Editing Nitrate

{{Short description|Polyatomic ion (NO₃, charge –1) found in explosives and fertilisers}}
{{About||the functional group  –{{chem|ONO|2}}|Nitrate ester|that functional group in medicine|Nitrovasodilator}}
{{Distinguish|text={{chem|NO|2|−}}, [[nitrite]]}}
{{About|the ion|the radical|nitrate radical}}
{{Otheruses}}
{{cs1 config|name-list-style=vanc|display-authors=6}}
{{Chembox
| ImageFile1 = Nitrate-3D-balls.png
| ImageAlt1 = [[Ball-and-stick model]] of the nitrate ion
| ImageFileR1 = <!-- filename lost in history?  (then L1/R1 together))-->
| ImageAltR1 = [[Space-filling model]] of the nitrate ion
| SystematicName = Nitrate
| IUPACName =
| OtherNames =
| Section1 = {{Chembox Identifiers
| CASNo = 14797-55-8
| CASNo_Ref = {{cascite|correct|CAS}}
| UNII = T93E9Y2844
| UNII_Ref = {{fdacite|correct|FDA}}
| PubChem = 943
| SMILES = [N+](=O)([O-])[O-] 
| ChemSpiderID = 918
| InChI = 1/NO3/c2-1(3)4/q-1
| InChIKey = NHNBFGGVMKEFGY-UHFFFAOYAI
| StdInChI = 1S/NO3/c2-1(3)4/q-1
| StdInChIKey = NHNBFGGVMKEFGY-UHFFFAOYSA-N
| RTECS = 
| MeSHName = 
| ChEBI = 17632}}
| Section2 = {{Chembox Properties
| Formula = {{chem|N|O|3|-}}
| N=1|O=3
| Appearance =
| Solubility = 
| ConjugateAcid = [[Nitric acid]]}}
| Section3 = {{Chembox Hazards
| MainHazards =
| FlashPt =
| AutoignitionPt = }}
}}
'''Nitrate''' is a [[polyatomic ion]] with the [[chemical formula]] {{chem|NO|3|−}}. [[salt (chemistry)|Salts]] containing this [[ion]] are called '''nitrates'''. Nitrates are common components of fertilizers and explosives.<ref name="Ullmann">{{Ullmann|first1=Wolfgang|last1=Laue|first2=Michael|last2=Thiemann|first3=Erich|last3=Scheibler|first4=Karl Wilhelm|last4=Wiegand|title=Nitrates and Nitrites|year=2006|doi=10.1002/14356007.a17_265|isbn=978-3527306732}}</ref> Almost all inorganic nitrates are [[solubility|soluble]] in [[water]]. An example of an insoluble nitrate is [[bismuth oxynitrate]].

== Chemical structure ==
{{Citation needed section|date=April 2024}}
[[File:Nitrate-ion-with-partial-charges-2D.png|thumb|150px|The nitrate ion with the partial charges shown]]
The nitrate [[anion]] is the [[conjugate acid|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|date=April 2024}} This charge results from a combination formal charge in which each of the three oxygens carries a −{{frac|2|3}} charge,{{Citation needed|date=April 2024}} whereas the nitrogen carries a +1 charge, all these adding up to formal charge of the polyatomic nitrate ion.{{Citation needed|date=April 2024}} This arrangement is commonly used as an example of [[Resonance (chemistry)|resonance]]. Like the [[isoelectronic]] [[carbonate]] ion, the nitrate ion can be represented by three resonance structures:

<div style="text-align: center">[[Image:Nitrate-ion-resonance-2D.png|400px|Canonical resonance structures for the nitrate ion]]</div>

== Chemical and biochemical properties ==
In the {{Chem2|NO3-}} 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 [[Oxidizing agent|oxidizer]] as evidenced by its [[explosive]] behaviour at high temperature when it is [[Detonation|detonated]] in [[ammonium nitrate]] ({{Chem2|NH4NO3}}), or [[Gunpowder|black powder]], ignited by the [[shock wave]] of a [[Explosive#Primary|primary explosive]]. In contrast to [[red fuming nitric acid]] ({{Chem2|HNO3/N2O4}}), or concentrated [[nitric acid]] ({{Chem2|HNO3}}), nitrate in [[aqueous solution]] at neutral or high [[pH]] is only a weak [[oxidizing agent]] in [[redox]] reactions in which the [[reductant]] does not produce hydrogen ions (such as mercury going to [[calomel]]). However it is still a strong oxidizer when the reductant does produce hydrogen ions, such as in the oxidation of hydrogen itself. Nitrate is stable in the absence of [[microorganism]]s or reductants such as organic matter. In fact, nitrogen gas is thermodynamically stable in the presence of 1 [[atmosphere (unit)|atm]] of oxygen only in very acidic conditions, and otherwise should combine with the oxygen to form nitrate. This is shown by subtracting the two oxidation reactions:<ref>{{cite book |last1=[[Marcel Pourbaix]] and N. [[Zubov|de Zoubov]] |editor1-last=Marcel Pourbaix |title=Atlas of Electrochemical Equilibria in Aqueous Solutions |date=1974 |pages=49, 497, 500 |url=http://sunlight.caltech.edu/aic/pourbaix.pdf |chapter=Nitrogen}}</ref>

