Sept., 1982
THERMODYNIhIIC PROPErZTIES AND PHASE
pearance potentials are 199 and 201 kcal./mole, respectively. These values are higher than that of 173k ~ a I . / m o l ebut , ~ ~agree with those determined from the ethylene oxide and propylene oxide study2 and the above study of 1,2-epoxy-3-rnethoxypropane. mle = 42.-CzH20+ is formed from epichloroC1. hydrin with the neutral products CH3 ANf+(Czl120)is calculated to be 191 kcal./mole. This is somewhat lower than values reported in the litera bure.* , 1 1 m/e = 49.-The only possible ion is CHsC1+, since from Fig. 1 it can be seen that the ratio of 149/151 is about 3. But the energetics do not
+
RELATIOM IN
IIYDROGEN-~I,4FYIVM SYSTEM
1657
+
distinguish between the neutral products GO CHa and C2Ha0. AHf+(CH2C1)is calculated to be 226-236 kcal./moIe. For the present, we tend to favor the greater value. The heat of formation for this ion, to our knowledge, has not been reported previously. The correspoiiding ion jn epibromohydrin was not observed. mle = 57.-This ion is formed by removal of halogen from the parent molecule-ion in both epichlorohydrin and epibromohydrin. AHff(C8H50) are 207 kcal./mole for epichlorohydrin and 216 kcal./mole for epibromohydrin. Although in fair agreement, these values are a little higher than that reported from a study of propylene oxide.2
THERMODYNAMIC PROPERTIES AND PHASE RELATIOSS I N THE SYSTEM HYDROGEN-HAFNIU&!11,2 BY RTJSSELL K. EDWARDS AND EWALD VELECKIS3 Department oaf Chemistry, Illinois Institute of Technology, Chicago 16, Illinois Received iMaich IO, 1962
Isothermal equilibrium pressure measurements were carried out as a function of composition for the hydrogen-hafnium system in the temperature range 251 to 872" and up to a pressure of one atmosphere. The proposed partial phase diagram indicates the existence of three homogeneous phases. In the primary solid solution phase ( a ) the solubility of hydrogen reaches nearly I 1 atom % hydrogen a t the highest temperature investigated but decreases considerably toward the lower temperatures. Conformance to Sieverts' law was established throughout the 0-phase; the temperature variation of the Reverts' law constant can be expressed by the equation log ( Z / = / N n ) = 3.820 - 1964T-', where P,, is the pressure of hydrogen and N E is the atom fraction of hydrogen in the condensed phase. The second phase (6) is homogeneous from a minimum icomposition of 24% a t the highest temperature to a maximum of about 64% a t lower temperatures. In the pressure range studied this latter boundary remains essentially constant with temperature. A very narrow two-phase region separating the &phase and a third phase (E) is inferred with the latter extending to HfHl.9, for our lowest temperature. The extrapolated phase boundaries show a good correlation with the room tem erature X-ray studies of Sidhu and McGuire. The relative partial molal and integral heats and entropies are presented for t i e composition range 0 to 57 atom % hydrogen. The heat and entropy of the a -+ 6 reaction are - 15.66 0.84 kcal./g.-atom H and - 12.34 =t 0.28 cal./deg. g.-atom H; for the 6 -+ B reaction the respective quantities are - 13.39 =k 0.21 kcal./g.-atom H and -20.64 =t 0.16 cal./deg. g.-atom H.
Introduction The system hydrogen-hafnium has been studied by Sidhu and ~o-workers,~ who made a thorough investjigation of phase and structural relations by metallographic and by X-ray and neutron diffraction techniques at room temperat'ure. 14hexagonal primary solid solution was found to t'ransform into a tetragonal (deformed cubic) phase in the range from 'A,
8il.!) 826.8
10. li I !I, 95 9.20 8 .G!l 7.55
23.0 28.0
.. ..
32.2 34.4 38.0
..
7%,5 77!1. 4
7'15.2 717.8
6.92
GR2.7
6.15 3.55
594.8 396.5 365.1
322. 0 297.9 271.6
..
.. .. .. ..
