Interaction of the Tetraethanolammonium Ion with Water as

Chem. , 1966, 70 (9), pp 2974–2980. DOI: 10.1021/j100881a042. Publication Date: September 1966. ACS Legacy Archive. Cite this:J. Phys. Chem. 1966, 7...
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D. F. EVANS, G. P. CUNNINGHAM, AND R. L. KAY

2974

Interaction of the Tetraethanolammonium Ion with Water as Determined from Transport Properties

by D. Fennel1 Evans, G. P. Cunningham, and Robert L. Kay' Mellon Institute, Pittsburgh, Pennsylvania

16615 (Received April

4, 1966)

Precise conductance measurements are reported for tetraethanolammonium iodide, (EtOH)*NI, in water a t 25" and for the bromide a t 25 and 45", as well as viscosity B coefficients from measurements on aqueous solutions at concentrations up to 0.1 M for the bromide at temperatures ranging from 0 to 65". This substitution of a hydroxyl group for a terminal methyl group in the Pr4N+ion is shown to result in the elimination of the effects that have been attributed to water structure enforcement around the side chains of the large tetraalkylammonium ions. This interpretation is shown to be consistent with existing conductance data for the Mez(EtOH)zN+ and the Me3(EtOH)N+ ions and their correct alkyl analogs. The number of carbon atoms in the side chains is shown to be a poor criterion for comparison of various quaternary ammonium ions since water structure enforcement does not commence significantly untiI the chain reaches the size of a propyl group. The concentration dependence of the tetraethanolammonium salts is compared with that for the tetraalkylammonium salts.

Introduction Recent conductance2 and viscositya measurements for the large tetraalkylammonium ions at 10, 25, and 45" in aqueous solution have demonstrated that t,he transport behavior of these ions is influenced to a significant extent by the enforcement of water structure about their hydrocarbon chains. This structure has been attributed to the inert and hydrophobic nature of the hydrocarbon side chains. The explanation generally given for this phenomenon is that water molecules situated a t t,he surface of these large ions will be oriented very little either by the ionic charge or the nonpolar side chains on its one side. Consequently, such water molecules can be oriented to a great,er extent than normal by their nearest water molecule neighbors and can in effect be oriented into favorable positions for the formation of water cages about the inert side chains. Although the water in these cages will be exchanging very rapidly with bulk water, the net effect will be an increase in the size of the ions involved as well as an increase in the local viscosity about the ions. Both effects decrease the ionic mobility, the effect being greater the lower the temperature. At the same time, this explanation accounts for The Journol of Physical Chemistry

the viscosity increase resulting from solution of these ions that is far greater than can be accounted for by their size a10ne.~ This viscosity increase is also very temperature dependent and disappears rapidly as the temperature increases.a This explanation suggests that if the nonpolar hydrocarbon chains were made polar by substitution of a hydroxyl group in place of the terminal methyl group, the water structural effects causing the abnormal transport properties of these large ions should disappear. We have verified this hypothesis by investigating the temperature dependence of tshe transport properties of the tetraethanolammonium ion in aqueous solution. This conclusion is in conflict with those reported from measurements on the lower homologs of (EtOH)4N+, namely, the (EtOH)J4ezN+ and (EtOH)lVesN+ * (1) To whom all correspondence is t o be addressed. (2) (a) D. F. Evans and R. L. Kay, J. Phys. Chem., 70, 366 (1966); (b) R. L. Kay and D. F. Evans, ibid., 70, 2325 (1966). (3) R. L. Kay, T. Vituccio, C. Zawoyski, and D. F. Evans, ibid., 70, 2336 (1966). (4) H. 5. Frank and W. Y. Wen, D ~ s ~ ~ s s i Faraday ons Soe., 24, 133 (1957). (5) J. Varimbi and R. M. Fuoss, J. Phya. Chem., 64, 1335 (1960). (6) H. 0. Spivey and F. M. Snell, ibid., 68, 2126 (1964).

