S. BRUCKENSTEIN AND I. M. KOLTHOFF
10
[CONTRIBUTION FROM
THE SCHOOL O F
CHEMISTRY, UNIVERSITY
Vol. i 8 OF
MINNESOTA]
Acid-Base Equilibria in Glacial Acetic Acid. 11. Spectrophotometric Determination of the Ionization and Dissociation Constants of p,p’-Dimethylaminoazobenzene and Pyridine. The Abnormal Effect of Water on Indicator Bases1 BY S. BRUCKENSTEIN~ AXD I. M. KOLTHOFF RECEIVED JUNE 22, 1955 The following constants were determined spectrophotometrically at 25’ : Kf (ionization constant of p,p’-dimethylaminoazobenzene, ( I ) ) = 0.100; K i (ion-pair dissociation constant of I ) = 5.0 X 10-6, K I (over-all dissociation constant of I ) = 4.6 X lo-’; KH%O = 8.4 X lo-”; KfY = 5.37 (Py = pyridine); KzY = 9.4 X IO-’; Kpy = 7.9 X lo-’. Upon addition of water to a solution of an indicator base I like p,p’-dimethylaminoazobenzenein acetic acid the color changes in favor of that of the acid species, even though water is a base. This abnormal effect of water is attributed to formation of ion quadruplets and triplets: IH-AcH30+Acwhere Q can dissociate into an ion triplet and a simple ion. Other bases instead of I like pyridine and N,K-diethylaniline exhibit a similar behavior with water. The formation constant of Q from IH+Ac- and water was found equal t o 2.0 f 0.1, and its dissociation constant 3.3 (fl) X The formation constant of Q from IH+Ac- and pyridine is reported equal to 1.1, and that from N,N-diethylaniline and water to 2.65 It 0.10.
+
&iF-ZOi(Q),
p,p’-Dimethylaminoazobenzene(DMAAB = I) the value of K fwhich yielded values of [IH+] such is a relatively strong base in glacial acetic acid. that the experimental data resulted in a straight The spectrum of the basic form of DMAAB and its line when plotted according to equation 2a. K i and acid (ionized) forms are easily distinguishable. The KH*O were calculated from the slope and intercept ionization and dissociation constants3K i = ( [ I H f ] of this line. [Ac-])/[IH+Ac-] and Kf = [IH+Ac-]/[I] for reFrom the effect of pyridine on the dissociation of DMAAB, the over-all dissociation constant of actions 1 and l a pyridine, Kp,, ( = ([PyH+] [Ac-l)/( [Pyl [PyH+I + HAC IH+Ac- (ionization) (1) A - 1 ) ) was calculated. The addition of pyridine IH+Ac- )r I H + Ac- (dissociation) (la) to a solution of DMAA4Bcauses a decrease in the were obtained from spectrophotometric measure- concentration of acid colored indicator species iniments of the concentration of acid forms ( 2 [ I H f ] = tially, and ultimately an increase. The decrease is [IH+] [IH+Ac-1) of the indicator base in solu- caused by repression of the dissociation of IH+Actions of different concentrations of DMA1AB in by acetate ions arising from the base, pyridine, acetic acid. I n the calculation of the constants al- while the subsequent increase is caused by formalowance must be made for the effect of trace tion of higher ionic species which are discussed later amounts of water in the solvent. Water is a base in this paper. I n a solution containing only DMAAB and pyriin acetic acid and in the calculations it is necessary to consider the acetate ions which arise from the dine, the rule of electroneutrality reduces to dissociation of water. Equation 2a is obtained [IH+] [PyH+] = [Ac-] (3) from the rule of electroneutrality ([H&+] can be provided the concentratioii of pyridine is low neglected) enough to neglect the formation of ion-triplets. [HzO+] [IH+] = [Ac-I (2) Using Kpy,Kf and Ki equation 3 is easily transand the expressions for the over-all dissociation formed into equation 3a, where CpY = [Py] constant of water [PyH’Ac-1. The method of calculating the dif-
+
+
+
+
+
+
