Polyhalogen ions

Last updated November 02, 2023

Polyhalogen ions are a group of polyatomic cations and anions containing halogens only. The ions can be classified into two classes, isopolyhalogen ions which contain one type of halogen only, and heteropolyhalogen ions with more than one type of halogen.

Introduction

Numerous polyhalogen ions have been found, with their salts isolated in the solid state and structurally characterized. The following tables summarize the known species.^[1]^[2]^[3]^[4]^[5]^[6]

Isopolyhalogen cations
Diatomic species	* [Cl₂]⁺, [Br₂]⁺, [I₂]⁺
Triatomic species	[Cl₃]⁺, [Br₃]⁺, [I₃]⁺
Tetraatomic species	[Cl₄]⁺, [I₄]²⁺
Pentaatomic species	[Br₅]⁺, [I₅]⁺
Heptaatomic species	† [I₇]⁺
Higher species	[I₁₅]³⁺

*[Cl₂]⁺ can only exist as [Cl₂O₂]²⁺ at low temperatures, a charge-transfer complex from O₂ to [Cl₂]⁺.^[2] Free [Cl₂]⁺ is only known from its electronic band spectrum obtained in a low-pressure discharge tube.^[3]

† The existence of [I₇]⁺ is possible but still uncertain.^[1]

Heteropolyhalogen cations
Triatomic species	[ClF₂]⁺, [Cl₂F]⁺, [BrF₂]⁺, [IF₂]⁺, [ICl₂]⁺, [IBrCl]⁺, [IBr₂]⁺, [I₂Cl]⁺, [I₂Br]⁺
Pentaatomic species	[ClF₄]⁺, [BrF₄]⁺, [IF₄]⁺, [I₃Cl₂]⁺
Heptaatomic species	[ClF₆]⁺, [BrF₆]⁺, [IF₆]⁺

Isopolyhalogen anions
Triatomic species	[Cl₃]⁻, [Br₃]⁻, [I₃]⁻
Tetraatomic species	[Br₄]²⁻, [I₄]²⁻
Pentaatomic species	[I₅]⁻
Heptaatomic species	[I₇]⁻
Octaatomic species	[Br₈]²⁻, [I₈]²⁻
Higher species	[I₉]⁻, [I₁₀]²⁻, [I₁₀]⁴⁻, [I₁₁]⁻, [I₁₂]²⁻, [I₁₃]³⁻, [I₁₆]²⁻, [I₂₂]⁴⁻, [I₂₆]³⁻, [I₂₆]⁴⁻, [I₂₈]⁴⁻, [I₂₉]³⁻

Heteropolyhalogen anions
Triatomic species	[ClF₂]⁻, [BrF₂]⁻, [BrCl₂]⁻, [IF₂]⁻, [ICl₂]⁻, [IBrF]⁻, [IBrCl]⁻, [IBr₂]⁻, [I₂Cl]⁻, [I₂Br]⁻, [AtBrCl]⁻, [AtBr₂]⁻, [AtICl]⁻, [AtIBr]⁻, [AtI₂]⁻
Pentaatomic species	[ClF₄]⁻, [BrF₄]⁻, [IF₄]⁻, [ICl₃F]⁻, [ICl₄]⁻, [IBrCl₃]⁻, [I₂Cl₃]⁻, [I₂BrCl₂]⁻, [I₂Br₂Cl]⁻, [I₂Br₃]⁻, [I₄Cl]⁻, [I₄Br]⁻
Hexaatomic species	[IF₅]²⁻
Heptaatomic species	[ClF₆]⁻, [BrF₆]⁻, [IF₆]⁻, [I₃Br₄]⁻
Nonaatomic species	[IF₈]⁻

Structure

Most of the structures of the ions have been determined by IR spectroscopy, Raman spectroscopy and X-ray crystallography. The polyhalogen ions always have the heaviest and least electronegative halogen present in the ion as the central atom, making the ion asymmetric in some cases. For example, [Cl₂F]⁺ has a structure of [Cl−Cl−F]⁻ but not [Cl−F−Cl]⁻.

