BETA ZEN
Ammonium sulfate
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| Names | |
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| IUPAC name
Ammonium sulfate | |
Other names
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| Identifiers | |
3D model (JSmol) |
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| Abbreviations | AS |
| ChEBI | |
| ChemSpider | |
| ECHA InfoCard | 100.029.076 |
| EC Number |
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| E number | E517 (acidity regulators, ...) |
| KEGG | |
PubChem CID |
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| UNII | |
CompTox Dashboard (EPA) |
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| Properties[1] | |
| (NH4)2SO4 | |
| Molar mass | 132.13 g·mol−1 |
| Appearance | White or brown orthorhombic crystals |
| Density | 1.77 g/cm3 |
| Melting point | 280 °C (536 °F; 553 K) (decomposes) |
| 76.4 g/100g | |
| −67.0×10−6 cm3/mol | |
| 79.2% (30 °C (86 °F)) | |
| Structure | |
| orthorhombic | |
| Thermochemistry[1] | |
Heat capacity (C) |
187.5 J⋅mol−1⋅K−1 |
Std molar entropy (S⦵298) |
220.1 J⋅mol−1⋅K−1 |
Std enthalpy of formation (ΔfH⦵298) |
−1180.9 kJ⋅mol−1 |
Gibbs free energy (ΔfG⦵) |
−901.7 kJ⋅mol−1 |
| Hazards | |
| GHS labelling:[2] | |
| H402 | |
| P273, P501 | |
| NFPA 704 (fire diamond) | |
| Lethal dose or concentration (LD, LC): | |
LD50 (median dose) |
2840 mg/kg, rat (oral) |
| Related compounds | |
Other anions |
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Other cations |
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Related compounds |
Ammonium iron(II) sulfate |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Ammonium sulfate (American English and international scientific usage; ammonium sulphate in British English); [NH4]2SO4, is an inorganic salt with a number of commercial uses. The most common use is as a soil fertilizer. It contains 21% nitrogen and 24% sulfur.
Uses
Agriculture
The primary use of ammonium sulfate is as a fertilizer for alkaline soils. In the soil, the ammonium ion is released and forms a small amount of acid, lowering the pH balance of the soil, while contributing essential nitrogen for plant growth. One disadvantage to the use of ammonium sulfate is its low nitrogen content relative to ammonium nitrate or urea, which elevates transportation costs.[4]
It is also used as an agricultural spray adjuvant for water-soluble insecticides, herbicides, and fungicides. There, it functions to bind iron and calcium cations that are present in both well water and plant cells. It is particularly effective as an adjuvant for 2,4-D (amine), glyphosate, and glufosinate herbicides.[citation needed]
Laboratory use
Ammonium sulfate precipitation is a common method for protein purification by precipitation. As the ionic strength of a solution increases, the solubility of proteins in that solution decreases. Being extremely soluble in water, ammonium sulfate can "salt out" (precipitate) proteins from aqueous solutions.[5][6] Ammonium sulfate precipitation provides a convenient and simple means to fractionate complex protein mixtures.[7]
In the analysis of rubber lattices, volatile fatty acids are analyzed by precipitating rubber with a 35% ammonium sulfate solution, which leaves a clear liquid from which volatile fatty acids are regenerated with sulfuric acid and then distilled with steam. Selective precipitation with ammonium sulfate, opposite to the usual precipitation technique which uses acetic acid, does not interfere with the determination of volatile fatty acids.[8][full citation needed]
Food additive
As a food additive, ammonium sulfate is considered generally recognized as safe (GRAS) by the U.S. Food and Drug Administration, and in the European Union it is designated by the E number E517.[9][10] It is used as an yeast nutrient in breads.[11]
Other uses
Ammonium sulfate is a precursor to other ammonium salts, especially ammonium persulfate ((NH4)2S2O8.
