July 1950
INDUSTRIAL A N D E N G I N E E R I N G CHEMISTRY
density were constant. The percentage gain in weight was found to be almost independent of air flow rates when samples were heated below ignition temperature. SUMMARY
The results of this investigation indicate t h a t the hydrides of tantalum, zirconium, columbium, and titanium will ignite in air and lose their hydrogen when heated sufficiently. This activity is not explosive when small quantities are tested. Each hydride has a characteristic ignition temperature which varies slightly as heating conditions vary. Tantalum hydride has the lowest and titanium hydride the highest ignition temperature. This ignition temperature is affected by the physical form of the hydride and by whether it is heated in static or dynamic air. These hydrides gain weight during heating essentially because of oxidation although it is possible some gain is due t o nitride formation. The rate of gain in weight is dependent on the rate of diffusion of oxygen into the pellet, a t least up to the ignition point. When other factors are constant, t,he percentage gain in weight
1383
decreases with increase in density and increases with increase in temperature. The reactivity of titanium hydride with dry dynamic air is reduced as the density increases, when the factors of temperature and air flow rate are constant. Temperature increase, when density and air flow rate are constant, results in increased reactivity. When the temperature and density are constant, reactivity will increase when the air flow rate is accelerated. LITERATURE CITED
(1) Gibb, T. R. P., Jr., Electrocham. Soc., 93,199 (1948). (2) Gibb, T. R. P., Jr., "Preparation and Properties of Hydrides,"
Metal Hydrides Incorporated, Beverly, Mass.
(3) Huttig, G. F., Z.angew Chem., 39,67 (1926). (4) Kirschfeld, L., and Sieverts, A , , Ber., 59, 2891 (1926). RECPIVEDOctober 3, 1949. This document is based on work performed under oontract for t h e United States Air Force by the NEPA Division, Fairahild Engine and Airplane Corporation at Oak Ridge, Tenn. Presented before the Division of Industrial and Engineering Chemistry at the 116th CHEMICAL SOCIETY,Atlantic City, N. J. Meeting of the AMERICAN
Iodine Heptafluoride Preparation and Some Properties WALTER C. SCHUMB AND MAURICE A. LYNCH, JR.' Massachusetts Institute of Technology, Cambridge, Mass. Iodine pentafluoride, as one of the interhalogen compounds, has chemical interest both of a theoretical and practical nature. It is unique in being the only simple binary compound in which sevenfold co-ordination is exhibited, and in common with other halogen fluorides it is a possible substitute for fluorine in certain fluorination processes. A detailed account of the preparation and purification of iodine heptafluoride is presented. No stable addition compounds of the type M(IF8) were obtained from the interaction of iodine heptafluoride and several alkali halides. The reactions between iodine heptafluoride and several chlorofluorocarbons are presented along with a brief synopsis of previous work.
M
OISSAN ( 4 ) observed that when iodine pentafluoride, prepared by the combustion of iodine in fluorine, was heated to 500" C . iodine w m liberated, and he concluded that another higher fluoride of iodine must simultaneously be formed. This observation led Ruff and Keim (6) to the discovery in 1930 of iodine heptafluoride. At ordinary temperatures, fluorine and iodine pentafluoride fail to react, but a t elevated temperatures the pentafluoride vapors react in the presence of excess fluorine to produce the heptafluoride, a white solid, which sublimes a t 4 " C. forming a colorless gas with a distinctive odor resembling neither fluorine nor iodine. Under pressures slightly above atmospheric the solid melts a t 5 ' to 6 O C. to a colorless, mobile liquid ( 2 ) . The compound has particular interest not only because it represents a unique example of a neutral binary compound with 7-co-ordination of the central atom, the molecular structure of which has recently been shown (3)to be pentagonal bipyramidal, likewise unique, but also because it may conceivably serve as a substitute for elementary fluorine in fluorination processes. Certain of the other halogen fluorides, such as chlorine trifluoride and bromine pentafluoride, have become commercially available for this purpose. 1
Present address, Linde Air Products Company, Tonawanda, li.Y.
