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The Halogen Fluorides[6]
The Halogen Fluorides'^
The formation o f interhalogen compounds has been known for well over a hundred years. Those containing fluorine are, however, relatively new and are still the subject o f consider able research. The following table shows the range o f compounds known at present. Type
AB
AB3
AB5
ClF(-lOr)
αΤ^ΪΪΤ)
BrF(20°)
BrF3(126^)
ci?^ BrF5(4r)
BrCl
IF3
IFsilOr)
ICl
ICI3
AB7
IF7(277%subl.)
IBr
T w o points are immediately apparent from this table. In the first place fluorine is able to excite the higher valencies o f the other halogens and, secondly, the halogen fluorides, the boil ing points o f which are shown in parentheses as far as they are known, are comparatively volatile. All o f these substances are formed by direct union o f the elements, with some variation in reaction conditions to form a particular compound when alternative combinations are possible. Thus chlorine monofluoride results when equimolar proportions o f the two elements are passed through a nickel tube at 250°. A n alternative method is to allow chlorine to react with the trifluoride. With a higher proportion o f fluor ine combination o f the elements will give the trifluoride, while the pentafluoride is obtained only by using a large excess o f fluorine at 350° and a pressure o f the order o f 250 atmos pheres.'^^ The formation o f the bromine fluorides occurs rather similarly, except that the pentafluoride is easier to 11
12
THE CHEMISTRY OF FLUORINE
prepare, in keeping with the more basic character o f the heavier halogen. Iodine monofluoride is unknown, though bands associated with the I F molecule are observed in the emission spectrum o f the pale green flame o f iodine burning in fluorine. The main product is the pentafluoride: the heptafluoride results when the vapour o f the pentafluoride is allowed to react with excess o f fluorine at 250-270°. Iodine trifluoride is reported as formed when fluorine is passed into a solution o f iodine in a Freon solvent at - 7 8 ° . It is unstable and disproportionates to iodine and the pentafluoride at above about - 3 0 ° . N o t surprisingly the interhalogen compounds show a gen eral similarity in their reactions to the halogens themselves. Those containing fluorine are extremely reactive as fluorinating agents and, in this respect, they are rather similar to the element. There are however difl'erences in reactivity between the various members, a tentative order o f decreasing reactivity being CIF3
> BrFs >
IF7
> CIF > BrFj >
IF5
> BrF.
Bond energy data are, however, largely lacking. The structures o f these molecules is a matter o f consider able interest. Chlorine monofluoride has a dipole moment o f 0.88 0.02 Debye, the dissociation energy being 60.2 kcal/ mole. There is believed to be a true σ-bond with the polarity C I — F . The dipole moment o f bromine monofluoride is higher (1.29 Debye), though the dissociation energy is unknown. Both chlorine trifluoride and bromine trifluoride are T-shaped molecules in the gas phase, the dimensions being as shown below 1.698Ä
F^^Cl 1.598Ä I
83^1
1.810Ä
F-^—Br 1.721Ä
F F These structures are derived from the trigonal bipyramid and there are two pairs o f electrons occupying vacant equatorial
THE HALOGEN FLUORIDES
13
sites. The dimensions in each case are based on the micro wave spectrum and the structures are supported by the results of infrared and Raman studies. The '^F nuclear magnetic resonance spectrum o f chlorine trifluoride shows that fluorine exchange occurs at room temperature and it seems likely that the mechanism involves a dimeric form o f the molecule. The structure of the solid is the same as that o f the gas, the lattice being an assembly o f CIF3 molecules. Chlorine pentafluoride, bromine pentafluoride, and iodine pentafluoride all have structures based on a square pyramid with fluorine at the apices and the heavier halogen atom somewhat below the basal plane containing four fluorine atoms (see Fig. 1).
Figure 1.
