THEBASICIONIZATION CONSTANT OF GLYCINE IN DIOXANE-WATER
June, 1943
[CONTRIBUTION FROM
THE
1117
DEPARTMENT OF CHEMISTRY OF YALEUNIVERSITY~
The Basic Ionization Constant of Glycine in Dioxane-Water Solutions BY HERBERT S. HARNED AND CLAIRM. BIRDSALL'
The ionization of weak electrolytes in aqueous solution and in water-organic solvent mixtures has been studied successfully by the use of cells without liquid junction.2 The present determination of the basic ionization constant of glycine in dioxane-water solutions supplements the work recently reported on the acidic ionization constant of this substance in the same The acidic and basic ionization constants of glycine have been determined in aqueous solution by Owen4 from cells without liquid junction. For the evaluation of the basic ionization constant cells of the type HZI Z(m,), NaZOH(m?), NaCl(ms), X q / D , YXH?O 1 Ag-AgC1
have been employed where Z designates the zwitterion, +NH&H2COO-, NaZOH the hydrated sodium glycinate, and X the per cent. by weight of dioxane in the solvent. By thermodynamic treatment of the ionization data the free energy, heat content, heat capacity and entropy of the ionization reaction can be evaluated. Outline of the Method-Extrapolation Function.-The ionization may be represented by 0 H Z Z - Z +OH
and the thermodynamic ionization constant for the reaction expressed as - UZ~OH B - - - = aOH2
mzmoHYzYou *lOHZYOHZ
(1)
where m represents molality, y the activity coefficient and 2 , OH and OH2 designate the ionic species. The apparent ionization constant KL may be expressed by
The thermodynamic equation for the cell is E --- Eo* k
- - log a H a C l
(3)
where k equals 2.3026 RT/F.6 (1) This contribution contains material from a dissertation presented by Clair M. Birdsall t o the Graduate School of Yale University in partial fulfillment of the requirements for the degree of Doctor of Philosophy, June, 1942. (2) Harned and Owen, Chem. Rev., 26, 31 (1939). (3) Harned and Birdsall, THISJOURNAL, 66, 54 (1943). (4) Owen, i b i d . , 66, 24 (1934). (5) The values of R,F and the ice point used in calculating k were 1 087 cal./deg.. 23059 cal. and 273.1°A., respectively.
By combining equations (1) and (3) with that for the thermodynamic ionization constant of water I \ ' ~=
mFffl?OHYHYOH aHIO
the following equation is obtained
log K{
Kw has been determined in 20, 45 and 70% dioxane-water solutions by Harned and Fallon,6 ma = m3 is known, mZ and mOHZ can be obtained by correcting the stoichiometric values ml and m2 for the amount of ionization, and Eo* has been determined by Harned and M o r r i s ~ n . ~ ~ ~ Measurement of E completes the data necessary for the evaluation of the left side of equation ( 5 ) designated as log Kfj. A preliminary value of log Kfj was obtained by neglecting the effect of the ionization on ml and m2. This value was substituted in equation ( 2 ) and a first approximation of calculated. ml and m2 were corrected by this amount and log Kf; again computed. The corrections in all cases were small and the procedure was unnecessary for the data obtained in the 70yo dioxane-water solutions. The extrapolation of log K{ against m yields the thermodynamic ionization constant a t zero ionic strength. Experimental Procedure and Observed Electromotive Forces.-The experimental technique has been described previously in detaiLY The stock solution was prepared by weighing glycine, standard sodium hydroxide, sodium chloride and water into a large flask. The glycine was crystallized three times from conductivity water and dried over solid potassium hydroxide in a vacuum desiccator. It was found to contain O.l6Y0 water. lo Carbonate free sodium hydroxide was prepared by diluting a saturated sodium hydrox(6) Harned and Fallon. THISJOURNAL, 61, 2374 (1!)39). (7) Harned and Morrison, ibid., 68, 1908 (1936) (8) The equations, based on the data of Harned and Morrison, by which the standard potentials were computed are given by Harned, Morrison, et n l . , ibid., 61, 49 (1939). (9) Harned and Morrison, A m . J . Sci., 33, 1ti1 (1937). (10) Harned and Birdsall, THIS JOURNAL, 66, 54 (1943).
