Difference between revisions of "Nitrate" - New World Encyclopedia
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[[Image:Nitrate-ion-elpot.png|thumb|right|200px|An [[electric potential|electrostatic potential]] map of the nitrate ion. Areas colored red are lower in energy than areas colored yellow.]] | [[Image:Nitrate-ion-elpot.png|thumb|right|200px|An [[electric potential|electrostatic potential]] map of the nitrate ion. Areas colored red are lower in energy than areas colored yellow.]] | ||
| − | In [[inorganic chemistry]], a '''nitrate''' is a [[salt]] of [[nitric acid]] with an [[ion]] comprised of one [[nitrogen]] and three [[oxygen]] atoms. In [[organic chemistry]], the [[esters]] of nitric acid and various [[alcohols]] are called nitrates. | + | In [[inorganic chemistry]], a '''nitrate''' is a [[salt]] of [[nitric acid]] with an [[ion]] comprised of one [[nitrogen]] and three [[oxygen]] atoms. In [[organic chemistry]], the [[esters]] of nitric acid and various [[alcohols]] are called nitrates. Nitrates in the soil are important as nutrients for plant growth, but their use in artificial fertilizers has led to pollution of groundwater and surface waters in various agricultural regions. |
| − | ==Occurrence and | + | ==Occurrence, history, and production== |
[[Image:AYool WOA surf NO3.png|thumb|right|250px|Annual mean sea surface '''nitrate''' for the [[World Ocean]]. Data from the [[World Ocean Atlas]] [http://www.nodc.noaa.gov/OC5/WOA01/ 2001].]] | [[Image:AYool WOA surf NO3.png|thumb|right|250px|Annual mean sea surface '''nitrate''' for the [[World Ocean]]. Data from the [[World Ocean Atlas]] [http://www.nodc.noaa.gov/OC5/WOA01/ 2001].]] | ||
| − | Solid nitrates are not very abundant in | + | Solid nitrates are not very abundant in nature as they are very soluble. They can appear where nitrogen-containing groundwater is evaporating (such as in soils of arid regions and on animal shed walls). [[Nitrification]] bacteria in the soil are also needed for the process. |
| − | Until the early part of the twentieth century, there were no known methods for the chemical synthesis of nitrates. Chile was a major exporter, and European countries | + | The first commercially exploited source was [[India]], providing the [[British Empire]] with a reliable supply. By contrast the European continental powers had to collect scrapings from walls and barns, install saltpeter farms (based on aging and leaching manure and urine). The chemist [[Lavoisier]] was also a tax collector and commissioner of the Saltpeter Administration. Later, the large deposits of [[#Sodium nitrate|sodium nitrate]] in the [[Atacama Desert]] of northern Chile acquired economic significance. |
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| + | Until the early part of the twentieth century, there were no known methods for the chemical synthesis of nitrates. Chile was a major exporter, and European countries were dependent on its nitrates for use as fertilizer to feed their people. Nitrates were needed to produce military explosives as well. These two uses influenced world history in significant ways. Had the Germans not devised the [[Haber process|Haber]] and [[Ostwald process|Ostwald]] processes for producing nitrate, they would not have been able to feed their civilian population and armies, nor continued to make explosives. The First World War might have ended as a direct result of embargo of essential raw materials. With the aid of organic chemistry, however, the war continued. Nowadays, most nitrates are produced from [[ammonia]] synthesized from atmospheric nitrogen. | ||
==Chemical properties== | ==Chemical properties== | ||
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The nitrate ion is the [[conjugate acid|conjugate base]] of nitric acid. A nitrate salt forms when a positively charged ion attaches to the negatively charged oxygen atoms of the ion, forming an ionic [[chemical compound|compound]]. Almost all nitrates are [[solubility|soluble]] in [[water]] at [[standard temperature and pressure]]. | The nitrate ion is the [[conjugate acid|conjugate base]] of nitric acid. A nitrate salt forms when a positively charged ion attaches to the negatively charged oxygen atoms of the ion, forming an ionic [[chemical compound|compound]]. Almost all nitrates are [[solubility|soluble]] in [[water]] at [[standard temperature and pressure]]. | ||
