Oxidation states of nitrogen
How many oxidation states does nitrogen have?
Nitrogen is an element in the 15ᵗʰ group (under the new classification) of the second period of the Period Table. It is encountered in nature in the form of two isotopes (atoms with identical atomic numbers, but different mass numbers) – nitrogen with the mass numbers 14 and 15. Nitrogen comprises around 78% of the volume of the air. This gas is one of the most abundant on Earth.
Physical properties and obtaining nitrogen
By its physical properties, nitrogen is a colorless gas without smell or taste. The boiling point of nitrogen is -195.75 ᵒC (-320.35 ᵒF). The gas is chemically inert. In the laboratory, nitrogen is obtained by the decomposition of ammonium nitrate NH₄NO₂:
NH₄NO₂ = N₂ + 2H₂O (the vessel must first be heated, then cooled, as the process is exothermic – it takes place with the release of heat, up to 335 kJ). The nitrogen released contains impurities – nitrogen and ammonium oxides.
Molecular nitrogen can also be obtained with the thermal decomposition of ammonium and nitrogen (I) oxide:
2NH₃ = N₂ + 3H₂;
2N₂O = 2N₂ + O₂.
In industrial conditions, nitrogen is obtained by distilling liquefied air. Nitrogen displays a considerable number of oxidation states, from -3 to +5.
Compounds with an oxidation state of nitrogen of -3
The 2 most common types of compounds in which the oxidation state of nitrogen is -3 is ammonium and ammonium salt (or ammonium hydroxide NH₄OH).
The two most common methods of obtaining ammonium are in the laboratory and industrially:
Industrially – 3H₂ + N₂ = 2NH₃ (in harsh conditions at high pressure and temperature, and in the presence of a catalyst);
In the laboratory – Ca(OH)₂ + 2NH₄Cl = 2NH₃ + 2H₂O + CaCl₂.
The main properties of ammonium are the following:
NH₃ + HCl = NH₄Cl (reaction with acids leads to the formation of ammonium salts);
2NH₃ + AgCl = [Ag(NH₃)₂]Cl (reaction with salts of heavy metals leads to formation of complexes; in this case silver (I) diamine chloride forms);
4NH₃ + 3O₂ = 2N₂ + 6H₂O (combustion reaction);
4NH₃ + 5O₂ = 4NO + 6H₂O (reaction of catalytic oxidation in the presence of platinum with heating);
2NH₃ + CuO = N₂ + 3Cu + 3H₂O (reduction reaction of copper from its oxide, carried out with heating).
Ammonium salts are solid crystalline substances. You can detect the presence of ammonium ions in a salt by adding an alkali to it:
NH₄Cl + KOH = NH₃ + NaCl + H₂O (the ammonium released in the reaction has a specific smell; because it is a base litmus paper turns blue when it is held up to the reaction flask);
Ammonium salts can also decompose under thermal impact. The precise products of reaction depend on the composition of the initial salt:
(NH₄)₂SO₄ = NH₃ + NH₄HSO₄;
NH₄NO₃ = N₂O + 2H₂O;
NH₄Cl = NH₃ + HCl.
Otherwise, the properties of ammonium salts repeat the properties of other salts soluble in water.
Properties of nitrogen with an oxidation state of 0
As a simple substance, nitrogen is an inert diatomic gas. It enters into many reactions only when heated:
- reaction with metals with heating (apart from lithium, which it reacts with at room temperature):
3Ca + N₂ = Ca₃N₂;
6Li + N₂ = 2Li₃N;
- reaction with acetylene:
С₂Н₂ + N₂ = 2HCN (takes place in a condensed discharge);
- reaction with hot coal:
2C + N₂ = (CN)₂ (dicyanogen forms);
- reaction with boron at high temperature:
2B + N₂ = 2BN;
- reaction with oxygen with formation of nitrogen (II) oxide:
N₂ + O₂ = 2NO (takes place with heating);
- reaction with hydrogen with heating (500 degrees ᵒC or 932 ᵒF), high pressure (350 atmospheres) and in the presence of a catalyst (for example Fe):
N₂ + 3H₂ = 2NH₃;
- reaction with sodium carbonate and reaction with coal with heating:
2Na₂CO₃ + 8C + 2N₂ = 4NaCN + 6CO;
- reaction with fluorine in electric discharge:
N₂ + 3F₂ = 2NF₃;
- reaction with calcium carbide at a temperature of 1000 ᵒC (1832 ᵒF):
CaC₂ + N₂ = CaCN₂ + C;
- reaction with hydrogen with heating (from 500 ᵒC or 932 ᵒF and higher depending on the demands for reaction speed and output), high pressure (around 350 atmospheres) and in the presence of a catalyst (for example, Fe or FeO):
N₂ + 3H₂ = 2NH₃.
Nitrogen does not react directly with sulfur or halogens. The gas also does not react with acids, alkalis or water.
Here you can find exciting experiments with nitrogen.
Nitrogen in the oxidation states +1, +2, +4
The nitrogen oxides N₂O and NO are non-salt-forming. Nitrogen oxide, in which nitrogen has an oxidation state of +1, has a sweet smell, and dissolves well in water. It can be obtained by the thermal decomposition of ammonium nitrate:
NH₄NO₃ = N₂O + 2H₂O (it is important to remember that high temperatures can cause decomposition of the oxide:
2N₂O = N₂ + O₂.
Nitrogen (II) oxide
There are several methods for obtaining nitrogen (II) oxide:
- Adding diluted nitric acid to copper (laboratory):
3Cu + 8HNO₃ = 2NO + 3Cu(NO₃)₂ + 4H₂O.
- Catalyst oxidation of ammonia on a platinum-rhodium catalyst at 700 ᵒC or 1292 ᵒF (industry):
4NH₃ + 5O₂ = 4NO + 6H₂O.
