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.
NOTES
1682
sociated with an actual experiment of this kind are shown to be negligible. Shartsis and Capps,l based on a literature survey of similar experiments, show that the heat of dilution of reaction 3 can be considered as negligible. In addition, their data for samples of different sizes are in good agreement, indicating the small effect of dilution on the heat of solution. I n order t o compute the heat of reaction 4 one must know the difference in the heat of solution of Na20.B20, (reaction 1) between 0 and 25'. Although in general the heat of solution does not vary appreciably over a small range of temperature, it may do so in this case because the end solu' may not be the same as a t tion in reaction 1 a t 0 25'. If one rejects this possibility and assumes the heat of solution for reaction 1 is the same a t 0 as a t 25' the heat of reaction 4 at 25' is -67.1 kcal./ mole. Using values of -99.4 and -305.3 kcal./ moleg for the heat of formation of NazO and BzOa respectively, the heat of formation for Na20.B203 2.8 kcal./mole a t 25'. The unceris -471.8 tainty given above for the heat of formation of sodium met'aborate does not include the error in assuming the heat of solution of reaction 1 is the same a t 0' as a t 25'. However, it is felt that this assumption will not introduce an additional uncertainty in the heat of formation of sodium metaborate of more than f 2 kcal./mole. The calculated value for the heat of formation of sodium metaborate of -468.3 kcal./mole from the data of Shartsis and Capps' is in good agreement with the value given above, but in disagreement with the literature value of -506 kcal./mole.2 It therefore seems reasonable to assume that calculations of the heats of formation for the other borates in the Na20-B203system involving only an interpolation of the heats of solution data of Shartsis and Cappsl will be in good approximation to the true values. Shartsis and Capps made only one set of measurements a t a composition corresponding to a pure borate. This compound was Na20.2BzO3, sodium tetraborate. Using their experimental value for the heat of solution, we have calculated the heat of formation to be -786.4 kcal./mole a t 25' (literature value = -777.79. In this calculation the same assumptions in regard to the heats of dilution were used. Interpolation of the data of Shartsis and Capps t o the composition Naz0.3B203 yields a heat of formation value of -1101.4 kcal./mole a t 25'. For the composition Na20.4BzO8, the interpolation gives a value for the heat of formation of - 1413.2 kcal./mole a t 25'.
*
.
(9) J. P. Coughlin, "Heats and Free Energies of Formation of Inorganic Oxides," Bureau of Mines Bulletin 542.
HEAT CAPACITY OF TITANIUM DIBORIDE FROM 30 TO 7'00' BY B. E. WALKER, C. T. EWING AND R. R. MILLER Inorganic and Nuclear Division Chemistry Division U. 8. Naval Research Laborator;, Washington 26, D . b. Receiued August 1 4 , I967
The purpose of this investigation was to attain
Vol. 61
some useful and accurate high temperature data on titanium diboride. It is hoped that this information will be useful in estimating high temperature properties of similar diboride substances. Experimental Apparatus and Method.-The calorimeter, furnace for heating the samples and the procedures used in these determinations have been described previous1y.l I n brief, the method consists in heating the sampIes in a container of known heat content to a known temperature in the furnace, then dropping it into the calibrated calorimeter and measuring the temperature rise of the calorimeter. Repetition a t a number of furnace temperatures provides heat content data, from which the heat capacity may be derived. The heat content measurements of TiBl were all adjusted to the same temperature, 30°, which is the operating point of the calorimeter. Small corrections, a maximum of 0.05% of the total heat measured, were made to account for the oxide formation on the container (monel, in these determinations) and for the argon gas in the bucket. I
Experimental Results The authors are indebted to Dr. F. Glaser of American Electro-Metals Corporation for preparing, finishing and grinding the TiBz sample t o fit the monel measuring bucket. Analysis of the sample indicated a purity of 99.7y0. The impurities consisted of approximately 0.2770 boron and 0.1% various metals, principally iron. No corrections were made to account for the impurities, since calculations showed that this would affect the enthalpy measurements by less than 0.05%.
0.14
Yii/liii
0
400 500 600 Temperature, "C. Fig. 1.-Heat capacity of TiB2.
100 200 300
700
The sample weighed 13.6355 g., in vacuo, and the monel container, 28.8336 g. The container was not sealed, but was thoroughly flushed with argon gas before being placed into the calorimetric system. A constant flow of argon is maintained in the system a t all times to provide an inert atmosphere and thereby minimize oxidation of the container. The heat content results for TiBz are listed in Table I and agree to about A2y0 with those computed from I