..
50.1 42.3
47.R 56.2
.. .. .. .. ..
..
1.1/11:11
1
il,.
i>,l,l.
1I 'Z,
52.1 55.1 5li.5 57.3
.. ..
58.5
.. ..
50.!l G1.5 64.0 64.5 65.1 G5.9 66.1 66.3
.. 64.2 63.7 64.2 64.0
250.7 64.2 Compositions in atom yo hydrogen.
5!1 . 3
roliili v l
qiiiirit
itics wporl.d.E
l'hc st,:iuid:ird Iii~iila l i i 1 (ml r q i y o1'fi~rin:iI,i~i11, All, :iid A S f , c ~ ~ r r ~ y ~ t,liv ~ i fnrmal~inn n ~ l (if on(: g r m i :itom of sohilion froin solid tinfniiini inid g:isciiiis di:rl.ninic hydr~igcn:it. a prcssurc (if I ;rl,in. Tho d:Aa beyond 57 atom 2) hy:Ir(igiw iirc i:nnsidcrcd too uncertain frnm ii st~:indpriiiil. n S phase boundnrics and from ari cxcccdiirgly sansit,ivc prcssrirc variation t,o marrant reporting nf thc complete thermodynamic results. Discussion The phase iliagrams of t,he systems 1%-Ti and 11-Zr arc rrgardcd' to h e of the eutectoid type. The I-I-Hf system would be expected to show an analogous behavior. The partial phase diagram of Jig. 2 does not exclude the possibility of having a eut,eetoid transformation outside the temperature ( 6 ) For a discairsion of the methods involved in the evaluation of similar thermodynamic data see, c.o., C. Wagner. "Thermodynamios of Allow," Addison-Wesley Pres. Ino., Cambridge, Mea% 1952. Chap. I. (7) A review of Literature on the H-Ti and H-Zr swtema may bc found in M. FIanscn and K. Anderko, "Constitution of Rinsry Alloys," MoGmw-Hill Rook Co.. Ine.. New York, N. Y..1958. PD. 799.808.
1660
TABLE II THER~MODYNAXIC QUANTITIES Atom % H
- (NH
- '/doad,
kcal./g.-atom H
FOR THE SYSTEM
HYDROGES-€~AFNIVM -&€If:
- @H - ' / I S " H d ,
cal./deg. g.-atom H
kcal./g.-atom
-
4&,4
caI./deg. g.-atom
0.00 2.00 4.00 6.00 8.00 8.69"
-cc 8.99 f 0.8 0 0 8.99 f .8 3.12 f 0 . 3 0.N 0.023 8.99 f .8 4.49 f .3 .35< 8.99 f .8 5.30 f . 3 .53< f 9.5% 8.99 f . 8 5.87 f . 3 ,711 ,317 6.04 f . 3 8.99 f .8 .781 .359 .. 12.34 f .3d 15.66 f . 8 d ... ... 34.40 15.19 f 1 . 2 11.95 f . 5 4.97 3.73 37.00 15.69 f 1.1 12.45 f . 4 5.39 4.07 38.00 12.65 f .4 15.88 f 1.1 5.55 4.20 39.00 12.90 f .4 16.13 f 1 . 0 5.72 4.34 13.20 i . 4 40.00 16.42 f 1 . 0 5.89 4.48 13.51 f . 4 41 .OO G.07 16.73 f 0 . 9 4.63 42.00 6.25 17.03 f .9 13.83 f .4 4.78 f 2.6% 1 4 . 1 4 2 ~ .4 6.44 17.32 f .9 43.00 4.94 14.48 f .4 44.00 6.64 17.62 f .9 5.11 18.86 f . 9 45.00 15.67 f . 3 6.84 f 5.6% 5.29 46.00 16.01 f . 3 7.06 19.15 f .8 5.48 47.00 16.41 f . 3 7.29 19.50 f .8 5.68 48 ' 00 16.78 f . 3 7.52 19.79 f .8 5.88 49.00 20.02 & . 6 17.11 f . 2 7.76 6.10 20.19 f .8 17.39 f .3 50.00 8.00 6.32 17.63 f .3 51 .OO 20.29 f .9 8.25 6.54 2 0 . 5 6 4 ~ .7 8.50 52.00 18.08 f . 3 6.77 18.00 f . 3 20.26 f .8 8.74 53.00 7.00 18.13 f .2 8.98 2 0 . 1 4 f .7 54.00 7.23 18.70 f . 3 9.22 7.48 20.42 f .8 55.00 19.05 f . 6 9.47 56.00 20.43 f 1.3 7.73 9.73 19.82 f 1 . 0 20.75 f 2 . 