Nitrate

From Wikipedia, the free encyclopedia
(Redirected from Nitrates)
For the functional group –ONO
2, see Nitrate ester. For that functional group in medicine, see Nitrovasodilator.
Not to be confused with NO−
2, nitrite.
This article is about the ion. For the radical, see nitrogen trioxide.
For other uses, see Nitrate (disambiguation).
Nitrate
Ball-and-stick model of the nitrate ion
Names
Systematic IUPAC name
Nitrate
Identifiers
3D model (JSmol)
ChEBI
ChemSpider
UNII
  • InChI=1S/NO3/c2-1(3)4/q-1
    Key: NHNBFGGVMKEFGY-UHFFFAOYSA-N
  • InChI=1/NO3/c2-1(3)4/q-1
    Key: NHNBFGGVMKEFGY-UHFFFAOYAI
  • [N+](=O)([O-])[O-]
Properties
NO−
3
Molar mass 62.004 g·mol−1
Conjugate acid Nitric acid
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).

Nitrate is a polyatomic ion with the chemical formula NO−
3. Salts containing this ion are called nitrates. Nitrates are common components of fertilizers and explosives.[1] Almost all inorganic nitrates are soluble in water. An example of an insoluble nitrate is bismuth oxynitrate.

Chemical structure[edit]

The nitrate ion with the partial charges shown

The nitrate anion is the conjugate base of nitric acid, consisting of one central nitrogen atom surrounded by three identically bonded oxygen atoms in a trigonal planar arrangement. The nitrate ion carries a formal charge of −1.[citation needed] This charge results from a combination formal charge in which each of the three oxygens carries a −2⁄3 charge,[citation needed] whereas the nitrogen carries a +1 charge, all these adding up to formal charge of the polyatomic nitrate ion.[citation needed] This arrangement is commonly used as an example of resonance. Like the isoelectronic carbonate ion, the nitrate ion can be represented by three resonance structures:

Canonical resonance structures for the nitrate ion

Chemical and biochemical properties[edit]

In the NO−3 anion, the oxidation state of the central nitrogen atom is V (+5). This corresponds to the highest possible oxidation number of nitrogen. Nitrate is a potentially powerful oxidizer as evidenced by its explosive behaviour at high temperature when it is detonated in ammonium nitrate (NH4NO3), or black powder, ignited by the shock wave of a primary explosive. However, in contrast to red fuming nitric acid (HNO3/N2O4), or concentrated nitric acid (HNO3), nitrate dissolved in aqueous solution at neutral or high pH is only a weak oxidizing agent and is stable under sterile, or aseptic, conditions, in the absence of microorganisms. To increase its oxidation power, acidic conditions and high concentrations are needed, under which nitrate transforms into nitric acid. This behaviour is consistent with the general theory of reduction-oxidation (redox) in electrochemistry: oxidizing power is exacerbated under acidic conditions while the power of reducing agents is reinforced under basic conditions. This can be illustrated by means of a Pourbaix diagram (Eh–pH diagram) drawn using the Nernst equation and the corresponding redox reactions. During the reduction of oxidizers, the oxidation state decreases and oxide ions (O2−) in excess released in water by the reaction are more easily protonated under acid conditions (O2− + 2 H+ → H2O) which drives the reduction reaction to the right according to Le Chatelier's principle. For the oxidation of reducing agents, the reverse occurs: as the oxidation state increases, oxide anions are needed to neutralise the surplus of positive charges born by the central atom. As basic conditions favor the production of oxide anions (2 OH− → O2− + H2O), this drives the chemical equilibrium of the oxidation reaction to the right.

Meanwhile, nitrate is used as a powerful terminal electron acceptor by denitrifying bacteria to deliver the energy they need to thrive. Under anaerobic conditions, nitrate is the strongest electron acceptor used by prokaryote microorganisms (bacteria and archaea) to respirate. The redox couple NO−3/N2 is at the top of the redox scale for the anaerobic respiration, just below the couple oxygen (O2/H2O), but above the couples Mn(IV)/Mn(II), Fe(III)/Fe(II), SO2−4/HS−, CO2/CH4. In natural waters, inevitably contaminated by microorganisms, nitrate is a quite unstable and labile dissolved chemical species because it is metabolised by denitrifying bacteria. Water samples for nitrate/nitrite analyses need to be kept at 4 °C in a refrigerated room and analysed as quick as possible to limit the loss of nitrate.

In the first step of the denitrification process, dissolved nitrate (NO−3) is catalytically reduced into nitrite (NO−2) by the enzymatic activity of bacteria. In aqueous solution, dissolved nitrite, N(III), is a more powerful oxidizer that nitrate, N(V), because it has to accept less electrons and its reduction is less kinetically hindered than that of nitrate.

