The enthalpies of formation of iodine pentafluoride and iodine heptafluoride

J. inorg nucl. Chem., Supplement I976. Pergamon Press. Printed in Great Britain

THE ENTHALPIES OF FORMATION OF IODINE PENTAFLUORIDE AND IODINE HEPTAFLUORIDE*t J. L. SETTLE, J. H. E. JEFFES,§ P. A. G. O'HARE and W. N. HUBBARD Chemical Engineering Division, Argonne National Laboratory, Argonne, IL 60439, U.S.A.

(Received 22 January 1974) Abstract--The energy of combustion of iodine in fluorine was measured in a bomb calorimeter and the standard enthalpy of formation of liquid iodine pentafluoride, AH~(IFs, 1,298.15 K), was calculated to be -(210.8~-+0.39) kcal/mole. The combustion products contained variable quantities of gaseous iodine heptattuoride (IF7) in addition to IFs. A linear least squares analysis of the standard energy of combustion of iodine as a function of the ratio of IF7 to IF~ enabled a value of -(229.80-+0.54) kcal/mole to be deduced for AH}(IF7, g, 298.15 K). two parts. For the part in which the objective was to make IF5 the major product, the combustions were conducted under relatively mild fluorinating conditions, and a precise value for the energy of formation (A U~) of liquid IF~ was obtained. For the part in which the objective was to make IF7 the major product, much more vigorous combustion conditions were employed. A least-squares analysis of all the combustion data, relating the corrected internal energy change per g of iodine to the number of moles of IF7 produced per g of iodine, yielded a value for the standard energy of fluorination of IF~(1) to form IF7(g).

INTRODUCTION

THE thermochemical properties of metallic iodides, for instance fission product iodides and those used in purification and vapor-plating of metals by the van Arkel-de Boer technique, are of substantial interest. Solution calorimetry has been used for thermochemical studies of many such materials, but fluorine bomb calorimetry provides an attractive alternative approach for iodides that are not readily amenable to solution techniques. Limited investigation of the reactions of metallic iodides with fluorine has shown that, even with low-pressure fluorine, significant amounts of iodine heptafluoride (IF7) are produced in addition to the major iodine pentafluoride (IF~) product. Therefore, a necessary prerequisite for fluorine bomb calorimetric studies of iodides is the establishment of reasonably precise values for the enthalpies of formation (AH~) of IF5 and IFT. The only existing literature value for the enthalpy of formation of IFs(I) is based on Woolf's[1] measurements of the enthalpy of hydrolysis. Bernstein and Katz[2] measured the equilibrium constant of IF7(g)~IFs(g) + F2(g)

EXPERIMENTAL Materials Fluorine of 99.99 mole % purity was prepared by distillation of a commercial sample in a low temperature still[4]. Iodine was purchased from Electronic Space Products, Inc., Los Angeles, California. Spark-source mass spectrometric, chemical, and activation analyses of the iodine detected the following impurities (in ppm by mass). AI, 1; B, 1; Cr, 2; Fe, 5; Ti, 10; C1, 6; P, 2; S, 10; C, 26; N, 10; and O, 2. The low oxygen content, obtained by activation analysis, indicated that the iodine had a negligiblewater content. The iodine was handled exclusively in a helium atmosphere glovebox (H:O < 5 ppm; O, N < 10 ppm).

(1)

as a function of temperature and calculated the enthalpy change for the reaction. Combination of the enthalpy of reaction (1) with the enthalpies of formation[l] and vaporization[3] of IF5 (1) yielded the enthalpy of formation of IF7 (g). Because of the potential use of the data for the iodine fluorides in other thermochemical studies, the values of their enthalpies of formation were checked by the independent method of fluorine bomb calorimetry. The calorimetric determination of the energy of combustion (A U°c) of iodine in fluorine was conducted in *Work performed under the auspices of the U.S. Atomic Energy Commission. tPresented in part at the 23rd Annual Calorimetry Con[erence, Midland, Michigan, August 13-15 (1968). §Permanent address: Royal School of Mines, Imperial College of Science and Technology, London, S.W. 7, England.

