May, 1962
PROPERTIES OF ELECTROLYTES IN HYDROGEN PEROXIDE SOLUTIOKS
existence of a maximum. in the curves of Fig. 1with respeot t o water content is also in keeping with similar maxima observed in the curves representing the rate of exchange of hydrogen atoms of isobutane with deuterated cracking catalysts.2 This again is consistent with the hydrogen atom production being associated with a Bronsted acid but unfortunately it does not pro've the existence of such acids on the catalyst surface. Taken as a .\vhole, the present radiation experiments suggest that part of the water content of the (14) 8. G. Hindin, A. G. Oblad, and G. A. Mills, J . A m . Chem. SOC., 77, 535, 538 (1955).
923
silica-alumina catalysts is present in some special form capable (like other inorganic acids) of yielding hydrogen atoms on irradiation with y-rays at - 1 9 . 5 O . Whether these atoms may be said to come from BrGnsted acid sitesol the catalyst or to result from special trappillg effects of Lewis acid sites cannot be stated .rvithcertainty at the present time, This hydrogen atom producing capability then is to be added to a long list of distinctive acid-type Properties Possessed to a markedly greater extent than by by gel or by alumina.
PROPERTIES OF ELECTROLYTES IK HYDROGES PEROXIDE-WATER SOLUTIONS. I. SOLVATION OF ALKALI NITRATES' BY MARTINE. EVER HARD,^ PAUL M. GROSS, J R . ,AND ~ JAMES W. TURNER^ Cobb Chemical Laboratorg of the L'niversity of Virginia, Charlottesville, Va., and the Department of Chemistry of W a k e Forest College, Winston-Salem, it'. C. Received December 18, 1061
The partial prlzssures of hydrogen peroxide and water in the systenls alkali nitrate-HzOz-HzO over the full range .of hydrogen peroxide-water compositions a t 50" have been determined. The results are discussed in terms of the preferential solvation of the smaller ions by water and the larger ions by hydrogen peroxide in agreement with the previously reported solubility of these salts in this solvent.
Introduction Aqueous solutions of hydrogen peroxide exhibit significantly drff erent solvent effects on the solutions of cations of differing size based on a series of solubility measurements of the alkali metal salts over the full range of hydrogen peroxide-water concentrations a t 0, 15, and 25°.4,5 Considering only the data obtained for nitrates, it was found that the solubilities of LiK03 and S a N 0 8 decreased on the addit,ion of hydrogen peroxide to water. On the other hand, the solubilit,iesof KN03, RbNOg, and C2s;lT08 were coixderably greater in hydrogen peroxide-rich solutions than in water. Furthermore, ILiNO3 forms hydrates while KX03 and RbKOs form hydroperoxidates. It was concluded that the smaller Li+ and Na+ ions were more easily solvated by water molecules than by hydrogen peroxide in the mixed solvent and that the larger K+, Rb+, and Cs+ ions were preferentially solvated by the hydrogen peroxide. A more quantitative approach to this preferential solvation might be obtained by determining the partial pressures of H202and H20 at, a fixed mole fraction of salt over the full hydrogen peroxide-water concentration range. To accomplish this it is necessary to haTTe available accurate vapor pressures and vapor compositions of this mixed solvent over the full range of hydrogen peroxidewater concentrations. Accurate vapor pressure
determinations were made earlier by Scatchard and eo-workers6 and reliable vapor compositions and pressures were measured by Floyd' in essentially the same equipment described below. The most convenient temperature a t which to make the vapor pressure determinations proved to be 50" since at lower t,emperatures excessive equilibration time was required and at higher temperatures decomposition of the H202occurred. Experimental
(1) Thls work wag supported in part by the Office of Ordnance, U S Army. (2) Phllip Francis d u Pont Fellows. Based in part on the Ph.D. Theses of M. E. E. and J. W. T. a t the University of Virginia. (3) To whom inquiiies should he addiessed at Wake Forest College. (4) J. D. Floyd and P. M. Gross, Jr , J . Am. Chem. Soc., '17, 1435
(6) G. Scatchard, G. M. Karanaugh, and L. B. Ticknor, J . Am. Chem. SOC.,'14, 3715 (1952). (7) J. D. Floyd, Ph.D. Thesis, Univ. of Virginia, 1955. (8) G. Scatchard, C . L Raymond, and H. €5. Gilmann, .I. *4m.Chem. Soc., 60, 1275 (1938). (9) Donated b y the Becco Division, E". M. C . Corp. (10) P. 11. Gross, Jr., and R. C . Taylor, J . Am. Chem. Soe., 7'2, 2075 (1950). (11) C . E, Huekaba and k.€4. Meyes, zbid.,70, 2578 (1~348)~
(1955). ( 5 ) M. E. Everhard a n d P. M Gross, Jr , J . Phus. Chem., 6 6 , 548 (1962).