ferent quantities in equation 3a involves the trial where C H ~ O = [HzO] [H30+Ac-]. Equation 2a and error method used in equation 2a. The value has the form of a straight line, y = mx b, where of Kf found to fit the data in equation 2a should y = [I]/[IH+]’, X = CH,O’[I],m = K H ~ o ’ ( K ~also [ - fit the data in the plot of equation 3a. The intercept of equation 3a is the same as that of equaK i ) and b = 1/ ( K : K @ . tion 2a, which should further substantiate the reThe water content of the solvent was found by liability of the experimental results and the titration with Karl Fischer reagent. [ I ] is the dif- method of calculations. Kpy was calculated from ference between the analytical concentration of the slope of equation 3a. DMAAB and the spectrophotometrically deterThe absorption spectrum of the ionized forms of mined quantity 2 [ I H + ] ,while [ I H + ] = Z [ I H + ] - pyridine in acetic acid is different from that of the K:[I]. A trial and error method was used to find non-ionized forms, I n perchloric acid solutions in acetic acid, pyridine is present entirely in the ion(1) From a thesis submitted b y S Bruckenstein t o t h e Graduate School of the University of Minnesota in partial fulfillment of the reized form thus permitting the determination of the quirements for t h e degree of Doctor of Philosophy, 1954 molar absorptivity of this form. By measuring (2) D u Pont Teaching Fellow, 1953-1954 the absorbance of solutions of pyridine in acetic M. Kolthoff and S Rruckensteln, THIS JOUR(3) For notations see I acid the concentration of the ionized form was NAL, 78, 1 (1956).
+
+
ABNORMAL EFFECTOF WATERON INDICATOR BASES
Jan. 5, 1956
found and from this the equilibrium constant (K?) of the reaction Py
+ HACJr PyH+Ac-
(4)
was calculated. The dissociation into ions was negligible a t the concentrations of pyridine used. I n studying the effect of small amounts of water upon the color equilibrium of various indicators in acetic acid Willman4observed that water acted as a base, more perchloric acid being required to produce a given concentration of acid-colored indicator species in the presence of water than in its absence, for crystal violet, malachite green, benzylideneacetophenone and p-naphtholbenzein among others. These compounds require the presence of free acid in order to exhibit the color characteristic of their acid forms. However, an “abnormal water effect” was found for indicators which were already par tially converted to their acid forms in acetic acid, e.g., metanil yellow, p,p’-dimethylaminoazobenzene, tropeolin 00 and thymolbenzein. This “abnormal water effect” was not understood and not reported in the literature. In the present paper evidence is given that the abnormal effect is caused by triple and quadruple ion formation as indicated by equations 5 , j a , and/or 5b. All the ionized forms of the indicator have the same molar absorptivity and
11
known, KL+’=when KS b u t not KB is known, while KfiB can be found when both K i and KB are unknown. Experimental Reagents. p,p’-Dimethy1aminoazobenzene.-An Eastman Kodak product was recrystallized from acetic acid and dried over potassium hydroxide in mcuo at 50’. Potentiometric titration with perchloric acid in acetic acid indicated a purity of 99.9%. N,N-Diethylanilhe.-An Eastman Kodak product was distilled. Titration in acetic acid with perchloric acid and crystal violet as indicator gave an assay of 100.3%. Pyridine.-A reagent grade material was analyzed using the same technique as for diethylaniline; the assay corresponded to 99.9% pyridine. Acetic Acid.-The water content of the acid used was 0.012 M . The notation “acetic acid” in the experimental refers to the acid with this water content. All experiments were carried out a t 25 f 1O . For further details on all other chemicals and instrumental methods see paper 1 3 of this series.
