1679
NOTES
Dec., 1957
THE THERMODYNAMICS OF HYDROCHLORIC ACID AND T H E IONIZATION CONSTANTS OF FORMIC, ACETIC AND PROPIONIC ACIDS I N 83 WEIGHT PER CENT. DIOXANE-WATER SOLUTIONS
-003
c
BY S. S. DANYLUK, H. TANIGUCHI AND G. J. JANZ Department of Chemistry, Rensselaer Polytechnic Institute, Troy, New Yorh Received June 88, 1967
The present communication reports corrected values for the thermodynamic properties of hydrochloric acid and the ionization constants of the simple aliphatic acids from a recalculation of experimental data report,ed by Harned and associates' elsewhere. A need for this revision was observed in the course of a study of the therniodyiiamics of hydrogen chloride in absolute ethano1.2 It was found that the extrapolation of the e.m.f. data for the cell Hz I HCI ( m ) ,Dioxane (82%), HzO (18%)I AgCI-Ag gave a value at 25' for EO of -0.0614 volt, rather
u 0
0 01
0 005 rn.
Fig. 1.-Extrapo1:tion of (EO' - E,) in 82y0 dioxanewater mixture at 25 : 0 , HCI as strong electrolyte; @, 0, Harned and associatesla and present calculations, assuming HC1 as an ion associated electrolyte.
than the value, -0.0415 volt reported earlier.lb Gronwall, LaMer, and Sandved extended term The method of extrapolation was that outlined by equations after Harned, el ul.,lb and the dissociaHarned and his associateslb for solvents of low di- tion constants of Owen and Watersa3 The extrapelectric constant (ie., 10) in which appreciable ion- olation Eo' - E, is illustrated in Fig. 1, together pair formation is present. It involves the law of with the results obtained employing the Debyemass action and dissociation constants obtained Hiickel function EO' and the extrapolation by by conductance measurements. Harned and associates.' It is seen that the value When the Debye-Hiickel theory and the Gron- for EO a t 25O, -0.0614, is about 30 mv. lower than wall, LaMer and Sandved extension are inade- that based on hydrogen chloride as a strong elecquate, due to ion-pair formation, as is the case in trolyte and aboht 20 mv. lower than that reported this system, the cell electromotive force may be by Harned, et al.' The slight tail-up in the corexpressed by the equation rected extrapolation curve has been shown, in the case of the hydrogen chloride-ethanol system,2 to E = EO - 2k log ormy, (1) need for a better choice of the a paramindicate where T~ is the activity coefficient a t the actual eter in athe theoretical treatment. The same ionic concentration am. Combination of (1) with values as used by Harned and associates* for the the deviation function dielectric constants and ion-size parameter were 2kA'dE used in the present calculations. Evaluation of the EO' - EO = E - EO + 2k log m - 1 ~+ B/,,//E f(m) influence of variation of this parameter on the (2) value of Ea requires e.m.f. data in the region of extreme dilution, Le., less than 0.005 m. leads to the expression The work in 82% dioxane-water solutions is the 2kA'dt Eo E (EO'- EO) -2k log a y o - 1 + BI.\/a = f(ln) only example of precise physico-chemical measurements of this type in a medium of low dielectric (3) constant. A recalculation of the results of Harned Rearranging (3) gives the expression for the ap- and associatesla-g was accordingly undertaken to parent standard potential, i.e., (EO' - E,), meet the need for this information with reference which on extrapolation to infinite dilution, ie., to the study of the properties of electrolytes and m = 0, leads t o the desired Eo since there f(m) is medium effects in such systems. zero. The values of a and yrrwere calculated by Results a series of successive approximations using the 1. Table I lists the corrected EO values in the (1) (a) H. El. Harned and B. B. Owen, "The Physical Chemistry of temperature range 5-45' for the cell Hz I HCl(m), Electrolytic Solutions," 2nd Ed., Reinhold Publ. Corp., New York. Dioxane (82%), H 2 0 (18%) J AgC1-Ag, as obN. Y.,1950,Chap. 1 1 and 15: (6)for extrapolation of I$', see: H. S. Harned, F. Walker and C. Calmon. J. A m . Chem. Sot., 61, 44 (1939): tained by the extrapolation method discussed ( c ) for thermodynamic properties, see: H. 8. Harned and F. Walker, previously. In addition, the deviations of the ibid.. 61, 48 (1939); (d) eummary and critique: H. S. Harned, J. 0. extrapolated values from the values calculated usMorrieon. F. Walker, J. 8. Donelson and C. Calmon, ibid., 61, 49 ing the equation (1939); (e) ionization constants of formic acid: H. 8. Harned and E
R. 8. Done, ibid., 63, 2579 (1941): (f) ionization constants of acetic acid: H. S. Harned and L. D. Fallon, ibid., 61, 2377 (1939); ( 9 ) ionization constants of propionic acid: H. S. Harned and T. R. Dedell, ibid., 63, 3308 (1941). (2) H. Taniguchi and G. J. Janz, THISJOURNAL, 61, 688 (1957).
