Nitrates: properties, structure, nomenclature, formation

Nitrates chemical structure
Nitrate anion is represented by a ball-and-stick model. Source: Benjah-bmm27 / Public domain.

The term ‘nitrates’ immediately refers to salts, fertilizers, and explosives. They are also related to meats and their curing processes to inhibit the growth of bacteria and thus delay their deterioration. Nitrates are also synonymous with vegetables, as they contain high amounts of these salts.

NO  (upper image) is the most oxidized form of nitrogen, being the final and assimilable product of its biological cycle. Nitrogen in the atmosphere undergoes various transformations, either by electric rays or by microbial actions, to become NH + and NO , both soluble in water and absorbable by plants.

Properties of nitrates


Nitrates are in principle neutral substances since NO  is the conjugate base of nitric acid, HNO 3, a strong acid. In water, it does not hydrolyze appreciably:

NO  + H 2 O ⇌ HNO 3 + OH 

In fact, it can be said that this hydrolysis does not occur at all, the amounts of OH  being negligible. Therefore, aqueous solutions of nitrates are neutral unless other substances are dissolved.

Physical appearances

Inorganic nitrates consist of salts whose general formulas are M (NO 3 ) n, where n is the valence of the positive charge of the metal cation. All these nitrates have crystalline brilliance and are whitish in color. However, transition metal nitrates can be colorful.

For example, copper (II) nitrate or cupric nitrate, Cu (NO 3 ) 2, is a bluish crystalline solid. Meanwhile, nickel (II) nitrate, Ni (NO 3 ) 2, is an emerald solid. Some other nitrates, such as those of iron, are faintly colored.

Thermal decomposition

Metal nitrates are sensitive to high temperatures, as they begin to break down according to the following chemical equation:

2M (NO 3 ) 2 (s) → 2MO (s) + 4NO 2 (g) + O 2 (g)

As seen, metal nitrate decomposes into an oxide, MO, nitrogen dioxide, and oxygen. This thermal decomposition does not occur at the same temperature for all nitrates; some resist more than others.

As a general rule, the larger and smaller the charge on the M + cation, the higher the temperature to which nitrate must be heated to decompose. On the other hand, when M + is small or has a large positive charge, nitrate decomposes at lower temperatures, being, therefore, more unstable.

For example, sodium nitrate, NaNO 3, decomposes at a lower temperature than barium nitrate, Ba (NO 3 ) 2, because although Ba 2+ has a higher charge than Na +, its size is much larger.

Oxidizing agent

NO  is a relatively stable anion. However, its nitrogen atom is strongly oxidized, with an oxidation state of +5 (N 5+ O 2- ), so it is “thirsty” for electrons. For this reason, nitrate is an oxidizing agent, which will seek to steal electrons from substances around it.

It is this lack of electrons in the nitrogen atom of NO  which makes NaNO 3 and KNO 3 strong oxidizing agents, used as components of gunpowder. Adding to this characteristic the fact that NO 2 and O 2 gases are released when decomposing, it is understood why it is part of many explosives.

When nitrate gains electrons or is reduced, it can transform into ammonia, NH 3, or nitric oxide, NO, depending on the reactants and conditions.


All inorganic nitrates, or what is the same, as metal and ammonium nitrates, NH 4 NO 3, are water-soluble compounds. This is because water molecules have a strong affinity for NO  since the crystal lattices of these nitrates are not very stable.

Structure of Nitrate anion

Nitrate anion

Nitrate anion
Nitrate resonance structures. Source: Benjah-bmm27 / Public domain.

The upper image shows the resonance structures of the nitrate anion. As can be seen, two negative charges are located on two oxygen atoms, which are delocalized between the three oxygen atoms. Therefore, each O has a charge of -2/3, while nitrogen has a charge of +1.

Thus, NO  interacts electrostatically, or forms coordination bonds, through any of its three oxygen atoms.


All inorganic nitrates are saline and ionic. Therefore, its structures are crystalline, which means that its ions, M + and NO , are arranged in an orderly space thanks to their electrostatic interactions. Depending on these interactions, your crystals will have different structures.

For example, the crystal structure of NaNO 3 is trigonal or rhombohedral, while that of KNO 3 is orthorhombic.

