Oxygen

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Oxygen
15.9994(3)



  O
8
1s22s22p4 16,2,p
[ ? ] Non-Metal:
Properties:
Odorless, tasteless, colorless gas. Very reactive
Compounds:
Oxides, O2, O3
Uses:
Steel manufacturing, syngas, respiration
Hazard:
Increases flammability

Oxygen (chemical symbol O), the third most abundant element in the universe (after hydrogen and helium), is the most abundant element of the lithosphere (about 47% of Earth's crust[1]), much of it oxides of silicon (e.g., silicon dioxide, SiO2: rocks, sand) and other elements. It is a constituent of water (H2O) and accounts for 20.95% of the dry air in the Earth's atmosphere. Oxygen gas may be diatomic (O2) or triatomic (O3, known as ozone). The most common form, by far, is the diatomic gas.

The O-atom has nuclear charge 8e, where e is the elementary charge, and 8 electrons, occupying the eighth position in the Periodic Table of Elements, i.e., its atomic number Z is eight. Atomic oxygen is very reactive, so that the element does not occur naturally in free form, but only in compounds.

Chemical properties

The oxygen atom in its quantum mechanical ground state has the electronic configuration 1s22s22p4. The ground state atom is in a 3P state, i.e., it is an electron-spin triplet (S = 1) and has orbital angular momentum quantum number L = 1. The O2 molecule in its ground state has the electron configuration 1σ2g2u2g2u4u2g2g. The highest occupied molecular orbitals, 1πg,x and 1πg,y, are directed perpendicular to the bond; they are singly occupied. The electrons occupying these two orbitals couple to a triplet spin. Hence the O2 molecule is a spin triplet in its ground state and is therefore paramagnetic. The binding energy of the molecule 16O2 is 493.5 kJ/mol with an equilibrium bond length re of 0.1207433 nm.[2]

The oxygen atom in compounds is very electronegative,[3] second only to fluorine in its ability to attract a pair of electrons bonding it with another atom. Oxygen is a strong oxidizer, so that many elements (metals and non-metals) and even organic compounds appear as oxides, with ionic or covalent bonds. The best-known oxides are: dihydrogen oxide (water), carbon dioxide (in aqueous solution "carbonic acid"), silica (silicon dioxide), nitrogen oxide and nitrogen dioxide, nitrate (NO3- in nitric acid and its salts), sulphur dioxide, sulphate (SO42- in sulphuric acid and its salts), and ferric oxide Fe2O3 (well-known in the form of rust, which is ferric oxide plus crystal water).

Oxygen gas (O2) has an excited singlet state about 94.2 kJ/mol above the triplet ground state, and the transition from the triplet to the singlet is strictly forbidden in the isolated molecule. As a consequence, singlet oxygen in the gas phase is extremely long lived (72 minutes). Interaction with solvents, however, reduces the lifetime to microseconds or even nanoseconds. The chemistry of singlet oxygen is quite different from that of triplet oxygen.

The ozone molecule (O3) is relatively stable. At normal temperature and pressure it decomposes slowly according to: 2O3 → 3O2 + 69 kcal/mol. In the outer atmosphere (stratosphere) ozone is formed under the influence of ultraviolet (UV) light.

Molecular oxygen can be prepared by the heating of several oxides, for instance the heating of mercury oxide gives: 2HgO → 2Hg + O2. This was the reaction used by Scheele and Priestley (see the history section).

The electrolysis of water is a well-known manufacturing process of very pure oxygen. Two electrodes are immersed into a vessel of water and are subjected to a voltage difference of at least 0.83V. At the cathode (the negative electrode) the following reaction occurs:

4H2O + 4e → 2H2 + 4OH.

The OH anions move to the anode (positive electrode) and decompose according to

4OH → O2 + 2H2O + 4e,

so that the net effect is electron transport from cathode to anode and the formation of gaseous hydrogen at the cathode and gaseous oxygen at the anode. This process, which requires electrical power, is energetically not very efficient.

Pure oxygen is most efficiently produced using the Linde process (developed by Carl von Linde between 1895 and 1901), which uses fractional distillation of liquefied air. This is the process widely applied by industry.

Physical properties

GNU Image
Phase diagram for crystalline oxygen. From Ref. [4]

At normal temperature (0 °C) and pressure (1 atm = 101.325 KPa) the oxygen molecule is a tasteless, colorless, and odorless gas of density 1.429 kg/m³. Its boiling point and freezing point at normal pressure are 90.188 K, and 54.361 K, respectively. The density of liquid O2 at boiling point and normal pressure is 1142 kg/m³. Critical quantities are: Tc = 154.581 K, pc = 5.0430 MPa, and critical density (Dc): 13630 mol/m³ = 436.2 kg/m³.