During the biological denitrification process, further nitrite reduction also gives rise to another powerful oxidizing agent: nitric oxide (NO). NO can fix on myoglobin accentuating its red coloration. NO is an important biological signaling molecule and intervenes in the vasodilation process, but it can also produce free radicals in biological tissues, accelerating their degradation and aging process. The reactive oxygen species (ROS) generated by NO contribute to the oxidative stress, a condition involved in vascular dysfunction and atherogenesis.[2]

Detection in chemical analysis[edit]

The nitrate anion is commonly analysed in water by ion chromatography (IC) along with other anions also present in solution. The main advantage of IC is its ease and the simultaneous analysis of all the anions present in the aqueous sample. Other methods for the specific detection of nitrate rely on its conversion to nitrite followed by nitrite-specific tests. The reduction of nitrate to nitrite is effected by a copper-cadmium material. The sample is introduced in a flow injection analyzer, and the resulting nitrite-containing effluent is then combined with a reagent for colorimetric or electrochemical detection. The most popular of these assays is the Griess test, whereby nitrite is converted to a deeply colored azo dye suited for UV-vis spectroscopic analysis. The method exploits the reactivity of nitrous acid derived from acidification of nitrite. Nitrous acid selectively reacts with aromatic amines to give diazonium salts, which in turn couple with a second reagent to give the azo dye. The detection limit is 0.02 to 2 μM.[3] Such methods have been highly adapted to biological samples.[4]

Occurrence and production[edit]

Nitrate salts are found naturally on earth in arid environments as large deposits, particularly of nitratine, a major source of sodium nitrate.

Nitrates are produced by a number of species of nitrifying bacteria in the natural environment using ammonia or urea as a source of nitrogen and source of free energy. Nitrate compounds for gunpowder were historically produced, in the absence of mineral nitrate sources, by means of various fermentation processes using urine and dung.

Lightning strikes in earth's nitrogen- and oxygen-rich atmosphere produce a mixture of oxides of nitrogen, which form nitrous ions and nitrate ions, which are washed from the atmosphere by rain or in occult deposition.

Nitrates are produced industrially from nitric acid.[1]

Uses[edit]

Agriculture[edit]

Nitrates are used as fertilizers in agriculture because of their high solubility and biodegradability. The main nitrate fertilizers are ammonium, sodium, potassium, calcium, and magnesium salts. Several billion kilograms are produced annually for this purpose.[1]

Firearms[edit]

Nitrates are used as oxidizing agents, most notably in explosives, where the rapid oxidation of carbon compounds liberates large volumes of gases (see gunpowder for an example).

Industrial[edit]

Sodium nitrate is used to remove air bubbles from molten glass and some ceramics. Mixtures of the molten salt are used to harden some metals.[1]

Photographic film[edit]

Nitrate was also used as a film stock through nitrocellulose. Due to its high combustibility, the film making studios swapped to cellulose acetate safety film in 1950.

Medicinal and pharmaceutical use[edit]

In the medical field, nitrate-derived organic esters, such as glyceryl trinitrate, isosorbide dinitrate, and isosorbide mononitrate, are used in the prophylaxis and management of acute coronary syndrome, myocardial infarction, acute pulmonary oedema.[5] This class of drug, to which amyl nitrite also belongs, is known as nitrovasodilators.

Toxicity and safety[edit]

The two areas of concerns about the toxicity of nitrate are the following:

Methemoglobinemia[edit]

Main articles: Blue baby syndrome and Methemoglobinemia

One of the most common cause of methemoglobinemia in infants is due to the ingestion of nitrates and nitrites through well water or foods.

In fact, nitrates (NO−3), often present at too high concentration in drinkwater, are only the precursor chemical species of nitrites (NO−2), the real culprits of methemoglobinemia. Nitrites produced by the microbial reduction of nitrate (directly in the drinkwater, or after ingestion by the infant, in his digestive system) are more powerful oxidizers than nitrates and are the chemical agent really responsible for the oxidation of Fe2+ into Fe3+ in the tetrapyrrole heme of hemoglobin. Indeed, nitrate anions are too weak oxidizers in aqueous solution to be able to directly, or at least sufficiently rapidly, oxidize Fe2+ into Fe3+, because of kinetics limitations.

