Phosphate
A phosphate is a salt that is prepared from orthophosphoric acid. Phosphates are usually sodium salts but can also include potassium, calcium, magnesium, and ammonium salts as well. The phosphates have a wide range of applications in the chemical industry. The mono-, di-, and tri-metallic salts of phosphoric acid, for example, are used extensively in many industrial applications. Polyphosphates are a group of phosphates that are created from polymeric chemical compounds. These compounds are produced from the condensation reactions of several phosphoric acid molecules. The chemical formula for the phosphate ion is PO43-. The phosphate ion is a polyatomic ion, in other words, a group of covalently bonded atoms that together carry an electrical charge. Polyatomic ions can combine with oppositely charged ions, through ionic bonding. When they do so, they form ionic compounds.
Because the phosphates have many industrial applications, they are produced in great quantities by the chemical industry. Phosphoric acid (HPO3) is required in the production of phosphates, as well as sodium, potassium, and calcium hydroxides [NaOH, KOH, and Ca(OH)2]. Phosphates can exist in a variety of hydrated forms, each with a different stability. The two most stable species of phosphates are the two extremes of hydrates, the anhydrous salt and the fully hydrated salt. These are the two phosphate species that are produced in the greatest quantities, although phosphates of every degree of hydration are available. Many grades of phosphates are also commercially available in terms of purity, from low-grade to "technical" and "for use in foods."
The production of phosphates is highly automated. The pH of the reaction vessel must be kept within certain specific ranges, depending on the properties and chemical nature of the phosphate being produced. The pH values correspond to the various stages of the neutralization of phosphoric acid involved in the production of phosphates. The pH is adjusted by adding alkali or ammonia (to raise the pH) as necessary and then crystallizing out the salt at the optimal temperature for crystallization. For example, both disodium and dipotassium hydrogen phosphates are prepared in a solution with a pH value of 8.5, attained by the addition of alkali to phosphoric acid. If the anhydrous form is being prepared, the solution is then concentrated to remove the water and crystallized. The crystals are separated from any remaining solution by centrifugation. In general, the production of sodium phosphates is much less expensive than the potassium phosphates. As a result, more sodium phosphates are produced by industry. Potassium phosphates are produced in quantities when they are specifically needed, for example, when working with bacterial cultures the potassium ion is much more biologically compatible than the sodium ion.
Phosphates are used in a wide range of applications. Phosphates can act as buffers to control the pH values of various solutions. They are also used in water treatment and the production of sugar solutions, the etching of metals, in polymerization processes, and in the textile industry. Phosphates can aid in the emulsification and protection of colloidal suspensions for the food industry as well. Certain phosphates, monoammonium dihydrogen phosphate and diammonium hydrogen phosphate, in particular, evolve non-flammable gases when heated and are used as flame retardants. These same two phosphates are also used as nutrients for cultures of microorganisms. Sodium, potassium, and calcium phosphates are all used in the chemical and pharmaceutical industries. Condensed phosphates, or polyphosphates, are the choice for water treatment and detergents because of their ability to form strong complexes with the unwanted ions in hard water.
Polyphosphates are prepared differently from other phosphates. Orthophosphoric acid is sometimes heated to dehydration to form polyphosphoric acids, which are then neutralized to form polyphosphates. More commonly, the salts of orthophosphoric acid are simply heated to form the polyphosphates. There are several specific polyphosphates which are produced in the chemical industry. For example, the tetra-alkali metal ion dimers sodium pyrophosphate (Na4P2O7) and potassium pyrophosphate (K4P2O7), which are used to make liquid detergents, to etch metal, to stabilize emulsions, and as food additives. Another polyphosphate that is produced commercially is the dimeric disodium salt, disodium dihydrogen pyrophosphate (Na2H2P2O 7 ), which is used in both the food and mining industries. Another group of polyphosphates which are useful industrially are the ultraphosphates, which have the general chemical formula (MePO3)n. Sodium hexametaphosphate (n=6) is one of the most important of the ultraphosphates, used in both detergents and dispersing agents. Another polyphosphate which is used in detergents is sodium tripolyphosphate, Na5P3O10. This polyphosphate acts as a support for silicates and bleaching agents in the detergents.
Phosphates are just as important biologically as they are industrially. The phosphate buffer system in the cytoplasm of our cells is extremely important in maintaining the pH of the cytoplasm between 6.4 and 7.4. The phosphate buffer system consists of the hydrogen phosphate ion (HPO42-) and the dihydrogen phosphate ion (H2PO4-). The DNA in our cells also relies on phosphates. Phosphate groups make up a significant part of the DNA backbone and provide the sites where H+ ion transfer can occur.
One of the most important applications of phosphates, biologically and economically, is their use in fertilizers. Until the middle of the nineteenth century, when the use of phosphates in fertilizers began, the amount of phosphorus compounds that could be used by plants was very low. The phosphorus content of soils was so low, in fact, that the crop yields were declining. Two reasons the phosphorus levels in the soil were so low is the fact that phosphorus only makes up about 0.1% of the Earth's crust, and naturally occurring phosphorus compounds, such as those found in manure, are generally insoluble in aqueous solutions. The development of phosphate fertilizers increased the yield of crops. The extra phosphorus available in the soil gave rise to stronger, healthier plants.
With these benefits of phosphates have come problems. Phosphates that leach out of the soil and drain into lakes, streams and ponds are causing major ecological disturbances. A build-up of phosphates in stagnant waters causes the eutrophication of the waters. Phosphates from fertilizers are not the only cause, either. Recent improvements in detergents have called for the addition of more phosphates which also find themselves in water systems. Legislation has been passed which requires the amounts of phosphates in effluents to be maintained below a maximum level. Research efforts have been established to find alternatives to phosphates, especially in detergents. Some possible alternatives include trisodium citrates, zeolites, or nitrilotriacetic acid (NTA), all of which could also cause hazardous health effects in large amounts. Despite these problems associated with phosphates, however, they are still considered some of the most important compounds in the chemical industry and the benefits of their use far outweigh their hazards.
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