1396
VOl. 58
SAMUEL KORMAN AND VICTOR K. LA MER
freed of water of hydration by drying at 170" for three hours. The silver salt, obtained by treatment of the sodium salt in aqueous solution with silver nitrate, separates directly without water of crystallization. The reaction between the sodium salt (10 g.) and acetyl chloride (50 ml.) is completed by boiling the mixture for one hour. The sodium chloride is filtered off, the acetyl chloride is evaporated, and the residue (3.4 9.) is crystallized from ether. Similar treatment of the silver salt yields the same compound (mixed melting point) which forms yellow needles that melt a t 77-78". Anal. Calcd. for C14Hlz06: C, 64.6; H, 4.6. Found: C, 64.7; H, 4.8. Ethyl 3-acetoxyindone-2-carboxylateis readily hydrolyzed, even on exposure to atmospheric moisture, to 1,3diketohydrindene. Benzoylation of Ethyl 1,3-Diketohydrindene-2-carboxylate.-The reaction of the silver salt of the diketo ester with benzoyl chloride in benzene is completed by warming for one hour, and the product is crystallized from benzene and ether. In agreement with the description of Hantzsch and G a j e w ~ k ithe , ~ benzoate forms yellow crystals that melt a t 146-148".
[CONTRIBUTION FROM
THE
The sodium salt of the diketo ester does not react with benzoyl chloride in boiling benzene. However, if the sodium salt (54 8.) is stirred with benzoyl chloride (60 9 . ) a t 80" for five hours, a reaction takes place. Separated from the sodium chloride with ether and crystallized from benzene, the organic product forms large yellow prisms that melt a t 140-141 '. Anal. Calcd. for C14HloO4: C!, 6 9 3 ; H, 4.1. Found: C, 69.4; H, 4.4. The substance depresses the melting point of the benzoate. Treated with sodium ethoxide in alcohol it gives a red solution which becomes yellow on acidification. Benzoic acid may be sublimed from the tarry residue obtained by evaporating this solution.
Summary Bromination is suggested as a method for cleaving 2-acyl derivatives of 1,3-diketohydrindene. Some reactions of 2-benzoyl diketohydrindene and of ethyl 1,3-diketohydrindene-2-carboxylateare described. MINNEAPOLIS, MI".
RECEIVED JUNE 8, 1936
DEPARTMENT OF CHEMISTRY, COLUMBIA UNIVERSITY ]
Deuterium Exchange Equilibria in Solution and the Quinhydrone Electrode' BY SAMUEL KORMAN AND VICTOR K. LA MER Introduction For the elucidation of the isotopic influence of deuterium upon various phenomena in heavy water (deuterium oxide), it is important to know with some degree of precision the magnitude of the equilibrium constants of the exchange processes set up between the protons and deutons in the solvent and solute molecules. For example, the interpretation of acid and base catalysis in DzOHzO mixtures requires a knowledge of the distribution of deuterium and protium between the various substrate, catalyst (e. g., H30+ and DaO+) and solvent molecules.2 The measurement of the e. m. f. of galvanic cells involving DzO furnishes a precise means of determining these exchange constants. To exclude the experimental complications and the theoretical uncertainties inherent in galvanic cells with liquid junction potentials, we have restricted our investigation to cells without transference. (I) From a thesis submitted by Samuel Korman in partial fulfilment of the requirements for the degree of Doctor of Philosophy in the Faculty of Pure Science of Columbia University. (21 Hamill and La Mer, -7. Chem. PhyS., 4, 294 (1936); 4, 395 (1938); Urey and Teal, Rev. Modern P h y s . , 7, 34 (1935). For a discussion of the importance of exchange reactions see pp. 52, 5!), 60.