- In the atmosphere the gas may form in lightning discharges:
N₂ + O₂ = 2NO.
By its properties it is a reducer. The chemical properties of nitrogen (II) oxide are the following:
2NO + O₂ = 2NO₂;
5NO + 3KMnO₄ + 2H₂SO₄ = 2MnSO₄ + 3KNO₃ + Mn(NO₃)₂ + 2H₂O;
2NO + 2SO₂ = 2SO₃ + N₂;
2NO + 2CO = 2CO₂ + N₂ (rhodium catalyst and heating required);
6NO + 4KOH = N₂ + 4KNO₂ + 2H₂O (takes place in a flux).
Nitrogen (IV) oxide
Nitrogen (IV) oxide is a brown gas which displays properties of an acidic oxide. It has the ability to dimerize.
This compound may be obtained by several methods:
- Cu + 4HNO₃ = Cu(NO₃)₂ + 2NO₂ + 2H₂O (concentrated hot water);
2Pb(NO₃)₂ = 2PbO + 4NO₂ + O₂ (decomposition of dried nitrate of a heavy metal at a high temperature);
2NO + O₂ = 2NO₂ (in industry).
2NO₂ + H₂O = HNO₃ + HNO₂ (disproportion reaction – the nitrogen atom is oxidized and reduced);
4NO₂ + 2H₂O + О₂ = 4HNO₃;
С + 2NO₂ = CO₂ + 2NO (combustion);
S + 2NO₂ = SO₂ + 2NO (combustion).
Nitrogen in the oxidation state of +3
Nitrogen has an oxidation state of +3 in the oxide N₂O₃ (a dark blue liquid with acidic properties) and nitrous acid HNO₂. Nitrogen (III) oxide can be obtained by the following reactions:
- Cooling a mixture of oxides (II) and (IV) at a temperature of -36 ᵒC or -32.8 ᵒF):
NO + NO₂ = N₂O₃.
- Dehydration of nitrous acid (dehydrating agent – phosphorus oxide P₄O₁₀):
2HNO₂ = N₂O₃ + H₂O.
It displays typical acidic properties:
N₂O₃ + H₂O = 2HNO₂;
N₂O₃ + 2NaOH = 2NaNO₂ + H₂O;
N₂O₃ + Na₂O = 2NaNO₂.
Nitrous acid is a weak acid, and concentrated solutions have a bluish tiny. It is obtained by dilution of the according oxide:
2HNO₂ = N₂O₃ + H₂O or in a mixture with nitric acid diluted with nitrogen (IV) oxide:
2NO₂ + H₂O = HNO₃ + HNO₂.
It displays typical acidic properties – it reacts with alkalis and base oxides with formation of salts. It can be an oxidizer and a reducer:
2HNO₂ + 2HI = I₂ + 2NO + 2H₂O (oxidizer);
NO₂ + Cl₂ + H₂O = HNO₃ + HCl (reducer).
It decomposes when heated. Salts of nitrous acid – for example nitrites of alkaline metals are highly soluble colorless (yellowish) crystalline substances.
Nitrogen in an oxidation state of +5
N₂O₅ is a solid unstable substance with acidic properties. It can be obtained by the reactions:
- Passing the according acid through a column with phosphorus (V) oxide:
4HNO₃ + P₄O₁₀ = 2N₂O₅ + 4HPO₃ (cooling to -10 ᵒC or 14 ᵒF is required).
- Ozone reacting with nitrogen (IV) oxide:
2NO₂ + O₃ = N₂O₅ + O₂.
- Decomposition into oxygen and nitrogen (4) oxide:
2N₂O₅ = 4NO₂ + O₂;
- Dissolving in water:
N₂O₅ + H₂O = 2HNO₃;
- Reactions with bases and base oxides:
N₂O₅ + 2NaOH = 2NaNO₃ + H₂O; N₂O₅ + Na₂O = 2NaNO₃.
Nitrous acid is a strong oxidizer (by its physical properties it is a colorless liquid with no smell). It can be obtained in the laboratory as follows:
KNO₃ + H₂SO₄ = KHSO₄ + HNO₃ (concentrated sulfuric acid is taken).
In industry a three-stage process is used:
- Oxidation of ammonia (platinum catalyst and heating to 500 ᵒC or 932 ᵒF required):
4NH₃ + 5O₂ = 4NO + 6H₂O.
- *Oxidation of obtained oxide in air to NO₂:
2NO + O₂ = 2NO₂.
- Dissolution of NO₂ in water in an excess of oxygen:
4NO₂ + О₂ + 2H₂O = 4HNO₃.
Main chemical properties (besides those typical for all acids):
4HNO₃ = 4NO₂ + 2H₂O + O₂ (decomposition of concentrated acid in light);
S + 4HNO₃(conc.) = SO₂ + 4NO₂ + 2H₂O (reaction with non-metals).
Concentrated and diluted nitric acid react with metals without release of hydrogen:
4HNO₃ + Cu = Cu(NO₃)₂ + 2NO₂ + 2H₂O (concentrated acid);
8HNO₃ + 3Cu = 3Cu(NO₃)₂ + 2NO + 4H₂O (diluted acid);
10HNO₃ + 8Na = 8NaNO₃ + N₂O + 5H₂O (concentrated acid);
10HNO₃ + 8Na = 8NaNO₃ + NH₄NO₃ + 3H₂O (diluted acid).
Concentrated nitric acid passivates iron, chromium, aluminum, gold, platinum and iridium. Many nitrates (salts of nitric acid) are soluble in water.
Nitrogen compounds are used in medicine (liquid ammonia, ammonia spirit), agriculture (fertilizers) and in industry for the synthesis of several organic compounds.