Ammonium sulfate is listed as an ingredient for many vaccines by the Centers for Disease Control.[12]
Ammonium sulfate has also been used in flame retardant compositions acting much like diammonium phosphate. As a flame retardant, it increases the combustion temperature of the material, decreases maximum weight loss rates, and causes an increase in the production of residue or char.[13][better source needed]
Preparation
Ammonium sulfate is made by treating ammonia with sulfuric acid:
- [NH4]2SO4 "}},"i":0}}]}'/>2 NH3 + H2SO4 → [NH4]2SO4
A mixture of ammonia gas and water vapor is introduced into a reactor that contains a saturated solution of ammonium sulfate and about 2% to 4% of free sulfuric acid at 60 °C (140 °F). Concentrated sulfuric acid is added to keep the solution acidic, and to retain its level of free acid. The heat of reaction keeps reactor temperature at 60 °C (140 °F). Dry, powdered ammonium sulfate may be formed by spraying sulfuric acid into a reaction chamber filled with ammonia gas. The heat of reaction evaporates all water present in the system, forming a powdery salt. Approximately 6,000 million tons were produced in 1981.[4]
Ammonium sulfate also is manufactured from gypsum (CaSO4·2H2O). Finely divided gypsum is added to an ammonium carbonate solution. Calcium carbonate precipitates as a solid, leaving ammonium sulfate in the solution.
- [NH4]2SO4 + CaCO3 "}},"i":0}}]}'/>[NH4]2CO3 + CaSO4 → [NH4]2SO4 + CaCO3
Ammonium sulfate occurs naturally as the rare mineral mascagnite in volcanic fumaroles and due to coal fires on some dumps.[14]
Ammonium sulfate is a byproduct in the production of methyl methacrylate (CH2=C(CH3)COOCH3).[15]
Properties
Ammonium sulfate becomes ferroelectric at temperatures below −49.5 °C (−57.1 °F). At room temperature it crystallises in the orthorhombic system, with cell sizes of a = 7.729 Å, b = 10.560 Å, c = 5.951 Å. When chilled into the ferrorelectric state, the symmetry of the crystal changes to space group Pna21.[16][better source needed]
Reactions
Ammonium sulfate decomposes upon heating above 250 °C (482 °F), first forming ammonium bisulfate. Heating at higher temperatures results in decomposition into ammonia, nitrogen, sulfur dioxide, and water.[17]
As a salt of a strong acid (H2SO4) and weak base (NH3), its solution is acidic; the pH of 0.1 M solution is 5.5. In aqueous solution the reactions are those of [NH4]+ and SO2−4 ions. For example, addition of barium chloride, precipitates out barium sulfate. The filtrate on evaporation yields ammonium chloride (NH4Cl).
Ammonium sulfate forms many double salts (ammonium metal sulfates) when its solution is mixed with equimolar solutions of metal sulfates and the solution is slowly evaporated. With trivalent metal ions, alums such as ferric ammonium sulfate (NH4Fe(SO4)2·12H2O) are formed. Double metal sulfates include ammonium cobaltous sulfate, ferrous diammonium sulfate ((NH4)2SO4·Fe(SO4)·6H2O), ammonium nickel sulfate which are known as Tutton's salts (M2M′(SO4)2(H2O)6) and ammonium ceric sulfate ((NH4)4Ce(SO4)4·2H2O).[4]
Anhydrous double sulfates of ammonium also occur in the Langbeinites family.
Airborne particles of evaporated ammonium sulfate comprise approximately 30% of fine particulate pollution worldwide.[18][relevant?]
It reacts with additional sulfuric acid to give triammonium hydrogen disulphate, (NH4)3H(SO4)2.[19]
See also
References
- 1 2 Haynes, William M., ed. (2016). CRC Handbook of Chemistry and Physics (97th ed.). Boca Raton, Florida: CRC Press. pp. 4–47, 5–4. ISBN 9781498754293.
- ↑ Sigma-Aldrich Co., Ammonium sulfate. Retrieved on 26 August 2026.
- ↑ "SDS - Ammonium sulfate" (pdf). www.fishersci.com. Rev. 6. ThermoFisher Scientific. 18 December 2025 [11 September 2009]. Retrieved 26 August 2026.
- 1 2 3 Zapp, Karl-Heinz (2012). "Ammonium Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a02_243. ISBN 978-3-527-30673-2.
