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Nitrate in Groundwater

  1. Thomas Addiscott

Published Online: 15 APR 2003

DOI: 10.1002/047126363X.agr174

Encyclopedia of Agrochemicals

Encyclopedia of Agrochemicals

How to Cite

Addiscott, T. 2003. Nitrate in Groundwater. Encyclopedia of Agrochemicals. .

Author Information

  1. Rothamsted Experimental Station, Herts, Harpenden, United Kingdom

Publication History

  1. Published Online: 15 APR 2003

1 Introduction

  1. Top of page
  2. Introduction
  3. Nomenclature
  4. Physical Properties
  5. Agricultural Uses
  6. Chemistry
  7. Environmental Fate
  8. Nitrate and Health
  9. Bibliography
  10. Further Reading

The nitrate ion is one of the more ubiquitous chemical substances on the planet and is nearly always found in water. Most of the water around us contains nitrate, but the water with which we are concerned here is groundwater, which is water accumulated in the saturated zones of certain rock formations, usually at depth. Most of this water has passed through the soil before it accumulates, so that activities at the soil surface, particularly agriculture, can have a strong influence on the concentrations of nitrate and other agrochemicals in groundwater. Despite its commonplace nature, nitrate has for at least two decades been a source of widespread concern because of its perceived effects on our environment and our health. As a result, the “nitrate problem” has been a major influence on agroecological research in the developed world during this period. Environmental concern has centered mainly on the formation of algal blooms and excessive growth of water plants in surface fresh waters and in the coastal areas of the sea. Worries about our health spring from fears that nitrate in potable water might cause stomach cancer in adults or methemoglobinemia (“blue-baby” syndrome) in infants. Recent medical research, however, suggests not only that nitrate is beneficial to our health but also that we produce it within our bodies. Water supplies are drawn from both ground and surface waters according to their availability. This article is concerned with nitrate in groundwater, which has health, rather than environmental, implications, but environmental issues are not ignored.

2 Nomenclature

  1. Top of page
  2. Introduction
  3. Nomenclature
  4. Physical Properties
  5. Agricultural Uses
  6. Chemistry
  7. Environmental Fate
  8. Nitrate and Health
  9. Bibliography
  10. Further Reading

“Nitrate” is the chemical name for the NO3 ion, and it is not known by any other. The practice of referring to “nitrates” in natural waters and water supplies is incorrect because, as in all dilute electrolyte solutions, the anions and cations are dissociated from each other. The species with which we are concerned is, therefore, the free nitrate ion, which is unique rather than plural.

2.1 Structural Formula

The nitrate ion, NO3, has a symmetrical planar trigonal structure in which the nitrogen atom has a formal positive charge. Two negative charges are shared between the three oxygen atoms in a resonance structure comprising three electronic conformations in which each of the oxygen atoms, in turn, is without charge. The uncharged atom has two electron pairs and is attached to the nitrogen atom by a π-bond, and the charged atoms have three electron pairs.

3 Physical Properties

  1. Top of page
  2. Introduction
  3. Nomenclature
  4. Physical Properties
  5. Agricultural Uses
  6. Chemistry
  7. Environmental Fate
  8. Nitrate and Health
  9. Bibliography
  10. Further Reading

3.1 Solubility

The salts formed by the nitrate ion are generally soluble, and calcium nitrate has such a high affinity for water that it is deliquescent, which means that it will pick up moisture from the air and dissolve in it. The main cations in groundwater are likely to be calcium, magnesium, potassium, sodium, iron, and aluminium, and the salts they form with nitrate are all very soluble (Table 1). Ammonium nitrate is also highly soluble. Calcium is usually the dominant cation in groundwater, and the nitrate concentration at the limit of solubility for calcium nitrate is 32,000 times greater than the U.S. limit for nitrate concentration in potable water and 28,000 times greater than the E.C. limit. Solubility cannot, therefore, limit nitrate concentrations in groundwater.

