This section covering phosphoric acid comprises these pages:
Phosphoric acid has become the most significant source of phosphate fertiliser production. World phosphoric acid production was 86 million tonnes in 2020, around 90% of which are used to produce fertilisers, mainly ammonium phosphates (DAP and MAP). There are two basic types of processes for the production of phosphoric acid; furnace processes and wet processes. Furnace processes include the blast-furnace process and the electric-furnace process. The blast furnace process has not been used commercially since 1938. The electric-furnace process is used extensively to make elemental phosphorus, most of which is converted to phosphoric acid for non-fertiliser uses. Since it is unlikely that furnace processes will be competitive for producing phosphoric acid for fertiliser use, except possibly in unusual circumstances, these processes are described only briefly.
Wet processes may be classified according to the acid used to decompose phosphate rock. Sulfuric, nitric, and hydrochloric acid are used in commercial processes. Processes using hydrochloric acid are not always competitive for fertiliser purposes, except under specific conditions. Processes using sulfuric acid are by far the most common means of producing phosphoric acid for fertiliser use (and sometimes for other uses).The scope of this resource precludes extensive detail of the production processes. For more detail, readers should consult Phosphoric Acid, edited by A. V. Slack , and other references listed below.
Chemistry of the wet process
The main chemical reaction in the wet (sulfuric acid) process may be represented by the following equation using pure fluorapatite to represent phosphate rock:
Ca10F2(PO4)6 + 10H2SO4 + 10nH2O à 10CaSO4nH2O + 6H3PO4 + 2HF
where n = 0, 1/2, or 2, depending on the hydrate form in which the calcium sulfate crystallises.
The reaction represents the net result of two stages. In the first stage, phosphoric acid reacts with the apatite forming monocalcium phosphate; in the second stage monocalcium phosphate reacts with sulfuric acid to form phosphoric acid and calcium sulfate. These two stages do not necessarily require two reaction vessels; they usually take place simultaneously in a single reactor.
Sign in or Register to proceed
The extensive Knowledge Base in FerTechInform is free to use, but we need to know something about its users. This information will help us further develop the resource and will not be shared with other organisations.
Phosphate rock contains many impurities, both in the apatite itself and in accessory minerals. Furthermore, as reserves of phosphate rock are utilised, the quality of what remains is decreasing. These impurities participate in numerous side reactions. Most phosphate rocks have a higher CaO:P2O5 ratio than pure fluorapatite, The additional CaO consumes more sulfuric acid and forms more calcium sulfate. The HF formed by the reaction reacts with silica and other impurities (Na, K, Mg, and Al) to form fluosilicates and other more complex compounds. A variable amount of the fluorine is volatilised as SiF4, HF, or both. The amount volatilised and the form depend on phosphate rock composition and process conditions.
As a result of side reactions, numerous impurity compounds (some of them very complex) are formed. For a complete discussion of the nature of impurities, see Phosphoric Acid by A. V. Slack .
Heat released in reaction
The reaction involved in producing phosphoric acid from fluorapatite and sulfuric acid by the dihydrate process may be represented by the following equation:
Ca10F2(PO4)6 (s) + 10H2SO4 (liq) + 20H2O (liq) à 10CaSO4•2H2O (s) + 2HF (aq) + 6H3PO4 (aq)
The heat of reaction may be calculated by using the heats of formation of reactants and products (Table 1) [3,4,5]. The heat of reaction so calculated is 256.94 kcal/gmol of apatite which is equivalent to 255 kcal/kg of apatite or about 600 kcal/kg of P2O5. The heat required to raise the temperature of the gypsum (Cp = 0.272) and phosphoric acid (30% P2O5; Cp = 0.703) from 25° to 82°C is calculated to be 197 kcal/kg of P2O5 [6,7]. Thus, about 403 kcal/kg of P2O5 remains to be dissipated, and most processes provide a means of removing the excess heat. In practice, some of the heat is lost by convection and conduction. On the other hand, some heat may be introduced by use of heated wash water, or if the wash water is not heated, some of the heat in the gypsum is transferred to the recycled weak acid and thus returned to the reaction. Additional heat will be generated by reaction of additional sulfuric acid with impurities in the rock. Most phosphate rock contains 10%- 20% more calcium than that required to form pure fluorapatite with the phosphorus in the rock, which may result from substitution of carbonate for phosphate in the apatite or presence of calcite or both. Reaction of this amount of calcium with sulfuric acid to form gypsum would increase the net heat of reaction per kilogram of P2O5 by about 11% -16%.
