THEF ; Q ~ . I L I B ROF~ ABISMUTH AND BISMUTH TRIIODIDE
March, 1961
52 1
is similar in many ways to the ruthenium-chlorine system: CrC13(g) and CrC14(g) are important vapor species, and reactions analogous to those in Table I occur. Doerner found standard entropies for TABLE I1 CrC13(g) and CrC14(g) t o be 86.1 and 88.3 e.u., CO&iPARISON O F AS0zss VALUES FOR THE RUTHENIUM- respectively, which may be compared with the CHLORINE SYSTEMA N D THE CHROMIUM-CHLORINE SYSTEM above values for RuC13(g) and RuCL(g). Table -AaSo298 (e.11.)I1 gives a comparison of AX OW^ values for analogous Reaction Ruthenium" Chromium b reactions involved in the two systems. In our hlCl,(s) = blC13(g) $64.7 f 4 . 0 +55.1 studies of the dissociation pressure of RuCla(s),6 1 MC&(s) + - Clz = MCl,(g) +32.2 rt 3 . 0 $30.7 we found the standard entropy for RuC13(s) to 2 R~(s)+ 3/2C12 MCl,(g) + 8.3 & 3.0 + 0.5 be 30.5 =t 2.5 e.u., which may be compared to a M(s) + 2c12 = MCl,(g) -24.0 f 4 . 0 -24.0 value of 31.0 e.u. for CrC13(s) found by Doerner. This work. H. A. Doerner.$ Acknowledgments.-The authors are indebted to The chromium-chlorine system at high tempera- R. C. Jensen, M. Tagami and R. E. Inyard for perture, which was thoroughly studied by D ~ e r n e r , ~forming part of the experimental work.
value for RuClr(g) seems reasonable; however, the value for RuCl,(g) is about 10 e.u. higher than one might expect.
THE EQUILIBRIUM 2/3Bi(1)
+ l/3Bi13(g) = BiI(g)'
BY DANIELCUBICCIOTTI Stanford Research Institute, Menlo Park, California Reeeiued September B6, 1960
The equilibrium 2/3 Bi(1)
+ 1/3 BiIs(g) = BiI(g) was studied by a transpiration technique over the temperature range
545 to 735' and the pressure range 0.1 to 50 mm. The enthalpy change for the reaction was 21 kcal. and the standard
entropy 16 e x . a t mid-temperature. An excess pressure of reduced Bi species in the higher pressure range vias intprpreted due to the formation of a polymer, B i A
:LP
Introduction The equilibria of liquid Bi with gaseous BiX3 and BiX, in which X was C1 or Br, have recently been ~tudied.~,3For those systems rat'her appreciable percentages of the monohalide were found in the equilibrium vapor, but no evidence of a dihalide (whether as BiXz or as Bi2X2). The iodide system was investigated as a companion study, with the results reported below. Experimental
periment were reduced to the pressures of BiI and Bi13 required to account for the amount and composition of the transpired material. Then these pressures were plotted as logarithm of pressure of BiI us. logarithm of pressure of BiI,. In such a plot consistency of results with a single equilibrium is indicated by the concordance of the points with a straight line, and the relative stoichiometry of the gaseons species in the equilibrium is given by the slope of the line. The log plots of the present data are given in The same transpiration method was used as in the chlo- in Fig. 1. For the lower half of the diagram (presride work*; however, i t was found necessary to envelop the exit tube of the apparatus with an inert atmosphere sures below about 3 mm. of Bi13) the points fall box because of the strong susceptibility of the products to reasonably well on the straight lines drawn at air oxidation. The Bi-I samples obtained were analyzed each temperature. Those lines mere drawn at a for Bi by conversion to BizOa as in the Bi-Br system.8 The Bi13 was made from Biz03and aqueous HI, dried and slope of 3.0, so that the equilibrium obtained in that region was doubly distilled in an NZstream. The transpiration experiments were made a t five different temperatures: 543-547; 592-596; 639-642; 683-687; 734-736', although the temperature of each run was constant to 1". The pressures finally calculated were adjusted (with very small corrections) to the temperatures: 545, 595, 640, 685 and 735', respectively. The flow rates were such that from 0.005 to 0.2 mole of Nz was passed in 8 hours and from 0.1 to 0.2 g. Bi-I sample collected. From the results of experiments to evaluate diffusion effects a correction of from 0 to 6% was made on the results, the larger corrections applying to the experiments made with smaller flow rates and higher pressures of Bi species in the gas stream.
Results and Discussions The data obtained were treated like those of the earlier systems2,3; that is, the results of each ex(1) This work was made possible by the financial support of the Research Division of the United States Atomic Energy Commission. (2) D. Cubicciotti, .I. Phvs. Chrm.. 64, 791 (1960). (3) D. Cubicciotti, ibid., 64, 1506 ( I S G O ) .