Table 1. Solubilities in Cold Water1 of the Salts of the Nitrate Ion (26)
CationSaltSolubility (g m−3)
SaltNitrate2
  • 1

    The temperatures at which the solubilities in the table (26) had been determined were not all the same and ranged from 0 K to 25 K.

  • 2

    The “nitrate” concentration is that corresponding to the solubility of the salt.

Ca++Ca(NO3)2 · 4H2O2.66 × 1061.40 × 106
Mg++Mg(NO3)2 · 6H2O1.25 × 1060.61 × 106
K+KNO30.32 × 1060.18 × 106
Na+NaNO30.92 × 1060.67 × 106
NH4+NH4NO31.18 × 1060.91 × 106
Fe++Fe(NO3)2 · 6H2O0.84 × 1060.36 × 106
Fe+++Fe(NO3)3 · 6H2O1.50 × 1060.80 × 106
Al+++Al(NO3)3 · 9H2O0.64 × 1060.32 × 106

3.2 Sorption

Nitrate, being an anion, is attracted to positively charged surfaces. Nearly all agricultural soils in the developed world are usually maintained at pH values that are not acid enough to permit the development of the positive charges that will retain nitrate. However, there are some soils, particularly highly weathered soils in the Tropics, which are sufficiently acid for nitrate retention to occur. In the absence of clear evidence that the soil is positively-charged, it will be advisable to assume that sorption, like solubility, does nothing to limit nitrate concentrations in groundwater.

3.3 Other Properties

Nitrate and its salts do not exert a vapor pressure. The melting point of each salt depends on the cation in the salt.

4 Agricultural Uses

  1. Top of page
  2. Introduction
  3. Nomenclature
  4. Physical Properties
  5. Agricultural Uses
  6. Chemistry
  7. Environmental Fate
  8. Nitrate and Health
  9. Bibliography
  10. Further Reading

Nitrate is used in agriculture solely as a constituent of fertilizers.

4.1 Formulations

Guano, the oldest form of nitrate fertilizer, is the accumulated excrement of sea-birds and is found most notably on the cliffs of the Peruvian coast and nearby islands. Because of the marine origin of this material, the main cation associated with the nitrate is sodium. Pure sodium nitrate would contain 16.5% of N, but guano contains a smaller and somewhat variable percentage of N.

Ammonium nitrate is widely used as a nitrogen fertilizer, particularly in Europe, because both the cation and the anion contain nitrogen, so the pure salt has 36% of N. The solid fertilizer contains fillers and stabilizers for quality and safety reasons (see also the discussion of Reactivity in the Chemistry section) and usually has a stated N content of 33% to 35.5%. Ammonium nitrate is also used in liquid fertilizer formulations, often in combination with urea (see below).

Calcium ammonium nitrate is also used as a fertilizer. It contains 26% to 28% of N, depending on the manufacturer.

We also need to note three other fertilizers, urea, ammonium sulphate, and directly injected ammonia, which do not contain nitrate but are transformed to nitrate by soil microbes.

Urea is a very useful fertilizer where transport is a problem because it contains 46% of N, more than any other solid N source, so that the least possible noneffective weight has to be carried. Nearly half the world's fertilizer production is as urea. Urea is converted to ammonium by the Urease enzyme, which is very widespread.

Ammonium sulfate used to be a popular fertilizer because it was a cheap by-product of gas production from coal. Because of the increased use of natural gas in many countries, its use has declined during the last 20 years.

Injected ammonia. Ammonia liquified under pressure can be taken to the field and injected directly into the soil, usually to a depth of about 100 mm. The machinery used cuts a slot in the soil with a disc. The nozzle feeding the liquified ammonia is directly behind the disc and is followed by a flat wheel, which closes the slot. This is a very efficient source in that ammonia is 82% N, but the specialized equipment needed for storing and injecting the ammonia tends to localize its use.

