Professional scientific experiment of the oxidation of nitrogen (II) oxide
What are the products of oxidation of nitric oxide?
Nitrogen (II) oxide is a non-salt-forming oxide, a poisonous gas without color or smell which dissolves poorly in water. Nitrogen (II) oxide is capable of partially dimerizing (a dimer is a compound consisting of two simpler identical molecules). Boiling point is -152 ᵒC or -242 ᵒF.
Obtaining nitrogen (II) oxide
NO is the only nitrogen oxide which can be obtained from simple substances:
N₂ + O₂ = 2NO (an electric charge or heating to 1200-1300 ᵒC or 2192-2372 ᵒF is required).
In the laboratory and in industry, other reactions are used. In the laboratory, diluted nitric acid is used:
3Cu + 8HNO₃ = 2NO + 3Cu(NO₃)₂ + 4H₂O.
In industry, catalyst oxidation of ammonium is used with a platinum-rhodium catalyst with heating:
4NH₃ + 5O₂ = 4NO + 6H₂O (necessary temperature around 700 ᵒC or 1292 ᵒF).
Nitrogen (II) oxide does not react with water and practically does not dissolve in it. By its nature it is non-salt-forming – salts cannot be obtained from this oxide by reactions. In compounds, the oxidation state of nitrogen is +2, and no nitrogen-containing acids correspond to this oxide (acid cannot be obtained from this oxide by diluting). Depending on the type of reaction, nitrogen (II) oxide may be an oxidizer or reducer. The most important chemical properties of the substance are the following.
On a rhodium catalyst, nitrogen (II) oxide may oxidize carbon monoxide to carbon dioxide:
2NO + 2CO = 2CO₂ + N₂:
It can oxidize sulfur dioxide to trioxide:
2NO + 2SO₂ = 2SO₃ + N₂.
Alkaline causes a disproportionation reaction to begin – “self-oxidizing-self-reducing” – some of the nitrogen atoms oxidizes, and others reduce:
6NO + 4KOH = N₂ + 4KNO₂ + 2H₂O (in a flux).
The compound may enter into combination reactions with halogens:
2NO + Cl₂ = 2NOCl (this reaction takes place in at a cold temperature).
Combination between nitrogen (II) and (III) oxides may also begin:
NO + NO₂ = N₂O₃.
Oxidation of nitrogen (II) oxide
In oxidation reactions, nitrogen (II) oxide is a reducer – for its oxidation by oxygen to begin, no special conditions are required.
2NO + O₂ = 2NO₂ (the gas turns brown, as the oxide with nitrogen at an oxidation state of is a brown gas).
Here you can find experiments with nitrogen dioxide.
In an acidic medium, oxidation of potassium permanganate is also possible:
5NO + 3KMnO₄ + 2H₂SO₄ = 2MnSO₄ + 3KNO₃ + Mn(NO₃)₂ + 2H₂O.
Gas only displays oxidation properties in reactions with better reducers that surpass its reduction ability. The first oxidation reaction of nitrogen (II) oxide by oxygen is necessary for the manufacture of nitric acid, as this type cannot be obtained directly from its oxide. Nitric acid can be obtained by nitrogen IV oxide by the following reaction (there must be a surplus of oxygen):
4NO₂ + О₂ + 2H₂O = 4HNO₃.
Often the nitrogen oxide required for this reaction is obtained from NO, as it is quite easy to oxidize this gas to the required NO₂ (special conditions are not required for this).
Often nitrogen (II) oxide NO is used for the manufacture of nitric acid. The gas has great biological and biochemical significance, as it takes part in the life activity process. The substance is also a reagent for conducting certain chemical reactions – for example obtaining nitric acid.