The experimental difficulties attendant upon the use of the deuterium gas electrode with small quantities of solution make an appropriate substitute highly desirable. We have found that the quinhydrone electrode3 establishes equilibrium quickly and is particularly well adapted for the micro technique necessary fur heavy water studies. To this end, the following cells were investigated Pt, Q.QHZ/HCl (0.01 M ) in HZO/AgCl, Ag (11) Pt, (Q.QH2 Q.QDz)/HCl-DCl (0.01 M ) in HZO-DzO/AgCl, Ag (111) Pt, Q.QH2/{HA(md NaA(mJ NaCl(m3) I in HZO/AgCl, Ag
(I)
+
+
+
where Q.QHz and Q.QDz represent light and heavy quinhydrone, and HA and DA are the light and heavy varieties of a weak acid. For (I) and (11) the cell reactions may be written (A) 2HC1 (0.01 M ) (B) 2DC1 (0.01 M)
+ Q(Hz0)2AgCl(s) + 2Ag(~) +=QHdHzO); EH + Q(DzO) + 2Ag(s) 2AgCl(s)
+ QDdDzO); ED =
(3) La Mer and Korman, THISJOURNAL, 67, 1510 (1935).
DEUTERIUM EXCHANGE EQUILIBRIA
Aug., 1936
1397
0.05). We thus have a method for obtaining the ratios of dissociation constants of weak acids in = 2HC1(0.01 M) QDz(D10) 4-Q(H20); (ED- EH) heavy and light water, which have proved to be Measurements of the solubility of quinone in valuable in the interpretation of acid catalyzed light and heavy water given below, permit (C) reactions in heavy water.2 to be rewritten as
Subtracting (A) from (8) (C) 2DC1(0.01 M) f QHz(Hn0) -I- Q(D20)
+
+
QHz(H20) = (D) ZDCl(0.01 M) 2HC1 (0.01 M ) 4-QD*(DzO); (ED - EH)
Then From this fundamental exchange constant, by appropriate combination with values obtained from other sources, a number of important exchange equilibrium constants will be derived. Furthermore, the data on cells (111) and (IV) can be used to obtain the ratio of dissociation constants of weak acids following the method of Harned and co-workers4 who used the hydrogen electrode. To maintain precision using a micro method, we again substitute the quinhydrone electrode for the gas electrode. The cell reaction then becomes (E) 2AgCl(s)
4-QHz
+
2NaA(ncl) = 2Ag(s) 4- 2HA(ml)
+ PNaCl(mJ 4- Q
Experimental
Materials.-Quinhydrone,o constant boiling hydrochloric acid solution,7 heavy waters and silver oxidev were prepared according to directions given previously. Acetic anhydride was purified by fractional distillation, while sodium acetate, sodium salicylate and sodium chloride were recrystallized from conductivity water. Salicylic acid was recrystallized from chloroform. Quinone was resublimed twice. Apparatus.-The micro cell has been described by La Mer and Armbruster.6 Their preliminary measurements established the reliability of the method as applied to the silver chloride electrode in heavy water, where the quantities of solution available for rinsing are severely limited. The vacuum technique for removing oxygen from the final cell solution, for rinsing the silver chloride electrodes. and for filling the cell, follows closely the method given by Harned and Wright.Io The cell shown in Fig. 1 was of 4cc. capacity. Space is provided for duplicate electrodes of each kind, and stopcocks arranged to separate the compartments for rinsing of the silver chloride electrodes.
for which (2)
E
- Eo
=
- ( R T / F ) bfafclmamcl
where Eo is ‘the normal electrode potential of cell (I). Since KHA
fHfAmHmA/fEAmHA
we substitute for mH in equation (2) and rewrite the expression
Fig. 1.-Diagram
P
of cell.
