Fluorine Polyhalides

Perhaps the reason for this apparent oversight of fluorine has been the aver- sion of chemists to the handling of fluorine and the fluorides. Accordin...
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FLUORINE POLYHALIDES * BY HAROLD SIMMONS BOOTH, CARL F. SWINEHART AND WILLIAM C. MORRIS

Introduction Many polyhalides of the various elements, chiefly of the alkali metals of the first group, have been prepared. Numerous methods have been used to prepare the polyhalides, and there is question as to whether or not some that are said to exist really do. But chemists seem to have forgotten that fluorine is a halogen; all of the studies have been made on the polyhalides containing chlorine, bromine, and iodine, in various combinations and proportions. Perhaps the reason for this apparent oversight of fluorine has been the aversion of chemists to the handling of fluorine and the fluorides. According to Sidgwick,‘ “Any halogen except fluorine can form part of a perhalide.” He assumes that fluorine does not follow the behavior peculiar to the polyhalide formation because it is the lightest halogen and only the heavier ones follow this rule. The purpose of this paper is to describe the preparation of the salts formed by the union of iodine trichloride and the fluorides of the alkali metals, cesium, rubidium, potassium, and ammonium, their properties and the experimental data for the preparation and analyses of the salts.

Historical It has long been known that iodine would dissolve in potassium iodide solutions, either alcoholic or aqueous, but it was not known whether it was merely a physical solution or whether a complex polyhalide of potassium was formed. There was some evidence for both lines of reasoning. BaudrimonP found that carbon disulfide removed the iodine dissolved in an aqueous solution of potassium iodide. This would indicate that no chemical compound was formed. JorgensenSsaid that carbon disulfide did not remove iodine from an alcoholic solution of potassium iodide containing two atomic weights of iodine for each atomic weight of potassium iodide. He also showed that an alcoholic soluti’on of potassium iodide completely removed the iodine from a solution of it dissolved in carbon disulfide. This would indicate that a compound was formed by the iodine uniting with the potassium iodide. Since this early work there have been many complex halides isolated containing not only iodine but also chlorine and bromine. All of the polyhalides

* Contribution from the Morley Chemical Laborator8 Western Reserve University. First ublished in art a8 a “Note to the Editor” J. Am. hem. SOC., 54, 2561 (1932). 1 &cIgwick: Electronic Theory of Valence,” 293 (1927). 2 Baudrimont: Compt. rend., 51, 825 (1860). 3 Jorgensen: J. prakt. Chem., 2, 347 (1870).

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that have beenprepared may be divided into two major classes; those containing three halogen atoms and those containing five halogen atoms. Table I contains a list of all of the alkali metal polyhalides, arranged in the order of increasing stability. This stability is based on the temperature to which it is necessary to heat the compound in an open tube before it whitens. I n general the stability is as follows: Cs > R b > NH4> K > Na > Li.

Trihalides By concentrating an aqueous or alcoholic iodine-potassium iodide solution over sulfuric acid in a desiccator Johnson’ obtained some large, dark blue, prismatic crystals which he claimed, from the result of his analyses, to be potassium triiodide. Wells2 and some of his co-workers repeated Johnson’s preparation but not the analysis. Abegg and Hamburgera attempted to determine the composition of all the alkali polyiodides that are stable a t 25OC. by shaking the solid iodide in varying amounts of a nearly saturated solution of iodine in benzene. They found no evidence of the compound KI3, Several phase rule studies of the system KI-IZ-HZO have been made but none of them even indicated that K13 existed. Bancroft4 made a phase-rule study of the system and found no evidence for the existence of KI3. On repeating Johnson’s work he obtained the same salt and, by using the same inaccurate method of analysis that Johnson used, he found that the iodine content of the salt corresponded with that required for K13. On making an X-ray study of the salt he found that there was no indication that the salt was potassium tri-iodide. Bancroft concluded from his study of the salt that it was merely potassium iodide that had taken up enough iodine to change its crystal structure slightly and to give the analysis corresponding to K13. Briggs and Geigle5made a freezing point study of various mixtures of potassium iodide and iodine and found that potassium iodide and iodine were the only solids that could exist in equilibrium with the melt. Serullas6 prepared potassium dichloro-bromide and potassium dibromochloride by the action of chlorine monobromide on potassium chloride and potassium bromide. Abegg and Hamburger (loc. cit) found evidence for the existence of RbI3, csI3, and “413. Foote and Chalker’ claim to have isolated the tri-iodides of cesium, rubidium, and potassium. Wells and Wheeler (loc. cit.) have prepared the following trihalides in addition to the potassium triiodides: CsBr12, CsBrJ, CsClBrI, CsC121, CsBro, CsC1Br2, CsClzBr, csI3, Rb13, RbBr21, RbCIBrI, RbBra, RbClZBr, RbCIBrz, KBr 21, and KClzI by dissolving the normal halide and the halogen in question in a warm water or Johnson: J. Chem. SOC.,31, 249 (1877). Welk: “Studies from the Chemical Laboratory, SheEeld Scientific School,” Vol. I, Inorganic Chemistry. 3 Abegg and Hamburger: Z. anorg. Chem., 50, 427 (1906). Bancroft: J. Phys. Chem., 35, 764 (1931). 6 Briggs and Geigle: J. Phys. Chem., 34,2250 (1930). 8 Serullas: Ann. Chim. Phys., 45, 190 (1830). 7 Foote and Chaker: Am. Chem. J., 39, 561 (1908). 1 2