This arrangement leaves one pair o f electrons to occupy the vacant octahedral site. The structure o f iodine heptafluoride is still somewhat un certain. Relatively early work on the Raman and infrared spectra o f the liquid and vapor, together with electron dif fraction studies, led to general agreement on a pentagonal bipyramidal structure with fluorine at the apices and iodine in a symmetrical central site. Quite recently,'*^ however, the crystal structure o f the solid has been subject to a critical examination. The data are stated to be in the best agreement with a model based on a dodecahedron in which two atoms at the end o f the 4 axis have been allowed to coalesce. This leads
14
THE CHEMISTRY OF FLUORINE
Figure 2.
to an arrangement shown diagrammatically in Fig. 2. A t o m F ' and the four atoms F " are believed all to be at 1.825 Ä from the iodine atom, though the four F ' ' atoms are not in one plane. The bonds to the F ' " atoms are both 1,97 A in length, and are therefore assumed to be weaker. Such a structure relates the heptafluoride very simply to the pentafluoride. The '^F nuclear magnetic resonance spectrum o f the liquid at room temperature gave unexpectedly broad lines, the simplest explanation o f which is that the '^F nuclei are in non-equivalent positions. Unfortunately, however, overlap o f individual lines prevented more detailed analysis. It seems from the above that there is a difference between the structure o f the solid and that o f the liquid and gas. Alternatively one or the other o f these structures must be wrong and the problem must be reexamined. In view o f the structural evidence outlined above it is reasonable to suppose that bonding in the halogen fluorides is covalent. It is surprising, therefore, to ñnd unambiguous evidence for self-ionization in several species. Iodine chlorides and its bromide have been known for many years to be ionized when dissolved in polar solvents and the monochloride and monobromide also conduct in the molten state. T h e actual modes o f ionization are as follows: 2 ICl ^
Γ
2 ICI3 ^
ICI2 +
+ ICljICI4-
THE HALOGEN FLUORIDES
15
These schemes are well supported by the usual solvo-acid and solvo-base studies. In the case o f the halogen fluorides it is impossible to study conduction in polar solvents o f the usual types as the halogen compounds are much t o o reactive. Since both fused silica and platinum are unattacked, conductivities of the pure liquids can be measured in specially constructed cells provided great care is taken to avoid contamination. This approach has led to values o f the specific conductivities that are listed below: Compound CIF3
Specific conductivity (ohm~' cm"') 6.5 X 10-'
BrF3 BrFs IF5
8.0 X 10-^(25'') 9 X 10-^25") 5.4 X 10-^25°)
By analogy with iodine trichloride, the mode o f selfionization o f bromine trifluoride may be postulated to be 2 BrFa ^
BrF2^ + BrF4-
There is clear evidence for this which will be described later and also less complete evidence that chlorine trifluoride ionizes similarly. For the pentafluorides the most likely mode o f ionization is 2 AB5 ^ AB4^ + A B ö " and this hypothesis also receives considerable support. What we must look for in order to establish the reality o f these ionization schemes is a series o f compounds containing the cation or anion characteristic o f the solvent system and capable in solution o f enhancing its conductivity. Chlorine trifluoride is found to react with alkaU metal chlorides, bro mides, or iodides at 100°. A solution is obtained from which, after evaporating excess o f the interhalogen com pound and free halogen, solids o f the type M^C1F4 are ob tained ( M * = K , R b , C s ) . There may be formulated as ionic compounds containing the C I F 4 - anion which is believed to be characteristic of the solvent. Similarly, chlorine trifluoride forms 1:1 adducts with S b F j , B F 3 , and A s F s , and these may
