Dioxygenyl

Last updated
Dioxygenyl
Dioxygenyl-ion-2D-dimensions.png
Names
IUPAC name
Dioxygenyl
Other names
Dioxyl
Identifiers
ChEBI
ChemSpider
  • InChI=1S/O2/c1-2/q+1
    Key: KMHJKRGRIJONSV-UHFFFAOYSA-N
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

The dioxygenyl(or dioxyl) ion, O+
2, is a rarely-encountered oxycation in which both oxygen atoms have a formal oxidation state of +1/2. It is formally derived from oxygen by the removal of an electron:

Contents

O2 → O+
2 + e−

The energy change for this process is called the ionization energy of the oxygen molecule. Relative to most molecules, this ionization energy is very high at 1175 kJ/mol. [1] As a result, the scope of the chemistry of O+
2 is quite limited, acting mainly as a 1-electron oxidiser. [2]

Structure and molecular properties

O+
2 has a bond order of 2.5, and a bond length of 112.3 pm in solid O2[AsF6]. [3] It is isoelectronic with nitric oxide and is paramagnetic. [4] The bond energy is 625.1 kJ mol−1 and the stretching frequency is 1858 cm−1, [5] both of which are high relative to most of the molecules.

Synthesis

Neil Bartlett demonstrated that dioxygenyl hexafluoroplatinate (O2PtF6), containing the dioxygenyl cation, can be prepared at room temperature by direct reaction of oxygen gas (O2) with platinum hexafluoride (PtF6): [6]

O2 + PtF6 → [O
2]+
[PtF
6]−

The compound can also be prepared from a mixture of fluorine and oxygen gases in the presence of a platinum sponge at 450 °C, and from oxygen difluoride (OF
2) above 400 °C: [6]

6 OF
2 + 2 Pt → 2 [O
2][PtF
6] + O
2

At lower temperatures (around 350 °C), platinum tetrafluoride is produced instead of dioxygenyl hexafluoroplatinate. [6] Dioxygenyl hexafluoroplatinate played a pivotal role in the discovery of noble gas compounds. The observation that PtF6 is a powerful enough oxidising agent to oxidise O2 (which has a first ionization potential of 12.2  eV) led Bartlett to reason that it should also be able to oxidise xenon (first ionization potential 12.13 eV). His subsequent investigation yielded the first compound of a noble gas, xenon hexafluoroplatinate. [7]

O+
2 is also found in similar compounds of the form O2MF6, where M is arsenic (As), antimony (Sb), [8] gold (Au), [9] niobium (Nb), ruthenium (Ru), rhenium (Re), rhodium (Rh), [10] vanadium (V), [11] or phosphorus (P). [12] Other forms are also attested, including O2GeF5 and (O2)2SnF6. [11]

The tetrafluoroborate and hexafluorophosphate salts may be prepared by the reaction of dioxygen difluoride with boron trifluoride or phosphorus pentafluoride at −126 °C: [12]

2 O2F2 + 2 BF3 → 2 O2BF4 + F2
2 O2F2 + 2 PF5 → 2 O2PF6 + F2

These compounds rapidly decompose at room temperature:

2 O2BF4 → 2 O2 + F2 + 2 BF3
2 O2PF6 → 2 O2 + F2 + 2 PF5

Some compounds including O2Sn2F9, O2Sn2F9·0.9HF, O2GeF5·HF, and O2[Hg(HF)]4(SbF6)9 can be made by ultraviolet irradiation of oxygen and fluorine dissolved in anhydrous hydrogen fluoride with a metal oxide. [13]

All attempts to prepare O+
2 with chloro anions like [O
2]+
[SbCl
6]−
 met with failure.

Reactions

The reaction of O2BF4 with xenon at 173 K (−100 °C) produces a white solid believed to be F–Xe–BF2, containing an unusual xenon-boron bond: [14]

2 O2BF4 + 2 Xe → 2 O2 + F2 + 2 FXeBF2

The dioxygenyl salts O2BF4 and O2AsF6 react with carbon monoxide to give oxalyl fluoride, C2O2F2, in high yield. [15]

Related Research Articles

<span class="mw-page-title-main">Hydrofluoric acid</span> Solution of hydrogen fluoride in water

Hydrofluoric acid is a solution of hydrogen fluoride (HF) in water. Solutions of HF are colorless, acidic and highly corrosive. It is used to make most fluorine-containing compounds; examples include the commonly used pharmaceutical antidepressant medication fluoxetine (Prozac) and the material PTFE (Teflon). Elemental fluorine is produced from it. It is commonly used to etch glass and silicon wafers.

