DISSOCIATION C O N S T A N T OF MORPHOLINIUM I O N
2869
Dissociation Constant of Morpholinium Ion and Related Thermodynamic Quantities from 0 to 50"
by Hannah B. Hetzer, Roger G. Bates, and R. A. Robinson National Bureau of Standards, Washington, D. C.
(Received February 88, 1966)
The thermodynamic dissociation constant of morpholinium ion (MH f) at 11 temperatures from 0 to 50" has been determined from emf measurements of hydrogen-silver chloride cells without liquid junction. The dissociation constant (K,) for the process H30+ is given as a function of T (OK) by the equation -log K , = R4H+ HzO $ 1 4 1663.29/T 4.1724 - 0.0042239T. At 25", -log K , is 8.492, AH" is 39,030 joules mole-', 48" is -31.7 joules deg-' mole-', and 4Cp0is 48 joules deg-' mole-'.
+
+
+
Introduction The measurement of the dissociation constant of piperidhie over a range of temperature has been describedl in an earlier communication. More recently, similar measurements for pyrrolidine have been reported.2 Pyrrolidine differs structurally from piperidine in that it has a five-membered ring with one CH2 group less than piperidine. Morpholine (tetrahydro1,4-oxasine) has a six-membered ring in which one CH2 group of piperidine is replaced by an oxygen atom. The three related bases have the following structures
0 NH
piperid'ine
n
NH pyrrolidine
NH morpholine
'
The pK, value of piperidinium ion a t 25" is 11.123 and that of pyrrolidinium ion is 11.305; hence the values of AGO for the dissociation process differ by 1040 joules mole-'. The values of AH" a t 25" differ by 1080 joules mole-', whereas those of AS" differ by less than 0.2 joule deg-1 mole-'. The thermodynamic quantities associated with the dissociation of piperidinium and pyrrolidinium ions are therefore quite similar in magnitude. On the contrary, literature value^^-^ for the p K , of morpholinium ion range from 8.3 to 8.5, corresponding to a value for AGO of about 48.5 kjoules mole-' for the dissociation process. This value is to be compared with 64.5 kjoules mole-' for pyrrolidjnium ion. It is of interest
therefore to make pK measurements for morpholinium ion over a range of temperature, in order to ascertain if the AH" and AS" values are also markedly different. We have now determined thermodynamic quantities H2O M for the dissociation process MHf H30+from 0 to 50" and have calculated pH values for a buffer mixture over this temperature range.
+
--f
+
Method The method used followed closely that for pyrrolidine,z except that the silver-silver chloride electrode was used instead of the silver-silver bromide electrode. The cell can be represented as P t ; Hdg, 1 iztm),
(md, C(C&)ZO(CHZ)~NH~C~
r(CHz)20(CH2)zNH (m2>,AgCl; Ag i
i
(1)
where m is molality. Three corrections sometimes needed in emf measurements of solutions of bases proved unnecessary in this case. One is the consideration of the partial pressure of the amine in correcting the pressure of hydrogen to 1 atm. The allowance in the case of pyrrolidine solu(1) R. G.Bates and V. E. Bower, J . Res. A'atZ. Bur. Std., 57, 153 (1956). (2) H. B. Hetzer, R. G. Bates, and R. A. Robinson, J. Phys. Chem., 67, 1124 (1963). (3) H. K.Hall, ibid., 60, 63 (1956). (4) H. K.Hall, J . Am. Cheni. SOC.,79, 5439 (1957). (5) A. R. Ingram and W. F. Luder, ibid., 64,3043 (1942). (6) A. Marxer, Hetv. Chim. Acta, 37, 166 (1954).
