3410
L. MUKHERJEE, J. KELLY,Mi. BARANETZKY, AND J. SICA
Equilibria in Pyridine. I.
Determination of Absolute pK Values of Several
Uncharged Acids and Investigation of a Few Typical Acid-Salt Mixtures'i2 by L. M. Mukherjee, John J. Kelly, William Baranetzky, and Jerry Sica Chemistry Department, Polytechnic Institute of Brooklyn, Brooklyn, N e w York 11201
(Received February 16, 1968)
Absolute pK values of two indicator acids, 2,4- and 2,5-dinitrophenuI,have been determined spectrophotoyielded the pK value of the metrically. Investigation of mixtures of these indicators with HBr and "03 latter acids. The dissociation constant of HI, in turn, was determined from differentialvapor pressure measurements (at 37'). Correlation of these results with our earlier hydrogen electrode measurements has made it possible to obtain the potential of the Zn(Hg)IZnClz(s)reference electrode. Potentiometric studies of "03LiN03and HC104-LiC104 mixtures suggest a simple common-ion effect, whereas those of HN03-LiC104mixtures indicate a metathetical reaction: "03 LiC104 $ HC104 LiN03.
+
Introduction Although pyridine appears to be an interesting solvent for certain analytical application^,^ systematic studies of equilibria concerning neutralization reactions in this solvent have not yet been attempted. In a preliminary comm~nication,~ the relative dissociation constants of several acids, vix., HC1, HBr, HI, HXO3, and HC104, in pyridine were reported from emf measurements using the cells Zn(Hg) lZnCln(s) ref electrodel 1 H X in pyridinelHz (1 atm), Pt (I) Zn(Hg) /ZnCl,(s) ref electrodel 1 HXI HX2 in pyridinejH2 (1 atm), Pt (11)
+
In this paper, spectrophotometric studies on 2,4dinitrophenol (2,4-DNP) and 2,s-dinitrophenol (2,5DXP) and their mixtures with nonabsorbing acids such as HBr and fIi?rTO3are presented together with those of differential vapor pressure (DVP) measurements of solutions of hydriodic acid. Correlation of these results with our earIier potentiometric measurements4 has yielded the potential of the Zn(Hg) ZnCla(s) reference electrode (us. a normal hydrogen electrode in pyridine) and the absolute dissociation constants of the diff went acids studied. In addition, mixtures containing (i) HiL'O3 and LiYOS, (ii) HC1O4 and LiC104, and also (E) HN03 and LiC104 have been investigated using the cell
I
Zn(Hg) IZnClz(s) ref electrodel I H X + LiX (or LiY) in pyridinelHz (1 atm), Pt
(111)
The difference between the emf of cell I containing pure H X solution and that of cell I11 containing a mixture of H X and LiX is found to agree with a simple commonion effect; the pK values of LiN03 and LiC104 relative to the corresponding acids have been obtained from The Journal of Physical Chemistry
+
these results. On the other hand, the data obtained from cell I11 on HNO3-LiC104 mixtures indicate a proLiC104 ;zt LiN03 ton-transfer equilibrium: HN03 HC104. The equilibrium constant of this reaction has been evaluated and the acidities of the HNOaLiC104 mixtures have been explained quantitatively.
