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Nitrate is apolyatomic ion with thechemical formulaNO− 3.Salts containing thision are callednitrates. Nitrates are common components of fertilizers and explosives.[1] Almost all inorganic nitrates aresoluble inwater. An example of an insoluble (inorganic) nitrate isbismuth oxynitrate.
In nature, nitrates are produced by a number of species ofnitrifying bacteria in the natural environment usingammonia orurea as a source of nitrogen and source of free energy. Nitrate compounds forgunpowder were historically produced, in the absence of mineral nitrate sources, by means of variousfermentation processes using urine and dung. Modern nitrate production is mostly focused on creation forfertilizer and chemical manufacturing for various applications, such as medicine synthesis, ceramics andpreservation of meat. Annually, about 195 millionmetric tons of synthetic nitrogen fertilizers are used worldwide, with nitrates constituting a significant portion of this amount.[2]
Because nitrates are soluble and easily can be swept away from the soil because of precipitation, excessive agricultural use has been associated withnutrient runoff,water pollution, and the proliferation ofaquatic dead zones. Direct exposure of nitrates for humans can have direct health consequences: the excess consumption of nitrates in cured meats is associated with intestinal cancers.[3]
The nitrateanion is theconjugate base ofnitric acid, consisting of one centralnitrogenatom surrounded by three identically bondedoxygen atoms in atrigonal planar arrangement.[4] The nitrate ion carries aformal charge of −1.[5] 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 ofresonance. Like theisoelectroniccarbonate ion, the nitrate ion can be represented by three resonance structures:
In theNO− 3 anion, theoxidation state of the central nitrogen atom is V (+5). This corresponds to the highest possibleoxidation number of nitrogen. Nitrate is a potentially powerfuloxidizer as evidenced by itsexplosive behaviour at high temperature when it isdetonated inammonium nitrate (NH4NO3), orblack powder, ignited by theshock wave of aprimary explosive. In contrast tored fuming nitric acid (HNO3/N2O4), or concentratednitric acid (HNO3), nitrate inaqueous solution at neutral or highpH is only a weakoxidizing agent inredox reactions in which thereductant does not produce hydrogen ions (such as mercury going tocalomel). 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 ofmicroorganisms, or reductants such as organic matter. In fact, nitrogen gas is thermodynamically stable in the presence of1 atm of oxygen only in very acidic conditions, and otherwise would combine with it to form nitrate. This is shown by subtracting the two oxidation reactions:[6]
N2 + 6 H2O → 2 NO−3 + 12 H+ + 10 e−
2 H2O → O2 + 4 H+ + 4 e−
giving:
2 N2 + 5 O2 + 2 H2O → 4 NO−3 + 4 H+
Dividing by 0.0118 and rearranging gives the equilibrium relation:
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 asammonia, which is then oxidized, eventually becoming nitrate. Nitrate does not accumulate to high levels in nature because it reacts with reductants in the process calleddenitrification (seeNitrogen cycle).
Nitrate is used as a powerful terminalelectron acceptor bydenitrifying bacteria to deliver the energy they need to thrive. Underanaerobic conditions, nitrate is the strongest electron acceptor used byprokaryotemicroorganisms (bacteria andarchaea) to respirate. Theredox coupleNO−3/N2 is at the top of theredox scale for theanaerobic 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 habitated by microorganisms, nitrate is a quite unstable and labile dissolved chemical species because it ismetabolised 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) iscatalyticallyreduced into nitrite (NO−2) by theenzymatic activity of bacteria. In aqueous solution, dissolved nitrite, N(III), is a more powerful oxidizer that nitrate, N(V), because it has to accept lesselectrons and itsreduction is lesskinetically 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 onmyoglobin, accentuating its red coloration. NO is an important biologicalsignaling molecule and intervenes in thevasodilation process. Still, it can also producefree radicals inbiological tissues, accelerating their degradation and aging process. Thereactive oxygen species (ROS) generated by NO contribute to theoxidative stress, a condition involved in vascular dysfunction andatherogenesis.[9]
The nitrateanion is commonly analysed in water byion chromatography (IC) along with other anions also present in the solution. The main advantage of IC is its ease and the simultaneousanalysis 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.[10]
Previously, nitrate determination relied onspectrophotometric andcolorimetric measurements after a specific reagent is added to thesolution to reveal a characteristic color (often red because it absorbs visible light in the blue). Because of interferences with the brown color ofdissolved organic matter (DOM:humic andfulvic acids) often present insoil pore water, artefacts can easily affect theabsorbance 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 (NO−3) often rely on its conversion tonitrite (NO−2) followed by nitrite-specific tests. Thereduction of nitrate to nitrite can be effected by acopper-cadmiumalloy, metalliczinc,[11] orhydrazine. The most popular of these assays is theGriess test, whereby nitrite is converted to a deeply red coloredazo dye suited forUV–vis spectrophotometry analysis. The method exploits the reactivity ofnitrous acid (HNO2) derived from the acidification of nitrite. Nitrous acid selectively reacts with aromatic amines to givediazonium salts, which in turn couple with a second reagent to give theazo dye. Thedetection limit is 0.02 to 2 μM.[12] Such methods have been highly adapted to biological samples[13] and soil samples.[14][15]
For direct online chemical analysis using a flow-through system, the water sample is introduced by aperistaltic pump in aflow injection analyzer, and the nitrate or resulting nitrite-containing effluent is then combined with a reagent for its colorimetric detection.