: {{chem2|N2 + 6 H2O → 2 NO3- + 12 H+ + 10 e-}} <math>\qquad E_0 = 1.246-0.0709 \text{ pH } + \frac {0.0591} {10} \log\frac{(NO_3^-)^2}{P_{N_2}}</math>

: {{chem2|2 H2O → O2 + 4 H+ + 4 e-}} <math>\qquad\qquad\qquad E_0 = 1.228-0.0591 \text{ pH } + \frac {0.0591} 4 \log{P_{O_2}}</math>

giving:

: {{chem2|2 N2 + 5 O2 + 2 H2O → 4 NO3- + 4 H+}} <math>\qquad0=0.018-0.0118 \text{ pH } + \frac {0.0591} {10} \log\frac{(NO_3^-)^2}{P_{N_2}}-\frac {0.0591} 4 \log{P_{O_2}}</math>

Dividing by 0.0118 and rearranging gives the equilibrium relation:

: <math>\log\frac{(NO_3^-)}{P_{N_2}^{1/2}P_{O_2}^{5/4}}=\text{ pH }-1.5</math>

However, in reality, nitrogen, oxygen, and water do not combine directly to form nitrate. Rather, a reductant such as hydrogen reacts with nitrogen to produce "fixed nitrogen" such as [[ammonia]], which is then oxidized, eventually becoming nitrate. Nitrate does not accumulate to high levels in nature because it reacts with reductants in the process called [[denitrification]] (see [[Nitrogen cycle]]).

Nitrate is used as a powerful terminal [[electron acceptor]] by [[denitrifying bacteria]] to deliver the energy they need to thrive. Under [[Hypoxia (environmental)|anaerobic conditions]], nitrate is the strongest electron acceptor used by [[prokaryote]] [[microorganism]]s ([[bacteria]] and [[archaea]]) to respirate. The [[redox]] couple {{Chem2|NO3-}}/{{Chem2|N2}} is at the top of the [[Redox gradient|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), {{Chem2|SO4(2-)}}/{{Chem2|HS-}}, {{CO2}}/{{CH4}}. In natural waters, inevitably contaminated by microorganisms, nitrate is a quite unstable and labile dissolved chemical species because it is [[metabolism|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 ({{Chem2|NO3-}}) is [[Catalysis|catalytically]] [[Redox|reduced]] into nitrite ({{Chem2|NO2-}}) by the [[Enzyme catalysis|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 [[electron]]s and its [[Redox|reduction]] is less [[Chemical kinetics|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 [[Cell signaling|signaling molecule]] and intervenes in the [[vasodilation]] process. Still, it can also produce [[Radical (chemistry)|free radicals]] in [[Tissue (biology)|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 [[Atherosclerosis|atherogenesis]].<ref name="Lubos2008">{{cite journal | vauthors = Lubos E, Handy DE, Loscalzo J | title = Role of oxidative stress and nitric oxide in atherothrombosis | journal = Frontiers in Bioscience | volume = 13 | issue = 13 | pages = 5323–5344 | date = May 2008 | pmid = 18508590 | pmc = 2617738 | doi = 10.2741/3084 | publisher = IMR Press }}</ref>

== Detection in chemical analysis ==
The nitrate [[anion]] is commonly analysed in water by [[ion chromatography]] (IC) along with other anions also present in the solution. The main advantage of IC is its ease and the simultaneous [[Analytical chemistry|analysis]] of all the anions present in the aqueous sample. Since the emergence of IC instruments in the 1980s, this separation technique, coupled with many detectors, has become commonplace in the chemical analysis laboratory and is the preferred and most widely used method for nitrate and nitrite analyses.<ref name="Michalski2006">{{cite journal | last1=Michalski | first1=Rajmund | date=2006 | title=Ion chromatography as a reference method for determination of inorganic ions in water and wastewater | journal=Critical Reviews in Analytical Chemistry | volume=36 | issue=2 | issn=1040-8347 | doi=10.1080/10408340600713678 | pages=107–127}}</ref>  