Experimental Results and Discussion Spectra of DMAAB.-DMAAB, which is red in acid solution, takes on an intermediate shade of orange in pure acetic acid, and approaches a yellow color in the presence of 0.01-0.02 M pyridine because of the repression of the dissociation in IH+Ac-. It appears that DMAAB is a strong enough base to react with acetic acid to give an apI H +Ac- + H 8 0‘Ac- F? I H +Ac-H30 +Ac(5) preciable amount of IH+Ac- so that even in the IH+Ac-H30+ Ac- (5a) 1H+Ac-H30’Ac{Ac-IH+Ac- + H 3 0 f (5b) presence of excess pyridine detectable concentrations of acid colored species exist. I n order to calabsorption spectrum. I n the present paper the culate the various constants involved in the ionizaformation constant of the quadrupole and its over- tion of DMAAB it is necessary to know the molar all dissociation constant have been determined absorptivities of the acid and basic forms a t suitable spectrophotometrically. lengths. Curve I in Fig. 1 gives the absorpThe pronounced tendency of hydronium acetate wave tion spectrum acid, curve I1 the spectrum to form higher ionic aggregates is not limited to in- in the presenceinofacetic an excess of perchloric acid and dicator bases, but also occurs with colorless bases, curve I11 the spectrum in the presence of 0.01 M e.g., pyridine and diethylaniline. The formation of in which the concentration of acid species higher ionic aggregates was established by the shift pyridine is a minimum. It is clear that all the acid species of the ultraviolet spectrum of these two bases in of the indicator have the same spectrum with a favor of that of the ionic forms upon addition of wa- maximum absorption a t 317 mp and the basic form ter. The principal reactions are similar to those given for a n indicator base. Formally the formation of the quadrupole by interaction of I with B can be written in the following three forms
+
+ +
IH+AcBH+AcQI.B + B +QI.B IH+Ac+ I + B +QPB
2.
(6) (6a) (6b)
4
where QI.B = IH+Ac-BH+Ac-. Defining K r B + = [QI.B]/([IH+Ac-] [BHfAc-I), K r B = [QI.B]/[IH+AC-]CB and KLB = [QI.B]/C~CB, where CI represents the sum of the concentration of the undissociated forms of I, and CB the same for those of B. K F B +is related to the other constants by equations 6c and 6d.
(mkc). KL+’B+ can be calculated when K f and KR are (4) A. Willman, Ph.D. Thesis, University of Minnesota, 1933
Fig. 1.-Spectrum of DMAAB (CDMAAB)~ = 1.03 X 10” M , 50 mm. cell used: I, in acetic acid containing 0.012 M water; 11, in excess HClO4; 111, in 0.01 M pyridine.
S. BRUCKENSTEIN AND I. M. KOLTHOFF
12
at 406 mp. It will be shown that the shoulder in curve 3 in the 517 mp region is caused by the presence of IH+Ac- in this solution and is not characteristic of the indicator base. This shoulder is not found in pure benzene solutions of DMAAB where ionized species do not exist. From experiments like those in curve I1 of Fig. 1 the molar absorptivity of the acid species was found to be 4.20 X lo4 liters/mole cm. a t 517 mp and 0.12 X lo4 liters/mole cm. a t 406 mp. Similarly, from experiments in the presence of base, the molar absorptivity of the basic form of DMAAB was found to be 2.31 X lo4 liters/mole cm. a t 406 mp, assuming the molar absorptivity of the base to be