<span class="mw-page-title-main">Xenon hexafluoroplatinate</span> Chemical compound

Xenon hexafluoroplatinate is the product of the reaction of platinum hexafluoride with xenon, in an experiment that proved the chemical reactivity of the noble gases. This experiment was performed by Neil Bartlett at the University of British Columbia, who formulated the product as "Xe+[PtF6]−", although subsequent work suggests that Bartlett's product was probably a salt mixture and did not in fact contain this specific salt.

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.

<span class="mw-page-title-main">Oxygen fluoride</span> Any binary compound of oxygen and fluorine

Oxygen fluorides are compounds of elements oxygen and fluorine with the general formula OnF2, where n = 1 to 6. Many different oxygen fluorides are known:

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

Xenon tetrafluoride is a chemical compound with chemical formula XeF
4. It was the first discovered binary compound of a noble gas. It is produced by the chemical reaction of xenon with fluorine:

<span class="mw-page-title-main">Xenon hexafluoride</span> Chemical compound

Xenon hexafluoride is a noble gas compound with the formula XeF6. It is one of the three binary fluorides of xenon that have been studied experimentally, the other two being XeF2 and XeF4. All known are exergonic and stable at normal temperatures. XeF6 is the strongest fluorinating agent of the series. It is a colorless solid that readily sublimes into intensely yellow vapors.

<span class="mw-page-title-main">Platinum hexafluoride</span> Chemical compound

Platinum hexafluoride is the chemical compound with the formula PtF6, and is one of seventeen known binary hexafluorides. It is a dark-red volatile solid that forms a red gas. The compound is a unique example of platinum in the +6 oxidation state. With only four d-electrons, it is paramagnetic with a triplet ground state. PtF6 is a strong fluorinating agent and one of the strongest oxidants, capable of oxidising xenon and O2. PtF6 is octahedral in both the solid state and in the gaseous state. The Pt-F bond lengths are 185 picometers.

<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.

<span class="mw-page-title-main">Xenon difluoride</span> Chemical compound

Xenon difluoride is a powerful fluorinating agent with the chemical formula XeF
2, and one of the most stable xenon compounds. Like most covalent inorganic fluorides it is moisture-sensitive. It decomposes on contact with water vapor, but is otherwise stable in storage. Xenon difluoride is a dense, colourless crystalline solid.

<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.

<span class="mw-page-title-main">Xenon oxytetrafluoride</span> Chemical compound

Xenon oxytetrafluoride is an inorganic chemical compound. It is an unstable colorless liquid with a melting point of −46.2 °C that can be synthesized by partial hydrolysis of XeF
6, or the reaction of XeF
6 with silica or NaNO
3:

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

Gold(V) fluoride is the inorganic compound with the formula Au2F10. This fluoride compound features gold in its highest known oxidation state. This red solid dissolves in hydrogen fluoride but these solutions decompose, liberating fluorine.

<span class="mw-page-title-main">Hypofluorous acid</span> Chemical compound

Hypofluorous acid, chemical formula HOF, is the only known oxyacid of fluorine and the only known oxoacid in which the main atom gains electrons from oxygen to create a negative oxidation state. The oxidation state of the oxygen in hypofluorites is 0. It is also the only hypohalous acid that can be isolated as a solid. HOF is an intermediate in the oxidation of water by fluorine, which produces hydrogen fluoride, oxygen difluoride, hydrogen peroxide, ozone and oxygen. HOF is explosive at room temperature, forming HF and O2:

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

Manganese tetrafluoride, MnF4, is the highest fluoride of manganese. It is a powerful oxidizing agent and is used as a means of purifying elemental fluorine.

Dioxygenyl hexafluoroplatinate is a compound with formula O2PtF6. It is a hexafluoroplatinate of the unusual dioxygenyl cation, O2+, and is the first known compound containing this cation. It can be produced by the reaction of dioxygen with platinum hexafluoride. The fact that PtF
6 is strong enough to oxidise O
2, whose first ionization potential is 12.2 eV, led Neil Bartlett to correctly surmise that it might be able to oxidise xenon (first ionization potential 12.13 eV). This led to the discovery of xenon hexafluoroplatinate, which proved that the noble gases, previously thought to be inert, are able to form chemical compounds.

A hexafluoride is a chemical compound with the general formula QXnF6, QXnF6m−, or QXnF6m+. Many molecules fit this formula. An important hexafluoride is hexafluorosilicic acid (H2SiF6), which is a byproduct of the mining of phosphate rock. In the nuclear industry, uranium hexafluoride (UF6) is an important intermediate in the purification of this element.