PREPARATION
The iodine heptafluoride needed both for the determination of the infrared and Raman spectra of this compound and for the study of its fluorinating action was prepared in quantity and in a condition of high purity by the following modification of the method employed by Ruff and Keim (6): The reactor em lo ed, shown in Figure 1, was constructed of copper, nickel, and'donel, as indicated. Quarter-pound samples of reagent grade iodine were introduced by removing the soft copper gasket, 4, a t the top of the column, or else samples of iodine pentafluoride, previously prepared, were introduced through the fluorine inlet valve, 3. (The preparation of the pentafluoride was carried out in a similar apparatus, except that the condenser assembly was omitted and a nickel stirrer, motor driven, passed centrally down through the column, I , to the bottom of the Monel sphere, 2.) In the former case, but not the latter, the Monel sphere, 2, was chilled in an ice bath. Fluorine was admitted at a rate of approximately 0.4to 0.5 mo!e per hour, a t which rate absorption was essentially complete. The fluorine was supplied from a commercial 70-ampere 100" C. cylindricaltype generator manufactured and donated for the authors' use by the Hanhaw Chemical Company of Cleveland, Ohio. The electrolyte was of composition approximating KF-1.9 HF, containing a small proportion of lithium fluoride which reduces the incidence of anodic polarization (7). Hydrogen fluoride was removed from the gas by passage through a copper trap cooled to -78' C. and through a copper tube containing sodium fluoride a t 100" C. Waste fluorine was absorbed by passage through a 6-foot length of 2-inch iron pipe packed with roll sulfur t o convert the gas t o a mixture of sulfur fluorides. After passage of the effluent gas through a 1-inch Monel tube heated t o about 400" C. in order t o yrolyze any disulfur decafluoride, the effluent gas was allowe~ftoescape into the air through a copper tube leading directly t o the roof. When the absorption of fluorine had appreciably decreased, the temperature of the column was raised t o 280' t o 290 O C. and the Monel flask was brought t o 70" t o 80" C. by means of a water bath, thereby increasing the concentration of iodine pentafluoride vapor passing up the column. The rate of flow of fluorine was reduced at the same time to 0.2 to 0.3 mole per hour. The iodine heptafluoride leaving the column was condensed out of the stream of excess fluorine in Monel traps cooled to -78' C. by a mixture of solid carbon dioxide and trichloroethylene. Based on the uantities of iodine used, estimates of the amount of heptafluori3e obtained indicated nearly quantitative yields of the product.
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INDUSTRIAL AND ENGINEERING CHEMISTRY
The purification of iodine heptafluoride was carried out by fractional sublimation. The saom le contained in a Monel trap (9, Figure 2) was cooled to -78 and evacuated t o remove any more volatile impurities and then sublimed at 0' C. from one trap to the other, which was cooled t o -78" C. Again, by changing the second trap it was possible t o obtain several fractions. The essential purity of the heptafluoride was indicated by the fact that no appreciable difference was noted in the infrared spectra among the various fractions taken from the purified and the unpurified samples.
Vol. 42, No. 7
fluoromethane was bubbled through iodine pentafluoride and the mixed vapors heated to 225' C. only a few per cent of dichlorodifluoromethane were formed. On the other hand, bromine trifluoride was found by Banks et al. (1) to be capable of converting 96% of trichlorofluoromethane t o dichlorodifluoromethane, and under pressure chlorotrifluoromethane in good yields could be obtained from the same reactants. I n the present work the reactions of iodine heptafluoride with Freon-12 and Freon-13 were studied. Analysis of the products formed was carried out by infrared spectroscopy. Samples of REACTIONS OF IODINE HEPTAFLUORIDE pure dichlorodifluoromethane and chlorotrifluoromethane were ABSENCE OF ADDITION COMPOUNDS WITH ALKALI FLUORIDES. used to obtain the percentage absorption of several of the bands for both compounds a t various values of their vapor pressures. Sharpe and Emeleus ( 8 ) have reported the preparation of addiThe complete spectra of both compounds checked those pretion compounds of some alkali fluorides with bromine trifluorideviously recorded (11) to within four, and for many bands within e.g., K(BrF4) and Ba(BrF&-and mentioned some preliminary one, wave number. The Pyrex glass No. 774 infrared cell emevidence for the existence of a similar compound with iodine