EO = -0.0611
- 25.73
X 10-4 ( t
-
to)
-
(t
9.65 X
- toP
(4)
(3) B. B. Owen and G. W. Waters, J. A m . Chem. Soc., 60, 2371 (1938).
.
1680
NOTES
Vol. 61
TABLE I TABLEI1 STANDARD POTENTIALS FOR THE CELLHz HCl(m), DIOXANE MEANACTIVITY COEFFICIENTS OF HYDROCHLORIC ACID I N 82% DIOXANE-WATER (82%,), HzO (18%)1AgCl-Ag
I
ED (extr.) (V.)
T, ' C .
50
m
A
0.001 .0015 .002 .003 .005 .007 .010 .015 .020 .030 .050 .070
(mv.)
5 -0.0130 -0.4 10 - .0246 0 15 - .0370 +0.7 20 - .0487 .3 25 - .0614 .3 30 - ,0738 - .4 35 - .0871 - .7 40 - .lo12 - .6 45 - .1172 .~ .8 Max. deviation = 0.8 mv.; av. deviation = 0.47 mv.
+ +
+
25
15O
350
450
0.3063 0.2843 0.2698 0.2540 0.2272 .2690 .2497 .2364 .2221 .1978 .2432 .2256 .2134 .Zoo1 .I776 .2067 .1914 .181g .1708 .1504 .1674 .1380 .1212 .1561 .147g .145g .1361 .1288 .119g .lo47 .1259 .I176 .1105 .lo23 .OS93 .lo63 ,0991 ,0858 .0744 .0930 .0946 ,0654 .0878 ,0822 .0757 .0804 .0741 .0691 .0633 .0547 .065g .0603 .0560 ,0513 .0441 ,0563 .0521 .0484 .0441 .0377 .os06 ,0465 ,0429 .0390 .0332 .0453 .0412 .0378 .0343 .0291 .0426 ,0386 .0353 ,0319 ,0269 .0407 .0367 .0333 .0297 .0248 ,0430 .0298 .0244 .0381 .0340
are given. The constants in equation (4) were de.loo termined by the method of Cox and M a t u ~ c h a k . ~ .150 The deviations in Table I are somewhat greater .ZOO than those observed by Harned, et al.,lb and result .300 because of the uncertainty in extrapolating the .500 curved portion to obtain the EO'S. This uncerTABLE I11 tainty will lead t o a somewhat greater error in the thermodynamic properties than that reported by PARAMETERS OF THE EQUATIONS FOR THE RELATIVE PARHarned. TIAL MOLARHEATCONTENT AND HEATCAPACITY OF HCI 2. The mean activity coefficients of hydroIN 82% DIOXANE-WATER chloric acid in 820/, dioxane-water mixtures listed , B Cz(25') >~(25') in Table I1 were recalculated for the temperature 0.001 5990 0.0948 2438 56.5 range 5-45' and up to 0.5 m using the relation .0015 6520 ,1019 2539 60.8 -01
E = EO
- 2k log- m .y
(5) . .
3. Values for the parameters of the equation for the relative partial molal heat content
Lz
= a
+
p ~ cal. 2
(6)
and relative partial molal heat capacity
fz
=
2pT cal.
(7)
and the thermal properties 1,and J2 for 25' are given in Table 111. Because of the greater uncertainty in extrapolation of the Eo's, the error in the recalculated values will be greater than in the original results of Harned, et al. IC 4. Corrected values of the ionization constants for formic, acetic and propionic acids in 82% IONIZATION CONSTANTS FOR T ,' C .