+
+
Theory Spectrophotometric Measurements. Solutions of Indicator Acids and Indicator-Nonabsorbing Acid Mixtures. An indicator acid, HIn, dissociates in pyridine according to HIn e H+In(acid color) (base color)
+
H+ In(base color)
(1)
and the (over-all) dissociation of a nonabsorbing acid H X can be represented as HX
e H + + X-
(2)
In the present case, solutions of 2,4-D1\'P and 2,sD N P and mixtures of 2,4-DNP and HBr and of 2,5DNP and HN03 have been investigated. Treatment of the absorbance data for both the pure indicator solutions and those of mixtures containing the indicator and a nonabsorbing acid have been given previously5(1) Presented at the NATO International Conference on Nonaqueous Solvent Chemistry held in June 1967, at McMaster University, Hamilton, Ont. (2) Taken, in part, from a thesis submitted by J. J. Kelly to the Graduate School of the Polytechnic Institute of Brooklyn, N. Y . ,in partial fulfillment of the requirements for the degree of ,Master of Science (June 1967), and also from the B.S. Senior theses (June 1967) submitted by 'CV. Baranetzky and J. Sica. (3) (a) Carbide and Carbon Chemicals Co., Works Laboratory Manual 31-9A-2(1956); (b) R. H. Cundiff and P. C. Markunas, Anal. Chem., 28, 792 (1956); (c) C. A. Streuli and R. R. Miron, ibid., 30, 1978 (1958); (d) S. 0. Thompson and G. Chesters, ibid., 36, 655 (1964); (e) M. S. Spritzer, J. M. Costa, and P. J. Elving, ibid., 37, 211 (1965); (f) J. E. Hickey, M. S. Spritzer, and P. J. Elving, A n a l . Chim. Acta, 35, 277 (1966). (4) L. 31. Mukherjee and J. J. Kelly, J . Phys. Chem., 71, 2348 (1967).
3411
EQUILIBRIA IN PYRIDINE the notations used here have the same significance as before. Differential Vapor Pressure Measurements of HI Solutions. Considering the activity coefficients of the undissociated forms to be unity and those of the ionic species equal, the equilibrium constant, KHX,for the over-all dissociadion of an acid HX can be expressed as
aH+
=
dl
KHXCHX KLiXCLiX KHXCHX
(6b)
+
Accordingly, the emf of cell I11 containing a mixture C H X molar in HX and C L ~ molar X in LiX is given by EHX,LiX
=
Eref
+ Elj +
(3) where Ern represents the total molarity as determined from the differential vapor pressure measurement for a Comparison of eq 7 with eq 4 obtained for cell I consolution of stoichiometric molarity m,, and f i denotes taining C H X molar H X yields the ionic activity coefficient which has been estimated from the n'larshall-Grunwald equation as bef01-e.~ E H X- E H X , L=~ X AE = Equation 3 has also been used for calculation of pK 0.02956 log (1 KLixCLix) (8) values of HC1 and HBr in ethylenediamine (EDA) from cryoscopic data.5 KHXCHX
+
Potentiometric Measurements.
I.
Standardization
of the Zn(Hg)IZnClz(s) Reference Electrode. For a weakly dissociated acid of concentration CHX, the emf of cell I a t 25" is given by
E H X=
Eref
+ E11 +
+ 0.02956 log CHX (4)
0.02956 log KHX
where Erefdenotes the potential of the reference electrode vs. a normal hydrogen electrode; El, is the liquid junction potential, and KHX represents the over-all dissociation constant of the acid. According to eq 4 the intercept of the plot of E H Xus. log CHXis equal to Eref El, 0.02956 log KHX. If El, is negligibly small, Eref can be calculated provided an acid of known KHX is available for measurements using cell I. (This assumption is immaterial so long as Eli remains constant and practically independent of the junction; in the latter case, i.e., if El, # 0, the value of Ere< as reported in this study will also include the value of ElJ.) I I . Relative pK of HX and LiX. Expressing the over-all dissociation constant of an acid H X using eq 2 and, similarly, that of its lithium salt IdX as
+
+
On rearranging terms in eq 8 one obtains KLiX 'HX _ _ _ -- (10A.3/0.029561) KHX CLix
(9a)
That is ~ K HX PKLix
=
log (10A.3/0.'-"2956 - 1)
+ 1%
2
(9b)
Equation 9b has been used to evaluate the difference in pK between H N 0 3 and LiN03, and that between HC104 and LiC104. III. HX-LiY Mixtures. In a mixture containing H X and LiY a metathetical reaction such as HX
+ LiY e HY + LiX
(10)
involving a proton transfer may conceivably occur. The equilibrium constant, K, of this reaction is related to the over-all dissociation constants of the various species involved in the following manner
(5) and assuming that in a mixture of H X and LiX the activity coefficients of all ionic species are equal and those of the uncharged species are unity, one obtains from the electroneutrality rule
If KHXand K L ~ are X small, the equilibrium concentrations [HX] and [LiX] can be replaced, respectively, by the corresponding total analytical concentrations CHX and C L ~ in X the given mixture. Thus, eq 6a can be rewritten as follows
Thus, in a particular case K can be evaluated from the values of KHX/KLiX and K H Y I K L ~ Y . Using the electroneutrality rule and assuming as before that the ionic activity coefficients are equal and that the activity coefficients of the uncharged species are unity, one obtains for an HX-LiY mixture
( 5 ) L. M. Mukherjee, S. Bruckenstein, and F. A. K. Badawi, J . Phys. Chem., 6 9 , 2537 (1965).
