22.5+Oxygen

=__Section 22.5 Oxygen__= == Oxygen is a reactive gas at room temperature. It has two allotropes, O 2 and O 3. It is paramagnetic and light blue in its liquid form. Oxygen is the second most electronegative atom after fluorine.

Oxygen forms strong bonds, so oxygen-containing compounds are typically very thermodynamically stable. Since the bond between the oxygen atoms in O 2 are also strong, reactions involving oxygen often have high activation energies. These reactions are also often highly exothermic and accelerate quickly once they achieve the proper activation energy. One example of this is combustion reactions.

Atomic Structure
Oxygen has six valence electrons, so it forms 2- ions or exists in a compound in -2 oxidation state. It either forms two single bonds or a double bond.

Ozone
Ozone is the less common allotrope of oxygen, O 3. It can be created by applying electricity to O 2. It has a bent molecular structure and exhibits resonance. Like molecules such as benzene, it has a pi bond delocalized over its entire structure. Ozone is unstable and readily decomposes into O 2. It is also a stronger oxidizing agent than O 2 and will oxidize all common metals except for gold and platinum. Though it is an important part of our upper atmosphere by absorbing ultraviolet rays, in the lower atmosphere it is both toxic and a damaging pollutant due to its oxidizing ability.

Oxides
An oxide is a compound containing oxygen in a -2 oxidation state.

Nonmetals form covalent oxides. These will react readily with water to form acids and are called acidic anhydrides (meaning without water) or oxides. One example of this is the reaction of sulfur dioxide with water to form sulfurous acid by the equation SO 2 + H 2 O --> H 2 SO 3. In fact, this particular reaction is responsible for the formation of acid rain when water vapor in the atmosphere encounters sulfur dioxide released as a pollutant.

Metals form ionic oxides. They react with water to form hydroxides, so they are basic anhydrides or oxides. The reaction of calcium oxide with water, CaO + H 2 O --> Ca(OH) 2, is an example. If a metal can form multiple oxides, the basicity of the oxide decreases as the metal's oxidation state increases. Because of their basic properties, even normally insoluble metal oxides can be dissolved in acidic solutions. In this case, the hydroxide ions that would normally be formed combine with the hydrogen ions to form water. For example, Fe 2 O 3 + 6 H + --> 2 Fe 3+ + 3 H 2 O.
 * = Oxide ||= Metal Oxidation State ||= Acid/Base Properties ||
 * = CrO ||= 2+ ||= Basic ||
 * = Cr 2 O 3 ||= 3+ ||= Amphoteric ||
 * = CrO 3 ||= 6+ ||= Acidic ||

Peroxides and Superoxides
Superoxides contain O-O bonds and the oxygen atoms in the compound are in the -1/2 oxidation state, such as in potassium superoxide, KO 2. They are very unstable and react with water to give off oxygen gas: 4 KO 2 + 2 H 2 O --> 4 K + + 4 OH - + 3 O 2. They are only ever formed by the most reactive metals like potassium, cesium, and rubidium.

Peroxides contain O-O bonds the oxygen atoms are in a -1 oxidation state, such as in hydrogen peroxide (H 2 O 2 ). They are unstable and are only formed by metals slightly less reactive than the ones named above, such as sodium, barium, calcium, and strontium. They, like superoxides, decompose to form oxygen gas.

Further Exploration
If you want to know more about oxygen, or any element really, watch this video and the others made by the same authors about the elements.

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