0 57.00 8.00 a Concentration a t the a - ( a 6 ) boundary for an arbitrarily selected tenqmaturc, 1052OK. b Concentration a t hhe (a 6) - 6 boundary for the same temperature. In the single phase regions the integral heats and entropies show 110 variation with temperature within the experimental error. However, because of the temperature dependence of Lhe phase boundaries, due regard must be paid in the application of these data for temperatures other than 1052'K. It is to be noted that, whereas the second and third columns list relative partial quantities, the particular nunibers 15.66 and 12.34 reported for the two phase region correspond instead to the heat and entropy for the reaction
+
+
xhrre r' arid 1"' :irca t h utviii ratios (kl/Hf) a t the were left in the table for convenience.
a
- ( a + 6) ~
rarigr of this study (Espagno, et ~ l . e, ~ . fouiid , evidence for a phase transformation a t -100'). The closure of the ( a 6) miscibility gap, suggested by the rapid increase in mutual solubilities of the a and &phases with temperature, is not likely because of the presumably different crystalline structures of these two phases. The two-phase region at 950" extending from 16.7 to 33.3 atom yohydrogen reparted by Espagno, et a1.,5 is wider than would be expected from an extrapolation of our data to that temperature. In view of a perceptible temperature dependence of composition in the plateau portion of their isobar, we feel that their data would permit narrower limits. The second heterogeneous region observed by Espagno, et ~ 1 . ~in5 their 1 atm. isobar (-400" and -63 atom % hydrogen) is in good agreement with our phase diagram (376" aiid -64%). The extrapolation o i thc phase boundaries toward the lower temperatures is rather arbitrary. ?;evertheless, the correspondence with the room
+
n (da
+ 6) - 6 boundaricb, rcspectively.
The numbers
tcnipcrature limits of Siclhrr, ot ~ , l . it,, ~rt:uboiral)le and aids in the slructural idcntification of the iiilermediate phases. Thus the a-phase would have the hexagonal structure of the primary solid soltition of hydrogen in a-hafnium. The 6-phase would correspond to the tetragonal (deformed cubic) or the face-centered cubic structures, assuming that these two structures actually represent a single phase of continuous distortion. The €-phase would correspond to the face-centered tetragonal structure. Neutron diffraction studies of metal-hydrogen solid phases have gone far in clarifying the sometimes ambiguous interpretation of the role of hydrogen in metals. Rundle's8 stressing of the importance of metal-metal bonds in the interstitial compounds points also to similar considerations in the metal-hydrogen systems. Sidhu, et have shown that significant structural changes take place in the hafiiium metal lattice as hydrogen is introduced. It is clear that the (8) R. E. Rundle, J . Am. Chem. Soc., 69, 1719 (1047).