Volume 7.9, Number 10 October 1968
L. MUKHERJEE, J. KELLY,W. BARANETZKY, AND J. SICA
3412
As is evident, calculation of H+ activity in a given mixture of H X and LiY would require the knowledge of K (cj. eq 11)) KHY,and the ratios K H X / K H Yand KLiy/' KHY. In the present case, the behavior of mixtures of HN08 and LiC104 has been investigated using cell 111. The equilibrium constant of the reaction
purity product-the observed melting points were 112-113 and 107-108", respectively. Salts. Reagent grade anhydrous lithium nitrate (Fisher) and lithium perchlorate (Columbia Organic Chemicals Co.) were kept overnight in a 110" oven before use. The sample of lithium nitrate was found to melt at 252" thus indicating the absence of LiN03. H N 0 3 LiC104 e HC104 LiN03 (13) 3Hz0. was calculated from the values of K H S O ~ / K L ~and N O ~ Zinc chloride (Baker and Adamson) was used in the form of ZnClz. 2C5HsN which was prepared according K H C ~ O ~ / K Lwhich ~C~O were , obtained from our emf to standard method. The product was recrystallized measurements described in the preceding section. twice from absolute ethanol and dried in vacuo a t Using this value of the equilibrium constant the equi80"; the melting point of the final product was 206". HC104, LiN03, and librium concentrations of "03, The purity was further checked by determining the LiC104 in a given mixture of "03 and LiC104 were zinc as well as the chloride. The solubility of the salt calculated neglecting the ionic concentrations which (ZnClz.2CsHsN) in pyridine was found to be 0.22 64 are presumably small. at 25". In subsequent calculations of the H + activity of such mixtures using eq 12, the best estimates of K H C ~ O ~ , The potassium salts of 2,4-DNP and 2,5-DNP were prepared by treating an excess of the phenols with K H N O ~ / K Hand C ~K O I~I, C ~ O ~ / KasLestablished ~ C ~ O ~ in the potassium hydroxide. The excess phenol was removed present work were used. by repeated washing with diethyl ether; the potassium salts thus obtained were finally recrystallized from Experimental Section methanol and dried in vacuo at 80". Potentiometric Technique. I . Hydrogen Electrode. Chemicals. Pyridine. Karl Fischer reagent grade Platinum wire lightly coated with platinum black was pyridine (Fisher) was kept over solid KOH (-20 g/kg) used in conjunction with pure and dry hydrogen gas.6 for 2 weeks. The supernatant liquid was then fracFor each experiment, a freshly platinized electrode was tionally distilled over Linde No. 5A Molecular Sieve and used. solid IIOH, at a reflux ratio of 1:20, using a Corad I I . Reference Electrode. Both Hg, HgClz(s),LiCl(s) vacuum-j acketed silvered column. The fraction boiland Zn(Hg), ZnClz(s) electrodes in pyridine proved ing a t 115.3" (cor 760 mm) was collected and stored in equally suitable as reference electrodes. All measurea 1-1. bottle fitted with an automatic buret from which ments reported here were made using the Zn(Hg), it could be dispensed as required, under pressure of dry ZnClz(s) reference electrode. carbon-dioxide-free nitrogen. This minimized the Preparation of Zn(Hg), Z n C l z ( s ) Reference Electrode. chance of contamination with atmospheric water and An excess of pure zinc (Fisher, 5-10 g) was added to carbon dioxide. The specific conductance of this triple-distilled mercury (-90 g) and stirred by shaking and 4.47 X product varied between 2.92 X in a glass-stoppered bottle; the amalgam was dispensed mho/cm a t 25", and the refractive index, n z 5 ~ was , into the reference electrode vessel in a glove box by found to be 1.5065. No water could be detected by filtering through a glass capillary, avoiding exposure to Marl Fischer titration. air. Acids. HC104, HN03, HI, HBr, and HC1 were used Alternatively, an aqueous solution of zinc chloride in the form of their corresponding salts with pyridine. containing a little hydrochloric acid, was electrolyzed The pyridinium salts were prepared by treating a using a mercury-pool cathode in a separatory funnel slight excess of pure pyridine