2078
RALPHP. SEWARD AND HARRY W. OTTO
Vol. 65
THE RATE OF DISSOCIATION OF PERCHLORATE 10s I N FUSED SODIUM HYDROXIDE1?' BY RALPHP. SEWARD -4ND HARRY 'Iv. OTTO Dept. of Chemistry, The Pennsylvania State University, Cniz'ersity Park, Penna. Receiced June 19, 1961
Sodium perchlorate in fused sodium hydroxide a t temperatures from 360 to 420" is found to decompose to chloride in two consecutive first-order reactions. Rate constants for conversion of perchlorate to chlorate and from chlorate to chloride have bwn evaluated, the first step being the faster. Respective activation energies of 47.3 and 53.2 kcal. mole-' have been calculated. Decomposition rates are much greater in fused sodium hydroxide than in fused sodium nitrate. It is suggested that the decomposition is promoted by hydroxyl ion through the formation of a peroxide intermediate.
When the present investigation was initiated no study of the rate of a homogeneous liquid phase reaction in a non-reacting fused salt solvent had appeared in the literature. It was thought that the decomposition of an oxygen containing anion in fused sodium hydroxide might be such a reaction and that its investigation would be worthwhile. Sodium hydroxide was chosen as the solvent because of its low melting point and an earlier interest in the nature of this material, although its use precluded the use of glass containers and thus introduced some experimental problems. A preliminary trial showed that perchlorate decomposed at a satisfactory rate a t 350-400O while bromate decomposed too rapidly for measurement and iodate disproportionated rapidly to periodate and iodide without loss of oxygen. Furthermore, the decomposition of the alkali metal perchlorates as pure salts in both solid and liquid state already had been in~estigated.~ The over-all progress of the reaction was followed by the determination of the concentration of chloride ion in samples taken a t suitable intervals. Determination of chloride rather than oxygen was choseii because it is known4 that fused sodium hydroxide reacts with oxygen gas, forming up to 3 wt. yosodium peroxide. After it was apparent that a considerable amount of chlorate accumulated during the reaction, experiments were carried out in which both chlorate and chloride were determined in the samples. As the disappearance of chlorate based on these aiialyses follon-ed a smooth first-order relation, it was concluded that no significant concentrations of chlorite or hypochlorite were present during the reaction. Hypochlorite could riot be detected by qualitative tests. Thus the concentration of undecomposed Perchlorate could lie calculated by subtracting the sum of chlorate and chloride concentrations from the total chloriile concentration. X trace of chlorine \vas found in the gaseous product but this was too small for quantitative determination. It nas found that the rate of decomposition of perchlorate in fused sodium hydroxide was some lo4times the rate Calculated by extrapolation of the rates observed for decomposition in pure liquid (1) This work was supported by the U. S Atomic Energy Cuiiim under Contract 4T(30-1)-1881. (2) From the Ph.D. thesis of Hariy Otto, Th? P e n n q l v a n i a stare Vniversity. J u n e 1061. (3) A. E. Harvey, RI. T Edmison, E I) Jonei, R A Seybeit a n d K. A. Catto, b. Am. Chem. S O C 7, 6 , 3270 (1954). (4) H Lux K. Kuhn and T Niedtrmaier, Z nnoig allgem C h e m , 298, 283 (19-a)
potassium perchlorate a t somewhat higher temperatures.3 The difference in the rates of decomposition of potassium perchlorate and of sodium perchlorate in fused sodium hydroxide being relatively insignificant, it is apparent that the hydroxyl ion is an important factor in determining the rate of decomposition of perchlorate ion. For this reason rates of decomposition of perchlorate 'Viere measured in various sodium nitrate-sodium hydroxide mixtures. Investigation of other factors which it was thought might influence the reaction rate included the addition of iiisoluble aluminum, barium oxide and magnesium oxide, soluble sodium peroxide, barium peroxide and water. The effect of substitution of lithium, and in part potassium, for sodium also was investigated. Allurninurncontainers nere used