Organic nitrates

Organic nitrates are represented by the general formula RONO 2, where R is an alkyl or aromatic group. These compounds are characterized by their R-ONO 2 bond and usually consist of nitric derivatives of polymers and other organic substances.


NO  coordinates with metal centers to form an M + -ONO  bond, being an interaction different from ionic. These complexes are essentially inorganic in nature. Nitrate can even coordinate using two of its oxygen atoms at the same time, M + —O 2 NO.


To name a nitrate, the words ‘nitrate of’ must first be written followed by the name of the cation or the R group that accompanies it in its respective chemical formula. The valence of the cation is specified in parentheses when it has more than one. Likewise, the suffixes –ico and –also can be used if preferred, following the traditional nomenclature.

For example, consider Fe (NO 3 ) 3 . Its name is iron (III) nitrate, because its valence is +3, or it can also be called ferric nitrate.

These rules also apply to organic nitrates, as long as their structures are simple. For example, CH 3 ONO 2 is called methyl nitrate, since the –CH 3 group becomes the R group that accompanies –ONO 2.



Nitrates are formed in nature as part of the biological nitrogen cycle. Because of this, soils, seas, and some streams have significant amounts of nitrates. Depending on the surrounding ions, different nitrate salts will be formed, with NaNO 3 and KNO 3 being the most common and abundant.


Nitrates are formed on industrial scales by neutralizing nitric acid, either with metal hydroxides, alcohols, polymers, etc. For example, calcium nitrate, Ca (NO 3 ) 2, can be prepared according to the following chemical equation:

Ca (OH) 2 + 2HNO 3 → Ca (NO 3 ) 2 + 2H 2 O

Similarly, various organic substances are attacked by HNO 3 under certain conditions to promote the substitution of some of its groups by -ONO 2 . This is what happens with the reaction of cellulose to transform into nitrocellulose or cellulose nitrate, an explosive polymeric material.


The NO  anion, and therefore inorganic nitrates, can be formed by the photocatalytic action of titanium oxide waste, TiO 2, using nothing more than nitrogen and oxygen from the air as raw material. This study assumes that where there are excesses of TiO 2, there will be unwanted amounts of NO , which affect the potability of the water and can even plague them with algae.

Applications of Nitrates

Curing of meats

Sausages are one of the processed meats that contain the most nitrates. Source: Pxhere.

Nitrates are added to various meats, especially sausages, to eliminate bacteria and thus delay their deterioration. Likewise, they react with their proteins to give them a more reddish color. The problem with these meats is that, when cooked at high temperatures, they produce nitrosamines: compounds linked to colon cancer.

This reaction is partly avoided if there are vitamins present, as is the case with vegetables, which, although rich in nitrates are not associated with carcinogenic pathologies.


Nitrates are a soluble source of nitrogen. Therefore, it serves as a fertilizer to provide nitrogen to the plants, and thus, favors their growth.


Nitrates have a special function in the body. When assimilated by enzymatic action, it is reduced to nitric oxide, NO. This molecule occupies volume and dilates the veins and arteries, allowing more blood flow. Therefore, nitrates are used as drugs to combat pathologies of the circulatory system.

Ironically and curiously, organic nitrates such as glyceryl trinitrate, nitroglycerin, isosorbide mononitrate, and pentaerythritol tetranitrate have been used for this purpose, all good candidates in the world of explosives.


Nitrates are used in explosive formulations, with gunpowder being the most symbolic example. As they are oxidizing agents, they favor the combustion of matter, in addition to contributing to the abrupt expansion of the volume due to its release of gases after decomposing.

Examples of nitrates

Throughout the previous sections, more than one example of nitrates has been mentioned. Finally, some others will be listed together with their respective formulas and names:

  • LiNO 3: lithium nitrate
  • RbNO 3: rubidium nitrate
  • Mg (NO3)2: magnesium nitrate
  • Sc (NO3)2: scandium nitrate
  • Cr (NO3)3: chromium (III) nitrate
  • Zn (NO3)2: zinc nitrate
  • Pb (NO3)2: lead (II) nitrate
  • AgNO3: silver nitrate
  • CH3CH2ONO2: ethyl nitrate
  • CH3(CH2)4 ONO2: amyl nitrate

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