Sodium sulfate
From Wikipedia, the
free encyclopedia
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Sodium sulfate
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Thenardite (mineral)
Glauber's salt (decahydrate) Sal mirabilis (decahydrate) Mirabilite (decahydrate) |
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Identifiers
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WE1650000
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Jmol-3D images
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Properties
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Na2SO4
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142.04 g/mol
(anhydrous)
322.20 g/mol (decahydrate) |
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Appearance
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odorless
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2.664 g/cm3 (anhydrous)
1.464 g/cm3 (decahydrate) |
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884 °C (anhydrous)
32.38 °C (decahydrate) |
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1429 °C (anhydrous)
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anhydrous:
4.76 g/100 mL (0 °C) 42.7 g/100 mL (100 °C)
heptahydrate:
19.5 g/100 mL (0 °C) 44 g/100 mL (20 °C) |
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1.468 (anhydrous)
1.394 (decahydrate) |
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Structure
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Hazards
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EU Index
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Not listed
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Irritant
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Non-flammable
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Related compounds
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Related compounds
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Phase behaviour
Solid, liquid, gas |
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Except where noted otherwise, data are given for materials in
their standard state (at 25 °C (77 °F), 100 kPa)
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Sodium sulfate is the sodium salt of sulfuric acid. When anhydrous,
it is a white crystalline solid of formula Na2SO4 known as the mineral thenardite; the decahydrate Na2SO4·10H2O is found naturally as the mineral mirabilite, and in
processed form has been known as Glauber's salt or, historically, sal mirabilis since the 17th century. Another solid is the heptahydrate, which
transforms to mirabilite when cooled. With an annual production of 6 million tonnes,
it is a major commodity chemical product.
Sodium sulfate is mainly used for the manufacture of detergents and in the Kraft process of paper pulping.
About two-thirds of the world's production is from mirabilite, the
natural mineral form of the
decahydrate, and the remainder from by-products of chemical processes such as hydrochloric acid production.
Contents
The hydrate of sodium sulfate is known as Glauber's Salt after
the Dutch/German chemist and apothecary Johann Rudolf Glauber(1604–1670), who discovered it in 1625 in Austrian spring water.
He named it sal mirabilis (miraculous salt), because of its medicinal
properties: the crystals were used as a general purpose laxative, until more sophisticated alternatives came
about in the 1900s.[1][2]
In the 18th century, Glauber's salt began to be used as a raw
material for the industrial production of soda ash
(sodium carbonate), by reaction with potash (potassium carbonate). Demand for soda ash increased and the supply of sodium
sulfate had to increase in line. Therefore, in the nineteenth century, the
large scale Leblanc process, producing synthetic sodium sulfate as a key intermediate,
became the principal method of soda ash production.[3]
Sodium sulfate is chemically very stable, being unreactive
toward most oxidizing or reducing agents at normal temperatures. At high temperatures, it can be
converted to sodium sulfide by carbothermal reduction:[4]
Na2SO4 + 2 C → Na2S + 2 CO2
Sodium sulfate is a neutral salt, which forms aqueous solutions with pH of
7. The neutrality of such solutions reflects the fact that sulfate is derived,
formally, from the strong acid sulfuric acid. Furthermore, the Na+ ion, with only a
single positive charge, only weakly polarizes its water ligands provided there
are metal ions in solution. Sodium sulfate reacts with sulfuric acid to give
the acid saltsodium bisulfate:[5][6]
Na2SO4 + H2SO4 ⇌ 2 NaHSO4
Sodium sulfate has unusual solubility
characteristics in water.[7] Its solubility in water rises more than
tenfold between 0 °C to 32.384 °C, where it reaches a maximum of
497 g/L. At this point the solubility curve changes slope, and the
solubility becomes almost independent of temperature. This temperature at
32.384 °C, corresponding to the release of crystal water and melting of the
hydrated salt, serves as an accurate temperature reference for thermometer calibration.
Sodium sulfate is a typical ionic sulfate, containing Na+ ions and SO42− ions. The existence of
sulfate in solution is indicated by the easy formation of insoluble sulfates
when these solutions are treated with Ba2+ or Pb2+ salts:
Na2SO4 + BaCl2 → 2 NaCl + BaSO4
Graph showing
solubility of Na2SO4vs. temperature.
Sodium sulfate displays a moderate tendency to
form double salts.