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FLUORINE POLYHALIDES

weak alcohol solution. They were unable to prepare the other trihalides necessary to complete the series. The colors range from brilliant black for CsIa to bright yellow for CsClnBr. They found that the trihalides containing iodine were more stable than those not containing iodine, but that the instability of the compound was not governed entirely by the volatility of the halogen, for they found that CsClzBr was more stable than CsCIBrz. Several of the ammonium trihalides have been prepared. They are very similar to the alkali metal trihalides and are prepared in the same way. Johnson1 obtained NH41s by passing ammonia gas into the mother liquor from which the potassium tri-iodide was crystallized. Besides the normal trihalide he also obtained two salts having the composition NH41a.7KIand 5NH413.KI respectively. Roozeboom*prepared ammonium tribromide, NH4Brs. Chataway3obtained the following trihalides, NH4C121,NH4ClBrI,and NHaBr3and found that the first one was the most stable. Ray and Sarkar' studied the formation of aqueous solutions of HCIIz, HClBrZ, and HBrIz by means of the distribution of halogen between an aqueous solution of the acid in question and chloroform, carbon disulfide, or benzene. They found that HBrIz was the most stable and that the other two were about as stable as the corresponding potassium salts. Cremer and Duncan6found evidence for the existence of HIBrz, HIC12, and HIBrCl in solution.

Pentahalidis of the Alkali Metals

I

Wells and Wheeler (loc. cit.) also prepared the following pentahalides: CsI6, CsBr6, CsCI.ICls, RbCI.ICla, KCI.ICls, NaCI.ICls.2H20, and LickICls.4HzO. The method of preparation is very similar to that used in the preparation of the trihalides. The first five are fairly stable while the last two are very unstable. According to Wells and Wheeler, when the pentahalides are heated they apparently lose halogen and go to the trihalide since the pentahalides all whiten a t approximately the same temperature as the corresponding trihalides. On that basis are they listed in Table 1. There are some higher polyhalides described in the literature but there is some doubt about their existence. Abegg and Hamburger found evidence for the existence of KI,, RbIT, Rb4, Cs1.1,and CsIs. Foote and Chalker claimed to obtain KI, but none of the other polyiodides with more than five atoms of iodine. Rae6 has prepared CsBr4 and CsBrs and thinks that the tetrabromide was the real compound discovered by Wells and Wheeler and called pentabromide by them. By running the isothermal diagram for the system CsI Iz - HzO, Briggs, Greenwald, and Leonard' found the CsIs and cSI4 but did not find any evidence for Cs15.

-

Johnson: J. Chem.

Soc., 33, 397 (1878).

* Roozeboom: Ber., 14,2398 (1881). * Chataway: J. Chem. SOC., 107, 105 (1915). Ray and Sarkar: J. Chem. Soe., 121, 1449-55 (1922). Cremer and Duncan: J. Chem. SOC., 133, 1857-66(1931). Rae: J. Chem. SOC.,133, 1578-81 (1931). ' Briggs, Greenwald, and Leonard: J. Phys. Chem., 34, 1951-60(1930).

