Oxygen

Oxygen (chemical symbol O) is the most abundant (46.6%) chemical element of the lithosphere (the external layer of earth, which constitutes plates). It is a constituent of water (H₂O), of rocks and sand in the form of silicium oxide (SiO₂) and part of the atmosphere (21%) in the form of the oxygen molecule (O₂). The O-atom has nuclear charge 8 e, where e is the elementary charge, and 8 electrons, which means that it occupies the eighth position in the Periodic system of elements. Atomic oxygen is very reactive, so that the element does not occur naturally in free form, but only in compounds.

History

We now know that the burning of wood and other fuels is a process of oxidation. It is a chemical reaction of fuel with oxygen from the ambient air, yielding mainly heat, carbondioxide (CO₂), and water (H₂O). Although this chemical reaction is the most important discovery in the history of mankind, for many centuries the combustion process was not at all understood. The first step in understanding came in the early 18^th century when Georg Ernest Stahl conjectured that in the combustion process a substance called phlogiston escapes from the flames. This idea was adopted by many chemists, among whom the independent discoverers of oxygen Carl Wilhelm Scheele (in 1771) and Joseph Priestley (in 1774). They discovered what they called dephlogisticated air. Soon after (1777) Antoine Laurent Lavoisier communicated to the French Academy that dephlogisticated air is 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).

Chemical properties

The oxygen atom in its quantum mechanical ground state has the electron 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 1σ²_g 1σ²_u 2σ²_g 2σ²_u 1π⁴_u 3σ²_g 1π²_g. The highest occupied molecular orbitals, 1π_g,x and 1π_g,y, are directed perpendicular to the bond; they are 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 is 117.93 kcal/mol = 493.42 kJ/mol with a vibrationally averaged bondlength r₀ of 0.1207 nm.

The oxygen atom in compounds is very electronegative, second only to fluorine. Oxygen is a strong oxidizer, so that many elements appear as an oxide. The best-known oxides are: dihydrogen oxide (water), carbon dioxide ("carbonic acid"), nitrogen oxide and nitrogen dioxide, sulphur dioxide, sulphate (SO₄ in salts and sulphuric acid), and ferric oxide Fe₂O₃ (well-known in the form of rust, which is ferric oxide plus crystal water).

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

The molecule O₃ (ozone) 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 influence of ultraviolet (UV) light.

Molecular oxygen can be prepared by the heating of several oxides, for instance heating of mercury oxide gives: 2HgO → 2Hg + O₂. This was the reaction used by Scheele and Priestley. A well-known manufacturing process is the electrolysis of water that gives pure H₂ and O₂. However, the most efficient process is the fractional distillation of liquefied air by the Linde process, invented by Carl von Linde in 1895. This is the process that is widely used in the industrial manufacturing of oxygen.

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.

Another application is in oxyfuel welding, where pure oxygen is used to burn a fuel, so that temperatures are obtained that are high enough for welding.

Liquid oxygen is used as an oxidator in rocket fuels.

Physical properties

At normal temperature (0 ⁰C) 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: T_c = 154.581 K, p_c = 5.0430 MPa, and critical density (D_c): 13630 mol/m³ = 436.2 kg/m³.

Crystalline oxygen has three phases: α (monoclinic), β (rhombohedral), and γ (cubic). Phase transitions are γ → β: 43.80 K, α → β: 23.89 K.

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.