Ammonium sulfate
Ammonium sulfate Lewis structure
Ball-and-stick model of two ammonium cations and one sulfate anion
IUPAC name
Ammonium sulfate
Other names
  • Ammonium sulphate
  • Ammonium sulfate (2:1)
  • Diammonium sulfate
  • Sulfuric acid diammonium salt
  • Mascagnite
  • Actamaster
  • Dolamin
3D model (JSmol)
ECHA InfoCard 100.029.076 Edit this at Wikidata
EC Number
  • 231-984-1
E number E517 (acidity regulators, ...)
  • InChI=1S/2H3N.H2O4S/c;;1-5(2,3)4/h2*1H3;(H2,1,2,3,4) checkY
  • InChI=1/2H3N.H2O4S/c;;1-5(2,3)4/h2*1H3;(H2,1,2,3,4)
  • O=S(=O)(O)O.N.N
Molar mass 132.14 g/mol
Appearance Fine white hygroscopic granules or crystals
Density 1.77 g/cm3
Melting point 235 to 280 °C (455 to 536 °F; 508 to 553 K) (decomposes)
70.6 g per 100 g water (0 °C)
74.4 g per 100 g water (20 °C)
103.8 g per 100 g water (100 °C)[1]
Solubility Insoluble in acetone, alcohol and ether
−67.0×10−6 cm3/mol
79.2% (30 °C)
GHS labelling:
GHS07: Exclamation markGHS09: Environmental hazard
H315, H319, H335
P261, P264, P270, P271, P273, P280, P301+P312, P302+P352, P304+P340, P305+P351+P338, P312, P321, P330, P332+P313, P337+P313, P362, P391, P403+P233, P405, P501
NFPA 704 (fire diamond)
NFPA 704 four-colored diamondHealth 2: Intense or continued but not chronic exposure could cause temporary incapacitation or possible residual injury. E.g. chloroformFlammability 1: Must be pre-heated before ignition can occur. Flash point over 93 °C (200 °F). E.g. canola oilInstability 0: Normally stable, even under fire exposure conditions, and is not reactive with water. E.g. liquid nitrogenSpecial hazards (white): no code
Flash point Non-flammable
Lethal dose or concentration (LD, LC):
2840 mg/kg, rat (oral)
Related compounds
Other anions
Ammonium thiosulfate
Ammonium sulfite
Ammonium bisulfate
Ammonium persulfate
Other cations
Sodium sulfate
Potassium sulfate
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).
checkY verify (what is checkY☒N ?)

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.


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. The main disadvantage to the use of ammonium sulfate is its low nitrogen content relative to ammonium nitrate, which elevates transportation costs.[2]

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.

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. Ammonium sulfate is extremely soluble in water due to its ionic nature, therefore it can "salt out" proteins by precipitation.[3] Due to the high dielectric constant of water, the dissociated salt ions being cationic ammonium and anionic sulfate are readily solvated within hydration shells of water molecules. The significance of this substance in the purification of compounds stems from its ability to become more so hydrated compared to relatively more nonpolar molecules and so the desirable nonpolar molecules coalesce and precipitate out of the solution in a concentrated form. This method is called salting out and necessitates the use of high salt concentrations that can reliably dissolve in the aqueous mixture. The percentage of the salt used is in comparison to the maximal concentration of the salt in the mixture can dissolve. As such, although high concentrations are needed for the method to work adding an abundance of the salt, over 100%, can also oversaturate the solution, therefore, contaminating the nonpolar precipitate with salt precipitate.[4] A high salt concentration, which can be achieved by adding or increasing the concentration of ammonium sulfate in a solution, enables protein separation based on a decrease in protein solubility; this separation may be achieved by centrifugation. Precipitation by ammonium sulfate is a result of a reduction in solubility rather than protein denaturation, thus the precipitated protein can be solubilized through the use of standard buffers.[5] Ammonium sulfate precipitation provides a convenient and simple means to fractionate complex protein mixtures.[6]

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.[7]

Food additive

As a food additive, ammonium sulfate is considered generally recognized as safe (GRAS) by the U.S. Food and Drug Administration,[8] and in the European Union it is designated by the E number E517. It is used as an acidity regulator in flours and breads.[9][10][11]

Other uses

In the treatment of drinking water, ammonium sulfate is used in combination with chlorine to generate monochloramine for disinfection.[12]

Ammonium sulfate is used on a small scale in the preparation of other ammonium salts, especially ammonium persulfate.