4.2 Compound Fertilizers

All the above sources of nitrogen are described as “straight” because they supply nitrogen alone and do not provide phosphate or potassium, the other two major nutrients. In compound fertilizers, nitrogen is mixed, usually in granules but sometimes in liquid form, with either or both of these nutrients and occasionally others. The nutrient composition of a compound depends on the crop for which it is manufactured. Cereal crops, for example, need a large proportion of nitrogen, whereas potatoes need more phosphate and potassium than other crops (see Biological role discussion in the Chemistry section). Ammonium phosphate is often a constituent of compound fertilizers. The diammonium phosphate seems to be the more widely used. Saltpeter, natural potassium nitrate, could have been an early compound fertilizer because it contains two of the three main nutrients needed by plants. However, it was not greatly used as a fertilizer in the past, almost certainly because it was more valuable as a preserving agent for meat and as a constituent of gunpowder. It is still not used on a wide scale, although it may be used in foliar applications.

Table 2 shows world fertilizer nitrogen consumption in 1994–1995 for the various sources of nitrogen. More recent figures were not found, but overall consumption of nitrogen fertilizer has increased by about 9% since then, probably without much change in the ratios between the sources.

Table 2. World Consumption in 1994/1995 of Nitrogen in Various Types of Fertilizer (27)
Type of FertilizerN Consumed (Tonnes)% of Total
Ammonium nitrate6.58 × 106 9.0
Calcium ammonium nitrate3.69 × 106 5.1
Urea31.57 × 106 43.3
Ammonium sulphate2.37 × 106 3.2
Injected ammonia4.20 × 106 5.8
N solutions3.79 × 106 5.2
Other straights9.05 × 106 12.4
Compound fertilizers11.68 × 106 16.0
Total72.93 × 106100.0

5 Chemistry

  1. Top of page
  2. Introduction
  3. Nomenclature
  4. Physical Properties
  5. Agricultural Uses
  6. Chemistry
  7. Environmental Fate
  8. Nitrate and Health
  9. Bibliography
  10. Further Reading

5.1 Reactivity

Nitrate is the most fully oxidized compound of nitrogen and is, therefore, stable to oxidation but potentially a strong oxidizing agent. Saltpeter (potassium nitrate) has long been the oxidizing constituent of constituent of gunpowder. Solid ammonium nitrate can explode because the nitrate moiety can oxidize the ammonium moiety. Mixed with aluminium powder it formed Ammonal, one of the most widely used explosives in the Second World War. There is, however, an important difference between the salts of nitrate in the solid and dissolved states. Because of the stability conferred by the resonance structure of the ion, nitrate in a near neutral dilute solution of its dissociated salts (as found in groundwater) is unreactive chemically. Its biological reactivity is discussed below.

5.2 Synthesis and Manufacture

Synthetic nitrate is manufactured (1) in two main stages, ammonia (NH3) being produced and then oxidized to nitrate.

Ammonia has long been synthesized from nitrogen and hydrogen in the Haber process in which the two elements are reacted over a catalyst at high temperature and pressure. Modern methods of production often involve the steam reforming of natural gas, in which the methane (CH4) from the natural gas and the steam (H2O) react with the air to give carbon dioxide (CO2) and hydrogen (H2). Production often involves a secondary reforming process. Nitrogen (N2) left from the air reacts with the hydrogen over a nickel catalyst to give ammonia (NH3). Sulpfur and oxygen compounds (particularly the CO2) have to be removed before the reaction over the catalyst, because they inhibit its activity. Other sources of carbon and hydrogen, such as naphtha, oil, and coal can be used but give poorer energy efficiency than natural gas.

The ammonia is oxidized to nitric acid over a platinum catalyst, which is alloyed with rhodium for strength, and the nitric acid is reacted either with ammonia to give ammonium nitrate or with the appropriate oxide, hydroxide, or carbonate to provide the nitrate salt required. The ammonia may also be reacted with phosphoric acid to give ammonium phosphates, which are also fertilizer materials.

