18.9 Occurrence, preparation, and compounds of oxygen  (Page 3/27)

The oxides of halogens, at least one of the noble gases, and metals with higher reduction potentials than copper do not form by the direct action of the elements with oxygen.

Reaction with compounds

Elemental oxygen also reacts with some compounds. If it is possible to oxidize any of the elements in a given compound, further oxidation by oxygen can occur. For example, hydrogen sulfide, H ₂ S, contains sulfur with an oxidation state of 2−. Because the sulfur does not exhibit its maximum oxidation state, we would expect H ₂ S to react with oxygen. It does, yielding water and sulfur dioxide. The reaction is:

It is also possible to oxidize oxides such as CO and P ₄ O ₆ that contain an element with a lower oxidation state. The ease with which elemental oxygen picks up electrons is mirrored by the difficulty of removing electrons from oxygen in most oxides. Of the elements, only the very reactive fluorine can oxidize oxides to form oxygen gas.

Oxides, peroxides, and hydroxides

Compounds of the representative metals with oxygen fall into three categories: (1) oxides , containing oxide ions, O ²⁻ ; (2) peroxides , containing peroxides ions, ${\text{O}}_2{}^{\mathrm{2-}},$ with oxygen-oxygen covalent single bonds and a very limited number of superoxides , containing superoxide ions, ${\text{O}}_2{}^{\text{−}},$ with oxygen-oxygen covalent bonds that have a bond order of $1\frac{1}{2},$ In addition, there are (3) hydroxides , containing hydroxide ions, OH ⁻ . All representative metals form oxides. Some of the metals of group 2 also form peroxides, MO ₂ , and the metals of group 1 also form peroxides, M ₂ O ₂ , and superoxides, MO ₂ .

Oxides

It is possible to produce the oxides of most representative metals by heating the corresponding hydroxides (forming the oxide and gaseous water) or carbonates (forming the oxide and gaseous CO ₂ ). Equations for example reactions are:

However, alkali metal salts generally are very stable and do not decompose easily when heated. Alkali metal oxides result from the oxidation-reduction reactions created by heating nitrates or hydroxides with the metals. Equations for sample reactions are:

With the exception of mercury(II) oxide, it is possible to produce the oxides of the metals of groups 2–15 by burning the corresponding metal in air. The heaviest member of each group, the member for which the inert pair effect is most pronounced, forms an oxide in which the oxidation state of the metal ion is two less than the group oxidation state (inert pair effect). Thus, Tl ₂ O, PbO, and Bi ₂ O ₃ form when burning thallium, lead, and bismuth, respectively. The oxides of the lighter members of each group exhibit the group oxidation state. For example, SnO ₂ forms from burning tin. Mercury(II) oxide, HgO, forms slowly when mercury is warmed below 500 °C; it decomposes at higher temperatures.

Burning the members of groups 1 and 2 in air is not a suitable way to form the oxides of these elements. These metals are reactive enough to combine with nitrogen in the air, so they form mixtures of oxides and ionic nitrides. Several also form peroxides or superoxides when heated in air.