Previously, nitrate determination relied on [[Spectrophotometry|spectrophotometric]] and [[Colorimetric analysis|colorimetric]] measurements after a specific reagent is added to the [[Solution (chemistry)|solution]] to reveal a characteristic color (often red because it absorbs visible light in the blue). Because of interferences with the brown color of [[Dissolved organic carbon|dissolved organic matter]] (DOM: [[Humic substance|humic]] and [[Humic substance|fulvic acids]]) often present in [[soil]] pore water, artefacts can easily affect the [[absorbance]] values. In case of weak interference, a blank measurement with only a naturally brown-colored water sample can be sufficient to subtract the undesired background from the measured sample absorbance. If the DOM brown color is too intense, the water samples must be pretreated, and inorganic nitrogen species must be separated before measurement. Meanwhile, for clear water samples, colorimetric instruments retain the advantage of being less expensive and sometimes portable, making them an affordable option for fast routine controls or field measurements. 

Colorimetric methods for the specific detection of nitrate ({{chem2|NO3-}}) often rely on its conversion to [[nitrite]] ({{chem2|NO2-}}) followed by nitrite-specific tests. The [[Redox|reduction]] of nitrate to nitrite can be effected by a [[copper]]-[[cadmium]] [[alloy]], metallic [[zinc]],<ref name="daAscencao2024">{{cite journal | last1=da Ascenção | first1=Wellington Diego | last2=Augusto | first2=Caroline Cristine | last3=de Melo | first3=Vitor Hugo Soares | last4=Batista | first4=Bruno Lemos | date=2024-05-23 | title=A simple, ecofriendly, and fast method for nitrate quantification in bottled water using visible spectrophotometry | journal=Toxics | volume=12 | issue=6 | issn=2305-6304 | pmid=38922063 | pmc=11209534 | doi=10.3390/toxics12060383 | doi-access=free | page=383}}</ref> or [[hydrazine]]. The most popular of these assays is the [[Griess test]], whereby nitrite is converted to a deeply red colored [[azo dye]] suited for [[Ultraviolet–visible spectroscopy|UV–vis spectrophotometry]] analysis. The method exploits the reactivity of [[nitrous acid]] ({{chem2|HNO2}}) derived from the acidification of nitrite. Nitrous acid selectively reacts with aromatic amines to give [[Diazonium compound|diazonium salts]], which in turn couple with a second reagent to give the [[azo dye]]. The [[detection limit]] is 0.02 to 2 μM.<ref name="Moorcroft2001">{{cite journal | vauthors = Moorcroft MJ, Davis J, Compton RG | title = Detection and determination of nitrate and nitrite: a review | journal = Talanta | volume = 54 | issue = 5 | pages = 785–803 | date = June 2001 | pmid = 18968301 | doi = 10.1016/S0039-9140(01)00323-X }}</ref> Such methods have been highly adapted to biological samples<ref name="Ellis1998">{{cite journal | vauthors = Ellis G, Adatia I, Yazdanpanah M, Makela SK | title = Nitrite and nitrate analyses: a clinical biochemistry perspective | journal = Clinical Biochemistry | volume = 31 | issue = 4 | pages = 195–220 | date = June 1998 | pmid = 9646943 | doi = 10.1016/S0009-9120(98)00015-0}}</ref> and soil samples.<ref name="Bremmer1965">{{Citation |last1=Bremner |first1=J. M. |title=Inorganic forms of nitrogen |date=1965 |work=Methods of Soil Analysis |pages=1179–1237 |url=https://acsess.onlinelibrary.wiley.com/doi/abs/10.2134/agronmonogr9.2.c33 |access-date=2025-03-03 |publisher=John Wiley & Sons, Ltd |language=en |doi=10.2134/agronmonogr9.2.c33 |isbn=978-0-89118-204-7|url-access=subscription }}</ref><ref name="Guiot1975">{{Cite journal |last=Guiot |first=J. |date=1975 |title=Estimation of soil nitrogen reserves by determination of mineral nitrogen |url=https://www.cabidigitallibrary.org/doi/full/10.5555/19751922999 |journal=Revue de l'Agriculture (Bruxelles) |volume=28 |issue=5 |pages=1117–1132 |via=CABI Databases}}</ref>

