficiently high to validate the above assumption. Values for the equilibrium constant for the ionization of I, corresponding to reaction 5 and eq 9 were calculated from the data in Table 111. The value is fairly constant,
CONC. I z ( M )
Figure 2. Electrical conductivity of pyridine-iodine solutions.
2X from 0.158 M Iz to 0.630 M I,, and then increases to 1 X lo-, M at an I, concentration of 1.576 M I,. The variation in equilibrium constant may be partially accounted for by the fact that concentrations are substituted for activities. It is likely, however, that the ionization of I, is complex and that our proposed mechanism requires modification. The variation in the value of x and K’with KI concentration observed in Tables I and I1 and noted above may be another indication of the complexity of the ionization. It is of interest to compare our equilibrium constant at the lower I, concentrations, 2 X lov3,with that of Nigretto and J o ~ e f o w i c zfor ~ ~reaction ~ 5, 1.4 X There is a difference of two orders of magnitude between the values. The reason for the discrepancy is not known. One consideration, however, is that our I, concentrations are in the range of 0.16-1.6 M whereas those of Nigretto and Jozefowicz were about lo-, M. I t is interesting to note that a discrepancy of the same order of magnitude occurs between the value of the equilibrium constant for the related ionization, reaction 4, obtained by Nigretto and JozefowM, and that obtained from conductivity icz,8p9 1.8 X measurement^,^ 4.6 X lo-* M.
