Sodium sulfate

Sodium sulfate is an inorganic compound with the chemical formula Na2SO4, which is a white solid and soluble in water. Its decahydrate has an annual production of 6 million tons and is a major chemical product. It is mainly used as a filler in the manufacture of powdered household laundry detergents, in applications such as the sulfate pulping method in pulp manufacture, and in the manufacture of highly alkaline sulfides.

Form

Anhydrous sodium sulfate; the mineral form is anhydrous mannite, also known as metamictan in high purity, fine particles, and crystals composed of the rhombohedral crystal system. It is widely used in various industrial fields: in the paper industry as a cooking agent for sulfate pulp; in the glass industry as a substitute for soda ash; in the chemical industry for the manufacture of sodium sulfide, sodium silicate; in the dyestuffs industry as a dyeing agent, and so on.

Sodium sulfate heptahydrate: a very rare form of orthorhombic or rhombohedral crystal system composed of crystals that can be generated from anhydrous sodium sulfate by absorbing water.

Sodium sulfate decahydrate: often referred to as mannite, is a monoclinic crystal system consisting of crystals. It is widely used in the chemical industry, pharmaceuticals, tanning, textile printing and dyeing.

History

Sodium sulfate decahydrate is known as Glauber's salt, named after the Dutch/German chemist and pharmacist Johann Glauber. He discovered the salt in a spring in Austria in 1625. Naming it sal mirabilis (miracle salt) because of the medicinal value of its crystals as a general purpose laxative, it wasn't until the 20th century that other drugs with similar effects were discovered.

In the 18th century, mannite began to be used as a raw material for the industrial production of soda ash (sodium carbonate). As the demand for soda ash increased, the supply of sodium sulfate had to increase accordingly. Therefore, in the 19th century, the large-scale Leblanc alkali production method to produce synthetic sodium sulfate as a key intermediate substance became the main method of soda ash production.

Chemical Properties

Sodium sulfate is a typical electrostatically bound ionic sulfate. When barium or lead salts are added to form insoluble sulfates, the presence of free sulfate ions in solution can be demonstrated:

Na2SO4 + BaCl2 → 2 NaCl + BaSO4

Sodium sulfate is unreactive to most oxidizing or reducing agents. At high temperatures it can be converted to sodium sulfide by carbothermal reduction:

Na2SO4 + 2 C → Na2S + 2 CO2

This reaction is used in the Leblanc process, an obsolete industrial route for the production of sodium carbonate.

Sodium sulfate reacts with sulfuric acid to give the acid salt sodium bisulfate:

Na2SO4 + H2SO4 ⇌ 2 NaHSO4

Sodium sulfate shows slightly the property of forming complex salts. The only alums formed with common trivalent metals are NaAl(SO4)2 (unstable above 39 °C) and NaCr(SO4)2, while potassium sulfate and ammonium sulfate form many stable alums.

Physical Properties

Solubility curve of Na2SO4 in water

Sodium sulfate has unusual solubility characteristics in water.  Its solubility in water rises more than tenfold between 0°C and 32.384°C to a maximum value of 49.7 g/100 mL, followed by a gradual decrease in solubility at a temperature corresponding to the release of water of crystallization and the melting of hydrated salts, which serves as an accurate temperature reference for thermometer calibration.

Framework

The crystals of decahydrate consist of [Na(H2O)6]+ ions with an octahedral shaped molecular configuration. These octahedra share sides, so that 8 of the 10 water molecules are bound to sodium and the other 2 are interstitial, hydrogen bonded to sulfate. These cations are connected to the sulfate anion by hydrogen bonds.The Na-O distance is about 240 pm.Crystallized sodium sulfate decahydrate is also unusual among hydrated salts in that it has a measurable cosentropy (entropy at absolute zero) of 6.32 J/(K-mol). This is due to its ability to distribute water more rapidly than most hydrates.

Production

World production of sodium sulfate, almost exclusively in the form of decahydrate, is about 5.5 to 6 million tons per year (Mt/a).In 1985 production was 450 Mt/a, half from natural resources and half from chemical production.After the 21st century it was at a steady level until 2006, when natural production increased to 4 Mt/a and chemical production decreased to 1.5 to 2 Mt/a. Total production was 5.5 to 6 Mt/a.[17][18][19][20] For all applications, naturally and chemically produced sodium sulfate are practically interchangeable.

Natural resources

Two-thirds of the world's decahydrate production comes from the mineral mannite, such as is found on lake beds in southern Saskatchewan.In 1990, Mexico and Spain were the world's leading producers of natural sodium sulfate (about 500,000 tons each), and Russia, the United States, and Canada each had about 350,000 tons. Natural resources are estimated at over 1 billion tons.

Natural minerals used to prepare sodium sulfate are mainly manganese, anhydrous manganese, calcium manganese three kinds, of which manganese and anhydrous manganese for the sodium sulfate and its hydrate, calcium manganese for the sulfate minerals containing sodium calcium, chemical formula for Na2Ca (SO4) 2, the theoretical composition of 22.29% Na2O, 20.16% CaO, 57.55% SO3, monoclinic crystal system. All three ores are mainly used in the industrial preparation of anhydrous sodium sulfate and sodium sulfide, and their deposits are often associated with each other.

chemical production

About one third of the world's sodium sulfate is produced as a by-product of other processes in the chemical industry. Much of this production is inherent in the chemistry of the primary process and is of only minor economic interest. The production of sodium sulfate as a by-product is declining.