The equilibrium still consists of a boiler with a resistance wire heater and a thermocouple well. The hot vapors are heated to avoid condensation and fractionation and then are condensed, dripping into a trap where they can be removed for analysis. An overflow from the trap returns excess condensate to the boiler. The volume of solution contained in the boiler is about 17 ml. while that in the vapor trap is about 1 ml. The apparatus for the measurement and regulation of pressure is similar to that used by Scatchards and values can be determined to better than 1%The temperature was measured by a four junction thermal of copper-constantan to better than 0.05". The temperature of the boiling solution mas regulated by adjustment of the total pressure by a needle valve. A 30-1. ballast vessel was contained in the system to avoid erratic pressure changes. The hydrogen peroxide used was 95% by weightQexcept for the most concentrated solutions, which were prepared by distillation following a procedure described by Gross and Taylor.1o The hydrogen peroxide was analyzed by the method of Huckaba and Keyes.l' The solids were analyzed by dry weight determination. For one of the systems discussed here, the thermody-
N. E. EVERHARD, P. M. G ~ oB, e YR., AND J. W.TURNER
924
namic consistency of the data was considered using the method of Boissonnas,'2 which requires the assumptions that the vapors are ideal and that the volume of the li uid is negligible compared to the volume of the vapor. %so Raoult's law is assumed to hold a t the origin of the partial pressure curves. The low pressures a t which these experiments were conducted makes these assumptions tenable. The average deviation of the partial pressure of water in the RbNO, system calculated in this way from the experimental value was 0.52 mm., indicating good internal consistency of the data. The maximum deviation was 1.54 mm.
Results The experimental data are presented in Tables I through IV. The RbN03 concentration was lower than the other salts due to the excessive decomposition found a t higher concentrations. TABLE I VAPORPRESSURE AND COMPOSITION OF THE LiNOa-HzOz H20 SYSTEM AT 50", MOLEFRACTION LiN03 = 0.0632 f 0.001 Total pressure (mm.)
-Mole In solvent
74.3 67.7 58.6 40.5 32.8 25.6 21.6 16.1 12.4 10.0 9.1
0.0424 .1167 .2124 .3628 .4469 .5376 ,6185 .7101 .a160 .9042 .9452
fraction HzOaIn vapor
0.0004 .0004 Very small .a221 ,0489 ,113 .154 .314 .583 .824 .893
TABLE 11
VAPORPRESSURE AND COMPOSITION OR THE NaNOs-H& H 2 0 SYSTEMAT 50", MOLEFRACTION NaN03 = 0.0647 f 0.001 Total pressure
(mmJ
82.0 81.5 79.8 73.0 61.2 47.8 37.2 35.0 29.1 18.4 13.6 11.0
10.5 9.5 9.0
,----Mole In solvent
0.0127 .0228 ,0451 ,1239 .2391 .3717 ,4734 .5017
,5785 ,7391 .8285 .as54 ,9250 .9638 .9949
fraction HzOzIn vapor
Very small 0.0002 .0002 ,0005 0036 ,0085 .0305 ,0330 .0502 .1884 ,4430 .6957 .7381 .8952 .9702 I
From these data the partial pressures of water and hydrogen peroxide were calculated using the equation P, = (P,)(N,'), where P, is the partial pressure of constituent x, Pt is the total pressure, and N,' is the mole fraction of this constituent in the vapor phase. The deviation of the partial pressure curves from that of ideal vapors is an indication of the solvation occurring in solution. According to Hildebrand,13 (12) C. B. Boissonnas, HeEv. Chim. Acta, 22, 541 (1939).