negligible a t 517 mp. The concentration, [I], of the non-ionized form of the indicator base in these solutions was found by calculating the total concentration of acid colored species from the observed absorbance a t 517 mp and the molar absorptivity of acid colored indicator species, and subtracting this concentration from the analytical concentration of DMAAB. Using these values of [I] and the observed absorbances a t 407 mp, the molar absorptivity of I a t this wave length was calculated. Ionization and Dissociation Constants of DMAAB and the Over-all Dissociation Constants of Water and Pyridine.-The absorbance of DMAAB a t 517 mp was studied as a function of total concentration of indicator in acetic acid containing 0.012 M water. The experimental results are given in Table I. TABLE I Z[IH+] AT VARIOUSCOXCENTRATIOXS OF DMAAB Total concn. of DMAAB, M X 106
2.02 4.03 6.04 8.06
ZIIH+l, x 106
M
Total concn. of DhZAAB, 'W x 106
Z[IH+l, M X 106
12.09 16.1 20.1
2.86 3.55 4.21
0.626 1.20 1.65 2.10
10
Fig. 2.-Determination
q+
20
30 x
lo-2
I
I
I
20
IO
CP,/
11
30
4c
Fig 3-Determination of KP,, ( C D V A 4 B ) t = 2 015 X lo-' M : I, K: = 0 100 and (CP,), assumed equal t o CP,, 11, K: = 0.100 and Cp, corrected for dissociation; 111, K: = 0.095; IV, K: = 0 105
I
4 1
L
15
r
Figure 2 represents a plot of these data according to equation 2a, adopting a value of Kf of 0.100. The slope of 3.38 X lo2 and the intercept of 2.00 X lo6 correspond to K i = 5.0 X and K H ~ = O 8.4 X 10-l1. The value of the over-all dissociation constant of DMAAB is calculated to be 4.6 X '0
Vol. 78
40
,
of KHZO and KA; k': taken t o lie 0.100
The variation of Z[IH+] with pyridine concentration was determined spectrophotometrically in a 2.015 X 10-5 Msolution of DMAAB and the experimental data are given in Table 11. Curve I of Fig.
3 is a plot of these data according to equation 3a using the value of Kf = 0.1000. The straight line obtained has a slope of 3.0 X lo6which corresponds to KP,, = 7.5 X The abscissa, Cpyl[I],was calculated by assuming that CP, is equal to the stoichiometric concentration of pyridine in solution, an assumption which is not quite correct under our experimental conditions. Although pyridine has a very small over-all dissociation constant the concentrations of pyridine used were so small that dissociation into ions cannot be neglected. Curve 11, Fig. 3, represents the straight line which is obtained when correction for this dissociation is made using the values Kf = 0.100, = 5 X and as a first approximation the over-all dissociation of pyridine, Kpy = 7.5 X lo-'. [Ac -1
y'"
('PY)t
KPy + [Ac-]
where (Cp,) is the stoichiometric concentration of pyridine and
+ K:sd a1
___-
jAc-1 = .\/KP,CPY
The slope of the line is 3.13 X lo6and the intercept is -0.2 X 10'. The scale of this figure is such that the intercept cannot be determined precisely, but the observed value agrees with the value pre-
ABNORMAL EFFECT OF WATERON INDICATOR BASES
Jan. 5, 1956
dicted by Ki = 5 X KP,. is calculated to be 7.9 X lo-' from the slope of curve 11. Using this value of Kp,. to calculate Cpy further approximations did not alter the above figures.
13
molar, indicating that no ionic aggregates exist which contain more than one IH+ unit.