In the [[dimethylphenol]] method, 1 mL of concentrated [[sulfuric acid]] ({{chem2|H2SO4}}) is added to 200 μL of the solution being tested for nitrate. Under strongly acidic conditions, nitrate ions react with 2,6-dimethylphenol, forming a yellow compound, [[Nitrophenol|4-nitro-2,6-dimethylphenol]]. This occurs through [[electrophilic aromatic substitution]] where the intermediate [[nitronium]] ({{chem2|+NO2}}) ions attack the [[aromatic ring]] of dimethylphenol. The resulting product ([[Nitrophenol|ortho- or para-nitro-dimethylphenol]]) is analyzed using [[UV-visible spectroscopy|UV-vis spectrophotometry]] at 345 nm according to the [[Lambert-Beer law]].<ref name="dimethylphenol">{{Cite web| author=US-EPA| title=Approved method for water and wastewater analysis, 40 CFR part 136; and drinking water, 40 CFR part 141.23.| url=https://downloads.regulations.gov/EPA-HQ-OW-2011-0413-0018/content.pdf| access-date=2025-04-14}}</ref><ref name="HachTNTplus">{{Cite web| author=Hach Company| title=TNTplus 835/836 nitrate method 10206. Spectrophotometric measurement of nitrate in water and wastewater| url=https://downloads.regulations.gov/EPA-HQ-OW-2011-0413-0018/content.pdf| year=2025| access-date=2025-04-14}}.</ref>

Another [[Colorimetric analysis|colorimetric method]] based on the [[chromotropic acid]] (dihydroxynaphthalene-disulfonic acid) was also developed by West and Lyles in 1960 for the direct [[Spectrophotometry|spectrophotometric]] determination of nitrate [[Ion|anions]].<ref name="WestLyles1960">{{Cite journal |last1=West |first1=Philip W. |last2=Lyles |first2=George L. |date=1960 |title=A new method for the determination of nitrates |url=https://linkinghub.elsevier.com/retrieve/pii/S0003267060800578 |journal=Analytica Chimica Acta |volume=23 |pages=227–232 |doi=10.1016/S0003-2670(60)80057-8 |issn=0003-2670|url-access=subscription }}</ref>

If [[formic acid]] is added to a mixture of [[brucine]]  (an [[alkaloid]] related to [[strychnine]]) and [[potassium nitrate]] ({{chem2|KNO3}}), its color instantly turns red. This reaction has been used for the direct [[Colorimetric analysis|colorimetric detection]] of nitrates.<ref name="Baker1967">{{Cite web |last=Baker |first=Aaron Sidney |date=1967-05-01 |title=Colorimetric determination of nitrate in soil and plant extracts with brucine |url=https://pubs.acs.org/doi/pdf/10.1021/jf60153a004 |access-date=2025-03-01 |website=ACS Publications |language=EN |doi=10.1021/jf60153a004}}</ref>

For direct online chemical analysis using a flow-through system, the water sample is introduced by a [[peristaltic pump]] in a [[Flow injection analysis|flow injection analyzer]], and the nitrate or resulting nitrite-containing effluent is then combined with a reagent for its colorimetric detection.

== Occurrence and production ==
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 [[industrial fermentation|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 acid|nitrous]] ions and nitrate ions, which are washed from the atmosphere by rain or in [[Precipitation#occult deposition|occult deposition]].

Nitrates are produced industrially from [[nitric acid]].<ref name=Ullmann/>

== Uses ==

=== Agriculture ===
Nitrate is a [[chemical compound]] that serves as a primary form of nitrogen for many plants. This essential nutrient is used by plants to synthesize proteins, nucleic acids, and other vital organic molecules.<ref>{{cite journal | vauthors = Zhang GB, Meng S, Gong JM | title = The Expected and Unexpected Roles of Nitrate Transporters in Plant Abiotic Stress Resistance and Their Regulation | journal = International Journal of Molecular Sciences | volume = 19 | issue = 11 | pages = 3535 | date = November 2018 | pmid = 30423982 | pmc = 6274899 | doi = 10.3390/ijms19113535 | doi-access = free }}</ref> The transformation of atmospheric nitrogen into nitrate is facilitated by certain bacteria and lightning in the nitrogen cycle, which exemplifies nature's ability to convert a relatively inert molecule into a form that is crucial for biological productivity.<ref>{{Cite journal | vauthors = Chuang HP |date=2018-11-26 |title=Insight on transformation pathways of nitrogen species and functional genes expression by targeted players involved in nitrogen cycle. |journal=Impact |language=en |volume=2018 |issue=8 |pages=58–59 |doi=10.21820/23987073.2018.8.58 |issn=2398-7073}}</ref>

Nitrates are used as [[fertilizer]]s in [[agriculture]] because of their high solubility and biodegradability. The main nitrate fertilizers are [[Ammonium nitrate|ammonium]], [[Sodium nitrate|sodium]], [[Potassium nitrate|potassium]], [[Calcium nitrate|calcium]], and [[Magnesium nitrate|magnesium]] salts. Several billion kilograms are produced annually for this purpose.<ref name="Ullmann" /> The significance of nitrate extends beyond its role as a nutrient since it acts as a signaling molecule in plants, regulating processes such as root growth, flowering, and leaf development.<ref>{{cite journal | vauthors = Liu B, Wu J, Yang S, Schiefelbein J, Gan Y | title = Nitrate regulation of lateral root and root hair development in plants | journal = Journal of Experimental Botany | volume = 71 | issue = 15 | pages = 4405–4414 | date = July 2020 | pmid = 31796961 | pmc = 7382377 | doi = 10.1093/jxb/erz536 | veditors = Xu G }}</ref>