Ammonium sulfate is listed as an ingredient for many United States vaccines per the Centers for Disease Control.[13]

A saturated solution of ammonium sulfate in heavy water (D2O) is used as an external standard in sulfur (33S) NMR spectroscopy with shift value of δ = 0 ppm.

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.[14] Its flame retardant efficacy can be enhanced by blending it with ammonium sulfamate.[citation needed] It has been used in aerial firefighting.

Ammonium sulfate has been used as a wood preservative, but due to its hygroscopic nature, this use has been largely discontinued because of associated problems with metal fastener corrosion, dimensional instability, and finish failures.


Ammonium sulfate is made by treating ammonia with sulfuric acid:

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. 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. 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.[2]

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)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.[15]


Ammonium sulfate becomes ferroelectric at temperatures below –49.5 °C. 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]


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 NH+
and SO2−
ions. For example, addition of barium chloride, precipitates out barium sulfate. The filtrate on evaporation yields ammonium chloride.

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 are formed. Double metal sulfates include ammonium cobaltous sulfate, ferrous diammonium sulfate, ammonium nickel sulfate which are known as Tutton's salts and ammonium ceric sulfate.[2] Anhydrous double sulfates of ammonium also occur in the Langbeinites family. The ammonia produced has a pungent smell and is toxic.

Airborne particles of evaporated ammonium sulfate comprise approximately 30% of fine particulate pollution worldwide.[18]

Legislation and control

In November 2009, a ban on ammonium sulfate, ammonium nitrate and calcium ammonium nitrate fertilizers was imposed in the former Malakand Division—comprising the Upper Dir, Lower Dir, Swat, Chitral and Malakand districts of the North West Frontier Province (NWFP) of Pakistan, by the NWFP government, following reports that they were used by militants to make explosives. In January 2010, these substances were also banned in Afghanistan for the same reason.[19]

See also


  1. ^ Lide, David R., ed. (2006). CRC Handbook of Chemistry and Physics (87th ed.). Boca Raton, FL: CRC Press. ISBN 0-8493-0487-3.
  2. ^ a b c Zapp, Karl-Heinz (2012). "Ammonium Compounds". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a02_243. ISBN 9783527303854.
  3. ^ 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.
  4. ^ Duong-Ly, Krisna C.; Gabelli, Sandra B. (2014-01-01). "Salting out of Proteins Using Ammonium Sulfate Precipitation". Laboratory Methods in Enzymology: Protein Part C. Vol. 541. pp. 85–94. doi:10.1016/B978-0-12-420119-4.00007-0. ISBN 9780124201194. ISSN 1557-7988. PMID 24674064.
  5. ^ 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. ISBN 978-0471140863. ISSN 1934-3655. PMC 4817497. PMID 18429073.
  6. ^ "Ammonium Sulfate Calculator". EnCor Biotechnology Inc. 2013. Archived from the original on January 26, 2016. Retrieved March 2, 2013.
  7. ^ ASTM Standard Specification for Rubber Concentrates D 1076-06
  8. ^ "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.
  9. ^ "Panera Bread: Menu & Nutrition: Nutrition Information Profile". Archived from the original on August 19, 2009. Retrieved March 2, 2013.
  10. ^ "Official Subway Restaurants U.S. Products Ingredients Guide". Archived from the original on August 14, 2011. Retrieved March 2, 2013.
  11. ^ Sarah Klein (May 14, 2012). "Gross Ingredients In Processed Foods". The Huffington Post. Archived from the original on May 18, 2012. Retrieved March 2, 2013.
  12. ^ McCool, Pat. "Feeding Ammonium Sulfate to Form Combined Chlorine Residual" (PDF). The Kansas Lifeline. Kansas Rural Water Association. Archived (PDF) from the original on 6 June 2019. Retrieved 6 June 2019.
  13. ^ "Vaccine Excipient & Media Summary" (PDF). Centers for Disease Control and Prevention (CDC). February 2012. Archived (PDF) from the original on February 5, 2011. Retrieved March 2, 2013.
  14. ^ 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.
  15. ^ "Mascagnite". Mindat. Archived from the original on January 19, 2013. Retrieved March 2, 2013.
  16. ^ 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.
  17. ^ 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.
  18. ^ "Where Does Air Pollution Come From?". 2019-02-15. Archived from the original on 2019-02-20. Retrieved 2019-02-20.
  19. ^ "PAKISTAN: 'Anti-terrorist' fertilizer ban hinders farmers". IRIN Humanitarian News and Analysis. 2010. Archived from the original on May 13, 2013. Retrieved April 24, 2013.

Further reading