The only alums formed with common
trivalent metals are NaAl(SO4)2 (unstable above 39 °C) and NaCr(SO4)2, in contrast to potassium sulfate and ammonium sulfate which form many stable alums.[8] Double salts with some
other alkali metal sulfates are known, including Na2SO4·3K2SO4 which occurs naturally as the mineral glaserite. Formation
of glaserite by reaction of sodium sulfate with potassium chloride has been used as the basis of a method for producing potassium sulfate, a fertiliser.[9] Other double salts
include 3Na2SO4·CaSO4, 3Na2SO4·MgSO4 (vanthoffite) and
NaF·Na2SO4.[10]
Crystals consist of [Na(OH2)6]+ ions with octahedral molecular geometry, as seen for many metallic sulfate salts. These cations are
linked to the sulfate anions via hydrogen bonds. The Na-O distances are 240 pm. Two molecules of water per formula unit are
not coordinated to Na+.[11] Crystalline sodium sulfate decahydrate is also
unusual among hydrated salts in having a measureable residual entropy (entropy at absolute zero) of 6.32 J·K−1·mol−1. This is ascribed to its ability to distribute water much more
rapidly compared to most hydrates.[12]
The world production of sodium sulfate, mostly
in the form of the decahydrate amounts to approximately 5.5 to
6 million tonnes annually (Mt/a). In 1985, production was
4.5 Mt/a, half from natural sources, and half from chemical production.
After 2000, at a stable level until 2006, natural production had increased to
4 Mt/a, and chemical production decreased to 1.5 to 2 Mt/a, with a
total of 5.5 to 6 Mt/a.[13][14][15][16] For all applications,
naturally produced and chemically produced sodium sulfate are practically
interchangeable.
Two thirds of the world's production of the
decahydrate (Glauber's salt) is from the natural mineral form mirabilite, for
example as found in lake beds in southern Saskatchewan.
In 1990, Mexico and Spain were the world's main producers of natural
sodium sulfate (each around 500,000 tonnes), with Russia, United States and Canada around
350,000 tonnes each.[14] Natural resources are
estimated as over 1 billion tonnes.[13][14]
Major producers of 200,000 to 1,500,000
tonnes/a in 2006 include Searles Valley Minerals (California, US), Airborne Industrial Minerals (Saskatchewan,
Canada), QuÃmica del Rey (Coahuila, Mexico), Minera de Santa Marta and
Criaderos Minerales Y Derivados, also known as Grupo Crimidesa (Burgos, Spain), Minera de Santa Marta
(Toledo, Spain), Sulquisa (Madrid, Spain), and in China Chengdu Sanlian
Tianquan Chemical (Sichuan), Hongze Yinzhu Chemical Group (Jiangsu), Nafine
Chemical Industry Group (Shanxi), and Sichuan Province Chuanmei Mirabilite
(Sichuan), and Kuchuksulphat JSC (Altai Krai, Siberia, Russia).[13][15] In Saskatchewan, one
of the major miners is Saskatchewan Minerals.
Anhydrous sodium sulfate occurs in arid
environments as the mineral thenardite. It
slowly turns to mirabilite in damp air. Sodium sulfate is also found as glauberite, a
calcium sodium sulfate mineral. Both minerals are less common than mirabilite.
About one third of the world's sodium sulfate
is produced as by-product of other processes in chemical industry. Most of this
production is chemically inherent to the primary process, and only marginally
economical. By effort of the industry, therefore, sodium sulfate production as
by-product is declining.
The most important chemical sodium sulfate
production is during hydrochloric acid production, either from sodium chloride (salt) and sulfuric acid, in the Mannheim process, or from sulfur dioxide in the Hargreaves process.[17][18] The resulting sodium
sulfate from these processes is known as salt cake.
Mannheim: 2 NaCl + H2SO4 → 2 HCl + Na2SO4
Hargreaves: 4 NaCl + 2 SO2 + O2 + 2 H2O → 4 HCl + 2 Na2SO4
The second major production of sodium sulfate
are the processes where surplus sulfuric acid is neutralised by sodium hydroxide, as applied on a large scale in the production ofrayon. This method is also a regularly applied and
convenient laboratory preparation.
In the laboratory it can also be synthesized
from the reaction between sodium bicarbonate and magnesium sulfate.
2NaHCO3 + MgSO4 → Na2SO4 + Mg(OH)2 + 2CO2
Formerly, sodium sulfate was also a by-product of the
manufacture of sodium dichromate, where sulfuric acid is added to sodium chromate solution
forming sodium dichromate, or subsequently chromic acid. Alternatively, sodium
sulfate is or was formed in the production of lithium carbonate, chelating agents, resorcinol, ascorbic acid, silica pigments, nitric acid,
and phenol.[13]
Bulk sodium sulfate is usually purified via the decahydrate
form, since the anhydrous form tends to attract iron compounds and organic compounds. The anhydrous form is easily produced from the hydrated form
by gentle warming.