Electrodes.-Reproducible quinhydrone electrodes were prepared by immersing several platinum spiral wires, carefully sealed in separate glass tubes filled with mercury RT In KEA contacts, in 0.01 N hydrochloric acid solution saturated F with quinhydrone, and “short circuiting” them together and a similar expression involving DzO, QDz, for about a half hour.I1 They then agreed to within 0.03 etc., for the deutero acid. The logarithm of the millivolt. The electrodes were kept in cleaning solution12 ratio of the activity coefficient products is zero a t until another pair was required, whereupon they were thoroughly with distilled water. infinite dilution and varies but little from zero rinsed The silver-silver chloride electrodes were the Type 2 a t low ionic strengths. The extrapolation of the described by Harned.13 Several electrodes made simulright side of equation (3) plotted against ionic taneously were washed repeatedly with distilled water, strength is linear a t temperatures over 15’ for immersed in 0.01 N hydrochloric acid solution, and ordinary water.4 Since the dielectric constants of “shorted” together for twenty-four to thirty hours. They HzO and DzO are almost identical,6 the same ex- agreed to within 0.05 millivolt. The electrodes so prepared were protected against light. t r a p o l a t i o n will be v a l i d i n heavy w a t e r . Thus
the values of K H Aand KDAmay differ, but their ratio will certainly remain constant to a very good approximation in the dilute range (- p = (4) Harned and Ehlers, THISJOURNAL, 54, 1350 (1932) ; 66, 2179 (1933). ( 5 ) Lewis, Olson and Maroney, ibid., 55, 4731 (1933); Muller, Phrsik. Z., 36, 1009-11 (1935); P. Abadie and G . Champetier, Comet. rend., 200, 1387 (1936).
(6) La Mer and Armbruster, THISJOURNAL, 67, 1510 (1935). (7) Bonner and Titus, i b i d . , 62, 633 (1930). (8) W. N. Baker and La Mer, J . Chem. P h y s . , 3, 406 (1935). (9) Helferich and Klein, Ann., 450, 225 (1926). (10) Harned and Wright, THISJOURNAL, 66, 4849 (1933). (11) L. E. Baker, Thesis, Columbia University, 1922. (12) Morgan, Lammert and Campbell, THIS J O U R N A L , 5S, 451 (1931). (13) Harned, ibid.. 61, 416 (1929); Guntelberg, Z . g h y s i k . Chem., 123, 199 (1926).
1398
SAMUEL
KORMAN AND VICTOR K. LA MER
Preparation of Cell Solutions.-The limiting factor in the precision of the measurements is the accuracy with which the composition of the cell solution is known. In the preparation of the buffer solutions, the equivalent amounts of sodium chloride and the sodium salt of the weak acid were transferred from weighing bottles to a 50-cc. ground glass-stoppered flask containing the water ~ a m p l e . 1 The ~ final concentration of the cell solution was corrected for the water lost in the evacuation process. Characteristics of the Micro Cell.-Equilibrium was attained within from twenty to forty-five minutes a t 25', and was maintained to within 0.05 millivolt for two to four hours, followed by a slow drift to higher values, due probably to the formation of chlorohydroquinone. There is a definite lag in attainment of equilibrium when the temperalure is raised, while equilibrium is reached almost immediately in proceeding from a higher to a lower temperature.
Results and Precision of the Measurements Hydrochloric Acid Solutions.--E (0.01) in 'I'able I is calculated on the basis of mols/55.3 mols D20-H20. Figure 2 is E (0.01) plotted
precision with the weak acid solutions is about the same, and corresponds to that reported by Harned and Eh1e1-s.~ TABLEI DATAFOR CELLS(I) AND (11) Molality HCI N D ~ O ( M / 5 5 . 3 mols) -&bsd.
0.0
,0515 ,1030" .3454 .5109 .6925 ,9240 .9640
0.2700
( 1.0000)
A.'d., mv.
6
mv.
-Eo
01
0.011182 0.24122 0.01 0.05 0.23548 ,010855 .23963 .03 .08 .23541 .011263 .23960 -03 .07 ,23349 .013103 .24738 .02 .05 ,23350 .01010 ,23605 .a555 .01010 ,23602 .23552 .01 .05 ,24409 ,014100 .26175 .009139 .23957 .02 .06 .24419 ,011913 ,25915 .02 .06 .25016 ,010762 .25398 .02 .07 .25021 .009528 ,25543 .05 .05 .25770 .010470 ,26022 .02 .06 .25766 .011524 .27414 .05 .02 .26685 .011022 ,27190 .04 .04 .26690 .03 .05 .26807 ,011870 .27687 .011878 .27694 .02 .07 ,26812 ( .2700)
The results, for NDXo= 0.1030, were obtained in 1934 by Dr. M. H. Armbruster in a preliminary investigation.