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TABLE I The Alkali Metal Polyhalides (Arranged in order of increasing stability, based on the temperature at which they whiten when heated in an open tube) Trihalides Formula Whiten a t Degrees C

RbCIBrz RbClzBr RbBr3 CsCIBrz CsCllBr CsBrs KBrJ RbClBrI KClzI KIs CsBrIz RbClJ RbBr21 RbI3 CSCIZI CsClBrI CsBrZI CSI8

80' I IO0

Formula

Pentahalides Whiten a t Degrees C

NaC1.1C13.2H20 LiC1.ICl3.4H20

I 15'

180'

I400 I500 I500 I 60'

CsBr5

180" 200'

215'

KCl.IC13

2259 260'

265'

RbCl.IC13

265' 270'

290'

CsCl.IC13

290'

320'

330' CSIs (The data on the pentahalides are uncertain)

Polyhalides containing Fluorine

At the time these studies were begun no evidence of the existence of polyhalides containing fluorine had been adduced. Cremer and Duncan' in studying the dissociation pressures and other properties tried the absorption of iodine bromide by cesium fluoride and found evidence for a compound having the formula CsFIBr. I t should be very similar in its composition and properties to the cesium chloride-iodine-monobromide. No description or analysis was given and apparently they did not isolate the salt. Preparation of Fluorine Polyhalides For the preparation of the fluorine polyhalides four procedures were used. I n the first method chlorine was passed into a hot saturated solution of the normal fluoride, containing one equivalent weight of iodine for each atom of fluorine, until all of the iodine color had disappeared. On cooling the solution, small orange-yellow, needle-like crystals were obtained, but these were not all

* Cremer and Duncan

J Chem. Sac., 133, 2243-54 (1931).

FLUORINE POLYHALIDES

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homogeneous. By careful recrystallization this salt gave a crop of homogeneous crystals. However, this method gave a very poor yield. The second method tried was practically the same as the first except that the solution of the fluoride was slightly acidified. This gave a much better yield and the crystals were homogeneous when examined under the microscope. In both the above methods of preparation the chlorination was accomplished in an Erlenmeyer flask of appropriate size, fitted with a reflux condenser to return both water and any volatile iodine-chloride formed. The flask was weighed before and after chlorinating and the amount of chlorine absorbed gave some idea as to the compound formed. I n the third or dry method, the normal fluoride was intimately mixed by grinding with one equivalent weight of iodine for each atom of fluorine, and then chlorinated to constant weight in a flask similar to the above with ahlorine dried by concentrated sulfuric acid. A dark brown pasty liquid first appeared which gradually changed over to a bright yellow solid. The intermediate product was formed fairly rapidly and probably was a eutectic mixture of iodine monochloride and the fluoride. The final product is formed very slowly. On recrystallizing this solid, very good yields of the same bright yellow, needlelike crystals previously described were obtained. However, this method is too slow. The fourth and best method found for the preparation of the fluorine polyhalides was to add a slight excess of a cold saturated solution of iodine trichloride to a cold, saturated, slightly acid solution of the normal fluoride. On mixing the two aqueous solutions the characteristic brilliant yellow crystals separated out immediately. After recrystallization, these were entirely homogeneous. A much better yield is obtained by this method than by any of the others. The iodine trichloride was made by subliming iodine from a glass retort into an ice-cooled flask similar to the other chlorinating flasks while a stream of chlorine in excess of that required to form IC13 was sent through it. The bright yellow iodine trichloride deposits on the cold bottom and sides of the flask. This method of forming IC13 is much better and quicker than that of chlorinating the solid iodine to the trichloride or to constant weight.

Preparation of the Cesium Fluorine Polyhalide Cesium chloride was converted to the fluoride by first heating it with an excess of sulfuric acid ; this bisulfate was then converted to cesium hydroxide by treating the hot solution of the sulfate with a hot saturated barium hydroxide solution, which precipitated the sulfate as barium sulfate and converted the cesium into cesium hydroxide. By careful addition of the barium hydroxide solution the point was reached a t which there was neither sulfate nor barium in solution. The barium sulfate was filtered off and the cesium hydroxide solution was evaporated until the cesium hydroxide or cesium carbonate began to separate. This was then neutralized with hydrofluoric acid. On mixing an aqueous solution of the iodine trichloride with a slightly acid solution of cesium fluoride, the brilliant yellow needles of the cesium polyhalides separated.

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HAROLD SIMMONS BOOTH, CARL F. SWINEHART, WILLIAM

c. MORRIS

The cesium polyhalide is much more stable than the potassium salt, but it does lose halogen slowly in the open. For this reason it is best kept, in tightly stoppered bottles. Analyses of this salt (Table 11) shows that it has the formula CsFIC13.

