Dec. 5, 1056
HEATOF
PRECIPITATION AND
drolysis of samarium nitride three additional DTATGA runs were made. I n the first, Sm(OH)3 was precipitated, washed, and vacuum dried a t room temperature; the weight and DTA curves were largely without character. In a second run the Sm(OH)3 was precipitated, washed and dried a t 110' for six hours. An X-ray powder diagram of the dried material showed a complex lattice of large cell dimension quite different from the hexagonal structure of the Sm(OH)3produced by hydrolysis of SmN. A DTATGA run on this material was in general aspect similar to the vacuum dried material and both had a large amount of water, Sm(OH)3.2.3Hz0 and Sm(OH)3.2.8H20, respectively. A third preparation similar to the second but from which almost all the water (Sm(OH)B. 0.12 HzO) had been removed showed the same hexagonal structure reported above and gave DTATGA curves very similar to the ones observed from the hydrolyzed SmN. That is, a sharp break in the weight curve a t 375" with DTA peaks corresponding to the initial endothermic decomposition of Sm(OH), and of a composition around SmO(0H). The Sm(OH)3.2.8H20 and other highly hydrated materials apparently have water incorporated within the crystal lattices. This water is completely lost only a t temperatures far above the decomposition temperature of Sm(OH)3, hence, the
[CONTRIBUTION FROM
THE
FORMATION OF SILVER MOLYBDATE
5989
latter decomposes rapidly without indicating the low molecular weight hydrated material. All the samples of EuN contained a trace contaminant exhibiting a complex powder diagram which was believed to be an oxide. Because of this impurity and the scarcity of europium, no additional work was done on EuN. Samples of YbN prepared by direct combination of the metal with nitrogen always contained some unreacted metal. Even a t lOOO', 200' above the melting point of ytterbium, the reaction did not go to completion. When these samples containing metal and YbN were distilled in a high vacuum mm.), both ytterbium metal and YbN were found in the distillate. A complete distillation occurred if sufficient time was allowed a t 1400". This may imply a high vapor pressure for YbN. SmN heated to 1600' did not exhibit any tendency to distil. This suggests a considerable thermal stability and low vapor pressure. Samples of YbN prepared by the reaction of YbH2 and N2 a t 1000" were monophasic. At lower temperatures conversion was not quantitative. Acknowledgment.-The authors wish to thank Miss Shirley Curtis for help in measuring the powder diagrams and in performing some of the calculations. IOWACITY,IOWA
COBB CHEMICAL LABORATORY, UNIVERSITY OF VIRGINIA]
The Heat of Precipitation and Formation of Silver Molybdate BY C. N. MULDROW, JR., AND L. G. HEPLER RECEIVED JULY 23, 1956 The heat of precipitation of silver molybdate has been determined from two series of calorimetric experiments. From these data and the results of earlier investigators, we calculate the heat of formation of ,4g2Mo04 to be -200.4 kcal./mole and the free energy and entropy of solution of AgzMoOa to be f15.6 kcal./mole and -9.4 cal./deg. mole, respectively,
As part of a systematic investigation of the thermochemistry of molybdenum, it was desired to have reliable values for the heats of formation and precipitation or solution of one or more slightly soluble molybdates. Therefore, the heat of precipitation of Ag2MoOl has been carefully determined, and from this quantity several thermodynamic calculations have been made. Experimental
excellent.394 All heats of reaction have been carried out a t 25.0 i 0.5".
Results The heat of reaction of crystalline NazMoOd with dilute solutions of AgN03 was measured to determine the heat of precipitation of AgzMoO,. The calorimetric reaction may be written as eq. 1
+
NaZMoOa(c) -j- AgNOa (excess dil. s o h ) = AgZMoOd(c) AgX03 (dil. s o h ) Nar\'Ot (dil. s o h ) AH1 (1)
+
The high-precision solution calorimeter used in this investigation has been previously described in some detail.' The silver nitrate was Mallinckrodt Analytical Reagent crystals. Fisher Certified Reagent sodium molybdate (NasMoOp) was dried for 24 hours a t 110" and found by analysis to be 46.77, Mo as compared with the calculated value of 46.6%. The analysis was done by the permanganate method.* The calorimeter used in this investigation has been checked periodically by determining the heat of solution of potassium chloride.' Agreement with reported values is
The results of these experiments are given in Table I. The heats were actually determined to the nearest 0.02 kcal./mole, but because of the spread of values they are reported only to the nearest 0.1 kcal./mole. All determinations were made using 050 cc. of solution. T o calculate the standard heat of formation of Ag2Mo04from the experimental values of AHl, it is necessary to know the heats of formation of Na2Mo04(c),I NaN03(aq)3and AgN03(aq).3 We
(1) R. L. G r a h a m and L. G. Hepler, THIS JOURNAL, 78, 4846 (1956). (2) W . F. Hiliebrand, G. E. F. Lundell, H. A. Bright and J. I. Hoffman, "Applied Inorganic Analysis," Second Edition, John Wiley a n d Sons, Inc., New York, N. Y.,1983,p . 307.