16
THE CHEMISTRY OF FLUORINE
be formulated as containing the C1F2^ cation (e.g., ClFj^SbFö", C1F2-'BF4-). They are in fact ionic in the solid state and also in chlorine trifluoride solution and the infrared spectrum o f the boron trifluoride adduct also shows the presence o f the BF4" anion j ^ ' Bromine trifluoride has been more fully studied in this con nection. Various "bases" containing the BrF4" anion have been prepared and analyzed [e.g., K B r F 4 , Ba(BrF4)2], to gether with a number o f acids [e.g,, B r F j S b F g , (BrF2)2SnF6]. Both acids and bases dissolve without decomposition in bromine trifluoride to give conducting solutions and it is also possible to carry out conductometric titrations o f acids with bases and so to observe a neutralization process such as 2KBrF4 + (BrF2)2SnF6 = K2SnF6 + 4 BrF3 Very recently'*®^ a 1:1 complex o f N O F and C I F has been isolated which has been shown to contain the linear CIF2" anion, presumably belonging to the system 2 C I F ^ C r + CIF2", though there is as yet no further evidence on this point. Both chlorine trifluoride and bromine trifluoride are also misciblewith anhydrous hydrogen fluoride.'^Ionization in such solutions is very limited but it is reasonable to suppose that it occurs according to the equations CIF3 + HF ^
C1F2^ + H F 2 -
BrF3 + HF ^
BrF2^ + H F 2 -
These ionization schemes are supported by the occurrence o f neutralization reactions which may be followed by conducto metric titrations, e.g. (BrF2^ + H F 2 - ) -f (H2F-' + SbFö") — BrFzSbFe + 3HF (BrF2^ + H F 2 - ) -f ( K ^ + BrF4-) — KHF2 + 2 BrF3 The study of halogen pentafluorides in this connection has so far been confined to iodine pentafluoride, which reacts with potassium fluoride to give K I F e and with antimony penta fluoride to give ISbFio. Both o f these solids dissolve in iodine pentafluoride and enhance its conductivity and a conducto-
THE HALOGEN FLUORIDES
17
metric titration may also be carried out which corresponds with the equation IF6- + I F / SbFe" — KSbF^ + 2 IF5 It may well be that chlorine and bromine pentafluorides will show a similar behavior. There is some further evidence in support o f the ionization postulated for iodine pentafluoride. Boron trifluoride increases its conductivity and, when passed into a solution of potassium fluoride in the pentafluoride, K B F 4 results. It seems that the acid IF4"*^BF4- must be an intermediate here, even though it has not so far been isolated. Adducts o f iodine pentafluoride with several other species ( N 2 O 5 , M 0 O 3 , W O 3 , K I O 4 ) have also been isolated but it is not yet clear whether they are in any way related to the solvent system under discussion. Iodine heptafluoride, which has an unusually narrow liquid range, appears not to form compounds with potassium, rubidium or cesium fluorides. There have, however, been rereports o f the addition compounds I F v - A s F j and IFt SSbFs.''^^ It is tempting to formulate the first o f these as IFe^ AsFó", i.e. as an acid in the system 2 IF7 ^ IF^^ + IFg" and some support for this view comes from the reaction with potassium fluoride, which may be written as IFö^ AsFö" 4- K F — KAsFö + IF7 Several other points relating to the halogen fluorides re main to be mentioned. A l l are strong fluorinating agents for both organic and inorganic compounds.^'^^ This property has, on the whole, been much less used than might have been ex pected in view o f the relative ease with which the compounds may be prepared and stored. A substance such as chlorine trifluoride, for example, will fill much the same role as fluorine itself in forming the higher fluorides o f transition metals and is also able to displace oxygen quantitatively from a number o f metallic oxides. In most cases, as far as available evidence goes, the product is a fluoride rather than a chlorofluoride. One exception, in the case o f chlorine monofluoride.
