Nitrogen is a colorless, odorless and non-toxic gas, which is very inert. Pure nitrogen gas can be prepared by the home chemist by mixing a nitrite and an ammonium salt in aqueous solution. Liquid nitrogen gas can be fun to play with, but from a chemical point of view, nitrogen gas is not really interesting for the home chemist.
Although the element is not really interesting for the home lab, several compounds of the elements are very interesting and some are a must have for a serious home lab.
Nitrogen exists in compounds in all oxidation states, ranging from -3 to +5. All nitrogen compounds in oxidation states, other than -3, 0 and +5 are reactive compounds, which are really interesting and allow many interesting experiments to be performed in aqueous media. Many nitrogen compounds also are capable of forming interesting complexes with transition metal ions and several experiments, described on this site are based on the properties of these nitrogen compounds.
The number of interesting nitrogen compounds for the home lab is too large to be given exhaustively. An overview of compounds, interesting for the home lab, is given here for each of the possible oxidation states of nitrogen.
Oxidation state +5: Nitric acid and nitrates are very common compounds of nitrogen. Nitric acid is interesting for the home chemist, both as a very reactive and strongly oxidizing acid in concentrated form, but also as an acid, which has an anion, which hardly forms any coordination complexes in dilute aqueous chemistry. The nitrate ion on its own is not really interesting in aqueous chemistry. At low concentration or at pH, larger than 1, the ion is quite inert. Many compounds are available as nitrates, because of the fact that the nitrate-ion is a counter-ion, which does not have a strong influence on the chemistry of the cation, which is the active part of the compound. For dry-chemistry experiments, however, nitrates can be quite interesting. At elevated temperatures, nitrates are powerful oxidizers and many interesting pyrotechnic experiments can be done with nitrates as oxidizers. If one wants to perform interesting pyro-experiments, then potassium nitrate is the best choice, because this is non-hygroscopic and can be obtained as a fertilizer at reasonable price at acceptable purity.
Concentrated nitric acid is a dangerous acid. It is very corrosive and a strong oxidizer, which reacts violently with many organic compounds and many metals, producing toxic nitrogen oxides. Because of its reactivity, it is nice to have some nitric acid at 50 - 70% by weight. Many interesting experiments can be done with this. Nitric acid must be obtained locally. International shipping of this acid, especially in the concentrated form, hardly is possible, due to postal and transport regulations.
Acid of higher concentration is not suitable for the home lab. Pure nitric acid is unstable and slowly decomposes, giving off oxygen and nitrogen dioxide and is exceedingly corrosive. The benefits of having highly concentrated acid at home do not outweigh the extreme risks in storing and handling.
Oxidation state +4: There are no compounds of nitrogen in this oxidation state, which are available commercially to the general public. An interesting but very corrosive and toxic compound, which can be made at home, when one has access to concentrated nitric acid (see above), is nitrogen dioxide. This is a dark brown gas. It exists in equilibrium with dinitrogen tetroxide, which is a colorless compound. Low temperatures and high pressure favor the formation of N2O4 and high temperature and low pressure favor the formation of NO2. The gas NO2 is very reactive and forms some interesting compounds with transition metals. The production of NO2 from a metal and nitric acid already is a nice experiment on its own.
Be warned, the gas NO2 is insidiously toxic. It attacks the lungs, but it has not a strong warning level when inhaled. Although it can be smelled easily, it is not really choking when inhaled in dangerous concentrations. Its adverse effects can be delayed with several hours! This is quite a remarkable difference, compared to e.g. hydrogen chloride, chlorine and sulphur dioxide, which are really choking on their first inhalation of any appreciable quantity. Because of the lack of sufficient and immediate warning, it is dangerous to work with this gas. As a rule of thumb, when one can see some brown color of NO2 escaping from a beaker or a test tube, then assume that inhaling the air from the room, in which the NO2 is released is potentially dangerous and ventilate well for some time and leave the room immediately. Experiments with NO2 should be conducted outside or in a good fume hood.