The most important chemical production of sodium sulfate is in the production of hydrochloric acid, from sodium chloride and sulfuric acid in the Mannheim process, or from sulfur dioxide in the Hargreaves process. The sodium sulfate produced in these processes is known as salt cake.

2 NaCl + H2SO4 → 2 HCl + Na2SO4

4 NaCl + 2 SO2 + O2 + 2 H2O → 4 HCl + 2 Na2SO4

The second major production process for sodium sulfate is the neutralization of the remaining sodium hydroxide with sulfuric acid. This method is also a frequently applied and convenient laboratory preparation.

2 NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2 H2O(l) ΔH = -112.5 kJ (highly exothermic)

In the laboratory, it can also be prepared from the reaction of sodium bicarbonate and magnesium sulfate by precipitating magnesium carbonate.

2 NaHCO3 + MgSO4 → Na2SO4 + MgCO3 + CO2 + H2O

It is not usually prepared in the laboratory because it is readily available from commercial sources.

In the past, sodium sulfate was also used as a by-product in the manufacture of sodium dichromate, where sulfuric acid was added to a solution of sodium chromate to form sodium dichromate, or subsequently chromic acid. In addition, sodium sulfate is or was formed during the production of lithium carbonate, chelating agents, resorcinol, ascorbic acid, silica pigments, nitric acid, and phenol.

Bulk sodium sulfate is usually purified through the decahydrate form because the anhydrous form tends to attract iron compounds and organic compounds. The anhydrous form is readily produced from the hydrate by gentle heating.

appliance

Commercial use

Sodium sulfate is a very cheap material. The largest use is as a filler in household laundry detergents, consuming about 50 per cent of world production. This use is declining as consumers gradually switch to compact or liquid detergents that do not contain sodium sulfate.

paper industry

Another major former use of sodium sulfate was in sulfate pulping for the manufacture of wood pulp. However, due to improvements in thermal efficiency in the kraft recycling process in the early 1960s, more efficient sulfur recovery was achieved and the need for sodium sulfate decreased dramatically. As a result, the use of sodium sulfate by the pulp industry in the United States and Canada declined from 1.4 million tons per year in 1970 to approximately 150,000 tons in 2006.

Glass Manufacturing

The glass industry is another important use of sodium sulfate, the second most important use in Europe. Sodium sulfate is used as a melting agent to help remove small air bubbles from molten glass. It makes the glass transparent and prevents the formation of slag in the molten glass during the refining process. Consumption by the glass industry in Europe remained stable at 110,000 tons per year from 1970 to 2006.

Textile Fabrics

Sodium sulfate is important in the production of textiles, especially in Japan. Sodium sulfate is added to increase the ionic strength of the solution, reducing the negative charge on the textile fibers and allowing the dye to penetrate evenly. Unlike the alternative sodium chloride, it does not corrode stainless steel containers used for dyeing.In 2006, about 100,000 tons were consumed in Japan and the U.S. for this application.

Food Industry

Sodium sulfate is used as a diluent for food coloring. Its E-number is E514.

thermal storage

The high heat storage capacity of sodium sulfate in the phase transition from solid to liquid, as well as the favorable phase transition temperature of 32 °C (90 °F), make this material particularly well suited for storing low-level solar heat for future release in space warming. In some applications, the material is incorporated into thermal tiles placed in attic spaces, while in others the salt is incorporated into batteries surrounded by solar-heated water. The phase change allows for a significant reduction in the mass of material required for effective heat storage (the heat of fusion of sodium sulfate decahydrate is 82 kJ/mol or 252 kJ/kg [25]), and the temperature remains consistent as long as there is enough material in the proper phase.

For cooling applications, mixtures with sodium chloride (NaCl) lower the melting point to 18 °C (64 °F).The heat of fusion for NaCl-Na2SO4-10H2O mixtures increases slightly to 286 kJ/kg.

small-scale application

Anhydrous sodium sulfate is widely used in the laboratory as an inert neutral desiccant for removing trace amounts of water from organic solutions. It is more effective than the similar magnesium sulfate, but slower acting. It is only effective below 30°C, but can be used for a wide variety of materials because it is chemically quite inert.

Sodium sulfate decahydrate is used as a laxative. It is effective in removing drugs such as paracetamol (acetaminophen) from the body, such as after an overdose.

In 1953, sodium sulfate was proposed to be used for thermal storage in passive solar heating systems. This took advantage of its unusual solubility properties and high heat of crystallization (78.2kJ/mol).

Other uses for sodium sulfate include defrosting windows, starch manufacturing, as an additive in carpet fresheners, and as an additive in cattle feed.

At least one company, Thermaltake, uses sodium sulfate decahydrate inside quilted plastic pads to make a laptop cooling pad (iXoft Notebook Cooler). The material slowly turns to liquid and is recirculated to equalize the temperature of the laptop and act as an insulator.

Security

Although sodium sulfate is generally considered non-toxic, it should be handled with care. The dust may cause temporary asthma or eye irritation, a risk that can be prevented by the use of eye protection and paper masks.