HERBERTS. HARNEDAND CLAIR31. RIKDSALL
11 1s
ide solution. The dilute solution was standardized against potassium acid phthalate. Baker analyzed sodium chloride was recrystallized and dried in a muffle a t 300'. The molalities in the buffer alution were: ml = mZ = 0.41914, m2 = mNaZOH = 0.38122, m3 = mNaC1 = 0.320:%. Thus ml/mz = 1.0994, m2/m3 = 1.3082. The cell solutions were prepared by weight dilution of portions of the stock solution with dioxane and water. The cells were prepared in triplicate a t each chosen ionic strength. After equilibrium had been attained a t 25' measurements were made a t five degree intervals from 0 to 50'. The original and final measurements a t 25?"were used in judging satisfactory cell behavior. The experimental electromotive forces can be expressed by the equation Et =
E26
+
it
- 25)
+ b(t - 25)'
(6)
The parameters of this equation obtained by a graphical methodll are given in Table I. The deviations of the calculated from the observed electromotive forces tabulated in column five of Table I show that the equation reproduces the experimental results within * 0.16 mv. This compares favorably with the agreement among the three cells measured. In general the cell behavior was less satisfactory as the dioxane content of the solvent was increased. Ionization Constants.-The values of log Kg computed a t the chosen ionic strengths were plotted on a large scale graph. The extrapolated value a t zero ionic strength is equal to log K B since the second term oil the right side I
5
Vol. (i3
TABLEI ELECTROMOTIVE FORCES OF
THE
CELLS
+ %a(tdioxane - 25) + by
Constants of the equation Et o= E26 h(t - 25)*; valid from 0 to 50 . x' = weight. A = average deviation between tromotive forces and those computed by 821 = 1.0994mz; mi = 1.3082~~8. a X 106
Eza
rill
-b
observed electhe equation,
X 108
A
x = 20 0.006946 .007920 .008587 .010222 .010717 .013638 ,015175 ,019072 ,019930
0.92147 .91847 ,91615 .91203 ,91057 .90472 ,90199 ,89650 .89&38
0.006251 ,006921 .008671 ,010950 ,013349 ,014709 ,019075 ,022095
0.90445 ,90187 ,89627 ,89055 ,88571 .88339 ,87712 ,87358
0.008770 ,013246 ,015137 ,020763
0.85970 ,85055 ,84739 ,84083
200 224 198 228 192 196 218 194 196
0.02 .12 .05 .09
82 90 104 124 146 154 174 180
124 120 135 120 116 128 134 140
0.08
70 386 416 424 450
70 78 62 80
0.16 .13 .06
98 90 81 69 67 44 38 14 10
. mi .04 .14 .02 .03
x = 45
x
.ll .12 .10 .05 .05 .03 .10
=
.09
of equation (5) then equals zero. Figure 1 shows the 25' extrapolations in the three dioxane-water solutions. In Table I1 are tabulated all the values of -log KBor pK, determined from similar plots.
I
TABLEI1 T m BASICIONIZATION CONSTANT OF GLYCINE IN DIOXANE-WATER SOLUTIONS - log K B = ~ K B .X = % dioxane by weight I
+ 0.500 2 0.420
c I
/A I
/
i
M
'
+ t.
0.01
0
0.02
0.03
P.
Pig. 1.-Estrapolations
a t 25: log K"Bus.
g;
PKB
x =20
0 4.9285 5 4.8804 10 4.8317 15 4.7910 20 4.7493 25 4.7129 30 4.6802 35 4.6495 40 4.6186 45 4.5938 50 4.5703
PKB
-0.4 +1.3
5.7475 5.6938 5.6436 5.5950 5.5467 5.5058 5.4674 5.4310 5.3863 5.3623 5.3298
-1.1 $1.1 -0.8 -0.5 +O 7 f1.2 -1.2 +0.1 +0.3
X=70
x=45
A X 10-3
A
x
10-3 P K B
-0.1 $0.3 +1.1 +0.5 -2.7 -1.1 +0.3 +1.3 +1.6 +0.4 -1.4
6.832 6.775 6.719 6.672 6.621 6.575 6.531 6.490 6.443 6.402 6.363
A X 10-r
+2 0 -3 +1 -2 0
+1 +4 -1 -1 0
diameter of
circles = 0.4 my. ( 1 I ) Harried arid Nirn.;, THISJ O U R N A L , 64, 423 (l!l32).