INTERACTION OF THE TETRAETHANOLAMMONIUM ION WITH WATER

ions. Both conclusions are based on a comparison of mobilities of the ethanol derivatives with the tetraalkylammonium ions. We have resolved the discrepancy by noting that such comparisons must be made with the correct tetraalkylammonium analog. Wen and Saito7 have measured activity coefficients and partial molar volumes of (EtOH)4NBr and (EtOH)4NF in aqueous solution at 25" and have found that these thermodynamic properties are more normal in the case of these salts than was found for their tetraalkylammonium analogs. Also, Price8 has reported that the introduction of hydroxyl groups into the tetraalkylammonium ions reduced the heat of transport significantly, a direction in keeping with less rather than more order in the solution.

Experimental Section All electrical equipment, cells, salt-cup dispensing device, and general techniques for the conductance measurements were the same as previously described. The modification required for the 45'2b conductance determination and the method employed for weighing hygroscopic sa1tsl1 have already been reported. The cell constant was determined at 25" and calculated for 45", the change amounting to less than O.Ola/o.zb The conductance baths were set at 25 i 0.003' and 45 f 0.007' with a calibrated platinum resistance thermometer. The viscosity measurements were carried out using a suspended-level Ubbelohde-type viscometer with a flow time of 500 see. No kinetic energy correction was found necessary at any temperature. The experimental techniques were the same as those previously d e ~ c r i b e d . ~ Tetraethano1;tmmonium bromide, (HOC2H4)4NBr, was prepared hy the method described by Wen and Saito' with a number of modifications; 0.8 mole of 2bromoethanol (Fisher Scientific Co.) was refluxed with 1.6 moles of triethanolamine (Fisher Scientific Co.) in 300 ml of n1eth:tnol for 24 hr. This reaction gives approximately equal amounts of (EtOH)4NBr and (EtOH)3NHBr, and it is the separation of these two products which presents the greatest difficulty in the synthesis of (EtOH),NBr. The reaction mixture was titrated with concentrated aqueous hydrobromic acid to pH 3 to convert the remaining free amine to (EtOH)3NHRr. The resulting (EtOH)3NHBr crystals were removed by filtration and the filtrate was taken to dryness by azeotroping off the last traces of water with ethanol under reduced pressure. The amine hydrobromide is only sparingly soluble in methanol, whereas (EtOH)4KBr is exceedingly soluble at temperatures as low as - 20". This property permitted further separa-

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tion of the mixture since only the hydrobromide precipitated on cooling a saturated methanol solution to -20'. The white, nonhygroscopic hydrobromide was detected by its sharp melting point of 187". Final purification of (EtOH)4NBr was effected by dissolving the salt in the minimum volume of methanol and adding four volumes of absolute ethanol. White crystals precipitated on slow cooling with periodic shaking. This recrystallizing procedure was repeated 12 times. Recrystallization from absolute ethanol produced an oil which completely solidified even on slow cooling. Owing to its hygroscopic nature, all recrystallizations and manipulations of (EtOH)4NBr were carried out in a drybox. The powdered salt was dried under vacuum at 56' for 15 hr. The purified compound melted at approximately loo",with decomposition. In fact, this salt appeared to undergo decomposition readily at temperatures above 80', and melting points taken on samples in evacuated sealed tubes gave the same results. Upon standing for 1 month, this salt showed some signs of decomposition, as indicated by the insoluble residue found when an old sample was dissolved in methanol. Our conductance results also confirmed this observation since measurements on month-old samples gave limiting conductances that were somewhat higher than those obtained using freshly prepared salt samples. However, the conductance parameters obtained from freshly recrystallized samples were shown to be independent of further recrystallization. All attempts to obtain the (EtOH)4NBrby titration of (Et0H)dNOH (97.8010,RSA Corp., New York, N. Y.) gave a 90% yield of (EtOH)3NHBr,but we could detect no (EtOH)4NBr. (EtOH)4NI was prepared from the purified bromide by ion exchange. When all free amine had been removed from reagent grade anionic exchange resin in the hydroxyl form by repeated washing with water, the resin was converted to the iodide form with KI. One pass through a tenfold excess of resin converted (EtOH)4NBr to the iodide. The only test performed on the final product to test for complete exchange was to note that the final product was not hygroscopic, whereas the bromide is extremely hygroscopic. Conductivity grade water2&was prepared by passing (7) W.Y.Wen and S. Saito, J. Phys. Chem., 69, 3569 (1965). ( 8 ) C . D. Price, Technical Report, Armed Services Technical Information Agency, AD 276280 (1961). (9) J. L.Hawes and R. L. Kay, J. Phys. Chem., 69, 2420 (1965). (10) D. F. Evans, C. Zawoyski, and R. L. Kay, ibid., 69, 3878 (1965). (11) R. L. Kay, C. Zawoyski, and D. F. Evans, ibid., 69, 4208 (1965).