ployed (8, in Figure 3) was 6-cm. long and had potassium bromide pentafluoride, possibly K(IF.5). The authors have carried out windows sealed to either end with gIypta1. By measuring the semiquantitative experiments with iodine heptafluoride in the total pressure of the gas sample within the cell and comparing its presence of anhydrous sodium fluoride, potassium fluoride, and spectrum with the standard curves, it was possible to estimate rubidium chloride, respectively. Weighed samples (0.1 to 0 3 within a few per cent the composition of the sample. This gram) of the well dried, finely divided alkali halides were placed method of analysis is more accurate and simpler than low in a weighed Pyrex glass No. 774 tube, closed a t one end and contemperature fractionation and chemical identification. taining a stopcock and ground joint a t the other end, leading in Spectroscopically pure chlorotrifluoromethane was mixed with turn by way of the joint to a Housekeeper metal-to-glass seal and iodine heptafluoride in the reaction tube (1, in Figure 3) containthenck t o a supply of the heptafluoride. Samples of liquid iodine heptafluoride (abobt 0.5 ml.) were condensed in this tube and mixed intimately with the solid alkali fluoride sample by shaking. The solubility of these alkali fluorides in liquid iodine heptafluoride was inappreciable. When the iodine heptafluoride was later allowed to evaporate a t room temperature and the Pyre\ glass No. 774 tube and contents reweighed, in no case was any sig-4 nificant gain in weight observed. I n the case of rubidium chloride, a loss in weight was noted to a partial conversion to fluoride I n no case was there any noticeable heat of reaction. From these results it was concluded that no stable addition compounds of the type M(IF8) had been formed a t room temperature. REACTIONS WITH CHLOROFLUOROCARBONS. Because of a certain degree of hazard connected with the transportation and handling of compressed fluorine gas, attention has been given recently to the use of halogen fluorides as fluorinating agents to replace elementary fluorine. In some instances the use of such halogen fluorides results in simultaneous chlorination or bromination as well as fluorination and this may or may not be a desirable consequence. Banks, EmelBus, and eo-workers ( 1 ) showed that carbon tetraiodide reacts with iodine pentafluoride to form good yields of trifluoroiodomethane, CIFP; while a t 90 O C. carbon tetrabromide was converted t o the extent of 83% to dibromodifluoromethane, with small amounts of the mono- and tetrafluoro compounds also present. The fluorination of carbon tetrachloride has been investigated by various experimenters both with elementary fluorine or with fluorine in the presence of a catalyst. For example, when carbon tetrachloride reacts a t 70" C. with fluorine in the presence of arsenic, carbon tetrafluoride and chlorotrifluoromethane are formed in 20 and 54% yields, respectively (6). When bromine, BS well as arsenic, is added to carbon tetrachloride and fluorine passed through, nearly pure trichlorofluoromethane is obtained Figure 1. Iodine Heptafluoride Reactor (p); and when antimony trifluoride is used as the catalyst, trichlorofluoromethane, dichlorodifluoromethane, and small quanti1. Column reactor, 24-inch length of '/s-inch nickel pipe ties of chlorotrifluoromethane are formed, but no carbon tetra2. Monel metal sphere, 5 inches in diameter 3. Hoke valve, blunt stem fluoride (10). With mercuric fluoride as the catalyst, fluorine 4. Soft copper gasket and a slight excess of dichlorodifluoromethane give high yields 5. Condenser housing, 12-inch length of 2.5-inch copper tubing of chlorotrifluoromethane at 340" to 370' C. but no appreciable 6. Nickel tubing and plate 7. Te%ongasket cmbon tetrafluoride (9). 8. Deville-type nickel condenser With iodine pentafluoride and carbon tetrachloride, Ruff 9. Fluorine inlet 10. Tap water inlet and outlet and Keim (6) obtained chiefly trichlorofluoromethane, with a 11. Copper outlet tube connected to Monel trap small amount of dichlorodifluoromethane. When trichloro12. Nickel clamp plate a i t h 6 bolts
8
July 1950
I N D U S T R I A L A N D E N G I N E E R I N G CHEMISTRY
1385