Auxiliary observations Preliminary experiments demonstrated that the reaction between iodine and fluorine was spontaneous. This necessitated the use of two-compartment reaction vessels so that the iodine could be protected from premature exposure to fluorine. The products of the preliminary experiments were identified by i.r. spectroscopy in the 1-15/zm range to be IF515] and IF716] only. In other experiments, the combustion products were collected by bulb-to-bulb distillation from the reaction vessels to a translucent Kel-F tube. Strong side-lighting of the tube enabled the color of the products to be observed. In experiments which employed an amount of fluorine which was about 25-50 per cent in excess of the stoichiometric requirement for IF5 formation, the products were sometimes pale yellow to light brown in color which indicated that some of the iodine had not been fluorinated. In experiments in which a 2-fold excess of fluorine was employed, the products were always colorless, which indicated that no iodine had survived the combustion and subsequent handling of the products. Therefore, at least a 2-fold excess of fluorine was employed in the calorimetric combustions, and the amount of

135

J . L . SETTLE et al.

136

iodine consumed in these experiments was based on the mass of iodine placed in the bomb. The nickel crucibles that contained the iodine never exhibited any weight change (detection limit 0.02 mg), so they neither reacted with the iodine nor were they significantly attacked by the fluorine. There was no observable sign of attack by iodine on metallic parts of the reaction vessels. With sufficient excess of fluorine to insure that all the iodine reacted, some IF7 was always produced. With fluorine excesses that ranged from 2 to 10-fold, the IFv yields ranged from a few tenths to over 65 per cent of the products. Therefore, provided the mixture of the two fluorides was not changing composition in the final rating period of the calorimetric experiments, and that adequate chemical analyses of the mixtures could be made, the enthalpies of formation of both IF5 and IF7 could be determined from a series of measurements in which the fluorine pressure was intentionally varied to form varying amounts of the two products. In our investigation, it was assumed that the composition of the products did not change significantly during the 2-3 hr period that elapsed between the end of a combustion and the separation of the excess fluorine from the products. This assumption is based on the work of Fischer and Steunenberg[7] who studied the kinetics of the reaction IFs(g) + F2(g) ~ IFT(g)

(2)

in the 323.8--368.2 K temperature range. Extrapolation of the data to 298.15 K led to a value of 2.3 × 10-4 mm -1 hr -~ for k, the second order rate constant. From this it was calculated that at room temperature and under our most severe fluorinating conditions, the maximum rate of conversion of IF5 to IF7 was 9.1 x 10-9 mole hr-', a negligible amount in the context of the present work. Chemical analysis of the aqueous hydrolysates of IF~ and IF7 was judged to be the most suitable method of determining the composition of the reaction products. It has long been known [8] that IF5 hydrolyzes quantitatively to give fluoride (F-) and iodate (IO3-) ions. However, there were conflicting reports as to the products of IF7 hydrolysis. The products of the reaction between IF7 and water were stated by Ruff and Keim [9] to be periodate ion (IO4-) and F-, and by Schleiffer and Adloff[10] to be IO3-, IO4and F-. Thus, in order to use the established procedure [ l l ] for the analysis of aqueous mixtures of IO3- and IO,-, it was necessary to demonstrate that hydrolysis of IF5 and IF7 yielded, respectively, only IO3- and only IO4-. A pure sample of IF~ was prepared by combustion of iodine in fluorine and removal of the IF7 formed by fractional distillation. It was shown that a weighed amount of this sample could be hydrolyzed and quantitatively titrated as IOn-. The best sample of IF7 available to us was shown by infrared analysis to contain a minor amount of IFs. Measurements of the vapor density of this sample and titration of the IO3- and IO, produced by hydrolysis could only be reconciled by assuming that the IF7 hydrolyzed exclusively to IO4-. Although our work was limited in scope and was performed on an impure sample, the results support the observation of Ruff and Keim[9] that IO4- is the only iodine-containing product of the reaction between If7 and water. The analytical method[11] depends upon the following facts: (1) in neutral or slightly alkaline solution 103 does not react with iodide ion (I-), whereas IO, undergoes the reaction IO4- + 21- + H20 ~ 2OH- + IO3- + I2;

(3)

and (2) in acid solution iodide ion reacts with both IO3 and 104

*atm= 760Torr= 101 325Nm -2.

according to: IO3- + 51- + 6H ÷~ 3H20 + 312

(4)

IO4- + 7I- + 8H + -->4H20 + 412.