with the appropriate acid until an amalgam of -3% Zn by weight was produced. in an ice bath. Concentrated reagent grade acid was The amalgam was transferred and repeatedly washed used in each case. The solid thus obtained was filtered with distilled water and later acetone, avoiding exposure and subsequently recrystallized twice from a mixture of to the atmosphere during all manipulations, and finally absolute ethanol and anhydrous diethyl ether and dried by passing a stream of pure and dry nitrogen finally dried a t 80" in a vacuum desiccator. The purity through it. of the final product was checked by noting the melting The Zn(Hg), ZnClz(s) electrode was set up by adding point and by determining the halogen content in apa saturated solution of ZnClz.2CsHsNin pyridine to the propriate cases. The observed melting points are given amalgam already placed in the electrode vessel. A in parentheses for the respective pyridinium salts : quantity of solid ZnClz.2CsH5Nwas added to the zinc C5HjS.HC1 (88"); CsH5N.HBr (235"); CjH&.HI (>190" dec); C5H5N.HNO3(117"); and CjHsN* chloride solution to ensure saturation. The reference HC104 (288"). The indicator acid 2,4-DNP was purified by recrys(6) S. Bruckenstein and L, M. Mukherjee, J . Phgs. Chem., 6 6 , 2225 tallizing twice from benzene; 2,5-DXP was a highest (1962).
+
The Journal of Physical Chemistry
+
EQUILIBRIA IN PYR1DIK.E
3413
electrode thus obtained was allowed to equilibrate for a t least 24 hr before use. A JIetrohm Herisau &Iode1E 388 pH meter was used for emf measurements. This instrument can be read directly to k0.2 mV. In view of the over-all experimental uncertainty, however, the emf values reported in this study are considered reliable to within k 2 mV. I n each run, hydrogen was allowed to disperse through the solution for about 5 min. After this initial period, the cell emf was read every 3 or 4 min until two successive readings agreed to within 1 mV. Each measurement took about 15-20 min. All measurements were made in an air bath maintained a t 25 f 0.5". In view of the small vapor pressure of pyridine (-20 mm) at 25", no pressure correction was made to the observed emf values. Spectrophotometmk Technique. Cary spectrophotometers, Models 14 and 15, were used to record the spectra. All measu.rements were made a t 25" using matched silica cells. The light path of the cells was checked periodically by observing the absorbance of standard potassium chromate solution in the presence of potassium hydroxide. Dij'erential VapoT Pressure Measurements. The technique has been described el~ewhere.~A Mechrolab Model 301 A osmometer which was set at 37" has been used, Biphenyl and l13-diphenylguanidine have been used as standard monomers. All manipulations including preparation of working solutions were carried out in a glove box free from water and carbon dioxide.
Results and Discussions The results of absorbance measurements on pure solutions of 2,4-DNP and 2,5-DNP in pyridine are given in Figure 1 in the form of the plot of the equation5 Eapp
= €In-
-
(€&PP
(€In-
-
-
E')2CtS12 6')KHIn
The values of (5' and E I ~ - for the appropriate wavelengths are given in the legend for both systems. It may be pointed out that 2,4-DNP gives two characteristic absorption maxima (at 374 and 432.5 mp) in solutions of the pure indicator, its completely ionized form (i-e., the potassium salt), and the indicator-nonabsorbing acid mixtures. At higher acid concentrations, however, the 374-mp peak appears to undergo a slight shift to a lower wavelength. Present calculations are based on the absorbance data corresponding to 374 and 432.5 mp. The indicator 2,5-DNP exhibits absorption maxima at 360 and 470 mp, but the molar absorptivities of pure solutions at only one of these two wavelengths ( i e . , a t 470 m,K) show concentration dependence in accordance with the above equation. (The data obtained a t 360 m p obey Beer's law.) Calculation of K H I n in this case has been based on the absorbance of solutions a t 470 mp. The values of K H Icalculated ~
I
.