0 2600
0
Vol. 58
0.2
0.4
0.6
0.8
1.0
ND~o. Fig. 2.-Quinhydrone electrode in heavy water.
against NDSO= 9.377 A S - 1.01 (AS)*, where A S is the excess specific gravity of the solvent over ordinary water. The fourth column gives the average deviation from the mean, and the fifth column gives the precision measure for each experiment. This takes into account the deviation of the mean, the weighing error and the difference in electrodes. It will be seen that the average 6 for this set of results does not exceed 0.06 millivolt, while the average deviation does not exceed 0.03 millivolt, a result which is about twice as precise as has been obtained with the hydrogen gas electrode, and ten times as precise as has been obtained with the deuterium gas e1ectr0de.l~ The (14) For further details in the preparation of cell solutions, see W. N. Baker and V. K . La Mer, J . Chcm. Phys., 8, 406 (1935). (15) Abel, Bratu and Redlich, z. physik. Chem., 8173,353 (1035).
The curve in Fig. 2 has an unexpected but definite minimum a t ND,o = 0.05, corresponding qualitatively to the flat minimum a t NDgO= 0.6 for the cell Pt,H2-D2/HC1-DCIin H20-D20/AgCl,Ag, studied by Abel, Bratu and Redlich.16 These data appeared during the progress of our investigation. No discussion or claim was given concerning the presence of a minimum in the curve for the gas electrode. The experimental difficulties inherent in the deuterium gas electrode prevented these authors from obtaining the precision which we have obtained with the quinhydrone electrode. They mention, however, that as a result of the large value of the constant for the exchange reaction (No. V, Table I1 below), the concentration of deuterium gas a t the electrode was greater than they had supposed a t the time the measurements were made. It is likely that their data for mixtures of HzO and D2O represent steady state values rather than true equilibrium values with the bulk of the solution since the exchange reaction occurs only on the platinum black at the surface of the electrode. This surface equilibrum will determine the e. m. f., but the e. m. f. will be disturbed depending upon the rateof convection and diffusion of acid and water from the bulk Of the solution in H20-Dz0 mix-
DEUTERIUM EXCHANGE EQUILIBRIA
Aug., 1936
1399
TABLEI1 EXCHANGE EQUILIBRIA IN SDLTJTIOh Process
NO.
Eo = (0.059in) log K
K
0,0345
14,64 0.96 15.30 1.30 11.80 11.26 28.58 2.42 0.40 .21 6.13
I 2DC1+ QHs = 2HC1+ QDi I1 QHz Do0 = QDz Hz0 I11 2DC1 Hz0 = 2HC1 Dz0
++
+
++
+
2DC1 &(g) = 2HC1 Dz(g) Dz(g) Hz0 c Hdg) Dz0 QHz Dzk) QDz Hzk) Hl(g) 2 H 2 0 Dz(g) 2NaOD = 2D20 HzO 2NaOD = Dz0 2NaOH IX HC1 NaOD = DC1 NaOH x 2NaOD HZ(g) = 2NaOH DZ(g) ODHz0 = H f OHDzO XI D ”
IV V VI VI1 VI11
++
+ + + + + + +
++
+ +
+
+ +
.0034
+ 2NaOH
.0431
+
.0233
Observer
L. and K. H. and L. 1-11 A., B. and R. 111-IV I-IV A.. B. and R. VII-V ‘/2 (VIII-111) VII-2v ‘/z (VI1 1V)
+
tures. The deuterium content of the gas phase was always greater than that determined by the exchange equilibrium set up on the surface of the electrode between the gaseous H2-D2 mixture and the HC1-DC1 and the H20-D20. Since an excess of hydrogen in the cell solution occurred in all the intermediate mixtures which were used, Abel, Bratu and Redlich’s measurements can be considered only as indicative of the presence and position of a minimum. Their value for pure DzO and the pure deuterium gas electrode obtained by short extrapolation, however, should not be affected appreciably by any disturbance of the equilibrium mixtures. The substitution of the quinhydrone electrode results not only in greater precision, but in contrast to the two-phase gas electrode, the exchange equilibria in cell (11), since they occur entirely in the cell solution, are homogeneous. This circumstance establishes our measurements in mixtures as true equilibrium values. Our results show conclusively that the e. m. f. curve of the quinhydrone electrode, when plotted against the DzO content of the solvent, has a minimum value, which we may reasonably suppose will be true for the analogous deuterium electrode, although the position of the minima will be different.