In general, the structures of most heteropolyhalogen ions and lower isopolyhalogen ions were in agreement with the VSEPR model. However, there were exceptional cases. For example, when the central atom is heavy and has seven lone pairs, such as [BrF₆]⁻ and [IF₆]⁻, they have a regular octahedral arrangement of fluoride ligands instead of a distorted one due to the presence of a stereochemically inert lone pair. More deviations from the ideal VSEPR model were found in the solid state structures due to strong cation-anion interactions, which also complicates interpretation of vibrational spectroscopic data. In all known structures of the polyhalogen anion salts, the anions make very close contact, via halogen bridges, with the counter-cations.^[4] For example, in the solid state, [IF₆]⁻ is not regularly octahedral, as solid state structure of [(CH₃)₄N]⁺[IF₆]⁻ reveals loosely bound [I₂F₁₁]²⁻ dimers. Significant cation-anion interactions were also found in [BrF₂]⁺[SbF₆]⁻, [ClF₂]⁺[SbF₆]⁻, [BrF₄]⁺[Sb₆F₁₁]⁻.^[2]

General structures of selected heteropolyhalogen ions
Linear (or almost linear)	[ClF₂]⁻, [BrF₂]⁻, [BrCl₂]⁻, [IF₂]⁻, [ICl₂]⁻, [IBr₂]⁻, [I₂Cl]⁻, [I₂Br]⁻
Bent	[ClF₂]⁺, [Cl₂F]⁺, [BrF₂]⁺, [IF₂]⁺, [ICl₂]⁺, [I₂Cl]⁺, [IBr₂]⁺, [I₂Br]⁺, [IBrCl]⁺
Square planar	[ClF₄]⁻, [BrF₄]⁻, [IF₄]⁻, [ICl₄]⁻
Disphenoidal (or seesaw)	[ClF₄]⁺, [BrF₄]⁺, [IF₄]⁺
Pentagonal planar	‡ [IF₅]²⁻
Octahedral	[ClF₆]⁺, [BrF₆]⁺, [IF₆]⁺, ¶ [ClF₆]⁻, [BrF₆]⁻, [IF₆]⁻
Square antiprismatic	[IF₈]⁻

‡[IF₅]²⁻ is one of the two XY_n-type species known to have the rare pentagonal planar geometry, the other being [XeF₅]⁻.

¶[ClF₆]⁻ is distorted octahedral as the stereochemical inert-pair effect is not significant in the chlorine atom.

The [I₃Cl₂]⁺ and [I₃Br₂]⁺ ions have a trans-Z-type structure, analogous to that of [I₅]⁺.

Higher polyiodides

The polyiodide ions have much more complicated structures. Discrete polyiodides usually have a linear sequence of iodine atoms and iodide ions, and are described in terms of association between I₂, I⁻ and [I₃]⁻ units, which reflects the origin of the polyiodide. In the solid states, the polyiodides can interact with each other to form chains, rings, or even complicated two-dimensional and three-dimensional networks.

Bonding

The bonding in polyhalogen ions mostly invoke the predominant use of p-orbitals. Significant d-orbital participation in the bonding is improbable as much promotional energy will be required, while scant s-orbital participation is expected in iodine-containing species due to the inert-pair effect, suggested by data from Mössbauer spectroscopy. However, no bonding model has been capable of reproducing such wide range of bond lengths and angles observed so far.^[3]

As expected from the fact that an electron is removed from the antibonding orbital when X₂ is ionized to [X₂]⁺, the bond order as well as the bond strength in [X₂]⁺ gets higher, consequently the interatomic distances in the molecular ion is less than those in X₂.

Linear or nearly-linear triatomic polyhalides have weaker and longer bonds compared with that in the corresponding diatomic interhalogen or halogen, consistent with the additional repulsion between atoms as the halide ion is added to the neutral molecule. Another model involving the use of resonance theory exists, for example, [ICl₂]⁻ can be viewed as the resonance hybrid of the following canonical forms:

Evidence supporting this theory comes from the bond lengths (255 pm in [ICl₂]⁻ and 232 pm in ICl(g)) and bond stretching wavenumbers (267 and 222 cm⁻¹ for symmetric and asymmetric stretching in [ICl₂]⁻ compared with 384 cm⁻¹ in ICl), which suggests a bond order of about 0.5 for each I–Cl bonds in [ICl₂]⁻, consistent with the interpretation using the resonance theory. Other triatomic species [XY₂]⁻ can be similarly interpreted.^[2]

Even though they have a reduced bond order, all three halogen atoms are tightly bound. The fluorine–fluorine bond of trifluoride, with bond order 0.5, has a bond-strength is 30 kcal/mol, only 8 kcal/mol less than the fluorine–fluorine bond in difluorine whose bond order is 1.^[7]

Synthesis

The formation of polyhalogen ions can be viewed as the self-dissociation of their parent interhalogens or halogens:

2 XY_n ⇌ [XY_n−1]⁺ + [XY_n+1]⁻
3 X₂ ⇌ [X₃]⁺ + [X₃]⁻
4 X₂ ⇌ [X₅]⁺ + [X₃]⁻
5 X₂ ⇌ 2 [X₂]⁺ + 2 [X₃]⁻

Polyhalogen cations

There are two general strategies for preparing polyhalogen cations:

By reacting the appropriate interhalogen with a Lewis acid (such as the halides of B, Al, P, As, Sb) either in an inert or oxidizing solvent (such as anhydrous HF) or without one, to give a heteropolyhalogen cation.