Infants younger than 4 months are at greater risk given that they drink more water per body weight, they have a lower NADH-cytochrome b5 reductase activity, and they have a higher level of fetal hemoglobin which converts more easily to methemoglobin. Additionally, infants are at an increased risk after an episode of gastroenteritis due to the production of nitrites by bacteria.[8]

However, other causes than nitrates can also affect infants and pregnant women.[9][10] Indeed, the blue baby syndrome can also be caused by a number of other factors such as the cyanotic heart disease, a congenital heart defect resulting in low levels of oxygen in the blood,[11] or by gastric upset, such as diarrheal infection, protein intolerance, heavy metal toxicity, etc.[12]

Drinking water standards[edit]

Through the Safe Drinking Water Act, the United States Environmental Protection Agency has set a maximum contaminant level of 10 mg/L or 10 ppm of nitrate in drinking water.[13]

An acceptable daily intake (ADI) for nitrate ions was established in the range of 0–3.7 mg (kg body weight)−1 day−1 by the Joint FAO/WHO Expert Committee on Food Additives (JEFCA).[14]

Aquatic toxicity[edit]

Sea surface nitrate from the World Ocean Atlas

In freshwater or estuarine systems close to land, nitrate can reach concentrations that are lethal to fish. While nitrate is much less toxic than ammonia,[15] levels over 30 ppm of nitrate can inhibit growth, impair the immune system and cause stress in some aquatic species.[16] Nitrate toxicity remains a subject of debate.[17]

In most cases of excess nitrate concentrations in aquatic systems, the primary sources are wastewater discharges, as well as surface runoff from agricultural or landscaped areas that have received excess nitrate fertilizer. The resulting eutrophication and algae blooms result in anoxia and dead zones. As a consequence, as nitrate forms a component of total dissolved solids, they are widely used as an indicator of water quality.

Dietary nitrate[edit]

A source of nitrate in the human diets arises from the consumption of leafy green foods, such as spinach and arugula. NO−
3 can be present in beetroot juice. Drinking water represents also a primary nitrate intake source.[18]

Nitrate ingestion rapidly increases the plasma nitrate concentration by a factor of 2 to 3, and this elevated nitrate concentration can be maintained for more than 2 weeks. Increased plasma nitrate enhances the production of nitric oxide, NO. Nitric oxide is a physiological signaling molecule which intervenes in, among other things, regulation of muscle blood flow and mitochondrial respiration.[19]

Cured meats[edit]

Nitrite consumption is primarily determined by the amount of processed meats eaten, and the concentration of nitrates in these meats. Although nitrites are the nitrogen compound chiefly used in meat curing, nitrates are used as well. Nitrites lead to the formation of carcinogenic nitrosamines.[20] The production of nitrosamines may be inhibited by the use of the antioxidants vitamin C and the alpha-tocopherol form of vitamin E during curing.[21]

Many meat processors claim their meats (e.g. bacon) is "uncured" – which is a marketing claim with no factual basis: there is no such thing as "uncured" bacon (as that would be, essentially, raw sliced pork belly).[22][better source needed] "Uncured" meat is in fact actually cured with nitrites with virtually no distinction in process – the only difference being the USDA labeling requirement between nitrite of vegetable origin (such as from celery) vs. "synthetic" sodium nitrite. An analogy would be purified "sea salt" vs. sodium chloride – both being the exact same chemical with the only essential difference being the origin.

Anti-hypertensive diets, such as the DASH diet, typically contain high levels of nitrates, which are first reduced to nitrite in the saliva, as detected in saliva testing, prior to forming nitric oxide.[18]

Domestic animal feed[edit]

Symptoms of nitrate poisoning in domestic animals include increased heart rate and respiration; in advanced cases blood and tissue may turn a blue or brown color. Feed can be tested for nitrate; treatment consists of supplementing or substituting existing supplies with lower nitrate material. Safe levels of nitrate for various types of livestock are as follows:[23]

Category %NO3 %NO3–N %KNO3 Effects
1 < 0.5 < 0.12 < 0.81 Generally safe for beef cattle and sheep
2 0.5–1.0 0.12–0.23 0.81–1.63 Caution: some subclinical symptoms may appear in pregnant horses, sheep and beef cattle
3 1.0 0.23 1.63 High nitrate problems: death losses and abortions can occur in beef cattle and sheep
4 < 1.23 < 0.28 < 2.00 Maximum safe level for horses. Do not feed high nitrate forages to pregnant mares

The values above are on a dry (moisture-free) basis.

Salts and covalent derivatives[edit]

Nitrate formation with elements of the periodic table:

See also[edit]

References[edit]