(5)

The amount of IO,- in one aliquot of the solution was determined by titrating the I2 liberated according to reaction (3) with arsenious acid solution. Another aliquot was made distinctly acidic and the total I2 liberated according to reactions (4) and (5) was titrated with sodium thiosulfate solution; subtraction of four times the amount of I2 released by the IO4 according to reaction (3) gave the amount of I2 released according to reaction (4) and, consequently, the amount of IO3- present. Analyses of known mixtures of potassium iodate, periodic acid and hydrofluoric acid established that IO,- and the total of IO3 and IO4- could be determined with an uncertainty of one per cent of the amount in solution. In the calorimetric experiments, the mass of recovered iodine, which was calculated from the analyses of the solution containing the hydrolyzed combustion products, was always less than the mass of iodine placed in the reaction vessel. It was assumed that the material lost was only IF, for the following reasons. Iodine heptafluoride is a gas at room temperature and it is easily transferred by bulb-to-bulb distillation. Iodine pentafluoride is a liquid with a vapor pressure[3, 12] of only 26 Tort* at room temperature. It is very difficult to obtain quantitative transfer of IF5 through metal equipment unless the containing vessel and transfer lines are strongly heated. In our experiments, the heating was limited and the surface temperature was only about 60°C. This encouraged distillation of the products but minimized the possibility of reaction between the iodine fluorides and the materials of construction. Therefore, it was probable that all the IF7 produced in the combustions was recovered and that all the missing product was IF5 that had not been transferred to the hydrolysis tube. Consequently, the mass of iodine placed in the reaction vessel was taken as the measure of iodine that reacted, the chemical analysis of the IO,- content of the aqueous hyrolysate was taken as the measure of IF7 product, and the amount of IF5 produced was calculated by difference.

Apparatus and procedure Two reaction vessels were employed during this investigation. The experiments intended to yield mostly IF5 were performed in a two-compartment reaction vessel (hereafter called I) which consisted of a thin-walled fluorine storage tank that fitted concentrically around a nickel combustion bomb[13]. The experiments intended to yield mostly IF7 were performed in a high-pressure two-compartment reaction vessel (hereafter called II)[14]. When the latter vessel was used, it was necessary to place an inverted nickel cup over the sample to prevent flames from impinging upon and igniting the Teflon sealing button of the valve that connected the two compartments. Both reaction vessels were used in conjunction with a rotating bomb calorimeter (laboratory designation ANL-R2). The rotating mechanism of the calorimeter was used to open the valve that connected the two compartments of the reaction vessels. Each calorimetric system was calibrated by combustions of benzoic acid (Standard Sample 39i, National Bureau of Standards) in oxygen. The temperature of the calorimeter was measured with a Hewlett-Packard Model 2801A quartzcrystal thermometer as previously described[15]. Calorimetric measurements followed standard procedures[16]. The values obtained for the energy equivalents, E(Calor), were (3369.80_+0.61) cal/deg (mean and standard deviation of the mean) and (3255.84-+ 0.29) cal/deg for the systems which included reaction vessels I and II, respectively.

