NOTES
286
D ~ T FIT A TO AdAdsorbent sorbate Ferric oxide gel Benzene Silica plui: Benzene Silica pluf: Cyclohexane Silica pluj: n-Hexane Silica-alumina Sitrogen gel Catalysis -660 Nitrogen Catalysis -680 Nitrogen Catalysis -970 Xitrogen
TABLE I BDDT EQUATION
THE
T (OIL) Vm C 323 0.081 27.0 299 7 7 . 0 18.35
299 299
53.9 54.6
78 76.07 78 9 8 . 8 78 88.5 78 -13.5
g
5.31 13.8 10.75 1 4 . 0 79.7 65.0 57.0 65.0
n
13.07 6 16.05 23 35 25
Ref. 4
5 5 5
8.58 11.5 6 8.58 7 7 8.58 7 7 8.58 7 7
Vol. 64
Clod-, F-, HF2- and K + to the acid. From their value of the activation energy they concluded the slow step to be one of diffusion of molecular H F through an effective boundary layer. The purpose of this study was (1) to extend the dissolution studies above 0.5 N H P using pressure measurements of evolved hydrogen for rate determinations, (2) to study the effect of temperature on rate, (3) to investigate the effect of noble salt additions to the acid, and (4) to explain the mechanism of dissolution.
Experimental Material and Apparatus.-A low-hafnium content zirconium was used with an average analysis: 0, 0.11%; N, 0.005%; Fe, 0.04%; Hf, 0.01% b.w. After rolling, the zirconium was given a stress relief by an annealing treatr ment in vacuo for 30 minutes a t 700". The apparatus operated on the principle of measuring a differential pressure between a selected standard atmosphere and the resulting pressure of the gaseous product formed in the unplasticized polyvinyl chloride reactor vessel. Use of a differential pressure between the reaction vessel and a ballast vessel of the same volume containing the same liquid reactant eliminated the necessity of corrections for changes in atmospheric pressure and temperature during a run and for vapor pressure effects of the liquid reactant. Furthermore, it permitted the study of reactions at pressures up to the operating pressure limit of the reactor (165 p.s.i.a.) and down to the equilibrium vapor pressure of the liquid reactant. Zirconium samples were mounted in unplasticized polyvinyl chloride in a metallographic mounting press a t 6000 ( 5 ) J. J Van Voorhis, R. G. Craig and F. E. Bartell, TIIISJVURNAL, p.s.i. and 130". The polished sample was attached to the 61, 1513 (1957). stirrer so as to expose approximately 1 cm.2 of surface area. (6) W. D. Harkins and G. Jura, J. Am. Chem. Soc., 66, 1366 (1944). Due to the relatively large volumes of liquid used (500 ml.), changes in temperature of the bulk solution did not exceed (7) €1. IC. Ries, Advances 2 % Catalysts, 4, 87 (1952). more than one degree up to 1.O N and two degrees to 3 A'. The Rates.-Since the velocity of the liquid across the sample surface influences the rate of dissolution, the effect of CHEMICAL KINETICS OF THE stirrer speed was studied on samples in 0.25 N H F at 30". ZIRCONIUM-HYDROFLUORIC ACID The rate of hydrogen evolution was found to be directly proportional to the stirrer speed, r = 0.7s 200 where r s! REACTIOXl in units of mm.3 om.+ m b - 1 and s is the stirrer speed in r.p.m. In this study a stirrer speed of 90 r.p.m. was arbiBY W. J. JAMES, W. G. CVSTEADA N D M. E. STRAUMANIS trarily selected in order to approximate the measured rates Departments of Chemical and ilfetallurgical Engineering, School of observed in a previous study a t lower concentrations .4 Mines and Metalluvgy, University of Missouri, Rolla, Missouri The differential pressures recorded on the chart were conReceived October 84, 1969 verted into mm.3 of Hz evolved. The rates were calculated Many reports have been published on the re- using the expression 1000 A V sistance of zirconium and its alloys to chemical y =A At attack, particularly in more recent years with respect to corrosion in water and steam a t elevated where r = rate in mn1.3 em.+' min.-*, V = cc. of Hz a t STP or temperatures. Published material pertaining to evolved in At minutes, A = area of the samplc in 2.035AV the corrosion of zirconium in hydrofluoric acid has y=---
agreement. In the other cases the agreement was intermediate between the results shown in Fig. 1 and 2. It should also be pointed out that variation of "g" from the calculated value did not improve the fit of the experimental curve. On the whole, it is felt that the theory adequately represents the experimental adsorption isotherm. In order to calculate adsorption isotherms at different temperatures, the BDDT paper assumes Q to be independent of temperature. Examination of equation 7 reveals this to be a necessary comequence of this equation. This interesting fact also means that Q can be evaluated at any temperature, where surface tension and density data are available, axid the results will be applicable at the adsorption temperature.