KAX
1010 (formic)
K.4 X 10l1 (acetic) (propionic)
KAX
THE
7345 8600 10050 10835 11350 11660 11800 11935 12060 12135 12235 12410 12590 12910 13630
.loo
.150 .200 .300 .500
.1115 .1259 .1433 .1535 .1608 .1654 .1677 .1704 .1735 .1754 .177g .1817 .1854 ,1925 .2075
2567 2592 2689 2811 2945 3044 3108 3213 3364 3458 3580 3743 3892 4203 4817
66.5 75.1 85.5 91.5 95.9 98.6 100.0 101.6 103.5 104.6 106.1 108.3 110.6 114.8 123.7
TABLE IV SIMPLEALIPHATICACIDSIN 82% DIOXANE-WATER
5
10
15
20
25
30
35
40
45
9.754 3.60 1.986
9.170 3.55 1.965
8.417 3.40 1.886
7.919 3.32 1.862
7.231 3.10 1.771
6.671 2.96 1.702
5.983 2.67 1.587
5.234 2.43 1.439
4.295 2.03 1.220
dioxane-water from 5-45' are summarized in Table IV. These were recalculated using the equation
. A graph of log K'
versus p gives a straight line which was extrapolated to obtain log K Aat p = 0. The corrected EO values were used in (8). 5. The corrected values for a comparison of the primary medium effectsla a t 25' in 82% dioxane (4) G . J. Cox a n d M.
.002 .0°3 .005 .007 .010 .015 .020 .030 .050 .070
C, Matuschak, THISJOURNAL, 45, 36 (1941).
are log YOHCI, 2.399; 1,og YOA for formic, acetic and propionic acids are 2.69, 2.88 and 2.94, respectively. ~~~~~~i~~~ of the present results with the earlier values1 shows that the generalizations based on the trends within any series are essentially unchanged. The present values however differ appreciably in magnitude from the earlier and are recommended as the best values, within the limits possible of the experimental data available. They should be used for any quantitative (5) I n private communication with Professor H. 6. Harned i t was learned that during the World War I1 period the original data and records on the 82% dioxane-water solutions were lost. It has n o t been possible to find a n explanation for the large error in the earlier calculations.
.
Dee., 1957
NOTES
calculations with the thermodynamics of hydrochloric acid, and the ionizatioii of the aliphatic acids in 82% dioxane-water solutions. Acknowledgment.-This work is part of a broader program of study on simple electrolytes. in polar organic solvents, supported by Contract AT(30-1)-1999, United States Atomic Energy Commission. The authors wish to acknowledge with thanks the continued interest of Professor H. S. Harned and Dr. Roger G. Bates in this work.
THE HEAT OF SOLUTION OF SODIUM METABORATE AT 0’ BY GEORGEGRENIERA N D DAVIDWHITE Contribution of the Cryogenic Laboratory, Department of Chemistry, The Ohio State University, Columbus 10, Ohio Received Julu 91,1967
Since there has been considerable interest of late in the thermodynamic properties of borates, we have decided to present the results of some measurements on the heat of solution of Naz0.B203from which the heat of formation of sodium metaborate can be calculated. Recently Shartsis and Cappsl determined the heat of solution in 2 N nitric acid of various mixtures of stable compounds in the Na20-B20asystem. Only in one case did the composition of the starting material correspond to a pure borate. This compound was Na20.2Bz03. By extrapolation of their data to the composition corresponding to sodium metaborate a value of the heat solution is obtained which leads to a heat of formation of this compound of -468.3 kcal./mole. This is not in agreement with the value of -506 kcal./mole2 given in the literature. Experimental Apparatus and Procedure.-The measurements were made with a modified Bunsen ice calorimeter, which has been described previously .3 The calorimeter was a modification of the calorimeters developed by the National Bureau of Standards.416 The procedure for heat of solution measurements using this ice calorimeter as well as its calibration is discussed in detail by Clifton and MacWood.3 After evacuation to remove traces of water, the sodium metaborate samples were sealed in approximately spherical glass bulbs in a dry nitrogen atmosphere. The samples were dissolved by shattering the bulbs in a standardized solution of 2.005 N nitric acid in the