Aluminium sulfate

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Aluminium sulfate
Aluminium sulfate hexadecahydrate
CAS number 10043-01-3 YesY
7784-31-8 (octadecahydrate)
PubChem 24850
ChemSpider 23233 YesY
UNII I7T908772F YesY
EC number 233-135-0
RTECS number BD1700000
Jmol-3D images Image 1
Molecular formula Al2(SO4)3
Molar mass 342.15 g/mol (anhydrous)
666.42 g/mol (octadecahydrate)
Appearance white crystalline solid
Density 2.672 g/cm3 (anhydrous)
1.62 g/cm3 (octadecahydrate)
Melting point 770 °C (1,420 °F; 1,040 K) (decomposes, anhydrous)
86.5 °C (octadecahydrate)
Solubility in water 31.2 g/100 mL (0 °C)
36.4 g/100 mL (20 °C)
89.0 g/100 mL (100 °C)
Solubility slightly soluble in alcohol, dilute mineral acids
Acidity (pKa) 3.3-3.6
Refractive index (nD) 1.47 [1]
Crystal structure monoclinic (hydrate)
Std enthalpy of
-3440 kJ/mol
MSDS External MSDS
NFPA 704
Related compounds
Other cations Gallium sulfate
Magnesium sulfate
Related compounds See Alum
Supplementary data page
Structure and
n, εr, etc.
Phase behaviour
Solid, liquid, gas
Spectral data UV, IR, NMR, MS
Infobox references

Aluminium sulfate, alternatively spelled either aluminum or sulphate, is a chemical compound with the formula Al2(SO4)3. It is soluble in water and is mainly used as a flocculating agent in the purification of drinking water[2][3] and waste water treatment plants, and also in paper manufacturing.

Aluminium sulfate is sometimes referred to as a type of alum. Alums are double sulfate salts, with the formula AM(SO
, where A is a monovalent cation such as potassium or ammonium and M is a trivalent metal ion such as aluminium.[4] The anhydrous form occurs naturally as a rare mineral millosevichite, found e.g. in volcanic environments and on burning coal-mining waste dumps. Aluminium sulfate is rarely, if ever, encountered as the anhydrous salt. It forms a number of different hydrates, of which the hexadecahydrate Al2(SO4)3•16H2O and octadecahydrate Al2(SO4)3•18H2O are the most common. The heptadecahydrate, whose formula can be written as [Al(H2O)6]2(SO4)3•5H2O, occurs naturally as the mineral alunogen.


Aluminium sulfate may be made by adding aluminium hydroxide, Al(OH)3, to sulfuric acid, H2SO4:

2 Al(OH)3 + 3 H2SO4 → Al2(SO4)3·6H2O

or by heating aluminum metal in a sulfuric acid solution:

2 Al(s) + 3 H2SO4 → Al2(SO4)3 + 3 H2 (g)


Aluminium sulfate is used in water purification and as a mordant in dyeing and printing textiles. In water purification, it causes impurities to coagulate into larger particles and then settle to the bottom of the container (or be filtered out) more easily. This process is called coagulation or flocculation. Research suggests that in Australia, aluminium sulfate used this way in drinking water treatment is the primary source of hydrogen sulfide gas in sanitary sewer systems.[5] Improper and excess application polluted the water supply of Camelford in Cornwall.

When dissolved in a large amount of neutral or slightly alkaline water, aluminium sulfate produces a gelatinous precipitate of aluminium hydroxide, Al(OH)3. In dyeing and printing cloth, the gelatinous precipitate helps the dye adhere to the clothing fibers by rendering the pigment insoluble.

Aluminium sulfate is sometimes used to reduce the pH of garden soil, as it hydrolyzes to form the aluminium hydroxide precipitate and a dilute sulfuric acid solution. An example of what changing the pH level of soil can do to plants is visible when looking at the Hydrangea macrophylla. The gardener can add aluminium sulfate to the soil to reduce the pH level which in turn will result in the flowers of the Hydrangea turning a different color (blue).

Aluminium potassium sulfate and another form of alum, aluminium ammonium sulfate are the active ingredients in some antiperspirants; however, beginning in 2005 the US Food and Drug Administration no longer recognized it as a wetness reducer.

Aluminium potassium sulfate is usually found in baking powder, where there is controversy over its use due to concern regarding the safety of adding aluminium to the diet.


It is also used in styptic pencils, and pain relief from stings and bites.


Chemical reactions[edit]

The compound decomposes to γ−alumina and sulfur trioxide when heated between 580 and 900 °C. It combines with water forming hydrated salts of various compositions.

Aluminium sulfate reacts with sodium bicarbonate to which foam stabilizer has been added, producing carbon dioxide for fire-extinguishing foams:

Al2(SO4)3 + 6 NaHCO3 → 3 Na2SO4 + 2 Al(OH)3 + 6 CO2

The carbon dioxide is trapped by the foam stabilizer and creates a thick foam which will float on top of hydrocarbon fuels and seal off access to atmospheric oxygen, smothering the fire. Chemical foam was unsuitable for use on polar solvents such as alcohol, as the fuel would mix with and break down the foam blanket. The carbon dioxide generated also served to propel the foam out of the container, be it a portable fire extinguisher or fixed installation using hoselines. Chemical foam is considered obsolete in the United States and has been replaced by synthetic mechanical foams, such as AFFF which have a longer shelf life, are more effective, and more versatile, although some countries such as Japan and India continue to use it.[citation needed]



  2. ^ Global Health and Education Foundation (2007). "Conventional Coagulation-Flocculation-Sedimentation". Safe Drinking Water is Essential. National Academy of Sciences. Retrieved 2007-12-01. 
  3. ^ Kvech S, Edwards M (2002). "Solubility controls on aluminum in drinking water at relatively low and high pH". WATER RESEARCH 36 (17): 4356–4368. doi:10.1016/S0043-1354(02)00137-9. PMID 12420940. 
  4. ^ Austin, George T. (1984). Shreve's Chemical process industries. (5th ed. ed.). New York: McGraw-Hill. p. 357. ISBN 9780070571471. 
  5. ^ Ilje Pikaar, Keshab R. Sharma, Shihu Hu, Wolfgang Gernjak, Jürg Keller, Zhiguo Yuan. "Reducing sewer corrosion through integrated urban water management". Retrieved 2014-08-25. 



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