Volume 70, Number 9 September 1966
H. B. HETZER,R. G. BATES,AND R. A. ROBINSON
2870
tions was negligible2 even at 50". Since morpholine (bp 128" (760 mm)) is even less volatile than pyrrolidine (bp 86.5" (760 mm)), the correction can be omitted with certainty. This conclusion was confirmed by measurements of the emf of two cells of type 1 containing a solution of composition m1 = -0.1, m 2 = -0.05 a t 25, 40, 45, 50" and then again at 25". One cell contained a single hydrogen saturator, and the other was provided with an extra triple saturator. Since the values for the two cells agreed within 0.05 mv at all temperatures, the cells with one saturator were deemed to be adequate for the determination of the dissociation constant. Secondly, the increase in the chloride ion concentration resulting from the solubility of silver chloride in the cell solution has to be considered. The logarithm of the stability constant of the morpholine-silver complex ion has been reported' to be 4.98 at 25" and was found in the present work to be 4.2 a t 50". The increase in chloride ion concentration calculated in the usual way for the morpholine buffer solutions used here was negligible even at 50", and the correction mas consequently omitted. Thirdly, in calculating pK from emf data, the ratio (ml- n a H moa)/(nzz W ~ H- OH) is required. In this instance, this ratio could be taken equal to ml/ m2 because neither neutral morpholine nor its protonated cation is appreciably hydrolyzed in the buffer mix tures studied.
+
+
6 weeks after opening. No evidence of decomposition of morpholine was observed in this last set of solutions. The ratio of salt to free base (ml/m2) in three sets of solutions was approximately 2 and in a fourth set approximately 1.2. The molality of free base was determined by weight titration with standard hydrochloric acid to the calculated equivalence point (pH 4.8 for an approximately 0.07 M solution of the salt), detected with a glass electrode.
Results The emf values of cell 1 for 19 solutions are recorded in Table I. The emf of the cell is related to the pK, value of morpholinium ion by the equation
1%
The Journal of Physical C h m k t r y
('YMH +'YCl -/?'MI
where k denotes (RT In 10)/F and i\'I stands for morpholine. Introducing the Debye-Huckel equation with zero ion-size parameter, we obtain pKa' = ( E
- E " ) / k + log (mlz/m2)- 2Aml"'
(2)
where A is the Debye-Huckel parameter on the molal scale. With the E" values of Bates and Bower,9 good straight-line extrapolations of pKa' against ml gave limiting values of pK, a t ml = 0, using a zero ion-size parameter. The standard deviation of the intercept was 0.001 at all temperatures. These pK. values are given in Table I1 along with those calculated from the equationlo
Materials Examination of a ''purified" commercial grade of morpholine by mass spectrometrys indicated about 99% purity, with about 0.1% of heavier materials and a small amount of water present. This material was distilled under reduced pressure through a Podbielniak column with platinum Heligrid packing, rated a t about 100 theoretical plates. Three intermediate fractions, interspersed with small samples for test, were removed and sealed under vacuum. Mass spectra of the distilled material were consistent with the structure of morpholine. There was some evidence of traces of water, but no trace of substances heavier than morpholine was detected. Gas-liquid partition chromatograms of the distilled morpholine, using four different substrates, revealed only single peaks. The buff csr solutions were prepared, under nitrogen, by first adding morpholine to hydrochloric acid of known composition and then diluting this stock solution. Three sets of solutions were prepared using morpholine from vials opened immediately before use, and a fourth set was made from one kept stoppered for
+ log (m?/mJ +
pKa = ( E -
pK,
=
AI/T
- A2 + A3T
with A1 = 1663.29, A2 = -4.1724, -0.0042239 (0°C = 273.15"K).
(3) and Aa
=
Discussion Equation 3 leads, by standard thermodynamic equations, to values of the enthalpy, entropy, and heat capacity changes for the dissociation process. Values at 25" for morpholinium ion and for the two closely related substances are as shown in Table 111. These values for piperidinium ion and pyrrolidinium ion are remarkably similar except, perhaps, for the heat capacity change. However, for morpholinium ion (7) R. J. Bruehlmann and F. H. Verhoek, J. Am. Chem. SOC.,70, 1401 (1948). (8) The authors are indebted to Mr. E. E. Hughes for the mass spectrometric analyses, to Dr. R. T. Leslie for the distillation, and to both Dr. Leslie and Mr. Hughes for the gas-liquid partition chromatographic analyses of morpholine. (9) R. G. Bates and V. E. Bower, J. Res. Natl. Bur. Std., 53, 283 (1954). (IO) H. S. Harned and R. A. Robinson, Trans. Faraday SOC.,36, 973 (1940).