Hydrogen fluoride is shown as a product of reaction in the first equation. It reacts with the silica present as an impurity in phosphate rock to form fluosilicic acid which, in tum, forms fluosilicates and other compounds with impurities in the rock. The thermal effect of these reactions is negligible.
The net heat of reaction is influenced appreciably by the concentration of the sulfuric acid used as indicated in Table 2. If the conditions are such that the calcium sulfate crystallises in the form of anhydrite or hemihydrate rather than gypsum, the excess heat to be dissipated is about 100 kcal/kg of P2O5 less than the values given above.
Types of wet processes
Commercial wet processes may be classified according to the hydrate form in which the calcium sulfate crystallises:
Anhydrite – CaSO4
Hemihydrate – CaSO4•0.5H2O
Dihydrate – CaSO4•2H2O
The hydrate form is controlled mainly by temperature and acid concentration, as shown in Figure 1. This graph presents only an approximation of production features because sulfuric acid excess concentration and impurities also have an influence.
At present there is no commercial use of the anhydrite process, mainly because the required reaction temperature is high enough to cause severe corrosion difficulties. Processes in commercial use are given in Table 3. From the very beginning, straight dihydrate processes have been by far the most popular worldwide because they are relatively simple and adaptable to a wide range of grades and types of phosphate rock.
Hemihydrate processes have the significant advantage of producing phosphoric acid with a relatively high concentration without using any concentration step. There is also some interest in two-staqe processes that increase the level of recovery of P2O5. These involve crystallisation in the hemihydrate form followed by re-crystallisation in the dihydrate form (or vice versa), with or without intermediate separation by filtration or centrifuging.
The various types of wet process can be categorised according to:
- The form of gypsum crystal produced.
- The number of crystallisation stages and filtrations – single or double.
- Relative process efficiency
- Strength of acid produced.
The resulting categorisation, along with the features of each process type, are shown in Table 4.
2. Phosphoric Acid. 1968. AV. Slack (Ed.), Marcel Dekker, Inc., New York, NY, U.S.A.
3. Farr, Thad D., and Kelly L. Elmore. 1962. ‘System CaO-P205-HF-H20: Thermodynamic Properties,’ Journal of Physical Chemistry, 66(2):315-318.
4. Rossini, Frederick D., et al. 1952. ‘Selected Values of Chemical Thermodynamic Properties,’ National Bureau of Standards, Circular No. 500, U.S. Department of Commerce, Washington, D. C., U.S. A.
5. Egan, E. P., Jr., and B. B. Luff. 1961. ‘Heat of Solution of Orthophosphoric Acid,’ Journal of Physical Chemistry, 65(3):523-526.
6. Kelly, K. K. 1934. ‘Contributions to the Data on Theoretical Metallurgy. II. High-Temperature Specific-Heat Equations for Inorganic Substances,’ Bureau of Mines Bulletin 371, U.S. Department of Commerce, U.S. Government Printing Office, Washington, D.C. U.S.A
Links to Related IFS Proceedings
67, (1961), Developments in Phosphoric Acid Manufacture, W C Weber, Frank W Edwards
81, (1964), Insoluble Phosphate Losses in Phosphoric Acid Manufacture by the Wet Process: Theory and Experimental Techniques, S M Janikowski, N Robinson, W F Sheldrick
112, (1969), New Phosphoric Acid Processes – Symposium, S M Janikowski, N Robinson
165, (1977), Phosphoric Acid Techniques to Match Raw Materials and Fertiliser Trends, A Davister
362, (1995), Forty Years with the Phosphate Industry – Nineteenth Francis New Memorial Lecture, A Davister
364, (1995), Partially Acidulated Phosphates – Production, Agronomic and Environmental Aspects, Y Pelovski, M K Garrett
492, (2002), Safety Legislation and the Fertiliser Industry, K D Shah
668, (2010), The Phosphate Life-Cycle: Rethinking the Options for a Finite Resource, J Hilton, A E Johnston, C J Dawson
726, (2013), New Process Route to Phosphoric Acid, D Fati, T Theys, O Schrevens
806, (2017), Capturing phosphoric acid know-how in a training simulator, A Durand and S Joao
821, (2018), Approaches to improving the quality of phosphoric acid, T Henry
Links to external sources
Becker, P. (1989) Phosphates and Phosphoric Acid: Raw Materials: Technology, and Economics of the Wet Process. Marcel Dekker, Inc., New York, NY, U.S.A.
Havelange, S. et al. (2022). Phosphoric Acid and Phosphates in Ullmann’s Encyclopaedia of Industrial Chemistry.
Need more information?
If the information you need on this topic is not on this page, use this button to access the relevant section of the forum, where this may have been provided. If not, you may ask the question.