Since 1916 it is known that crystalline oxygen has three low-pressure phases: α (monoclinic), β (rhombohedral), and γ (cubic), which occur for increasing temperatures (see the diagram). Phase transition temperatures[4] (at p = 0) are α → β: 23.880 K, β → γ: 43.801 K, and γ → liquid: 54.4 K. After 1979 three high-pressure phases δ (between 9.6 and 10 GPa), ε (between 10 and 96 GPa), and ζ (above 96 GPa) were discovered. Except for the ζ phase, all the known phases are shown in the diagram.

Isotopes

The standard atomic mass—the average over different isotopes weighted by abundance—of oxygen is 15.9994 u. (The unit u is the unified atomic mass unit). The following are the atomic masses of the stable isotopes with their abundance in brackets:

16O    15.994 914 6221 (99.757%)
17O    16.999 131 50 (0.038%)
18O    17.999 160 4 (0.205%)

Besides these stable isotopes, a number of short-lived (i.e., radioactive) artificial isotopes of oxygen are known. The longest living of the non-stable isotopes is 15O, which has atomic mass 15.0030656 u and has a half-life of 122.24 seconds.

Biological role

See Oxidative stress

Almost all organisms need oxygen for their cellular metabolism from which they obtain their energy. By aerobic respiration, the required oxygen is imported from the air and some waste products, such as carbon dioxide, are released. For mammals, red blood cells containing hemoglobin are an important mediator in the transport of oxygen to the tissues. A human can survive for only a few minutes without oxygen, because irreversible brain damage occurs if there is a lack of oxygen for a longer period.

Large quantities of oxygen are removed daily from the atmosphere by combustion of fossil fuels by industry, automobiles, and home furnaces. The respiration of aerobic organisms also consumes oxygen. The only source of replenishing oxygen is the process of photosynthesis, which occurs in green plants. In fact, it is generally assumed that this process is responsible for the existence of gaseous oxygen in the earth's atmosphere and that higher forms of life only became feasible after sufficient oxygen had been supplied by photosynthesis.

Industrial applications

Most industrial oxygen is used in the basic oxygen steelmaking process. In this method pure oxygen is blown into a bath of molten blast-furnace iron and scrap. The oxygen initiates the oxidation of such impurities as carbon, silicon, phosphorus, and manganese, yielding relatively pure steel. Industrial oxygen is also used to make a mixture of carbon monoxide and hydrogen called synthesis gas, used for the synthesis of methanol and ammonia.

Another application is in oxyfuel welding, where pure oxygen is used to burn a fuel (often acetylene C2H2), so that temperatures sufficiently high for welding are obtained.

Liquid oxygen is used as an oxidator in rocket fuels.

History

Today we know that the burning of wood and other carbonaceous fuels in air occurs through the process of oxidation—the stripping of electrons from the fuel source molecules and their transfer to oxygen. The chemical reaction of fuel with oxygen from the ambient air—combustion—yields mainly heat, carbon dioxide (CO2), and water (H2O). Although this chemical reaction is perhaps the most important discovery in the history of mankind, for many centuries the combustion process was not at all understood. The alchemists of the seventeenth century recognized that air contained an essential ingredient, an 'elixir of life'. The first step in understanding came in the early eighteenth century when Georg Ernest Stahl conjectured that in the combustion process a substance called phlogiston escapes from the fuel in the flames and into the air. This idea was adopted by many chemists, among whom were the independent discoverers of oxygen Carl Wilhelm Scheele (in 1771) and Joseph Priestley (in 1774). They discovered what Priestley called dephlogisticated air, i.e., air in which the postulated phlogiston was absent. Soon after (1777) Antoine Laurent Lavoisier communicated to the French Academy that dephlogisticated air is actually a separate gas, a constituent of several acids and hence must be a chemical element. Because Lavoisier believed (erroneously) that the presence of oxygen in an acid was essential, he proposed to replace the name dephlogisticated air by oxygen meaning generator of acidity. (The old-Greek word for wine vinegar being oxys—όξύς, from the Greek word for sharp, also όξύς). Lavoisier understood that oxygen was an important part of air, and this fact was proved beyond doubt when oxygen was separated from liquid air by distillation (1883).

Reference

  1. Abundances of the Elements in the Earth's Crust Physics and Astronomy Dept., Georgia State University
  2. W. Steinbach and W. Gordy Physical Review vol. A 11, p. 729 (1975)
  3. Electronegativity
  4. 4.0 4.1 Yu.A. Freiman and H.J. Jodl, Solid oxygen, Physics Reports, vol. 401, pp. 1-228 (2004).