MERCURY (11) HALIDE MIXEDCOMPLEXES IN SOLUTION
Sept., 1962
entrance of hydrogen brings about disturbances of the *3!-M bonds. Some of the changes show lengt!henirig of some M-RII bond classes along with a fihorteiiing of other M-M bond classes. The hexagonal structure of pure hafnium is of lion-ideal axial ratio. Since initial directional preferences are present in the pure metal, it is not surprising to see the distortion which is met in the face-centered cubic structure. One can infer that the distortion of the M-M bonds by hydrogen is itself endothermic since the sign of the relative partial molal enthalpy of hafnium is positive. Both deformation and energetics indicate that the process is much more than a simple filling of interstitiall holes. The fact that diatomic hydrogen enters a number of metals exothermally while a t the same time undergoing dissociation has always seemed somewhat startling from an energetic point of view to those thinking in terms of a simple “solution” of a gas in a solid. The bonding of a hydrogen atom to its surrounding metal atoms is thus
1661
even stronger than its covalent bonding to another
hydrogen atom in diatomic hydrogen. Sidhu, et u E . , ~ have stated that the M-H bonds are stronger than AI-31 bonds, but this is misleading. The important factor i s that the hydrogen atoms within the interstitial positions yield additional bonding beyond the normal &I-AI bonding. That is, due to the small size of the hydrogen atom, one gains 11-H bonding while retaining most of the &AI bonding. If the assumption is made that the hydrogen bonding energy is divided among four bonds associated with its four nearest metal neighbors, use of the thermodynamic data here obtained, the dissociation energy of diatomic hydrogen, and the sublimation energy of hafnium leads to a greater energy for the AI-31 bond than for the 31-H bond. Acknowledgment.-The support of the ONR and AFOSR during the course of this study is gratefully acknowledged. We wish to thank Prof. S.E. Wood for valuable discussions.
MERCURY(I1) HALIDE MIXED COMPLEXES I N SOLUTION. V. COMPARISON O F CALCULATED AKD EXPERIMESTAL STABILITY CONSTANTS1 BY Y. MARCUS AND I. ELIEZER Israel Atomic Energy Commission Laboratories, Rehovoth, Israel Received March 10,1961
The equililiriuni constants for the formation of mercury( 11) halide ternary (mixed ligand) complexes from the parent binary complexes have been calculated on the basis of a “polarized ion” model. The values obtained agree fairly well with the experimental results available.
Introduction an extent the experimental results can be explained The stability of ternary complexes MA& as by applying theoretical consideratioils along the comparcd with that of the binary complexes MA, lines mentioned Theory, a. Definitions.-We can write for and MB,, where n = i j , has not been studied much as yet but the basis for its theoretical treat- the formation of the mixed complex from the ment was laid by Bjerrum in his study of the ratio parent complexes hctween consecutive formation constants of binary Bjerrum divided the factors influcnciiig the complex formation constant’s into a “statistical effect” and a “ligand effect” further subequilibrium co~is t:in t f< If (1) dividing the latter into an “electrostatic effect” and a “resit effect.” Of late Kida has discussed Using thr coiivciitional10over-all stability coiistants some of the above factor^.^ ,B one obtains The mixed complexes of mercury@) with C1, Br, and I have been thoroughly investigated by one K M = Pi, X X Pori-'/'' (11) of while very recently Hume and Spiro8 Lct us now analyze KRZsomewhat similarly to studied spectrophotometrically the uncharged mixed mercury halides confirming the results ob- Bjerrum’s ideas. We can write tained in ref. 5 . We have tried to ascertain to what log KM = log &tat log K i log KR (111)
+
(1) Presented at the 7th International Conference on Coordination Chemistry, Stockholm, June, 1962.
(2) J. Bjerrum, “Metal Ammine Formation in Aqueous Solution,” P. Haase & Sons, Copenhagen, 1957. (31 S. Kida, Bull. Chem. Soc. J a p a n , 34,962 (1961). (4) Y . Marcus, Acta C h e k . Scand., 11, 329 (1957). ( 5 ) I’. Marcus, ibzd., 11, 599 (1957). (6) Y . Marcus, ibid., 11, 610 (1957). (7) Y . Marcus, ibid., 11, 811 (1957). (8) T. 0.Spiro a n d D. N. Hume, J . Am. Chem. Soc., 83, 4305 (1961).
+
+
where K s t a t = the value of K Mif formation of the mixed complex proceeds statistically; K,I = the stabilization constant of the mixed complex due to the electrostatic effect; K R = any additional stabilization, which Bjerrum calls the rest effect. Y.Marcus, Bull. Res. C o m c d Israel, IOA, N o 3 , 2 (1961). 110) J. BJertum, G bchmaraenbach, a n d L G Sill6n “Stab~lity Conhtants ” The Chemical Society, London, 1958. (9)