While nitrate is beneficial for agriculture since it enhances soil fertility and crop yields, its excessive use can lead to nutrient runoff, water pollution, and the proliferation of aquatic dead zones.<ref>{{cite book | vauthors = Bashir U, Lone FA, Bhat RA, Mir SA, Dar ZA, Dar SA | chapter = Concerns and Threats of Contamination on Aquatic Ecosystems |date=2020 | title = Bioremediation and Biotechnology |pages=1–26 | veditors = Hakeem KR, Bhat RA, Qadri H |place=Cham |publisher=Springer International Publishing |language=en |doi=10.1007/978-3-030-35691-0_1 |isbn=978-3-030-35690-3 }}</ref> Therefore, sustainable agricultural practices that balance productivity with environmental stewardship are necessary. Nitrate's importance in ecosystems is evident since it supports the growth and development of plants, contributing to biodiversity and ecological balance.<ref>{{cite journal | vauthors = Kirchmann H, Johnston AE, Bergström LF | title = Possibilities for reducing nitrate leaching from agricultural land | journal = Ambio | volume = 31 | issue = 5 | pages = 404–408 | date = August 2002 | pmid = 12374048 | doi = 10.1579/0044-7447-31.5.404 | bibcode = 2002Ambio..31..404K }}</ref>

=== Firearms ===
Nitrates are used as [[oxidizing agent]]s, most notably in [[explosive]]s, where the rapid [[Redox|oxidation]] of carbon compounds liberates large volumes of gases (see [[gunpowder]] as an example).

=== Industrial === 
Sodium nitrate is used to remove air bubbles from molten [[glass]] and some [[ceramic]]s. Mixtures of [[molten salt]]s are used to harden the surface of some metals.<ref name=Ullmann/>

=== Photographic film ===
Nitrate was also used as a [[Film base#Nitrate|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 ===
In the medical field, nitrate-derived organic [[ester]]s, such as [[Nitroglycerin (medication)|glyceryl trinitrate]], [[isosorbide dinitrate]], and [[isosorbide mononitrate]], are used in the prophylaxis and management of [[acute coronary syndrome]], [[myocardial infarction]], [[Pulmonary edema|acute pulmonary oedema]].<ref>{{Cite journal | vauthors = Soman B, Vijayaraghavan G |title=The role of organic nitrates in the optimal medical management of angina | journal =  e-Journal of Cardiology Practice | date = April 2017 |volume=15 | issue = 2 |url=https://www.escardio.org/Journals/E-Journal-of-Cardiology-Practice/Volume-15/The-role-of-organic-nitrates-in-the-optimal-medical-management-of-angina |access-date=2023-10-30 }}</ref> This class of drug, to which [[amyl nitrite]] also belongs, is known as [[nitrovasodilator]]s.

== Toxicity and safety ==
The two areas of concerns about the toxicity of nitrate are the following: 
* nitrate reduced by the microbial activity of [[nitrate reducing bacteria]] is the precursor of [[nitrite]] in water and in the [[lower gastrointestinal tract]]. Nitrite is a precursor to [[carcinogenesis|carcinogenic]] [[nitrosamine]]s, and;
* via the formation of nitrite, nitrate is implicated in [[methemoglobinemia]], a disorder of [[hemoglobin]] in [[red blood cell]]s susceptible to especially affect infants and toddlers.<ref>{{cite journal | vauthors = Powlson DS, Addiscott TM, Benjamin N, Cassman KG, de Kok TM, van Grinsven H, L'Hirondel JL, Avery AA, van Kessel C | title = When does nitrate become a risk for humans? | journal = Journal of Environmental Quality | volume = 37 | issue = 2 | pages = 291–295 | year = 2008 | pmid = 18268290 | doi = 10.2134/jeq2007.0177 | s2cid = 14097832 | bibcode = 2008JEnvQ..37..291P | url = https://digitalcommons.unl.edu/agronomyfacpub/102 | url-access = subscription }}
</ref><ref>{{cite web|work=The Merck Veterinary Manual|title=Nitrate and Nitrite Poisoning: Introduction|url=http://www.merckvetmanual.com/mvm/index.jsp?cfile=htm/bc/212300.htm|access-date=2008-12-27}}</ref>

=== Methemoglobinemia ===
{{Main Article|Blue baby syndrome|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 ({{Chem2|NO3-}}), often present at too high [[concentration]] in drinkwater, are only the precursor chemical species of [[nitrite]]s ({{Chem2|NO2-}}), the real culprits of methemoglobinemia. Nitrites produced by the [[Denitrification|microbial reduction of nitrate]] (directly in the drinkwater, or after ingestion by the infant, in his digestive system) are more powerful [[Oxidizing agent|oxidizers]] than nitrates and are the chemical agent really responsible for the [[Redox|oxidation]] of Fe<sup>2+</sup> into Fe<sup>3+</sup> 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 Fe<sup>2+</sup> into Fe<sup>3+</sup>, because of [[Chemical kinetics|kinetics]] limitations.