5.3 Biological Nitrate Production

Because nitrate is chemically stable and cannot be oxidized further, it is the end product of a key biological nitrogen chain in the soil (2, 3). The topsoil (first 250 mm of the soil) contains large quantities of nitrogen, often of the order of 5,000 kg ha−1, in organic forms (“organic” is used here in its original chemical sense of “pertaining to the special chemistry of carbon” rather than in that of recent farming philosophy). The organic carbon and nitrogen come from the debris of green plants, including dead roots, and dead tops where they are not harvested, exudates from roots, and animal excreta. This organic matter is colonized by soil organisms of various sizes, ranging from earthworms, through springtails and mites, to bacteria and fungi. They form a chain in which the largest organisms make the organic matter more available to the smallest but also predate on them. The process is described as mineralization, because the end products are the most highly oxidized forms of carbon and nitrogen, carbon dioxide and nitrate, which are in the realm of inorganic—or mineral—chemistry. The final stage of the process for nitrogen, the oxidation of ammonium to nitrate, is important where ammonium fertilizers are used. It proceeds in two stages (2, 3), each of which is effected by a chemoautotrophic bacterium. First, ammonium is oxidized to nitrite by Nitrosomonas species:

  • equation image(1)

and then the nitrite is oxidized to nitrate by Nitrobacter species:

  • equation image(2)

Some nitrous oxide (N2O) is formed during the second stage.

The mineralization of organic nitrogen is an entirely natural process, and it cannot be controlled to more than a limited extent because the nitrate is produced in the soil without any human intervention. Measurements made at Rothamsted from 1877 to 1915 and summarized more recently (4) show that, even back in the 1870s, nitrate concentrations in water draining from an uncropped, unploughed soil that had received no nitrogen fertilizer for at least 10 years exceeded the present day U.S. and E.C. limits for potable water of 44 g and 50 gm−3. During the 38 years of the study, the soil lost more than 1,000 kg ha−1 of nitrogen from its organic matter, all of which emerged as nitrate in the drainage from the soil. Losses of nitrate produced by mineralization need to be considered in any discussion of groundwater issues because there is evidence (5) that they are usually greater than direct losses from fertilizer.

5.4 Adventitious Nitrate Production

Many industrial processes emit substantial quantities of oxides of nitrogen to the atmosphere. Further emissions of this kind come from vehicles with internal combustion engines, which cause a reaction between the nitrogen and oxygen of the atmosphere by compressing them and subjecting them to high temperature and a spark. Industry and motor traffic each generates about half of these man-made nitrogen oxides. Nitrogen oxides are also produced naturally by lightning (6). Other activities, notably farming, emit ammonia to the atmosphere, and this is readily converted to nitrate, particularly when it reaches the soil.

Between 1877 and 1915, during the experiments mentioned in the previous section, about 6 kg ha−1 of mineral nitrogen, as ammonium and nitrate, was deposited in rainfall at Rothamsted each year. By 1990, measurements at four sites in southeast England showed an annual deposition of 35–40 kg ha−1 (7), but these measurements included deposition of nitrate on particulate matter and dry deposition of nitrogen oxides, nitric acid, and ammonia in addition to ammonium and nitrate in rain. A more recent estimate (8) suggests that about 37 kg ha−1 of nitrogen is deposited annually on bare soil and 48 kg ha−1 on soil carrying the extra deposition area supplied by a winter wheat crop. The latter amount is one-quarter of the average application of nitrogen fertilizer in England and Wales. This deposited nitrogen probably contributes 10–15 kg ha−1 of nitrate-N to annual losses from the soil.

5.5 Biological Role

Nitrogen is vital to the growth of plants. It is part of all the essential constituents of cells, including the chlorophyll needed for photosynthesis; the DNA and RNA, which encode the plant's program for growth and development; the proteins—including the enzymes, which catalyze all biochemical processes; and the cell walls, which do more than just hold the cell together. All plant nutrients increase the growth and yield of crop plants but nitrogen has the largest effect, except in plants that form large storage organs such as potato tubers. These plants store large amounts of phosphorus and potassium in their storage organs and, therefore, have a larger demand for these nutrients.