for the sodium hydroxide fusions in this work. When aluminum is immersed in fused sodium hydroxide, gas evolution from the surface of the metal occurs but this lasts only a few seconds. It is proposed that the protective coating consists of a layer of an insoluble sodium aluminate since if the metal is removed, washed and then returned to the melt, the brief attack occurs again. While nickel is satisfactory in resistance to corrosion by fused sodium hydroxide at moderate temperatures, the unfortunate tendency of the liquid to creep up the walls of the container, to solidify n7hen it reaches :t cooler spot, is much more noticeable in nickel than in alumiiium. Corrosion of the aluminum containers did occur but slowly enough so that they could bc used for many hours with only a few milligrams l o ~ in s weight. Experimental Apparatus.-React ions were carried out in aluminum cups inch wall thickness. Constancy of temperature TVRS achieved by fitting the reaction ciips into holes in an :duminum cylinder, 4 inches in diameter and 7 inches in hr4ght. The cylinder had also B thermocouple w l l and a well for a Fenwall Thermoswitch bimetallic temperature controller. In use, the alumiriurn cylinder fitted into a four-inch iron pipe which was surrounded on the sides and bottom by about 3 inches of insulation. A Translte cover reduced heat loss a t the top. A main heater of electrical miring kept the cylinder nearly up to the debired temperature and a smaller intermittent heater operating through the Thcrmosdtch kept the temperature constant to about 1 1 ' . Durmg runs the temperature control -was improved to zttD.5' hy manual adjnstment of the heating current. Temperatures were measured with it. Chromel-Alumel thermocouple Trhich had been calibrated ifith S.R.S. samples of +.in, lead and zinc. The thermorouplr. e m.f. n a s mcapiired to iI microvolt with a Lceds and Korthrup Type K2 potentiometfr. On immersion of the thermocouple in thc I cwting mixture$, teniprratiireu narc fniintl of ttiv ortler ol I ' 1owi.r th:an that 1 inch in diameter and 2.75 inches t d l with
Nov., 1961
RATEOF DISSOCIATION OF PERCHLORATE IONIN FUSED SODIUM HYDROXIDE 2079
of the aluminum cylinder and the recorded temperatures corrected for this difference. This effect is attributed t o heat loss through the aluminum sampling pipets which protruded through holes in the cover, this heat loss more than compensating for the slightly exothermic nature of the reaction. I n carrl ing out a reaction, the reaction cup was first preheated to 470 =t10" in a separate furnace, about 15 g. of sodium hydroxide pellets added, and the molten hydroxide kept a t this temperature for an hour to remove water. The cup then was transferred to the aluminum cylinder and, after giving it time to come to the cylinder temperature, preheated samples of the perchlorate or chlorate were added, and the mixture briefly stirred manually to ensure homogeneity. During reaction the solution was adequately stirred by the evolution of oxygen gas. Approximately one-g. samples of the melt were removed a t various intervals by means of aluminum pipets, dropped on a cool metal plate, weighed after solidifying and cooling, and dissolved in water. Chloride then was determined gravimetrically as silver chloride. Since the densities of the reacting solutions were not knoivn, all concentrations were calculated as moles per kg. of solution. When chlorate also was to be detrrmined, a separate aliquot of the dissolved sample MYLS acidified with sulfuric acid, sodium bisulfate added to reduce the chloratt., and chlorate and chloride combined determined as silver chloride. In samples containing nitrate, reduction of chlorate in alkaline solution by zinc was found more satisfactory. The chemicals used were Reagent Grade commercial products except lithium perchlorate which was prepared from lithium carbonate and perchloric acid. The perchlorates were recrystallized and dried by heating t o 150" or higher a t less than 5 mm. pressure. No chloride or chlorate could be detected in them. Sodium hydroxide pellets were dried as described above and contained up t o 0.5% sodium carbonatc. Other chemicals were untreated before use except for thorough drying.
