OXYGEN

Atomic Number: 8
Mass Number: 15.9994
Electron Configuration: 1s²2s²2p⁴
Boiling Point:-183^oC
Melting Point: -218.4^oC
First Ionization Energy: 1314 kJ/mol
Atomic Radius: 74 pm
Ionic Radius: 140 (O^2-)
Electronegativity: 3.4

Oxygen exists as three isotopes as summarized below. By far the most common isotope is ¹⁶O. The second, ¹⁷O, is NMR active and has a nuclear spin of 5/2. The third isotope, ¹⁸O, is commonly used as a nonradioactive `tracer for the element.

A an element, oxygen occurs as a diatomic element in which the oxygen atoms are connected by a double bond. The bond distance is 121 pm. Oxygen gas is colorless and odorless, but is pale blue in color as a liquid. Oxygen is paramagnetic, a property which can only be explained through molecular orbital theory. Oxygen combines with nearly all other elements on the periodic table. Oxygen will react with most metals to form oxides, with hydrogen to form water, and with nonmetals such as sulfur and fluorine. Oxygen is the second most electronegative element on the periodic table, next to fluorine. In most covalent bonds with oxygen the bonding electrons are sharred unequally, resulting in a partial negative charge on the oxygen atom and a partial positive charge on the bonded atom.

A second form of oxygen is the element ozone, which consists of triatomic molecules The two oxygen-oxygen bonds of the ozone molecule are the same and have a length of 128 pm. The longer bond length represents a lower bond order than in the oxygen diatomic molecule. As might be expected due to the heavier molecular weight, ozone boils at a higher temperature than molecular oxygen. Ozone molecules have a bent shape, a property which could be predicted from the dot structure. Unlike oxygen, ozone has a pungent odor. Ozone is much more chemically reactive than oxygen. Diatomic oxygen and ozone are examples of allotropes, or different forms of the same element.

As one might predict by the position of oxygen on the periodic table, oxgen can gain two additional oxygen atoms to form the oxide ion, and many compounds containing the oxide ion are known. Oxygen also forms the peroxide and superoxide ions. The formulas for these species are listed in the table below. The peroxide can be viewed as an oxygen molcule with two extra electrons. According to molecular orbital theory, these two electrons go into antibonding orbitals, reducing the bond order to one. Metals that can form ionic peroxides include the alkali metals, calcium, barium, and strintium. The superoxide molecule can be viewed as an oxygen molecule with an extra single electron and has a bond order of one and a half. Metals that can form ionic superoxides include potassium, rubidium, and cesium.

Magnesium, calcium, strontium, and barium oxides adopt a rock-salt structure, in which the oxide ions form a face-centered cubic array and the metal ions fill the octahedral holes. Lithium, sodium, and potassium oxides, in contrast, adopt the antifluorite structure. In this type of structure, the oxide ions form a face-centered cubic array and the metal ions fill half the tetrahedral holes. This is considered to be the inverse of the fluorite structure, where the cations form the face-centered cubic array and the anions fit into the tetrahedral holes.

Compounds with Hydrogen

Hydrogen peroxide is a colorless, viscous liquid with a melting point of - 0.41 C and a boiling point of 150.2 C. Dilute solutions of hydrogen peroxide (approximately 3%) are commonly used as an antiseptic. Hydrogen peroxide is unstable and decomposes to give water and oxygen; this serves as a potential laboratory preparation of oxygen. The decomposition is catalyzed by manganese dioxide, as well as several other metal ions. Hydrogen peroxide is a strong oxidizing agent, and will oxidize the iodide ion to iodine. In alkaline solution, hydrogen peroxide can also serve as a mild reducing agent.

Compounds with Sulfur

The two most common molecular compounds between sulfur and oxygen are sulfur dioxode and sulfur trioxide. Sulfur dioxide has a boiling point of -10 C and is a melting point of -75.5 C. It has a pungent, choking odor and is the product formed when sulfur burns. Like ozone, the sulfur dioxide molecule has an angular shape, and the the bond angle is 119 degrees. Liquid sulfur dioxide is a good solvent. Sulfur dioxide dissolves in water to produce sulfurous acid, a weak diprotic acid.

Sulfur trioxide is formed from the oxidation of sulfur dioxide. In this particular molecule, the sulfur is the central atom and supports an expanded octet. Sulfur trioxide has a melting point of 16.9 C and a boiling point of 44.6 C. Sulfur trioxide reacts vigorously with water to produce sulfuric acid.

Compounds with Carbon

Two common compounds are carbon dioxide and carbon monoxide. Carbon dioxide is produced in vast quantities by combustion of carbon-containing molecules. Carbon dioxide is a colorless, odorless gas. In the carbon dioxide molecule, both carbon-oxygen bonds are double bonds and the shape of the molecule is linear. Dry ice is composed of solid carbon dioxide. Dry ice sublimes, or passes from a solid directly to a gas, at -78.5 C.

_{Carbon monoxide is often formed when carbon-containing molecules are burned in a limited supply of oxygen. Carbon monoxide has a boiling point of -190 C and is a gas at room temperature. In the carbon monoxide molecule, the atoms are joined by a triple bond. Carbon monoxide is highly toxic because it is absorbed by hemoglobin many times better than oxygen, tying up the hemoglobin so that it can no longer combine wtih oxygen.}

_{Compounds with Nitrogen}

_{Oxygen forms an entire series of molecular compounds with nitrogen. Some properties of these compounds are summarized below.}

In addition to the molecular compounds listed above, oxygen forms an extensive series of oxoanions with many of the nonmetals. The chemical formulas and names of some nitrogen, oxygen, and chlorine oxoanions are listed in the table below.

_{Oxygen is highly reactive. Metals tend to react with oxygen to form oxides. For example, magnesium burns with a brilliant flame in oxygen to produce magnesium oxide. Sulfur burns with a blue flame in oxygen to produce sulfur dioxide.}