2/3 Bi(1)
+ 1/3 EiIa(g) = BiT(g)
(1)
The equilibrium constants for that reaction were calculated from the data on the straight lines and these were plotted vs. reciprocal of absolute temperature, as in Fig. 2, to give the enthalpy change for that reaction. Since a straight line represents the data well, the enthalpy change reaction 1 mas constant in the temperature range studied and equal to 21.1 f 0.8 kcal. The standard free energy changes for the reaction derived from the equilibrium constants at 545 and 735' were 7.80 f 0.1 and 4.76 f 0.1 kcal., respectively. Thus the standard entropy change for reaction 1 at the midtemperature (640') was 16 1 e.u. Kelley4 gives values for the absolute entropies of Bi(1) and Bi(4) I(.K. Kelley, U. S. Bur. &lines BuU, No. 477 (1950) and No. 584
*
(1960).
522
DANIEL CUBICCIOTTI
Vol. 65
either iodide on which to base a calculation of the other. The dissociation energies of BiI(g) reported5 40 by Gaydon (2.5 f 1 e.v.) and Herzberg (2.7 e.v.) are too doubtful to use for this purpose and there are no 20 measured values for Bi13. In the higher pressure region (above about 3 mm. Bi13) it was apparent that 10 the results were departing 8 systematically from the 6 lines representing the equilibrium of equation 1 (see Fig. 1). The departures ; 4 were in the direction of an excess of lower-valent Bi I -" over that expected from equation 1. They increased m 2 E as the pressure increased 3 and, when expressed as BiI 10 r w in excess of equilibrium 1, a 10 became as large as 25% of the BiI pressure. This excess of lower-valent Bi was felt to be an indication of polymerization of 04 the BiI to form one or more new species of formula Bi&. Therefore, the difference between the experimental value of the pressure of BiI and the value calculated from the equilibrium constant derived from the low pressure values (ie., the value given by the straight lines in Fig. 1) was plotted on a log-log plot. The rePRESSURE 881 - mrn. Fig. 1.-Logarithmic plot of pressure of BiII us. BiI at 545, 595, 640,685 and 735". The sults were quite scattered since they represented small straight lines were drawn with slope 3.0 through the lower pressure values. differences between large numbers. Because of the scatter it was not possible to decide unequivocally on a value for n, the degree of polymerization. A value of three seemed best to the author but values of two and four were not unreasonable. This degree of polymerization could probably be derived better from some other type of experiment. Attempts to fit the higher pressure data with equilibria involving BiIz as well as BiI and BiI3 were unsuccessful. At a given temperature it was possible to account for the amount of Bi and I in each sample on this basis; however, it was not possible to fulfill the requirement of constancy of two concentration quotients among the three species from one sample to another. An experiment to observe the species in the vapor 1.0 1.1 1.2 with a mass spectrometer was made. Bi13 was i o 3 / ~ 0K. Fig. 2.-Ilogarithm of the equilibrium constant of reaction 1 passed over a heated Bi surface and then through us. reciprocal of absolute temperature. a small orifice into the ionization chamber of a Bendix mass spectrometer. The temperature of I(g); sal the absolute entropy of Bi13(g) at 640' can be calculated to be 126 f 5 e.v. At present ( 5 ) See T. L. Cottrell, "The Strengths of Chemical Bonds," there are no reliable enthalpies of formation for Butterworthel London, 2nd Ed,, 1958. BO
60
W
(0
March, 1961
HEAT
O F SOLUTION OF
the Bi and the pressure of the Bi13were both varied. Peaks corresponding to Bi13+, Bi12+, BiIf and Bi+ were observed. The relative intensities of these peaks for &I3only were: weak, medium, strong, medium, respectively; and for BiIa passed over hot Bi they were: weak, medium, very strong and strong. The BiI+ and Bi+ peaks increased relative to the Bi13+ as the temperature of the Bi in the cell was increased. There were no peaks at masses greater than 600 so that no polymers of BiI were observed, although the mass spectrometer was capable of detecting masses as great as 1000. Thus the experiment confimed the presence of BiI in the vapor, and its increase as the Bi temperature was increased. It was felt that the polymer of BiI was not observed in this experiment because the vapor did not come to equilibrium with the Bi metal; therefore, the pressure of BiI was much lower than the equilibrium pressure, and that of the polymer was so much lower as not to be observable. It was estimated that the Bi13 in this experiment was in contact with the Bi for only one hundredth the time of contact found necessary in the transpiration experiments, so that presumably the equi-
ORTHOPHOSPHORIC ACID
523
librium concentration of BiI and BiJn did not have enough contact time to be formed. A comparison of the three Bi-halide systems studied shows some regularities. The equilibrium constants for the formation of the monohalide (cf. reaction 1) tend to increase slightly in going from C1 to Br and from Br to I (;.e., a t 600': Kci = 0.011; K B = ~ 0.014; KI = 0,019). This increase is due to a progressively more favorable enthalpy change for the reaction which just barely overrides the progressively less fayorable entropy change. In the iodide system there was evidence for a polymer of the monohalide a t pressures of BiI3 above about 3 mm. In the chloride and bromide systems no evidence for polymers was observed; however, those systems vere investigated a t lower BiX3 pressures. It may well be that a t higher pressures measurable amounts of polymers may be formed. Acknowledgments.-The author is indebted to Mr. William E. Robbins, who performed the Milne, transpiration experiments, to Dr. Thomas -4. who did the mass spectrometer experiment, and to Dr. Francis J. Keneshea, Jr. for fruitful discussions.