TABLE I1 Results of Analyses of Salts

Cesium salt, Specimen 1 Cesium (perchlorate method) : 34.27%,

Theoretical for CsFIC13 34,33% 34.49% c s

Cesium salt, Specimen d Cesium (by gentle ignition) : 34.59%, Average (both specimens) Chlorine: 27.61%, 28.06% Average ,, Iodine: 33.38%, 32.81%

34.56% 34.44% 27.84% 33.09%

34.49% cs 27.62% C1 32.96% 1

95.37 4.63%

4.93%F

Fluorine by difference :

Rubidium salt Rubidium: ~ 5 . 2 7 % ~ ~ 5 . 4 3 7 ~ Average 1, Chlorine: 3 1 . 4 6 % ~ 31.52% ,> Iodine: 3 7 . ~ 2 % ~ 38.05%

13.31% 36.46% 43.67%

Average 1, 9,

13.34 36.53 43.62 93.49 6.51

Fluorine by difference

Ammonium salt NH3 calc. to NHaF: 1 3 . 5 7 % ~ Chlorine: 39.39%, Iodine: 46.7970,

25.35 31.49 37.78 94.62 5.38

Fluorine by difference

Potassium sa& Potassium: 13.37%, Chlorine: 36.60%~ Iodine: 43.57%,

Theoretical for RbFIC13 25.30%Rb

31.50%C1 37.58%1

5.62% F

Theoretical for KFIC13 13.42% K 36.51%C1 43.75%1

6.32%"F

Theoretical for NHhFIC13 13.53% 39.32% 46.82%

Average 1,

7,

13.55% 39.36% 46.81% 99.72%

13.68% 39.44%CI 46.90% 1

FLUORINE POLYHALIDES

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Preparation of the Rubidium Fluorine Polyhalide Starting with rubidium chloride, the rubidium fluorine polyhalide was prepared in the same way as the cesium salt. The rubidium fluorine polyhalide is similar to the cesium salt. I t is much more stable than the corresponding potassium salt but not quite as stable as the cesium salt. Analyses (Table 11) show it has the formula RbFIC13. Preparation of the Potassium Fluorine Polyhalide The potassium fluoride used was obtained from the pure acid fluoride either by fusing and driving off hydrogen fluoride leaving the normal fluoride, or by neutralizing the acid fluoride in solution with potassium carbonate. The polyhalides of potassium were prepared by all four methods described above although method 4 was best. The salt is so unstable that at first no consistent analyses could be obtained. However, it was found that the loss of halogen occurred during separation of the crystals from the mother liquor and drying. This was obviated by collecting the crystals in a covered Jena fritted glass Gooch crucible and immediately centrifuging at high speed still keeping the potassium salt covered in the crucible. The salt was immediately transferred to tiny weighing bottles which were almost completely filled by the sample (to avoid loss of Ic13). Excellent analyses, establishing the formula as KFICl,, were obtained in this fashion as shown in Table II. Preparation of the Ammonium Fluorine Polyhalide Pure ammonium bifluoride was almost completely neutralized with ammonia and to this cold solution a cold saturated solution of iodine trichloride was added. The yellow crystals were collected and separated from mother liquor centrifugally as in the case of the potassium salt. The analyses (Table 11) showed the salt to have the formula NH,FICl,. Trihalides containing Fluorine In addition to the fluorine pentahalides, an attempt was made to prepare the fluorine trihalides of the alkali metals by mixing a solution of the fluoride with a solution of iodine monochloride, by treating the solid fluoride with liquid iodine monochloride, and by chlorinating a suspension of iodine in an aqueous solution of the fluoride, both hot and cold. If any compound was formed i t was too unstable to be isolated. Methods of Analysis On all the salts a complete analysis was made, save for fluorine which was determined by difference. In the first cesium fluorine polyhalide specimen prepared the cesium was determined as cesium perchlorate and in the second as cesium fluoride. It was found that the same or better results could be obtained by the latter method of gently igniting the polyhalide, whereupon all the cesium was left as cesium fluoride. This latter method was also used on the rubidium and potassium salt. In no case did the residue after ignition contain even a trace of chlorine or iodine.