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

Neptunium(VI) fluoride (NpF6) is the highest fluoride of neptunium, it is also one of seventeen known binary hexafluorides. It is an orange volatile crystalline solid. It is relatively hard to handle, being very corrosive, volatile and radioactive. Neptunium hexafluoride is stable in dry air but reacts vigorously with water.

Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others the highest oxidation states of oxides and fluorides are always equal.

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

Chlorine oxide trifluoride or chlorine trifluoride oxide is a corrosive liquid molecular compound with formula ClOF3. It was developed secretly as a rocket fuel oxidiser.

<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. ↑ Michael Clugston; Rosalind Flemming (2000). Advanced Chemistry, Oxford University Press, ISBN   0-19-914633-0, ISBN   978-0-19-914633-8, p. 355.
  2. ↑ Foote, Christopher S.; Valentine, Joan S. (1995). Active oxygen in chemistry. Joel F. Liebman, A. Greenberg. Springer. ISBN   0-412-03441-7.
  3. ↑ Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. ISBN   978-0-08-037941-8. p. 616
  4. ↑ Cotton, F. Albert; Wilkinson, Geoffrey; Murillo, Carlos A.; Bochmann, Manfred (1999), Advanced Inorganic Chemistry (6th ed.), New York: Wiley-Interscience, ISBN   0-471-19957-5
  5. ↑ J. Shamir; J. Binenboym; H. H. Claassen (1968). "The vibrational frequency of the O2+ cation". J. Am. Chem. Soc. 90 (22): 6223–6224. doi:10.1021/ja01024a054.
  6. 1 2 3 Bartlett, Neil; Lohmann, D. H. (1962). "Fluorides of the Noble Metals. Part II. Dioxygenyl hexafluoroplatinate(V), [O
    2    ]+
    [PtF
    6    ]−
    ". J. Chem. Soc. 115: 5253–5261. doi:10.1039/jr9620005253.
  7. ↑ Bartlett, Neil (1962). "Xenon hexafluoroplatinate(V), Xe+
    [PtF
    6    ]−
    ". Proc. Chem. Soc. : 197–236. doi:10.1039/PS9620000197.
  8. ↑ Young, A. R.; Hirata, T.; Morrow, S. I. (1964). "The Preparation of Dioxygenyl Salts from Dioxygen Difluoride". J. Am. Chem. Soc. 86 (1): 20–22. doi:10.1021/ja01055a006.
  9. ↑ Nakajima, Tsuyoshi (1995). Fluorine-carbon and fluoride-carbon materials: chemistry, physics, and applications. CRC Press. ISBN   0-8247-9286-6.
  10. ↑ Vasile, Michael J.; Falconer, Warren E. (1975). "Vapour transport of dioxygenyl salts". J. Chem. Soc., Dalton Trans. 1975 (4): 316–318. doi:10.1039/DT9750000316.
  11. 1 2 Holleman, Arnold F.; Wiberg, Egon (2001). Inorganic chemistry. Academic Press. p. 475. ISBN   0-12-352651-5.
  12. 1 2 Solomon, Irvine J.; Brabets, Robert I.; Uenishi, Roy K.; Keith, James N.; McDonough, John M. (1964). "New Dioxygenyl Compounds". Inorg. Chem. 3 (3): 457. doi:10.1021/ic50013a036.
  13. ↑ Mazej, Zoran; Goreshnik, Evgeny (2020-02-03). "Syntheses of Dioxygenyl Salts by Photochemical Reactions in Liquid Anhydrous Hydrogen Fluoride: X-ray Crystal Structures of α- and β-O2Sn2F9, O2Sn2F9·0.9HF, O2GeF5·HF, and O2[Hg(HF)]4(SbF6)9". Inorganic Chemistry. 59 (3): 2092–2103. doi: 10.1021/acs.inorgchem.9b03518 . ISSN   0020-1669. PMC   7307900 . PMID   31942804.
  14. ↑ Goetschel, C. T.; Loos, K. R. (1972). "Reaction of xenon with dioxygenyl tetrafluoroborate. Preparation of FXe-BF2". Journal of the American Chemical Society. 94 (9): 3018–3021. doi:10.1021/ja00764a022.
  15. ↑ Pernice, H.; Willner, H.; Eujen, R. (2001). "The reaction of dioxygenyl salts with 13
    CO Formation of F13
    C(O)13
    C(O)F". Journal of Fluorine Chemistry. 112 (2): 277–590. doi:10.1016/S0022-1139(01)00512-7.
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.