Volume 70,Number 9 September 1986

D. F. EVANS, G. P. CUNNINGHAM, AND R. L. KAY

2976

!-

I

I

I

I

I

1 - 1 Table I : Equivalent Conductances in Aqueous Solution 104c

0.1 +/C''2

A

104c

A

(EtOH)*NBr, 25"lO'~0

0

I

I 01

11.991 24.631 35.153 44.402 54.642 65.980 75.307 87.613

I

I

02

0.3

c"2

Figure 1. Plots of eq 1 for (EtOH)4NBr in aqueous solution a t various teniperatureh.

distilled water through a 4-ft mixed-bed ion-exchange column.

Results The density increments for the volume concentrations and viscosity measurements were obtained by direct measurement on 0.06 M solutions. The 0 value in the density equation, d = do 6@z, where rTz is the concentration in moles per kilogram of solution, was found to be 0.098 at 25 and 45" for the bromide and was assumed to be constant in the temperature range 0-65" used in the viscosity measurement. The value of 0 was assumed to be 0.03 higher for the iodide in keeping with previous experience.2a The viscosity data for (EtOH)JXBr are given in graphical form in Figure 1 and can be seen to conform to the Jones-Dole equaticn12

+

+/C"

=

A

+ BC"'

(1)

where $ =- (q/vo) - 1. The intercept A = 0.008 is identical with the value calculated from the Falkenhagen equation13 within the precision of the measurements. It can be seen that B = 0.32 f 0.02 is independent of temperature within the stated precision. The small amount of spread in the points in Figure 1 can be attributed almost entirely to the temperature dependence or the bromide ion that has been shown to have ionic B values3 that vary from -0.08 to -0.01 in the temperature range 0-65", respectively. The measured equivalent conductances and corresponding concentrat,ion in moles per liter are given in , specific conductances of the Table I along with K ~ the solvent. Conductance parameters in Table I1 were obtained from the Fuoss-Onsager conductance equation14 -1 = .io - SC'"

+ EC log C + (J - B&)C

102.31 101.02 100.17 99.54 98.92 98.28 97.83 97.25

1 0 7 ~=~

10.665 23.891 37.560 49.757 62.286 73.549 84.532 98.814

1.9 102.47 101.03 99.97 99.18 98.46 97.89 97.38 96.77

A

-(EtOH),NBr, 45'107% = 3.2 11.358 146.40 23.570 144.54 31.423 143.60 41.746 142.53 51.758 141.62 61.643 140.80 70.911 140.09 80.652 139.42

(EtOH)aNI, 25' 1.3 1 0 7 ~= 1 . 5 7.415 101.48 7.887 101.52 17.570 100.23 17.489 100.20 26.910 99.30 27.366 99.27 36.716 98.49 36.063 98.57 97.94 44.026 98.01 44.803 53.101 97.39 52.889 97.44 61.330 96.90 61.858 96.92 69.680 96.44 71.445 96.41 1 0 7 ~=~