Enthalpies of formation of IF~ and 1F7 The unit of energy is the defined thermochemical calorie equal to exactly 4.184 absolute J. Prior to the calorimetric experiments with iodine and fluorine, the reaction vessels were preconditioned by several noncalorimetric combustions. During pretreatment and thereafter, the vessels were disassembled in the glovebox for insertion of the iodine into the combustion compartment. The iodine samples, contained in nickel dishes, were weighed on a Sartorius semi-micro balance in the glovebox. Because of the volatility of iodine, the combustion compartments could not be evacuated after insertion of the samples, so an atmosphere of helium was present during all the combustion experiments. For experiments using reaction vessel I, the storage compartment was charged with fluorine to a pressure of 7.1 atm.; after expansion, there was about 2.6atm. fluorine pressure in the combustion compartment. When reaction vessel II was used, the storage compartment was charged with fluorine to a pressure of 20arm.; after expansion, there was about 15.7atm. fluorine pressure in the combustion compartment. After each combustion experiment, the sealed reaction vessel was attached to a manifold consisting of stainless steel tubing, Teflon-packed stainless steel valves, three nickel U-traps in series, and a 0-2000 Torr Heise bourdon tube gauge which measured pressures to ±l Torr. In order to minimize premature hydrolysis of the combustion products, the manifold was evacuated, treated with a mixture of IF~, IF7 and F2, heated, and again evacuated. The contents of the reaction vessel were then admitted to the manifold and collected in the U-traps immersed in liquid nitrogen. Helium and excess fluorine were removed by pumping. The mixture of iodine fluorides was distilled from the U-traps to a liquid nitrogen-cooled KeI-F trap containing 15 cm3 of air-free, deionized water. Hydrolysis of the iodine fluorides took place as the contents of the traps were allowed to warm to room temperature. After neutralization with sodium hydroxide, the amounts of IF7 and IF5 were determined. After recovery of the combustion products, the sealed reaction vessel was transferred to the glovebox and inspected for attack on the interior surface of the combustion compartment. RESULTS Combustion

experiments

The results for the combustions of iodine in fluorine are given in Table 1. Symbols in the table are taken principally from Ref.[17], which also gives detailed procedures for the calculations of the standard state corrections. The calorimetric samples of iodine were found to evaporate into the glovebox atmosphere at the rate of 0.15 mg min -~. The elapsed time for transfer of a weighed sample to the reaction vessels was (2.0+0.5) rain, so a correction of - ( 0 . 3 + 0.1) mg was applied to the mass of sample weighed for each experiment. The quantity m' in Table 1 includes this correction. Auxiliary data used in the calculation of item 4 were taken from the following sources: Cp (Ni) from Ref, [18]; C,(I2) and Cv(He, F2, IF5 gas, and IF7 gas) from Ref. [19]; the vapor pressure, density, and Cp of IFs(1) from Ref.[3]; and densities of Ni and I2 from Ref. [20]. For the calculation of item 5, values of/~(atm -1) in the equation P V = n R T ( 1 - t~P) and of ( O U / O P ) r (cal/atm. mole) were derived from the force constants of F2121], He[22] and IF5 (calculated from the data in Ref. [3]), and estimated force constant of IF7. These quantities, as functions of composition at 298.15 K, are given by Eqns

137

(6) and (7), in which x~, xz, x3, and x4 are the mole fractions of F2, He, IF5 and IF7, respectively, in the gas mixtures: /~ = 0.0008 xl 2- 0.0005 x22 + 0.1640 xa z + 0.0101 x~2 - 0.0013 x,x,, +0.0316 x~x3 + 0.0128 x~x~ - 0.0003 x2x3 - 0.0013 x2x4 +0.1331 x3x~ -(OU/OPb-

(6)

= 1.780 xL2- 0.007 x f + 204.684 x~"

+ 14.409 x42+0.414 xlx2 + 32.839 x~x~ + 16.129 x~x4 + 4.306 x2x3+2.286 x2x4 + 138.950 x3x4. (7) Previous experiments, in which gaseous products were formed by the reaction between fluorine and solid combustibles, have shown that the gaseous products were entirely confined to the combustion compartment of the two-compartment reaction vessels[23,24]. A similar situation was assumed to prevail for the present experiments. Because a small part of the IF~ product was in the vapor phase, a correction (item 6) calculated from the enthalpy of vaporization AHv(IFs) = 9.96 kcal/mole[3], was applied for the condensation of the vapor to the liquid state. Observations of the (P, V, T) properties of the combustion products in the collection manifold indicated that the solubility of IF7 in liquid IF5 was negligible unless the quantity of IF7 was large enough to saturate the available vapor space. Based upon the vapor pressure equation[25], it was calculated that the amount of IF7 produced did not exceed 55 per cent of the quantity needed to saturate the combustion compartment volume. It was therefore assumed that the IF7 did not dissolve in the IF5 to any significant extent. The use of two-compartment reaction vessels necessitated the application of a correction, A U~,.k (item 7), which has been described in detail elsewhere[26]. This correction is the mean of several determinations made with each reaction vessel. The correction, h Uimpurities, was based on the assumption that the impurities were uncombined. The products of combustion of the impurities were taken as their most stable fluorides, except for nitrogen and oxygen which were assumed to form gaseous N2 and 02. The net correction for impurities was estimated to be (0.70 - 0.23) cal/g of sample. The quantity A U~F7 is the correction applied for the conversion of IFT(g) to IFs(I) according to the reaction IF7(g) ---)IFs(I) + F2(g ).

