SODIUM SULFATE (NA2SO4): STRUCTURE, PROPERTIES, USES, OBTAINING - CHEMISTRY - 2025

The sodium sulfate is an inorganic salt having the chemical formula Na ₂ SO ₄. It consists of a white solid, present in three forms: anhydrous, the heptahydrate (scarcely available) and the decahydrate (which is known as Glaubert's salt); the latter is the most abundant form of sodium sulfate.

Sodium sulfate decahydrate, Na ₂ SO ₄ · 10H ₂ O, was discovered in 1625 by Glaubert in spring water, who named it mirabilis salt (miraculous salt) due to its medicinal properties.

Watch glass with a sample of sodium sulfate. Source: Walkerma via Wikipedia.

Sodium sulfate has numerous applications in the textile and paper industry, as well as in glass manufacturing. Its uses are extended to thermal applications, which include the supply of ambient heat and cooling of laptops.

Sodium sulfate is a compound of low toxicity and its harmful actions are mainly mechanical and non-chemical. For crystallographic reasons, this salt, like its potassium counterpart, K ₂ SO ₄, has lattice and polymorphic structures.

Structure

Anhydrous salt

Anhydrous sodium sulfate ions. Source: Claudio Pistilli

The formula Na ₂ SO ₄ indicates at once that the Na ⁺ and SO ₄ ^2- ions are in a 1: 2 ratio in the salt crystals; that is, for every two Na ⁺ cations there is an anion SO ₄ ^2- interacting with them through electrostatic attraction (upper image).

Of course, this applies to anhydrous Na ₂ SO ₄, with no water molecules coordinated with the sodium within the crystals.

Sodium sulfate

Despite being an apparently simple salt, its description is structurally complex. Na ₂ SO ₄ presents polymorphism, having up to five crystalline phases: I, II, III, IV and V, whose transition temperatures are 180, 200, 228, 235 and 883 ºC, respectively.

Although there are no references to certify it, Na ₂ SO ₄ I must be the one with a hexagonal crystalline structure, denser compared to the orthorhombic Na ₂ SO ₄ III, in whose crystals Na ⁺ forms tetrahedra (NaO ₄) and coordination octahedra (NaO ₆); that is, it can be surrounded by four or six SO ₄ ^2- anions.

Dehydrated salt

Meanwhile, the monoclinic crystalline structure of its most important hydrate, Na ₂ SO ₄ · 10H ₂ O, is simpler. In it, it is practically the water molecules that interact or coordinate with the Na ⁺ in the Na (H ₂ O) ₆ ⁺ octahedra, with the SO ₄ ^2- just providing enough stability to the crystal so that it exists in the solid phase.

However, its melting point (32.38ºC) much lower than that of the anhydrous salt (884ºC) shows how water molecules and their hydrogen bonds weaken the stronger ionic interactions in Na ₂ SO ₄.

Properties

Names

-Sodium sulfate (IUPAC)

-Glauber's salt (decahydrate)

-Miraculous salt (decahydrate)

-Disodium sulfate.

Molar mass

142.04 g / mol (anhydrous)

322.20 g / mol (decahydrate)

Physical appearance

White hygroscopic crystalline solid

Odor

Toilet

Taste

Bitter and salty

Density

2,664 g / cm ³ (anhydrous)

1,464 g / cm ³ (decahydrate)

Note how the water molecules inside the crystals cause them to expand and, therefore, decrease their density.

Melting point

884 ºC (anhydrous)

32.38 ºC (decahydrate)

Boiling point

1,429 ºC (anhydrous)

Water solubility

4.76 g / 100 ml (0 ºC)

13.9 g / 100 ml (20 ° C)

42.7 g / 100 ml (100 ° C)

All solubility values ​​correspond to the anhydrous salt, which is quite soluble in water at all temperatures.

The solubility increases abruptly between 0ºC and 38.34ºC, observing that in this temperature range the solubility increases more than 10 times. However, from 32.38ºC the solubility is independent of temperature.

It happens that at a temperature of 32.8 ºC, the sodium sulfate decahydrate dissolves in its own crystalline water. An equilibrium is thus reached between the decahydrate salt, the anhydrous salt and the saturated solution of sodium sulfate.

As long as the three-phase condition is maintained, the temperature will remain constant, which allows the temperature of the thermometers to be calibrated.

On the other hand, the solubilities for the heptahydrated salt are:

19.5 g / 100 ml (0 ºC)

44.0 g / 100 ml (20 ° C)

Note that at 20 ° C the heptahydrate salt is three times more soluble than the anhydrous one.

Refractive index

1,468 (anhydrous)

1.394 (decahydrate)

Stability

Stable under recommended storage conditions. Incompatible with strong acids and bases, aluminum and magnesium.

Decomposition

When heated to decomposition it emits toxic smoke of sulfur oxide and sodium oxide.

pH

A 5% aqueous solution has a pH of 7.