- ↑ Duong-Ly, Krisna C.; Gabelli, Sandra B. (2014-01-01). "Salting out of Proteins Using Ammonium Sulfate Precipitation". In Lorsch, Jon (ed.). Methods in Enzymology. Laboratory Methods in Enzymology: Protein Part C. Vol. 541. Academic Press. pp. 85–94. doi:10.1016/B978-0-12-420119-4.00007-0. ISBN 9780124201194. PMID 24674064.
- ↑ Wingfield, Paul T. (2017-05-05). "Protein Precipitation Using Ammonium Sulfate". Current Protocols in Protein Science. 13 (1): A.3F.1–8. doi:10.1002/0471140864.psa03fs13. ISSN 1934-3655. PMC 4817497. PMID 18429073.
- ↑ "Ammonium Sulfate Calculator". EnCor Biotechnology Inc. 2013. Archived from the original on January 26, 2016. Retrieved March 2, 2013.
- ↑ ASTM Standard Specification for Rubber Concentrates D 1076-06
- ↑ "Select Committee on GRAS Substances (SCOGS) Opinion: Ammonium sulfate". U.S. Food and Drug Administration. August 16, 2011. Archived from the original on February 11, 2012. Retrieved March 2, 2013.
- ↑ "Food and Feed Information Portal Database | Ammonium Sulphate (E517)". ec.europa.eu. European Commission. Retrieved 26 August 2026.
- ↑ "Official Subway Restaurants U.S. Products Ingredients Guide". Archived from the original on August 14, 2011. Retrieved March 2, 2013.
- ↑ "Vaccine Excipient & Media Summary" (PDF). Centers for Disease Control and Prevention (CDC). February 2012. Archived from the original (PDF) on February 5, 2011. Retrieved March 2, 2013.
- ↑ George, C. W.; Susott, R. A. (April 1971). "Effects of Ammonium Phosphate and Sulfate on the Pyrolysis and Combustion of Cellulose". Research Paper INT-90. Intermountain Forest and Range Experiment Station: USDA Forest Service. OL 16022833M.
- ↑ "Mascagnite". mindat.org. Retrieved 2 March 2013.
- ↑ Bauer, Jr., William (2002). "Methacrylic Acid and Derivatives". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a16_441. ISBN 978-3-527-30673-2..
- ↑ Okaya, Y.; K. Vedam; R. Pepinsky (1958). "Non-isomorphism of ferroelectric phases of ammonium sulfate and ammonium fluoberyllate". Acta Crystallographica. 11 (4): 307. Bibcode:1958AcCry..11..307O. doi:10.1107/s0365110x58000803. ISSN 0365-110X.
- ↑ Liu Ke-wei, Chen Tian-lang (2002). "Studies on the thermal decomposition of ammonium sulfate". Chemical Research and Application (in Chinese). 14 (6). doi:10.3969/j.issn.1004-1656.2002.06.038.
- ↑ "Where Does Air Pollution Come From?". www.purakamasks.com. 2019-02-15. Archived from the original on 2019-02-20. Retrieved 2019-02-20.
- ↑ 4</sub>)<sub>3</sub>H(SO<sub>4</sub>)<sub>2</sub>"},"date":{"wt":"1978"},"last1":{"wt":"Suzuki"},"first1":{"wt":"S."},"last2":{"wt":"Makita"},"first2":{"wt":"Y."},"journal":{"wt":"Acta Crystallographica Section B Structural Crystallography and Crystal Chemistry"},"volume":{"wt":"34"},"issue":{"wt":"3"},"pages":{"wt":"732–735"},"bibcode":{"wt":"1978AcCrB..34..732S"}},"i":0}}]}'/>Suzuki, S.; Makita, Y. (1978). "The crystal structure of Triammonium hydrogen Disulphate, (NH4)3H(SO4)2". Acta Crystallographica Section B Structural Crystallography and Crystal Chemistry. 34 (3): 732–735. Bibcode:1978AcCrB..34..732S. doi:10.1107/S0567740878003969.
External links
- Clegg, Simon. "Surface tension calculator for aqueous solutions". www.aim.env.uea.ac.uk. University of East Anglia. Retrieved 26 August 2026.
- Clegg, Simon. "Densities of H+
- NH+
4 - Na+
- SO2−
4 - NO−
3 - Cl−
- H
2O". www.aim.env.uea.ac.uk. University of East Anglia. Retrieved 26 August 2026.
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