k
4
I
I
1
lO000~
3oool 50?9?\
I ' ,,
\
2,5 - DNP
I
1000
I .03
I %\*
.02
0.0 I
J
- e')ZCtf,Z ' [(eapp - e')2/e~n- - d]Ctfi2 (eapp
€Inet
Figure 1. Plot of eapp us. for pure indicator solutions: 0, 2,4-DNP a t 432.5 m p ; 0 , 2,PDNP at 374 mp; A, 2,5-DNP a t 470 mp. Least-squares lines are shown in the figure; the constants are given in the format: indicator, X (mp), slope (standard deviation), intercept (standard deviation): 2,4DNP, 374, - 1.555 X 104 (&0.0724 X lo4),1.694 X lo4 (=t0.0095 X lo4); 2,4-DNP, 432.5, -2.189 X lo4 (i0.0760 X lo4), 1.930 X lo4 (f0.0114 X lo4); 2,5-DNPJ 470, -1.717 X l o 6 (f0.0913 X l o b ) , 5.031 X l o 3 (rk0.0462 X 108). (The experimental values6 of 61,- and e' are: €I,,- = 1.635 X lo4 (374 mp), 1.73 X lo4 (432.5 m p ) ; e' = 0.69 X lo4 (374 mp), 0.45 X lo4 (432.5 mp) for 2,4-DNP and €I,,- = 5.43 X IOa (470 mp), e' = 0.173 X lo3 (470 mp) for 2,5-DNP.)
from slopes of the plots (Figure 1) are: 6.43 X (k2.994 X at 374 mp, 4.57 X (k1.586 X 10-6) at 432.5 mp for 2,4-DNP and 5.82 X (k3.166 X lo-') at 470 mp for 2,5-DXP. Carey's* estimates of pK for these two indicators in pyridine are: 4.00 and 5.15. However, the values of K H I n as calculated from the slope :intercept ratio of the equation5
-[HX] - C H In
KHIn ___
+------KHX fi2[In-12 K2HIn
CHIn
KHX
following a trial and error treatment of the absorbance data of the HIn-HX mixtures (Tables I and 11) are 5.16 X (k4.991 X a t 374 mp and 4.025 X (55.213 X lo-') at 432.5 mp for 2,4-DNP and 1.704 X 10-6 (*3.595 X lo-*) at 470 mp for 2,5-DNP. The corresponding values of KHIJKHX used for the best fit in these calculations are 1.00 for 2,4-DNP-HBr (7) J. F. Coetzee and R. M. Lok, J . Phys. Chem., 69, 2690 (1966). (8) E. J. Corey, J . Amer. Chem. Soc., 7 5 , 1172 (1953).
Volume 78. Number 10 October 19bV
3414
L. MUKHERJEE, J. KELLY,W. BARANETZKY, AND J. SICA
Table I : Spectrophotometry of Mixtures of 2,bDinitrophenol and Hydrobromic Acid in Pyridine (25') Concn
105
Concn (M) of HX X 106
0.812
3.08
( M )of HIn X
3.94
4.01
3.94
4.01
4.74
3.34
eapp
CHIn
X
10-4
1.33a 1.304b 1.23" 1.14b 1,220 1.14b 1,240 1.14b
[HXI/CHIn f i 1 [In -1%
fic
0,908" 0 .908b 0.8815 0 .884b 0.882" 0 . 884b 0.879a 0 . 883b
3 . 760a 3.779 1.018" 1,020b 1.027" 1.022b 0.707a 0 . 702b
x
10-4
10.605 11.18b 4 , 32a 5.12b 4,525"
Table 111: Results of Differential Vapor Pressure Measurements on Solutions of Hydriodic Acid in Pyridine Concn of HI (ma)
zm
0.0227 0 0328 0.0448 0.06075
KHI~
0.0310 0.0422 0.0605 0.0765
I
x io-4
5.08b 3.375" 4 . 32sb
8.22 x 10-4 6.08 x 10-4 1. .03 x 6.72 x 10-4 Av 7.84 ( h 1 . 4 3 ) .
a At 37"; -log fi = 7.712/;/(1 where y = Zm m,.