tances, calculated on the basis of X " ( B T - ) ~= ~ ~78.22 and X0(I-)250 = 76.98, are ho[(EtOH)4N+]2p = 27.07 from the bromide and 26.87 from the iodide. Considering the problems of stability of these salts and the extreme hygroscopic nature of the bromide, this agreement is entirely acceptable. At the higher temperature, a Xo[(EtOH)4T+]450= 40.08 is obtained from XO (Br-)4ae = 110.69.2 The d parameter given in Table I1 for (EtOH)4NBr at 25" is in good agreement with the values 1.6, 1.86, and 1.8 reported2 for Pr4NBr at 10, 25, and 45") respectively, and the value at 45" of 2.41 is not too different. On the other hand, the average d value for (EtOH)4NIof 1.3 reported here is substantially higher than the 0.1, 0.3, and 0.4 obtained for Pr4T\IIat 10, 25, and 45", respectively.2 I n the case of P r 4 x I , the conductance data analyzed for a small amount of association. The same analyses of the data for (EtQH),N\'I detected a small amount of association in the first run but none in the second.

Discussion In t'he absence of any change in the interaction with the solvent, replacement of a terminal methyl by a

(2)

using a least-square computer analysis. The values 0.8903, 0.5963, 78.38, and 71.51 were used for the viscosity in centipoise and dielectric constant of water at 252a and 45°,2brespectively. The limiting ionic conducT h e Journal of Physical Chemistry

= 1.6

104c

(12) G. Jones and M. Dole, J. Am. Chem. Soc., 51, 2950 (1929). (13) H. S. Harned a n i B. B. Owen, "The Physical Chemistry of Electrolytic Solutions, 3rd ed, Reinhold Publishing Gorp., New Tork, N. T., 1958, p 240. (14) R. M. Fuoss and F. Accascina, "Electrolytic Conductance," Interscience Publishers, Inc., New York, N. Y . , 1959.

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INTERACTION OF THE TETRAETHANOLAMMONIUM IONWITH WATER

Table 11: Conductsxe Parameters in Aqueous Solution a t 25 and 45" d

do

(Et0H)aNBr

25

(Et0H)aNI

45 25

105.30f0.006 105.27 & 0.01 150.77 f 0 . 0 1 103.92 f 0 . 0 2 103.84 0.005

hydroxyl group is expected to result in little change in the size of a quaternary ammonium ion. Consequently, the alkyl analog of the (EtOH)4N+ ion is the Pr4N+ ion, and the two ions should have thesamehydrodynamic properties if size is the only criterion. However, in aqueous solution a t 25°,2eXO (Pr4N+) = 23.22, some 15% lower than the corresponding value for (EtOH)X+ as reported here. This mobility difference suggests that the enforcement of water structure around the hydrocarbon side chains detectedzbin the Pr4N+ ion is absent in the case of the (EtOH)4N+ ion, owing to the presence of the polar groups in the otherwise inert side chains. This conclusion is substantiated by the effect of temperature on the mobilities involved, as is shown in Figure 2. The Xoqo product for the (EtOH)4N+ ion does not show an increase with temperature as do the larger quaternary ammonium ions around which water structure enforcement is prevalent.2b In contrast to its alkyl analog, the Walden product for (EtOH)4N+ shows a very slight decrease with temperature, a behavior much more characteristic of structure-breaking ions such as Me4N+.Zb Further evidence that (EtOH)4N+does not enforce water structure in its immediate vicinity comes from the fact that its Walden product (0.241) for aqueous solution is almost identical with those for Its alkyl analog in acetonitrilez8 (0.240), nitr~methane'~(0.245), and methanolz& (0.251), in which the hydrodynamic properties of the Pr4N+ ion should be essentially unaffected by solvent interaction. Further evidence that cages of water do not form around the side chains of the (EtOH)4N+ion can be obtained from the viscosity data. Viscosity B coefficients for (EtOH)4NBr are compared with those for the tetraalkylammonium bromides3 at temperatures between 0 and 65" in Figure 3. The "melting" of such cages around the side chains of both the Pr&+ and BudN+ ions is evident3 from the rapid decrease in B for the bromides of these ions with increased temperature. I n contrast, the B coefficients for (Et0H)dNBr are independent of temperature and much lower than those for Pr4NBr. It is possible that, at temperatures high enough for the elimination of all water structural effects, the B coefficients for these two salts would be