18
THE CHEMISTRY OF FLUORINE
is the formation o f S F 5 C I from S F 4 and C I F and other similar reactions o f the monofluoride might well be found. Bromine trifluoride has been particularly useful in pre paring a number o f unusual fluoro complexes. Examples o f such reactions are given below. Excess o f bromine trifluoride is used in each case. 2 NOCÍ + SnF4 — 2NOBrF4 -f BrF2SnF6 ^ N 2 O 4 + Sb203 — N 0 2 B r F 4 + BrF2SbF6 ^ Ag
LiF KCl
-I- A u
+ VF5 ^
(NO)2SnF6 + 2BrF3 N02SbF6 + 2BrF3
AgBrF4 + BrF2AuF4
~ * A g A u F 4 + 2BrF3
LiBrF4 + BrF2VFe
— LiVFe + 2BrF3
+ Ru — KBrF4 -f BrF2RuF6
K R u F ö + 2BrF3
The equations are set out to show the way in which the bromine trifluoride first fluorinates the reactants to give inter mediates that are either acids or bases in the bromine tri fluoride solvent system. A neutralization reaction then leads to the products shown, the final step being the removal in vacuum o f excess o f bromine trifluoride and other volatile products. Some o f the intermediates (e.g., NOBrF4, BrF2 V F ö ) have not been isolated but there seems to be ample evidence for postulating their formation. A second point that merits some mention is the relationship between the ionic species derived from the halogen fluorides and the so-called polyhalide ions.^'"*^ It seems best to regard the ions containing fluorine as part o f the general pattern. Thus ICl2^ and ICU" are matched by C1F2^ and C I F 4 - . The two main diff*erences are that ions such as I 5 " have no counterpart and there appear to be few, if any, cations or anions with fluorine and one other halogen both bonded to a heavier halogen [e.g., ( B r F 2 C l 2 ) ~ ] . This may, however, be a question o f finding a suitable preparative method. Neutral interhalogen compounds containing fluorine with two other halogens [^.g., BrCUF(5_;,)] have also not been made. Here again it may be a question of finding the right preparative method. There is a considerable similarity between the halogen fluorides and compounds with both fluorine and oxygen bonded to a heavier halogen.^'^^ This falls into line with what
THE HALOGEN FLUORIDES
19
has been said earlier about the ability o f fluorine and oxygen jointly to excite a high valency state o f a third element. The known halogen oxyfluorides are listed below with their com mon names and their formulas, and serve to illustrate this point. CIO2 F CIO3F
Br02F IO2F IOF3
IO3F IOF5
chloryl fluoride perchloryl fluoride bromyl fluoride iodine dioxyfluoride iodine oxytrifluoride iodine trioxyfluoride iodine oxypentafluoride
b.p. - 6" b.p.-46.8° m.p.-35° — — — m.p. - 10° t o - 2 0 °
The simplest chlorine oxyfluoride, C l O F , has not been made, though there are indications that it might result in the con trolled hydrolysis o f chlorine trifluoride. Chloryl fluoride is readily made by the interaction o f bromine trifluoride and potassium Perchlorate. One o f its more interesting reactions is the formation o f adducts with Lewis-acid fluorides (e.g., SbFs): these are almost certainly salts o f the C102^ ion. Bromyl fluoride is obtained from bromine pentafluoride and potassium brómate at - 5 0 ° and is less stable than its chlorine analogue, decomposing at ca. 50° to bromine trifluoride, bromine, and oxygen. Perchloryl fluoride, which may be considered as a derivative o f the Perchlorate ion is a tetrahedral molecule centered on the chlorine atom. It is made quite readily by the interaction o f fluorosulfonic acid and potassium Perchlorate. It is much more stable than the chloryl fluoride and is potentially useful as an oxidizing agent. A s might be expected the four iodine oxyfluorides are all relatively very unstable and, perhaps for this reason, little is known about their chemistry. The first, I O 2 F , is formed by the reaction of fluorine with iodine pentoxide, and the second, I O F 3 , from iodine pentoxide and pentafluoride. Reaction o f fluorine with a solution o f periodic acid in anhydrous hydro gen fluoride gives iodine trioxyfluoride, and the last member of the series, I O F 5 , is produced when iodine heptafluoride reacts with silica. There is little o f a systematic nature in these preparative methods and no doubt other routes, which may
20
THE CHEMISTRY OF FLUORINE
well be more convenient, will be found in time. For the present purposes, however, the main point is to note how oxygen can replace fluorine in maintaining the higher oxida tion states o f the heavier halogens and also how these species are related to the oxyhalide anions {e.g., C I O 2 F to CIO3" and C I O 3 F to C I O 4 - ) .