-
+
2.303 X 7.712/-)"',
a Refers to 374 my. Refers to 432.5 mp. -Log f i = 8.191\G, where p is calculated according to eq 21 given in ref 5.
Table I1 : Spectrophotometry of Mixtures of 2,5-Dinitrophenol and Nitric Acid in Pyridine a t 470 my (25") Concn (M)of HIn X
Concn (M)of HX X
106
106
6.07 6.07 6.06 6.04
1.53 3.06 6.14 13.75
eapp X
~
9.90 9.205 8.41 6.95
0.940 0.940 0.935 0.924
a -Log fi = 8 . 1 9 1 4 , where eq 21 given in ref 5.
IHxlx
loa
CHIn
fi"
p
2.88 6.30 13.90 38.80
CHIn x 10-6
f12[1n-]z
6.62 7.92 10.10 17.70
is calculated according to
Based on the above results and our earlier hydrogen electrode measurements, absolute pK values of all the different acids studied have been calculated. A summary of these pK values is presented in Table IV. Wherever possible, the pK values reported by other workers have also been incorporated for the purpose of reference only. No attempt has been made in this paper to explain any discrepancy between the present and the previous values. In the last column, the pK values obtained in EDAS for several of the systems studied here are presented for comparison. Table IV: Summary of the pK Values of Acids r-.
Acid
and 0.02 for 2,5-DSP-HN03. As is evident, in the case of 2,4-DNP the pK values obtained from measurements of the pure indicator solutions agree fairly closely with those obtained from the data of 2,4-DNPHBr mixtures. However, the pK values of 2,5-DTiP, as determined by the two different methods, are found to differ by -0.5 pK unit. No explanation can be offered for this discrepancy. Under the circumstances, for both 2,4-DNP and 2,5-DNP, the K H Ivalues ~ calculated from the pure indicator data and those obtained from the HIn-HX mixtures have been weighted inversely as the square of their standard deviations. These weighted-average KHI,, values are 4.16 X and 1.76 X for 2,4-DNP and 2,5-DNP, respectively. Combining these values with the corresponding value of the KHIJKHX ratio mentioned above finally yields 4.38 for P K H Band ~ 4.06 for ~ K H N o ~ . Table I11 presents the results of DVP measurements on solutions of hydriodic acid. Using eq 3 the pK value of H I is calculated (work done a t the University of Alberta, Edmonton) to be 3.11 (at 37"). Because of the general lack of our knowledge of the temperature coefficient of dissociation constants in nonaqueous solvents, this pK value is considered of special interest in relation to the ones obtained a t 25". T h e Journal of Physical Chemistry
HC104 HI
HNOi
HBr HC1 2,4-DNP 2,5-DNP
In pyridine
3.26a,b 3.12' 3.39agb 3 . 23e 3.11a$c 4.065~~ 4 . 30e 4.36ar 4.38atd 5.66a8h 6 . 14e 4 . 38atd 4 .OOOtd 5. 76a*d 5.15'3'
In EDA
3.10 2.97
3.20 3.73 4.01
Reference 4, using cells I and 11. DVP a Present work. measurements (37'). Spectrophotometry. Other workers (cf. ref 8 and 9b-d).
As is evident, the order of dissociation constants in the case of the inorganic acids is HC104 > H I
> HN03 > HBr > HC1
confirming the earlier observations. Except for the relative order of HI and HC104, the order of strength of the other acids in this series is the same as that found
EQUILIBRIA IN PYRIDINE
3415
. Table V : Results on HX-LiX Type of Mixtures Using Cell I11 CHX, M
HX "03
HCIOd
a
LiX
CLiX,
AE, V
M
KHX/KL~X'
6 . 5 7 x 10-3 1.75 x 10-3 1.75 X 8.81 x 10-4 8.81 x
LiN03
1.04 X 10-1 5 . 9 5 x 10-2 8.53 x 7.23 X 1.11 x 10-1
0.014 0.024 0.023 0.030 0.037
1.66 x 10-3 1.66 x 10-3 1.66 x 10-3
LiC10,
6.84 x 10-3 2.66 x 10-2 6.66 x
0.003 0,010 0 I020
8.00 6.23 9.75 8.75 7.51 Av 8 . 0 j i C 0 . 9 6 15.61 13.60 10.66 __ Av 1 3 . 3 iC 1.42
Equation 9a.