Nitrate salts are found naturally on earth in arid environments as large deposits, particularly ofnitratine, a major source ofsodium nitrate.
Nitrates are produced by a number of species ofnitrifying bacteria in the natural environment usingammonia orurea as a source of nitrogen and source of free energy. Nitrate compounds forgunpowder were historically produced, in the absence of mineral nitrate sources, by means of variousfermentation processes using urine and dung.
Lightning strikes in earth's nitrogen- and oxygen-rich atmosphere produce a mixture of oxides of nitrogen, which formnitrous ions and nitrate ions, which are washed from the atmosphere by rain or inoccult deposition.
Nitrate is achemical 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.[20] 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.[21]
Nitrates are used asfertilizers inagriculture because of their high solubility and biodegradability. The main nitrate fertilizers areammonium,sodium,potassium,calcium, andmagnesium salts. Several billion kilograms are produced annually for this purpose.[1] 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.[22]
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.[23] 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.[24]
Sodium nitrate is used to remove air bubbles from moltenglass and someceramics. Mixtures ofmolten salts are used to harden the surface of some metals.[1]
via the formation of nitrite, nitrate is implicated inmethemoglobinemia, a disorder ofhemoglobin inred blood cells, especially affecting infants and toddlers from ingesting nitrates in drinking water.[26][27]
One of the most common cause ofmethemoglobinemia in infants is due to the ingestion of nitrates and nitrites throughwell water or foods.
In fact, nitrates (NO−3), often present at too highconcentration in drinkwater, are only the precursor chemical species ofnitrites (NO−2), the real culprits of methemoglobinemia. Nitrites produced by themicrobial reduction of nitrate (directly in the drinkwater, or after ingestion by the infant's digestive system) are more powerfuloxidizers than nitrates and are the chemical agent really responsible for theoxidation of Fe2+ into Fe3+ in thetetrapyrroleheme ofhemoglobin. Indeed, nitrate anions are too weak oxidizers inaqueous solution to be able to directly, or at least sufficiently rapidly, oxidize Fe2+ into Fe3+, because ofkinetics limitations.
Infants younger than four months are at greater risk given that they drink more water per body weight, they have a lowerNADH-cytochrome b5 reductase activity, and they have a higher level offetal hemoglobin which converts more easily tomethemoglobin. Additionally, infants are at an increased risk after an episode ofgastroenteritis due to the production ofnitrites bybacteria.[28]
However, other causes than nitrates can also affect infants and pregnant women.[29][30] Indeed, theblue baby syndrome can also be caused by a number of other factors such as thecyanotic heart disease, acongenital heart defect resulting in low levels of oxygen in the blood,[31] or by gastric upset, such as diarrheal infection, protein intolerance, heavy metal toxicity, etc.[32]
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).[34]
Infreshwater orestuarine systems close to land, nitrate can reach concentrations that are lethal to fish. While nitrate is much less toxic than ammonia,[35] levels over 30 ppm of nitrate can inhibit growth, impair the immune system and cause stress in some aquatic species.[36] Nitrate toxicity remains a subject of debate.[37]
In most cases of excess nitrate concentrations in aquatic systems, the primary sources are wastewater discharges, as well assurface runoff from agricultural orlandscaped areas that have received excess nitrate fertilizer. The resultingeutrophication and algae blooms result inanoxia anddead zones. As a consequence, as nitrate forms a component oftotal dissolved solids, they are widely used as an indicator ofwater quality.
Human impacts on ecosystems through nitrate deposition
Excessive nitrate andphosphate concentrations measured in the Pacific Ocean[38]
Nitrate deposition into ecosystems has markedly increased due toanthropogenic activities, notably from the widespread application of nitrogen-richfertilizers in agriculture and the emissions fromfossil fuel combustion.[39] Annually, about 195 millionmetric tons of synthetic nitrogen fertilizers are used worldwide, with nitrates constituting a significant portion of this amount.[2] 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.[2]
A source of nitrate in the human diets arises from the consumption of leafy green foods, such asspinach andarugula.NO− 3 can be present inbeetroot juice. Drinking water represents also a primary nitrate intake source.[42]
Nitrate ingestion rapidly increases theplasma 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 ofnitric oxide, NO. Nitric oxide is a physiologicalsignaling molecule which intervenes in, among other things, regulation of muscle blood flow and mitochondrial respiration.[43]
Nitrite (NO−2) consumption is primarily determined by the amount ofprocessed meats eaten, and the concentration of nitrates (NO−3) added to these meats (bacon,sausages…) for their curing. Althoughnitrites are the nitrogen species chiefly used inmeat curing, nitrates are used as well and can be transformed into nitrite by microorganisms, or in the digestion process, starting by their dissolution insaliva and their contact with themicrobiota of the mouth. Nitrites lead to the formation ofcarcinogenicnitrosamines.[44] The production of nitrosamines may be inhibited by the use of theantioxidantsvitamin C and thealpha-tocopherol form ofvitamin E during curing.[45]
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).[46][better source needed] "Uncured" meat is in fact actually cured with nitrites with virtuallyno 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.
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:[47]
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.
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^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):1955–1962.Bibcode:2007EnvTC..26.1955R.doi:10.1897/07-144r.1.PMID17705664.S2CID19854591.
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