The nitrate ion is usually the main form of nitrogen taken up by the plant's roots from nonacid aerobic soils, although ammonium ions can also contribute to its uptake. The ammonium ion predominates in anaerobic soils such as those found in rice paddies and may be the main form taken up there. The form of nitrogen supplied has a considerable effect on the cation-anion balance of the plant and, hence, its growth (9). As might be expected, supplying nitrogen as the NH4 cation rather than the NO3 anion lessens the uptake of other cations, particularly potassium. The apparent preference of many plants for nitrate rather than ammonium in aerobic soils probably reflects not only the effects of charge but also the ubiquity of nitrate in soils and the sorption by nonacid soils of ammonium but not nitrate.

6 Environmental Fate

  1. Top of page
  2. Introduction
  3. Nomenclature
  4. Physical Properties
  5. Agricultural Uses
  6. Chemistry
  7. Environmental Fate
  8. Nitrate and Health
  9. Bibliography
  10. Further Reading

The fate of nitrate and the type of environmental problem it causes depend on two main processes, leaching and denitrification.

6.1 Leaching

Nitrate is leached from the soil because neither solubility nor sorption withholds it from the water passing through the soil, so, whenever water moves, nitrate tends to move with it. The only restriction on such movement is afforded by the structure of the soil, which holds some water in aggregates and larger structural units (10). This water does not move appreciably, so any nitrate within it is temporarily safe from leaching. Nitrate may be in the immobile water because it was produced within the structural unit or because it diffused into it. Inward diffusion would result from a greater nitrate concentration on the outside than on the inside and could be the result of fertilizer application. Once the concentration outside the structural unit has fallen below that on the inside, nitrate diffuses out into water that may move downwards, but the “hold-back” provided by the soil structure can help usefully to restrict nitrate concentrations in water draining from the soil. The effect is greatest in the soils with the best-defined soil structure, which are usually clay soils.

The environmental problems caused by nitrate leaching spring from the fact that it is not only land plants that use nitrogen for extra growth. Plants growing in water respond too, and their extra growth is usually not welcome. Increased nitrate concentrations caused by leached nitrate in rivers and lakes encourage reeds to grow to excess, narrowing waterways and potentially overloading and damaging banks. The proliferation of underwater plants fouls the propellers of boats, entangles the tackle of fishermen, and blocks water supply conduits, thereby damaging machinery.

The large plants are usually not as great a problem as algal blooms. The blue-green algae are very small single-celled plants of the Cyanobacteria species that grow on the surface of practically anything, including water (11). Some of them are toxic, and others are a problem because of buoyancy conferred by the gas vesicles they contain, which enable them to rise to the surface of the water during calm conditions. The resulting “bloom” or “scum” is often blown by even gentle breezes to the edge of the lake or river where it is particularly unpleasant—and a hazard if it is toxic. A further problem is that when algae die, the bacteria that consume them use oxygen to do so, and this lessens the supply to fish and other desirable organisms, which may die as a result. Algal blooms and other problems can also occur in the sea, particularly in partially enclosed water bodies such as the Gulf of Mexico and the Baltic Sea, which are not flushed by strong currents. Algal blooms are usually triggered in fresh water by the phosphate concentration to which they are sensitive over five orders of magnitude of concentration (11, 12), but they also depend on the presence of nitrate. They are probably more sensitive to nitrate in the sea.

6.2 Denitrification

Nitrate is, as noted above, the most fully oxidized compound of nitrogen, and, when certain microbes in the soil need oxygen and are unable to get enough, they take it from nitrate ions. The reduction is effected by facultative anaerobic bacteria, mainly of the Pseudomonas and Bacillus species once the partial pressure of oxygen has become low (<0.004 bar). The pathway for the process can be summarized as:

  • equation image(3)

The first stage occurs under aerobic conditions (and corresponds to Equation 2), but the subsequent stages are all anaerobic. The extent to which the reduction proceeds is strongly influenced by the pH and temperature of the soil. At a low temperature and a pH less than 5, the ratio of N2O to N2 is about 1, but, at a temperature of 25 °C or more and a pH greater than 6, most of the N2O is reduced to N2.

