Phosphorus Fertilizers And Eutrophication

In solution, the mono- and dihydrogen phosphate ions, HPO4" and H2PO4 are the predominant forms of phosphate at the usual pH values (since H3PO4 is a strong acid), and it is in these forms that phosphate is taken up by organisms. The orthophosphate ion PO3" will exist in significant concentration in solution only at very high pH values.

Polyphosphates can be formed from the heat-induced condensation polymerization of simple orthophosphate units, for example,

which occurs readily with pure orthophosphoric acid. The product here is pyrophosphoric acid, which is made of two tetrahedra with a shared oxygen making a common corner (Figure 10-10).


OO FIGURE 10-10 Pyrophosphoric acid.

Many polyphosphate compounds can be prepared by heating simpler phosphates in reactions analogous to reaction (10-16). Other polyphosphates may be cyclic such as trimetaphosphate (Figure 10-11).

The linear tripolyphosphate used in detergents was discussed in Section 9.5.6. Other condensed phosphates that have been used include pyrophosphate, tetraphosphate, and hexametaphosphate. In dilute solutions, polyphosphates are hydrolyzed to the orthophosphate, although the reactions are not always fast. Inorganic phosphate concentration is limited by solubility, since many metal phosphates are insoluble. For example, Ksp for FePO4 is 10~23. Aluminum phosphate, AlPO4, and hydroxyapatite, Ca5OH(PO4)3, are also very insoluble. Ligands that form strong complexes with the metal, or protonation of the phosphate ions, can increase the effective solubility, however. Typically, maximum availability of phosphate in soils is around pH 7.

Natural waters may contain organophosphorus compounds that can make up a significant fraction of the total soluble phosphorus content. These compounds are generally of unknown composition and are derived from biological products or possibly P-containing insecticides such as the phosphate esters discussed in Chapter 8, or phosphono compounds such as phosphonomethyl-glycine [Glyphosate: (OH)2P(O)CH2NHCH2CO2H] used as a herbicide.

Although phosphorus compounds in a variety of other oxidation states are known in chemistry [especially P(-III) in PH3 and derivatives, P(III) in phosphorous acid, PQ3 etc.], the inorganic forms of these oxidation states have no general importance under environmental conditions. The natural phosphorus cycle consists of weathering and leaching of phosphate from rocks and soils, and runoff to the oceans, which serve as a sink. Biological cycling intervenes in

FIGURE 10-11 Cyclic trimetaphosphate.

this process, but there are no atmospheric steps except as particulates. The low phosphate levels in water systems are increased easily by anthropogenic processes such as the use of fertilizers and phosphate-containing detergents. Runoff from cattle feedlots is another important source of environmental phosphate as well as nitrogen compounds.10

Phosphorus is an essential nutrient material for plants and animals, being required in biological synthesis and energy transfer processes. The overall photosynthesis reaction in aquatic organisms (70% of all photosynthesis takes place in the ocean) results in the eventual formation of biological material that has an overall C:N:P ratio of approximately 100:16:1. The N:P ratio in ocean water is about 13:1 (Table 11.1). In freshwater lakes, the ratios vary considerably, but N is usually in excess; typical ranges are N, 10-4 to 10-5 M; P, 3 x 10-5 to 3 x 10-7 M. (There is evidence that natural processes act to make phosphorus the most important algal nutrient in any event.11) One milligram of phosphorus results in production of 0.1 g of organic material if N is in excess, and in systems where phosphorus is the limiting nutrient, additional phosphorus input can greatly increase the biological yield. The 0.1 g of organic material produced requires more than 0.1 g of oxygen for decomposition, and if this is not replaced, the water becomes depleted of dissolved oxygen, with consequent decrease in the abundance of the higher life-forms and production of undesirable products of anaerobic reactions: NH3, H2S, and CH4. The enhanced growth and decay of algae caused by increased phosphorus levels is the cause of the concern over phosphate pollution and its effects on eutrophication processes in freshwater lakes (Section 9.1). Large anoxic regions in the oceans, for example, the Gulf of Mexico, the Black Sea, and the Kattegat Strait, have been linked to agricultural fertilizer use, but nitrate rather than phosphorus.12

Phosphate fertilizers are produced commercially from insoluble phosphate rock, which has the formula Ca2(PO4)2CaX, where X can be CO2-,2(OH)-, and others, but is usually 2F-. Soluble phosphates are produced by an acid displacement reaction:

Ca3(PO4)2 • CaX + 3H2SO4 ^ Ca(^PO4)2 + 3CaSO4 + H2X (10-17)

If fluoride is present, the HF produced13 will react with silica in the rock to produce SiF4 and fluorosilicate salts (M2SiF6). Residual fluoride impurities may remain in the product, along with other elements such as As or Cd that may be present in the ore.

10M. A. Mallin, Impacts of industrial animal production on rivers and estuaries, Am. Sci., 88, 26 (2000).

11D. W. Schindler, Evolution of phosphorus limitation in lakes, Science, 195, 260 (1977).

12D. Ferber, Keeping the Stygian waters at bay, Science, 291, 968 (2001).

13If X = 2F-, H2X in reaction (10-17) represents 2HF.

The Ca(H2P04)2— CaS04 mixture is often sold as such as "superphosphate." The sulfate itself is a useful component, since sulfur is also an important nutrient. A higher phosphate content can be obtained in so-called triple superphosphate, produced by the process

Ca3(P04)2 • CaX + 6H3PO4 ^ 4Ca(H2P04)2 + H2X (10-18)

Treatment of superphosphates with ammonia solution produces ammoniated phosphate fertilizers. Phosphoric acid is formed from excess H2S04 on Ca3(P04)2, and can serve as the source material for other phosphate chemicals through neutralization and condensation reactions.

As with nitrogen, phosphate fertilizers are subject to leaching from soils, adding to the nutrient content of runoff water and lakes, and contributing to eutrophication. Because of its higher charge, however, phosphate is held more strongly than nitrate by ion exchange, as well as through the formation of insoluble phosphate salts.

Besides phosphorus and nitrogen, a third major component of fertilizers is a soluble salt of potassium. Most common lawn fertilizers consist of a mixture of these three materials, and the composition specifications (e.g., 6.10.6) refer to the percentage composition by weight in terms of N, P20s, and K20.

Figure 10-12 shows the use of commercial fertilizers over time; the large increase in nitrogen fertilizer use since 1960 is notable. Fertilizer use

Eutrophication Fertilizers








FIGURE 10-12 Global fertilizer use. Used by permission of International Fertilizer Industry Association (IFA);

worldwide has decreased in recent years. European use began to decrease in the 1970s for economic reasons and with changing agricultural practices; there has been a large decrease in use in the former Soviet Union, but demand in the United States and developing countries continues to increase.

In addition to phosphorus, potassium, and nitrogen, other fertilizer elements are needed in particular cases. Important examples of these are calcium, magnesium, and sulfur. Special fertilizers are produced to meet the need of some crops for available trace metals, such as zinc. Unfortunately, in some cases waste materials such as dusts from cement kilns or steel recycling furnaces are used in the production of these fertilizers. These may include other, toxic, heavy metals including cadmium, lead, and mercury which may further add to the buildup of these materials in soils. At this time, there are no national standards for heavy metal content of fertilizers in the United States although there is in some other nations.14

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