The formation o f interhalogen compounds has been known for well over a hundred years. Those containing fluorine are, however, relatively new and are still the subject o f consider able research. The following table shows the range o f compounds known at present. Type
AB
AB3
AB5
ClF(-lOr)
αΤ^ΪΪΤ)
BrF(20°)
BrF3(126^)
ci?^ BrF5(4r)
BrCl
IF3
IFsilOr)
ICl
ICI3
AB7
IF7(277%subl.)
IBr
T w o points are immediately apparent from this table. In the first place fluorine is able to excite the higher valencies o f the other halogens and, secondly, the halogen fluorides, the boil ing points o f which are shown in parentheses as far as they are known, are comparatively volatile. All o f these substances are formed by direct union o f the elements, with some variation in reaction conditions to form a particular compound when alternative combinations are possible. Thus chlorine monofluoride results when equimolar proportions o f the two elements are passed through a nickel tube at 250°. A n alternative method is to allow chlorine to react with the trifluoride. With a higher proportion o f fluor ine combination o f the elements will give the trifluoride, while the pentafluoride is obtained only by using a large excess o f fluorine at 350° and a pressure o f the order o f 250 atmos pheres.'^^ The formation o f the bromine fluorides occurs rather similarly, except that the pentafluoride is easier to 11
12
THE CHEMISTRY OF FLUORINE
prepare, in keeping with the more basic character o f the heavier halogen. Iodine monofluoride is unknown, though bands associated with the I F molecule are observed in the emission spectrum o f the pale green flame o f iodine burning in fluorine. The main product is the pentafluoride: the heptafluoride results when the vapour o f the pentafluoride is allowed to react with excess o f fluorine at 250-270°. Iodine trifluoride is reported as formed when fluorine is passed into a solution o f iodine in a Freon solvent at - 7 8 ° . It is unstable and disproportionates to iodine and the pentafluoride at above about - 3 0 ° . N o t surprisingly the interhalogen compounds show a gen eral similarity in their reactions to the halogens themselves. Those containing fluorine are extremely reactive as fluorinating agents and, in this respect, they are rather similar to the element. There are however difl'erences in reactivity between the various members, a tentative order o f decreasing reactivity being CIF3
> BrFs >
IF7
> CIF > BrFj >
IF5
> BrF.
Bond energy data are, however, largely lacking. The structures o f these molecules is a matter o f consider able interest. Chlorine monofluoride has a dipole moment o f 0.88 0.02 Debye, the dissociation energy being 60.2 kcal/ mole. There is believed to be a true σ-bond with the polarity C I — F . The dipole moment o f bromine monofluoride is higher (1.29 Debye), though the dissociation energy is unknown. Both chlorine trifluoride and bromine trifluoride are T-shaped molecules in the gas phase, the dimensions being as shown below 1.698Ä
F^^Cl 1.598Ä I
83^1
1.810Ä
F-^—Br 1.721Ä
F F These structures are derived from the trigonal bipyramid and there are two pairs o f electrons occupying vacant equatorial
THE HALOGEN FLUORIDES
13
sites. The dimensions in each case are based on the micro wave spectrum and the structures are supported by the results of infrared and Raman studies. The '^F nuclear magnetic resonance spectrum o f chlorine trifluoride shows that fluorine exchange occurs at room temperature and it seems likely that the mechanism involves a dimeric form o f the molecule. The structure of the solid is the same as that o f the gas, the lattice being an assembly o f CIF3 molecules. Chlorine pentafluoride, bromine pentafluoride, and iodine pentafluoride all have structures based on a square pyramid with fluorine at the apices and the heavier halogen atom somewhat below the basal plane containing four fluorine atoms (see Fig. 1).
Figure 1.