  1. ^ a b c d Laue W, Thiemann M, Scheibler E, Wiegand KW (2006). "Nitrates and Nitrites". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a17_265. ISBN 978-3527306732.
  2. ^ Lubos, Edith (2008). "Role of oxidative stress and nitric oxide in atherothrombosis". Frontiers in Bioscience. 13. IMR Press: 5323. doi:10.2741/3084. ISSN 1093-9946.
  3. ^ Moorcroft, M.; Davis, J.; Compton, R. G. (2001). "Detection and determination of nitrate and nitrite: A review". Talanta. 54 (5): 785–803. doi:10.1016/S0039-9140(01)00323-X. PMID 18968301.
  4. ^ Ellis, Graham; Adatia, Ian; Yazdanpanah, Mehrdad; Makela, Sinikka K. (1998). "Nitrite and Nitrate Analyses: A Clinical Biochemistry Perspective". Clinical Biochemistry. 31 (4): 195–220. doi:10.1016/S0009-9120(98)00015-0. PMID 9646943.
  5. ^ Soman, Biji; Vijayaraghavan, Govindan. "The role of organic nitrates in the optimal medical management of angina". www.escardio.org. Retrieved 2023-10-30.
  6. ^ Powlson, David S.; Addiscott, Tom M.; Benjamin, Nigel; Cassman, Ken G.; De Kok, Theo M.; Van Grinsven, Hans; l'Hirondel, Jean-Louis; Avery, Alex A.; Van Kessel, Chris (2008). "When Does Nitrate Become a Risk for Humans?". Journal of Environmental Quality. 37 (2): 291–5. doi:10.2134/jeq2007.0177. PMID 18268290. S2CID 14097832.
  7. ^ "Nitrate and Nitrite Poisoning: Introduction". The Merck Veterinary Manual. Retrieved 2008-12-27.
  8. ^ Smith-Whitley, Kwiatkowski. Nelson Textbook of Pediatrics. Elsevier Inc. pp. 2540–2558.
  9. ^ Addiscott, T.M.; Benjamin, N. (2006). "Nitrate and human health". Soil Use and Management. 20 (2): 98–104. doi:10.1111/j.1475-2743.2004.tb00344.x. S2CID 96297102.
  10. ^ A. A. Avery: Infant Methemoglobinemia - Reexamining the Role of Drinking Water Nitrates, Environmental Health Perspectives, Volume 107, Number 7, July 1999.
  11. ^ MedlinePlus Encyclopedia: Cyanotic heart disease
  12. ^ Manassaram DM, Backer LC, Messing R, Fleming LE, Luke B, Monteilh CP (October 2010). "Nitrates in drinking water and methemoglobin levels in pregnancy: a longitudinal study". Environmental Health. 9 (1): 60. doi:10.1186/1476-069x-9-60. PMC 2967503. PMID 20946657.
  13. ^ "4. What are EPA's drinking water regulations for nitrate?". Ground Water & Drinking Water. Retrieved 2018-11-13.
  14. ^ Bagheri, H.; Hajian, A.; Rezaei, M.; Shirzadmehr, A. (2017). "Composite of Cu metal nanoparticles-multiwall carbon nanotubes-reduced graphene oxide as a novel and high performance platform of the electrochemical sensor for simultaneous determination of nitrite and nitrate". Journal of Hazardous Materials. 324 (Pt B): 762–772. doi:10.1016/j.jhazmat.2016.11.055. PMID 27894754.
  15. ^ Romano N, Zeng C (September 2007). "Acute toxicity of sodium nitrate, potassium nitrate, and potassium chloride and their effects on the hemolymph composition and gill structure of early juvenile blue swimmer crabs(Portunus pelagicus Linnaeus, 1758) (Decapoda, Brachyura, Portunidae)". Environmental Toxicology and Chemistry. 26 (9): 1955–62. doi:10.1897/07-144r.1. PMID 17705664. S2CID 19854591.
  16. ^ Sharpe, Shirlie. "Nitrates in the Aquarium". About.com. Archived from the original on July 24, 2011. Retrieved October 30, 2013.
  17. ^ Romano N, Zeng C (December 2007). "Effects of potassium on nitrate mediated alterations of osmoregulation in marine crabs". Aquatic Toxicology. 85 (3): 202–8. doi:10.1016/j.aquatox.2007.09.004. PMID 17942166.
  18. ^ a b Hord NG, Tang Y, Bryan NS (July 2009). "Food sources of nitrates and nitrites: the physiologic context for potential health benefits". The American Journal of Clinical Nutrition. 90 (1): 1–10. doi:10.3945/ajcn.2008.27131. PMID 19439460.
  19. ^ Maughan, Ronald J (2013). Food, Nutrition and Sports Performance III. New York: Taylor & Francis. p. 63. ISBN 978-0-415-62792-4.
  20. ^ Bingham SA, Hughes R, Cross AJ (November 2002). "Effect of white versus red meat on endogenous N-nitrosation in the human colon and further evidence of a dose response". The Journal of Nutrition. 132 (11 Suppl): 3522S–3525S. doi:10.1093/jn/132.11.3522S. PMID 12421881.
  21. ^ Parthasarathy DK, Bryan NS (November 2012). "Sodium nitrite: the "cure" for nitric oxide insufficiency". Meat Science. 92 (3): 274–9. doi:10.1016/j.meatsci.2012.03.001. PMID 22464105.
  22. ^ "Is There a Difference Between Cured and Uncured Bacon?". 9 December 2022.
  23. ^ "Nitrate Risk in Forage Crops - Frequently Asked Questions". Agriculture and Rural Development. Government of Alberta. Retrieved October 30, 2013.

External links[edit]

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