Infants younger than 4 months are at greater risk given that they drink more water per body weight, they have a lower [[Nicotinamide adenine dinucleotide|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 [[nitrite]]s by [[bacteria]].<ref>{{Cite book | vauthors = Smith-Whitley K, Kwiatkowski JL | chapter = Chapter 489: Hemoglobinopathies | veditors = Kliegman RM |title=Nelson Textbook of Pediatrics | edition = 21st | date = 2000 |publisher=Elsevier Inc.|pages=2540–2558 | isbn = 978-0-323-52950-1 }}</ref> 

However, other causes than nitrates can also affect infants and pregnant women.<ref>{{Cite journal | doi=10.1111/j.1475-2743.2004.tb00344.x|title = Nitrate and human health| journal=Soil Use and Management| volume=20| issue=2| pages=98–104|year = 2006| vauthors = Addiscott TM, Benjamin N | s2cid=96297102 }}</ref><ref name="pmid10379005">{{cite journal | vauthors = Avery AA | title = Infantile methemoglobinemia: reexamining the role of drinking water nitrates | journal = Environmental Health Perspectives | volume = 107 | issue = 7 | pages = 583–6 | date = July 1999 | pmid = 10379005 | pmc = 1566680 | doi = 10.1289/ehp.99107583 | bibcode = 1999EnvHP.107..583A }}</ref> 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,<ref name="MedlinePlusEncyclopedia Cyanotic heart disease">{{MedlinePlusEncyclopedia|001104|Cyanotic heart disease}}</ref> or by gastric upset, such as diarrheal infection, protein intolerance, heavy metal toxicity, etc.<ref>{{cite journal | vauthors = Manassaram DM, Backer LC, Messing R, Fleming LE, Luke B, Monteilh CP | title = Nitrates in drinking water and methemoglobin levels in pregnancy: a longitudinal study | language = En | journal = Environmental Health | volume = 9 | issue = 1 | pages = 60 | date = October 2010 | pmid = 20946657 | pmc = 2967503 | doi = 10.1186/1476-069x-9-60 | doi-access = free | bibcode = 2010EnvHe...9...60M }}</ref>

=== Drinking water standards ===
Through the [[Safe Drinking Water Act]], the [[United States Environmental Protection Agency]] has set a maximum contaminant level of 10&nbsp;mg/L or 10 ppm of nitrate in drinking water.<ref>{{Cite web|url=https://safewater.zendesk.com/hc/en-us/articles/211401718-4-What-are-EPA-s-drinking-water-regulations-for-nitrate-|title=4. What are EPA's drinking water regulations for nitrate?|website=Ground Water & Drinking Water|date=20 September 2016 |language=en-US|access-date=2018-11-13}}</ref>

An acceptable daily intake (ADI) for nitrate ions was established in the range of 0–3.7&nbsp;mg (kg body weight)<sup>−1</sup> day<sup>−1</sup> by the Joint FAO/WHO Expert Committee on Food Additives (JEFCA).<ref>{{cite journal | vauthors = Bagheri H, Hajian A, Rezaei M, Shirzadmehr A | title = 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 = Journal of Hazardous Materials | volume = 324 | issue = Pt B | pages = 762–772 | date = February 2017 | pmid = 27894754 | doi = 10.1016/j.jhazmat.2016.11.055 | bibcode = 2017JHzM..324..762B }}</ref>