XY_n + MY_m → [XY_n−1]⁺ + [MY_m+1]⁻

By an oxidative process, in which the halogen or interhalogen is reacted with an oxidizer and a Lewis acid to give the cation:

Cl₂ + ClF + AsF₅ → [Cl₃]⁺[AsF₆]⁻

In some cases the Lewis acid (the fluoride acceptor) itself acts as an oxidant:

3 I₂ + 3 SbF₅ → 2 [I₃]⁺[SbF₆]⁻ + SbF₃

Usually the first method is employed for preparing heteropolyhalogen cations, and the second one is applicable to both. The oxidative process is useful in the preparation of the cations [IBr₂]⁺, [ClF₆]⁺, [BrF₆]⁺, as their parent interhalogens, IBr₃, ClF₇, BrF₇ respectively, has never been isolated:

Br₂ + IOSO₂F → [IBr₂]⁺[SO₃F]⁻

2 ClF₅ + 2 PtF₆ → [ClF₆]⁺[PtF₆]⁻ + [ClF₄]⁺[PtF₆]⁻

BrF₅ + [KrF]⁺[AsF₆]⁻ → [BrF₆]⁺[AsF₆]⁻ + Kr

The preparation of some individual species are briefly summarized in the table below with equations:^[1]^[2]^[3]^[4]

Synthesis of some polyhalogen cations
Species	Relevant chemical equation	Additional conditions required
[Cl₂]⁺ (as [Cl₂O₂]⁺)	Cl₂ + [O₂]⁺[SbF₆]⁻ → [Cl₂O₂]⁺[SbF₆]⁻	in anhydrous HF at low temperatures
[Br₂]⁺	Br₂ (in BrSO₃F) + 3 SbF₅ → [Br₂]⁺[Sb₃F₁₆]⁻ (not balanced)	at room temperature
[I₂]⁺	2 I₂ + S₂O₆F₂ → 2 [I₂]⁺[SO₃F]⁻	in HSO₃F
[Cl₃]⁺	Cl₂ + ClF + AsF₅ → [Cl₃]⁺[AsF₆]⁻	at a temperature of 195 K (-78 °C)
[Br₃]⁺	3 Br₂ + 2 [O₂]⁺[AsF₆]⁻ → 2 [Br₃]⁺[AsF₆]⁻ + 2 O₂
[I₃]⁺	3 I₂ + S₂O₆F₂ → 2 [I₃]⁺[SO₃F]⁻
[Cl₄]⁺	2 Cl₂ + IrF₆ → [Cl₄]⁺[IrF₆]⁻	in anhydrous HF, at a temperature below 193 K (-80 °C)
[I₄]²⁺	2 I₂ + 3 AsF₅ → [I₄]²⁺[AsF₆]−2 + AsF₃	in liquid SO₂
[Br₅]⁺	8 Br₂ + 3 [XeF]⁺[AsF₆]⁻ → 3 [Br₅]⁺[AsF₆]⁻ + 3 Xe + BrF₃
[I₅]⁺	2 I₂ + ICl + AlCl₃ → [I₅]⁺[AlCl₄]⁻
[I₇]⁺	7 I₂ + S₂O₆F₂ → 2 I₇SO₃F
[ClF₂]⁺	ClF₃ + AsF₅ → [ClF₂]⁺[AsF₆]⁻
[Cl₂F]⁺	2 ClF + AsF₅ → [Cl₂F]⁺[AsF₆]⁻	at a temperature below 197 K
[BrF₂]⁺	5 BrF₃ + 2 Au → 3 BrF + 2 [BrF₂]⁺[AuF₄]⁻	with excess BrF₃ required
[IF₂]⁺	IF₃ + AsF₅ → [IF₂]⁺[AsF₆]⁻
[ICl₂]⁺	ICl₃ + SbCl₅ → [ICl₂]⁺[SbCl₆]⁻
[IBr₂]⁺	Br₂ + IOSO₂F → [IBr₂]⁺[SO₃F]⁻
[ClF₄]⁺	ClF₅ + SbF₅ → [ClF₄]⁺[SbF₆]⁻
[BrF₄]⁺	BrF₅ + AsF₅ → [BrF₄]⁺[AsF₆]⁻
[IF₄]⁺	IF₅ + 2 SbF₅ → [IF₄]⁺[Sb₂F₁₁]⁻
[ClF₆]⁺	‡ Cs₂[NiF₆] + 5 AsF₅ + ClF₅ → [ClF₆]⁺[AsF₆]⁻ + Ni[AsF₆]₂ + 2 Cs[AsF₆]
[BrF₆]⁺	[KrF]⁺[AsF₆]⁻ + BrF₅ → [BrF₆]⁺[AsF₆]⁻ + Kr
[IF₆]⁺	IF₇ + BrF₃ → [IF₆]⁺[BrF₄]⁻^{[ dubious – discuss ]}