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HAROLD SIMMONS BOOTH, CARL F. SWINEHART, WILLIBM C. MORRIS

The chlorine and iodine were determined by the method of Gooch. In this method two sets of samples are run, one in .which the chlorine and iodine are both precipitated as the silver salts, and in the other set only the chlorine is precipitated as the silver salt. In both cases the samples were gently heated with dilute ammonium hydroxide solution to convert the chlorine and iodine into the chloride and iodide. The first set of samples were then acidified while the second were treated with dilute sulfuric acid and potassium nitrite t o expel the iodine. The solution was boiled gently until the disappearance of the iodine color. The silver halides were precipitated and the chlorine determined in the usual way. If equal weight samples are used in both cases, the difference in weight in the silver salts represents the weight of silver iodide in the former case. This method is not quite as satisfact,ory as might be desired, due to the fact that there is danger both of not expelling all of the iodine and also of driving off some of the chlorine. The ammonium salt was analyzed for NHz by distilling with excess alkali, absorbing in standard acid and backtitrating; and the NHs was reported as NH4F. Properties The cesium, rubidium, potassium, and ammonium fluorine-iodinetrichlorides have very much the same physical properties; they all have the same orange yellow color, crystalline form, and general properties but differ in their stability and the temperature a t which they decompose. When the normal fluoride solution is mixed with the solution of iodine trichloride, the pentahalide separates out as brilliant, orange-yellow, needle-like, tetragonal prisms exhibiting parallel extinction under crossed nicols in two positions and isotropic in the third. The cesium salt is the most stable. When heated in an open tube it melts and then decomposes well above 3ooOC.; although there is a slight decomposition below that temperature. When heated in a sealed tube it melts at 194' but on cooling there seems to have been no decomposition. The specific gravity of the CsFIC13 is 3. j6.5. The rubidium salt is much more stable than the corresponding potassium salt but is not quite as stable as the cesium salt. When heated in an open tube the rubidium pentahalide melts and decomposes a t approximately 300'. When heated in a closed tube, RbFIC13 melts without decomposition at 172'. The specific gravity of the rubidium polyhalide is 3.1 59, The potassium salt is the least stable. When the mother liquor from which the alkali metal fluorine polyhalide is precipitated is evaporated to dryness there is very little residue left. This shows that the fluorine pentahalides are not very soluble in water. The polyhalides are not soluble in and do not react with carbon tetrachloride and are insoluble in benzene but react with it if allowed to stand in contact with it for a long period of time.

FLUORINE POLPHALIDES

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Structure The structure of IC13 may be represented by

:c1: . X

:I :a: .. . X

c1: ..

:

in which the central I atom has ten electrons. However, according to Sidgwick, a decet is not a particularly stable arrangement, and it is not surprising that I c l , tends to combine with the alkali fluorides to form pentahalides. In this process the fluorine probably acts as a donor and the iodine as an acceptor increasing the electronic environment of the iodine to the more stable duo-decet, as follows:

:el: .x

x

: c1: . X

. X

* X

:el:

:el:

..

The great stability of the duo-decet is revealed in such a compound as SFS. As would be expected, the most stable polyhalide is formed with the alkali of highest atomic number, where the greater distance of the valence electrons from the nucleus facilitates the donation of two of the fluorine electrons. Undoubtedly fluorine polyhalides of this same type with other metallic elements could be prepared by similar methods. Thermal analysis of the system HF-ICl, may also reveal new compounds. Of this same type would be compounds formed by the union of BrFa with MX, yielding MFBrF3, MClBrF,, etc. Iodine pentafluoride should combine with alkali halides to form KXIF, in which two of the electrons of the iodine atom would be inert as in

,. .. :F: :F:

. X

. X

:F: :F:

.. ..

Even iodine heptafluoride might combine to yield compounds of the type MXIF8, since the stable shell of sixteen electrons is known in such a compound as OSFS. Investigation of these predicted compounds is now well under way in this laboratory, and will be reported shortly.

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HAROLD SIMMONS BOOTH, CARL F. SWINEHART, WILLIAM C . MORRIS

Summary I t was found that the most satisfactory method of preparing the fluorine polyhalides of the type formula MFICl3was to mix a saturated solution of the alkali fluoride with a saturated solution of iodine trichloride, both solutions being cold a t the time of mixing. The cesium and rubidium salts are both fairly stable, the ammonium salt less, and the potassium least, though over a long period of time the cesium and rubidium salt lose halogen unless kept in tightly stoppered bottles or in a desiccator containing a few crystals of iodine trichloride. The fluorine polyhalides are insoluble in benzene and carbon tetrachloride but react with the former slowly. The solubility in water is limited but is greater with the ammonium and potassium salts. The four polyhalides prepared form orange-yellow tetragonal crystals exhibiting parallel extinction. It is established that fluorine can form part of a polyhalide. Cleveland. Ohio.