(8)

This quantity was determined by least-squaring the combustion data and fitting to it the linear equation y = a + bx

(9)

where y is the corrected internal energy change per gram of iodine (the sum of items 3 through 8 divided by item 1); x is the number of moles of IF7(g) produced per gram of

[AU~]M(Iodine)](cal/g)

98.86 0.071 1.9905~ 0.97329 -3279'79 -1"53t -0.30§ -4.03 1.44 1.39 1.35 -1648.52

95.38 0.647 2.04319 1.00092 -3372.90 -l'60t -0.29§ -4.03 1.44 1.43 12.29 -1646.28

92.05 97.84 96.32 0.056 0.932 0.226 2.28367 1.08268 2.0052, 1.11956 0.53371 0.98258 -3772"69 -1798'50 -3311-10 -1-77t -0"78t -1'53t -0.34§ -0.15§ -0-19§ -4.03 -4-03 -4.03 1.44 1.44 1.44 1.60 0.76 1.40 1.06 17.70 4.29 -1652.92 -1647.36 -1650.5, Mean A U~/M(Iodine) = -1649.47 cal/g Standard deviation of the mean = _ + 0 , 8 4 cal/g

Table 1. Results of iodine combustions in fluorine

*cal -=4.184 J. tContents of the bomb included 10.13 g nickel, 0.0125 mole helium and, before combustion, 0.065 mole fluorine. ~Contents of the bomb included 60.8 g nickel, 0.0054 mole helium and, before combustion, 0.4345 mole fluorine. § V'(gas) = 0-5407 din3; W(gas) = 0.5402 dm3. IIW(gas) = W(gas) = 0.6444 dm 3.

I 2 3 4 5 6 7 8 9 I0

Iodine recovered (mole %) IF7 found (mmole) m' (g) Ate (deg) ,(Calor)(-htc) (cal*) AU ........ (cal) AU ~ (cal) -A Uvap(IFs)(cal) AUbl~, (cal) AUimp~n,,~(cal) A UI~7(cal) 98.97 0.116 2.26468 1.10886 -3736.64 -1"79t -0.33§ -4.03 1.44 1"59 2.20 -1650.37

- 1.74

3.00 0-71 106-25 -1646.39

- 1.74

3.00 0.71 97.27 -1649.9~

-0.12"

91.50 5.595 1'00815 0"54155 -1763.20 -4.83~ -0.0ql

83"15 5.122 1'01167 0.54164 - 1763.49 -4-83~

-0"10" -1.74 3.00 0.7O 95.35 -1652.9o

-4.81~

89'72 5.021 1.00413 0.53815 -1752.13

t-~ t~ ....

.t"

Enthalpies of formation of IG and IG iodine reacted; a is the zero mole IF7 intercept and, thus, the standard energy of combustion to form IFdl); b is the slope of the line and, thus, the standard internal energy change for Eqn (8). The value of a was found to be -(1649.47 +-0.8,) cal/g of iodine. The value of b was found to be (18.99 +-0.27)kcal/mole of IF7. This latter result was used to calculate item 9. Item 10, the standard energy of combustion of iodine in fluorine, A U~/M (Iodine), is the sum of items 3-9 divided by item 1, and refers to the idealized reaction 1/212(c) + 5/2Fffg)---) IFd 1)

(lO)

with reactants and product in their respective standard states at 298.15 K. The values of AU°dM (Iodine) exhibit no significant trend with respect to sample mass, amount of IF7 recovered, amount of product recovered for analysis, or reaction vessel used. Derived data