Results and Discussion Homogeneity of the Reaction.-Although vi&al observation showed sodium perchlorate to be readily dissolved in fused sodium hydroxide, a portion of the binary phase diagram mas investigated to show what the limits of homogeneity were. By recording cooling curves, it was found that a eutectic exists a t 243" and 15 mole yo sodium perchlorate. At the temperatures a t which the rate experiments were done, the solubility of sodium perchlorate is much greater than 15 mole %. From the freezing points in the dilute perchlorate region, the heat of fusion of sodium hydroxide was calculated as 1530 cal./mole, in good agreement with the calorimetric value of 1520 cal. obtained by Douglas and Dever,j thus indicating that the solutions are essentially ideal. The solubility of sodium perchlorate iii sodium nitrate (1n.p. 309") also was investigated. In this system a eutectic was found at 228" and 38 mole yo sodium perchlorate. At 309" the liquid phase is 50 mole % sodium perchloratr. Retortillo and Moles6 have shown that the sodium hydroxide-sodium nitrate system is completely liquid at all compositiorip when above 320". Several decomposition rate experimcntq were carried out with aluminum turnings added to the reaction cups, thus increasing the area of aluminum surface by as much as a factor of four. KO detectable change in the reaction rate was found, indicating that the reaction occurs in the liquid phase with no significant contribution from the metal surface. ( 5 ) T.€3. Iloiiglas a n d I L. Deier, J . Research Xatl. Bur. S t a n d w d s , 63, 81 ( I V i l , . (6) N. 31 Retortillo and E Moles, Anales S O L . espan fis. quim., 31, 830 (1933).
Decomposition rates in the presence of added magnesium oxide and barium oxide were measured. Both oxides are insoluble in sodium hydroxide although any moisture present would form some soluble barium hydroxide in the case of the barium oxide addition. A small decrease in rate was found with magnesium oxide and a small increase with barium oxide. No reason can be advanced for these effects which, in any case, are too small to have any great significance. The Over-all Decomposition Rate.-The rate of production of chloride ion from sodium perchlorate in fused sodium hydroxide was measured a t various constant temperatures from 338 to 416". In all cases plots of the logarithm of total chlorine coiicentration less the chloride concentrations, as a measure of the concentration of unreacted material, zersus time showed initial curved portions of iiicreasing slope followed by straight lines. The reaction thus appeared, after an initial induction period, to be first order. Specific rates, calculated from the linear portions of the plots, were independent of the initial perchlorate concentration as shown in Table I. X plot of the logarithms of the rate constants iersus reciprocal temperature for twenty individual experiments over a 60" range of temperature was linear within the expected precision. It is estimated that uncertainties in individual reaction rate constants may be as much as +15%, the largest uncertainty being due to the large temperature coefficient. X least squares treatment of the data gave
+ 15.18
log k (min.-I) = -11.44 ( 1 0 3 ) / T
(1)
for relating rate constants to temperature. The Rate of Decomposition of Chlorate.-The concentration w s u s time curves for perchlorat,, deconiposition in sodium hydroxide, as described in the previous section, suggested that the decomposition occurred in tn-o steps, the specific rate of the hecond step being somewhat smaller than that of thc first. As a likely intermediate \vas chlorate, measurements of the decomposition rate starting with sodium chlorate were made. In these runs hoth chlorate and chloride concentrations in the samples were determined. On subtracting the comhincd chlorate and chloride from total chlorine, it was found that at no time was there as much as 0..5% of the total iii the form of perchlorate. Plots of the logarithm of chlorate conceiitration z'ersus time were linear with no suggestion of the induction period observed \Then starting with perchlorate. lJirst-order rate constants Ivere evaluated from such plots at four different temperatures and found to be related to temperature by log k'10;
(min.-I) = -ll.G2(103)/T
+ 15.53
(2)
Rate constants calculated from equation 2 a t 630 to 690°K. are 17 to 23y0 higher than those calculated from equation l. While this discrepancy is slightly higher than the supposed uncertainty of 15%, it was concluded that both equations relate to the same reaction, namely the decomposition of the chlorate ion. The Rate of Decomposition of Perchlorate.Reaction rates again were determined with sodium perchlorate as the starting material, but with analy-
2080
RALPHP. SEWARD AND HARRY W. OTTO
TABLEI RATECONSTANTS FOR SODIUM PERCHLORATE DECOMPOSITION Temp., OC. 387 387 386 386 386 387 0.0142 0,0167 0.0198 0.0209 0.0246 Initial mole ratio NaC104/NaOH 0.0103 lo3k, m h - 1 7.03 6 62 5 84 5.93 6.22 6.41
ses for both chloride and chlorate. Perchlorate concentrations were obtained by subtracting the sum of chlorate and chloride concentrations from total chlorine. Perchlorate concentration was found to decrease rapidly while chlorate concentration rose to a maximum and then decreased. Plots of the logarithm of perchlorate concentration versus time were linear and first-order rate constants were evaluated. Equation 3 was obtained from the constants a t three different temperatures.