TABLE I1 EFFECT OF PYRIDINE ON DMAAB (CI)+. = 20.15 X 10-6Af Pyridine (Cpy)t,
A617
2.5 X 0.953 5.1 .674 8.0 .618 16.1 .557 32.1 .518 .482 64.2 1 . 2 8 X l o w 3 .454 2.57 .440 5.13 ,430 10.25 ,423 20.50 .422 41 .O .423 79.7 ,437 1 . 5 8 X 10-' .466 ,527 3.13 6.10 ,562
Z[IH+] x lo6
[IH+l
3.44 3.08 2.82 2.54 2.37 2.20 2.07 2.00 1.96 1.93 1.93 1.94 2.00 2.16 2.47 2.70
0.106 ,0802
[I]
[Pyridine-IH I
+-
aggregates]
[IH +Ac-1
.0629
.0443 .0334 .0223 ,0142 ,0099 .0072 .0060 .0060
- 1.0
0.060 .060 .066 ,099 .132 ,395 ,544
Curves I11 and IV of Fig. 3 were calculated using
K fequal to 0.095 and 0.105 and illustrate that the method used is quite sensitive to small variations in the assumed value of K:, deviations from linearity being observed with slight changes in the value of Kf. It is not necessary to correct any of these calculations for acetate ion resulting from the dissociation of water since even a t the lowest concentration of pyridine the correction for water content is only about 2%. Ion-Triplet and Quadruple Formation. DMAaddition of pyridine to AB and Pyridine.-The DMAAB in acetic acid results in a decrease in Z [IH+] initially, but a t concentrations of pyridine larger than 0.02 M, Z [IH+] increases because of the formation of ion-triplets and quadruple ions. The last column in Table I1 presents the ratio [Pyridine-IH+ Aggregates]/ [IH+Ac-] which was calculated from the equation [Pyridine-IH+ Aggregates] = Z[IH+] - [IHfAc-1, where [IH+Ac-] = Kf[I]. Figure 4 is a plot of log [Pyridine-IH+ Aggregates]/ [IH+Ac-] against log Cp,. If the solution had contained only quadruple ions a slope of one should have been found, while a slope of 0.500 should have been found if only triple ions were present. I n these solutions the dissociation of quadruple ions into ion-triplets is suppressed by acetate and pyridinium ions furnished by the dissociation of pyridine. The slope of the line obtained is slightly less than one, 0.96, indicating that +Ac-. the primary species present in solution is PyH Ac-IH+ The value of KkA'py(eq. 6a) is 1.1 in 0.1 M pyridine. I n a separate experiment in 0.8 M pyridine i t was found that the ratio [Pyridine-IH+ Aggregates]/ [IH+Ac-] was independent of the concentration of DMAA4Bin the concentration range 1-5 X
Fig. 4.-Formation
-
C 0
0.5 Log c p y .
of pyridine-IH+ aggregates: = 2.015 X M.
(CDMAAB)t
DMAAB and Water.-When water is added to a solution of DMAAB in acetic acid a regular increase in the red acid color occurs accompanied by a corresponding decrease in the yellow base color. The effect of water on a 5.64 X M solution of DMAAB is shown in Table 111. An isosbestic TABLE 111 QUADRUPOLE DISSOCIATION COXSTANT OF IH +Ac-HaO+Ac(CI), = 5.64 X
M. KdQ =
[QI.wI
where [TI =
[IHfAc-H30+] and [ Q I . ~ ]= [IH+Ac-H30+Ac-]
0.0120 .0307 ,0486 .OS47 ,158
,303 ,591
0.497 0 . 3 0 7 ,485 .324 ,480 .342 .370 ,475 ,447 .404 ,487 ,415 .648 ,335
4.23 4.12 4.07 4.02 3.77 3.47 2.74
1.40 1.48 1.56 1.69 1.84 2.22 2.96
1.60 2.21 2.65 3.34 4.37 5.95 8.32
(-0.35) 0.18 .34 .61 .91 1.37 2.22
0.010
..