ence of a roll of copper screening a t loo", 200', and 300' C. In some preliminary trials the yields of chlorotrifluoromethane obtained were small and the copper screening was found to be considerably attacked, presumably through reaction with iodine and subsequent conversion t o copper fluorides. (It was observed that under similar conditions, with fluorine instead of iodine heptafluoride in the reactor, the copper screening received only a slight surface coating of copper fluoride.) With the copper screening removed, no explosions were encountered and conversions of dichlorodifluoromethane t o chlorotrifluoromethane of from 4 to 9%, based on the iodine heptafluoride used, were obtained a t 340' to 370' C. When a bed of mercuric fluoride powder as a catalyst was placed along the length of the reactor, yields of 62 to 68% were realized. Mercurio Figure 2. Purification and Sampling Apparatus fluoride alone was found capable of replacing 1. Aluminum phosphate Raman tube 8. Copper to Pyrex glass No. 774 (House2. Threaded braas sleeve attachment keeper) seal one chlorine with fluorine, but the yields 3. Standard flare fitting, s/s-inch brass 9. Monel traps were lower than when the mercuric fluoride 4. Brass infrared cell 10. To mercury manometer 5 . Brass square yoke attachment 11. To Hyvac pump was used in conjunction with iodine hepta6. Teflon gasket 12. Nitrogen purge line 7. Silver soldered copper sweat coupling fluoride. Table I summarizes the data a n d results obtained with the fluorination ing a roll of copper screening (not shown) maintained a t 100" C. of dichlorodifluoromethane by means of iodine heptafluoride intervals up to 400' C., above which the decomposition into fluowith mercuric fluoride as a catalyst. The percentage yields rine and iodine pentafluoride is appreciable. The rate of inare calculated, first, from the weight of chlorotrifluoromethane obtained when iodine heptafluoride is converted troduction of the chlorotrifluoromethane was determined by a flowmeter and varied from 2 to 10 ml. per minute. The iodine to iodine, and secondly, from the volume of liquid chloroheptafluoride WM entrained in a stream of nitrogen, dried by passage over phosphorus pentoxide (13, in Figure 3). The rate of entry of the heptafluoride was regulated by immersing the Monel trap, 11, in an ice-water bath and by controlling the rate of flow of the nitrogen through the sulfuric acid bubbletube, 12. The duration of the reactions was generally 6 hours; the weight of heptafluoride used in this period varied from 1.5 to 10.0 grams. After leaving the reactor, 1, the vapors were washed in 1 N sodium hydroxide solution and the products were condensed separately and successively distilled in the series of glass traps, 21. The volumes of liquid condensate were measured in a calibrated 15mm. Pyrex glass No. 774 tube, 9. A liquid density of 1.6 grams per cubic cm. was used to estimate the quantity of product obtained. The weight of iodine heptafluoride used was determined by the loss of weight of the Monel trap, 11. OI' The spectra of the condensates collected by cooling the traps in liquid Figure 3. Apparatus for Reactions of Iodine Heptafluoride with Some Freons air showed no traces of carbon tetra1. Reactor, 12-inch length of 10-inch nickel pipe 10. Mercurv seal fluoride, indicating that no fluorina2. Thermooouple well 11. Monel irap containing iodine heptafluoride tion of chlorotrifluoromethane had been 3. Teflon seal 12. Sulfuric acid bubbler for nitrogen 4. Brass yoke clamp 13. Phosphorus pentoxide drying tube achieved by this reagent. 5. Copper t o Pyrex glass No. 774 (House14. Flowmeter for Freoqs keeper) seal Dichlorodifluoromethane, analyzed 15. Rubber sleeves for joining Pyrex glass No. 6. Pyrex glass No. 774 spherioal joints 8 18/9 774 and copper tubing spectroscopically and shown to contain 7. Sodium hydroxide scrubbing tower full of 16. Tank of Freon Pyrex glass No. 774 beads 17. Liquid nitrogen trap less than 0.3'% of chlorotrifluoro8. Pyrex glass No. 774 infrared tube with po18. hlitrogen inlet tassium bromide window 19. To mercury manometer methane, w&snext employed with iodine 9. Calibrated tube for estimating yielda of 20. To H vac pump heptafluoride in the reactor in the presFreons 21. Fractronation traps ,
U
1
0
-
7
o\
13C6
INDUSTRIAL AND ENGINEERING CHEMISTRY
REACTIONS OF DICHLORODIE'LI OROWITII IODlNE HEPTAFLUORIDL
TABLEI. DATA FOR METHANE
Run No. 1
2
3 4 5
Duration, Hr. 6 6 5 5 5 6
(MeiLuric fluoride catalyst) CCiIFz CCiI, Iodine RecovRecovHeptaReaction ered. ered, fluoride %yield Based on Temp., T'ol. Lit[., Vol. Liq., Used, CChF, C. hfl Mi, G. IFr+Ie useda 3 0 340 2.0 9.97 8.6 40 350 360 350 360
1.8 6.0 5.3 1.8 0.9 2.0 1.0
1 .? 1 2 1 0
2 5 2.5
25.95h 7.35 5.95 1.95 1.62
...
42
4.5 67
15 5g 75
7.1
14
6 67 370 6 0.6 24 7c 6 370 0 9 46 6 370 8d a The apparent pairing of the yields i n this ooliiiiin is t o be largely explained by a similarity in flow rate for t h e CCLFZ. b A marked increase in the rate of flow of IF, occiirred in t h e last few minutes of this run. C The same sample of HgFz wits ii-ed for this run as was used i n reactions 5 and 6. d A fresh sample of HgF2 was used for thiq run.