calculated from other data by use of the following equations’7
Exchange Equilibrium Constants
KQH2has the value 10-23.85.22Since K , H , / K Q D , = 11.26, then K Q D , = 10-24.70. Processes I and
K
of equation (I) may now be evaluated, and may be combined with constants derived in a similar way from the data on the deuterium electrode,I6and from the exchange reaction (F) QHz +I-Dz0 = QDz
+ Hi0
studied by Hamill and La Mer.I6 Table I1 lists the values of the constants obtained in this way. The value of the constant for No. V is in excellent agreement with the values which have been (16) Hamill and La Mer, unpublished results: for an outline of the experimental method, see Hamill and Freudenberg, THISJOURNAL, ST, 1427 (1936).
(G) HZ
+ Dz = 2HD;
K I = (HD)*/(Hz)(Dz) = 3.27
From Table 11, K4’ = 11.8; therefore our value of K3’ = (11.8)’”= 3.44. The values of K3’obtained from other sources are as follows. TABLEI11 VALUESOF K,’AT 25” Kd
Catalyst
3.8
Pt black
2.95 3.24
Pd black B. coli ( e . m. f . )
3.44
-
Observer
Bonhoeffer and Rummel18 Horiuti and Polanyi’s Farkas and FarkasZ0 Farkas, Farkas and YudkinZ1 La Mer and Korman
Av. 3.36
The exchange constants listed in Table I1 may also be interpreted as ratios of dissociation constants. For example, process VI is the ratio for the dissociation equilibria (M) QHks) = Q(s) (N) QD& = Q(s)
+ Hz(d; 4- Dn(d
KQH* KQD,
XI, Table 11, can be treated as ratios of dissociation constants for weak acids. If we assume that the activity of the chloride ion is the same in both (17) A. Farkas, “Light and Heavy Hydrogen,” Cambridge Univ. Press, Cambridge, England, 1935, p. 182 ff. (18) Bonhoeffer and Rummel, Naturwiss., 42, 45 (1934). (19) Horiuti and Polanyi, Nafure, 184, 819 (1933): 188, 142 (1934); lS4, 377 (1934). (20) Farkas and Farkas, Trans. Faraday Soc., SO, 1071 (1934). (21) Farkas, Farkas and Yudkin, Proc. Roy. Soc. (London), B l l S , 373 (1934). (22) V. K. La Mer and L. E. Baker, THISJOURNAL. 44, 1954 (1922).
’1400
SAMUEL
KORMAN AND VICTOR K.LA MER
TABLEIV DISSOCIATION CONSTANTS OF ACETIC AND (/.I E 0.05) Salicylic Acid I , QC
- &ba
1.
-E0
Molajity acid NDtO
25.0
0.17189 .17222 ,18381
0.47745
25.0
0.17095 ,16871
0.50885
SALICYLIC
Vol. 58
ACIDS
Molality Na salt
Molality NaCl
0.021811 .021855 .019770
0.019379 .019416 .021610
K
0
0.008304 .008321 .010504
N D t O = 0.917 0.008748 0.023766 ,008664 .021751 KBA/KDA= 4.09
0.020657 .02 1002
0.98 X 10-3 0.96 1.01 0.26 0.214
Acetic Acid NDzO
12.5 25.0 37.5 12.5 25.0 37.5 12.5 25.0 37.5
0.07773 .05945 .04020 ,09425 ,07658 .05807 ,08365 .06507 .04585
0.47906 .47745 ,47679
0.008682 .008682 .008682 .017179 .017179 .017179 .010568 .010568 .010568
12.5 25.0 37.5 12.5 25.0 37.5
0.11661 ,09242 ,07520 ,11301 ,08856 .07125
0.51046 .50885 .50819
0.027605 ,027605 .027605 ,025765 ,025765 .025765
= 0
0.035988 ,035988 ,035988 ,036280 .036280 ,036280 ,035661 .035661 .035661
0.0189991 .OB9991 .0189991 .019488 .019488 .019488 .019864 .019864 .019864
1.79 x lo-’ 1.85 1.80 1.74 1.79 1.73 1.77 1.80 1.71
0.043260 ,043260 .043260 .043571 .043571 .043571
0.023060 .023060 .023060 .021463 -021463 .021463
0.754 .613 ,634 .756 .611 .634
ND, = 0.917
I, o c .