.025 .040
1.6 2.3 .068 2 . 9 .119 3 . 3 ,210 3 . 9 .323 5 . 7
point is obtained a t 455 mp which is identical to the isosbestic point obtained when perchloric acid is added to DMAAB in acetic acid, proving that the spectrum of the species produced by the reaction of DMAAB and water is the same as that of 1"Clod-. The total concentration of DMAAB, (C&, can be calculated by solving the set of simultaneous equations 7, 7a and 7b for ( C I ) t . I n these equations
+
( C I ) ~= [I] ZlIH+l (7) A W = ~ s ' ~ ; . bf [ I ]a : ~ + 6 Z [ I H + ]= a:$,+bL:[IH+] (7a) A m = ~:o,b[I] a:E+6Z[IHf] (7b)
+
subscripts refer to wave lengths in mp and superscripts to chemical species, while A represents the measured absorbance and b the cell length, 5.00 cm. Table 111 contains values for [I]and Z [IH+] calculated from equations 7a and lib, using the molar absorptivities reported earlier in this section. The calculated average value of (CI)~,5.65 X is in excellent agreement with the amount taken,
S. BRUCKENSTEIN AND I. M. KOLTHOFF
14
5.64 X This justifies the earlier assumption that the shoulder on the DMAAB spectrum in the presence of small amounts of a colorless base (curve 3, Fig. 1) is caused by the presence of IH+Ac- and is not part of the spectrum of I. It is possible to determine the quadrupole formation constant, KIQ+.W, if sufficient pyridine is added to acetic acid to repress the dissociation of both IHfAc- and IH+Ac-H30+Ac-. Table IV contains spectrophotometrically determined values of [I] and Z[IH+] as a function of the concentration of water in solutions with an initial analytical concentration of 3.63 X 10-6molar DMAAB and 0.153 M pyridine. The concentration of water-IH+ quadrupoles, [QI.w],is given by [ Q I . ~= ] Z[IH+l - ([IH+Ac-] [Pyridine-IH + Aggregates]), where [IH+Ac-] = Kf [I]. The concentration of pyridine-IH + aggregates was calculated, assuming that the ratio [Pyridine-IH+ Aggregates]/'[IH+Ac-] is constant and equal to 0.28. This value, 0.28, is obtained by assuming that [ Q I . ~is] negligibly small in the solution which is 0.01 M' in water. We thus find that [ Q I . ~= ] Z[IH+] - 0.128 [I], since Kr = 0.100. The average value of KL*'W, calculated from the data in Table IV, is 2.0 f 0.1. TABLE IV THEQUADRUPOLE FORMATIONCONSTAST OF DMAAB ASD
Vol. 78
all acid colored species present. Starting with the spectrophotometrically determined values of 8[ I H + ] and [I], and the expression Z[IH+] = [IH+] [IH+Ac-] [IHfAc-H30+] [ Q I . ~ ]where , [IHfAc-] = Kf[I] and [ Q I . ~ ]= K r w [ [ I H + A c - ] C ~ , ~the , sum ( [ I H + ] [IH+Ac-H30+]) can be found. Neglecting [HzAcf], the electro-neutrality rule, [IH+] [IH+Ac-H30+] = LAC-] - [H30+] = LAC-] - CH~O [KH~O/(KH*O [Ac-]) 1, yields the acetate ion concentration. The equilibrium concentration of the indicator ion, [ I H + ] = K:[IH+Ac-]/[Ac-] is subtracted from the sum ( [ I H + ] [IH+Ac-H30+]) to give the concentration of ion triplets. Example.