Vol. 66
TABLE I11
VAPORPRESSURE AND COMPOSITION OB THE KNO~-HZO~HzO SYSTEM AT 50°, MOLEFRACTION "01
=
0.0646 f
0.001 Total pressure (mm.)
84.1 79.1 75.3 58.5 35.3 34.7 34.5 25.7 16.1 15.2 11.8 11.7 10.0 9.5
---Mole In solvent
0.0217 .lo22 .1527 .3313 .5617 .5670 .5712 .6789 .8$43 .8417 .go16 .9246 ,9579 .9669
fraction HtO2-In vapor
Very small Very small Very small 0.0011 .0395 .0365 ,0276 .OS43 .236 .344 .530 .452 .832 .916
TABLE IV VAPORPRESSURE AND COMPOSITION OF THE RbNOa-HzOaHzO SYSTEM AT 50°, MOLEFRACTION RbNOs = 0.0397
*
0.001 Total pressure (mm.1
87.1 76.8 51.6 30.6 24.7 15.6 10.8 10.2 9.84
-Mole In solvent
o.ooO0 ,1293 .3763 .5926 ,6480 .8303 .9179 ,9282 .9421
fraotion HzOeIn vapor
0.0000 .0022 .0108 ,1219 ,1509 .3611 ,7235 .goo0 .a601
negative deviatiQns from Raoult's law and increased solubilities tend to occur when the components of a solution have an attraotion for one another which leads to the formation of solvates. I n order to help separate the solvent-solvent interactions from the solute-solvent interactions, the difference between the total pressure of the salt solutions from that of the pure solvent was calculated. The results are 3 h o m in Table V. Sinoe the vapor pressure of water is about ten times greater than that of hydrogen peroxide a t the temperature studied, the effects noted are largely those exhibited by water. These data will be used in the following discussion. In the case of the LiNOasystem, the partial pressure or the escaping tendency of water is decreased over the full range of concentration while that of the hydrogen peroxide is decreased only up to about 0.9 mole fraction hydrogen peroxide. The negative deviations from ideality are due in part to the interaction of the hydrogen peroxide and the mater. This causes the change in $he total pressure to be always negative, see Table V. A consistent picture of the solution structure would he one in which the solute-solvent solvation is largely with water molecules and only involves solute-hydrogen peroxide bonding to a small extent, possibly in the second solvation layer. The slight increase in the (13) J. H. Hildebrand, "Solubility," Reinhold Publ. C o w . , New York, N. Y.,1936,p. 113.
May, 1962
PROPERTIES OF ELECTROLYTE^ IN HYDROGEN PEROXIDE SOLUTIONS
925
RbN03 system as compared to that in the KNOs TABLE V DIFFERENCE IN THE TOTAL PRESSURE IN THE PURE syntem also is evident from the formation of a hydroperoxidate at lower HtOe concentration in the MIXEDSOLVENT FROM THAT IN THE E L ~ C T R O SOLUTION LYT~ former ~ystem.4~5In the RbNO3 system the A P LiNOa AP NaNOs AP Kxon A P RbNOB Mo!e system, system, system, wstem. fraction solute-water coordination would be expected to be mm. mm. mm. HzOz mm. slight as it is when water is the only solvent. The 5 . 6 13.2 9 . 8 7 . 5 0.0 solvation that does occw would be expected to be -2.7 ,-12.4 -2.8 -6.8 .1 mostly by hydrogen peroxide. ,-12.2 -0.2 -5.0 -0.2 .2 Vapor pressure measurements for the CsN03 sys-4.4 .-11.9 0.2 0.7 .3 tem were attempted but excessive decomposition, I . 3 .-11.0 0 . 4 3 . 7 .4 even at low solute concentrations, prevented the -3.0 1.3 0.6 -9.6 .5 completion of these experiments. From the solu2 . 0 1 . 1 1 7 . 6 .o .6 bility data5 it appears that the CsSOa is solvated -1.4 1.2 0.7 -5.3 .7 preferentially by hydrogen peroxide. This solva0 . 4 1 . 6 0 . 6 3 . 5 .8 tion would not be expected to be as strong as in the -0.9 -0.6 0.3 -1.4 .9 case for the smaller ions since a hydroperoxidate 1 . 0 1 . 4 -1.1 -1.2 1 .o was not found. These results support those found earlier from escaping tendency of the hydrogen peroxide, at solubility consideration^.^^^ I n these previous about 0.9 mole fraction hydrogen peroxide, indicates that very little solute-hydrogen peroxide in- studies, it was reported that the smaller cations were teraction occurs. That is, as the water content be- solvated better by water and formed solid hydrates comes small, very little water is “left-over” from while the larger ions were solvated better by hydrothe solute-water interaction to bond with hydrogen gen peroxide and formed solid hydroperoxidates. peroxide. Consequently, the hydrogen peroxide is Similarly the increase in solubility of the larger freer than if the solute were not present and shows ions as the solvent became richer in hydrogen peroxide was taken as evidence that these ions were an increase in its partial pressure. The partial pressures in the NaN03 system shorn better solvated by hydrogen peroxide than by water. a slightly smaller negative deviation than in the The converse was found to be true for the smaller LiNOS system The negative deviations of the