Aluminium sulfate is a chemical compound with the formula Al2(SO4)3. It is soluble in water and is mainly used as a coagulating agent (promoting particle collision by neutralizing charge) in the purification of drinking water and waste water treatment plants, and also in paper manufacturing.
aluminum salt (3:2)
3D model (JSmol)
|E number||E520 (acidity regulators, ...)|
CompTox Dashboard (EPA)
|Molar mass||342.15 g/mol (anhydrous) |
666.44 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)
|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|
Refractive index (nD)
Std enthalpy of
|Safety data sheet||See: data page|
|US health exposure limits (NIOSH):|
IDLH (Immediate danger)
|Supplementary data page|
|Refractive index (n),|
Dielectric constant (εr), etc.
|UV, IR, NMR, MS|
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
|what is ?)(|
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 is sometimes called alum or papermaker's alum in certain industries. However, the name "alum" is more commonly and properly used for any double sulfate salt with the generic formula XAl(SO
2O, where X is a monovalent cation such as potassium or ammonium.
In the laboratoryEdit
- 2 Al(OH)3 + 3 H2SO4 → Al2(SO4)3 + 6H2O
or by heating aluminum metal in a sulfuric acid solution:
- 2 Al + 3 H2SO4 → Al2(SO4)3 + 3 H2↑
From alum schistsEdit
The alum schists employed in the manufacture of aluminium sulfate are mixtures of iron pyrite, aluminium silicate and various bituminous substances, and are found in upper Bavaria, Bohemia, Belgium, and Scotland. These are either roasted or exposed to the weathering action of the air. In the roasting process, sulfuric acid is formed and acts on the clay to form aluminium sulfate, a similar condition of affairs being produced during weathering. The mass is now systematically extracted with water, and a solution of aluminium sulfate of specific gravity 1.16 is prepared. This solution is allowed to stand for some time (in order that any calcium sulfate and basic ferric sulfate may separate), and is then evaporated until ferrous sulfate crystallizes on cooling; it is then drawn off and evaporated until it attains a specific gravity of 1.40. It is now allowed to stand for some time, and decanted from any sediment.
From clays or bauxiteEdit
In the preparation of aluminum sulfate from clays or from bauxite, the material is gently calcined, then mixed with sulfuric acid and heated gradually to boiling; it is allowed to stand for some time, the clear solution drawn off.
When cryolite is used as the ore, it is mixed with calcium carbonate and heated. By this means, sodium aluminate is formed; it is then extracted with water and precipitated either by sodium bicarbonate or by passing a current of carbon dioxide through the solution. The precipitate is then dissolved in sulfuric acid.
It is sometimes used in the human food industry as a firming agent for food starch, where it takes on E number E520, and in animal feed as a bactericide. Aluminum sulfate may be used as a deodorant, an astringent, or as a stiptic for superficial shaving wounds.
Aluminium sulfate is used in water purification and as a mordant in dyeing and printing textiles. In water purification, it causes suspended 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. An improper and excess application incident in 1988 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 Hydrangea macrophylla. The gardener can add aluminium sulfate to the soil to reduce the pH which in turn will result in the flowers of the Hydrangea turning a different color (blue). The aluminium is what makes the flowers blue; at a higher pH, the aluminium is not available to the plant.
In limnology (lake & pond science), Alum is often effectively employed to control levels of soluble ortho-phosphate (PO4) and to agglomerate suspended dirt particles for improved clarification. The product can be applied as a granular product or as "Liquid Alum" which is water + dissolved ALum (48.6% Alum dry basis). pH must be monitored in lakes with valuable fish when considering an Alum treatment. Alum is quite acidic and a lake with low pH (below 7) or a lake lacking a natural carbonate buffering capacity may find that the pH falls below 6 and to level when fish as harmed or killed. In such cases, a buffering agent - such as lime or Soda Ash or Sodium Bicarbonate - may be applied simultaneously with the Alum to prevent a precipitous drop in pH.
3 + 3 Pb(CH
2 → 2 Al(CH
3 + 3 PbSO
3 + 2 Pb(CH
2 → Al
4 + 2 PbSO
- 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.
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