‡ In this reaction, the active oxidizing species is [NiF₃]⁺, which is formed in situ in the Cs₂[NiF₆]/AsF₅/HF system. It is an even more powerful oxidizing and fluorinating agent than PtF₆.

Polyhalogen anions

For polyhalogen anions, there are two general preparation strategies as well:

By reacting an interhalogen or halogen with a Lewis base, most likely a fluoride:
[(CH₃CH₂)₄N]⁺Y⁻ + XY_n → [(CH₃CH₂)₄N]⁺[XY_n+1]⁻
X₂ + X⁻ → [X₃]⁻
By oxidation of simple halides:
KI + Cl₂ → K⁺[ICl₂]⁻

The preparation of some individual species are briefly summarized in the table below with equations:^[1]^[2]^[3]^[4]

Synthesis of some polyhalogen anions
Species	Relevant chemical equation	Additional conditions required
[Cl₃]⁻, [Br₃]⁻, [I₃]⁻	X₂ + X⁻ → [X₃]⁻ (X = Cl, Br, I)
[Br₃]⁻	Br₂ + [(CH₃CH₂CH₂CH₂)₄N]⁺Br⁻ → [(CH₃CH₂CH₂CH₂)₄N]⁺[Br₃]⁻	in 1,2-dichloroethane or liquid sulfur dioxide. [Br₃]⁻ does not exist in solution and is only formed when the salt crystallizes out.
[Br₅]⁻	2 Br₂ + [(CH₃CH₂CH₂CH₂)₄N]⁺Br⁻ → [(CH₃CH₂CH₂CH₂)₄N]⁺[Br₅]⁻	in 1,2-dichloroethane or liquid sulfur dioxide, with excess Br₂
[ClF₂]⁻	ClF + CsF → Cs⁺[ClF₂]⁻
[BrCl₂]⁻^[8]^: v1p294	Br₂ + Cl₂ + 2 CsCl → 2 Cs⁺[BrCl₂]⁻
[ICl₂]⁻^[8]^: v1p295	KI + Cl₂ → K⁺[ICl₂]⁻
[IBr₂]⁻^[8]^: v1p297	CsI + Br₂ → Cs⁺[IBr₂]⁻
[AtBr₂]⁻, [AtICl]⁻, [AtIBr]⁻, [AtI₂]⁻	AtY + X⁻ → [AtXY]⁻ (X = I, Br, Cl; Y = I, Br)
[ClF₄]⁻	NOF + ClF₃ → [NO]⁺[ClF₄]⁻
[BrF₄]⁻	6 KCl + 8 BrF₃ → 6 K⁺[BrF₄]⁻ + 3 Cl₂ + Br₂	excess BrF₃ needed
[IF₄]⁻	2 XeF₂ + [(CH₃)₄N]⁺I⁻ → [(CH₃)₄N]⁺[IF₄]⁻ + 2 Xe	the reactants were mixed at 242 K, then warmed to 298 K for the reaction to proceed
[ICl₄]⁻^[8]^: v1p298	KI + ICl₃ → K⁺[ICl₄]⁻
[IF₅]²⁻	IF₃ + 2 [(CH₃)₄N]⁺F⁻ → [(CH₃)₄N⁺]₂[IF₅]²⁻
[IF₆]⁻	IF₅ + CsF → Cs⁺[IF₆]⁻
[I₃Br₄]⁻	Ph ₄P]⁺Br⁻ + 3 IBr → [Ph₄P]⁺[I₃Br₄]⁻
[IF₈]⁻	IF₇ + [(CH₃)₄N]⁺F⁻ → [(CH₃)₄N]⁺[IF₈]⁻	in acetonitrile

The higher polyiodides were formed upon crystallization of solutions containing various concentrations of I⁻ and I₂. For instance, the monohydrate of K⁺[I₃]⁻ crystallizes when a saturated solution containing appropriate amounts of I₂ and KI is cooled.^[8]^: v1p294

Properties

Stability

In general, a large counter cation or anion (such as Cs ⁺ and [SbF₆]⁻) can help stabilize the polyhalogen ions formed in the solid state from lattice energy considerations, as the packing efficiency increases.