Soils are not usually uniform, and, even in mainly aerobic soils, there are usually small localized anaerobic zones in which denitrification occurs (13). This spatial variability, together with the variable ratio of N2O to N2, makes denitrification very difficult to measure. It also implies that nitrification and denitrification may be occurring simultaneously close to each other in the soil.

Denitrification is beneficial to groundwater because it decreases nitrate concentrations in it. The process is not restricted to the soil and can occur in groundwater, provided both the microbes and a suitable carbon substrate are available. The origin of the aquifer usually determines how much natural denitrification is likely (14, 15). The rock surface in Triassic Sandstone, for example, may hold enough substrate for measurable denitrification to occur, but rocks formed under hotter conditions will have less substrate available. The process is used in industrial plants for removing nitrate from water in which methanol is added as a substrate—subject to controls that ensure that the microbes exhaust the methanol before the nitrate. There are also reports that glucose and methanol are being introduced into aquifers for this purpose.

Denitrification is harmless to the rest of the environment if it proceeds to N2 because 78% of the atmosphere is N2. However, N2O is not beneficial in the atmosphere because it is a greenhouse gas and about 200 times as effective in this role as CO2.

6.3 Fate of Fertilizer Nitrogen

Nitrogen fertilizer obviously contributes to some extent to nitrate losses to groundwater, and we need to know how much it contributes and when. The answer depends on many variables, of which the crop grown is among the most important. It is not possible to answer the question for all crops in all countries, but the answer for winter wheat illustrates the approach involved. This is the crop on which most research has been done in England and Wales because it is the most widely grown crop and also represents the group of crops that are sown in autumn and have well-established root systems by the time they receive applications of nitrogen fertilizer in spring. Winter wheat is also a parsimonious crop; only sugar beet allows less nitrate to escape its roots.

In the U.K., nitrogen fertilizer for winter wheat is applied in spring, usually in late March or April. The fate of nitrogen applied to winter wheat in spring was determined in studies in which nitrogen from fertilizer was labeled (or “tagged”) with the isotope 15N. Several such studies were made, and they were summarized and reviewed in (5). Table 3 outlines the results from one of the experiments. The losses of the 15N depended strongly on the amount of rain in the 3 weeks after the fertilizer was applied, and the weather determined whether denitrification or leaching predominated in a particular season. On average, nearly two-thirds of the loss was by denitrification and most of the rest by leaching. Direct leaching losses from fertilizer were relatively small, commonly as little as 5% to 6%, of which 4% to 5% would be lost before harvest and the remaining 1% left in the soil at harvest and leached subsequently. Sugar beet was even more efficient than winter wheat at restricting nitrate losses, but some other crops, notably potatoes, were much less efficient.

Table 3. Fate of 15N-labeled Nitrogen Fertilizer Applied in Spring to Winter Wheat1
Fate of Fertilizer% of Amount Applied
  • 1

    In one of the experiments reviewed in (5).

  • 2

    Safe from leaching, but may be remineralized.

  • 3

    Vulnerable to leaching; the ammonium is likely to be nitrified.

In grain or straw of crop69
Left in organic matter in soil216
Left as ammonium or nitrate3 1
Lost14

The relatively small direct leaching loss cited for winter wheat can increase sharply if too much nitrogen fertilizer is applied. This implies that one of the keys to managing the nitrate problem is the correct assessment of the amount of nitrogen needed. This is a topic that has attracted great attention, and computer models are playing an increasingly important part in decision-support for fertilizer application.

7 Nitrate and Health

  1. Top of page
  2. Introduction
  3. Nomenclature
  4. Physical Properties
  5. Agricultural Uses
  6. Chemistry
  7. Environmental Fate
  8. Nitrate and Health
  9. Bibliography
  10. Further Reading

Nitrate has been blamed for stomach cancer in adults and methemoglobinemia (“blue-baby syndrome”) in infants, and this section would have been entitled “Toxicology” were it not for recent medical research. This has shown that, far from being toxic, nitrate almost certainly helps to prevent several diseases from which people die every day. Even if nitrate were responsible for stomach cancer, the toxic agent would be the nitrite to which nitrate can be reduced rather than nitrate itself.

