Oxidation states of oxygen and its characteristics
Some properties of oxygen and its compounds
Oxygen is an element of the 6ᵗʰ group (according to the new classification an element of the 16th group) of the main subgroup of the periodic table. It is a member of the group of chalcogens (besides oxygen, they include sulfur, selenium, tellurium and polonium). In translation, “chalcogen” means “giving birth to ore”
Physical properties of oxygen, its presence in nature
Oxygen is a colorless diatomic gas in normal conditions. It has no taste or smell, but its presence can be identified with a smoldering stick: when it is placed in an atmosphere of oxygen it flares up, as oxygen supports combustion well. In an atmosphere of oxygen, breathing and decomposition are also possible.
In nature it is in encountered in the form of three isotopes (elements with an identical number on the periodic table, but with different atomic numbers). Oxygen can be encountered most frequently with the atomic numbers 16, 17 or 18.
In nature, the maximum quantity of oxygen is contained in the air (around 21% of its volume), in water and in the earth’s crust (up to 47% oxygen by mass).
Molecular oxygen has the allotropic modification O₃ (ozone). Ozone is a poisonous gas of a bluish color with a specific smell. If it is compressed, it acquires a rich blue color. In a solid state, the substance forms dark blue, almost black crystals.
It forms from the breakdown of molecular oxygen into atomic oxygen (this takes place with an electrical charge, harsh ultraviolet irradiation of the air or the breakdown of peroxides). The O₃ molecule is rather unstable, so for a relatively short time (several tens of minutes), it breaks down into molecular oxygen. To prevent this from happening, it must be cooled to -78 ᵒC or -108.4 ᵒF in a hermetic glass or metal container, or have a little nitric acid added to it. Ozone is a very reactive substance, and surpasses molecular oxygen by its reactivity and oxidation properties.
Click here for experiments with pure oxyden.
There are numerous methods for obtaining oxygen:
distillation of liquid air (its rectification);
electrolysis of water according to the following equation:
2H₂O = O₂ + 2H₂ (presence of an alkali is necessary);
- electrolysis of salts of acids containing oxygen:
2CuSO₄ + 2H₂O₂ = 2Cu + 2O₂ + 2H₂SO₄;
breakdown of oxidizers:
2KClO₃ = 3O₂ + 2KCl (with heating);
2KMnO₄ = O₂ + MnO₂ + K₂MnO₄ (with heating);
- potassium nitrate:
2KNO₃ = O₂ + 2KNO₂ (with heating);
- hydrogen peroxide:
2H₂O₂ = O₂ + 2H₂O (in the presence of manganese (IV) oxide MnO₂ or with heating);
- reaction of potassium peroxide with carbon dioxide: 2Na₂O₂ + 2CO₂ = O₂ + 2Na₂CO₃.
In nature, oxygen is obtained by photosynthesis from carbon dioxide and water in light:
mCO₂ + nH₂O = mO₂ + Cₘ(H₂O)ₙ (besides oxygen, carbohydrates also form when illuminated with light).
Oxidation states of oxygen
Being a strong oxidizer, oxygen most frequently displays an oxidation state of -₂ in compounds (other chalcogens have a lower oxidation state). Typical examples of compounds where oxygen has an oxidation state of -2 are H₂SO₄, H₂O, KNO₃ and CaO (in the hydroxonium ion H₃O⁺ the oxidation state of oxygen is also -2). In compounds of this type (apart from the hydroxonium ion), the valence of oxygen (its ability to form a certain number of chemical bonds) is equal to two. For H₃O⁺, the valence of oxygen is equal to three, as oxygen forms three bonds – two covalent and one donor-acceptor.
In peroxides, the oxidation state of oxygen is -1. The formation of peroxides is characteristic for hydrogen (hydrogen peroxide H₂O₂) and some metals (for example, sodium peroxide Na₂O₂, barium peroxide BaO₂, calcium peroxide CaO₂ etc.).
In a free state, oxygen has the oxidation state of 0, like other simple substances.