| − | In [[organic chemistry]] a nitrate is a [[functional group]] with general chemical formula RONO<sub>2</sub> where R stands for any organic residue. They are the [[esters]] of nitric acid and [[alcohols]] formed by | + | In [[organic chemistry]] a nitrate is a [[functional group]] with general chemical formula RONO<sub>2</sub>, where R stands for any organic residue. They are the [[esters]] of nitric acid and [[alcohols]] formed by the process known as ''nitroxylation''. Examples are '''methyl nitrate''' formed by reaction of [[methanol]] and nitric acid <ref>''Organic Syntheses'', Coll. Vol. 2, p.412 (1943); Vol. 19, p.64 (1939) [http://www.orgsynth.org/orgsyn/pdfs/CV2P0412.pdf Link]</ref>, the nitrate of [[tartaric acid]],<ref>''Organic Syntheses, Coll. Vol. 3, p.471 (1955); Vol. 22, p.65 (1942) [http://www.orgsynth.org/orgsyn/pdfs/CV3P0471.pdf Link]</ref> and the inappropriately named [[nitroglycerin]]. |
== Effects on aquatic life == | == Effects on aquatic life == | ||
| − | In [[freshwater]] or [[estuary|estuarine]] systems close to land, nitrate can reach high levels | + | In [[freshwater]] or [[estuary|estuarine]] systems close to land, nitrate concentrations can reach high levels, potentially causing the death of fish. Although the nitrate ion is much less toxic than ammonia or nitrite, levels over 30 parts per million (ppm) of nitrate can inhibit growth, impair the immune system, and cause stress in some aquatic species. |
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| + | In most cases, high nitrate concentrations in aquatic systems are the result of [[surface runoff]] from agricultural or [[landscape]]d areas that have received excess nitrate fertilizer. These levels of nitrate can also lead to algal blooms, and when nutrients (such as potassium, phosphate, or nitrate) become limiting, [[eutrophication]] can occur. Besides leading to water [[anoxia]], these blooms may cause other changes to [[ecosystem]] functions, favoring some groups of organisms over others. Consequently, as nitrates form a component of [[total dissolved solids]], they are widely used as indicators of [[water quality]]. | ||
== Specific nitrates == | == Specific nitrates == | ||
Revision as of 14:58, 22 March 2007
In inorganic chemistry, a nitrate is a salt of nitric acid with an ion comprised of one nitrogen and three oxygen atoms. In organic chemistry, the esters of nitric acid and various alcohols are called nitrates. Nitrates in the soil are important as nutrients for plant growth, but their use in artificial fertilizers has led to pollution of groundwater and surface waters in various agricultural regions.
Occurrence, history, and production
Solid nitrates are not very abundant in nature as they are very soluble. They can appear where nitrogen-containing groundwater is evaporating (such as in soils of arid regions and on animal shed walls). Nitrification bacteria in the soil are also needed for the process.
The first commercially exploited source was India, providing the British Empire with a reliable supply. By contrast the European continental powers had to collect scrapings from walls and barns, install saltpeter farms (based on aging and leaching manure and urine). The chemist Lavoisier was also a tax collector and commissioner of the Saltpeter Administration. Later, the large deposits of sodium nitrate in the Atacama Desert of northern Chile acquired economic significance.
Until the early part of the twentieth century, there were no known methods for the chemical synthesis of nitrates. Chile was a major exporter, and European countries were dependent on its nitrates for use as fertilizer to feed their people. Nitrates were needed to produce military explosives as well. These two uses influenced world history in significant ways. Had the Germans not devised the Haber and Ostwald processes for producing nitrate, they would not have been able to feed their civilian population and armies, nor continued to make explosives. The First World War might have ended as a direct result of embargo of essential raw materials. With the aid of organic chemistry, however, the war continued. Nowadays, most nitrates are produced from ammonia synthesized from atmospheric nitrogen.