The polyhalogen cations are strong oxidizing agents, as indicated by the fact that they can only be prepared in oxidative liquids as a solvent, such as oleum. The most oxidizing and therefore most unstable ones are the species [X₂]⁺ and [XF₆]⁺ (X = Cl, Br), followed by [X₃]⁺ and [IF₆]⁺.

The stability of the [X₂]⁺ salts (X = Br, I) are thermodynamically quite stable. However, their stability in solution depends on the superacid solvent. For example, [I₂]⁺ is stable in fluoroantimonic acid (HF with 0.2 N SbF₅, H₀ = −20.65), but disproportionates to [I₃]⁺, [I₅]⁺ and I₂ when weaker fluoride acceptors, like NbF₅, TaF₅ or NaF, are added instead of SbF₅.^[4]

14 [I₂]⁺ + 5 F⁻ → 9 [I₃]⁺ + IF₅

For polyhalogen anions with the same cation, the more stable ones are those with a heavier halogen at the center, symmetric ions are also more stable than asymmetric ones. therefore the stability of the anions decrease in the order:

[I₃]⁻ > [IBr₂]⁻ > [ICl₂]⁻ > [I₂Br]⁻ > [Br₃]⁻ > [BrCl₂]⁻ > [Br₂Cl]⁻

Heteropolyhalogen ions with a coordination number larger than or equal to four can only exist with fluoride ligands.

Color

Most polyhalogen ions are intensely colored, with deepened color as the atomic weight of the constituent element increases. The well-known starch-iodine complex has a deep blue color due to the linear [I₅]⁻ ions present in the amylose helix.^[4] Some colors of the common species were listed below:^[3]

fluorocations tend to be colorless or pale yellow, other heteropolyhalogen ions are orange, red or deep purple^[4]
compounds of [ICl₂]⁺ are wine red to bright orange; while that of [I₂Cl]⁺ are dark brown to purplish black
[Cl₃]⁺ is yellow
[Cl₄]⁺ is blue^[2]
[Br₂]⁺ is cherry red
[Br₃]⁺ is brown
[Br₅]⁺ is dark brown
[I₂]⁺ is bright blue
[I₃]⁺ is dark brown to black
[I₄]²⁺ is red to brown
[I₅]⁺ is green or black, the salt [I₅]⁺[AlCl₄]⁻ exists as greenish-black needles, but appears brown-red in thin sections
[I₇]⁺ is black, if its existence in the compound [I₇]⁺[SO₃F]⁻ has been firmly established
[I₁₅]³⁺ is black^[5]
[ICl₂]⁻ is scarlet red
[ICl₄]⁻ is golden-yellow
polyiodides have very dark colors, either dark brown or dark blue

Chemical properties

The heteropolyhalogen cations are explosively reactive oxidants, and the cations often have higher reactivity than their parent interhalogens and decompose by reductive pathways. As expected from the highest oxidation state of +7 in [ClF₆]⁺, [BrF₆]⁺ and [IF₆]⁺, these species are extremely strong oxidizing agents, demonstrated by the reactions shown below:

2 O₂ + 2 [BrF₆]⁺[AsF₆]⁻ → 2 [O₂]⁺ [AsF₆]⁻ + 2 BrF₅ + F₂

Rn + [IF₆]⁺[SbF₆]⁻ → [RnF]⁺[SbF₆]⁻ + IF₅

Polyhalogen cations with lower oxidation states tend to disproportionate. For example, [Cl₂F]⁺ is unstable in solution and disproportionate completely in HF/SbF₅ mixture even at 197 K:

2 [Cl₂F]⁺ → [ClF₂]⁺ + [Cl₃]⁺

[I₂]⁺ reversibly dimerizes at 193 K, and is observed as the blue color of paramagnetic [I₂]⁺ dramatically shifts to the red-brown color of diamagnetic [I₂]⁺, together with a drop in paramagnetic susceptibility and electrical conductivity when the solution is cooled to below 193 K:^[2]

2 [I₂]⁺ ⇌ [I₄]²⁺

The dimerization can be attributed to the overlapping of the half-filled π* orbitals in two [I₂]⁺.