+
been more limited. However, interest in the aqueous processing of zirconium-uranium reactor fuels has stimulated more research into the zirconium hydrofluoric acid reaction and is reflected in the recent literature.2-6 In part'icular Smith and Hill3 using radioactive Zrg5as a tracer measured the rates of dissolution in HF-HC1 mixtures up to approximately 0.5 .V. They found the reaction t o be first order with respect to un-ionized HF, independent of oxygen concentration in the vapors above the acid and independent of small additions of NOS-, C1-,
(1) This work supported by t h e U. S. Atomic Energy Commission. (2) J. C. Baumrucker, Dissolution of Zirconium in Hydrofluoric Acid, ANL-5020 (March 31, 1950). (3) T. Smith and G. R. Hill, J . Electrochem. Sac., 105, 117 (1958). (4) M . E. Straumanis, W. J. James and A. 9. Neiman, Corrosion, 16, 286t (1959). (5) E. hl. Vander Wall and E. M. Whitner, Ind. Eng. Chem., 61, 51 (1959).
A At
where r is now in units of mg. of Zr cm.-2 n 1 h - I . The reaction rates were observed for an acid concentration range of 0.1 to 3.0 N a t a temperature of 30' in aqueous HF. A t concentrations above 3.0 N the reaction was so rapid that accurate measurements could not be made. The rate equation for the dissolution of zirconium in these concentrations was assumed to be
A plot of log rate versus log acid concentration (stoichiometric) is shown in Fig. 1 . In concentrations up to approximately 0.5 N the reaction obeys a first-order rate law. The deviation from linearity a t higher concentrations would appear to indicate that either the order of the reaction is changing or that the rate is not proportional to H F only ?t these higher concentrations. Rates also were observed in the same concentration range in mixtures ?f H F and 1. N HCl to control the equilibrium concentrations of fluonde complexes. The rate of hydrogen evolution in the HF-HCl mixtures
Feb., 1960 was higher but similar in shape to that of pure HF, Fig. 1. Thus, the effect of added HCl apparently was to shift the equilibrium further in the direction of un-ionized HF thereby increasing the rate of hydrogen evolution. Influence of Tern erature on Rate.-The equipment was similar to that used \y Straumanis, et aL4 The experiments were made a t seven concentrations of acid from 0.01 to 0.25 N . Duplicate runs were made at 30,40,50 and 60' in each of the seven concentrations. The temperature was controlled to within zkO.2'. A stirring speed of 200 r.p.m. was maintained throughout all runs. From plots of initial rate versus H F concentrations a t each temperature, specific reaction rate constants were calculated. From a graph of log IC versus 1 / T an activation energy of 3.8 kcal./mole was obtained. The increase in rate above 0.5N suggested the possibility of higher temperatures a t the liquid-metal interface than in the bulk solution. Accordingly the change in temperature ( A T ' )was investigated by attaching an iron-constantan thermocouple to the back of a mounted zirconium specimen immersed in various concentrations of H F stirred a t 200 r.p.m. (Fig. 2). Effect of Noble Salt Additions-Gold chloride, platinum chloride and silver nitrate additions were made after sufficient time had elapsed to allow the rates to become constant in pure acid. A rate increase was observed in the first ten minutes following addition after which there was a rather rapid and continuous drop for the remainder of the runs. All samples were covered with loose deposits identified by X-ray analysis as the respective reconstituted noble metals. For equivalent amounts of added salts the platinum was most effective in passivating the metal surface (93%) with silver nest (18%) and gold (lG%). A plot of rate versus time is shown for a platinum chloride addition (Fig. 3).
Discussion and Conclusions At 30.0" and a concentration of 0.25 N H F the value of k expressed in g(Zr) min.-l (mole HF)-' liter is 4.3 =t0.4 X On the basis of an activation energy of 3.8 kcal./mole (Smith and Hill, 3.34) (Baumrucker, 4.2)2 (Vander Wall and Whitner, 6.6)6 and the fact that the rate increases a t high H F concentrations (excluding the slow step as one of adsorption) it is likely that the slow step is diffusion of molecular H F to the metal surface. The high value of activation energy obtained by Vander Wall and Whitner was observed in HF-HiY03 solutions. The high concentration of "03. a strongly oxidizing acid, would favor the formation of oxide layers which in turn would further hinder the diffusion of HF molecules resulting in lower rates and a higher activation energy. In our studies we have also observed that the addition of oxidizing agents such as Cr04- and MnO4- reduces the rate of reaction considerahly. Up to approximately 0.5 N H F the reaction is In concenfirst order with respect to (HF).,. trations of H F or HF-HC1 mixtures up to approximately 0.4 S,a black film, ZrHP, is present on the metal surface.4v6 In the vicinity of 0.4 S the hydride is quite soluble in HF and this may in part account for the increased rate since the H F rieed no longer diffuse through the porous hydride. In addition, Fig. 2 suggests that a temperature gradient exists in the boundary layer which could account further for some of the observed increase in the rate. However, the temperature coefficient is far too sinal1 t o account for all of the increase. It is interesting to consider the reactions which might OCC'III' in the sample surface. Even high (ti) W. J . Jaiiies and hl. E. Stiaunianis, J . Eledrochem. Soc., 106, 631 (195'3).