-In 0.591 A4 water solution [I] = 2.74 X and Z [ I H + ] = 2.96 X M. [IH+Ac-] = 0.100 (2.74 X = 0.274 X and [ Q I . ~= ] 2.0 (0.274 X (0.591) = 0.32 X so that [ I H f ] [IH+Ac-H30+] = 2.37 X = [Ac-] - 0.591 [8.4 X l0-l1/([Ac-] S.4 X 1O-l1j], and [Ac-] = 8.32 X [IH+] = 5 X (0.274 X 10+)/(8.32 X = 0.15 X 1OW6and[TH+Ac-H30+] = 2.37 X l o p 6 0.15 X = 2.22 X The average value of K2 ([IH+Ac-H30+][Ac-]/[Q~.w]jis 3.3 X 10-j. Kf has a distinct trend to increase as CH,O increases. This approach neglects the presence of the speWATER cies Ac-IH+Ac-, and is justified, since the prod(CI)t 3.63 X l o w 6M; (Cp,)t = 0.153 llf uct [IHfA4c-] [Ac-] in the above experiment is [I1 Z[IH'I [QI.w~ X 108 X 108 KQ always smaller than the same product for the exAs17 X 106 Water, M Am periments which are presented in Table 11. I n 0.01 0.380 0.086 3 . 2 2 0.41 0.00 the latter experiment, KpJ was obtained and no evi1.06 .69 1 . 9 .302 .211 2.54 1.43 dence for the species Ac-IH+Ac- was found. .95 1 . 9 1.26 2.11 ,271 .264 2 . 3 8 I n order to test the reliability of the constants of 1.49 1.22 2 . 1 2.78 ,244 ,312 2.05 the various equilibria involving DMAAB and wa1.62 1.38 2 . 1 3.43 ,219 ,340 1 . 8 5 ter, the spectrum of DMAAB was determined a t The data in Table I11 can be interpreted, mak- a high concentration of the indicator, 8.52 X ing use of the known values of the'constants K f , M , in water, using a cell with a path length of apK:, K H ~and O KI +.w to calculate the concentration proximately 0.07 cm. The ratio Ab17'A466 was of acetate ion and the individual concentrations of found to be 0.243. From the constants found above the equilibrium concentrations of the various species are calculated to be: [ I ] = 7.52 X [ I H + ] = 1.94 X lop5, [IH+Ac-] = 7.32 X lov5, [IH+Ac-H30+] = 0.39 X lop5, and [QI.w] = 0.21 X 10-5 or Z [ I H + ] = 1.01 X Using u' these values and equations 7a and 'ib the calculated 0 ratio Aj17lA406 is 0.242 in excellent agreement with a 21.50 the experimental value of 0.243. Colorless Bases and Water.-Diethylaniline solutions in acetic acid are colorless. Curve I of Fig. 5 represents the ultraviolet spectrum of the free base in acetic acid and curve I1 that in the presence of excess perchloric acid. The molar absorptivity of diethylaniline a t 306 mp was found to be 21.1 liters/mole cm., while that of the perchlorate was negligible. The absorbance a t this wave OS0 length must be due to the non-ionized form of the base. 280 300 320340 The addition of water to diethylaniline solutions qualitatively produces the same effect as the addi(mA. tion of perchloric acid a t all wave lengths investiFig. 5.-Spectrum of S , ~ - d i e t h ~ l a n i l i ~ i I, e : 0.0531 AM gated, illustrating that the abnormal effect of water is not restricted to colored bases. The absorbances S N-diethylaniline in acetic acid; 11, 0.17 Id diethylaniline measured a t 306 mp upon addition of water to a perchlorate; 20 mm. cell used.
+
+
+
+
+
+
+
+
t
\
I
LhZLLJ
+
+
ASSOCIATION OF SULFATE Iox
Jan. 5, 1956
0.0531 M solution of diethylaniline are presented in Table V and the average value of the quadrupole formation constant between diethylaniline and water, KEeW = [QD.w]/(CDCH,O),was found to be 2.65 f 0.10, (D = diethylaniline.)
WITH
TRIPOSITIVE COBALTAMMINE IONS
15
identical a t these two wave lengths. These results show that more than 80% of the pyridine is present in the ion-pair form. Adopting the mean value of KFy of 5.37 and Km = 7.9 X Kp = 9.4 x 10-7.
TABLE V
EFFECT OF WATERUPONDIETHYLANILINE (D) (C D ) ~ = 0.0531 M 2.00 cm. quartz cells Water, hl
0.274 0.551 1.09 2.08
Asan
1.39 0.960 0.580 0.380
KD.W.