[Cl₄]⁺ in [Cl₄]⁺[IrF₆]⁻ is structurally analogous to [I₄]²⁺, but decomposes at 195 K to give Cl₂, and salts of [Cl₃]⁺ instead of [Cl₂]⁺.^[2]

Attempts to prepare ClF₇ and BrF₇ by fluorinating [ClF₆]⁺ and [BrF₆]⁺ using NOF have met with failure, because the following reactions occurred:^[3]

[ClF₆]⁺[PtF₆]⁻ + NOF → [NO]⁺[PtF₆]⁻ + ClF₅ + F₂

[BrF₆]⁺[AsF₆]⁻ + 2 NOF → [NO]⁺[AsF₆]⁻ + [NO]⁺[BrF₆]⁻ + F₂

The anions are less reactive compared to the cations, and are generally weaker oxidants than their parent interhalogens. They are less reactive towards organic compounds, and some salts are of quite high thermal stability. Salts containing polyhalogen anions of the type M⁺[X_mY_nZ_p]⁻, where m + n + p = {3, 5, 7, 9...}, tend to dissociate into simple monohalide salts between M⁺ and the most electronegative halogen, so that the monohalide has the highest lattice energy. An interhalogen is usually formed as the other product. The salt [(CH₃)₄N]⁺[ClF₄]⁻ decomposes at about 100 °C, and salts of [ClF₆]⁻ are thermally unstable and can explode even at −31 °C.^[4]

Related Research Articles

The halogens are a group in the periodic table consisting of six chemically related elements: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), astatine (At), and tennessine (Ts), though some authors would exclude tennessine as its chemistry is unknown and is theoretically expected to be more like that of gallium. In the modern IUPAC nomenclature, this group is known as group (XVII) or group (VII).

In chemistry, noble gas compounds are chemical compounds that include an element from the noble gases, group 18 of the periodic table. Although the noble gases are generally unreactive elements, many such compounds have been observed, particularly involving the element xenon.

In chemistry, an interhalogen compound is a molecule which contains two or more different halogen atoms and no atoms of elements from any other group.

<span class="mw-page-title-main">Chlorine pentafluoride</span> Chemical compound

Chlorine pentafluoride is an interhalogen compound with formula ClF₅. This colourless gas is a strong oxidant that was once a candidate oxidizer for rockets. The molecule adopts a square pyramidal structure with C_4v symmetry, as confirmed by its high-resolution ¹⁹F NMR spectrum. It was first synthesized in 1963.

<span class="mw-page-title-main">Cobalt(III) fluoride</span> Chemical compound

Cobalt(III) fluoride is the inorganic compound with the formula CoF₃. Hydrates are also known. The anhydrous compound is a hygroscopic brown solid. It is used to synthesize organofluorine compounds.

<span class="mw-page-title-main">Silver(II) fluoride</span> Chemical compound

Silver(II) fluoride is a chemical compound with the formula AgF₂. It is a rare example of a silver(II) compound. Silver usually exists in its +1 oxidation state. It is used as a fluorinating agent.

An inorganic nonaqueous solvent is a solvent other than water, that is not an organic compound. These solvents are used in chemical research and industry for reactions that cannot occur in aqueous solutions or require a special environment. Inorganic nonaqueous solvents can be classified into two groups, protic solvents and aprotic solvents. Early studies on inorganic nonaqueous solvents evaluated ammonia, hydrogen fluoride, sulfuric acid, as well as more specialized solvents, hydrazine, and selenium oxychloride.

Antimony pentafluoride is the inorganic compound with the formula SbF₅. This colourless, viscous liquid is a strong Lewis acid and a component of the superacid fluoroantimonic acid, formed upon mixing liquid HF with liquid SbF₅ in 1:1 ratio. It is notable for its strong Lewis acidity and the ability to react with almost all known compounds.

<span class="mw-page-title-main">Tetrafluoroborate</span> Anion

Tetrafluoroborate is the anion BF⁻
₄. This tetrahedral species is isoelectronic with tetrafluoroberyllate (BeF²⁻
₄), tetrafluoromethane (CF₄), and tetrafluoroammonium (NF⁺
₄) and is valence isoelectronic with many stable and important species including the perchlorate anion, ClO⁻
₄, which is used in similar ways in the laboratory. It arises by the reaction of fluoride salts with the Lewis acid BF₃, treatment of tetrafluoroboric acid with base, or by treatment of boric acid with hydrofluoric acid.