Major sodium sulfate by-product producers of 50–80 Mt/a in 2006
include Elementis Chromium (chromium industry, Castle Hayne, NC, US), Lenzing
AG (200 Mt/a, rayon industry, Lenzing, Austria), Addiseo (formerly Rhodia,
methionine industry, Les Roches-Roussillon, France), Elementis (chromium
industry, Stockton-on-Tees, UK), Shikoku Chemicals (Tokushima, Japan) and
Visko-R (rayon industry, Russia).[13]
Sodium sulfate used to
dry an organic liquid. Here clumps form, indicating the presence of water in
the organic liquid.
By further application
of sodium sulfate the liquid may be brought to dryness, indicated here by the
absence of clumping.
With US pricing at $30 per tonne in 1970,6 up to $90 per tonne
for salt cake quality and $130 for better grades, sodium sulfate is a very
cheap material. The largest use is as filler in powdered home
laundry detergents, consuming approx. 50% of world production. This use is
waning as domestic consumers are increasingly switching to compact or liquid
detergents that do not include sodium sulfate.[13]
Another formerly major use for sodium sulfate, notably in the US
and Canada, is in the Kraft process for the manufacture of wood pulp. Organics present in the "black
liquor" from this process are burnt to produce heat, needed to drive the reduction of sodium sulfate tosodium sulfide. However, this process is being replaced by newer processes;
use of sodium sulfate in the US and Canadian pulp industry declined from 1.4
Mt/a in 1970 to only approx. 150,000 tonnes in 2006.[13]
The glass industry provides another
significant application for sodium sulfate, as second largest application in
Europe. Sodium sulfate is used as a fining agent,
to help remove small air bubbles from molten glass. It fluxes the glass, and
prevents scum formation of the glass melt during refining. The glass industry
in Europe has been consuming from 1970 to 2006 a stable 110,000 tonnes
annually.[13]
Sodium sulfate is important in the manufacture of textiles, particularly in Japan, where it is the
largest application. Sodium sulfate helps in "levelling", reducing
negative charges on fibres so that dyes can penetrate evenly. Unlike the
alternative sodium chloride, it does not corrode the stainless steel vessels used in dyeing. This application in Japan and US
consumed in 2006 approximately 100,000 tonnes.[13]
The high heat storage capacity in the phase change from solid to
liquid, and the advantageous phase change temperature of 32 °C (90 °F) makes
this material especially appropriate for storing low grade solar heat for later
release in space heating applications. In some applications the material is
incorporated into thermal tiles that are placed in an attic space while in
other applications the salt is incorporated into cells surrounded by
solar–heated water. The phase change allows a substantial reduction in the mass
of the material required for effective heat storage (the heat of fusion of
sodium sulfate decahydrate is 82 kJ/mol or 252 kJ/kg[19]), with the further advantage of a consistency
of temperature as long as sufficient material in the appropriate phase is
available.
For cooling applications, a mixture with common sodium chloride salt (NaCl) lowers the melting point to 18 °C (64 °F). The heat
of fusion of NaCl·Na2SO4·10H2O, is actually increased slightly to 286 kJ/kg.[20]
In the laboratory, anhydrous sodium sulfate is widely used as an
inert drying agent, for removing traces of water from organic
solutions.[21] It is more efficient,
but slower-acting, than the similar agent magnesium sulfate. It is only effective below about 30 °C, but it can be
used with a variety of materials since it is chemically fairly inert. Sodium
sulfate is added to the solution until the crystals no longer clump together;
the two video clips (see above) demonstrate how the crystals clump when still
wet, but some crystals flow freely once a sample is dry.
Glauber's salt, the decahydrate, was historically used as a laxative. It is effective for the removal of certain
drugs such as paracetamol(acetaminophen)
from the body, for example, after an overdose.[22][23]
In 1953, sodium sulfate was proposed for heat storage in passive solar heating systems. This takes advantage of its unusual solubility
properties, and the high heat ofcrystallisation (78.2 kJ/mol).[24]
Other uses for sodium sulfate include de-frosting windows, in
carpet fresheners, starch manufacture, and as an
additive to cattle feed.
Lately, sodium sulfate has been found effective in dissolving
very finely electroplated micrometre gold that is found in gold electroplated
hardware on electronic products such as pins, and other connectors and switches.
It is safer and cheaper than other reagents used for gold recovery, with little
concern for adverse reactions or health effects.[citation needed]
At least one company, ThermalTake, makes a laptop computer chill
mat (iXoft Notebook Cooler) using sodium sulfate decahydrate inside a quilted
plastic pad. The material slowly turns to liquid and recirculates, equalizing
laptop temperature and acting as an insulation.
Although sodium sulfate is generally regarded as non-toxic,[25] it should be handled with care. The dust can
cause temporary asthma or eye irritation; this risk can be prevented by using
eye protection and a paper mask. Transport is not limited, and no Risk Phrase or Safety Phrase apply.[26]
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