DISSOCIATION CONSTAST
2871
OF n/IORPHOLINIUM I O N
Table I : Electromotive Force of the Cell: Pt; Hz ( 9 , 1 atm), (CHz)zO(CHz)2NH4C1( m l ) , (CHZ)ZO(CHZ)ZNH ( m d , AgC1; Ag from 0 to 50" (in v ) I
5"
00
mi
7raz
0.10082 0.09024 0.08599 0.07955 0.07858 0.07271 0.06962 0.06542 0.05966 0.05513 0.05066 0.04989 0.03965 0.03475 0.02905 0.02068 0.017231 0.014732 0.009693
0.04334 0.07330 0.04446 0.03420 0,03860 0,03760 0.05655 0.03213 0.02565 0.02851 0.02488 0.04053 0,017047 0.017968 0,014267 0.008890 O.OO8910 0,007236 0 007873
Table 11:
Values of pK, from 0 to 50"
200
15O
25O
30'
0.77859 0.77881 0.77878 0.77859 0.77818 0.77752 0.77686 0.79542 0.79593 0,79621 0.79635 0.79627 0.79599 0.79555 ... 0.78541 ... ... ... ... ... 0.78279 0.78309 0.78318 0.78303 0.78271 0.78218 0.78156 0.78622 0.78657 0.78669 0.78663 0.78637 0.78590 0.78533 ... ... 0.78876 0.80000 0.80068 0.80105 0,80127 0.80130 0.80102 0.80081 0.78943 0.78990 0.79007 0.79011 0.78995 0.78956 0.78904 0.78800 0.78842 0.78852 0.78855 0.78829 0.78787 0.78736 ... 0.79441 ... ... ... 0.79424 0.79479 0.79509 0.79523 0.79515 0.79484 0.79437 0.80628 0.80715 0.80772 0.80806 0.80821 0.80814 0.80798 0,79581 0,79637 0.79674 0.79687 0.79683 0.79656 0,79623 ... ... ... 0.80409 ... , . . 0.80511 0.80585 0.80626 0.80655 0.80667 0.80658 0.80639 0,80875 0.80954 0.81014 0.81052 0,81069 0.81070 0.81047 .. , 0.81940 0.81861 0.81962 0,82031 0.82091 0.82129 0.82141 0.82147 0.83949 0,84087 0,84202 0.84295 0.84369 0.84418 0.84460 .
OC
PKS (exptl)
(calcd)n
0 5 10 15 20 25 30 35 40 45 50
9.108 8.978 8.850 8.727 8.608 8.492 8,380 8.268 8.161 8.056 7.955
9.108 8.977 8.851 8.728 8.608 8.492 8.379 8.269 8.161 8.057 7.955
1,
a
100
pK, = 1663.29/T
I
.
350
40'
450
50"
0.77588 0.79497 0.78403 0.78068 0.78442 0.78750 ... 0.78825 0.78654 0.79323 0.79368 0.80757 0.79562 0.80341 0.80585 0.81012 0.81911 0.82114 0.84475
0.77482 0.79420 0.78308 0.77974 0.78353 0.78661
0.77360 0.79332 0.78200 0,77857 0.78240 0.78555
0.77218 0.79224 0.78062 0.77727 0.78114 0.78432
...
...
...
0.78736 0.78554 0.79251 0.79290 0.80704 0.79489 0.80284 0.80529 0.80958 0.81873 0.82077 0.84477
0.78635 0.78444 0.79145 0.79178 0.80632 0.79397 0.80204 0.80450 0.80887 0.81822 0.82026 0.84459
0.78508 0.78319 0.79035 0.79057 0.80544 0.79290 0.80111 0.80359 0.80808 0.81752 0.81961 0.84426
pyrrolidinium ion at 25' but the difference changes with temperature in a manner consistent with the differences in the other thermodynamic quantities. The pK, values of morpholinium ion decrease with increased temperature less rapidly than do those of pyrrolidinium ion. Thus the difference in pK, is only 2.601 a t 50". While mixtures of morpholine and morpholinium hydrochloride are not proposed as pH standards, they may have some uses in biochemical work when a medium of controlled pH is needed. We have, therefore, interpolated in Table I to obtain values of p ( a ~ y c l (=-log ) aH+ycl-) and have derived the corre-
P&
+ 4.1724 - 0.0042239T.