trifluoromethane divided by the sum of the volumes of dichlorodifluoromethane and chlorotrifluoromcthane obtained. The validity of such a method of estimation is dependent upon the assumption that the two liquids have nearly equal densities
Vol. 42, No. 7
are riot very conclusive. Likewise, in its reaction with dichlorodifluoromethane, iodine heptafluoride appears t'o be even mow potent than fluorine itself; but it is possible that this effect may result from the cat,alyt,ic action of iodine upon the substitutioii react ion, 4. The fact that carbon tetrafiuoride appeared as a produvt in the fluorination of carbon tetrachloride by means of fluorine (6),but was absent in the reaction of fluorine with dichlorodifluoromethane (9), is but one indication of the increasing strength of the C C I bond in the more completely fluorinated members of the series of compounds derived from carbon tet,r:tchloride by successive fluorination, reaching a maximum i n chlorotrifluoromethane, as referred to in the second point abovo. These facts may be interpreted to imply that the mechanism of t.he suhstitution of fluorine for chlorine in this case may not be a simple, st,epwise process, hut ma proceed by way of formation of free radicals, which produce czains of carbon atoms, subsequerit,ly hroken in the fluorination reaction. ACKNOWLEDGMENT
This investigation was materially assisted by it grant-in-aid received from the Harshaw Chemical Company of Cleveland, Ohio. The authors wish to express their appreciation for this assistance. %he authors also wish to acknowledge the aid oi' Kinetic Chemicals, Inc., of Wilmington, Del., from whom the spectroscopically pure samples of Freon-12, CC12Fz, and Freon-13, CCIF,, were obtained.
CONCLUSIONS
From the results of earlier investigators previously referrell to and from the results of the present study, the following conclusions may be drawn: 1. The substitution of fluorine for chlorine i n carbon tetrachloride is accomplished more effectively by fluorine than by iodine pentafluoride. 2. It is recognized that the single chlorine atom in chlorotrifluoromethane is less susceptible to replacement than thoqc of dichlorodifluoromethane. The replacement of chlorine in dichlorodifluoromethane appears t o require a catalyst, such as mercuric fluoride (9). 3. From the facts that iodine pentafluoride has been shov,n able to convert a few per cent of trichlorofluoromethane to dichlorodifluoromethane, while iodine heptafluoride ronverts a few per cent of dichlorodifluoromethane to chlorotrifluoromethane, it may be argued that the heptafluoride is more potent as a fluorinating agent than the pcntafluoride; but the results
LI'I'EHATUHE CITED
H m k s . FmelBus, Haszeldine, and Kerrigan, J . Chem. Soc., 1948, 2188. Booth and Pinkston, Chern. Rets., 41, 421 (1947). Lord, Lynch, Sohumb, and Slowinski, J . Am. C h e m Suc., 72. 522 (1950). JIoiasan, Comp. rend., 135, 563 (1903). Ruff and Keini, 2. anorry. ZL. allyem. Chem., 193, 176 (1930). Ibid.. 201, 245 (1931). Schumb, Young, and Radimer, 1x1).EXG.CHEW,39, 244 (1947). Yharpe and EmelBus, J . Chern. Soc., 1948, 2135. Simons. Bond. and h1cArthur. J . Am. Chmn. Suc., 62, 3477 (1940).
Smith, A n n . R e p f s . on Pioyresu Chem. (Chem. SOC. I J o n d n r l ) , 44, 86 (1947). 'Thompson and T e m p l e , J . CherrL. Soc., 1948, 1422.
K E C E I V ENovember ~ 2 5 , 1949.
FRACTIONATOR CALIBRATION SYSTEM Binary vs. Afulticomponent Test Mixtures C. B. KINCANNON AND EARL 3IANNINC.
JH.
Shell Oil Company, H o u s t o n , T P X . T h e effect of the number of components present in the calibration system used in the calculation of plate efficiency of a distillation column was investigated on a 30-plate 1inch inside diameter glass Oldershaw column. Analyses were performed by use of a mass spectrometer. The plate efficiencies were calculated from a multicomponent fraction and compared with those for a binary system. Results from both mixtures show an average efficiency- value of 50% (fifteen theoretical plates) over a load range of 800 to 3000 ml. per hour overhead rate. It appears that the efficiencies of laboratory columns, evaluated with systenls of like molecular weight and type, are independent of the number of components present in the test mixture.
T IS ofteii desirable to l c n o ~the efficiencyof a fractionating column when it is io productive service and operating on a multicomponent system. However, the efficiency is usually determined by experimental data and calculations based on a binary test mist,ure. The differences in the number wf degrees of freedom and the rate of approach to thermal and phase equilibrium could make efficiency values based on multicomponent systems appreciably different from those based on binary systems. Therefore, an investigation of the efficiency of a laborat,ory column was iriitiat,ed in \vhicli all variables could be held constant except the number of components in the system. The column selected for use in the study of the effect of calibration mixture was a %plate 1-inch inside diameter glass per-