This arrangement leaves one pair o f electrons to occupy the vacant octahedral site. The structure o f iodine heptafluoride is still somewhat un certain. Relatively early work on the Raman and infrared spectra o f the liquid and vapor, together with electron dif fraction studies, led to general agreement on a pentagonal bipyramidal structure with fluorine at the apices and iodine in a symmetrical central site. Quite recently,'*^ however, the crystal structure o f the solid has been subject to a critical examination. The data are stated to be in the best agreement with a model based on a dodecahedron in which two atoms at the end o f the 4 axis have been allowed to coalesce. This leads
14
THE CHEMISTRY OF FLUORINE
Figure 2.
to an arrangement shown diagrammatically in Fig. 2. A t o m F ' and the four atoms F " are believed all to be at 1.825 Ä from the iodine atom, though the four F ' ' atoms are not in one plane. The bonds to the F ' " atoms are both 1,97 A in length, and are therefore assumed to be weaker. Such a structure relates the heptafluoride very simply to the pentafluoride. The '^F nuclear magnetic resonance spectrum o f the liquid at room temperature gave unexpectedly broad lines, the simplest explanation o f which is that the '^F nuclei are in non-equivalent positions. Unfortunately, however, overlap o f individual lines prevented more detailed analysis. It seems from the above that there is a difference between the structure o f the solid and that o f the liquid and gas. Alternatively one or the other o f these structures must be wrong and the problem must be reexamined. In view o f the structural evidence outlined above it is reasonable to suppose that bonding in the halogen fluorides is covalent. It is surprising, therefore, to ñnd unambiguous evidence for self-ionization in several species. Iodine chlorides and its bromide have been known for many years to be ionized when dissolved in polar solvents and the monochloride and monobromide also conduct in the molten state. T h e actual modes o f ionization are as follows: 2 ICl ^
Γ
2 ICI3 ^
ICI2 +
+ ICljICI4-
THE HALOGEN FLUORIDES
15
These schemes are well supported by the usual solvo-acid and solvo-base studies. In the case o f the halogen fluorides it is impossible to study conduction in polar solvents o f the usual types as the halogen compounds are much t o o reactive. Since both fused silica and platinum are unattacked, conductivities of the pure liquids can be measured in specially constructed cells provided great care is taken to avoid contamination. This approach has led to values o f the specific conductivities that are listed below: Compound CIF3
Specific conductivity (ohm~' cm"') 6.5 X 10-'
BrF3 BrFs IF5
8.0 X 10-^(25'') 9 X 10-^25") 5.4 X 10-^25°)
By analogy with iodine trichloride, the mode o f selfionization o f bromine trifluoride may be postulated to be 2 BrFa ^
BrF2^ + BrF4-
There is clear evidence for this which will be described later and also less complete evidence that chlorine trifluoride ionizes similarly. For the pentafluorides the most likely mode o f ionization is 2 AB5 ^ AB4^ + A B ö " and this hypothesis also receives considerable support. What we must look for in order to establish the reality o f these ionization schemes is a series o f compounds containing the cation or anion characteristic o f the solvent system and capable in solution o f enhancing its conductivity. Chlorine trifluoride is found to react with alkaU metal chlorides, bro mides, or iodides at 100°. A solution is obtained from which, after evaporating excess o f the interhalogen com pound and free halogen, solids o f the type M^C1F4 are ob tained ( M * = K , R b , C s ) . There may be formulated as ionic compounds containing the C I F 4 - anion which is believed to be characteristic of the solvent. Similarly, chlorine trifluoride forms 1:1 adducts with S b F j , B F 3 , and A s F s , and these may
