Nitrate is a nitrogen oxoanion formed by the loss of hydrogen from nitric acid with a charge of -1, one nitrogen atom, and three oxygen atoms.
- It is a polyatomic ion with the molecular formula NO3– that is commonly used in the form of salts in fertilizers and explosives.
- Nitrate is a monovalent inorganic anion and a conjugate base of nitric acid. It can form salts with various metals.
- Nitrate ion has a trigonal planar structure with a central nitrogen atom that is surrounded by three identically bonded oxygen atoms.
- The overall charge of the ion is -1, where each oxygen atom carries a -2/3 charge while nitrogen carries a charge of +1.
- Nitrates can acts as oxidizing agents as the nitrogen atom in the molecule has an oxidation state of +5.
- Nitrates are also essential in the diet of living beings which is consumed in the form of inorganic nitrates present in leafy vegetables. Drinking water is also a dietary source of nitrates.
- Nitrate salts occur on earth in the form of deposits and rocks. These are produced from nitrogen compounds by nitrifying bacteria.
- The most important use of nitrate is as fertilizers in agriculture as a result of their high solubility and biodegradability.
- In industries, nitrates are used as oxidizing agents during the production of metal products.
- Some examples of nitrates include sodium nitrate, aluminium nitrate, magnesium nitrate, etc.
Nitrite is a polyatomic anion of nitrogen formed by the loss of hydrogen atoms from nitrous acid with the charge of -1, one nitrogen atom, and two oxygen atoms.
- Nitrites are primarily produced as intermediate products during the oxidation of ammonia to nitrate.
- The molecular formula of nitrite is NO2– and it is the conjugate acid of nitrous acids, HNO2.
- However, the term nitrite might also be used to refer to organic compounds with the –ONO group that are esters of nitrous acid.
- Nitrites can act both as oxidizing and reducing agents as the nitrogen atom in the molecule has the oxidation state of +3.
- The structure of the nitrite ion has a bent geometry with the O-N-O bond angle around 120°C.
- The ion has a symmetrical structure where both of the N-O bonds have equal lengths and bond angles. The ion exists in two distinct forms that are mirror images of each other, and thus, it has a resonance hybrid.
- Nitrite is an essential intermediate in the nitrogen cycle as it is reduced to nitric oxide or ammonia in the presence of nitrifying bacteria.
- It is also used in azo dyes and other colorants via diazotization. Some nitrites like sodium nitrite are also used in food preservation.
- The presence of nitrite above the desired levels in food and water samples can result in various diseases in humans. Nitrites in diet form N-nitrosamines which are likely to cause stomach cancer.
13 Key Differences (Nitrate vs Nitrite)
|Definition||Nitrate is a nitrogen oxoanion formed by the loss of hydrogen from nitric acid with a charge of -1, one nitrogen atom, and three oxygen atoms.||Nitrite is a polyatomic anion of nitrogen formed by the loss of hydrogen atoms from nitrous acid with the charge of -1, one nitrogen atom, and two oxygen atoms.|
|Oxidation State||The oxidation state of nitrogen in nitrates is +5.||The oxidation state of nitrogen in nitrite is +3.|
|Structure||Nitrates have a trigonal planar structure.||Nitrites have a bent molecular structure.|
|Molecular Formula||The molecular formula of nitrate is NO3–.||The molecular formula of nitrite is NO2–.|
|Molar Mass||The molar mass of nitrate is 62 g/mol.||The molar mass of nitrite is 46.01 g/mol.|
|Oxidizing Nature||Nitrates act as oxidizing agents.||Nitrites can act as both oxidizing and reducing agents.|
|Atoms||Nitrates consist of a nitrogen atom and three oxygen atoms.||Nitrites consist of a nitrogen atom and two oxygen atoms.|
|Conjugate Acid||The conjugate acid of nitrates is nitric acid.||The conjugate acid of nitrite is nitrous acid.|
|Chemical Process||Nitrates are reduced to form nitrites.||Nitrites are oxidized to form nitrates.|
|Stability||Nitrates are comparatively more stable.||Nitrites are comparatively less stable than nitrates.|
|Hazardous Concentrations||The lethal dose of nitrates is usually high as these are less hazardous.||The lethal dose of nitrite is usually high as it is more hazardous.|
|Uses||Nitrates salts are used in fertilizers and explosives.||Inorganic nitrites are used as food preservatives.|
|Examples||Some examples of nitrates include sodium nitrate, aluminium nitrate, magnesium nitrate, etc.||Some examples of nitrites include sodium nitrite, potassium nitrite, calcium nitrite, etc.|
Examples of Nitrate
- Magnesium nitrate is an inorganic nitrate salt of magnesium with the molecular formula- Mg(NO3)2.
- Magnesium nitrate is a hygroscopic substance that quickly forms hexahydrate if left in the air. The hygroscopic form of the compound is Mg(NO3)2.6H2O.
- It exists in the form of a white crystalline solid and produces toxic oxides of nitrogen when heated.
- It can be naturally extracted from mines and caverns, but the commercial production of magnesium nitrate is by the reaction of nitric acid and magnesium salts.
- Magnesium nitrate readily reacts with alkali metal hydroxides to form alkali nitrates.
- As it is hygroscopic and has a high affinity for water, combustion of magnesium nitrate results in the decomposition into magnesium oxide, oxygen, and other oxides.
Examples of Nitrite
- Sodium nitrite is an inorganic sodium salt with the molecular formula NaNO2.
- It exists in the form of a yellowish-white crystalline solid that is commonly used as a food preservative.
- It is non-combustible by itself but might accelerate the combustion of other materials. In large quantities, however, it might result in fire and explosions.
- Sodium nitrite has multiple industrial applications and is often considered the most important nitrite salt.
- It is used as a precursor in various products like medicines, dyes, and fertilizers, but its most common use is as a food additive.
- Despite its applications in industries, it can be lethal to humans if consumed at a dose of more than 71 mg/kg.
- Even though it doesn’t occur naturally in vegetables, high levels of sodium nitrite have been observed in such products due to its overuse as fertilizers.
- Gautum SD, Pant M and Adhikari NR (2016). Comprehensive Chemistry, Part 2. Sixth Edition. Heritage Publishers and Distributors Pvt. Ltd
- National Center for Biotechnology Information. PubChem Compound Summary for CID 943, Nitrate. https://pubchem.ncbi.nlm.nih.gov/compound/Nitrate. Accessed Mar. 26, 2021.
- National Center for Biotechnology Information. PubChem Compound Summary for CID 25212, Magnesium nitrate. https://pubchem.ncbi.nlm.nih.gov/compound/Magnesium-nitrate. Accessed Mar. 26, 2021.
- National Center for Biotechnology Information. PubChem Compound Summary for CID 23668193, Sodium nitrite. https://pubchem.ncbi.nlm.nih.gov/compound/Sodium-nitrite. Accessed Mar. 26, 2021.