The uncertainty in the numerical value of log KH due to scattering of the experimental points
TIIEBASICIONIZATION CONSTANT OF GLYCINE IN DIOXANE-WATER
June, 1943
1119
TABLEI11 Constants of Equations (7) to (13) Inclusive.
X
=
% Dioxane by Weight
- log Kg 3.9203 20 4.4350 45 5.0433 70 (5.039)t t Since the results in the 70% dioxane solutions are less accurate than the others, these derived values of Te and -log KO are obviously in error. This is due to the difficulty in determining the parameter C', which when C' is small greatly affects the results (see equations (12) and (13)). This uncertainty also invalidates the value of ACii in Table IV. X
A*
O1*
1307.30 1519.8866 1364.9011 895.267
D*
C*
2.5629 3.3800 1.1999 -3.328
0.008038 .010046 ,0071393 ,0008182
5980.95 6953.7851 6244.6955 4096.02
may be *0.0016 in the 20 and 45% solutions and as much as *0.005 in the 70% solutions. Thermodynamic Quantities.-Although plots of PKB versus T do not show as much curvature as similar plots for @KA,the equation of Harned and Robinson12 log K =
- A* - -+ D" - C*T
(7)
is suitable for the expression of log KB as a function of the temperature. This is obvious from Table 11 in which A x = PKB (observed) fiKB The differences are well within the range of the estimated uncertainty in all but a few cases. The thermodynamic functions derived from equation (7) are as follows
AE = A' - D'T
+ C'T2
AH; = A' - C'T2 ACii = -2C'T A X = D' - 2C'T
(8) (9) (10) (11)
in which A' = 2.3026 RA* D' = 2.3026 RD* C' = 2.3026 RC*
The temperature a t which the maximum ionization occurs, T, and the value of the ionization constant a t that temperature, log KO,are calculated by use of
d m
Te = log Ke = D* - 2
4
m
D'
AI
(12) (13)
Table I11 contains the values of these constants obtained by the method of least squares. The thermodynamic quantities were calculated by equations (8) to (11) using the numerical (12) Harned and Robinson, Trans. Favaduy Soc., 36, 973 (1940). (13) Computed by Dr. R. A. Robinson from the data of Owen, Trns JOURNAL, 56, 24 (1934).
C'
11.7254 15.4642 5.4898 - 15.227
TO
0.036673 ,045964 ,032663 ,003743
403.3 389.0 437.2 ( 1046)t
values for the constants given in Table 111. The values of the thermodynamic quantities a t 25' are tabulated in Table IV. Data calculated from the measurements made by Owen13 on the basic ionization of glycine in aqueous solutions have been included in Table I11 and Table 1V. TABLE IV THERMODYNAMIC QUANTIT~ES AT 25". X = BY WEIGHT X 0'8 20 45
A??
AH:.
cal./mole
cal /mole
5753 6428 7511 8968
2713 2869 3342 3763
ACL,> cal./'-mole
-21 9 -27.4 -19.5 ( - 2 2)*
70 * See note (Table 111).
qc DIOXANE
AS:, cal /O-mole
-10.2 -11.9 -14.0 -17.5
Summary
I
Cells of the type Hz 2 (rnl), NaZOH(mz), NaCl(rns), X%D, Y'%H20 Ag-AgCI, have been employed to investigate the basic ionization of glycine in solutions containing 20, 45 and 70% dioxane by weight. The measurements were made at five degree intervals from 0 to 50'. The thermodynamic ionization constants, evaluated by graphical extrapolation of the experimental data, can be expressed over the temperature range by the equation A*
I
logK= --+D*-C*T T
(7)
The numerical values of the constants in this equation, obtained by the method of least squares, have been tabulated. The thermodynamic quantities of the ionization-reaction a t 25' are tabulated. Data based on the work of Owenla have been included. NEWHAVEN,CONNECTICUTRECEIVED MARCH13, 1943