287
NOTES lo; 6
3 2
6 104
3 N
Lj5
26
3
B + c,
2 -
.3
2 103 7 5 3 2 102
0.1
0.2 0.3 0.5 0.7 1.0 2.0 3.0 HF, N . Fig. 1.-Plot of hydrogen evolution rate versus HFoconcentration for zirconium dissolving in H F a t 30 .
15
? 10
5
1 .o
0
2.0 3.0 HF, N . Fig. 2.-Differcnce in temperature between bulk solution and metal-liquid interface as a function of H F Concentration for zirconium dissolving in HF. 1600
- 1400 e 1300 8 1000 ,
.3
N
7
e
800
E 600
-
2 400 i rj.
200 0 0 20 40 60 80 100 120 140 160 Time, min. Fig. 3.-Rate of hydrogen evolution 2;ersus time upon adding PtCl, a t the 70th minute to Zr dissolving in HF.
purity zirconium is covered with a thin tenacious oxide layer. Thus, the reaction probably begins
288
NOTES
by the attack of H F on the oxide layer with suhsequent formation of the hydride layer up to 0.4 N HF. Once the oxide layer is removed or traversed by HF the reaction of zirconium can take place by direct chemical attack of un-ionized HF. The increase in rate of Zr dissolving in HF-€IC1 mixtures can be attributed to a shift iii equilibrium to un-ionized HF, increase in conductivity of the solution and an increase in the solubility of the reaction products. The small initial rate increases observed with noble metal salt additions are difficult to explain. They may be attributed to increased cathodic areas in local cells (as postulated in electrochemical dissolution). The subsequent decrease of reaction rate is then due mainly to passivation of the surface by the plating of noble metal and in part due to the increasing anodic current. T'ander Wall and Whitner6 added silver nitrate and reported no effect. However, it should be noted that the addition was made in the presence of strongly oxidizing nitric acid, thus preventing the reduction of Ag+. It should be emphasized that this study is riot just typical of the particular sample used since almost identical results were obtained with zirconium samples containing 3% hafnium. Studies now in progress on hafnium, titanium, and solid solutions ZrO,, TiO,, are resulting in activation energies of the same magnitude lending further support to the idea that these reactions are not controlled by a chemical activation step.
Vol. 64
A more convenient method is outlined here which makes maximum utilization of available equilibrium data on systems for which (Ho2g8- HOO)and SOZ~~ are not known but high temperature heat contents have been measured and tabulated as (HOT - H0298) and (SOT - S0298).a Consider the free energy function
For a chemical reaction, the change of free energy function is A P T
A.H029*
T
DETERMINATION OF AF0298, AH0298 XSD AS"JS8 FROM EQUILIBRIUM DATA AT VARIOUS TEMPERATURES
and, when 4, is plotted versus 1/T, the curve should and intercept a t 1/T be a straight lineof slope At T = 298°K. - AFo2gX,and = 0 equal to ASozg8. BYJOHN L. MARGRAVE this value can be used to check AHo298and ASOZSS. Department of Chemistry,, University of Wisconsin, Madison, Ti'zsconszn Analytically, one may obtain independent values Received September P1, 1060 of A H 0 2 ~and 8 ASo298 from every pair of equilibrium The usefulness of the free energy function for measurements. This treatment yields a highly reliable AHOm treatment of equilibrium data to yield AH0298 hns been widely demonstrated.' In some cases, one is definitely better than can be obtained by ignoring prevented from developing reliable thermodynamic available heat capacity data above 298'K. or by data a t 298'K. by the lack of low temperature heat simply extrapolating a log KT 21s. 1/T plot back to capacities and adequate entropies. There are, 298OK. The approach should find application in however, often available heat capacity data abovc high temperature calculations, and might also be 298'K. where measurements are somewhat easier. useful for converting solution equilibrium data a t These data are normally utilized for calculation of temperatures other than 298OK. to the standard thermodynamic properties a t a desired tempera- reference temperature. When necessary, values of ture by empirical heat capacity equations and ACp.2 A p should be somewhat easier to estimate than A (1) (a) Various papers in the literature, especially high temperature (POT,; since C l p is a nearly linear function vapor pressure measurements; (b) J. L. Margrave, J Chem. Ed., 38,
Horn)
520 (1955). (2) G. N. Lewis and M. Randall, "Thermodynamics and the Free Energy of Chemical Substances," hIcGraw-Hi11 Book Co., New York, N. Y., 1923, pv. 102-105 and 173-175.
at 298OIi. and abovc for many substances. (3) For example, see K. K. Kelley, U. S. Buieau of Mines, Bulletin 476, 1949.