7.1 Stomach Cancer

The theory that nitrate causes stomach cancer was based on possibilities rather than hard evidence. It was suggested that nitrate (NO3) in our vegetables and water could be changed to nitrite (NO2) by microbes in the mouth that needed an oxygen source. This nitrite could react in the stomach with secondary amines released in the digestion of protein to produce carcinogenic N-nitrosamines, which could cause stomach cancer. This string of possibilities fell apart in 1985 and 1986 when epidemiologists looked directly for the connection between stomach cancer and nitrate. They measured nitrate and nitrite in saliva samples from healthy people from two areas with a high risk of stomach cancer and two with a low risk, finding that those in the high-risk areas had less nitrate and nitrite in their saliva than those in the low-risk area (16). The nitrate-stomach cancer theory suggested they should have had more. Another study (17) related deaths from stomach cancer to nitrate concentrations in drinking water and found a negative relation rather than the positive one suggested by the theory. The trend, showing that during the past 30 years nitrate concentrations in water have increased but the incidence of stomach cancer has gone down, also lends no support to the theory.

One interesting piece of evidence found by epidemiologists involved workers in a factory producing ammonium nitrate fertilizer (18). Visitors to such factories can often taste ammonium nitrate in their saliva within a minute or two of arriving, suggesting that workers in fertilizer factories could have an enhanced risk of stomach cancer from nitrate. But an epidemiological study showed this was not so. Fertilizer workers had the same death rate from stomach cancer as workers with comparable jobs in the same area. And the workers in the fertilizer factory were not just no more prone to stomach cancer than the others, they were actually healthier. Their death rate from heart disease was significantly lower (p < 0.05), and that from respiratory disease was measurably (but not significantly) lower.

7.2 Gastroenteritis

This hint that nitrate could, after all, be good for our health is supported by recent medical research. This revealed that nitrate is converted to nitrite in the mouth by microbes similar to those that convert nitrate to nitrite in the soil (19). It showed further that the metabolic system actively secretes nitrate into saliva and will even convert the amino-acid L-arginine into nitrate for this purpose if the intake of nitrate in water and vegetables is insufficient.

The conclusion that the body pumps nitrate into the saliva and keeps microbes in the mouth to convert it to nitrite, which goes down into your stomach, is apparently alarming because it has so much in common with the hypothesized link between nitrate and stomach cancer. However, the proposed reaction between the nitrite and a secondary amine to give a carcinogenic N-nitrosamine probably does not get the chance to occur because the strong acid in the stomach decomposes it straight away. In doing so, it gives off nitric oxide and other gases, which have a key role in the stomach; they kill Salmonella, Escherichia coli, and other unwelcome bacteria in the stomach. This is the body's defense system against bacterial gastroenteritis, and nitrate is the fuel that powers it (20, 21).

Every year, 3 to 5 million people die worldwide from gastroenteritis. This statistic arises mainly from poverty, malnutrition, and bacterially polluted water, but it also raises two important questions. Why is nitrate such a chemical pariah when it is the fuel of the system that protects us from this killer disease? And why, if nitrate helps to stop gastroenteritis, are authorities in the developed world rushing to remove it from our food and water?

7.3 Heart Disease and Skin Problems

Gastroenteritis is not the only reason for taking a fresh look at nitrate. The low death rates from heart disease among fertilizer workers may have occurred because nitrate lessens platelet function and clotting and could, therefore, have a direct effect on mortality from heart attacks (20). Heart disease is a major killer in the developed world.

Nitrite even benefits the skin. It reacts on the skin to release nitric oxide, and this may be a defense mechanism against skin infections (20). Nitrite mixed with an organic acid is an effective treatment for athlete's foot (Tinea pedis). Maybe we can see too why both animals and humans instinctively lick their wounds; the nitrite in saliva helps to prevent infection getting in. It also helps to protect our teeth from dental caries.