KEA/KDA
12.5 25.0 37.5
2.34 2.96 2.75
waters, No. I may be considered as the difference of the dibasic acid dissociations
+
(0) QHz = Q2Hf (in HzO); K H (P) QDz = Q - + 2 D + ( i n D z O ) ; K D
and 37.5’, we employ the equation of Harned and Wright lo (4) Eo = -0.47745 4-(91.02 X
t-25)z
(3.008 X 10- )(
In the absence of a similar equation for the corresponding heavy water cell, we have assumed that the coefficients for the temperature terms are the same for both waters. The Eovalues are given in Table IV. Figure 3 shows R log ( K H A ~ / K D A = ) plotted against (l/T). The slope of the curve is a measure of the difference (AHH- AHD) of the heats of dissociation between proto- and deutero(R) Dz0 = D + OD-; Kw(Dz0) acetic acids. From these preliminary measureand equals 6.13 a t 25’. ments, there appears to be a maximum in the Salicylic and Acetic Acid Solutions.-The data curve. The indication that this quantity changes for cells (111) and (IV) for salicylic and acetic sign, as a function of temperature, suggests that acids are given in Table IV, with the ratios KH*/ the temperature coefficient of the dissociation conK D A . To obtain this ratio for acetic acid a t 12.5 stant of deutero-acetic acid be subjected to a more (23) S. E. Sheppard, Trons. A m . Elecfrockem. Soc., Preprint 39 detailed investigation to eliminate the assump(1821); La Mer and Parsons, J . B i d . Ckem.. 67, 613 (1923).
where K H , the average dissociation constant, equals (K’K’”’’’ = [(1.75 X 10-lo) (4 X for ordinary water.23 = 2.64 X Log (KH/KD) = 0.0345/0.05911 = 0.5837, or (KH/’KD) = 3.84. No. XI is the ratio of the acid dissociations (Q) Hz0 = H C OH-; K w ( H z 0 )
+ +
Aug., 1936
DEUTERIUM EXCHANGE EQUILIBRIA
tion of equality of the temperature coefficients of the DCl and HC1 cells.
+ I 2 0.900
%
0.700
6
0
The Solubility of Quinone, Hydroquinone and Deuteroquinone in HzO and DzO.-The errors in the determination of solubility in heavy water by ordinary analytical methods, eriz., titration, become greatly magnified when the quantity of solvent is severely limited, so it is scarcely feasible when dilution with ordinary water is undesirable. Furthermore, distillation of the solvent and weighing of the residue is not permissible with volatile substances like quinone. Although the Zeiss interferometer has long been used for the accurate determination of the concentration of dilute solutions encountered in precision cryoscopic work, the capabilities of the instrument apparently have never been exploited for the precise determination of solubility on small samples. We present some illustrative data necessary for the interpretation of the e. m. f. results upon the quinhydrone electrode in heavy water. The Method.-A Zeiss-Lowe-Haber portable liquid interferometer was employed. The description of this instrument and the principle upon which its use m determining concentrations of binary mixtures is based is given by Adams. 24 Deuteroquinone was prepared by dissolving an excess of hydroquinone in 92% DzO and separating by a vacuum distillation. The dry sample was kept in a vacuum desiccator. About 5 g. of heavy water is weighed into a glass-stoppered 25-cc. flask and a quantity of solute just insufficient for saturation is added. About 0.15 CC. of the solution is suflicient for a reading in the interferometer when a special 10mm. cell containing a 9-mm. glass block, producing a film of solution 1 mm. thick, is employed. Pure solvent is placed in the comparison chamber. For solutions more colored than saturated benzo(24) L. H.Adams. Tms JOURNAL, 87, 1181 (1916).