Only in one case can oxygen have a positive oxidation state of +2 – in a compound with fluorine, with the formula OF₂. As fluorine is a more electrically negative element than oxygen, it acquires the negative oxidation state (-1) in the compound. The compound is called “oxygen fluoride II”. It is obtained in the combustion of water in an atmosphere of fluorine, and also by the reaction:
2F₂ + 2NaOH = OF₂ + 2NaF + H₂O (ozone and hydrogen peroxide can also form in the reaction).
There is a compound with a positive oxidation state of oxygen (+1), O₂F₂ (oxygen monofluoride, obtained in the reaction of molecular oxygen and fluorine). It is an unstable compound.
Properties of molecular oxygen
In the majority of cases, oxygen displays oxidizing properties in reactions, both with simple and complex substances. If the oxidation of oxygen takes place violently and rapidly, this reaction is called a combustion reaction, during which oxygen oxidizes any substance with the violent release of energy (usually in the form of heat and light). If the reaction with oxygen takes place gently and slowly (often with a catalyst), then this process is called oxidation. Which process takes place in a specific case often depends on the conditions in which the reaction is carried out and the properties of the oxidizing substance.
With non-metals, oxygen reacts with the formation of acidic oxides (with heating):
S + O₂ = SO₂;
4P + 5O₂ = 2P₂O₅ (phosphorus pentoxide exists in the form of a dimer in the composition of P₄O₁₀); https://commons.wikimedia.org/wiki/Category:Phosphorus_pentoxide#/media/File:Phosphorus-pentoxide-3D-balls.png
- C+O₂ = CO₂.
With metals, oxygen reacts with the formation of basic oxides or peroxides (the reaction takes place with heating):
4Li + O₂ = 2Li₂O;
2Zn + O₂ = 2ZnO;
2Na + O₂ = Na₂O₂.
With many complex substances, oxygen also reacts (in all cases in the second reagent there is an element which is not in the highest oxidation state – it is oxidized by oxygen to a higher oxidation state, or to the highest possible state for it). Oxygen also oxidizes organic and non-organic compounds (the most significant reactions are given, which reflect the oxidizing properties of oxygen):
СН₄ + О₂ = НСОН + Н₂О (incomplete oxidation of methane is possible with the formation of formaldehyde with the use of catalysts);
СН₄ + 2О₂ = СО₂ + 2Н₂О (complete oxidation of methane – combustion);
C₂H₂OH + 3O₂ = 2CO₂ + 3H₂O (complete oxidation of alcohol);
C₂H₂OH + O₂ = CH₃COOH + H₂O (gentle oxidation of alcohol by oxygen to acetic acid);
2BaO + O₂ = 2BaO₂ (barium oxide absorbs oxygen, forming peroxide);
Na₂O₂ + O₂ = 2NaO₂ (when reacting with peroxides, oxygen oxidizes them to superoxides with an oxidation state of oxygen of -1/2);
4NH₃ + 3O₂ = 2N₂ + 6H₂O (combustion);
4NH₃ + 5O₂ = 4NO + 6H₂O (catalytic oxidation with heating in the presence of platinum);
2SO₂ + O₂ = 2SO₃ (reaction takes place with heating and adding a catalyst);
H₂S + 3O₂ = 2SO₂ + 2H₂O (sulfur (IV) oxide forms in an abundance of oxygen);
2H₂S + O₂ = 2S + 2H₂O (molecular sulfur forms in the reaction with a lack of oxygen);
4Fe(OH)₂ + O₂ + 2H₂O = 4Fe(OH)₃;
4FeS₂ + 11O₂ = 8SO₂ + 2Fe₂O₃ (reaction takes place with heating with the formation of two oxides);
4CH₃NH₂ + 9O₂ = ₄CO₂ + 2N₂ + 10H₂O.
Qualitative reaction to ozone:
O₃ + 2KI + H₂O = I₂ + 2KOH + O₂ (molecular oxygen is released and a dark violet sediment of molecular iodine forms).
Oxygen has found wide application in industry (for example in smelting cast-iron and steel). It is also used for cutting metals (with acetylene) and in medicine. Ozone is used for bleaching fabrics and decontaminating water.