=== Aquatic toxicity ===
[[File:Annual mean sea surface nitrate (World Ocean Atlas 2009).png|right|thumb|Sea surface nitrate from the [[World Ocean Atlas]]]]
In [[freshwater]] or [[estuary|estuarine]] systems close to land, nitrate can reach concentrations that are lethal to fish. While nitrate is much less toxic than ammonia,<ref>{{cite journal | vauthors = Romano N, Zeng C | title = 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) | journal = Environmental Toxicology and Chemistry | volume = 26 | issue = 9 | pages = 1955–1962 | date = September 2007 | pmid = 17705664 | doi = 10.1897/07-144r.1 | bibcode = 2007EnvTC..26.1955R | s2cid = 19854591 }}</ref> levels over 30 ppm of nitrate can inhibit growth, impair the immune system and cause stress in some aquatic species.<ref>{{cite web | url=http://freshaquarium.about.com/od/watercare/a/nitrates.htm | title=Nitrates in the Aquarium | work=About.com | access-date=October 30, 2013 | author=Sharpe, Shirlie | archive-date=July 24, 2011 | archive-url=https://web.archive.org/web/20110724135608/http://freshaquarium.about.com/od/watercare/a/nitrates.htm | url-status=dead }}</ref> Nitrate toxicity remains a subject of debate.<ref>{{cite journal | vauthors = Romano N, Zeng C | title = Effects of potassium on nitrate mediated alterations of osmoregulation in marine crabs | journal = Aquatic Toxicology | volume = 85 | issue = 3 | pages = 202–208 | date = December 2007 | pmid = 17942166 | doi = 10.1016/j.aquatox.2007.09.004 | bibcode = 2007AqTox..85..202R }}</ref>

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

== Human impacts on ecosystems through nitrate deposition ==
[[File:Nitrate and Phosphate Pacific Ocean.jpg|thumb|Excessive nitrate and [[phosphate]] concentrations measured in the Pacific Ocean{{Citation needed|date=September 2024}}]]
Nitrate deposition into ecosystems has markedly increased due to [[Anthropogenic effect|anthropogenic]] activities, notably from the widespread application of nitrogen-rich [[fertilizer]]s in agriculture and the emissions from [[fossil fuel]] combustion.<ref>{{cite journal | vauthors = Kanakidou M, Myriokefalitakis S, Daskalakis N, Fanourgakis G, Nenes A, Baker AR, Tsigaridis K, Mihalopoulos N | title = Past, Present and Future Atmospheric Nitrogen Deposition | journal = Journal of the Atmospheric Sciences | volume = 73 | issue = 5 | pages = 2039–2047 | date = May 2016 | pmid = 32747838 | pmc = 7398418 | doi = 10.1175/JAS-D-15-0278.1 | bibcode = 2016JAtS...73.2039K }}</ref> Annually, about 195 million [[Tonne|metric tons]] of synthetic nitrogen fertilizers are used worldwide, with nitrates constituting a significant portion of this amount.<ref name=":2">{{Cite web |title=Global fertilizer consumption by nutrient 1965-2021 |url=https://www.statista.com/statistics/438967/fertilizer-consumption-globally-by-nutrient/ |access-date=2024-04-20 |website=Statista |language=en}}</ref> In regions with intensive agriculture, such as parts of the U.S., China, and India, the use of nitrogen fertilizers can exceed 200 kilograms per hectare.<ref name=":2" />

The impact of increased nitrate deposition extends beyond plant communities to affect [[Soil microbiology|soil microbial populations]].<ref>{{cite journal | vauthors = Li Y, Zou N, Liang X, Zhou X, Guo S, Wang Y, Qin X, Tian Y, Lin J | title = Effects of nitrogen input on soil bacterial community structure and soil nitrogen cycling in the rhizosphere soil of <i>Lycium barbarum</i> L | journal = Frontiers in Microbiology | volume = 13 | pages = 1070817 | date = 2023-01-10 | pmid = 36704567 | pmc = 9871820 | doi = 10.3389/fmicb.2022.1070817 | doi-access = free }}</ref> The change in soil chemistry and nutrient dynamics can disrupt the natural processes of [[nitrogen fixation]], [[nitrification]], and [[denitrification]], leading to altered microbial community structures and functions. This disruption can further impact the [[Nutrient cycle|nutrient cycling]] and overall [[ecosystem]] health.<ref>{{cite journal | vauthors = Melillo JM | title = Disruption of the global nitrogen cycle: A grand challenge for the twenty-first century : This article belongs to Ambio's 50th Anniversary Collection. Theme: Eutrophication | journal = Ambio | volume = 50 | issue = 4 | pages = 759–763 | date = April 2021 | pmid = 33534057 | pmc = 7982378 | doi = 10.1007/s13280-020-01429-2 | bibcode = 2021Ambio..50..759M }}</ref>

== Dietary nitrate ==
A source of nitrate in the human diets arises from the consumption of leafy green foods, such as [[spinach]] and [[arugula]]. {{chem|NO|3|−}} can be present in [[beetroot]] juice. Drinking water represents also a primary nitrate intake source.<ref name="Hord2009"/>

Nitrate ingestion rapidly increases the [[Blood plasma|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.<ref>{{Cite book| vauthors = Maughan RJ |title=Food, Nutrition and Sports Performance III|publisher=Taylor & Francis|year=2013|isbn=978-0-415-62792-4|location=New York|pages=63}}</ref>