in EDA;6 however, the observed ApK values in pyridine (D = 12.3) for the acid pairs HC1-HBr, HBr-HI, HC1-HN03, and HN03-HC104 are substantially larger with the exception of the case of HBr-HS03, for which the ApK value is found to be somewhat smaller than in EDA (D = 12.9). The tendency for an acid to be less dissociated in pyridine (Table IV) than in EDA may be a general phenomenon associated with the relative degree of interaction of the undissociated acid molecules with the solvents concerned. Being a much stronger base, EDA will probably react with the undissociated acid molecules to a larger extent and facilitate the formation of the intermediate "ion pairs" causing, as a result, a larger over-all dissociation as compared to pyridine. The potentials of the Zn(Hg)[ZnC12(s) reference electrode in pyridine based on the pK values of HBr, HN03, and HC104 as established in this work are calculated to be 0.789, 0.787, and 0.789 V, respectively, yielding an average value of 0.788 V us. a normal hydrogen electrode. Results of potentiometric studies of mixtures containing H N 0 3 LiXO3 and HC104 LiC10, using cell I11 are given in Table V. The values of PKLiNOa - PKHiiOa and 'pKLicioc - P K ~ c i oas~ calculated from eq 9b are, respectively, 0.91 and 1.12, indicating that the lithium salts are less dissociated in pyridine compared to the corresponding acids. This trend has also been observed in EDA.68'0 As mentioned earlier, mixtures of HNO, and LiC104 are considered to involve a proton-transfer reaction (cf. eq 13). The equilibrium constant (K) of the reaction as calculated from the pK difference of the lithium saltacid pairs, uix., LiN03-HN03 and LiC1Os-HC1O4 (cf. Table V), is found t o be 0.605, which is interestingly of the same order of magnitude as the values reported by Price and G r u i ~ w a l dfor ~ ~ similar reactions in glacial acetic acid (D == 6.15). Using the above value of 0.605 for K and 5.495 X 0.1585, and 13.3, respectively, for KHC~O,,
+
+
KHNO~IKHCIO,, and KHCI O ~ / K L ~(Tables C I O ~ IV and V), the H + activity of each of the HN03-LiC104 mixtures studied was calculated using eq 12. Substituting these calculated H + activities and the value 0.788 V for the potential of the Zn(Hg) IZnC12(s) reference electrode, emf values of cell I11 containing mixtures of H N 0 3 and LiC104 have been calculated from the relation E = E r e f 0.05916 log a H t
+
(assuming Elj is negligible) at 25". A comparison of the observed and the calculated emf values is presented in Table VI. The excellent agreement found in this case strongly suggests the validity of our postulated proton-transfer reaction. Table VI: Comparison of Calculated and Observed Emf Values of HIL'03-LiClOa Mixtures (Cell 111) Conon ( M ) of "08 x 104
Conon ( M ) of LiClO4
x 108
Obsd
Calcda
9.22 9.22 9.22
1.50 7.41 17.1
0.597 0.5945 0.589
0.593 0,593 0.589
Q
E -, -V
Equation 12.
Aclcnotuledgment. Acknowledgment is made to the donors of the Petroleum Research Fund, administered by the American Chemical Society, for partial support of this research. (9) (a) A. Hantzsch and K. S. Caldwell, 2. Phys. Chem. (Frankfurt am Main), 61, 227 (1908); (b) M. M.Davies, Trans. Faraday Soc., 31, 1561 (1935); (c) D.S. Burgess aud C. A. Kraus, J. Amer. Chem. SOC.,70, 706 (1948); (d) H. Angerstein, Roc2. Chem., 30, 865 (1956). (10) S. Bruckenstein and L. M. Mukherjee, J . Phys. Chem., 64, 1601 (1960). (11) E. Price and E. Grunwald, ibid., 68, 3876 (1964).
Volume 73. Number 10 October 1068