<span class="mw-page-title-main">Selenium tetrafluoride</span> Chemical compound

Selenium tetrafluoride (SeF₄) is an inorganic compound. It is a colourless liquid that reacts readily with water. It can be used as a fluorinating reagent in organic syntheses (fluorination of alcohols, carboxylic acids or carbonyl compounds) and has advantages over sulfur tetrafluoride in that milder conditions can be employed and it is a liquid rather than a gas.

Bromine compounds are compounds containing the element bromine (Br). These compounds usually form the -1, +1, +3 and +5 oxidation states. Bromine is intermediate in reactivity between chlorine and iodine, and is one of the most reactive elements. Bond energies to bromine tend to be lower than those to chlorine but higher than those to iodine, and bromine is a weaker oxidising agent than chlorine but a stronger one than iodine. This can be seen from the standard electrode potentials of the X₂/X⁻ couples (F, +2.866 V; Cl, +1.395 V; Br, +1.087 V; I, +0.615 V; At, approximately +0.3 V). Bromination often leads to higher oxidation states than iodination but lower or equal oxidation states to chlorination. Bromine tends to react with compounds including M–M, M–H, or M–C bonds to form M–Br bonds.

Iodine can form compounds using multiple oxidation states. Iodine is quite reactive, but it is much less reactive than the other halogens. For example, while chlorine gas will halogenate carbon monoxide, nitric oxide, and sulfur dioxide, iodine will not do so. Furthermore, iodination of metals tends to result in lower oxidation states than chlorination or bromination; for example, rhenium metal reacts with chlorine to form rhenium hexachloride, but with bromine it forms only rhenium pentabromide and iodine can achieve only rhenium tetraiodide. By the same token, however, since iodine has the lowest ionisation energy among the halogens and is the most easily oxidised of them, it has a more significant cationic chemistry and its higher oxidation states are rather more stable than those of bromine and chlorine, for example in iodine heptafluoride.

The thallium halides include monohalides, where thallium has oxidation state +1, trihalides in which thallium generally has oxidation state +3, and some intermediate halides containing thallium with mixed +1 and +3 oxidation states. These materials find use in specialized optical settings, such as focusing elements in research spectrophotometers. Compared to the more common zinc selenide-based optics, materials such as thallium bromoiodide enable transmission at longer wavelengths. In the infrared, this allows for measurements as low as 350 cm⁻¹ (28 μm), whereas zinc selenide is opaque by 21.5 μm, and ZnSe optics are generally only usable to 650 cm⁻¹ (15 μm).

There are three sets of Indium halides, the trihalides, the monohalides, and several intermediate halides. In the monohalides the oxidation state of indium is +1 and their proper names are indium(I) fluoride, indium(I) chloride, indium(I) bromide and indium(I) iodide.

The polyiodides are a class of polyhalogen anions composed entirely of iodine atoms. The most common and simplest member is the triiodide ion, I⁻
₃. Other known larger polyiodides include [I₄]²⁻, [I₅]⁻, [I₆]²⁻, [I₇]⁻, [I₈]²⁻, [I₉]⁻, [I₁₀]²⁻, [I₁₀]⁴⁻, [I₁₁]³⁻, [I₁₂]²⁻, [I₁₃]³⁻, [I₁₄]^4-, [I₁₆]²⁻, [I₂₂]⁴⁻, [I₂₆]³⁻, [I₂₆]⁴⁻, [I₂₈]⁴⁻ and [I₂₉]³⁻. All these can be considered as formed from the interaction of the I^–, I₂, and I⁻
₃ building blocks.

The tetrafluoroammonium cation is a positively charged polyatomic ion with chemical formula NF⁺
₄. It is equivalent to the ammonium ion where the hydrogen atoms surrounding the central nitrogen atom have been replaced by fluorine. Tetrafluoroammonium ion is isoelectronic with tetrafluoromethane CF^₄, trifluoramine oxide ONF^₃ and the tetrafluoroborate BF⁻
₄ anion.

<span class="mw-page-title-main">Bismuth pentafluoride</span> Chemical compound

Bismuth pentafluoride is an inorganic compound with the formula BiF₅. It is a white solid that is highly reactive. The compound is of interest to researchers but not of particular value.

<span class="mw-page-title-main">Astatine compounds</span>

Astatine compounds are compounds that contain the element astatine (At). As this element is very radioactive, few compounds have been studied. Less reactive than iodine, astatine is the least reactive of the halogens. Its compounds have been synthesized in nano-scale amounts and studied as intensively as possible before their radioactive disintegration. The reactions involved have been typically tested with dilute solutions of astatine mixed with larger amounts of iodine. Acting as a carrier, the iodine ensures there is sufficient material for laboratory techniques to work. Like iodine, astatine has been shown to adopt odd-numbered oxidation states ranging from −1 to +7.

