Oxygen

Key points: Oxygen has two allotropes, dioxygen and ozone. Dioxygen has a triplet ground state and oxidizes hydrocarbons by a radical chain mechanism. Reaction with an excited state molecule can produce a fairly long-lived singlet state that can react as an electrophile. Ozone is an un stable and highly aggressive oxidizing agent. Dioxygen is a biogenic gas (that is, one that has been produced by the action of organisms) and most of it is the result of pho to synthesis, although some is produced by the action of ultra violet radiation on water. It is colourless, odourless, and soluble in water to the extent of 3.08 cm3 per 100 cm3 water at 25C and atmospheric pressure. This solubility falls to below 2.0 cm3 in seawater but is still sufficient to support aerobic marine life. The solubility of O₂ in organic solvents is approximately ten times greater than in water. The high solubility of O2 makes it necessary to purge all solvents used in the synthesis of oxygen sensitive compounds. Oxygen is readily available as O2 from the atmosphere and is obtained on a massive scale by the liquefaction and distillation of liquid air. The main commercial motivation is to recover O₂ for steel making, in which it reacts exothermically with coke (carbon) to produce carbon monoxide. The high temperature is necessary to achieve a fast reduction of iron oxides by CO and carbon (Section 5.16). Pure oxygen, rather than air, is advantageous in this process because energy is not wasted in heating the nitrogen. About 1 tonne (1 t=10³ kg) of oxygen is needed to make 1 tonne of steel. Oxygen is also required by industry in the production of the white pigment TiO₂ by the chloride process:

Oxygen is used in many oxidation processes, for example the production of oxirane (ethylene oxide) from ethene. Oxygen is also supplied on a large scale for sewage treatment, renew al of polluted waterways, paper-pulp bleaching, and as an artificial atmosphere in medical and submarine applications. Liquid oxygen is very pale blue and boils at -183C. Its col our arises from electronic transitions involving pairs of neigh bouring molecules: one photon from the red yellow green region of the visible spectrum can raise two O₂ molecules to an excited state to form a molecular pair. Under high pressure the colour of solid oxygen changes from light blue to orange and then to red at approximately 10 GPa.

The molecular orbital description of O₂ implies the exist ence of a double bond; however, as we saw in Section 2.8, the outermost two electrons occupy different antibonding π orbitals with parallel spins; as a result, the molecule is para magnetic (Fig. 16.1). The term symbol for the ground state is ³∑g^-, and henceforth the molecule will be denoted O2 (3∑g^-)

when it is appropriate to specify the spin state.1 The singlet state 1∑g^- , with paired electrons in the same two π* orbitals as in the ground state, is higher in energy by 1.61 eV (155 kJ mol^-1), and another singlet state ¹∆_g (‘singlet delta’), with electrons paired in one π* orbital, lies between these two terms at 0.95 eV (91.7 kJ ^1-) above the ground state. Of the two singlet states, the latter has much the longer excited state lifetime (its return to the ground state is spin-forbidden) and O₂ (¹∆_g) survives long enough to participate in chemical reactions. When it is needed for reactions, O₂ (¹∆_g) can be generated in solution by energy transfer from a photoexcited molecule. Thus [Ru(bpy)₃]²⁺ can be excited by absorption of blue light (452 nm) to give an electronically excited state, denoted *[Ru(bpy)₃] ⁺² (Section 20.7), and this state transfers energy to O₂ (³∑_g^-):

Another efficient way to generate O₂ (¹∆_g) is through the thermal decomposition of an ozonide:

In contrast to the radical character of many O₂(³∑_g^-) reactions, O2 (1∆g ) reacts as an electrophile. This mode of reaction is feasible because O₂(¹∆_g) has an empty π* orbital, rather than two that are each occupied by a single electron. For example, O₂ (¹∆_g)

¹Thesymbols ∑, π, and are ∆ used for linear molecules such as dioxygen in place of the symbols S, P, and D used for atoms. The Greek letters represent the magnitude of the total orbital angular momentum around the internuclear axis.

adds across a diene, thus mimicking the Diels Alder reaction of butadiene with an electrophilic alkene:

Singlet oxygen is implicated as one of the biologically hazardous products of photochemical smog.

The other allotrope of oxygen, ozone, O₃, boils at -112C and is an explosive and highly reactive endoergic blue gas (∆_fG^O 163 kJ mol^-1). It decomposes into dioxygen

but this reaction is slow in the absence of a catalyst or ultraviolet radiation. Ozone has a pungent odour; this property is reflected in its name, which is derived from the Greekozein, to smell. The O3 molecule is angular, in accord with the VSEPR model (13) and has bond angle 117; it is diamagnetic. Gaseous ozone is blue, liquid ozone is blue black, and solid ozone is violet black. Ozone is produced from electrical discharges or ultraviolet radiation acting on O₂. This second method is used to produce low concentrations of ozone for the preservation of foodstuffs. The ability of O₃ to absorb strongly in the 220-290 nm region of the spectrum is vital in preventing the harmful ultraviolet rays of the Sun from reaching the Earth’s surface (Box 17.3). Ozone reacts with unsaturated polymers, causing undesirable cross-linking and degradation.

Reactions of ozone typically involve oxidation and transfer of an O atom. Ozone is very unstable in acidic solution and much more stable in basic conditions:

Ozone is exceeded in oxidizing power only by F₂, atomic O, the OH radical, and perxenate ions (Section 18.5). Ozone forms ozonides with Group 1 and 2 elements (Sections 11.8 and 12.8). They are prepared by passing gaseous ozone over the powdered hydroxide, MOH or M(OH)₂, at temperatures below 10C. The ozonides are red brown solids that decompose on warming:

The ozonide ion, O^-₃, is angular, like O₃, but with the slightly larger bond angle of 119.5.

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مواضيع ذات صلة

Hydrogen peroxide

Water

Selenium, tellurium, and polonium

Sulfur

Reactivity of oxygen

Condensed phosphates

Oxoanions of phosphorus, arsenic, antimony, and bismuth

Oxides of phosphorus, arsenic, antimony, and bismuth

Nitrogen(V) oxides and oxoanions

Oxides and oxoanions of nitrogen

Oxohalides

Halides of the heavy elements