(3) "Selected Values of Chemical Thermodynamlc Properties," Circular 500, National Bureau of Standards 1952 (4) P. H Spedding and C F Mlller THISJ O U R N A L , 74, 3158
(1952).
C. N . MULDIWW, JK.,
5990 'I'ABLEI
OF Na2Mo04(c) NanMoOd, moles
011 IfoOi liroles
0.04201 .OK257
0 .008577 .11086l14 , I IOBOOG .lil0767
,06955 ,08364
~ ' I T I ISa?nioOn(aq)
H e a t absorbed Heat absorbed minus heat of (Cd.)
soln. AgNOj
116.8 217.0 300.3 300. 1
-109.0 -119.0
- 73.2 - 149.0
Values obtained iii this Laboratory for the heat of solution of AgN03 agree well with the heat calculated from the Bureau of Standards data.3 AS with the first series of determinations, heats of dilution were ignored in making calculations. Because dilute solutions were used, the heats of dilution are again small and also tend to cancel each other. In view of the final uncertainty of several tenths of a kilocalorie and the lack of pertinent heats of dilution, ignoring these small heats is acceptable. The heat absorbed in dissolving the excess solid AgNOs is calculated from Bureau of Standards d a h 3 This heat is subtracted froin the total measured heat to give the heat of thc precipitation reaction. From the experimental heats of reaction 2 and the previously mentioned h a t s of formation, the heat of formation of Rg2MoOa(c) has been calculated to be -200.7 kcal./inole with an average deviation of 0.6 kcal./mole. Discussion of Results Britton and German5 have shown that normal alkali molybdates react with silver nitrate to give (.j) IT
'r
S Rr-itton aud \I7 I.. ( k r m a i i , J Cheiii. Sor., 11.>D (1113*).
AND
L. C. HEPLEK
Vol. 7s
normal silver molybdate. Every precipitation in the preseiit investigation was carried out with ex cess silver nitrate to prevent formation of polymolybdates. However, some slight formation of polymolybdate may be responsible for the fact that deviations from the average are considerably larger than the uncertainty in any one run. Several earlier investigations of AfiAIoOd arc related to ours. Ricci and Link$ have tleterinined thc solubility of Xg2Ll-00~over a range of temperatures. Seglecting small hydrolysis effects, we have calculated from their data therniodynamic solubility products a t several temperatures. n'ith the van't Hoff equation and these cquilibriuni constants we calculate the heat of solution of AIg:,hloOdto be +12 C, kcal./niole. This heat atld the heats of foriiiatioii of .Ig+(:tq)j and lIoOi= (aq)' give -200.2 O.(j kcal./rnole as the heat of formation of Ag2h100d. Pan7 has investigated .Ig2Mo04 with a silversilver rriolybdate electrode and has reported values for the standard free energy and heat of solution of ,\g?M004. Froiii his results and our recently deteriiiined heat of formation of aqueous niolybdatc ion,' we calculate that the heat of formation of .ig2MoO4(c) is -200.5 kcal. mole. I t is difficult to estimate the uncertainty of this value, but it is certainly fairly small. Prom our two series of determinations of tllc lieat oi precipitation of .IgL1\Io04and froiii our calcu1,itions based on Ricci and Linke's wluhility wori,'l ancl on PLtn'se.ii1.f. work,' we take -200. I i 0.2 kcal./mole ds the best value for the heat of foriiiation of Xg211004(c). Pan's7 free energy of solution of , i g ~ \ l o Oand ~ that calculated from Ricci and Linke'sGsolubility work a t 23" agree fairly well. Ricci and Linke also discuss carlicr determinations of the sohbility. IVe take the best value of the free energy of solutioii and the solubility product a t 25' to be Jrl5.G kcal 'mole and 2.S x 10-12, respectively. From this free energy of solution and the heats of formation of Xg2Mo04(d, Agf(aq)3 and l1004= (aq),' we calculate the entropy of solution of Ag2i\1oO4to be -9.4 cal. /deg. mole. Determination of the entropy of ,lg&IoO4(c) from heat capacity ineasurenicrits and the Third Law would thus make i t possible to calculate the entropy of -21004=(aq). The above entropy of solution and that calculated from estimated entropies? of Xo04-(aq) and .Ig$UoO1(c) and the bnown u i tropy ~ i.Ig+(aq) ' 'ire about the ~ i i n e . Acknowledgments. I\-e arc ple,ised to e q i r m our gratitude to Llr. I
