Aluminium sulfate
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| Names | |
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| IUPAC name
Aluminium sulfate
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| Other names
Cake alum
Filter alum Papermaker's alum Alunogenite Sulfuric acid, aluminum salt (3:2) | |
| Identifiers | |
3D model (JSmol)
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| ChemSpider | |
| ECHA InfoCard | 100.030.110 |
| EC Number |
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| E number | E520 (acidity regulators, ...) |
PubChem CID
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| UNII | |
CompTox Dashboard (EPA)
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| Properties | |
| Al2(SO4)3 | |
| Molar mass | 342.15 g/mol (anhydrous) 666.42 g/mol (octadecahydrate) |
| Appearance | white crystalline solid hygroscopic |
| Density | 2.672 g/cm3 (anhydrous) 1.62 g/cm3 (octadecahydrate) |
| Melting point | 770 °C (decomp, 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 |
| Acidity (pKa) | 3.3-3.6 |
Refractive index (nD)
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1.47 [1] |
| Structure | |
| monoclinic (hydrate) | |
| Related compounds | |
Other cations
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Gallium sulfate Magnesium sulfate |
| Supplementary data page | |
| Aluminium sulfate (data page) | |
Except where otherwise noted, data are given for materials in their standard state (at 25 °C [77 °F], 100 kPa).
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Aluminium sulfate, alternatively spelt aluminum sulfate, aluminium sulphate, or aluminum sulphate; is a chemical compound with the formula Al2(SO4)3. Aluminium sulfate 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 incorrectly referred to as alum but alums are closely related compounds typified by KAl(SO4)2.12H2O. 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.
Preparation
Aluminium sulfate may be made by adding aluminium hydroxide, Al(OH)3, in sulfuric acid, H2SO4:
- 2 Al(OH)3 + 3 H2SO4 → Al2(SO4)3·6H2O
Uses
Aluminium sulfate is used in water purification and as a mordant in dyeing and printing textiles. In water purification, it causes impurities to coagulate which are removed as the particulate settles to the bottom of the container or more easily filtered. This process is called coagulation or flocculation.
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.
Aluminium sulfate is the active ingredient of some antiperspirants; however, beginning in 2005 the US Food and Drug Administration no longer recognized it as a wetness reducer.
Aluminium 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.
In construction industry it is used as waterproofing agent and accelerator in concrete. Another use is a foaming agent in fire fighting foam.
It is also used in styptic pencils, and pain relief from stings and bites.
It can also be very effective as a molluscicide, killing spanish slugs.
Chemical reactions
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]
References
Footnotes
- ^ Pradyot Patnaik. Handbook of Inorganic Chemicals. McGraw-Hill, 2002, ISBN 0-07-049439-8
- ^ Global Health and Education Foundation (2007). "Conventional Coagulation-Flocculation-Sedimentation". Safe Drinking Water is Essential. National Academy of Sciences. Retrieved 2007-12-01.
- ^ 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.
Notations
- Pauling, Linus (1970). General Chemistry. W.H. Freeman: San Francisco. ISBN 0-486-65622-5.