=== Cured meats ===
''Nitrite'' ({{chem2|NO2-}}) consumption is primarily determined by the amount of [[processed meat]]s eaten, and the concentration of nitrates ({{chem2|NO3-}}) added to these meats ([[bacon]], [[sausage]]s…) for their curing. Although [[nitrite]]s are the nitrogen species chiefly used in [[Curing (food preservation)|meat curing]], nitrates are used as well and can be transformed into nitrite by microorganisms, or in the digestion process, starting by their dissolution in [[saliva]] and their contact with the [[microbiota]] of the mouth. Nitrites lead to the formation of [[Carcinogenesis|carcinogenic]] [[nitrosamine]]s.<ref name="pmid12421881">{{cite journal | vauthors = Bingham SA, Hughes R, Cross AJ | title = Effect of white versus red meat on endogenous N-nitrosation in the human colon and further evidence of a dose response | journal = The Journal of Nutrition | volume = 132 | issue = 11 Suppl | pages = 3522S–3525S | date = November 2002 | pmid = 12421881 | doi = 10.1093/jn/132.11.3522S | doi-access = free }}</ref> The production of nitrosamines may be inhibited by the use of the [[antioxidant]]s [[vitamin C]] and the [[alpha-tocopherol]] form of [[vitamin E]] during curing.<ref name="pmid22464105">{{cite journal | vauthors = Parthasarathy DK, Bryan NS | title = Sodium nitrite: the "cure" for nitric oxide insufficiency | journal = Meat Science | volume = 92 | issue = 3 | pages = 274–279 | date = November 2012 | pmid = 22464105 | doi = 10.1016/j.meatsci.2012.03.001 }}</ref>

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).<ref>{{cite web | url=https://www.tastingtable.com/1132614/is-there-a-difference-between-cured-and-uncured-bacon/ | title=Is There a Difference Between Cured and Uncured Bacon? | date=9 December 2022 }}</ref>{{Better source needed|date=August 2023}} "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 exactly the same chemical with the only essential difference being the origin.

[[Antihypertensive drug|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]] (NO).<ref name="Hord2009">{{cite journal | vauthors = Hord NG, Tang Y, Bryan NS | title = Food sources of nitrates and nitrites: the physiologic context for potential health benefits | journal = The American Journal of Clinical Nutrition | volume = 90 | issue = 1 | pages = 1–10 | date = July 2009 | pmid = 19439460 | doi = 10.3945/ajcn.2008.27131 | doi-access = free }}</ref>

== Domestic animal feed ==
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:<ref>{{cite web | url=http://www1.agric.gov.ab.ca/$department/deptdocs.nsf/all/faq8911 | title=Nitrate Risk in Forage Crops - Frequently Asked Questions | publisher=Government of Alberta | work=Agriculture and Rural Development | access-date=October 30, 2013}}</ref>

{| class="wikitable"
|-
! Category !! %NO<sub>3</sub> !! %NO<sub>3</sub>–N !! %KNO<sub>3</sub> !! Effects
|-
| 1 || <&nbsp;0.5 || <&nbsp;0.12 || <&nbsp;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 || <&nbsp;1.23 || <&nbsp;0.28 || <&nbsp;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 ==
Nitrate formation with elements of the periodic table:
{{nitrates}}

== See also ==
* [[Ammonium]]
* [[Eutrophication]]
* [[f-ratio (oceanography)|f-ratio in oceanography]]
* [[Frost diagram]]
* [[Nitrification]]
* [[Nitratine]]
* [[Nitrite]], the anion {{chem|NO|2|−}}
* [[Nitrogen oxide]]
* [[Nitrogen trioxide]], the neutral radical {{chem|NO|3}}
* [[Peroxynitrate]], {{chem|OONO|2|–}}
* [[Sodium nitrate]]

== References ==
{{Reflist}}

== External links ==
{{commons category|Nitrate ion}}
* [https://www.atsdr.cdc.gov/csem/csem.html ATSDR – Case Studies in Environmental Medicine – Nitrate/Nitrite Toxicity] ([https://web.archive.org/web/20100304080534/http://www.atsdr.cdc.gov/csem/nitrate/no3cover.html archive])

{{Nitric oxide signaling}}
{{nitrogen compounds}}
{{Functional groups}}
{{Nitrates}}
{{Authority control}}

[[Category:Nitrates| ]]
[[Category:Curing agents]]
[[Category:Functional groups]]
[[Category:Garde manger]]
[[Category:Nitrogen cycle]]
[[Category:Non-coordinating anions]]
[[Category:Nitrogen oxyanions]]
[[Category:Water quality indicators]]