calorimeter. In the early experiments the ratio of solute t o solvent was regulated so that i t would correspond to the ratio used by Shartsis and Capps. However, it was discovered that there was insufficient solvent to dissolve the sample completely when this ratio was used. Since our calorimeter operated a t Oo, 25“ lower than the one used by Shartsis and Capps, the solubility of the metaborate had decreased sufficiently to make their ratio of solute to solvent impractical. The working ratio used in this series of determinations was 0.001 mole of sodium metaborate to 25 cc. of solvent, 2.005 N nitric acid. Materials.-The sodium metaborate was prepared by removing the water of hydration from the tetrahydrate of the metaborate by a process of simultaneous heating and evacua(1) L.Shartsis and W. Capps, J . A n . Ceram. Soc., 37, 27 (1954). Values of Chemical Thermodynamic Properties,” National Bureau of Standards Circular 500. (3) D. G. Clifton and G. E. MacWood, THISJOURNAL, 60, 309 (1956). (4) D. C. Ginning8 and R. J. Corruccini, J . Research Natl. Bur. Standards, 38, 583 (1947). B(5)ID. C. Ginnings, T. B. Douglas and A. F. Ball, ibid., 46, 23 (1 950). (2) “Selected
1681
tion. This tetrahydrate, which is sold commercially under the trade name of “Kodalk,” was purified by recrystallla* tion from distilled water. Final purification was completed by placing the anhydrous metaborate in a furnace and heating above its melting point followed by slow crystallization from the melt. The material was stored in a vacuum desiccator.
Results and Discussion The results of the heat of solution measurements are listed in Table I. The average of five determinations gives a value of -20.43 & 0.36 kcal./ mole, a t 0’. TABLE I HEATOF SOLUTION OF SODIUM METABORATE AT 0’ N NITRICACIDSOLUTION Sample wt.
Heat measured
IN
2.005
-AH8
Run
(P.)
(cal.)
(kcal./mole)
1 2 3 4 5
0.1814 .1886 .1328 .1159 .1404
27.48 29.20 21.23 17.90 21.83
19.94 20.38 21.05 20.32 20.47
The heat of formation of Na*O.B2O3a t a certain temperature can be calculated from the heats of the following reactions a t the same temperature
+
Naz0.B203(s) solvent I = end soln. 1 B203(s) solvent I1 = end soln. 1 Na20(s) solvent I11 = end soln. 2
+
+
(1) (2)
(3)
Solvent I and I11 are solely 2.005 N nitric acid solution. Solvent 11, which is identical in composition to end solution 2, consisted of 0.001 mole of Na20 in 25 cc. of 2.005 N nitric acid. This ratio makes the end solution of reaction 2 identical in composition to the end solution of reaction 1. If the heat of reaction 1 is subtracted from the sum of reactions 2 and 3, the heat of the following reaction is obtained NazO(s)
+ Bz03(s) = NazO.BzOa(s)
(4)
From the heat of reaction 4 one can calculate the heat of formation of Naz0.B203(s),if the heats of formation of the pure oxides are known. The heat of reaction 1 was measured experimentally a t 0’ where solvent I was 2.005 N nitric acid. The value is -20.43 kcal./mole. The heat of reaction 2 could not be measured because the heat effect for this reaction was too small to.be measured with any reasonable precision on the existing apparatus. However, the heat of solution of BzOZ has been measured in nitric acid solutions as well as those containing borates a t 25°.1#637 It is evident from these results that heat effects sccompanying changes in concentration of the solvent are m i t e small. Thus. it can be safelv assumed that (he heat of reaction 2 using solvint I1 is -3.64 kcal./mole at 25°.1fj57 The heat of solution for reaction 3 a t 25’ is calculated to be -83.9 kcal./mole.2~8This calculation was made from data in the literature using the heat of neutralization of a strong base by a strong acid and the heat of solution of NazO in water. The calculation is valid, if the heats of dilution as(6) J. C. Southard, J. Am, Chem. Sou., 63, 3147 (1941).
(7) E. R. Van Artsdalen and K, P. Anderson, ibid., 73, 579 (1951). (8) F. R. Bichowsky and F. D. Rossini, “Thermochemistry of Chemical Substances,” Reinbold Publ, Corp,, New York, N. Y., 1936.