Oxidation state +3: The most common compounds with nitrogen in this oxidation state are the nitrites. Sodium nitrite and potassium nitrite are really interesting chemicals and allow many interesting experiments to be performed. In aqueous environments at room temperature, nitrites are much more reactive than nitrates. They can act both as oxidizer or as reductor. Many interesting compounds can be made with nitrites, e.g. coordination compounds, nitrogen monoxide, nitrosyl-compounds and organic nitrites. On this site there are descriptions of several experiments with nitrites. Sodium nitrite can be ordered from some photography raw chemical suppliers at fairly low prices. Sodium nitrite and potassium nitrite are very pale yellow crystalline hygroscopic solids.
Oxidation state +2: This oxidation state is encountered in nitrogen monoxide (NO). This gas cannot be purchased easily by the general public, but it can be prepared easily when one has access to sodium nitrite or potassium nitrite. It is a colorless gas, which is oxidized by oxygen from the air at once to brown nitrogen dioxide. For some experiments, the gas NO is interesting. When prepared it must be used immediately. For NO, the same warning holds as for NO2, because of the fact that NO is quickly converted to NO2 in air.
Oxidation state +1: This is encountered in the colorless gas dinitrogen dioxide (laughing gas). It is mildly anesthetic, but otherwise non-toxic. It supports combustion of many flammable compounds. Preparation of this gas in the home lab is not really easy. It can be prepared by heating a mixture of a nitrate and an ammonium salt, but that reaction is not without risk and may also produce impurities, due to side reactions. The gas is rather inert at room temperature in aqueous chemistry. It is not really interesting for the home chemist.
Other compounds of the +1 oxidation state exist, e.g. H2N2O2 and its salts, but these are very unstable and play no role in the home lab.
Oxidation state 0: As already mentioned, the element itself is a colorless and very inert gas, which is not interesting from a chemical point of view in the home lab.
Oxidation state –⅓: A special group of compounds are formed by the azides, which contain the anionic species N3–. This is a colorless ion. The only salt of this, which can be used in the home lab is sodium azide, NaN3. Sodium azide is a white solid. This can only be obtained from chemical supply houses, and from the canisters of solid, present in air bags of cars. The latter source of sodium azide is impure, usually mixed with sodium carbonate and silica.
Sodium azide is a dangerous compound. It is very toxic (close to toxicity of sodium cyanide) and it is quite unstable. With transition metals it can form interesting coordination complexes, but with certain metals, also extremely unstable and highly explosive precipitates can be formed. With acids, the explosive and extremely poisonous hydrazoic acid, HN3, is formed. This acid must be handled with extreme care.
Experimenting with azides absolutely is not something to start with and should only be conducted when one has quite some experience in safely handling chemicals, this is due to both toxicity and explosive properties.
Oxidation state –1: This oxidation state is represented by a group of compounds, derived from hydroxylamine, NH2OH, or its protonated ion NH3OH+. Free hydroxyl amine is not suitable for use in the home lab and it also is very hard to obtain for the general public. Free hydroxyl amine is notoriously unstable and is very risky on storage. Salts of hydroxyl amine, however, are stable and can be kept well without introducing serious risks. Hydroxyl amine and to a much lesser extent its salts are very reactive compounds and due to their reactivity are interesting candidates for a nice home lab. Hydroxyl amine sulfate or hydroxyl amine chloride can be ordered at photography raw chemical suppliers. These salts are white solids. Like nitrites, these compounds have an extensive aqueous chemistry, both in redox reactions and in formation of coordination complexes. Addition of strong base to an hydroxyl amine salt, releases the free compound, present at a strong dilution in alkaline aqueous solution. These dilute hydroxyl amine solutions have interesting properties, which allow many interesting experiments to be performed. Such dilute solutions of hydroxyl amine can be handled safely.
Oxidation state –2: This oxidation state is represented by a group of compounds, derived from hydrazine, H2NNH2, or its protonated ions H2NNH3+ (usually written as N2H5+) and +H3NNH3+ (usually written as N2H62+). For hydrazine, the same holds as for hydroxyl amine. The free compound or its hydrate is not suitable for the home lab, but its salts can be interesting. If, however, one has access to hydroxyl amine salts, then the purchase of hydrazine salts is not that interesting anymore. From an experimenter's point of view, the properties of hydroxyl amine and hydrazine are quite similar. Hydrazine and its salts are known carcinogens, hence the use of salts of hydroxyl amine in general is preferred.