HEAT OF SOLUTION OF ORTHOPHOSPHORIC ACID BY EDWARD P. EGAN,JR.,AND BASILB. LUFF Diwision of Chemical Development, Tennessee Valley Authority, Wilson Dam, Alabama Recezved Octobei 6 , 1960
From measurements of the heat, of solution of orthophosphoric acid over the range 0 to 89.13% H3P04, tables were derived relating the relative apparent molal heat content ( 6 ~and ) the partial molal heat contents (12, to concentration a t intervals of 5y0H,PO,. The molal heat of formation from the elements at 25" was calculated as a function of moles of water in solution.
zl)
The heat of solution of phosphoric acid in water was summarized by the National Bureau of Standards' from scattered values appearing before 1915. Another summary tabulatioq2 which drew upon unpublished work by TVA, had among its shortcomings a gap between 35 and 45% H2P04. This paper describes a redetermination of the heat of solution over the range 0 to 89.13% H3P04. Measurements a t higher concentrations would have heen complicated by the presence of non-ortho forms of acid,3with their heats of hydrolysis. Materials and Apparatus.-Reagent grade phosphoric acid was recrystallized twice as the hemihydrate. Stock solutions were prepared from the drained, unwashed crystals. A solution calorimeter from earlier work4,6 was modified in :t few details. The capsule-type platinum resistance thermometer was exposed directly to the solution. The head of the thermometer was sealed into a small glass support tube with Apiezon W wax. The thermometer was suspended in the glass draft tube of the stirrer with the tip of the capsule a few millimeters above the impeller. A 100-ohm heater was wound directly on the outside of the draft tube. (1) National Bureau of Standards Circular 500, U. S. Govt. Printing Office, Washington, D. C., 1952. (2) T. D. Farr, Tennessee Valley Authority, C h e n . E n g . Rept., N o . 8 ( 1950).
T.
(3) E. P. Egan, Jr., and Z . Wakefield, J . Phys. Chem., 61, 1500 (1957). (4) Egan, Jr., B. B. Luff and Z. T. Wakefield, ibid., 62, 1091 ( 1958). ( 5 ) E. P. Egan, Jr.. Z . T. Wakefield and K . Id,Elmora. .I. Am. Chem. Sac, 78, 1811 (1866),
E. P.
Assembly and disassembly of the calorimeter were facilitated by attaching the head to the body by means of six small parallel-jawed spring clamps of a type used in sheet metal assembly. The clamps were completely covered by the water-bath and were without observable effect on the heat leak. The initial bulk charge of liquid for each measurement (850 ml.) was weighed. Each incremental addition to the bulk liquid in the calorimeter was suspended inside the stirrer shaft in a thinwalled glass bulb which was crushed against the bottom of the Dewar to start the solution period. The calorimeter system was calibrated electrically immediately before and after each measurement. One defined calorie was taken as 4.1840 abs. j . Temperatures were recorded to four decimal places, as small differences were important. The conditions of measurement ensured a temperature rise of at least 0.15' for every observation. The temperature at the end of each measurement was 25 5 0.05", and no temperature corrections were necessary. As the water-bath around the calorimeter was held a t 26 f 0.02", heat leaks were always in the same direction.
Heats of Solution.-The concentration range 0 to 83.6 molal (89.13Oj,) &Po4 was covered in five series of measurements: Series 1 : H2@plus successive increments of 75.14yo H3P04; final concentration, 49.4y0 &Pod; Series 2 : same; Series 3: H,O plus successive increments of 43.86% H3P04; final concentration, 17.0% H3P04; Series 4 : 44.58y0 plus successive increments of 89.13y0 HaPod; final concentration, 75y0 H3P04; Series 5: 89.13% H31'04 plus successive increments of HzO; final concentration, 727, &Po4. Incremental