Vol. 65
406 0.0150 21.0
406 0.0214 21.8
ortho nitrate ions in the melt. If such ions are formed, the sodium hydroxide mole fraction would not be a true measure of hydroxyl ion concentration in the mixture. The Addition of Water, Chloride and Peroxide.The rate of decomposition of perchlorate in sodium hydroxide was measured with the perchlorate added as hydrated sodium perchlorate and also with sodium hydroxide which had been subjected t o less rigorous drying. As no significant change in rates was observed, it was concluded that log kclo, (min.-I) = -10.35(10-3)/T 14.78 (3) Comparison of equations 3 and 2 shows that, in the rate is not sensitive to small amounts of water. As the rate of decomposition of potassium perthe temperature range of interest, the specific rate for perchlorate decomposition is 12 to 18 times that chlorate in sodium hydroxide had been found to be for chlorate decomposition. Kinetic quantities as unaffected by addition of potassium chloride, no calculatcd from equations 2 and 3 are given in further experiments were done involving chloride addition. Table 11. Rate experiments were done in the presence of TABLEI1 added peroxide since qualitative tests showed that KINETIC QUANTITIES FOR THE DECOMPOSITION OF CHLO- small amounts of peroxide always were formed RATE AND PERCHLORATE IN FUSED SODIUM HYDROXIDE when perchlorate or chlorate decomposed in sodium Activation hydroxide, accumulating to a maximum concentraActivation Arrheniua entropy, tion and then gradually disappearing. The results Decomposing energy, frequency oal. mole-' ion kcal. mole-' factor, sec. -1 deg. -5 of these tests were inconclusive. At 360 and 384" Chlorate 53 2 f 2 . 7 5.58 X loL3 2.73 the rate of decomposition of sodium perchlorate was Perchlorate 47.3 i 2 . 4 5.96 X lo1% -0.70 not altered by the addition of sodium peroxide, Relation of Rate of Dissociation to Hydroxyl Ion which is readily soluble in sodium hydroxide. At Concentration.-As the rate of dissociation of 402O, however, peroxide additions caused a sigperchlorate ion in sodium hydroxide turned out to nificant increase in the rate. Barium peroxide, in be so much larger than the rate predicted from contrast, reduced the decomposition rate. Decomposition rates for potassium perchlorate in measurements on pure liquid perchlorates, it was of interest to investigate the effect of a variation in sodium hydroxide were measured while the effect hydroxyl ion concentration. For this purpose rate of peroxide was being considered, since more perexperiments were run in various mixtures of sodium oxide is formed by reaction of potassium hydroxide nitrate and sodium hydroxide. The results of ex- with oxygen than is formed in the sodium ~ y s t e m . ~ perimenh done a t 400 f 1" are shown in Table 111. Here, too, there was a significant increase a t 402", (30% greater rate than for sodium perchlorate, but TABLEI11 a t 380 and 386" no difference between sodium and potassium perchlorates. THEDECOMPOSITION OF SODIUM PERCHLORATE I N SODIUM NITRATE-HYDROXIDE MIXTURES The Rate of Decomposition of Lithium Perchlorate in Lithium Hydroxide-Lithium Nitrate Mole%NaOH 16 24 48 66 78 100 Solution.-Experiments with a system having only 103k(min.-') 1.94 2.18 2.33 3 05 5.67 15.2 lithium cations were done to avoid the presence of In pure sodium nitrate a t 400" the decomposi- peroxide which is reported not to be formed from tion of sodium perchlorate is too small to measure. oxygen and lithium hydroxide. As the melting As shown by gas evolution, sodium chlorate does point of lithium hydroxide is higher than the temdecompose in pure sodium nitrate. Thus the perature a t which the work on the sodium cation hydroxyl has a large effect on the decomposition of system had been done, its melting point was lowered perchlorate and a significant but lesser effect on the by adding lithium nitrate. These experiments decomposition of chlorate. The constants of Table yielded concentration time curves quite like those I11 were calculated as described in the section on the obtained with the sodium cation system. The deover-all decomposition and should be essentially composition rate constants a t 400" are greater, those for chlorate decomposition. The constants however, in the lithium system than in the sodium of Table I11 do not show any simple dependence on system, 5.3 X lov3 compared with 3.3 X the hydroxyl concentration. Retortillo and Moles6 min.-' for a solvent 38 mole% hydroxide, and 35 X found maxima in the temperature-composition min.-l in a solvent compared with 5.6 x diagram for the sodium hydroxide-sodium nitrate 73 mole hydroxide. These observations suggest system corresponding to the compounds NaOH. that a change in cation may be more important NaN03 and (Na0H)z. NaN03. They assumed than the presence or absence of peroxide. these conipounds to be evidence for the existence of Mechanism of Perchlorate and Chlorate De-
+
THERMODYNAMICS IN ALUMINUM-PRODCCIXG ELECTROLYTES
Nov., 1861
composition.-Harvey, Edmison, Jones, Seybert and Catto3 found the activation energy for the decomposition of pure potassium perchlorate to be 70 kcal., which coincides with the energy needed to break a chlorine-oxygen bond with formation of potassium chlorate and a gaseous oxygen atom. For the sodium salts a t 298"K., the corresponding energies are7 NaC104(s) = NaClOs(s) NaC103(s) = NaCIOz(s)
+ O(g) + O(g)
AH = 65.6 kcal. AH = 72.2 kcal.