ions, that is the poorer solvation of these ions by total pressure also are somewhat less, indicating less hydrogen peroxide resulted in a decreased solubility extensive solvation than in the LiN03 system. It as the solvent became richer in hydrogen peroxide. is expected that the solvation in the NaKOS system It seems likely that the hydrogen peroxide molecule is quite similar to the LiN03 system but, due to the is capable of replacing two water molecules in the smallcr charge density of the Na+ ion as compared coordination of the cation since hydrogen peroxide has two oxygens compared to one for water. Howto the Lif ion, the solvation is somewhat weaker. ever, since the exact number of molecules in the When the K S O 8 system is considered, some striking changes are evident. The solute has such solvation cluster cannot be determincd by the exa strong affinity for hydrogen peroxide that in periments reported here, it cannot be concluded that solutions containing high concentrations of water each hydrogen peroxide in the cluster replaces two the hydrogen peroxide is bonded to the solute so water molecules. The ability of the hydrogen peroxide molecule to that the expected hydrogen peroxide water clusters are broken down. The result 1s an increase in the solvate large ions can be attributed to the increased escaping tendency of the water molecules, as shown distance separating the non-bonding orbitals on by the positive values of the total pressure change in adjacent oxygens in conjunction with the large Table V. As more hydrogen peroxide is introduced dipole moment of the molecule. The increased disinto the system, the hydrogen peroxide-water bond- tance presumably allows the dipole to orient itself ing increases and results in a decreased escaping into a more favorable energy position than the water tendency for the water. The solvation sheath is molecule. smaller than in the cases discussed above due to I n order to compare the diameters of the cations the decrease in charge density, however, for K+ions with the distance separating the non-bonding oxythe hydrogen peroxide would be expected to be pres- gen orbitals of hydrogen peroxide, an approximate ent in the primary solvation layer. distance of separation was calculated. The angle The RbN03 system shows the same trends as between the non-bonding oxygen orbital and the ~ using the noted for the KN03 system discussed above. The 0-0 axis was calculated by S c h a t ~ . ’ By escaping tendency for water is somewhat higher length of the OH bond minus the radius of the than for the ideal case a t lower concentrations of hydrogen atom, a reasonable, minimum length for hydrogen peroxide. As more hydrogen peroxide is the non-bonding orbital can be found. From this added to the system, the escaping tendencv de- information the length of the skew line between the creases, as in the KKOs system. The change in the non-bonding orbitals on adjacent 0 atoms was caltotal pressure, TabIe V, is even more positive than d a t e d to be 2.1 A. This would be a minimum in the KNOS system. This indicates even stronger distance since the non-bonding orbitals actually solute-hydrogen peroxide bonding in the RbN03 would extend out an infinite distance from the oxysystem even though the solute concentration was gens. This calculated distance is smaller than the lower. This lower concentration was necessary due diameter of the potassium ion, 2.66 A,,and conto the decomposition a t higher concentrations. The (14) P. N. Schatz, personal communication, Univ. of Virginia, stronger solvation by hydrogen peroxide in the Charlottesville, Va. THE
926
TEDB, FLANAGAN AND CHANQ HWANKIM
siderably !mailer than the diameter of the rubidium ion, 2.96 A., but larger than the 1.90 A. of the K a + ion. It thus seem8 that ions smaller than potassium are too small to coordinate effectively with the
Vol. 66
hydrogen peroxide molecule; however, the larger charge densities of the smaller ions lead to very effective coardination with the smaller water molecule.
THE EFFECT OF IRRADIATIOS UPOS THE KINETICS OF AS ENDOTHERMIC SOLID REACTIOS. THE DEHYDRATIOS OF hllAKGASOUS OXALATE DIHYDRATE1 B Y TEDB.
FLhNAGAN’ AND CHAICG H ~ A E N( I M ~
Brookhaven iyational Laboratory, Upton, L. I.,
New Yorlc
Receired December 20, 1961
The kinetics of dehydration of virgin and reactor-irradiated manganous oxalate dihydrate have been examined with the aid of a quartz helix balance. The maximum rate of dehydration was noted to increase by a factor of approximately 3 times (70’) while the activation energy was reduced from 22.3 0.7 to 17.3 =t0.7 kcal./mole after a total neutron dose of 6.05 X 1018 n. cm.?.