Chemical properties
The nitrate ion is a polyatomic ion with the empirical formula NO3− and a molecular mass of 62.0049. It consists of one central nitrogen atom surrounded by three identical oxygen atoms in a trigonal planar arrangement. The nitrate ion carries a negative one formal charge and can be represented as a hybrid of the following resonance structures:
The nitrate ion is the conjugate base of nitric acid. A nitrate salt forms when a positively charged ion attaches to the negatively charged oxygen atoms of the ion, forming an ionic compound. Almost all nitrates are soluble in water at standard temperature and pressure.
In organic chemistry a nitrate is a functional group with general chemical formula RONO2, where R stands for any organic residue. They are the esters of nitric acid and alcohols formed by the process known as nitroxylation. Examples are methyl nitrate formed by reaction of methanol and nitric acid [1], the nitrate of tartaric acid,[2] and the inappropriately named nitroglycerin.
Effects on aquatic life
In freshwater or estuarine systems close to land, nitrate concentrations can reach high levels, potentially causing the death of fish. Although the nitrate ion is much less toxic than ammonia or nitrite, levels over 30 parts per million (ppm) of nitrate can inhibit growth, impair the immune system, and cause stress in some aquatic species.
In most cases, high nitrate concentrations in aquatic systems are the result of surface runoff from agricultural or landscaped areas that have received excess nitrate fertilizer. These levels of nitrate can also lead to algal blooms, and when nutrients (such as potassium, phosphate, or nitrate) become limiting, eutrophication can occur. Besides leading to water anoxia, these blooms may cause other changes to ecosystem functions, favoring some groups of organisms over others. Consequently, as nitrates form a component of total dissolved solids, they are widely used as indicators of water quality.
Specific nitrates
Ammonium nitrate
Ammonium nitrate (NH4NO3) is commonly used in agriculture as a high-nitrogen fertilizer. It can also be used as an oxidizing agent in explosives, especially improvised explosive devices. These uses have raised two types of concerns:
- The pollution of groundwater by excess nitrate from fertilizers.
- The potential use of ammonium nitrate for terrorist activities.
Potassium nitrate
Potassium nitrate (KNO3) is a naturally occurring mineral source of nitrogen. Its common names include saltpeter (saltpetre), nitrate of potash, and nitre. It is used in the production of nitric acid, model rocket propellants, and several types of fireworks. In addition, it is a fertilizer and food preservative. Although also used in gunpowder, it is not combustible or flammable by itself.
Sodium nitrate
Sodium nitrate (NaNO3) is a type of salt that has long been used as an ingredient in explosives and solid rocket propellants, in glass and pottery enamel, and as a food preservative (such as in hot dogs), and has been mined extensively for these purposes. It is also variously known as caliche, Chile saltpeter, saltpeter, and soda niter. Chile has the largest reserves of caliche. It can also be manufactured synthetically.
Related materials
- Nitrates should not be confused with nitrites, the salts of nitrous acid.
- Organic compounds containing the nitro (NO2) functional group are known as nitro compounds.
See also
Notes
ReferencesISBN links support NWE through referral fees
- Addiscott, T.M., A.J. Gold, C.A. Oviatt, N. Benjamin, and K.E. Giller. 2005. Nitrate, Agriculture and the Environment. CABI Publishing. ISBN 0851999131 and ISBN 978-0851999135.
- Chang, Raymond. 2006. Chemistry. 9th ed. New York: McGraw-Hill Science/Engineering/Math. ISBN 0073221031.
- Cotton, F. Albert, and Geoffrey Wilkinson. 1980. Advanced Inorganic Chemistry. 4th ed. New York: Wiley. ISBN 0-471-02775-8.
- Razowska-Jaworek, Lidia, and Andrzej Sadurski. 2005. Nitrates in Groundwater. International Association of Hydrogeologists Selected Papers. Leiden, the Netherlands: A.A. Balkema. ISBN 9058096645 and ISBN 978-9058096647.
External links
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