1401
quinone, 1 mm. is about the maximum thickness which will pass sufficient light. If desired the aliquot may be returned to the flask with a micro-pipet. In any case the flask and solution remaining are reweighed, a small but accurately vPeighed additional quantity of solute dissolved and a new reading made with the interferometer. Firially, when a saturated solution has been produced, a considerable excess of solute is added. The solution is momentarily heated several degrees above 25", and returned to the thermostat with constant swirling. Aliquots are removed a t convenient (ten-minute) intervals, until an almost constant reading with the interferometer is obtained indicating an equilibrium state. The solution is then momentarily cooled below 25' and the procedure is repeated at convenient intervals. The average of the asymptotes of the two converging time-concentration curves is taken as the interferometer reading for the saturated solution. Obviously these observations can be repeated very readily upon the same sample if satisfactory results are not obtained at first.
0.010 0.015 G. quinone/g. solvent. Fig. 4.Solubility of quinone in HsO and D20.
0.005
The tables and figures show that the interferometer readings are linearly proportional to the concentration a t least near the saturation point, and that the intersection of the time-concentration asymptotes and the concentration calibration curve yield a value for the concentration of the saturated solution which is precise to about 0.1 mg. per g. of solvent. For a 3-g. sample this
1402
KORMAN AND VICTORK. LA MER
SAMUEL
is approximately the error produced by evaporation during transfer of solute. Figure 4 shows the curves for quinone in light and heavy water;
Vol. 58
TABLEVI1 SOLUBILITY OF QUINONE, HYDROQUINONE AND DEUTEROQUINONE AT 25' G./g. H20
G.fg. Dz0
Mole/ mole Hz0
Mole/ mole Dz0
Mole ratio of soly. HnO/ DrO
Quinone 0.0151 0.0151 0.0025 0.0028 0.900 Hydroquinone .0802 . . . . .0131 . . . . ... . Deuteroquinone . . . .0611 . . . , .0109 1.204
.
Dissociation Constants of Weak Acids in HzO-DZO Mixtures.-The preliminary experiments with acetic and salicylic acids in 91.7y0 DzO have established the feasibility of the microtechnique and the quinhydrone electrode in heavy water, as a method for determining dissociation constants of light and heavy acid mixtures. To extrapolate the dissociation constant obtained in H20-D20 mixtures to pure DzO, we define a constant for weak acid solutions in H20-D20 mixtures by summing the various activities stoichiomet ricallyZ6
/
700
600 0.050
0.062
0.077
G. hydroquinone/g. solvent.
The fractions of deuterium ion and undissociated deutero-acid are
Fig. 5.-Solubility of hydroquinone in HzO and D:O.
Fig. 5 those for hydroquinone in H2O and for deuteroquinone in 99.8% D20. TABLEV SOLUBILITY OF QUINONEIN HzO G. Q l g . Hz0
0.00774 .01003 .01099 .01189 .01298
10 ' Reading, c0ncn.f corr. reading
94.0 120.0 130.0 141.0 155.7
Av. for satd, soh. .0151 178.5
823 836 845 844 834
AVD
G. 1219. D20
0.00646 ,00707 ,00865 .01103 .01191 .01282 .01353 .01451 .0151
Then
99.8% D& 108 Reading, concn./ corr. reading
64.2 ,. 71.9 88.4 .. 133.3 8275 143.9 8278 166.3 7710 176.4 7664 189.7 7646 196.5 Av. s. soh.
..
To evaluate the fractions a t various values of NDz0,we utilize the constants for the exchange equilibria (S)
HA
+ HDO = DA + H20; K
KHA 1 -X KDA
7.1
By so doing, we take account of the essential exchanges which must be postulated for HzO-DzO, HA-DA mixtures. Then
TABLE VI HYDROQUINONE IN H20 DEUTEROQUINONE IN 99.8% D 2 0
SOLU BILITY
OF
G. QHi/g. Reading, 108 concn./ G. QD?/g. Hz0 corr. D20 reading
0.06583 ,06592 .06959 ,07261 .07589 .07649 .07793 .0802
765.3 765.4 809.1 840.0
885.2 892.1 908.9 934.0
AND
Reading, 108 concn./ con. reading
8601 0.05336 661.7 8613 .05696 699.9 8597 .05861 721.0 8644 .05861 735.8 8573 ,05991 752.1 8574 8574 Av, for satd. s o h . Av. satd. .0611 765.9 soh.