The experimental activation energies for perchlorate and chlorate decomposition in sodium hydroxide are 47.3 and 53.2 kcal., less by 18.3and 19.0 kcal, respectively. This suggests that the influence of the fused salt medium is essentially the same for perchlorate (7) W.&I.Latimer, "Oxidation Potentials," Prentice-Hall Book Co.9 New York, N. Y., 1952.
208 1
and for chlorate decomposition. That the hydroxyl ion is the important factor is indicated by the reduction in rate attendant on diluting the hydroxyl with nitrate. A mechanism which would be consistent with the observations is C10,(210s-
+ OH- --+ ClOa- + HOz+ OH- +C102- + HOn-
followed by a rapid conversion of (3102- to Cl-. It not likely that any significant concentration of Ha2 could accumulate as equilibrium should lie toward the right in HOZ-
+ OH-
=
Hi0
+
02-
and the presumably rapid reaction 202'
+ 2H20 = 02 + 40H-
would cause the eventual disappearance of peroxide.
THERMODYNAMIC CONSIDERATIONS IN THE ALUMINUM-PRODUCISG ELECTROLYTE BY W. B. FRANK Aluminum Company of America, Alcoa Research Laboratories, Physical Cheniistry Division, h'ew Kensington, Pennsylvania Received June 19. I081
The thermodynamic values for NaaL41F6,NaF and AlFs appearing in the literature are corrected for an apparent error in AlF3 temperature measurement. The free energy change for the postulated dissociation mechanism NasAIFe e 3NaF is calculated a t 1300'K. Dissociation according to this scheme is absent or very slight. Thermochemical properties are developed for undissociated li uid cryolite and molten sodium tetrafluoroaluminate. The use of these functions substantiates the dissociation meclanism of molten cryolite as Na3AlFB + 2NaF NaAIF,. The possibility of the formation of Eiodium aluminate by dissolution of aluminum oxide in fused cryolite is considered. To a nominal alumina content of about 4 weight % this reaction is thermodynamically possible a t 1300'K.: 2NasA1Fo 2&03 = 3Na;1102 3NaAlF4.
+
+
+
Introduction In the industrial production of aluminum, an electrolyte consisting primarily of aluminum oxide dissolved in molten cryolite is used. Despite the intensive experimental and theoretical interest in this system, there is disagreement in the literature concerning the ionic constitution of the solvent and the interaction of the solvent with the solute. It has been established that cryolite dissociates appreciably at high temperatures. However, the nature and degree of dissociation remain controversial. IVhile most investigators accept chemical interaction between alumina and molten cryolite, in contrast t o physical solution, numerous aluminates and oxyfluoroaluminatcs have been proposed as reaction products. Fairly complete and reliable thermochemical data are now available for the pertinent constituents of the electrolyte to undertake a thermodynamic study of the dissociation mechanism and a trcatmciit of one of the proposed schemes of solution. The high temperature heat content measurements on cryolite by O'Brien and Kelley' do not agree with the measurements of Albrighk2 The discrepancy in the two data sets results in a significant difference in derived thermochemical functions a t high temperatures. For example, the use of the data of O'Brien and Kelley results in a value (1) C. J. O'Brien and
K. I