1.91 f 0.02 1.97 f 0.03 2.41 f 0.03 1.08 =!z 0.08 1.4d*00.02

0.35

0.30 A0310

UA

J

0,008 0.01 0.015 0.03 0.006

92.3 95.4 181.7 44.4 68.1

t1

1

i (Et OH), N*

25

45 TOC

Figure 2. Temperature dependence of the Walden product for the (EtOH)pN+ ion as compared to the tetraalkylammonium ions in aqueous solution.

the same as theory predicts for ions of the same size. Again, from viscosity measurements it would appear that the transport properties of the (EtOH)4N+ion are affected very little by interaction with water. We conclude from these results that the substitution of an ethanol for a propyl group in a quaternary ammonium ion results in an increase in mobility in aqueous solution due to the disrupting effect of the polar group on the water-structure enforcement normally around a propyl group. This conclusion would appear to be in conflict with the results reported in previous investiga~ , the ~ inclusion of an t,ions of this type of i ~ n since 1223 R. (15) (1963). L. Kay,

S. C. Blum, and H. I. Schiff, J. Phvs. Chem., 67,

vo'olume 70,Number 9

Septmber 1966

2978

D. F. EVANS, G. P. CUNNINGHAM, AND R. L. KAY

I

I

I

' 1

I

25i 20

r

MedNBr 4

L

I 0

10

20

I 30 T*C

I

I

I

40

50

60

1

Figure 3. Temperature dependence of the viscosity B coefficients for (EtOH)aNBras compared to the tetraalkylammonium bromides.

ethanol group in a quaternary ammonium ion was claimed to have little or no effect on its mobility in aqueous solution. We have resolved this conflict by noting that the alkyl analogs used for comparison had the same number of carbon atoms but did not have the correct geometry. This can be seen in Figure 4 where the limiting conductances for a number of quaternary ammonium ions at 25" in aqueous solution are shown as a function of the number of carbon atoms in the side chains. The values used for the limiting conductances, after recalculation on the basis of eq 2, are given in Table I11 and were taken from sources already cited and from data quoted by Kraus and co-workers.16-18 The results of previous investigations6v6of quaternary salts containing ethanol side chains were compared to a smooth curve passing through the points for the symmetrical R4N+ ion. Since the point for Mes(EtOH)N+ was very close to this line and that for Me2(EtOH)2N'+was not, too far removed from it, it was concluded that either the addition of an ethanol group affects the mobility to the same extent as a propyl group5 or, if the ethanol group was interacting with the water molecules in its vicinity, there were compensating factors that masked the effect on the mobility.6 Spivey and Snel16 attributed the low mobility of the hleoPrN+ ion, as compared to that for the symThe Journal of Physieal Chemistry

4

6

8

10

12

14

16

18

20

No. of Carbon Atoms

Figure 4. Limiting conductances for various quaternary ammonium ions in water a t 25" as a function of the total number of carbon atoms or their equivalent in the side chains: A, predicted value for the Me*Pr*N+ion.

metrical quaternary cations, to its asymmetry. However, the Me3(EtOH)N+ ion is equally as asymmetric and has a much higher mobility than the Me3PrN+ion. Furthermore, (EtOH)4N+ and Pr4N+ are both symmetric ions with equivalent numbers of carbon atoms, but the mobility of the latter is 15% lower. It would appear from these observations that symmetry, as well as number of carbon atoms, is not the deciding factor on which to base a comparison of the mobility for these large ions. The mobilities of the ions given in Table I11 fall into three separate families depending on the number of side chains about which there is enforcement of water structure. We have used a different approach in order to compare the mobility of these large ions. The larger symmetrical tetraalkylammonium ions fall into one category, as shown by the line through the points for Pr4X+,Bu4N+,and n-Am4N+in Figure 4. The Et4N+ ion, however, does not fit on the extension of this line to the Me4Y+ion owing to its inability to enforce water structure to the same extent per carbon atom as the ions containing larger side chains. This suggests that there is a critical size of hydrocarbon chain in a quater(16) H. M. Daggett, E. J. Bair, and C. A. Kraus, J . Am. Chem. Soc., 73, 799 (1951). (17) M. J. McDowell and C. A. Kraus, ibid., 73, 2170 (1951). (18) E. J. Bair and C. A. Kraus, ibid., 73, 1129 (1951).