Oxidation state –3: This oxidation state, together with oxidation state +5 is the most common oxidation state of nitrogen in its compounds. The best known compound with this oxidation state is ammonia, NH3. This is a colorless gas with a penetrating odor, which dissolves in water exceedingly well. Concentrated ammonia solutions are very noxious and one should be careful not to breathe too much of the gas. However, the gas has a good warning level. Before the gas really becomes toxic, the stench and irritation of the gas becomes unbearable, so one automatically will leave the area, when too much ammonia is present. Ammonia is encountered in every day life frequently, especially when one is living in a rural area. Ammonia solutions are interesting for the home lab. Dilute solutions, up to approximately 10% by weight, can be purchased in supermarkets and drugstores, more concentrated solutions can be purchased at better equipped drugstores.
Ammonia has a protonated ion, called ammonium, NH4+. Many compounds with an interesting anion are available as the ammonium salt, because ammonium ion is a rather inert ion in neutral to acidic environments and ammonium salts in general can be dissolved in water easily. Examples of these are ammonium dichromate, ammonium persulfate, ammonium metavanadate. For the home lab it might be interesting to have an ammonium salt with an inert anion at hand, such as ammonium chloride or ammonium sulfate. These ammonium salts can be purchased at drugstores and they can be ordered at photography raw chemical suppliers.
Other compounds of nitrogen in its -3 oxidation state are amides, containing NH2– ions and substituted ammonia/ammonium compounds, such as methyl amine or its salts. The amides and the amines are not very suitable for the home lab. Amides have no aqueous chemistry, they immediately react with water, forming ammonia and hydroxide. Amides also do not keep well, because of their sensitivity to moisture. Organic amines might be interesting, but if one wants to experiment with them, then usually it is sufficient to take only one of them. The sulfate salt of ethylenediamine (or better, 1,2-diaminoethane) can be purchased by the public from photography raw chemical suppliers. Methylamine and methylammonium salts unfortunately are not easily obtained, although they are fairly benign. The reason of this is that they are abused much for making drugs, and this has resulted in strong regulation of methylamine and salts, derived from this.
Mixed oxidation state (+3, +1): Nitrogen also forms mixed oxidation state compounds, of which the nitramides are most noteworthy. These are derived from ammonia, in which H-atoms are replaced by nitro groups. A compound, which is known for a long time already is H2NNO2. This is a reactive and not easily handled compound, not suitable for home chemistry. However, in 1971, the anion of disubstituted nitramide, HN(NO2)2, was discovered. The disubstituted nitramide has acidic properties. The acid itself is unstable, but the anion is remarkably stable. Nowadays, compounds like K[N(NO2)2] and NH4[N(NO2)2] are available. The ion N(NO2)2– is called dinitramide ion, and the salts are called dinitramides. This ion contains one nitrogen atom in the +1 oxidation state and two nitrogen atoms in the +3 oxidation state. It usually is simply written as N3O4–. Currently, this is not easily obtained by the general public, but in the near future, this may become a widespread compound of nitrogen. Lots of research is carried out at the moment for making production of dinitramide salts cheaper and easier. The dinitramide ion is a very nice oxidizing agent at higher temperatures, with stability and performance comparable to that of perchlorate. Current research is driven by environmental concerns. Perchlorate produces lots of toxic smoke (chlorinated compounds) when used as part of rocket driver, dinitramides only produce non-toxic or less toxic products, when used as rocket driver.
Other 'mixed oxidation state compounds' of nitrogen are ammonium salts of azide, nitrate or nitrite. Hydroxylamine-derived and hydrazine-drived salts of nitrate also exists. These seemingly mixed oxidation state compounds are not true mixed oxidation state compounds. They contain anions and cations, which both contain nitrogen, but the ions exist on their own. The best known of this class of compounds is ammonium nitrate (a common fertilizer).