Reactivity

Sodium sulfate dissociates in aqueous solution into 2 Na ⁺ and SO ₄ ^2-, which allows the sulfate ion to combine with Ba ²⁺ to precipitate barium sulfate. It practically helps displace barium ions from water samples.

Sodium sulfate is converted to sodium sulfide by reacting at elevated temperatures with coal:

Na ₂ SO ₄ + 2 C => Na ₂ S + 2 CO ₂

Glaubert's salt, NaSO ₄.10H ₂ O reacts with potassium carbonate to produce sodium carbonate.

Applications

Paper industry

Sodium sulfate is used in the manufacture of paper pulp. It is used in the production of Kraft paper, which does not contain lignin and is not subjected to the bleaching process, which gives it great resistance. In addition, it is used in the manufacture of cardboard.

Detergents

It is used as a filler material for synthetic household detergents, being added to detergent to reduce surface tension.

Glasses

It is used in glass making to reduce or eliminate the presence of small air bubbles in molten glass. Additionally, it eliminates the formation of slag during the refining process of the molten glass.

Textile industry

Sodium sulfate is used as a mordant, since it facilitates the interaction of dyes with the fibers of fabrics. Sodium sulfate decahydrate is used in the dye test.

In addition, sodium sulfate is used as a dye diluent and dye printing auxiliary agent; such as direct dyes, sulfur dyes, and other agents that promote cotton staining. It is also used as a retarding agent for direct silk dyes.

Medicine

Sodium sulfate decahydrate is used as a laxative, since it is poorly absorbed in the intestine, and therefore remains in the lumen of the intestine causing an increase in volume. This stimulates the increase in peristaltic contractions which induce the expulsion of intestinal contents.

Sodium sulfate is an antidote to control barium and lead salt poisoning. Glaubert's salt is effective in eliminating certain excessively ingested medications; for example, paracetamol (acetoaminophen).

In addition, it is used to supply deficient electrolytes present in isoosmotic solutions.

Drying agent

Sodium sulfate, being an inert reagent, is used to eliminate water from solutions of organic compounds.

Raw material

Sodium sulfate is used as a raw material for the production of numerous substances, including: sodium sulfide, sodium carbonate, and ammonium sulfate.

Obtaining

Sodium sulfate is obtained by mining extraction and by chemical reactions.

Mining extraction

There are three ores or minerals that are exploited with commercial yield: thenardite (Na ₂ SO ₄), mirabilite (Na ₂ SO ₄ · 10H ₂ O) and glaubarite (Na ₂ SO ₄ · CaSO ₄).

In Spain, thenardite and mirabilite deposits are exploited by underground mining of galleries and pillars. Meanwhile, the glauberite is obtained in the open, by means of large rafts that are placed on the mineral deposit.

The land is prepared with low intensity blasting to produce a porosity that allows the leaching of sodium sulfate. The production phase occurs with freshwater sprinkler irrigation of glauberite, the leaching of which spreads downward.

The sodium sulfate brine is collected, leaving the calcium sulfate residue as filling.

Chemical production

Sodium sulfate is obtained during the production of hydrochloric acid by two processes: the Mannheim process and the Hardgreaves process.

Mannheim Process

It is carried out in large steel furnaces and with a 6 m steel reaction platform. The reaction occurs between sodium chloride and sulfuric acid:

2 NaCl + H ₂ SO ₄ => 2 HCl + Na ₂ SO ₄

Hardgreaves Process

It involves the reaction of sodium chloride, sulfur oxide, oxygen, and water:

4 NaCl + 2 SO ₂ + O ₂ + 2 H ₂ O => 4 HCl + Na ₂ SO ₄

Others

Sodium sulfate is produced in the neutralization reaction between sodium hydroxide and sulfuric acid:

2 NaOH + H ₂ SO ₄ => Na ₂ SO ₄ + H ₂ O

Sodium sulfate is a by-product of the production of numerous compounds. It is extracted from the liquid waste discharged during the production of viscose and cellophane. Also in the production of sodium dichromate, phenols, boric acid and lithium carbamate.

Risks

Sodium sulfate is considered a low toxicity compound. However, it may cause some harm to the person who uses it improperly.

For example, contact can cause eye irritation, redness and pain. On the skin it can cause irritation and allergy in some people. Its ingestion can cause irritation of the digestive tract with nausea, vomiting and diarrhea. And finally, its inhalation produces irritation in the respiratory tract.

References

1. Shiver & Atkins. (2008). Inorganic chemistry. (Fourth edition). Mc Graw Hill.
2. Wikipedia. (2019). Sodium sulfate. Recovered from: en.wikipedia.org
3. National Center for Biotechnology Information. (2019). Sodium sulfate. PubChem Database. CID = 24436. Recovered from: pubchem.ncbi.nlm.nih.gov
4. BN Mehrotra. (1978). The crystal structure of Na ₂ SO ₄ III. Recovered from: rruff-2.geo.arizona.edu
5. Glauberite-Thenardite (sodium sulfate).. Recovered from: igme.es