INTERACTIONS OF THE TETRAETHV~NOLAMMONIUM ION WITH WATER

2979

~

~~

Table 111: Limiting Ionic Conductances for Aqueous Solutions a t 25'" Ion

Me4N Et4N Pr4N + BUN + n-AmN

Ref

Ion

Ion

Ref

44.42 32.22 23.22 19.31 17.38

1 1 1 1 16

38.21 33.58 27.0

5, 6 5 This work

MesEtN + MeaPrN+ MerBuN Merhexyl-N + Meroctyl-N + Me-decyl-N + Mebdodecyl-N + Medetradecyl-N +

40.86 36.64 33.54 29.52 26.49 24.30 22.40 20.96

5, 17 6 17 17 17 17 17, 18 17

XO

+

+

+

Mea(EtOH)N+ Me,(EtOH)zN (EtOH)aN+

+

+

+

a The cation limiting conductances are based on the following anion limiting conductances:z h(Cl-) = 76.39, Xo(Br-) = 78.22, &(I-) = 76.98, and Xo(NOs-) = 71.57.

nary ammonium ion below which water structure enforcement drops off rapidly but above which such structure is cooperatively formed with relative ease. Thus, in these ions, the methyl and ethyl groups are poor structure markers, but the propyl group has excellent structure-making properties. Further evidence for this critical size and cooperative aspect can be seen in the ionic conductances for the trimethylalkylammonium ions, as shown in Figure 4. On the basis of number of carbon atoms, they fall on a separate and higher line than do the R4N+ ions; that is, they act as ions with fewer carbon atoms. Thus, the three methyl groups are ineffective as structure markers and may well be reflecting some st ructure-breaking properties. This suggests, therefore, that the logical comparison for the nle3(EtOH)N+ ion is the Me3PrN+ ion and that the correct comparison for the RSe3(EtOH)zN+ion is not the Et4N+ion but rather the Me2Pr2N+ion. A calculation shows that the increase in the limiting conductances for nle3(EtOH)N+ and (EtOH)4N+ over their correct alkyl analogs is 4.0 f 0.2% per ethanol group. On this basis, the mobility to be expected for the n!IezPr2N+ion is 31.0 or 8% lower than that for Mez(EtOH)zN+. This point is designated by a solid triangle in Figure 4 and, as expected for an ion with water structure enforcement about two side chains, it is between the lines for ions about which there is water enforcement about one side chain and about four side chains. The obvious conclusion resulting from these data is that an ethanol group does not enforce water structure when substituted in a quaternary ammonium ion. On the other hand, ethanol itself, when added in small amounts to pure water, is known to enforce water structure,1g*20 presumably by being hydrogen bonded into the water cages.21 However, there is no evidence that the ethanol groups in the (EtOH)4N+ ion are incorporated into the "flickering of structured

water since a substantial decrease in mobility would result. No such decrease in mobility has been detected here on the basis of any realistic comparison. We conclude that the side chains of the (EtOH)4N+ ion are not incorporated into the surrounding water structure. In a previous investigation' of the tetraalkylammonium salts in aqueous solution, it was shown that the decrease of the conductance with concentration became greater the larger the anion and cation, contrary to the predictions of theory.14 The possibility of attributing the effect to ionic association and to changes in solvent structure due to overlap of ionic cospheres has been considered in some detail.z The same effect was observed by Skinner and F U O SinS their ~ ~ measurements on i-AmsBuNBr in aqueous solution, but the effect is particularly evident in the iodides. However, the conductance of (EtOH)4NI is about 0.5% greater at C = 0.007 M than that which would be required for the same concentration dependence as Pr4NI. This is a significant amount considering the agreement in io+ from the bromide and iodide. It indicates that (EtOH)4NI acts more like hfe4NI than like Pr4T\'Iand suggests that the abnormally large decrease in conductance for Pr4NI and Bu4NI is due to water structure considerations. Opposed to this conclusion is the fact that sodium tetraphenylboride in aqueous solution has a normal concentration dependence22 although recent measurementsz3indicate that water structure enforcement around the aryl groups is about the same as that around a butyl group. These results indicate that all (19) F. Franks, Ann. N . Y . Acad. Sci., 125, 277 (1965). (20) R. L. Kay and T. Vituccio, to be published. (21) A. D. Potts and D. W. Davidson, J. Phys. Chem., 69, 996 (1965). (22) J. F. Skinner and R. M. Fuoss, ibid., 68, 1882 (1964). (23) G. P. Cunningham, D. F. Evans, and R. L. Kay, to be published.