8064 8139 8128 7965 7967
The water ratios may be calculated from the equilibrium (U) H10 DzO = 2HDO; K = 3.27 If, then, we stipulate a value for KH*/K=A,we may determine the behavior of the constant K H , D over the range of H20-D20 concentrations. The values for K H , D / K H A given in equation (5) are
+
(26) Abel, Bratu and Redlich (ref. 15) have employed a similar definition for the ion product of water in mixtures.
Aug., 1936
DEUTERIUM EXCHANGE EQUILIBRIA
1403
plotted in Fig. 6 as a function of KDA/KHA.It will be noted that the extrapolation is practically linear above 90% D20 for most ratios. Since measurements must frequently be made in DSO concentrations less than loo%, a reasonably accurate value for pure DzO should be obtained by a method of successive approximations. Thus K H A / K H , D = 2.96 for acetic acid in 92% D 2 0 becomes K H A / K D A = 3.3 for pure DzO. Since for most acids so far investigated this ratio lies between 3 and 6, the above method justifies our belief that extrapolations of K H , D from measurements over 90% DzO are fairly reliable.
points in the acid molecule other than those corresponding to the measured ionization. Thus, the exchange which occurs in the hydroxyl group of the salicylate ion may effectively alter the character of the ion, depending upon the isotopic composition of the medium.
ratio.
NEWYORK,N. Y.
Summary
The free energy changes for a number of important deuterium exchange reactions in aqueous solution have been investigated, using the quinhydrone, silver chloride-silver electrodes and a semi-microtechnique on hydrochloric acid and weak acid buffer systems. The dissociation constants of the weak acids salicylic, acetic, hydroquinone and water in light and 92% heavy water have been measured. With the exception of salicylic acid, the ratio of these dissociation constants in light and heavy water decreases as the constant increases. An equation has been developed which affords a controlled extrapolation of this ratio to 100% DzO from a single measurement aEove 90% D2O. The quinhydrone electrode yields reliable results of greater precision than have been obtained with the deuterium gas electrode. The superior0 ity of the quinhydrone electrode is due not only to 0 0.5 1.0 the simplicity of manipulation, but to the fact ND,o. that the equilibrium between the proto- and deuFig. 6.-Dissociation constant, ratios in tero-forms of quinhydrone, of water and of acid H20-D20 mixtures. is established promptly in a homogeneous solution. HalpernZehas predicted that K H A / K D A , the With the gas electrode the exchange reactions are ratio of the dissociation constants of a weak acid established only a t the. surface of the electrode, in pure HzO and pure DzO, should increase as which equilibrium can be disturbed by convecK H decreases. Our studies upon hydroquinone tion and diffusion from the bulk of the solution, and acetic acid, together with the ratio of the if the gas has been generated by the complete water constants, indicate that this prediction is decomposition of a portion of the water employed. probably correct. The somewhat anomalous The temperature coefficients of the e. m. f . indiratio for salicylic acid will serve to indicate that cate that the difference in the heats of ionization certain conditions must be satisfied before a rigor- of the proto- and deutero-acetic acids may change ous comparison of weak acids can be attempted. sign in the neighborhood of 25'. The constitution of the anion must not be affected The solubilities of quinone, hydroquinone and by the change of solvent medium for HzO to DzO. deuteroquinone in small quantities of light and A complicating factor may be introduced by the heavy water, necessary for the interpretation of exchange of protium and deuterium atoms at the e. m. f. data, have been determined precisely (26) Halpern, J . Chcm. Phys.. 8, 456 (1935). Halpern's treatby an interferometer method. Hydroquinone is ment can have the status only of an approximate theory, since he 20% more soluble on a mole ratio basis than deuconsciously omits a detailed consideration of the specific contributions of moments of inertia, mass ratios and interaction effects, teroquinone a t 25', whereas quinone suffers no leaving the contribution of the vibrational frequencies of the proton change in solubility. and deuton bonds as the only significant factor in determining the RECEIVED JUNE 6, 1936