7.4 Methemoglobinemia

Methemoglobinemia (“blue-baby syndrome”) can kill infants less than one year old. The first report of the problem in the U.S. in 1945 referred to “well-water methemoglobinemia” (22), and all the cases identified in a 1991 book on the nitrate problem (23) were caused by water from wells, 98% of which were described as “privately dug.” Privately, and perhaps inexpertly, dug wells could have been too close to sewage tanks or manure heaps, and this may be why, in a number of reports, the water was polluted with bacteria as well as nitrate. The publisher of the original report (22) was well aware that both bacteria and nitrate were involved in the condition, but the possible role of the bacteria seems to have been played down subsequently. However, a recent American review (24) concludes that gastroenteritis may have been responsible for many cases of infantile methemoglobinemia previously attributed to nitrate in water. If gastroenteritis is the main factor in methemoglobinemia, nitrate may need to be seen as potentially beneficial, and clinical studies to resolve this question are urgently needed.

7.5 Diabetes

Not all medical researchers agree about the benefits of nitrate. A group in the U.K. (25) has found evidence in a statistically based study in part of Yorkshire that nitrate in drinking water is a factor in insulin-dependent diabetes (IDD) in young people. Their findings suggest that the threshold for the effect is 15 mg L−1 of nitrate. This is a worrying result because this threshold is about one-third of the U.S. limit for nitrate in potable water, and less than one-third of the E.C. limit. It is also not achievable in drainage water in many areas where there is arable agriculture, possibly even if nitrogen fertilizer is never applied again. This threshold, measured in Yorkshire, is already comprehensively exceeded in many areas of the U.K. and other countries, but without reports of increased incidence of IDD in young people, which suggests that it remains open to question.

Bibliography

  1. Top of page
  2. Introduction
  3. Nomenclature
  4. Physical Properties
  5. Agricultural Uses
  6. Chemistry
  7. Environmental Fate
  8. Nitrate and Health
  9. Bibliography
  10. Further Reading
  • 1
    G. C. Lowrison, in Fertilizer Technology, Ellis Horwood series in applied science and industrial technology, 1989, p. 139.
  • 2
    E. A. Paul and F. E. Clark, Soil Microbiology and Biochemistry, Academic Press, San Diego, CA, 1989.
  • 3
    R. E. White, Principles and Practice of Soil Science: The Soil as a Natural Resource, 3rd ed., Blackwell Science Ltd., Oxford, UK, 1997.
  • 4
    T. M. Addiscott, Soil Use Manage. 4: 9195 (1988).
  • 5
    D. S. Powlson, Proc. Fertil. Soc. 402: 141 (1997).
  • 6
    D. S. Jenkinson, Soil Use Manage. 6: 5661 (1990).
  • 7
    K. W. T. Goulding, Soil Use Manage. 6: 6163 (1990).
  • 8
    DOE, Report of the United Kingdom Review Group on Impacts of Atmospheric Nitrogen, Department of the Environment, London, 1994.
  • 9
    K. Mengel and E. A. Kirkby, Principles of Plant Nutrition, 3rd ed., International Potash Institute, Worblaufen-Bern, 1982, pp. 136138.
  • 10
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Further Reading

  1. Top of page
  2. Introduction
  3. Nomenclature
  4. Physical Properties
  5. Agricultural Uses
  6. Chemistry
  7. Environmental Fate
  8. Nitrate and Health
  9. Bibliography
  10. Further Reading
  • Stephenson, F. J., ed., Nitrogen in Agricultural Soils, American Society of Agronomy, Madison, WI, 1982.
  • Addiscott, T. M., Fertilizers and nitrate leaching, in R. E. Hester and R. M. Harrison, eds., Agricultural Chemicals and the Environment, Issues in Environmental Science and Technology, No. 5, The Royal Society of Chemistry, Cambridge, UK, 1996, pp. 126.
  • Spalding, R. F. and Exner, M. E., J. Environ. Qual. 22: 392402 (1993).
  • Wilson, W. S., Ball, A. S., and Hinton, R. H., eds., Managing the Risks of Nitrates to Humans and the Environment, Royal Society of Chemistry, Cambridge, UK, 1999.
  • Benjamin, N., Annales de Zootechnie 49: 207216.