Introduction The dehydration of niaiiganous oxalate dihydrate has been studied by Topley and Smith4 and Volmer and Seyde1.j Indeed manganous oxalate dihydrate is of some historic interest because the “Topley-Smith effect,” i.e., an unusual dependence of the rate of dehydration upon the surrounding water vapor pressure, first was observed with this compound (e.g., see ref. 6). Topley and Smith found an activation energy of 24.3 kca1.l mole for the dehydration; their rate constants mere determined somewhat arbitrarily, as the rate of loss of water vapor per decigram of manganous oxalate dihydrate at 20% dehydration. While there have been many recent studies of the effects of irradiation, e . g . , p r a y and neutron, upon exothermic solid decompositions, (see especially ref. 7 to 11), there has been a lack of‘ such studies for endothermic r e a c t i o i i ~ . Manganous ~~~~~ oxalate dihydrate was chosen since the dehydration of virgin material had been investigated previously and because the subsequent thermal decomposition of the anhydride also could be investigated.14 Experimental Materials.-Nanganous oxalate dihydrate 1%as prepared by the addition of potassium permanganate to a stirred oxalic acid solution at 8Oo1j and also by addition of manga(1) Work performed under the auspices of the E. S. Atomic Energy Commission. ( 2 ) T o whom inquiries should be sent regarding this work; present address: Chemistry Department, University of Vermont, Burlington, Vermont. (3) Participant summer student program, Brookhaven National Laboratory. ( 4 ) B. Topley and M.L. Smith, J . Chem. Soc., 321 (1935). ( 5 ) M. Volmer and G. Seydel, 2. physilc. Chem., A179, 153 (1937). ( 6 ) W. E. Garner, in “Chemistry of the Solid State,” ed. Garner, Butterworths, London, 1955, Chap. 8. (7) E. G. Prout, J . Inorg. R. Nuclear Chern., 7 , 368 11958). (8) E. G. Prout a n d M. J. Sole, ibid., 9, 232 (1959). (9) P J. Herley and E. G. Prout, ibid., 16, 16 (1960). (10) E. G. Prout, LVature,183,884 (1959). (11) P. J. Herley and E. G. Prout, J . Chem. SOC.,3300 (1959). (12) T. B. Flanagan, Trans. Faraday Sac., 55, 114 (1959). (13) P. J. Herley and E. G. Prout, J . Am. Chcm. Soc., 82, 1540 (1960). (14) T. B. Flanagan, to be published. (1.5) R. W. Coltman, l a d . Eny. Chsm., 16, 606 (1924).
nous carbonate to a hot, stirred oxalic acid solution. Microscopic examination revealed that the dihydrate resulting from both preparations was predominately in the form of rhombohedral plates; in addition, however, some long, prismatic crystals were noted. Several preparations using the manganous carbonate and oxalic acid procedure failed to eliminate the long prisms completely.15 Finally a preparation with a great majority of crystals in the rhombohedral platelet form was employed for the dehydration studies. It was observed, however, that there was negligible difference between the dehydration-time curves of material from the two preparative procedures. The particle size was less than 5 p . Elemental analysis gave: C = 13.42% and H = 2.21% compared to the theoretical values of C = 13.4170 and H = 2.257. for the dihydrate. Apparatus.-The dehydration was studied with the aid of a quartz helix balance (sensitivity 1 crn./l mg., Microchemical Specialties, Berkeley, California). One to 2mg. samples were employed. In the study of endothermic solid state reactions self-cooling often is present16 and consequently the temperature of the sample may be lower than that measured in the furnace. Self-cooling corrections can be made if the sample’s area is known accurately, i . e . , if large single crystals are employed.16 To eliminate selfcooling when employing fine particles, hydrogen gas may be added to the reaction vessel to aid in heat transfer to the sample during dehydration. Topley and Smith4 used this procedure. Helium gas was employed in the present investigation and it was observed that the rate was sensitive to small variations in the helium pressure. The effect of the helium pressure on the rate of dehydration a t both 85 and 95” was investigated. In both cases a maximum in the rate of dehydration occurred at approximately 0.15 mm. of helium; hence, this pressure was employed for determination of the activation energy. -4 self-cooling curve also was determined for an irradiated sample (6.05 X 1018 n. it was found that self-cooling was also of importance in this case, but the rate was not so sensitive to the helium pressure (70”)as was the rate of virgin samples. d pressure of 0.15 mm. of helium also was chosen for the determination of the activation energy of the irradiated sample. Irradiation Procedure.-Small samples, 10 to 40 mg., were placed in quartz containers ( 5 cc.) and evacuated to approximately 10-8 mm. before sealing off in vacuo. The samples then were irradiated in a water-cooled “hole” of the Brookhaven graphite research reactor. In the absence of any appreciable radiation heating due to the presence of a large sample or sample container, the measured temperature of the “hole” is 40 to 50’. The sample vessels were equipped with a break-off seal and could be opened into the high vacuum system to determine the gas evolved after irradiation. One sample was irradiated in air. This (16) M. L. Smith and B. Topley, Proc. Ray. Soc. (London), 8134, 224 (1931).