<span class="mw-page-title-main">Radon compounds</span>

Radon compounds are compounds formed by the element radon (Rn). Radon is a member of the zero-valence elements that are called noble gases, and is chemically not very reactive. The 3.8-day half-life of radon-222 makes it useful in physical sciences as a natural tracer. Because radon is a gas at standard conditions, unlike its decay-chain parents, it can readily be extracted from them for research.

References

1 2 3 4 King, R. Bruce (2005). "Chlorine, Bromine, Iodine, & Astatine: Inorganic Chemistry". Encyclopedia of Inorganic Chemistry (2nd ed.). Wiley. p. 747. ISBN 9780470862100.
1 2 3 4 5 6 7 8 9 Housecroft, Catherine E.; Sharpe, Alan G. (2008). "Chapter 17: The group 17 elements". Inorganic Chemistry (3rd ed.). Pearson. p. 547. ISBN 978-0-13-175553-6.
1 2 3 4 5 6 7 Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 835. ISBN 978-0-08-037941-8.
1 2 3 4 5 6 7 8 Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred (1999). Advanced Inorganic Chemistry (6th ed.). Wiley. ISBN 978-0471199571.
1 2 Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic Chemistry. Academic Press. pp. 419–420. ISBN 0-12-352651-5.
↑ Sonnenberg, Karsten; Mann, Lisa; Redeker, Frenio A.; Schmidt, Benjamin; Riedel, Sebastian (2020-02-04). "Polyhalogen and Polyinterhalogen Anions from Fluorine to Iodine". Angewandte Chemie International Edition. 59 (14): 5464–5493. doi:10.1002/anie.201903197. ISSN 1433-7851. PMID 31090163. S2CID 155093006.
↑ Braïda, Benoît; Hiberty, Philippe C. (2004). "What Makes the Trifluoride Anion F₃^– So Special? A Breathing-Orbital Valence Bond ab Initio Study" (PDF). J. Am. Chem. Soc. 126 (45): 14890–14898. doi:10.1021/ja046443a. PMID 15535716. S2CID 23159174.
1 2 3 4 5 Brauer, G., ed. (1963). Handbook of Preparative Inorganic Chemistry (2nd ed.). New York: Academic Press.

This page is based on this Wikipedia article
Text is available under the CC BY-SA 4.0 license; additional terms may apply.
Images, videos and audio are available under their respective licenses.

[Encyclo-1] 1 2 3 4 King, R. Bruce (2005). "Chlorine, Bromine, Iodine, & Astatine: Inorganic Chemistry". Encyclopedia of Inorganic Chemistry (2nd ed.). Wiley. p. 747. ISBN 9780470862100.

[InorgChem-2] 1 2 3 4 5 6 7 8 9 Housecroft, Catherine E.; Sharpe, Alan G. (2008). "Chapter 17: The group 17 elements". Inorganic Chemistry (3rd ed.). Pearson. p. 547. ISBN 978-0-13-175553-6.

[Greenwood-3] 1 2 3 4 5 6 7 Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 835. ISBN 978-0-08-037941-8.

[AdvInorgChem-4] 1 2 3 4 5 6 7 8 Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred (1999). Advanced Inorganic Chemistry (6th ed.). Wiley. ISBN 978-0471199571.

[wiberg2001-5] 1 2 Wiberg, Egon; Wiberg, Nils; Holleman, Arnold Frederick (2001). Inorganic Chemistry. Academic Press. pp. 419–420. ISBN 0-12-352651-5.

[6] Sonnenberg, Karsten; Mann, Lisa; Redeker, Frenio A.; Schmidt, Benjamin; Riedel, Sebastian (2020-02-04). "Polyhalogen and Polyinterhalogen Anions from Fluorine to Iodine". Angewandte Chemie International Edition. 59 (14): 5464–5493. doi:10.1002/anie.201903197. ISSN 1433-7851. PMID 31090163. S2CID 155093006.

[7] Braïda, Benoît; Hiberty, Philippe C. (2004). "What Makes the Trifluoride Anion F₃^– So Special? A Breathing-Orbital Valence Bond ab Initio Study" (PDF). J. Am. Chem. Soc. 126 (45): 14890–14898. doi:10.1021/ja046443a. PMID 15535716. S2CID 23159174.

[Brauer-8] 1 2 3 4 5 Brauer, G., ed. (1963). Handbook of Preparative Inorganic Chemistry (2nd ed.). New York: Academic Press.

[1]

[2]

[3]

[4]

[5]

[6]

[7]

[8]