16
THE CHEMISTRY OF FLUORINE
be formulated as containing the C1F2^ cation (e.g., ClFj^SbFö", C1F2-'BF4-). They are in fact ionic in the solid state and also in chlorine trifluoride solution and the infrared spectrum o f the boron trifluoride adduct also shows the presence o f the BF4" anion j ^ ' Bromine trifluoride has been more fully studied in this con nection. Various "bases" containing the BrF4" anion have been prepared and analyzed [e.g., K B r F 4 , Ba(BrF4)2], to gether with a number o f acids [e.g,, B r F j S b F g , (BrF2)2SnF6]. Both acids and bases dissolve without decomposition in bromine trifluoride to give conducting solutions and it is also possible to carry out conductometric titrations o f acids with bases and so to observe a neutralization process such as 2KBrF4 + (BrF2)2SnF6 = K2SnF6 + 4 BrF3 Very recently'*®^ a 1:1 complex o f N O F and C I F has been isolated which has been shown to contain the linear CIF2" anion, presumably belonging to the system 2 C I F ^ C r + CIF2", though there is as yet no further evidence on this point. Both chlorine trifluoride and bromine trifluoride are also misciblewith anhydrous hydrogen fluoride.'^Ionization in such solutions is very limited but it is reasonable to suppose that it occurs according to the equations CIF3 + HF ^
C1F2^ + H F 2 -
BrF3 + HF ^
BrF2^ + H F 2 -
These ionization schemes are supported by the occurrence o f neutralization reactions which may be followed by conducto metric titrations, e.g. (BrF2^ + H F 2 - ) -f (H2F-' + SbFö") — BrFzSbFe + 3HF (BrF2^ + H F 2 - ) -f ( K ^ + BrF4-) — KHF2 + 2 BrF3 The study of halogen pentafluorides in this connection has so far been confined to iodine pentafluoride, which reacts with potassium fluoride to give K I F e and with antimony penta fluoride to give ISbFio. Both o f these solids dissolve in iodine pentafluoride and enhance its conductivity and a conducto-
THE HALOGEN FLUORIDES
17
metric titration may also be carried out which corresponds with the equation IF6- + I F / SbFe" — KSbF^ + 2 IF5 It may well be that chlorine and bromine pentafluorides will show a similar behavior. There is some further evidence in support o f the ionization postulated for iodine pentafluoride. Boron trifluoride increases its conductivity and, when passed into a solution of potassium fluoride in the pentafluoride, K B F 4 results. It seems that the acid IF4"*^BF4- must be an intermediate here, even though it has not so far been isolated. Adducts o f iodine pentafluoride with several other species ( N 2 O 5 , M 0 O 3 , W O 3 , K I O 4 ) have also been isolated but it is not yet clear whether they are in any way related to the solvent system under discussion. Iodine heptafluoride, which has an unusually narrow liquid range, appears not to form compounds with potassium, rubidium or cesium fluorides. There have, however, been rereports o f the addition compounds I F v - A s F j and IFt SSbFs.''^^ It is tempting to formulate the first o f these as IFe^ AsFó", i.e. as an acid in the system 2 IF7 ^ IF^^ + IFg" and some support for this view comes from the reaction with potassium fluoride, which may be written as IFö^ AsFö" 4- K F — KAsFö + IF7 Several other points relating to the halogen fluorides re main to be mentioned. A l l are strong fluorinating agents for both organic and inorganic compounds.^'^^ This property has, on the whole, been much less used than might have been ex pected in view o f the relative ease with which the compounds may be prepared and stored. A substance such as chlorine trifluoride, for example, will fill much the same role as fluorine itself in forming the higher fluorides o f transition metals and is also able to displace oxygen quantitatively from a number o f metallic oxides. In most cases, as far as available evidence goes, the product is a fluoride rather than a chlorofluoride. One exception, in the case o f chlorine monofluoride.