[QD.w.]/CD
Q
0.705 1.47 3.09 5.24 Av. =
2.58 2.67 2.85 2.62 2.65 f 0.10
Pyridine solutions in acetic acid are also colorless and the spectrum of pyridine in acetic acid (curve I) and in excess perchloric acid (curve 2) are given in Fig. 6. The molar absorptivities in pure acetic acid were found to be 4.62 X lo3 a t 255.5 m p and 3.12 X lo3 at 261.0 mp, while in 0.1 fM perchloric acid these values a t the same wave lengths are 5.48 X lo3and 3.74 X lo3. The ratio of the molar absorptivities of pyridine in acetic acid a t 255.5 mp and 261.0 mp is 1.49, while this ratio is 1.47 in 0.1 M perchloric acid This suggests that the two spectra are identical and are caused by similar species. It is reasonable to assume that the base (Py) does not absorb a t either of these wave lengths and the observed absorption is caused by PyH+Acin acetic acid and by PyH+C104- in the presence of excess perchloric acid. Experiments using 1.01 and 2.03 X M pyridine in the presence and absence of excess perchloric acid yielded an average KPY of 5.52 f 0.04 at 255.5 mp and 5.21 0.09 a t 261.0 mp, assuming that the molar absorptivities of PyH fAc- and PyH +Clod- are
*
[CONTRIBUTION FROM
THE
Fig. 6.-Spectrum of pyridine: (CP,)~ = 2.03 X IO-' molar: I, in acetic acid; 11, in 0.05 M HClO4; 20 mm. cell used.
It was noted that when water was added to pyridine the effect on the spectrum was similar to that produced by perchloric acid, but no quantitative measurements were made to determine the species formed. I t may be stated that the "abnormal water effect'' is a general phenomenon and i t is not restricted to indicator bases. MINNEAPOLIS, MINN.
GEORGE HERBERT JONES LABORATORIES O F THE UNIVERSITY O F CHICAGO]
The Outer Sphere Association of Sulfate Ion with Tripositive Cobaltammine Ions BY FRANZ A. POSEY AND HENRYTAUBE RECEIVED SEPTEMBER 12, 1955 The equilibrium quotients, K, for the outer sphere association of cO(lW3)6++' and Co( KHa)jH20f++with SO,- have been measured a t several temperatures and ionic strengths by following the changes in optical density in the ultraviolet region of the absorption spectra of the cations as a function of sulfate ion concentration. Values of AFO, AH0 and A S 0 a t 25" are -4.53 kcal. mole-', 0.40 kcal. mole-', and 16.6 e.u. for the association of Co(NH3),+++ with SO4-. The equilibrium quotients calculated are independent of concentrations, of the wave length of light used, and vary with ionic strength in conformity with a Debye-Huckel equation. The quotient for the equilibrium ratio, (CO(NH~)~H~O+++.SO~-)/(CO(KH~)~SO~+), in 1 M NaCIOl a t 25' has been measured as 0.90 by following changes in the substitution equilibrium using light of wave length 560 mb. These experiments have also yielded a value for K in agreement with that measured using the instantaneous changes in the ultraviolet extinctions as SO1' is added to solutions of CO(NH,)~H~O+++.
Linhard' has demonstrated that the presence in solution of certain anions can cause marked changes in the absorption spectra of tripositive Co(II1) and Cr(II1) complex ions in the wave length region of the strong ultraviolet band. I n an earlier paper2 we reported that this effect occurs when SO4= is added to a solution containing Co(NH3),H20+++. The changes in question occur immediately on mix(1) M.Linhard, Z. Elekfrochem., 50, 224 (1944). (2) H.Taube and F. A. Posey, THIS J O U R N A L , 75, 1463 (1953).
ing the reagents, in contrast to changes in the two bands in the wave length region of visible light, which for the ions in question take place quite slowly at room temperature, concomitant with substitution in the inner sphere of coordination of the central metallic ion. Therefore, if the changes of spectrum in the ultraviolet region are caused by the association of the cation with the anions, this interaction must be of such nature that the anion does not become equivalent to the six ligands in the