Volume 70, Number 9 September 1966

2980

D. T. PETERSON AND J. A. STRAATMANN

the significant factors affecting the concentration dependence of conductance in dilute solutions are not known a t the present time.

Acknowledgment. This work was sponsored by the Office of Saline Water, U. s. Department of the Interior, under Contract No. 14-01-0001-359.

Lanthanum-Lanthanum Hydride Phase System'

by D. T. Peterson and J. A. Straatmann Contribution No. 1874 from the Institute for Atomic Research and Department of Metallurgy, Iowa State University, Ames, Iowa (Received April 6 , 1966)

The phase relationships between lanthanum and lanthanum hydride have been determined. The addition of hydrogen increased the melting point and decreased the fcc-tobcc transformation temperature of lanthanum to a eutectoid a t 773" and 23.5 at. % hydrogen. The solubility of lanthanum hydride in fcc lanthanum ranged from 2.83 at. % at 375" to 21.3 at. % hydrogen a t 773". The enthalpy of solution was calculated to be 5.76 f 0.17 kcal. Above 773" the solubility of lanthanum hydride in bcc lanthanum increases rapidly, and complete solubility occurs at temperatures above 960". Hydrogen depressed the low-temperature transformation in lanthanum by 27 ".

Introduction The lanthanum-hydrogen system has been investigated by Mulford and Holley2 by determining pressurecomposition isotherms in the temperature range between 600 and 800". They found a significant solubility of hydrogen in the metal which increased with increasing temperature while the lower hydrogen concentration limit of the hydride decreases with increasing temperature. Pressure-composition isotherms are limited to a small temperature range because the hydrogen dissociation pressure becomes either too small or too large to measure with reasonable accuracy. General relationships between the properties of metallic hydrides and their structure have been reviewed by lib ow it^.^ The heat of formation of lanthanum hydride is -49.7 kcal/mole of Hz. The face-centeredcubic (fcc) fluorite structure of lanthanum hydride exhibits a composition range from below lanthanum dihydride to nearly lanthanum trihydride. Neutron diffraction studies indicate that the tetrahedral holes are filled by hydrogen first, and further absorption of hydrogen results in the filling of the octahedral holes The Journal of Physical Chemistry

in the structure. The effect of hydrogen on the melt ing point and structural transformations of lanthanum metal has not been investigated. In order to determine these effects and more accurately establish the solubility limits over a wider range of temperatures, this system was investigated by differential thermal analysis, isothermal equilibration, X-ray diffraction, electrical resistivity, and dilatometer techniques. The melting point and the allotropic transformation temperatures of pure lanthanum have been measured by a number of investigators. Love4 reports a lowtemperature transformation from a hexagonal-closepacked (hcp) structure to an fcc structure at 310" and a transformation from the fcc structure to a body-

(1) This work was performed in the Ames Laboratory of the Atomic Energy Commission. (2) R. N. R. Mulford and C . E. Holley, Jr., J. Phys. Chem., 59, 1222 (1955). (3) G. G. Libowitz, "The Solid-state Chemistry of Binary Metal Hydrides," W. A. Benjamin, Inc., New York, N. Y., 1965. (4) B. Love in "Metals Handbook," Vol. 1, American Society for Metals, Novelty, Ohio, 1961, pp 1230, 1231.