18
THE CHEMISTRY OF FLUORINE
is the formation o f S F 5 C I from S F 4 and C I F and other similar reactions o f the monofluoride might well be found. Bromine trifluoride has been particularly useful in pre paring a number o f unusual fluoro complexes. Examples o f such reactions are given below. Excess o f bromine trifluoride is used in each case. 2 NOCÍ + SnF4 — 2NOBrF4 -f BrF2SnF6 ^ N 2 O 4 + Sb203 — N 0 2 B r F 4 + BrF2SbF6 ^ Ag
LiF KCl
-I- A u
+ VF5 ^
(NO)2SnF6 + 2BrF3 N02SbF6 + 2BrF3
AgBrF4 + BrF2AuF4
~ * A g A u F 4 + 2BrF3
LiBrF4 + BrF2VFe
— LiVFe + 2BrF3
+ Ru — KBrF4 -f BrF2RuF6
K R u F ö + 2BrF3
The equations are set out to show the way in which the bromine trifluoride first fluorinates the reactants to give inter mediates that are either acids or bases in the bromine tri fluoride solvent system. A neutralization reaction then leads to the products shown, the final step being the removal in vacuum o f excess o f bromine trifluoride and other volatile products. Some o f the intermediates (e.g., NOBrF4, BrF2 V F ö ) have not been isolated but there seems to be ample evidence for postulating their formation. A second point that merits some mention is the relationship between the ionic species derived from the halogen fluorides and the so-called polyhalide ions.^'"*^ It seems best to regard the ions containing fluorine as part o f the general pattern. Thus ICl2^ and ICU" are matched by C1F2^ and C I F 4 - . The two main diff*erences are that ions such as I 5 " have no counterpart and there appear to be few, if any, cations or anions with fluorine and one other halogen both bonded to a heavier halogen [e.g., ( B r F 2 C l 2 ) ~ ] . This may, however, be a question o f finding a suitable preparative method. Neutral interhalogen compounds containing fluorine with two other halogens [^.g., BrCUF(5_;,)] have also not been made. Here again it may be a question of finding the right preparative method. There is a considerable similarity between the halogen fluorides and compounds with both fluorine and oxygen bonded to a heavier halogen.^'^^ This falls into line with what
THE HALOGEN FLUORIDES
19
has been said earlier about the ability o f fluorine and oxygen jointly to excite a high valency state o f a third element. The known halogen oxyfluorides are listed below with their com mon names and their formulas, and serve to illustrate this point. CIO2 F CIO3F
Br02F IO2F IOF3
IO3F IOF5
chloryl fluoride perchloryl fluoride bromyl fluoride iodine dioxyfluoride iodine oxytrifluoride iodine trioxyfluoride iodine oxypentafluoride
b.p. - 6" b.p.-46.8° m.p.-35° — — — m.p. - 10° t o - 2 0 °
The simplest chlorine oxyfluoride, C l O F , has not been made, though there are indications that it might result in the con trolled hydrolysis o f chlorine trifluoride. Chloryl fluoride is readily made by the interaction o f bromine trifluoride and potassium Perchlorate. One o f its more interesting reactions is the formation o f adducts with Lewis-acid fluorides (e.g., SbFs): these are almost certainly salts o f the C102^ ion. Bromyl fluoride is obtained from bromine pentafluoride and potassium brómate at - 5 0 ° and is less stable than its chlorine analogue, decomposing at ca. 50° to bromine trifluoride, bromine, and oxygen. Perchloryl fluoride, which may be considered as a derivative o f the Perchlorate ion is a tetrahedral molecule centered on the chlorine atom. It is made quite readily by the interaction o f fluorosulfonic acid and potassium Perchlorate. It is much more stable than the chloryl fluoride and is potentially useful as an oxidizing agent. A s might be expected the four iodine oxyfluorides are all relatively very unstable and, perhaps for this reason, little is known about their chemistry. The first, I O 2 F , is formed by the reaction of fluorine with iodine pentoxide, and the second, I O F 3 , from iodine pentoxide and pentafluoride. Reaction o f fluorine with a solution o f periodic acid in anhydrous hydro gen fluoride gives iodine trioxyfluoride, and the last member of the series, I O F 5 , is produced when iodine heptafluoride reacts with silica. There is little o f a systematic nature in these preparative methods and no doubt other routes, which may
20
THE CHEMISTRY OF FLUORINE
well be more convenient, will be found in time. For the present purposes, however, the main point is to note how oxygen can replace fluorine in maintaining the higher oxida tion states o f the heavier halogens and also how these species are related to the oxyhalide anions {e.g., C I O 2 F to CIO3" and C I O 3 F to C I O 4 - ) .