Derived standard data for the formation of IFdl), IFdg), and IF7(g) are given in Table 2. The standard entropies of I2(c), F2(g), IFdl), IFdg), and IF7(g) were

139

taken as 27.757119], 48.44119], 53.74[3], 79.96[27], and 83.08 [27] cal/deg mole, respectively. The atomic weight of iodine was taken to be 126.9045 [28]. The uncertainties given in Table 2 are uncertainty intervals[29] equal to twice the combined standard deviations arising from all known sources. DISCUSSION Our result for AH~(IFs, I), -(210.8, -+0.39) kcal/mole, is significantly more negative than that reported by Woolf, -(204.7---1.0) kcal/mole[1]. However, Woolrs result is based on auxiliary thermochemical data that have since been largely superceded. We have recalculated his value from the thermochemical cycle and enthalpies of reaction and dilution given in Table 3. The value for AH~(IG, 1) found from solution calorimetry, -(210.77_0.41) kcal/mol, is thus seen to be in excellent agreement with the present determination, and this is due, almost entirely, to the recently redetermined AH~(F-, aq)[30]. Our result for the enthalpy difference between gaseous IF5 and IFT, (28.95 +-0.67) kcal/mole, agrees very well with the value of (28.5+-2.0) kcal/mole obtained by Bernstein and Katz[2].

Table 2. Derived data at 298.15 K

au~

aS~

AC~

(cal/deg mole)

(kcal/mole)

-(209.33---0.39) -(210.8,_+0.39) -(81.26±0.1,) -(199.96± 0.4o) -(200.8s- 0.40) -(55.02± 0.1~) -(228.32-+0.5,) -(229.8o± 0.54) -(100.3, ± 0.19)

-(186.59± 0.3~) -(184.42± 0.40) -(200.79± 0.5,)

(kcal/mole) IFdl) IFdg) IFT(g)

an~ (kcal/mole)

Table 3. Thermochemical cycle and enthalpies of reaction at 298.15 K for recalculationof AH~(IG, 1)* Reaction 1 2 3 4 5 6 7

IFdl) + 1570H20(1) = 5(HF.261.17H20) + HIO3.261.17H20 (H+ +IOf).~ H20 = 1/2I~(c)+3/202(~,+ 1/2 Hz(g) + ~H20 HIO~.261.17H:O + (~ - 261.17)H20(1)= (H÷ + IO3-).ooH20 5(H+ + F-)'~H20 = 5/2 Hdg) + 5/2 F2(g)+ ~H20(1) 5(HF.261.17H20)+(~-1305.85) H20(1) = 5(H+ + F ).ooH20 3Hdg)+ 3/20dg) = 3H20(1) IFdl) = I/2Idc)+5[2Fdg)

AH° (kcal) -(21.71 ±0.06)+ (52.51 ±ODDS -(0.45 ±0.05)§ (401.10 ±0.35)II -(15.74±0.15)" -(204.94±0.03)¶ (210,77- 0.41

*All species are in the aqueous phase except otherwise designated. tWoolf's result at 292 K converted to 298.1 K by means of heat capacity data of Osborne et a/.[3] for IF~, Roth et a/.[31] and Kolesov et al.[32] for HF.nHzO, and Randall and Taylor[33] for HIO3.nH:O. :~Enthalpy of formation of IO3-(aq) from Johnson et al. [34]. §Enthalpy of dilution from Stern and Passchier[351. I~Enthalpyof formation of F-(aq), and enthalpy of dilution of HF(aq) from Johnson et al.[30]. ¶Enthalpy of formation of H20(/) from Wagman et al.[19].

J . L . SETTLE et al.

140

Acknowledgements--The advice and assistance of Herbert H. Hyman during the formative phase of fluorine bomb calorimetry at Argonne and over the course of the intervening years are gratefully acknowledged. We wish to thank G. K. Johnson for providing us with distilled fluorine; J. A. Carter, A. M. Essling, G. G. Mapolo and Z. Tomczuk for performing special analyses; H. M. Feder, D. W. Osborne and L. Stein for helpful discussions; and J. Royal,/~I. Ader and T. Cramer for editorial assistance. REFERENCES

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