Table IV: Values of ~ ( U H ~and C I ~) C Lfor E a Buffer Solution Composed of Morpholine Hydrochloride (0.1 m ) and Morpholine (0.05 m ) from 0 to 50"
Table I11 t,
AP,
ACp',
PKS
joules mole-1
joules deg -1 mole-'
joules deg -1 mole-'
8,492 11.123 11.305
39,030 53,390 54,470
-31.7 -33.9 -33.7
48 88 68
AH',
Morpholinium Piperidinium Pyrrolidinium
there are marked differences both in pK. and in the enthalpy change. Thus, not only is there a large difference (2.813) in pK, between morpholinium ion and
OC
0
5 10 15 20 25 30 35 40 45 50
9.068 8.934 8.809 8.687 8.567 8.453 8.341 8.231 8,125 8.021 7.921
8.963 8.828 8.702 8.579 8.458 8.343 8.231 8.120 8.013 7.908 7.806
Volume YO, Number 9 September 1966
WARRENV. SHERMAN
2872
sponding va,lues of PCZH ( z - l o g UH+)by application of the Bates-Guggenheim” convention
The values obtained for the PUH of one buffer mixture from 0 to 50” are given in Table IV. (11) R. G. Bates and E. A. Guggenheim, Pure Appl. Chem., 1, 163 (1960).
The 7 Radiolysis of Liquid 2-Propanol. II.la The Reaction of Solvated Electrons with Mono- and Disubstituted Benzenes
by Warren V. Sherman’b Soreq Nuclear Research Centre, Yavne, Israel
(Receised March 1, 1966)
The effect of a number of mono- and disubstituted benzenes on the gaseous products of the radiolysis of solutions of nitrous oxide in 2-propanol has been studied. The substituted benzenes decrease the nitrogen yield, and this is interpreted as being due to competition with nitrous oxide for solvated electrons. From the measurements of G(N,) the relative reactivities of the substituted benzenes have been calculated. The site of attack of the solvated electron appears to be the aromatic ring since the substituent effects correlate well with those found generally in aromatic bimolecular nucleophilic substitution. In addition, satisfactory correlation is observed between reactivity and the Hammett cParafunctions. The p value for the monosubstituted benzenes is 3.1. The p value remains approximately constant for a given series of disubstituted benzenes in which one substituent is kept unchanged. Hoa~ever,the u values for disubstituted benzenes are not additive.
Introduction Solvated electrons are produced in the radiolysis of aliphatic alcohols.2 A knowledge of the subsequent reactions of this species is necessary for the complete understanding of the radiation chemistry of alcohols and alcoholic solutions. It has been shown1*that the molecular nitrogen produced in the radiolysis of dilute solutions of nitrous oxide in 2-propanol is a convenient measure of the solvated electrons which are amenable t o solute scavenging. Utilizing this technique, the measurement of the relative rate constants of a number of mono- and disubstituted benzenes are reported.
Experimental Section Materials. 2-Propanol was purified as described previously.18 The mono- and disubstituted benzenes The Journal of Phhyskal Chemistry
were of reagent grade and used without further purification. Procedure. The preparation and irradiation of the samples, and the analysis of gaseous products were as described previously.la The mean dose rate during the course of the present series of experiments was 3.7 X 101’ ev ml-’ mink’. All solutions were irradiated for 15 min.
(1) (a) Part I: W. Y. Sherman, J . Phys. Chem.. 70, 667 (1966). A preliminary communication of some of the results reported here appears in W. V. Sherman, J. Am. Chem. Soc., 88, 1567 (1966). (b) The Radiation Laboratory, University of Notre Dame, Notre Dame, Ind. (2) M. C. Sauer, S. Arai, and L. M. Dorfman, J . Chem. Phys., 42, 708 (1965),and references therein.
